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Chemistry STEMscopes Georgia Student Notebook

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Chemistry Notebook

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GEORGIA Student Notebook – Chemistry ISBN: 978-1-946725-81-3

Published by Accelerate Learning Inc., 5177 Richmond Ave, Suite 800, Houston, TX 77056. Copyright © 2017, by Accelerate Learning Inc. All rights reserved. No part of this publication may be reproduced or distributed in any form or by any means, or stored in a database or retrieval system, without prior written consent of Accelerate Learning Inc., including, but not limited to, in any network or other electronic storage or transmission, or broadcast for distance learning. To learn more, visit us at www.acceleratelearning.com

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GEORGIA

This Student Notebook is designed to be used as a companion piece to our online curriculum. The pages of this book are organized and follow the 5E model.

Student Handout

ENGAGE

A short activity to grab students’ interest

Student Journal

EXPLORE

A hands-on activity in which students get experience with the concept being taught

STEMscopedia

EXPLAIN

A reference material that includes parent connections, technology, and science news

Reading Science A reading passage about the concept that includes comprehension questions

ELABORATE Math Connection

A set of grade-level appropriate math problems that address the concept

Writing Science

EVALUATE

A writing prompt based on the concept studied in which the students can demonstrate what they have learned

Only student pages are included in this book and directions on how to use these pages are found in our online curriculum. Use the URL address and password provided to you by your district to access our full curriculum.

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Chemistry Student Notebook Table of Contents

Lesson

Page

C1ABC

C1D

C1EF

C1G

C2AB

117

C2C

139

C2D

159

C2E

187

C3AB

221

Models and Properties of Atoms

Nuclear Fusion

Electron Movement and Configuration

Electron Patterns and Properties of Elements

Inter- and Intra- Molecular Forces

Molecular Structure in Designed Materials

Bonding Models

Naming Compounds

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Chemistry Student Notebook Table of Contents

Lesson

Page

C3C

253

C3DE

285

C4ABC

315

C4D

341

C5AB

363

C5C

393

C6ABE

425

C6CD

451

C6FGH

475

The Mole and Avogadro’s Number

Stoichiometry

Reaction Rates

Equilibrium

Thermal Energy in Chemical and Physical Processes Gas Laws

Solutes and the Process of Dissolving

Concentrations

Properties and Models of Acids and Bases

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Chemistry Student Notebook Table of Contents

Lesson

Page Appendix

Glossary

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High School Chemistry

C1ABC

Models and Properties of Atoms

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C1ABC Models and Properties of Atoms

Student Handout Name:

Date:

Bag #1 Observations

Guess

Actual Object

Guess

Actual Object

Guess

Actual Object

Guess

Actual Object

Guess

Actual Object

Bag #2 Observations

Bag #3 Observations

Bag #4 Observations

Bag #5 Observations

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C1ABC Models and Properties of Atoms

Student Handout Questions: 1.

How did the bag hinder you from being able to fully describe the object?

How does this relate to early scientists?

What type of experiments/tests (without opening the bags) could you do with the bags to further figure out properties of the objects?

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C1ABC Models and Properties of Atoms

Student Journal Name:

Date:

Group:

Part I: Your Research and Presentation In this section, record your research. Name of scientist and model

Draw a picture of your atomic model. Briefly describe the experiment that your scientist used. Describe the protons for your model (charge, location). Describe the electrons for your model (charge, location). Describe the neutrons for your model (charge, location). What was discovered by this scientist?

What limitations did this model have?

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C1ABC Models and Properties of Atoms

Student Journal Part II: Student Debate When presentations are ready, participate in your group’s presentation. Watch and listen to the other presentations. Fill in the data tables below as the information is revealed. Scientists and Models Dalton & Model

Thomson & Model

Rutherford & Model

Bohr & Model

Model Name

Picture

Protons

Neutrons

Electrons

Merits

Limitations

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C1ABC Models and Properties of Atoms

Student Journal Reflections and Conclusions 1.

Predict the chronological order of the models.

Which model has the best representation of the proton?

Which model has the best representation of the electron?

Which model has the best representation of the neutron?

On another sheet of paper, use all of the following terms to develop a graphic organizer based on what you have learned in this activity.

Terms: John Dalton, J.J. Thomson, Ernest Rutherford, Niels Bohr, atom, postulates, Bohr model, nuclear model, electron, proton, neutron, atomic theory, nucleus, energy level, atomic mass

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C1ABC Models and Properties of Atoms

Student Journal Name:

Date:

Group:

Part I: Build an Atom PhET Simulation Using “Build an Atom,” play with the parts of atoms to discover the following. 1.

What parts go in the center of the atom? What is the center called?

Fill in the table to identify four examples. What is in your nucleus?

Draw your nucleus.

What element is it?

1 2 3 4 3.

What did you find determines the element? Test your idea by identifying the element for the three cases. Write down the information you use to determine the element.

Example 1

The atom has... # of protons: 6 # of netrons: 6 # of electrons: 6 # of protons: 7 # of netrons: 7 # of electrons: 7 # of protons: 8 # of netrons: 8 # of electrons: 8

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What element is it?

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C1ABC Models and Properties of Atoms

Student Journal Part I: Build an Atom PhET Simulation, continued 4.

Now make the number of protons and electrons unequal. Play until you discover some good rules about the charge of your atom or ion. What is a rule for making: •

a neutral atom that has no charge?

•

a positive ion that has a positive charge?

•

a negative ion that has a negative charge?

Write about how you used the tools in the simulation to help you decide whether the atom had a positive, negative, or zero charge. (A box on the lower right of the screen states the current charge.)

Use the table below to identify three examples of atoms and ions (one neutral with zero extra charges, one with a positive charge, and one with a negative charge) that show your rules for charge work and include a drawing of your atom. All your examples should also have a stable nucleus. What is in your atom? 1

Draw your atom ion.

What is your charge?

Is it a neutral atom, positive atom, or negative ion?

# of protons: # of netrons: # of electrons: # of protons: # of netrons: # of electrons: # of protons: # of netrons: # of electrons:

Play until you discover some good rules about the mass of your atom or ion. What is a rule for determining the mass?

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C1ABC Models and Properties of Atoms

Student Journal 8.

Using all of your rules, figure out what changes for each of these modifications to an atom or ion. Use this table and make predictions, and then test your ideas with the simulation. If you have new ideas, then rewrite your rules. Make the change

What also changes? Element name, charge, mass?

Add a proton Remove a neutron Remove an electron Add an electron 9.

Design a positive ion with a charge of +2. Include a drawing. •

What element is your ion? ____________________

•

What mass is your ion? _____________________

•

Is the nucleus of your ion stable or unstable? _______________

10. Design a neutral, stable atom with a mass of eight. Include a drawing. •

What element is your atom?_______________

•

What is the charge of your atom?____________

11. What does the tool called Symbol tell you about what parts are in an atom or ion?

12. What rules can you use to tell how many protons, neutrons, and electrons make up an atom or ion?

Reflection. Make sure you know working definitions for the following terms: nucleus, proton, neutron, electron, atom, ion, charge, neutral, atomic mass, and element. Games. Play the games to check your understanding of these concepts.

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C1ABC Models and Properties of Atoms

Student Journal Part II: Isotopes Observe the data chart and answer the related questions.

The relative abundance of the two isotopes of Bromine, Br-79 and Br-81, is shown above. The atomic mass of Bromine on the periodic table is 79.98. Explain how scientists could have determined this to be the correct mass.

Natural uranium is 98% U-238 and 2% U-235. Uranium used in a nuclear bomb is 99% U-235 and 1% U-238. How would the average atomic mass of samples from these uranium combinations compare?

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STEMscopedia: MODELS AND PROPERTIES OF

ATOMS C1ABC

Reflect Imagine a piece of aluminum foil about the size of a sheet of paper. If you cut this piece of foil in half, have you changed the identity of the matter making up the foil? What happens if you cut the foil in half again? And if you cut this sample in half, what do you have? You know that cutting aluminum foil in half does not change its chemical identity. You still have aluminum foil every time you cut a piece in half. But suppose you could keep cutting the foil in half indefinitely. Would you ever reach a point where you could not cut any further without destroying the aluminum’s identity? The Greeks hypothesized that matter is not infinitely divisible. Can matter be divided into infinitely smaller portions? Long before the microscope was invented, the ancient Greeks debated this question. In 435 BCE, the Greek philosopher Leucippus reasoned that matter must be composed of finite particles. If something is finite, it has a limit; in other words, it cannot be divided into smaller and smaller pieces forever. His student, Democritus, developed this idea. Democritus hypothesized Democritus used the idea of tiny, indivisible particles called atoms to that all matter and explain the properties and behavior of matter. (In Greek, atom literally its behavior could be means “uncuttable.”) Democritus hypothesized that these particles explained by the nature of are always moving, and therefore their arrangement in space must be tiny indivisible particles that constantly changing. According to Democritus, the movements and comprise it. changing arrangements of atoms determine all of the phenomena that we observe in the natural world. Neither of these men had any experimental evidence supporting their ideas. They based their ideas on philosophical thought and reasoning. John Dalton presented the first experimental evidence to support the concept of atoms. Many centuries passed before anyone was able to find experimental evidence for the existence of atoms. John Dalton, an English schoolteacher, carried out experiments that provided indirect evidence of atoms. He showed that oxygen combined with another gas, nitric oxide, in specific volume and weight ratios. In other words, no matter how much oxygen and nitric oxide he combined, the two always reacted with one another in a constant weight ratio. Dalton concluded that this could occur only if the gases were made up of atoms reacting with one another in whole number combinations. He proposed his atomic theory in 1803 as a series of statements called postulates: • All matter is composed of indivisible particles that cannot be created or destroyed. Like Leucippus and Democritus, Dalton hypothesized that the smallest unit of matter is an indivisible particle. He reasoned that these particles—which he also called atoms—are present in a fixed number in the universe. Atoms cannot be created or destroyed.

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•

Atoms of the same element are identical, but atoms of different elements have different properties. Scientists were well aware of different elements—including iron, mercury, gold, and silver—at the time Dalton developed his atomic theory. Dalton used his concept of atoms to explain the differences between elements. He proposed that each element is composed of the same type of atoms, but that the atoms in one element differ from the atoms in another element. A chief difference is atomic weight. Different types of atoms have different weights. • Atoms group together in whole number ratios to form compounds. Dalton’s experiments confirmed what other scientists had discovered about chemical compounds—they are made up of two or more different elements. He also confirmed that in any compound, the same weight ratio of the elements is always present. For example, carbon monoxide always has one gram of carbon for every 1.333 g of oxygen. However, carbon dioxide always has one gram of carbon for every 2.666 g of oxygen. Dalton used this information as evidence for his atomic theory. According to Dalton, the fixed weight ratios of elements in a compound result from the Dalton’s theory explains fixed ratios of atoms of those elements. the observation that carbon and oxygen can Chemical reactions are the result of the rearrangement of atoms. combine in two ways. By the time Dalton began his work, several scientists had already published results of their experiments involving chemical reactions. These results indicated that two substances combined in constant proportions. Dalton used his concept of atoms to explain that the constant proportions indicate a specific whole number of atoms of one reactant rearranging with another whole number of atoms of a second reactant. Although Dalton’s first postulate overlapped with the ideas put forth by the Greeks 2,000 years earlier, his ideas carried more weight because they were backed by experimental evidence. Dalton was cautious in that regard, relying on investigational data rather than philosophical thinking as he developed his atomic theory. J. J. Thomson discovered the first subatomic particle: the electron. Almost 100 years after Dalton proposed his atomic theory, English scientist J. J. Thomson made another discovery about the nature of matter. By this time, scientists had accepted Dalton’s postulates about the atom as the fundamental particle. However, nothing was known about the makeup of the atom itself. In the late 1800s, J. J. Thomson was using a device called a cathode ray tube to explore some mysterious rays that caused a fluorescent glow. A cathode ray tube is a glass tube with two metal wires at each end. When Thomson used a power supply to apply a high voltage across the two wires, he could observe a greenish glowing light inside the tube. 14

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A cathode ray tube uses high voltage to produce a light beam inside a glass tube.

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STEMscopedia: MODELS AND PROPERTIES OF

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Scientists at that time were not sure about the nature of this light. They didn’t know if it was some kind of light wave or a stream of particles. A number of scientists ran experiments to try to uncover the nature of the mysterious rays. J. J. Thomson succeeded in devising a series of experiments that solved the mystery. Thomson built several new cathode ray tubes, one of which is shown in the diagram to the right. He designed this tube to test whether the cathode rays could be bent by an external electric field. He added two plates to the interior of the tube: one with a positive charge and the other with a negative charge. These plates were positioned above and below the path of the cathode rays. When he tested his newly built tube, Thomson found that the cathode rays bent away from the negatively charged plate. Thomson concluded the rays were composed of particles that carried a negative charge. More experiments with other tubes enabled Thomson to estimate the size of these particles. His work established that they were much smaller than atoms. Thompson did not use the term electron to describe the negatively charged particles he discovered. However, Thompson did hypothesize that these particles represented parts of an atom. Thompson proposed that the structure of an atom could be likened to an English dessert known as plum pudding. He envisioned a model in which the negatively charged particles are embedded in a sphere of positive charge much like raisins are embedded in a plum pudding. A hunch led Ernest Rutherford to discover that atoms contain a nucleus. Thomson’s plum pudding model of the atom was later discarded in favor of a model in which electrons occupy the space around a positively charged nucleus. One of Thompson’s former students, Ernest Rutherford, came up with the nuclear model. Rutherford had worked in Thompson’s lab in England in the 1890s before becoming a professor at another English university.

J. J. Thomson’s cathode ray tube experiments demonstrated that cathode rays are composed of negatively charged particles much smaller than atoms. Later scientists would call these negative particles “electrons.”

J. J. Thomson developed the “plum pudding model,” in which negatively charged particles are embedded in a positively charged sphere.

There, he worked with his own research students to study the nature of a type of radioactivity known as alpha radiation. Alpha radiation results from the emission of large particles, called alpha particles, from unstable atoms. In 1908, Rutherford directed his students to conduct some experiments involving the emission of alpha radiation at squares of thin, gold foil. He wondered what kind of backscatter his students might find when the alpha radiation hit the gold atoms. His students carried out the experiments and reported that most of the alpha particles passed through the foil while a small proportion was deflected backward.

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STEMscopedia: MODELS AND PROPERTIES OF

ATOMS

From these results, Rutherford concluded that atoms contain extremely small, dense, and positively charged nuclei. If an alpha particle collided with an atom’s nucleus, the particle bounced off the gold foil. Because most particles passed through the foil, however, Rutherford concluded that the area around the nucleus is mostly empty space with a few negative electrons. He published these conclusions in 1911, refuting the plum pudding model. Ernest Rutherford’s gold foil experiment demonstrated that the atom could not be described by the plum pudding model.

What Do You Think? What results would Rutherford and his students have observed if the plum pudding model had been accurate? Note that in this model, the matter making up the atom was thought to be spread more evenly within the atom and not concentrated into a small space at the center. Niels Bohr refined the atomic model to explain how electrons were positioned around the nucleus. Rutherford’s work shed light on the presence of an atomic nucleus, but it didn’t reveal any new information about the positions of electrons in an atom. Niels Bohr, a student in Rutherford’s lab, worked on extending Rutherford’s model to explain how electrons were arranged in the atom.

Bohr’s model of the hydrogen atom placed electrons in specific orbits around the nucleus.

At that time, scientists knew that different elements emitted different patterns of light when they were heated to high temperatures. They also knew that the pattern of light emitted by an element could be described by a mathematical relationship. Bohr took this idea one step further and hypothesized that the mathematical relationship indicated the presence of different energy levels for electrons within the atom. From this, Bohr developed a model of an atom with a central nucleus surrounded by electrons moving in orbits. Moving from the nucleus outward, each orbit represented higher energy than the previous one. In Bohr’s model, electrons could move from one orbit to another. If an electron moved from a lower energy level to a higher one, it had to absorb energy to make the move. If an electron moved from a higher energy level to a lower one, it lost energy in the form of light. This explained the light emission spectrum of a hydrogen atom.

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However, scientists gradually realized that hydrogen is the only element whose light emission spectrum fits Bohr’s mathematical calculations. Despite this serious flaw, the Bohr model of the atom continued to be important to scientists. The idea of specific energy levels for electrons became an important one that endured even after the model itself was found to be inadequate. All of these scientists were important in the development of the atomic theory, the design of our current atomic model, along with the role of the different subatomic particles discovered through their many experiments. Atomic Structure All atoms have the same general arrangement. An atom contains a nucleus, which is in the center of the atom. An atom also has an area of space surrounding the nucleus called an electron cloud. Subatomic particles, the small particles within the atom, are known as protons, neutrons, and electrons. They differ in mass, charge, and location in the atom. • •

•

Neutron (no): Does not have an electrical charge. They are neutral. Neutrons are found in the nucleus of the atom and are represented by the green spheres in the atomic diagram. Proton (p+): Protons are positively charged particles (+1). They are found in the nucleus of an atom and are represented by the green spheres in the atomic diagram to the right. Since protons are the only charged particle in the nucleus, an atom’s nucleus is always positively charged. The number of protons in the nucleus determines the identity of the atom. This means that each element has a unique number of protons (atomic number Z) in the nucleus of each atom. For example, all carbon atoms contain six protons, all gold atoms contain 79 protons, and all lead atoms contain 82 protons. Electron (e-): Electrons are negatively charged particles (-1). They travel around the nucleus in orbitals, within the electron cloud surrounding the nucleus. Electrons are constantly moving. The blue spheres in the atomic diagram represent electrons.

Classic model of a carbon atom with atomic number of 6 (six protons). In the nucleus are six protons and six neutrons.

Different atoms have different combinations of subatomic particles. However, there are some general rules regarding the electrical charge of an atom. When the number of protons equals the number of electrons, the atom is neutral because the positive nucleus balances with the negatively charged electrons. If there are more protons than electrons, the atom becomes a positively charged ion. If there are more electrons than protons, the atom becomes a negatively charged ion.

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STEMscopedia: MODELS AND PROPERTIES OF

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What Do You Think? Why do you think the number of neutrons in an atom does not affect the overall charge of the atom? Hint: Think about the characteristics of protons and electrons. Atomic structure affects the properties of elements. Chemical reactions involve either the transfer or the sharing of electrons between atoms. Therefore, the chemical reactivity properties of an element primarily depend on the number of electrons in an atom of that element. Protons are significant because the tendency for an atom to either lose, gain, or share electrons depends on the total charge of the protons in the nucleus. The chemical reactivity of an atom depends on the number of electrons and protons, and is independent of the number of neutrons. The mass and radioactive properties of an atom depend on the number of protons and neutrons in the nucleus. The Masses of Subatomic Particles Protons and neutrons have similar masses. Each proton and neutron has a mass of approximately 1.67 x 10–27 kilograms. The mass of one electron is even smaller: 9.11 x 10–31 kilograms. Protons and neutrons located in the atom’s nucleus represent where most of the atom’s mass is located. Atomic mass is the sum of the mass of all of the protons and neutrons in the atom. Electrons are so small that they not do change the mass of the atom significantly. They do, however, take up a large amount of the volume of the atom. The following formula can be used to ﬁnd the atomic mass where P = number of protons and N = number of neutrons: P + N = atomic mass. Atoms of the same element can have different numbers of neutrons. Elements in the periodic table are ordered by their atomic number: the number of protons in the atom’s nucleus. The number of protons defines the type of element, yet atoms of the same element may have different numbers of electrons and neutrons. Because both atoms in the diagrams below have seven protons, they are atoms of the same element: nitrogen (N). They are also isotopes. What are isotopes? Isotopes are atoms of the same element that have different numbers of neutrons. Let’s consider another example. There are two natural isotopes of helium (He). One isotope has one neutron, and the other isotope has two neutrons. Both isotopes have two protons. Therefore, the mass number—the total number of an atom’s protons and neutrons—of the first helium isotope is three (two protons plus one neutron). The mass number of the second helium isotope is four (two protons plus two neutrons). Helium has only two isotopes, but most elements have more. Iron (Fe) has more than 25 isotopes, but many of these isotopes are not very stable. The most abundant iron isotope has 26 protons and 30 neutrons. Therefore, this isotope has a mass number of 56.

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Look Out There are several ways to write an isotope of a particular element. One way is to write the mass number as a superscript before the elemental symbol. Following this method, the two isotopes of helium are written 3He and 4He. Another way is to use a hyphen to separate the name of the element and its mass number: for example, helium-3 and helium-4. The average atomic mass of an element depends on its isotopes. On the periodic table, each element is listed with its atomic number above the chemical symbol. Below the atomic symbol is a decimal number that is the element’s average atomic mass. An element’s average atomic mass is the weighted average of the mass numbers of all the naturally occurring atoms and isotopes of that element. Most elements have more than one naturally occurring isotope. Do not confuse an isotope’s mass number with the average atomic mass of the element. For example, the element boron (B) has two naturally occurring isotopes: boron-10 and boron-11. To calculate a weighted average, you need to know the abundance of each isotope: How frequently does it occur in nature? The abundance of boron-10 is 19.9 percent, and the abundance of boron-11 is 80.1percent. (In other words, approximately four out of every five boron isotopes in nature have five protons and six neutrons. The remaining boron isotopes in nature have only five neutrons.) Average atomic mass is calculated by multiplying the mass number of each isotope by the abundance of that isotope, and then taking the sum. For example, here is how to calculate the average atomic mass of boron: Average atomic massB = (10)(19.9/100) + (11)(80.1/100) = 1.99 + 8.811 = 10.801 We just calculated boron’s average atomic mass to be 10.801. However, the periodic table lists boron’s average atomic mass as 10.811. What happened? In fact, an isotope’s mass number does not exactly equal the sum of its protons and neutrons. The mass number of boron-10 is actually closer to 10.013 atomic mass unit, and the mass number of boron-11 is actually closer to 11.009 atomic mass unit. (Atomic mass is measured in atomic mass units: amu.) If you use these numbers in your calculation, you will get an average atomic mass closer to 10.811 atomic mass units.

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The atomic number of boron (B) is 5. The average atomic mass of boron is 10.811.

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STEMscopedia: MODELS AND PROPERTIES OF

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What Do You Know? Atoms are made of protons, neutrons, and electrons. Read the characteristics of these subatomic particles in the box below. Decide whether each characteristic describes protons, neutrons, or electrons. Then write each characteristic in the correct section of the Venn diagram. Characteristics of Subatomic Particles • Negatively charged • Positively charged • Have no charge • Smallest mass • In the nucleus • In orbitals around the nucleus • Subatomic particles • Discovered by J. J. Thomson • Make up the mass of the atom • When equal, atom is neutral

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Connecting With Your Child Determining Averages at Home To help your child practice the calculations necessary to determine average atomic mass, calculate the average masses of everyday objects such as eggs in an egg carton. You may use a small kitchen scale to make the measurements; all eggs should be roughly the same size (small, medium, large, or extra large). Use two cartons, each with different-sized eggs (small, medium, large, or extra large). One carton should contain a dozen eggs, all the same size; the other carton should contain a half dozen eggs, all the same size. Have your child create a two-column table with these headings: “Carton 1” and “Carton 2.” Include 12 rows, one for each egg in carton 1; you will use only the first six rows for carton 2. Students should record their measurements in their tables as they measure each egg. After your child measures all the eggs in the first carton, add their masses and divide by 12 to determine the average mass of these eggs. Have your child repeat the experiment for the second carton, dividing by six rather than 12 to determine the average mass of those eggs. Finally, determine the weighted average mass of all the eggs, using what you know of the abundance of each size of egg. For example, if the average mass of the large eggs in carton 1 is 55 grams, then 67 percent of your eggs (12 of 18) have an average mass of 55 grams. If the average mass of the small eggs in carton 2 is 45 grams, then 33 percent of your eggs (6 of 18) have an average mass of 45 grams. You can calculate weighted average mass as follows: Average masseggs = (55 g)(67/100) + (45 g)(33/100) = 36.85 g + 14.85 g = 51.7 g The weighted average mass of the eggs is 51.7 grams. Here are some questions to discuss with your child: 1. Why is average atomic mass an important value for each isotope? What can average atomic mass tell you about an element? 2. How does the average egg mass relate to the average atomic mass? How does the individual mass of each egg relate to an atom’s mass number? 3. How do the different cartons relate to isotopes?

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C1ABC Models and Properties of Atoms

Reading Science Name:

Date:

Group:

Protons and Electrons 1.

Our current model of how the atom is made involves protons in a nucleus and orbiting electrons. This understanding was developed over many years and was built on the work of many scientists. A major step was made in 1911 by Ernest Rutherford. Building on the work of Eugen Goldstein, Wilhelm Wien, and J. J. Thomson, Rutherford discovered the nucleus. Rutherford is often called the father of nuclear science.

Rutherford studied Thomson’s model of the atom and wanted to prove that it was right. Thomson’s model said that electrons were mixed in the sphere of the atom, like raisins in plum pudding. While testing this theory, Rutherford did the “gold foil” experiment.

In the gold foil experiment, Rutherford discovered that the mass of the atom is not placed evenly throughout the atomic sphere, as Thomson thought. In this experiment, Rutherford shot radiation particles at a thin sheet of gold. He expected them to go straight through. Instead, a few particles bounced off the foil instead of passing through it. He figured out that the center of an atom is made of a dense, positively charged middle surrounded by negatively charged electrons.

Rutherford’s nuclear model of the atom triggered other advances, including the naming of the proton and the realization of its importance. The word proton comes from the Greek word meaning “first.” Goldstein and Wien had discovered the proton years before. However, no one realized what it was until Rutherford’s work on the nucleus. Antonius van den Broek soon proposed that protons determined an element’s identity. He then suggested that elements be placed into the periodic table based on the number of protons the element has.

Physicist Niels Bohr improved Rutherford’s nuclear model of the atom. He recognized that electrons move around the nucleus in different energy levels. Part of his theory says that only valence electrons occupy the outermost shell of the electron cloud of an atom. Scientists had recognized that the chemical properties of elements are related to their electron configurations, which affects an element’s reactivity. Bohr and G. N. Lewis found that valence electrons determine the way the atom will react with other atoms. In other words, valence electrons give an atom its chemical properties.

As with most scientific theory, many scientists have contributed through the years. They have come up with theories, performed experiments, and shared findings to help us understand how atoms are put together.

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C1ABC Models and Properties of Atoms

Reading Science 1.

Which of the following lists correctly places the scientists in order from earliest to latest based on their contribution to our understanding of atomic structure described in this passage? A.

Lewis, Wien, Goldstein

Rutherford, Goldstein, Bohr

Bohr, Rutherford, Thomson

Thomson, van den Broek, Lewis

Thomson’s model suggested that electrons were mixed into the sphere of an atom much like raisins in _____________. A.

muffins

plum pudding

a box

toast

What did Rutherford’s experiment prove about the atomic nucleus? A.

It is like a raisin muffin, with a positively charged core.

It is densely clustered in the middle of the atom.

It was unknown by the means that Rutherford had access to at the time of the experiment.

It is made of electrons and neutrons.

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C1ABC Models and Properties of Atoms

Reading Science

Based on the descriptions in the passage, which scientist proposed a model of the atom that matches the diagram? A.

Ernest Rutherford

Niels Bohr

J. J. Thomson

Antonius van den Broek

What is the main point of the passage? A.

Scientific theories often come about as a result of many people.

Atoms have only been identified in the past 100 years.

Gold foil is the only type of material that can be used to find the mass of atoms.

Scientists do not use others’ work in forming their theories.

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C1ABC Models and Properties of Atoms

Math Connections Name:

Date:

Group:

Atomic Structure 1.

Using your periodic table, complete the chart below. The first row for nitrogen is completed for you as an example. Element

Atomic Number

Protons (+1)

Neutrons (0)

Electrons (-1)

Mass Number

Nitrogen

Boron Fluorine Potassium Carbon-12

Silicon-31 22 10 Ne 140 58 Ce 127 -1 53 I 24 1+ 11 Na 84 2+ 40 Zr

Write an equation that provides a general way to determine the net charge of an atom.

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C1ABC Models and Properties of Atoms

Math Connections Use the following element for questions 3–7: 3.

The number of electrons in X is _____.

The number of neutrons in X is _____.

The number of protons in X is ______.

The atomic mass of X is ______.

The above element is _______.

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46 -2 20 X

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C1ABC Models and Properties of Atoms

Writing Science Name:

Date:

Group:

LOOK

THINK atoms are made up of three subatomic particles that give them certain properties. Protons and neutrons are found within the nucleus, while electrons orbit around the nucleus. Each element on the periodic table contains different numbers of all three of these subatomic particles, with no two elements containing the exact same numbers. Two students are discussing the atom and are trying to decide which subatomic particle is most important when describing the type of element an atom is. One student argues that the number of electrons is what determines the atom, while the other believes it is the number of protons. WRITE which particle is most important to the identity of an atom, and why?

Be sure to • clearly state your central idea; • organize your thoughts; • develop your essay in detail; • choose your words carefully; and • use correct spelling, capitalization, punctuation, and grammar.

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C1ABC Models and Properties of Atoms

Writing Science

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High School Chemistry

C1D

Nuclear Fusion

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C1D Nuclear Fusion

Student Handout Name:

Date:

Symphony of Science Directions: Answer the following questions. 1.

How did atoms that come from a star become part of us?

The song says “stand in the middle.” Where is the middle?

What is the song trying to tell us when it says “everything both ways”?

What does it mean to cook light elements into heavy elements?

What does the song mean when it says “enriched guts”?

How does this song relate to the elements that are within you? Write a claim and evidence from the song to prove your answer. Claim: Evidence:

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STEMscopedia: Nuclear Fusion C1D

Reflect If elements in our universe were forged in stars, does that mean we are made of stardust? If the simplest elements are hydrogen and helium, how did heavier elements form that make up the matter on Earth and the rest of the universe? Where did everything in the universe come from? Finding the answers to these questions requires an understanding of the events that began our universe. Element formation begins with nucleosynthesis. Elements formed during the Big Bang and throughout the life cycle of stars. Nearly all observable matter in the universe is either hydrogen or helium, formed in the first minutes after the Big Bang. Elements other than these remnants of the Big Bang form within stars.

nucleosynthesis: creation of elements with heavier nuclei from hydrogen and helium inside stars

Nuclear fusion within stars produces all atomic nuclei lighter than and including iron, and the process releases the energy seen as starlight. Heavier elements are produced when certain massive stars achieve a supernova stage and explode. The formation of elements is revealed in the history of the universe through spectral analysis of stars in various stages of their life.

Big Bang Nucleosynthesis The Big Bang led to the initial production of hydrogen, helium, lithium, and beryllium. Big Bang nucleosynthesis refers to the transformation of some of the hydrogen, which formed first, into deuterium (hydrogen atoms containing one proton and one neutron), helium, lithium, and beryllium. Big Bang nucleosynthesis occurred within the first three minutes of the beginning of the universe and is responsible for much of the abundance in the universe of 2H (D, deuterium), 3He (helium-3), 4 He (helium-4), and some of the 7Li (Lithium) and 9Be (Beryllium) that exist today.

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STEMscopedia: Nuclear Fusion Stellar fusion and alpha decays continue to produce some 4He, and high-energy impacts and certain types of radioactive decay continue to produce trace amounts of 1H. Scientists believe the nuclei of these elements, along with some 7Li and 7Be, formed between 100 and 300 seconds after the Big Bang when the primordial quarkgluon plasma froze out to form protons and neutrons.

quark: elementary particle that forms composite particles called hadrons (protons and neutrons), the components of atomic nuclei

For the first 380,000 years after the Big Bang, it was so hot that electrons were unable to orbit around the hydrogen, helium, and lithium nuclei. During that period, the universe was a hot, opaque, expanding plasma. Plasma is a gas with electrons not orbiting atomic nuclei.

gluon: messenger particle of the strong nuclear force, which binds the subatomic particles known as quarks

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STEMscopedia: Nuclear Fusion About 380,000 years after the Big Bang, matter had cooled enough that electrons could combine with nuclei to form neutral atoms. The absorption of free electrons caused the universe to become transparent. The light released at this time is detectable today in the form of radiation from the cosmic microwave background. The cosmic microwave background radiation is a remnant of the Big Bang. This NASA image shows minute temperature variations (shades of blue and purple) linked to slight density variations in the early universe, which led to clusters of galaxies as well as vast empty regions.

Look Out The Big Bang did not occur at the center of the universe. In fact, there really is no center of the universe. The Big Bang did not occur in one place and spread out through space; it created space and time. Space and time did not exist before the Big Bang. In other words, the universe is not moving into pre-existing space. Rather, as it expands, it creates more space. There is a geometric center of the universe, but there is nothing special there other than empty expanding space. Galaxies are bound together in groups called clusters of galaxies. Those clusters group together in superclusters of galaxies. As the universe expands, the space between superclusters of galaxies expands, moving the superclusters apart. Consider a breaddough-with-raisins analogy. The bread (space) expands, which causes the raisins (superclusters of galaxies) to move farther apart.

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STEMscopedia: Nuclear Fusion Stellar Nucleosynthesis 400 million years after the Big Bang, clumps of gas collapsed enough to form the very first stars and galaxies. Within the cores of stars, high temperatures and pressure caused the fusion of nuclei to form new elements in a process called stellar nucleosynthesis.

Hydrogen Burning: The formation of elements during the life cycle of a main sequence star, such as the Sun, begins with a process called “hydrogen burning,” which is a stellar process that results in the nuclear fusion of four protons to form a nucleus of helium-4. This should not be confused with the chemical combustion of hydrogen in an oxidizing atmosphere. The predominant process for hydrogen fusion in the Sun is called the Proton-Proton Chain Reaction. Proton-Proton Chain Reaction: In the cores of lower mass main sequence stars, such as the Sun, tremendous heat and pressure in the core causes a process called the proton-proton chain reaction. This creates a helium-4 nucleus through a sequence of chain reactions that begin with the fusion of two hydrogen protons to form a nucleus of deuterium. The proton-proton chain reaction requires temperatures of 10 million degrees Kelvin, where zero degrees Kelvin is the coldest it can be anywhere in the universe. The freezing point of water is 273 K, and 300 K is roughly room temperature. The Sun’s core, where fusion is occurring, is the innermost 25% of the Sun’s radius.

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STEMscopedia: Nuclear Fusion Once a star has converted nearly all the hydrogen in its core into helium, fusion stops there. The core begins to slowly contract, heating up the hydrogen just outside of the core. Hydrogen burning occurs in a shell surrounding the inert helium core. This fusion causes the outer layers of the star to expand. It becomes a red giant. The core temperature increases until it exceeds one hundred million degrees K. At that point helium burning in the core begins with a runaway thermal process called the helium flash, with hydrogen burning continuing in a thin shell surrounding the now active helium core. Helium Burning: Hydrogen burning continues in the star’s shell, producing helium that falls into the core. Helium fusion produces the new elements beryllium, carbon, and oxygen: He2+4He2

He2+8Be4

He2+12C6

Be4 C6

Eventually the star runs out of helium, and He-burning stops. For stars of less than 2–4 solar masses, the star contracts but cannot restart burning. Instead it contracts, cools, and becomes a white dwarf star with a density of one million g/cm3.

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STEMscopedia: Nuclear Fusion Element Formation Up to Iron: Stars of greater than about 4 solar masses go into further successive burning stages, producing oxygen, neon, magnesium, silicon, and iron in their cores and in shells around their cores, and become red supergiant stars: carbon burning (8×108 K) produces O, 20Ne10, 24Mg12 neon burning (1.5×109 K) produces O and Mg oxygen burning (2×109 K) produces Mg to 32S silicon burning (3×109 K) produces elements up to 56Fe26

Supernovae Nucleosynthesis Iron Core Collapses: In massive stars, a complex series of nuclear reactions leads to the production of iron in the core, and the star has wrung all the energy it could out of nuclear fusion. Fusion reactions that form elements heavier than iron actually consume energy rather than produce it. The star no longer has any way to support its own mass, and the iron core collapses. In just a matter of seconds the core shrinks from roughly 5,000 miles across to just a dozen, and the temperature spikes 100 billion degrees or more.

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STEMscopedia: Nuclear Fusion What Do You Think? Astronomers have imaged the Crab Nebula in our Milky Way galaxy. You can see the gas and dust remnants of a dying star with a pulsar at its center. A pulsar is the tiny, condensed remnant of a once massive star. At its core is the rapidly spinning remains of the original star. What processes do you think led up to the star’s demise? Rebound Releases Energy and New Elements The outer layers of the star initially begin to collapse along with the core, but rebound with the enormous release of energy and are thrown violently outward. Supernovae release an almost unimaginable amount of energy. For a period of days to weeks, a supernova may outshine an entire galaxy. Likewise, all the naturally occurring elements and a rich array of subatomic particles are produced in these explosions. On average, a supernova explosion occurs about once every hundred years in a typical galaxy. Supernova explosions result when the cores of massive stars have exhausted their fuel supplies and burned everything into iron and nickel. The star then collapses and rebounds, slamming into the infalling outer layers, causing them to reverse direction and speed outward as supernovae. Elements heavier than iron are forged in these supernova events. Nuclei with mass heavier than nickel (gold, silver, lead, uranium, etc.) form in the explosions. These elements form over seconds, compared to the lighter ones that took billions of years to form and are much rarer.

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STEMscopedia: Nuclear Fusion Heavy Elements Become Part of New Formations: These conditions exist for only a short time, but long enough to form the highest mass nuclei that are blown off into space. This material will later make its way into other nebulas to become incorporated into other stars (where the same cycle of events will be repeated) and into planets, like the Earth. Each cycle uses up more of the H and He from the early universe and creates greater amounts of the higher mass elements.

What Do You Think? 1. Match these phrases associated with the formation of elements with their description: Big Bang nucleosynthesis

A. Nuclei with mass heavier than nickel (gold, silver, lead, uranium, etc.) form in this process.

Stars with the mass of the Sun

Stars with 1.4 times the mass of the Sun

B. The proton-proton cycle is the dominant process in these stars for hydrogen burning. C. Initial large-scale production of the lightest elements.

Supernovae nucleosynthesis

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STEMscopedia: Nuclear Fusion 2. Describe the evidence scientists have of the Big Bang background radiation.

3. Why does solar fusion in our Sun not form the heaviest elements?

Connecting With Your Child Internet, Library, Educational TV Support To help your child learn more about the formation of elements that occurred during the Big Bang, during proton-proton reaction, during stellar formation, and during supernovae explosions, give your child access to the internet, local libraries, and educational television. If you do not have adequate access to the internet at home, consider visiting local libraries or asking friends or family to use their computer access. Using a Google search will provide a multitude of internet sites that offer current information and interactive website that will engage your child in meaningful thinking skills. Your local library also has a variety of books in chemistry, physics, and astronomy that could serve as great resources for delving into deeper explanation on the formation of elements. After she or he has spent time in research, ask your child what was most interesting. Ask for explanations of difficult concepts. Encourage further research on remaining questions on the formation of elements. Local Astronomy Club Many cities have local astronomy clubs where adults offer viewing through their telescopes at a dark sky location open to the public. Most club members are knowledgeable about the Big Bang and can offer spectacular images of the Crab Nebulas and other nebulas that are the remnants of dying stars that seeded the universe with new elements.

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C1D Nuclear Fusion

Math Connections Name:

Date:

Group:

Charge #

Mass #

Particle

Penetration Level (low, middle, high)

Stopped by

4 2

0 −1

0 0

Use the following information to answer Questions 4–18. Equations can be written to represent the reactants and products in a nuclear reaction during alpha or beta emissions. A balanced nuclear equation shows an equal sum of the atomic masses and the atomic numbers on both sides of the equation. Balance the following reactions: 4.

+−10e

32 15

N+−10e

14 7

5. 6.

238 92

141 56

Th+

234 90

+−10e

Ir+42He

181 77

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C1D Nuclear Fusion

Math Connections 9.

14 6

10.

234 90

11.

234 911

12.

3 1

13.

9 4

14.

15 6

15.

27 13

+147N

+−10e

+23492U

+32He

Be+11H

He+

4 2

C+1cn

Al+21H

He+

4 2

+01n

16.

Cu+11H

17.

63 29

18.

233 92

63 30

K+42He

42 19

Zn+

Th+

229 90

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C1D Nuclear Fusion

Writing Science Name:

Date:

Group:

LOOK

THINK Our Sun and all stars are the birthplace of elements. The more massive a star is, the greater the amount and type of elements it produces. Our Sun will produce elements ranging from helium to carbon during its lifespan, but a massive star can produce elements all the way up to iron on the periodic table. Once a massive star goes supernova, all other elements can be created from its blast. This means that stars are responsible for all elements beyond helium that are found in the universe, including the elements that make up you. WRITE What is the element-creating process that takes place in stars and during supernova? Be sure to • clearly state your central idea; • organize your thoughts; • develop your essay in detail; • choose your words carefully; and • use correct spelling, capitalization, punctuation, and grammar.

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C1D Nuclear Fusion

Writing Science

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High School Chemistry

C1EF

Electron Movement and Configuration

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C1EF Electron Movement and Configuration

Student Handout Name:

Date:

Discovering Spectra Emission spectra are like fingerprints. Every element has a unique spectrum. When only atoms from one element are present in the gas, an emission spectrum characteristic of that element is produced. When atoms from two or more elements are present, their emission spectra add together and produce a combined spectrum. Use the handheld visual spectroscope to examine the emission spectra of at least two different elements. Sketch each spectrum below. Be sure to label the wavelengths. Element _______________________________________________________________________ Emission Spectrum

Element _______________________________________________________________________ Emission Spectrum

Now sketch the spectrum that would result from mixing the two elements. Elements ______________________________________________________________________ Emission Spectrum

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C1EF Electron Movement and Configuration

Student Journal Name:

Date:

Group:

Background: Energy Levels 1.

Similar to the partially completed diagram in your Student Guide, draw a diagram illustrating the six lowest energy levels of the hydrogen atom. Make sure the distances between the levels are properly scaled to represent the energy of the level. Label your diagram with the proper “n” number and energy (eV). Draw arrows to represent the transition of the electron to a lower energy level. Draw the transition of electrons from higher energy levels to n=1, then to n=2, and, lastly, to n=3 using a different colored pencil for each set. Use additional paper, if needed.

Use simple math to calculate the amount of energy emitted for each transition you illustrated for hydrogen in question 1. Analyze your calculations to identify the ones that produce visible wavelengths and complete the table below.

From Energy Level (n)

Dropped to Energy Level (n)

Photon Emission (eV)

Color Category

1.889

Red

In comparing the emission spectra of hydrogen and carbon, what can you infer about their transitions between energy levels?

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C1EF Electron Movement and Configuration

Student Journal A Renaissance Painting or Modern Forgery? Analysis In the space below, draw the emission spectra for each sample and any elements that show key characteristics providing evidence for your conclusion regarding a pigment match. Blue Pigment Sample

Mark your conclusion:

☐

Smalt

☐

Azurite

Red Pigment Sample

Mark your conclusion:

☐

Vermilion

☐

Cadmium red

Green Pigment Sample

Mark your conclusion:

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☐ Chrysocolla ☐ Viridian

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C1EF Electron Movement and Configuration

Student Journal Reflections and Conclusions 1. Explain how you analyzed the evidence to infer when the painting was produced.

2. What did each of the emission spectra you observed have in common? How did they differ?

3. How were you able to identify different elements by the emission lines? In your explanation, include how this relates to energy levels of an element.

4. Using all of the following terms, develop a graphic organizer based on what you have learned in this Explore. Use another sheet of paper, if needed. Terms: element, emission spectrum, energy state, spectroscope, photon, emission, atom

Critical Thinking and You 1. In this activity, you used spectroscopy and emission spectra to analyze pigments. What are some other applications of spectroscopy you can think of? What kinds of careers rely on spectroscopy using those applications? Use additional paper, as needed.

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C1EF Electron Movement and Configuration

Student Journal Name:

Date:

Group:

Part I: Principal Energy Levels and Subshells You will use the periodic table image on this page throughout this Explore activity, so make sure that you are very neat with your work as you go. Start by following the instructions outlined below. Use the provided Student Reference Sheet as a guide. • • • •

Color the entire s orbital block yellow. These are all of the group 1 (also known as group 1A) and 2 (2A) elements. The element helium is considered an s-orbital element, not a p-orbital element. Color the entire p orbital block green. These are all of the group 13-18 (3A-8A) elements. Color the entire d orbital block red. These are all of the group 3-12 (group 1B-8B, or transition elements). Color the entire f orbital block blue. These are the lanthanide and actinide series elements. Make sure to note where these groups begin on your periodic table.

n=1

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C1EF Electron Movement and Configuration

Student Journal Part I: Principal Energy Levels and Sublevels, continued This may all seem very confusing to you at this point, but do not worry. Follow the step-by-step instructions found on this page in exactly the order that they are given. By the time you are finished, you will be able to write and understand any electron configuration. Go back to the colored periodic table you have just finished. On the left hand side of the table, write in the principal energy levels at the beginning of each period. Label each as n = 1, n = 2, etc. Remember that the letter n represents the principal energy level. The first one has been completed for you. Do not label the lanthanide or actinide series section. Next, “block out” each electron orbital section. For example, in the 2s section, place a 2s between lithium and beryllium, and then draw arrows to each side to show where the 2s orbital block is. Complete the periodic table in this way. Now use your completed periodic table to fill in the first two columns of the table below. After the next few activities, you will add the number of orbitals within each subshell, the number of electrons that each subshell can hold, and the maximum number of electrons that each principal energy level can hold.

Principal energy level, quantum number (n)

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Number and type of electron subshell (s, p, d, f)

Number of orbitals within each subshell

Total number of electrons each subshell can hold

Maximum number of electrons each principal energy level can hold

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C1EF Electron Movement and Configuration

Student Journal Part I: Principal Energy Levels and Sublevels, continued Use the information from the periodic table created by the class to help you answer the following questions. 1.

What atomic subshells can be found in the representative elements? These are also known as the group A elements or main group elements.

What atomic subshell is found in the transition elements?

What atomic subshell is found in the lanthanide and actinide series elements?

The element nitrogen is found in the second period of the periodic table. In what principal energy level (n) is nitrogen (N)? What about potassium (K)?

Describe the pattern that you see on the class periodic table. How are the subshells arranged, or ordered?

There was a break in the pattern starting with principal energy level 4 and the element scandium (Sc). What was this break? Why do you think this break may have occurred? Hint: It has to do with electron energies.

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C1EF Electron Movement and Configuration

Student Journal Part II: Orbital Filling Diagrams 1.

Complete the orbital filling diagram below.

ELEMENT

2p x

2p y

2p z

3p x

3p y

3p z

Hydrogen (H) Helium (He) Lithium (Li) Beryllium (Be) Boron (B) Carbon (C) Nitrogen (N) Oxygen (O) Fluorine (F) Neon (Ne) Sodium (Na) Magnesium (Mg) Aluminum (Al) Silicon (Si) Phosphorus (P) Sulfur (S) Chlorine (Cl) Argon (Ar)

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C1EF Electron Movement and Configuration

Student Journal Part II: Orbital Filling Diagrams, continued 2.

You have filled in the periodic table in Part I and the orbital filling diagram for the first three principal energy levels in Part II. Fill in the maximum number of electrons that each sublevel can hold, using the correct spin for each electron. s-orbitals

p-orbitals

d-orbitals

f-orbitals

Once you have completed the diagram above, complete the table in Part I. You will identify the number of orbitals that can be found in each subshell. Using that information, you can find the maximum number of electrons that each principal energy level can hold by calculating the number of electrons that can be found in each orbital of the subshells found in each principal energy level. 3.

The diagrams on this page are similar to the orbital filling diagram that you completed previously. Describe the information that you can obtain from diagrams such as this. Be as specific as you can. You should list at least three things.

Which shape belongs to which atomic sublevel? Write your answer below the shape.

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C1EF Electron Movement and Configuration

Student Journal Part III: Writing Electron Configurations 1.

You are now ready to write the electron configurations for each element. For the first two examples, use the orbital filling diagram for the elements found in Part II of the Student Journal and write the complete electron configuration for that element below.

Oxygen:

Phosphorus: 2.

Now that you have had some practice, use what you know to write the complete electron configurations for the first three noble gases. Notice that these elements may be found in bold in the orbital filling diagram you completed in Part II.

Helium:

Neon:

Argon: 3.

Next write both the complete electron configuration and the electron configuration using a noble gas core for the following elements.

Carbon:

Nickel:

Bromine:

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C1EF Electron Movement and Configuration

Student Journal Part IV (Extension): Lewis Valence Electron Dot Structures and Electron Configurations 1.

Draw the Lewis valence electron dot structure for each of the following elements in the space above the element’s name.

bromine neon oxygen phosphorus 2.

The electron configuration for an element is 1s22s22p63s23p64s2. What are the valence electrons? How do you know they are the valence electrons?

Group 2 (group 2A) is known as the alkaline earth metals. What would the Lewis valence electron dot structures look like for each member of the alkaline earth metals? Explain your answer.

Why are Lewis valence electron dot structures important in chemistry?

Why is the maximum number of dots that may be drawn around any single element eight?

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C1EF Electron Movement and Configuration

Student Journal Part V: Effects of Chemical Properties from Configurations 1.

Observe the following elements and their configurations. Sodium (Na): 1s22s22p63s1 Magnesium (Mg): 1s22s22p63s2 Aluminum (Al): 1s22s22p63s23p1

What type of elements are they?

Observe their valence electrons. What do they have in common?

If they are trying to get to eight valence electrons, would it be more energetically favorable to gain more electrons or lose electrons?

What general conclusion can you determine about metals when they are attempting to reach their eight electrons?

Observe the following elements and their configurations. Nitrogen (N): 1s22s22p3 Sulfur (S): 1s22s22p63s23p4 Chlorine (Cl): 1s22s22p63s23p5

What type of elements are they?

Observe their valence electrons. What do they have in common?

If they are trying to get to eight valence electrons, would it be more energetically favorable to gain more electrons or lose electrons?

10. What general conclusion can you determine about nonmetals when they are attempting to reach their eight electrons?

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C1EF Electron Movement and Configuration

Student Journal Reflections and Conclusions 1.

According to the Aufbau principle, which principal energy level would have higher energy, n = 2 or n = 5?

Which element has the following complete electron configuration: 1s22s22p63s23p64s23d5?

Draw the Lewis valence electron dot structure for the element represented by the following electron configuration: [Ne]3s23p4.

Why do the electrons within a specific orbital have different spins?

What connection do you notice regarding an element’s atomic number and the number of electrons represented in that element’s complete electron configuration?

Write the complete electron configuration for the element gold (Au) in the space below. Use the diagram from Part II of your Student Guide to help you.

When looking at the representative elements (group A or main group elements), the number of electrons in the highest occupied energy level will be the same as the elements’ group number, if using the 3A-8A group numbers. These electrons are found in the s and p subshells. What is another name for these types of electrons?

Using all of the following terms, develop a graphic organizer based on what you have learned in this Explore. Use additional paper, if needed.

Terms: atom, electrons, electron orbital, principal energy level, sublevel, periodic table, periods, representative element, transition elements, lanthanide and actinide series, s, p, d, f, Lewis valence electron dot structure, valence electron, electron configuration © Accelerate Learning Inc. - All Rights Reserved

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STEMscopedia: ELECTRON MOVEMENT AND CONFIGURATION C1EF

Reflect Think about all the things in the world around you that radiate light energy—that is, things that glow or produce visible light. Some examples are light bulbs, ﬁre, neon signs, lightning, and the stars in the sky (including the Sun). Though it may not always be obvious to us, the glow of a radiating object contains special information about its identity. Why do some substances radiate light? What does this radiated light tell us about the substance? Electrons and Energy Levels in Atoms Recall that the core of an atom is its nucleus. This is where positively charged protons are located, as well as neutrons, which have zero charge. Negatively charged electrons are located outside the atomic nucleus, where they occupy speciﬁc energy levels, or orbitals. The energy of an orbital is measured in electron volts, or eV. The Bohr model of the atom chlorine, shown at right, symbolically represents these orbitals. Electrons are typically located in the lowest available energy level, also known as the ground state. In the diagram on the right, the ground state corresponds to the level n = 1. However, sometimes an electron will gain energy and jump up to another energy level. This could happen if an electromagnetic wave strikes the electron. Alternatively, it could happen if the atom collides with another atom moving at high speed.

The Bohr model of a chlorine atom shows the different energy levels at which an electron can orbit the atomic nucleus. The ring closest to the nucleus is n = 1, the middle ring is n = 2, and the outer orbital is n = 3. When the electron gains energy, it moves to a higher level; when it loses energy, it drops to a lower level.

In order for an electromagnetic wave to increase the energy of an electron, the energy of the wave must be exactly equal to the energy difference between the electron’s current energy level and another energy level. For example, consider a single electron in a hydrogen atom that exists at the ground state. The difference in energy between the levels n = 1 and n = 2 is 10.2 electron volts. If an electromagnetic wave with exactly 10.2 electron volts strikes the atom, the electron will absorb this energy and move to level n = 2. (If the amount of energy that strikes the atom is 10.1 electron volts or 10.3 electron volts, the electron will not move to level n = 2.) Alternatively, the difference in energy between levels n = 1 and n = 3 is 12.1 electron volts. So if an electromagnetic wave with exactly 12.1 electron volts strikes the atom, the electron will move to level n = 3. Note that the energy of an electromagnetic wave is directly proportional to its frequency and inversely proportional to its wavelength. So shorter waves have more energy than longer waves. These waves actually travel in “wave packets” or bundles called a photon. Each photon is composed of waves of just one wavelength.

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STEMscopedia: ELECTRON MOVEMENT AND CONFIGURATION

After an electron has moved to a higher energy level, it typically remains there for a very short time before spontaneously dropping back to a lower energy level, typically the ground state. As it drops, it releases the electromagnetic energy it had absorbed. So if the electron had absorbed energy of light corresponding to the red wavelength of visible light, it would emit red light when it returned to the ground state. If it drops to an intermediate energy level, it emits radiation with energy equal to the difference between the higher and the intermediate level. All radiation occurs when electrons drop from a higher energy level to a lower level and emit an electromagnetic wave. The electromagnetic spectrum consists of all wavelengths of electromagnetic radiation. The wavelengths that humans can see make up the visible light part of the spectrum. As you can see, each color of visible light corresponds to a different wavelength. Red light waves have the longest wavelengths and lowest energies of all visible light. The range of different waves that can be released when electrons fall from a higher energy level to a lower energy level in an atom is called the emission spectrum for that atom. Consider billions of hydrogen atoms in a region. A scientist could excite their electrons to different energy levels. In this case, there would be a speciﬁc set of light waves that these atoms would emit as the various electrons in different excited states eventually cascade down to their ground states. Scientists use spectroscopes to observe these waves. The emission spectrum for hydrogen is displayed below.

A spectroscope (or spectrometer) is a device that causes radiated light to separate into different wavelengths. So the light emitted from hot hydrogen gas would travel through a spectroscope, and the different wavelengths of light emitted from the atoms would bend (or diffract) at different angles onto the spectrum shown. Note that the visible bands of light shown on the spectrum are well defined, and their positions correspond to their specific wavelengths of light, measured in nanometers (nm). Each element on the periodic table has a unique emission spectrum. This is because the particular ways in which electrons can fall to lower energy levels depends on the specific configuration of the atom. So a scientist can pass the light emitted by any glowing gas (or very hot object) through a spectroscope and observe its unique emission spectrum. In this way, emission spectra can be used to determine the identities of unknown elements on Earth or in space.

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STEMscopedia: ELECTRON MOVEMENT AND CONFIGURATION

In order to identify the different atoms in a substance, the mass spectroscopist must ionize the substance. This can be done in many ways, but one way is by heating the substance to knock electrons out of orbit. This gives each atom a positive or negative charge. Charged atoms are ions. Then the mass spectroscopist separates the ions by charge, placing them in a chamber with an electromagnetic beam; the extent to which an ion deﬂects from the beam depends on its mass. In this way, the mass spectroscopist can measure and analyze masses and charges of unknown ions in a substance to determine the identity of the substances. One application of this mass spectroscopy is urine testing. If unusual values show up in the mass spectrum for a urine sample, the spectroscopist can compare this with the known mass spectra for different drugs to determine whether a drug is present in the urine. What do you know? The emission spectra of ﬁve unknown elements are shown below. Study these spectra, then match each emission spectrum to the appropriate element in the table. (Note: Not every wavelength shown in the spectra is included in the table.)

Element Neon (Ne) Mercury (Hg) Sodium (Na) Hydrogen (H) Helium (He)

Emission Wavelength (nm) 660, 651, 640, 627, 616, 607, 588, 585, 540 579, 576, 546, 492, 436 589 656, 486, 434, 410, 397, 389, 384 668, 588, 492, 471, 447, 444, 439

Letter

Electrons are located in energy levels. The orbitals that hold electrons are in groups known as energy levels. Each group is a different distance from the nucleus. Electrons closer to the nucleus are lower in energy than electrons farther from the nucleus. The principal energy levels are denoted by the variable n, where n = 1, 2, 3, and so on. Energy levels are divided into four sublevels: the s, p, d, and f levels. These sublevels increase in energy: s < p < d < f. Each sublevel corresponds to a particular location on the periodic table. Likewise, each sublevel can contain a maximum number of electrons. A good way to remember the order or the sublevels is with the phrase: “some pigs do fly.”

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STEMscopedia: ELECTRON MOVEMENT AND CONFIGURATION

• • • •

s: The s sublevel corresponds to the first two groups (1A and 2A) on the periodic table. Only two electrons can occupy this sublevel. Each energy level has an s sublevel (1s, 2s, 3s, and so on). p: The p sublevel corresponds to groups 3A–8A on the periodic table. This sublevel can contain up to six electrons. Energy level n = 1 does not have a p sublevel (only 2p, 3p, and so on). d: The d sublevel corresponds to the d-block of the transition elements (groups 1B–10B). This sublevel can contain up to 10 electrons. Only the energy levels n = 3, 4, 5, and 6 have d sublevels (3d, 4d, 5d, and 6d). f: The f sublevel corresponds to the inner transition elements (lanthanide and actinide series). This sublevel can contain up to 14 electrons. Only the energy levels n = 4 and 5 can have f sublevels (4f and 5f).

These energy levels and sublevels can be used to write an electron configuration for each element. An electron configuration is a written way to show the location and number of electrons in an atom. To describe the location of an electron, you must name both the energy level (n = 1, 2, 3, ...) and the sublevel (s, p, d, or f) of the electron. The sublevels are filled in order from lowest to highest energy. This is referred to as the Aufbau principle: The 1s orbital is filled first, followed by the 2s orbital, the 2p orbital, the 3s orbital, and so on. Let’s look at an atom of chlorine (Cl), which has an atomic number of 17. Therefore, a neutral chlorine atom contains a total of 17 electrons. (Remember, in a neutral atom, the number of positive charges and negative charges are equal.) Because orbitals are filled sequentially from lower energy to This diagram shows the higher energy, the electron configuration for chlorine is written as 1s2, order in which electrons fill 2s2, 2p6, 3s2, 3p5. The superscripts represent the number of electrons orbitals. Starting at the top in that energy level. If you add the superscripts for this particular of the diagram (1s), follow atom, you get 17 (2 + 2 + 6 + 2 + 5 = 17). This is a way to check to each arrow down and to see if you completed the electron configuration correctly. Your atomic the left to determine when a number or number of electrons in a neutral atom should add up all particular orbital is filled. your superscripts in your electron configuration. For example, a ground state chlorine atom has an atomic number of 17; when you add the superscripts together, you will also have a sum of 17. Remember, each energy sublevel has a maximum number of electrons it can contain.

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STEMscopedia: ELECTRON MOVEMENT AND CONFIGURATION

What Do You Think? Here are electron configurations for three different atoms. Can you identify each atom based on its electron configuration? Place the name of the element in the blank next to its electron configuration. 1. 1s2, 2s2, 2p6, 3s1 ___________________________________ 2. 1s2, 2s2, 2p6, 3s2, 3p3 ___________________________________ 3. 1s2, 2s2, 2p6, 3s2, 3p6, 4s2, 3d5 ___________________________________ 2 4 4. [Ne] 3s , 3p ___________________________________ 5. K+ (write the electron configuration of this potassium ion)

Look Out! Isotopes are atoms of the same element that have different numbers of neutrons. There are two natural isotopes of helium (He). One isotope has one neutron, and the other isotope has two neutrons. Both isotopes have two protons. Therefore, the mass number— the total number of an atom’s protons and neutrons—of the first helium isotope is three (two protons plus one neutron). The mass number of the second helium isotope is four (two protons plus two neutrons). Helium has only two isotopes, but most elements have more. Iron (Fe) has more than 25 isotopes, but many of these isotopes are not very stable. The most abundant iron isotope has 26 protons and 30 neutrons. Therefore, this isotope has a mass number of 56. Lewis Valence Electron Dot Structure Remember, only the valence electrons of an atom are responsible for interactions with other atoms. For this reason, scientists use the Lewis valence electron dot structure as a model to represent the valence electrons in an atom. In this model, the chemical symbol of an element is surrounded by dots. Each dot represents a valence electron. In general, atoms with eight electrons in their valence shell are the most stable. (This is called the octet rule.) Exceptions include hydrogen and helium; atoms of these elements can have at most two valence electrons. In a Lewis dot structure, the electrons, or dots, surround the chemical symbol in pairs. Each “side” of the chemical symbol (above, below, left, right) can contain The Lewis dot structure an electron pair, for a total of eight valence electrons. Each of the first four for chlorine (Cl) shows valence electrons occupies an unpaired space. To the right you can see three electron pairs and an example of Lewis dot structure. one unpaired electron, for a total of seven electrons.

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STEMscopedia: ELECTRON MOVEMENT AND CONFIGURATION

Atomic Structure Because each element in a group of the periodic table has the same number of valence electrons, the Lewis dot structure of all elements in a group is the same. For example, both fluorine (F) and chlorine (Cl) are found in Group 7A on the periodic table, so their Lewis dot structures contain seven electrons, as you see on the previous page. Also, elements with the same number of valence electrons will tend to act or react the same. They will also tend to have similar chemical properties. The electrons shown in Lewis dot structures represent electrons in the s and p orbitals, so Lewis dot structures are not often used to model transition elements. What Do You Know? 1. Electrons occupy the lowest energy orbital first, then move to the next one and so on (the Aufbau principle). 2. Orbitals are considered to be in the same shell if they have the same first number (no matter the order of the filling). 3. An atom will gain or lose electrons in order to have eight electrons in its outer shell (the octet rule). 4. The outer shell is the highest numbered shell that has electrons in it. Only s and p orbitals are part of the outer shell. An atom has the tendency to lose electrons (to another atom) or to gain electrons (from another atom) in order to make the outer shell complete with eight electrons. Atoms with a complete outer shell (eight electrons) are considered stable. Some atoms naturally have eight electrons in their outer shell and are very stable. (Helium is the exception, being stable with two electrons in its outer shell.)

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STEMscopedia: ELECTRON MOVEMENT AND CONFIGURATION

Complete the following chart: Element Helium (He)

Atomic Number

Carbon (C)

Electron Configuration

Lewis Dot Structure

1s2, 2s2,2p6, 3s2, 3p6

Calcium (Ca)

Ca [Ar] 4s2, 3d6

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STEMscopedia: ELECTRON MOVEMENT AND Connecting With Your Child

CONFIGURATION

Constructing a Spectrometer To help your child learn more about emission spectra, have him or her construct a basic spectrometer to observe the emission spectra of light in everyday life. To do this, you will need the following materials: • Empty cereal box (taped shut) • Sharp knife or razor • Compact disc • Masking tape • Protractor The compact disc has diffraction grating on the shiny bottom surface. When light strikes this surface, it spreads out into the different wavelengths emitted from the light source. This allows an observer to see the emission spectra from a particular light source. How to Construct the Spectrometer 1. Have your child cut a five centimeter by five centimeter square into the top of the box, on the right corner. This will be the hole through which you will observe the emission spectra. 2. Then assist in cutting an eight centimeter diagonal slit from the right corner toward the center of the box along the front face of the box. The slit should make a 30 degree angle with the right side of the box. An identical slit should be cut from the right corner toward the center of the box along the back face of the box. 3. Slide the compact disc into the slits so that the shiny surface faces upward. You should be able to see this shiny surface when you look down through the hole cut into the top of the box. 4. Next, have your child cut a five centimeter by five centimeter hole into the top-left side of the box, opposite the compact disc that was inserted on the right side of the box. When you look through this hole, you should be able to see the compact disc at the other end of the box. 5. Then use masking tape to cover this hole so there is a tiny slit, one millimeter in width, through which light can enter the cereal box. How Does It Work? The spectrometer works when light from a source passes through this tiny slit and then diffracts along the compact disc. Your child can then observe the emission spectra by looking through the hole in the top of the box and observing the pattern along the compact disc. Have your child hold the slit-end of the box to different light sources, such as an incandescent light, ﬂuorescent light, plasma screen television, computer screen, neon signs, and so forth. Here are some questions to discuss with your child: 1. Why does light spread out into different colors along the compact disc surface? 2. What do these colors reveal about the light source? 3. How does the emission spectrum differ from one light source to another? 4. Did you observe any light sources that show only one wavelength of light through the spectrometer? If not, why do you think this is? If so, how is this source different from others? 78

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C1EF Electron Movement and Configuration

Reading Science Name:

Date:

Group:

Astronomical Spectroscopy 1

Humans are fascinated by the possibility of life on other planets. Scientists think that a planet with life must meet several criteria. An atmosphere containing appropriate amounts of oxygen, carbon, and water is important. It is easy to examine Earth’s atmosphere directly. Scientists can simply collect samples at various altitudes. Each sample is then analyzed. So, how do scientists analyze the atmosphere of distant planets?

Spectroscopy measures the radiation produced by a light source. Absorption spectroscopy measures the absorption of radiation after the radiation interacts with an unknown sample. The image above is the spectrum of visible light emitted from the Sun. The dark lines are the absorption lines. The position and thickness of each absorption line corresponds to a particular ion of one element. Absorption spectroscopy is especially useful when scientists can’t analyze an unknown sample directly.

Absorption spectroscopy can be used on Earth as well as in space. Remote analysis of automobile exhaust is an example. Remote sensing analyzes the exhaust from many vehicles traveling along a road. During remote sensing, an infrared source shines a continuous infrared beam across a roadway. The detector is located directly across the road. It is connected to a computer. Hydrocarbon and carbon monoxide emissions can be measured. Data from such remote sensing provides information on pollution from cars traveling along a road. A video camera can take a picture of a car’s license plate. Police use this information to issue a ticket to a particular polluting vehicle.

Absorption spectroscopy can be used to analyze the atmosphere around planets. It was used recently to analyze a planet orbiting HD 209458. HD 209458 is a star in the constellation Pegasus. A planet named HD 209458b orbits this star. It was discovered in November 1999. The planet has an atmosphere. The planet periodically transits in front of the star. It partially blocks the star’s light from reaching Earth. The atmosphere of the planet allows some wavelengths to pass through and absorbs other wavelengths.

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C1EF Electron Movement and Configuration

Reading Science Continued 5

The Hubble Space Telescope (HST) records the spectrum of light coming from HD 209458 under two different conditions. The first condition occurs when the light is not affected by the planet. The second occurs when the planet transits in front of the star as seen from Earth. Comparing the two spectra allows scientists to determine the composition of the atmosphere around HD 209458b. The presence of water in the atmosphere of the planet is currently disputed. Studies by two different scientists disagree on whether water is present in this planet’s atmosphere. Further data collection and analysis will likely resolve the controversy.

The HST’s location is critical for absorption spectroscopy. Earth’s atmosphere contains water, carbon dioxide, oxygen, and ozone—all of which absorb radiation. A spectroscope located on Earth cannot use wavelengths absorbed by these molecules. Instead, it has to use “atmospheric windows.” Atmospheric windows are defined as ranges of wavelengths transmitted through the atmosphere. The HST is located above Earth’s atmosphere. Studies of stars make use of the unique location of the HST. The Space Telescope Imaging Spectrograph (STIS) is one part of the HST. The STIS can record spectra from stars without interference by Earth’s atmosphere.

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C1EF Electron Movement and Configuration

Reading Science 1 Scientists think that a planet with life has to – A be exactly the same size as Earth. B have an atmosphere with oxygen, carbon, and water. C have at least one moon orbiting it. D have no volcanic activity at any time.

2 Absorption spectroscopy requires – A only a radiation source and a spectroscope. B only a spectroscope. C a radiation source, a sample between the radiation source and the spectroscope, and a spectroscope. D a radiation source, a spectroscope, and a sample covered by the spectroscope box.

3 Which of the following is an example of remote sensing? A A technician analyzing the car exhaust of the car connected directly to his emission testing machine. B A chemistry student measuring the pH of a solution using a probe dipped into the liquid. C A physics student timing the drop of two spheres with different mass using a photogate. D A scientist analyzing airplane exhaust from all airplanes landing at an airport.

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C1EF Electron Movement and Configuration

Reading Science 4 To analyze the atmosphere of the planet HD 209458b, the Hubble Space Telescope collected data – A when the planet was between its star and Earth and when it was not. B when the star was between the planet and Earth and when it was not. C when our Moon was between the star and Earth and when it was not. D when the star, the planet, and our Sun aligned.

5 What is the disagreement regarding the atmosphere of HD 209458b? A Whether or not it has clouds B Whether it is thicker than three miles or not C Whether or not it contains water D Whether or not it is symmetrical

6 The STIS is well suited for analyzing the spectra of stars because it is – A closer to the stars than a spectroscope on Earth. B not blocked by lunar eclipses. C more sensitive than spectroscopes on Earth. D outside Earth's atmosphere.

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C1EF Electron Movement and Configuration

Math Connections Name:

Date:

Group:

Part I: Light and the Electromagnetic Spectrum The electromagnetic spectrum consists of a continuous range of wavelengths and frequencies, from radio waves at the low-frequency end to gamma rays at the high-frequency end. Electromagnetic energy is characterized by a frequency, a wavelength, and an amplitude. The speed of a wave is equal to the wavelength multiplied by its frequency. Wavelength x Frequency=Speed λ (m)×v (s−1)=c (m/s) where c=3.00×108 m/s (speed of light) 1. Using the equation above, solve for wavelength.

2. Using the equation above, solve for frequency.

3. The light blue glow given off by mercury streetlamps has a wavelength of 436 nm. What is its frequency in Hertz? (1 Hz=1 s−1)

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C1EF Electron Movement and Configuration

Math Connections 4. What is the frequency of a gamma ray with a wavelength of 3.56×10−11 m? What is the frequency of a radar wave with a wavelength of 10.3 cm?

5. What is the wavelength (in meters) of an FM radio wave with frequency v=102.5 MHz? What is the wavelength of a medical X-ray with v=9.55 x 1017 Hz?

Part II: Colors and Energies of Electromagnetic Radiation Atoms give off light when heated, thus providing clues as to their atomic makeup. The table below gives the wavelengths and colors of light in the visible spectrum. The energy of a photon of light can be calculated using the de Broglie equation. E=hf=(hc)/λ E=energy of a photon of light h=Planck’s constant=6.626×10−34 J c=speed of light=3.00×108 m/s λ=wavelength Table 1. The colors and wavelengths of light in the visible spectrum.

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Color

Wavelength range (nm)

Violet

390–455

Blue

455–492

Green

492–577

Yellow

577–597

Orange

597–622

Red

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C1EF Electron Movement and Configuration

Math Connections Use the information above to solve the following problems. 6. A photon of wavelength 500 nm is emitted by a traffic light. What are the energy and frequency of the photon? What is the color of the light?

7. Four sources emit light of the following wavelengths: 400 nm, 580 nm, 650 nm, and 475 nm. What are the colors of light emitted by the sources? Which source emits photons with the highest energy? Which emits photons with the lowest energy?

8. I have a glass tube filled with hydrogen gas. I shine white light onto the tube. The spectrum I then measure has an absorption line at a wavelength of 474 nm. What are the frequency and energy of the photon?

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C1EF Electron Movement and Configuration

Writing Science Name:

Date:

Group:

LOOK

THINK Think about the relationship of light emission and the movements of electrons to element identification. The development of the modern atomic theory would not have been possible without the work of Max Planck. The Planck constant was first described as the proportionality constant between the energy (E) of a photon and the frequency (f ) of its associated electromagnetic wave. Atomic structure is where chemistry meets physics. Niels Bohr used Planck’s mathematical expression to calculate the distance between energy levels of atoms. It was believed that orbiting electrons could be “bumped” between energy levels by free electrons if enough energy was transferred, which would cause a change in energy, allowing the atom to emit light during an excited state. WRITE Describe how emission spectra lines are produced, what they indicate about elements, and how they can be used to identify elements. Be sure to • clearly state your central idea; • organize your thoughts; • develop your essay in detail; • choose your words carefully; and • use correct spelling, capitalization, punctuation, and grammar.

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C1EF Electron Movement and Configuration

Writing Science

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High School Chemistry

C1G

Electron Patterns and Properties of Elements

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C1G Electron Patterns and Properties of Elements

Student Journal Name:

Date:

Group:

Background 1.

What is a periodic trend?

To what can the similar properties of the elements in a family be attributed?

What are the three concepts that can be used to explain the different periodic trends?

Explain the two types of correlations.

If you worked at a job for nine dollars per hour, what type of correlation is the number of hours worked to your take-home pay?

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C1G Electron Patterns and Properties of Elements

Student Journal Part I: It’s Just a Trend 1.

What is the trend with atomic numbers of the elements as you move from left to right across a period on the periodic table? How does this sequence continue to the next period?

What is the trend for the number of valence electrons moving from left to right across a period on the periodic table? What is the pattern or trend within a group?

Na+

Cl-

Atomic radius = 186 pm

Ionic radius = 95 pm

Atomic radius = 99 pm

Ionic radius = 181 pm

Use the diagrams above to answer questions 3 and 4. 3.

What happens to the radius of the sodium atom when it forms a positive ion (cation)? What happens to the electrons?

What happens to the radius of the chlorine atom when it forms a negative ion (anion)? What happens to the electrons?

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C1G Electron Patterns and Properties of Elements

Student Journal Part I: It’s Just a Trend, continued 5.

Match each of the four trends listed to its description by writing the correct letter of the description next to the trend’s name.

Trends _____ ionic radius _____ electronegativity _____ ionization energy

_____ atomic radius

Descriptions A.

the amount of energy needed to remove an electron from a neutral atom

the radius of a neutral atom

the radius of a positive or negative ion

the ability of an element to attract electrons when forming a chemical bond

Use the diagrams shown below to answer questions 6 and 7. 6.

Which diagram represents the periodic trend for ionization energy? __________________

Which diagram represents the periodic trend for electronegativity? ___________________

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Increasing Increasing

Increasing

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C1G Electron Patterns and Properties of Elements

Student Journal Part II: Putting It All Together Use the Student Reference Sheet and your graphs to help you answer the following questions. Do you recognize any trends from this information alone? 1.

Which element below has the smallest atomic radius? ________________________________ beryllium (Be)

oxygen (O)

sodium (Na)

rubidium (Rb)

potassium (K)

Which element below has the lowest electronegativity? _______________________________ calcium (Ca)

fluorine (F)

Which element below has the smallest ionic radius? _________________________________ lithium (Li)

sodium (Na)

gallium (Ga)

selenium (Se)

bromine (Br)

Which element below has the highest ionization energy? ______________________________ strontium (Sr)

beryllium (Be)

magnesium (Mg)

calcium (Ca)

Which element has the greatest electronegativity on the periodic table? __________________

Which representative element has the largest atomic radius? ___________________________

Place the elements below in order of decreasing ionization energy. aluminum (Al)

chlorine (Cl)

magnesium (Mg)

sulfur (S)

[largest] _____________ _____________ _____________ _____________ [smallest] 8.

Place the elements in order of increasing atomic radius. gallium (Ga)

nitrogen (N)

indium (In)

tellurium (Te)

[smallest] _____________ _____________ _____________ _____________ [largest] 9.

Choose which of the two is larger, the atom or its ion. A. B.

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magnesium atom (Mg) oxygen atom (O)

OR OR

magnesium ion (Mg2+) _____________ oxygen ion (O2-) _______________

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C1G Electron Patterns and Properties of Elements

Student Journal Part II: Putting It All Together, continued Interpret this graph of electronegativities and answer the questions below.

10. What do you notice about the order of the elements on this graph? How do the rows on this graph relate to the periodic table?

11. What element shown on this graph has the greatest electronegativity? ___________________ 12. Which element shown on this graph has the lowest electronegativity? ___________________ 13. Explain how the information shown on this graph can help you identify the periodic trend of electronegativity on the periodic table.

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C1G Electron Patterns and Properties of Elements

Student Journal Part II: Putting It All Together, continued 14. Now that you have completed your graphs, analyze them to draw vertical arrows on the right and horizontal arrows on the top of the periodic table below to show the direction of the trends. Follow these instructions. • • • •

Draw blue arrows for the increasing atomic radii trend. Draw red arrows for the increasing ionic radii trend. Draw black arrows for the increasing first ionization energy trend. Draw brown arrows for the increasing electronegativity trend.

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C1G Electron Patterns and Properties of Elements

Student Journal Reflection and Conclusions 1.

Across a period, the ionic radii of the anions is larger than that of the cations. Use what you know about ions and electric charges to help explain this trend.

Fluorine is the most electronegative element on the periodic table. Use what you know about the halogens to help explain why fluorine is such a highly electronegative element. (Hint: It has to do with electrons.)

The ionization energy of an element relates to the amount of energy that is required to remove an electron from a neutral atom. The first ionization energy is the amount of energy it takes to remove the first electron. Use the trends you have learned to explain why the first ionization energy for lithium is 520 kJ/mol and the first ionization energy for oxygen is 1,314 kJ/mol. What characteristics of these two neutral atoms could account for the large difference in energy?

On another sheet of paper, use all of the following terms to develop a graphic organizer based on what you have learned in this Explore.

Terms: periodic table, elements, properties, groups, periods, column, row, trends, electronegativity, ionization energy, atomic radius, ionic radius, increasing, decreasing, shielding effect, neutral, cation, anion © Accelerate Learning Inc. - All Rights Reserved

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STEMscopedia: ELECTRON PATTERNS AND

PROPERTIES OF ELEMENTS C1G

Reflect In your study of science, you have probably learned about many different cycles, or patterns that repeat over time. All organisms have life cycles, for example. A frog is born from an egg, grows into a tadpole, and finally becomes an adult frog capable of reproducing by making new eggs. This pattern repeats itself throughout each generation of frogs. The life of a frog is cyclical. Another term that means cyclical is periodic. You are probably already familiar with this term from your study of chemical elements, which are arranged into rows (periods) and columns (groups) on the periodic table. This arrangement is periodic because each row or column contains elements with similar properties, and these properties follow predictable patterns as you move across the table. What are some properties of elements you have already learned about? Can you think of any trends, or patterns involving these properties? (You can review the periodic table at the top of the next page.) The Periodic Table To date, scientists have confirmed the discovery of at least 114 chemical elements. (The discoveries of several more elements are awaiting confirmation.) Based on an element’s location on the periodic table, you can draw several conclusions and make several predictions about the element. Some versions of the periodic table contain more information than others. However, nearly all versions provide three basic pieces of information about each element: • Chemical symbol: Each element is given a unique one- or two-letter abbreviation. The chemical symbol for hydrogen is H, while the chemical symbol for helium is He. • Atomic number: This whole number reveals the number of protons in an atom of the element. Each element is identified by its unique atomic number. • Atomic mass: An atom’s mass is roughly equivalent to the sum of its protons and neutrons. (Electrons are extremely small, so their masses are negligible.)

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An atom of helium (He) has two protons and an atomic mass of approximately 4.

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STEMscopedia: ELECTRON PATTERNS AND

PROPERTIES OF ELEMENTS

As you can see, the atomic number of elements increases as you move from left to right across each period (or row) of the periodic table. Periods and groups also provide information about the number of electrons that fill an element’s energy valence shell: an levels. As a general rule, elements in the same period have the same atom’s outermost number of energy levels. Elements in the same group have the same energy level number of electrons in their valence shells, which means they tend to behave similarly during chemical reactions. covalent bond: the connection between Trends in Atomic Size two atoms that share We can describe additional trends based on an element’s location in electrons the periodic table. The size of an atom determines its atomic radius, or one-half the distance between atoms in a covalent bond. Two factors determine the size of an atomic radius: the number of energy levels in an atom and the number of protons in an atom’s nucleus. • Energy levels: Atoms farther down in a group have more energy levels than atoms farther up in the same group. Because each energy level increases an atom’s atomic radius, atoms at the top of a group have smaller radii than atoms at the bottom of the group. You can see this trend in the figure on the next page.

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STEMscopedia: ELECTRON PATTERNS AND

PROPERTIES OF ELEMENTS

• Number of protons: Protons have positive charges and electrons have negative charges, and opposite charges attract. Therefore, the nucleus of an atom with more protons exerts a greater attractive force on the atom’s valence electrons, pulling the electrons closer to the nucleus. Atoms farther to the right in a period have more protons than atoms farther to the left in the same period. Therefore, atomic radius decreases when moving from left to right across the periodic table. You can see this trend in the figure at right. Trends in Ionic Size An ion is an atom that has acquired an electric charge by gaining or losing electrons. The size of an ion depends on the ion’s charge. A positively charged cation is always smaller in size than its parent atom. This is because a In general, atomic radius decreases as one cation forms when an atom loses electrons; the moves from left to right across the periodic resulting cation has more protons than electrons table and increases as one moves from top and therefore a net positive charge. This to bottom. increase in positive charge pulls the cation’s valence electrons toward its nucleus; as a result, the ionic radius is smaller than the radius of the parent atom. The greater the positive charge, the greater the pull on the valence electrons and the smaller the ionic radius. In contrast, an anion—a negatively charged ion—is larger in size than its parent atom. An anion has more electrons than protons. Because like charges repel, the stronger negative charge pushes the electrons away from each other, giving an anion a larger ionic radius than its parent atom’s radius. Anions with greater negative charges have larger ionic radii. Finally, ions toward the bottom of a group have greater ionic radii than ions toward the top of the same group. This is because ions farther down in a group have additional electron shells.

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STEMscopedia: ELECTRON PATTERNS AND What Do You Think?

PROPERTIES OF ELEMENTS

The following diagram outlines the top six periods of the periodic table. Based only on their locations on the periodic table, atoms of which element have the largest atomic radii: A, B, or C?

Trends in Electronegativity Some elements are able to attract electrons more easily than others—these elements have greater electronegativities. An American scientist named Linus Pauling established an electronegativity scale for each element, as shown below. Atoms with larger numbers according to Pauling’s scale have greater electronegativities and more easily attract electrons.

These values suggest two trends. First, the electronegativity of atoms typically decreases as you move down a group. Second, the electronegativity of atoms increases as you move left to right across a period.

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STEMscopedia: ELECTRON PATTERNS AND

PROPERTIES OF ELEMENTS

Trends in Ionization Energy Energy is needed to remove an electron from the outer electron shell of a neutral atom. This energy is the atom’s ionization energy because a neutral atom that loses an electron becomes an ion. The first electron removed from a neutral atom is referred to as the first ionization energy. For example, here is the chemical equation representing the first ionization of a sodium atom: [Ionization Energy]+Na

Na++e−

The ionization energy converts the sodium atom (Na) into a sodium cation (Na+) by removing an electron (e−). Protons in an atom’s nucleus keep electrons in orbit around the nucleus. Moving down a group on the periodic table, an atom’s valence shell gets farther from the atom’s nucleus. As a result, less energy is necessary to remove a valence electron from atoms farther down in a group. In other words, as you move down a group on the periodic table, the first ionization energy typically decreases.

Ionization energy is the energy necessary to remove a valence electron from an atom.

Ionization energy typically increases, however, as you move from left to right across a period. This is because atoms farther to the right on the periodic table have more protons. As a result, their larger nuclei exert greater attractive forces on their electrons.

Look Out First ionization energy refers to the removal of only the first electron from a neutral atom. Trends relating to ionization energy are slightly different for ions. For example, a neutral sodium atom (Na) has only one electron in its valence shell. Once this electron is removed, the resulting sodium cation (Na+) has eight electrons in its valence shell, as shown in the diagram. According to the octet rule, atoms with eight valence electrons are less likely to gain or lose more electrons. As a result, much more energy is necessary to remove an electron from Na+ than from Na. A neutral sodium atom (Na) has 11 electrons distributed among 3 energy levels.

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Scientists in the Spotlight: Linus Pauling The American chemist Linus Pauling (1901–1994) won the 1954 Nobel Prize in chemistry for his work on chemical bonds. Pauling challenged the contemporary view of how elements are connected. Pauling used an interdisciplinary approach to describe his concept of electronegativity. For example, he borrowed from the study of physics to describe the structure of atoms and molecules, including the angles at which chemical bonds form and the distances between atoms in compounds. In addition to his work on chemical bonds, Pauling expanded his knowledge into many different areas of science. He used X-ray crystallography to understand protein structure, discovering that proteins often contain two different types of folds: regions that twist into coils (known as alpha helixes) and regions that fan out like folded paper (known as beta sheets). Pauling also studied medicine and health, developing our understanding of anesthesia, sickle-cell anemia, and the importance of vitamin C in the human diet. Linus Pauling used his popularity as a Nobel-winning scientist to publicize his stance against nuclear testing and proliferation. For this work, Pauling won the Nobel Peace Prize in 1962. He is the only person to have individually won two Nobel prizes.

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Pauling discovered the structure of hemoglobin, a protein in the blood, is made of many alpha helices.

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STEMscopedia: ELECTRON PATTERNS AND

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What Do You Know? Compare ionization energy with electronegativity. Read the list of terms in the box below. Write each term in the correct place on the Venn diagram. Ionization Energy vs. Electronegativity • decreases going down a column

• is described by this equation: X X++e−

• is the energy needed to remove an electron

• is the greatest for the element fluorine

• increases going across a row

• is the greatest for the element helium

• is a value determined from an atom’s ability to attract an electron

• is the smallest for the element cesium

• was identified by Linus Pauling

Ionization Energy Electronegativity

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Connecting With Your Child Modeling Periodic Trends Your child can learn more about trends in the periodic table by creating a three-dimensional model of one of the four trends discussed in this companion: atomic radius, ionic radius, ionization energy, and electronegativity. First, you will need to locate a copy of the periodic table that describes the trend you wish to model. (You should be able to find such a copy online.) Your child may use pipe cleaners, sticks, or straws of various sizes to illustrate the magnitude of each property and represent it in three dimensions. For example, here are the atomic radii, measured in picometers (pm), of the elements in Group 1 (sometimes labeled 1A), from top to bottom: • Lithium (Li): 157 pm • Sodium (Na): 186 pm • Potassium (K): 231 pm • Rubidium (Rb): 244 pm • Cesium (Cs): 262 pm Your child may represent each 100 pm with 1 inch of pipe cleaner. In this case, the pipe cleaner for Li would measure approximately 1.6 in. Once your child has measured the appropriate lengths for each element, your child may compare them to see the relative size of each element, as measured by atomic radius. Here are some questions to discuss with your child: • How can models help you to better understand periodic trends? Explain. • Could you use your three-dimensional model to identify other trends in the periodic table? Explain. • Based on your model, what are the trends going down a group or across a period for the property you chose? Does this trend hold for each group and/or period? Explain.

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C1G Electron Patterns and Properties of Elements

Reading Science Name:

Date:

Group:

Ionization Energy

The periodic table is an amazing reference that gives a wealth of information at a glance. You can identify elements based on their chemical symbol, atomic number, and atomic mass. You can also classify elements by the patterns that are found within the table’s structure. Periodic trends are patterns that occur across a row (from left to right) or down a column of the periodic table. They can be used to predict certain properties of elements in their atomic or ionic form. These trends are based in large part on the number of protons and electrons found in the atoms. Periodic trends include such properties as atomic radius, electronegativity, and ionization energy.

As you may know, each element has a unique number of protons in the nucleus. The number of protons is the same as the atomic number of each element. Each of these protons has a positive charge, and the more protons that are in the nucleus of an atom, the greater the positive charge. Each neutral atom will have the same number of electrons as protons. Therefore, for neutral atoms, the atomic number will equal the total number of electrons in that atom. However, it is the valence electrons of the atom that become important. Remember that the valence electrons are the electrons found only in the outermost electron shell of an atom. These valence electrons will not only contribute to the possible charge of the atom, but they are also responsible for many of the periodic properties and trends of elements.

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Reading Science Continued 3

Period 1 elements can hold two electrons in their electron orbital, giving those atoms one or two valence electrons. The elements from periods 2 and 3 can hold from 1–8 valence electrons. The elements from periods 4 and 5 can have up to 18 outer-shell electrons. The elements from periods 6 and 7 can have up to 32 outer-shell electrons. When an electron shell contains all of the electrons it can hold, it is said to be “full” and is in its most stable state.

As electrons are removed or gained, the atom will develop a charge based on the number of electrons removed or gained. Ions with a positive or negative charge will be created. As you can imagine, it takes energy to move electrons. Ionization energy is one type of energy required to move electrons. Ionization energy refers to the amount of energy required to remove an electron from a neutral atom. The energy required to remove the first electron is called the first ionization energy. It is represented by the graphs at the beginning of this passage.

Consider the element neon (Ne), a noble gas found in group 8A, with 10 protons and 8 valence electrons. As this element has all of the valence electrons its outermost electron orbital can hold, it is in its most stable state. Neon does not “want” to lose its electrons. All noble gases have very high ionization energies, meaning it is very difficult to remove any of their valence electrons. This is one of the reasons why noble gases are considered nonreactive. On the other hand, sodium (Na) in the next period has 11 protons but only 1 valence electron. It does not take much energy to remove this electron from this atom. Sodium, therefore, has a very low ionization energy, as do all of the group 1A elements. The element to the right of neon, magnesium (Mg), has 12 protons and 2 valence electrons. There is more of a positive force pulling on the valence electrons of this atom. This extra pull makes it more difficult to remove an electron from magnesium than sodium. These examples demonstrate that the trend for ionization energy increases, moving from left to right across a period.

There is another dimension to this periodic trend. As you move down a column (group) on the periodic table, each element will contain more electron shells. This means that there will be more of a buffer between the positive protons and the negative valence electrons in the reactive outer shell. For example, argon (Ar) is just below neon (Ne) in group 8A. Compared to neon, argon has an additional electron shell between its protons and valence electrons. Therefore, argon has a lower ionization energy than neon. The same goes for the element calcium (Ca). Calcium has an additional electron shell between its protons and valence electrons. This gives calcium a lower ionization energy than magnesium. Therefore, the trend for ionization energy decreases as you move down a group on the periodic table.

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C1G Electron Patterns and Properties of Elements

Reading Science 1 Paragraph 2 discusses some of the properties of the elements that are found on the periodic table. What atomic component contributes to many of the trends of elements? A The protons B The total number of electrons C The valence electrons D The neutrons

2 Place the following elements in order of increasing ionization energy. Their elemental symbols are listed below. Argon (Ar), Magnesium (Mg), Phosphorus (P), and Chlorine (Cl) A Mg, P, Cl, Ar B Ar, Cl, P, Mg C Cl, P, Ar, Mg D Not enough information is given

3 Which of the following is not a trend found on the periodic table? A Atomic radius B Electronegativity C Radioactivity D Ionization energy

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Reading Science 4 One of the trends, or patterns, that occurs on the periodic table is that of ionization energy. What does ionization energy refer to? A The amount of energy required to add an electron to an atom. B The amount of energy required to remove a proton from an atom. C The amount of energy required to remove an electron from a neutral atom. D All of the above.

5 As you move from left to right across a period on the periodic table, there is an increase in the positive charge of the nucleus of the atom, which contributes to the increase in the first ionization energy of atoms. Why is there a decrease in ionization energy as you move down a column (group) on the periodic table? (Refer to paragraph 6.) A The valence electrons are not as negative. B The protons become negative. C There are fewer electron orbitals between the protons and the valence electrons. D There are more electron orbitals between the protons and the valence electrons.

6 Place the following elements in order of decreasing ionization energy. Their elemental symbols will be listed below. Barium (Ba), Neon (Ne), Calcium (Ca), and Bromine (Br) A Ne, Br, Ca, Ba B Ne, Ca, Br, Ba C Ba, Ca, Br, Ne D Ba, Br, Ca, Ne

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C1G Electron Patterns and Properties of Elements

Math Connections Name:

Date:

Group:

Atomic Radius and Electronegativity of Group 2: Alkaline Earth Metals 1. Plot atomic radius and electronegativity versus atomic number on the graph below. Use this graph to answer the following question. Element

Atomic Radius (pm)

Electronegativity

Beryllium

112

1.57

Magnesium

145

1.31

Calcium

194

Strontium

219

0.95

Barium

253

0.89

Radium

283

0.9

2. E xplain the relationship between atomic radius and electronegativity as you move down a group in the periodic table. Provide rationale for your response using data from the table and graph.

Electronegativity

Atomic radius (pm)

Atomic Radius and Electronegativity of Alkaline Earth Metals

Atomic Number

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C1G Electron Patterns and Properties of Elements

Math Connections Periodicity – Atomic Radius and Electronegativity of Period 4 Elements 3. Plot atomic radius and electronegativity versus atomic number on the graph below. Use this graph to answer questions 4–6. Element

Atomic Radius (pm)

Electronegativity

Potassium

243

0.82

Calcium Titanium Vanadium Chromium Manganese Cobalt Nickel Copper

194 176 171 166 161 152 149 145

1 1.54 1.63 1.66 1.55 1.88 1.91 1.9

Zinc

142

1.65

Germanium Arsenic Bromine Krypton

125 114 94 88

2.01 2.16 2.96 3

Electronega vity

Atomic radius (pm)

Atomic Radius and Electronega vity of Elements in Period 4

Atomic Number 112

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C1G Electron Patterns and Properties of Elements

Math Connections 4. What is the relationship between the size of atomic radii and electronegativity as you move across the periodic table from left to right?

5. Write an inequality statement that compares the electronegativity of the atoms. a. Mg, Al b. K, Fr c. Tl, Ga d. Pt, Lu 6. What is the trend in the size of atomic radii as you move from right to left and top to bottom on the periodic table?

7. Order the following atoms from the greatest to the smallest atomic radius. In, V, W, Re, B, K, Rb, Cs, Cl, I, H, Be

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Math Connections Trends in First Ionization Energy of Noble Gases Ionization energy is the the amount of energy required to remove an electron from a neutral atom.

Element

Atomic Number

1st Ionization Energy (K/kJ mol-1)

Atomic Radius (pm)

Helium

2373.3

Neon Argon Krypton Xenon Radon

10 18 36 54 86

2080.6 1520.4 1350.7 1170.4 1037.1

38 71 88 108 120

8. Use the data from the table to chart the relationship between the atomic number and the first ionization energy of the element.

Atomic Radii

1st Ionization Energy

Trends in Ionization Energy within the Same Chemical Family

Atomic Number 9. Describe the trend among chemical families you see in the graph.

10. Look at the period across the periodic table. Describe the trend in ionization energies.

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C1G Electron Patterns and Properties of Elements

Writing Science Name:

Date:

Group:

LOOK

THINK The periodic table can be used to identify and explain periodic trends in elements, such as atomic and ionic radii, ionization energy, and electronegativity. When creating the first periodic table of elements, Dmitri Mendeleev predicted that the elements could be arranged in an order that would predict their properties. After he analyzed the trends, he was even able to predict the placements of elements that had not yet been discovered. This illustration depicts the trends identified in the arrangement. The term “electron affinity” refers to an element’s electronegativity. Notice that, as the atomic radii decrease, the ionization energy and electronegativity of the element increase, which forms an inverse relationship. An element’s classification as metal or non-metal also shows trends in the table. Metallic characteristics increase with the atomic radii, while non-metallic ones increase as the atomic radii decreases. WRITE Analyze the trends of the periodic table, and then explain the relationship that exists between an element’s ionization energy and its electronegativity. Be sure to • clearly state your central idea; • organize your thoughts; • develop your essay in detail; • choose your words carefully; and • use correct spelling, capitalization, punctuation, and grammar.

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C1G Electron Patterns and Properties of Elements

Writing Science

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High School Chemistry

C2AB

Inter- and IntraMolecular Forces

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C2AB Inter- and Intra- Molecular Forces

Student Handout Name:

Date:

Place a penny on the table.

With a dropper, drop and count how many drops of water can be added to the surface of the penny without spilling over.

Repeat step 2 with ethanol, using a different penny and dropper.

Repeat step 2 with hexane, using a different penny and dropper.

Data Liquid

Drops

water ethanol hexane

Analysis Compare and contrast the three different liquids used in the experiment and their molecular structure. Give an explanation for your observations.

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C2AB Inter- and Intra- Molecular Forces

Student Handout Structure of Substances

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C2AB Inter- and Intra- Molecular Forces

Student Journal Name:

Date:

Group:

Step 1. Question

Step 2. Relevance

Step 3. Variables, if applicable Independent variable _________________________________________________________ Dependent variable __________________________________________________________ Control variable(s) ___________________________________________________________ Step 4. Hypothesis

Step 5. Materials

Step 6. Safety considerations

Step 7. Procedure

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C2AB Inter-and Intra- Molecular Forces

Student Journal Data and Results

Analysis

Conclusion and Scientific Explanation Write a scientific explanation to describe the data you gathered in this investigation.

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C2AB Inter- and Intra- Molecular Forces

Student Journal Name:

Date:

Group:

Station 1: London Dispersion Force Comparison 1.

Fill in the table. Sample

State of Matter

Strength Order

Predicted Substance

A B C 2.

Look up the following molecules and draw their structural formula. A.

propane (C3H8)

hexane (C6H14)

paraffin (C18H38)

Post-Activity Questions 3.

Explain your reasoning for the predicted substances.

Does the size of the molecule affect the strength of the London dispersion force? How?

How are London dispersion forces caused?

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C2AB Inter-and Intra- Molecular Forces

Student Journal Station 2: Hydrogen Bonding Comparison 1.

Fill in the table. Sample

Swirl

Shake

Prediction

D E F 2.

Look up the following molecules and draw their structural formula. A.

ethanol (C2H5OH)

water (H2O)

glycerin (C3H5(OH)3)

Post-Activity Questions 3.

Explain your reasoning for the predicted substances.

What characteristics determine if a molecule will form hydrogen bonds?

Identify the number of hydrogen bonds each molecule above can make.

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C2AB Inter- and Intra- Molecular Forces

Student Journal Station 3: Evaporation Comparison 1.

Fill in the table. Sample

Rate of Evaporation

Prediction

G H I 2.

Look up the following molecules and draw their structural formula. A.

isopropyl alcohol (C3H7OH)

water (H2O)

methanol (CH3OH)

Post-Activity Questions 3.

Explain your reasoning for the predicted substances.

What are the intramolecular and intermolecular forces in each compound?

What can be concluded about the strength of intermolecular forces and evaporation?

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C2AB Inter-and Intra- Molecular Forces

Student Journal Reflection and Conclusions 1.

List the intermolecular forces in increasing order.

What physical properties were used to identify the intermolecular forces in the substance? Why are these properties useful?

You are given an unknown white crystal. Give some examples of tests you can run to determine if it is an ionic or covalent compound and what intermolecular forces occur.

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STEMscopedia: INTER- AND INTRA-

MOLECULAR FORCES C2AB

Reflect Aubrey purchased a new car and wants to make sure she has the necessary tools to keep it clean. She buys a bucket, car soap, tire cleaner, washer fluid additive, and a sponge. She loves to wash her car in the driveway. She knows when it’s time for another wash because the additive she uses causes water to bead on the windows. When dew or rain sticks to the windshield, it’s time to wash the car again. How does the additive she uses cause the water to bead and fall off the car so easily? Why does water stick to the car when the additive needs to be replaced? Intramolecular and Intermolecular Forces Intramolecular forces create the bonds that hold a molecule together. Just as intramural sports involve competition of teams from one school, intramolecular forces involve the forces that bind atoms together into one molecule or formula unit. When atoms bond together into a compound, they acquire new properties that are different from the properties of the original atoms. Intramolecular force influences a substance’s chemical properties.

Chlorine is a poisonous gas at room temperature.

Consider the elements sodium and chlorine. Sodium is highly reactive and can explode in water. Chlorine is a poisonous gas. They must both be handled with extreme caution. When chlorine and sodium atoms bond together to form the ionic compound, sodium chloride, they are no longer considered hazardous to handle. Sodium chloride is ordinary table salt that is added to food every day. When you put it in water, there is no explosion or violent chemical reaction. All it does is dissolve in water. Sodium is a soft gray metal that is highly reactive in water. Salt crystals are sodium chloride.

Intermolecular forces act between molecules and influence the physical properties. Intermolecular bonds form between different molecules, much as interstate highways connect different states. These forces form bonds that are weaker than those bonds that are intramolecular. Intermolecular bonds are responsible for melting and boiling points, surface tension, friction, viscosity, state of matter, and many other physical properties.

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MOLECULAR FORCES

What do you know? Compare and contrast the intramolecular and the intermolecular properties for three forms of water in the table below: Substance

Intramolecular (chemical properties)

Intermolecular (physical properties)

Ice

Liquid Water

Steam

Chemical and Physical Properties of H2O Questions: 1. Which type of force, intramolecular or intermolecular, has the most similar properties for all the forms of water? Justify your answer.

2. Which type of force, intramolecular or intermolecular, has the least similar properties for all the forms of water? Justify your answer.

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STEMscopedia: INTER- AND INTRA-

MOLECULAR FORCES

Connecting With Your Child Plan and carry out these investigations with your child to infer the strength of intermolecular and intramolecular forces. Activity 1 Materials needed: • Glass of water • Salt • 2 Large paper clips • 9-volt battery Procedure: Watch this video on YouTube to see how to set up the materials, and then try this at home: How to Separate Hydrogen and Oxygen From Water https://www.youtube.com/watch?v=7CHSvcdoVGU Activity 2 Materials needed: • 2 Small glasses of water • 2 Droppers • 2 Pennies • Dish soap Procedure: • Put two drops of dish soap in one of the glasses of water. • Place the two pennies on the table, face up. • Use one dropper to add as many drops of water to the penny as you can until it overflows. Be sure to count the drops. • Repeat the last step with the soapy water and the other penny.

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STEMscopedia: INTER- AND INTRA-

MOLECULAR FORCES

Number of Drops of Water on Penny

Number of Drops of Soapy Water on Penny

Draw a picture of what you observed.

Questions: 1.

Which activity broke down intramolecular forces? Explain.

Which activity broke down intermolecular forces? Explain.

Which type of force seemed to be more difficult to break apart and why?

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C2AB Inter- and Intra- Molecular Forces

Reading Science Name:

Date:

Group:

Carbon 1

Carbon is one of the oldest recognized elements on Earth. In fact, the word “carbon” is derived from the Latin word “carbo” for coal. It is a second period element with the chemical symbol C and with an atomic number of 6 (meaning that it has 6 protons in the nucleus). Carbon is a nonmetal and forms covalent bonds with other elements. However, carbon is an unusual element. In fact, there would be no life on Earth without the element carbon. If you think about it, carbon is everywhere. Carbon is found in rocks, in the atmosphere, and in every biological cell. So what makes this single element so critical and so unusual? It has to do with the way carbon bonds to other elements.

Carbon has four valence electrons in the second period of the periodic table, which opens this element to all sorts of bonding capabilities. Carbon can form not only single bonds but double and triple bonds as well, giving it a wide variety of compound combinations. Carbon has the ability to form long chains, rings, and branches. Molecules containing carbon might only contain two atoms or there might be thousands. It is the building block of the four main biological molecules (carbohydrates, lipids, proteins, and nucleic acids), as well as many of the inorganic molecules found in rocks and gases. In fact, almost 80 percent of the known molecules on Earth contain some form of the element carbon. This is over 20 million compounds!

So why is carbon so good at bonding with other atoms? Carbon has a ground state electron configuration of 1s2 2s2 2p2. Carbon has four available valence electrons for bonding. This gives it something called tetrahedral geometry. What does this mean? The four valence electrons and tetrahedral geometry give carbon an equal chance of bonding with elements on all four sides. It can bond to a wide variety of other nonmetals, making it extremely versatile as well as thermodynamically stable. Therefore, the fact that carbon can form multiple stable covalent bonds gives it unique bonding abilities that are different from the other nonmetals.

Carbon can also bond to other carbon atoms in long rings or chains. This is perhaps one of the most critical bonding properties of this element. As a result, it can form very short (two) or very long (thousands) carbon chains. The unique chains and rings that carbon can form are what makes it critical to biological organisms. Carbon is a major component in the organic molecules necessary for life. Again, carbon is an element that is found in every living cell. This is why biological organisms are called carbon-based life forms. This speaks volumes not only for its importance, but also for its prevalence. Lipids, or the fats in biological organisms, are made of long chains of carbon compounds known as hydrocarbons. Amino acids, the building blocks of proteins, contain central carbon atoms. Nucleotides, the building blocks of both DNA and RNA, contain a central carbon-nitrogen ring. Carbohydrates, or sugars, also contain either rings or chains of carbon. From this, it is easy to see that without the element carbon and its many bonding properties, there would be no life on Earth.

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Reading Science 5

However, there is another interesting aspect of carbon and its bonding capabilities. Carbon can exist in three distinct forms based on its molecular structure, or the way that the atoms of the element arrange themselves in space. The three forms of carbon that you may find in nature include graphite (the lead in your pencil), fullerenes, and diamonds. All three forms are pure carbon. Due to their different molecular structures, they have very different and distinct properties. Diamond is colorless, transparent, and very hard. Graphite is gray, opaque, and extremely soft. These properties of graphite are what allow you to write with it. Fullerenes are found in soot. The carbon atoms bond to form tiny, hollow balls or tubes. These different structures of pure carbon are called allotropes. In fact, allotropes occur when any pure substance has multiple structures.

The different molecular structures that may occur in carbon have to do with electron pair geometries. Graphite has a flat, linear structure containing layers of carbon bonds due to its electron pair geometry. The carbons in fullerenes are arranged like a single sheet of graphite that has rolled up upon itself to form a closed structure. Diamonds are clear and extremely hard. In fact, diamonds are the hardest substance on Earth, with a Mohs scale rating (hardness rating) of 10. The carbon in diamonds forms tetrahedral molecular structures based on its electron pair geometry. Therefore, along with its unique bonding capabilities, the different molecular structures that may occur with the pure substance carbon also make it an extremely versatile element.

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C2AB Inter- and Intra- Molecular Forces

Reading Science 1.

Which of the following is a unique quality of carbon? A.

It is found in 80 percent of the known compounds on Earth.

It can form both long chains and rings.

It has four available valence electrons in the second period.

All of the above.

Carbon is a major component in most of the organic molecules necessary for life and is an element that is found in every living cell. In which type of biological molecule is carbon the building block of both DNA and RNA? A.

Lipids

Amino acids

Nucleic acids

Carbohydrates

On Earth, you can find two very distinct forms of carbon: graphite and diamonds. There are many differences between these two forms of carbon. Choose the best answer to explain why these two forms of pure carbon differ in so many ways. A.

They have different colors.

They have different molecular structures.

Coal contains additional elements.

One form is organic, and the other is inorganic.

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Reading Science 4.

Carbon has a ground state electron configuration of 1s2 2s2 2p2. This particular configuration is what gives carbon many of its unique bonding capabilities. What does this electron configuration “do” for the element carbon? A.

It gives carbon four available valence electrons.

It gives carbon six protons in the nucleus.

It can give carbon tetrahedral geometry.

Both A and C.

Atoms and molecules may have the same chemical formula but different internal or molecular structures. What do you call a pure substance that can occur in multiple forms, such as carbon? A.

Allotropes

Equivalent compounds

Groups

Tetrahedrons

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C1B Chemical and Physical Properties

Reading Science Name:

Date:

Group:

Archimedes and the Golden Crown 1

According to legend, Archimedes, a famous mathematician and scientist who lived in ancient Greece, was given the task of determining if a goldsmith had embezzled gold from King Hiero. The King had supplied the goldsmith with pure gold so he could make a royal crown. However, the King suspected that the goldsmith had greedily stolen some of the gold to keep for himself and instead made the crown out of a less valuable metal. After all, it would have been easy to disguise a crown made out of another metal by topping it with a golden exterior. Archimedes had to figure out if the crown was really pure gold, or if the goldsmith had been dishonest and used a cheaper metal to make the crown.

Archimedes knew that gold was a very heavy metal. It would be easy to find out if the crown was genuine by calculating its density, or mass per unit of volume. To do this, he’d have to melt the crown, mold it into a cube with a known volume, and measure its mass. Archimedes knew that the King would not approve of this method, since the crown would be destroyed. He had to find another solution! How could he determine the density of the crown?

One day while taking a bath, the observant Archimedes noticed that the water level in the tub rose as he immersed his body in the water. When his body went under the surface of the water, it pushed the water out of the way and took up some of the space where the water had been. Archimedes realized that this same effect could be used with the golden crown. He could figure out the volume of the irregularly shaped crown by measuring the volume of the water that was pushed out of the way. Excited by his discovery, Archimedes jumped out of the tub and ran naked through the town shouting, “Eureka! Eureka!” or “I’ve found it!”

Archimedes conducted the test and was able to figure out how much space the crown took up—that is, its volume. Once he knew the crown’s volume, he compared the mass of the crown to the mass of a solid gold cube that had the same volume as the crown. Since the crown had less mass than the pure gold cube, Archimedes concluded that the sneaky goldsmith had indeed tried to trick the King by making the crown out of another metal. Archimedes’ discovery was very important in helping to determine the volume and density of oddly shaped objects. In fact, the principles of his method continue to be used by scientists today.

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Reading Science 1 In Paragraph 1 of this passage, what does the word embezzled mean? A tricked B hidden C lost D stolen

2 Students of science have studied this story for more than 2,000 years. Which of these statements describes the most important impact of his discovery? A He solved a problem for an ancient king. B He discovered a way to save time by working and taking a bath. C He gave scientists who came after him a way to measure the volume of oddly shaped objects. D It’s a funny story, and people throughout history have enjoyed reading it.

3 What is the main point of this passage? A How a scientist solved a problem using the concept of density B How to cleverly steal gold from a King C Why people shouldn’t steal things that belong to others D How to catch a thief

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C1B Chemical and Physical Properties

Reading Science 4 In paragraph 3 of this passage, what clues help you know the meaning of the word immersed? A While taking a bath B Under the surface of the water C Figure out the volume D Took up some of the space

5 The diagram shows what happened when an irregular solid was immersed in water. This irregular solid has the same mass as a gold block that is 2 cm wide by 2 cm high by 5 cm long. Using the method described in the passage, the density of the unknown compared with the density of gold – A is equal to gold. B is less than gold. C is greater than gold. D cannot be determined.

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C2AB Inter- And Intra- Molecular Forces

WRITING SCIENCE Name:

Date:

Group:

LOOK

THINK Matter can be described by its chemical and physical properties. A substance’s physical property is any characteristic that can be observed or measured without changing the sample’s composition, such as density, color, taste, and melting point. Physical properties can be described as intensive and extensive. Intensive properties include density, color, conductivity, malleability, and luster. Extensive properties include mass, volume, and length. A chemical property is any characteristic that changes when in contact with another substance, for example, the ability to react or not react with another substance. The following table shows some of the properties of two elements before the elements are mixed to create a chemical reaction. Iron

Sulfur

Malleable and solid

Powdered form

Magnetic

Non-magnetic

Gray color

Yellow color

WRITE Using the table above, describe the physical properties (intensive and extensive) of three different elements of your choice, and compare them to their chemical properties. Be sure to • clearly state your central idea; • organize your thoughts; • develop your essay in detail; • choose your words carefully; and • use correct spelling, capitalization, punctuation, and grammar. © Accelerate Learning Inc. - All Rights Reserved

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WRITING SCIENCE

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High School Chemistry

C2C

Molecular Structure in Designed Materials

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C2C Molecular Structure in Designed Materials

Student Handout Name:

Date:

Match each material to its type: wooden stick ceramic tile metal paper clip/wire plastic bag molecular solid

Describe how the wooden stick bent and then broke. Use the terms your teacher mentioned.

Describe how the paper clip/wire bent and then broke. Use the terms your teacher mentioned.

Describe how the tile bent and then broke. Use the terms your teacher mentioned.

Describe how the bag bent and then broke. Use the terms your teacher mentioned.

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C2C Molecular Strucre in Designed Materials

Student Journal Name:

Date:

Group:

Part I 1. What is the molecular structure of PVC?

How does the molecular structure of PVC compare to the molecular structure of a metal like copper wire?

What is the relationship between molecular structures of PVC and copper and each material’s ability to conduct electricity?

Part II 1. What are some uses for neoprene?

What are some uses for norsorex?

What do the uses for neoprene and norsorex tell you about their molecular structures?

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Student Journal Part III 1. What is the molecular structure of cooking oil?

What is the molecular structure of water?

How is viscosity affected by molecular structure?

Based on molecular structure, explain why cooking oil is more viscous than water.

Analysis 1. Give a basic summary of the relationship between molecular structure and physical properties.

If you were responsible for designing a cellphone screen, what molecular and physical properties would you deem most important? Why?

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STEMscopedia: MOLECULAR STRUCTURE IN DESIGNED MATERIALS C2C

Reflect What do Nylon pantyhose, Teflon nonstick pan coating, and plastic water bottles have in common? Believe it or not, they are all made of plastic. Plastics are composed of hydrocarbons, which are long chain molecules made up of carbon and hydrogen. Hydrocarbons can be manipulated to form different structures for different purposes. Even the earliest people figured out how to manipulate objects to their benefit, like using rocks to sharpen objects. With today’s technology, we can manipulate things on a molecular level. Chemical engineers are scientists that specialize in designing new materials by arranging molecules in various patterns. Many industries benefit from this, including the petroleum industry, pharmaceutical development, water treatment plants, biotechnology, and more. To design these materials, scientists need to understand how different molecular structures will interact will each other, since these different interactions will affect structural aspects such as boiling point, melting point, flexibility, durability, hardness, and conductivity.

Intermolecular Forces Molecular structure depends on intermolecular forces. Intermolecular forces are attractive and repulsive forces acting between molecules. Attractive forces pull molecules closer together, while repulsive forces repel and separate molecules. Remember, there are several types of intermolecular forces: London dispersion forces (also known as Van der Waals), dipole-dipole forces, ion-dipole forces, and hydrogen bonding. Polymers Polymers are long chain molecules made up of repeating structures called monomers. Cellulose, rubber, silk, and cotton are examples of natural polymers. Synthetic, or man-made polymers, include PVC, polystyrene, nylon, and polyethylene.

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Polymers have many uses and are usually classified by their physical and chemical properties and their applicable uses in the household and industry. Plastics The petroleum industry is one of the most widespread industries in the world. Many people depend on this industry for jobs, income, and its products. Some might only recognize items like gas, oil, and diesel as Most building blocks are petroleum products, when in fact so much of our everyday life comes made out of acrylonitrile in direct contact with petroleum products. Plastics are so prevalent and butadiene styrene, a type useful due to the pliability of the polymers they are made from. These of plastic that can be polymers are flexible and soft and can be molded into many shapes, then molded into any shape hardened to retain that shape. Plastics can be modified to change their and is very sturdy when density, color, flexibility, and many other characteristics. hardened. Depending on the components of the polymer chain, some plastics can be melted and remolded over and over, making certain plastics ideal for recycling. Other types, once hardened, cannot be broken down and reformed. Electric forces are caused by subatomic particles. Static electricity remained a puzzle until the basic atomic model was developed at the turn of the 20th century. To explain static charges clearly, we must refer to the pictured model. Every atom is made up of smaller particles. A neutral atom—or one without an electric charge—includes equal numbers of electrons and protons, and a similar number of neutrons. An atom’s neutrons carry no charge and are not involved in electrostatics. Protons have a positive (+) charge and are packed tightly into the nucleus. Electrons have an equal but opposite negative (−) charge and circle the nucleus from a large distance, relative to the size of a nucleus. Strong nuclear forces tightly bind protons, but electrons can be easily removed from or added to the atoms of many elements. “Free-range” electrons not associated with individual atoms may also be added to a material. Not all atoms are neutral. If an atom has more electrons than protons, it has a negative charge. If an atom has more protons than electrons, it has a positive charge. Similarly, an object with more negatively charged atoms has a negative charge, and an object with more positively charged atoms has a positive charge. Both positive and negative charges are created when a molecule gains or loses electrons (never protons). Electrical charges exert forces. Remember, equal charges repel each other, and opposite charges attract each other. Conductive Material Electricity is conducted via valence electrons of different materials. As electrons are passed along from one atom to the next, the current of electricity will flow. Good conductors of electricity have few valence electrons, meaning that particular atom is willing to give away those valence electrons.

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STEMscopedia: MOLECULAR STRUCTURE IN DESIGNED MATERIALS

An element’s location on the Periodic Table of Elements will indicate the number of electrons available to pass along for electrical currents. The majority of the known “good” conductors of electricity are metals. This coincides with where metals fall on the periodic table. Pharmaceuticals The human body is an amazing chemical factory. Everyday biomolecules such as enzymes, carbohydrates, and nucleic acids are working nonstop to keep this factory in optimal working condition. Sometimes, the body needs help. Sickness, traumatic injury, or even inherited genetic mutations can cause certain processes in the body to not function properly. Biomedical science and engineering has made great discoveries in the field of pharmaceutical and therapeutic drugs. A quick trip down the aisle of any drugstore will show you a wonderful selection of medications to treat anything from upset stomachs to itchy rashes. Sometimes these medications are not specific enough to the problem, and they not only treat the issue but also cause any number of side effects. Therefore, treating your headache will sometimes come with the cost of an upset digestive system. This is where molecularly designing receptor sitespecific drugs can come in. Now, instead of a drug flowing through the body randomly, it can travel directly to a targeted location.

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The active site of a neuron is usually specific to what it will interact with. If a drug is designed for a specific receptor, it will only act on that receptor. This targeted response can have many benefits. Only through understanding of the molecular design of the molecules, can we create such wonderfully engineered products. Application Match the following terms with their definitions: Polymer Outermost shell of electrons Monomer Large molecule made of repeating units Valence electrons Attractive forces between molecules Receptor site-specific Basic building block of large polymers Intermolecular forces Molecule shaped to fit a particular active site

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STEMscopedia: MOLECULAR STRUCTURE IN DESIGNED MATERIALS

Connecting With Your Child Help students observe polymers and intermolecular forces by making “Gak” silly putty together. Materials • 8-oz bottle of Elmer’s Glue • Mixing bowl • Food color of choice • Plastic cup • Teaspoon • Measuring cup • Borax laundry detergent • Stirring spoon • Water Instructions 1. Pour the 8-oz bottle of Elmer’s glue into the mixing bowl. 2.

Fill the empty glue bottle halfway with warm water to rinse out any excess glue and pour into the mixing bowl.

Stir this mixture while adding your desired food coloring.

Fill your plastic cup halfway with water and add 1 teaspoon of Borax.

Stir this mixture in the plastic cup until some of the Borax has dissolved.

Slowly pour and stir the Borax mixture into the glue-and-water mixture in the mixing bowl. The mixture should start to solidify to a silly putty consistency.

*Note: The more Borax and water added to the glue, the more it will solidify. *Note: Do not let a child put this in their mouth because Borax is a detergent and should not be digested. Discuss the following questions with your student at the end of the investigation. 1. What are some physical properties that you observed between the glue, Borax, and water? 2.

How do these physical properties change when you mix all the ingredients together?

What happens when you add more Borax?

Manipulate the end product into a ball and bounce it on the ground. How high does it bounce?

Add more Borax to the mixture and bounce the ball again. Did you notice any difference?

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C2C Molecular Structure in Designed Materials

Reading Science Name:

Date:

Group:

Plastics: Helpful or Harmful Polymers? 1

Molecular structure interactions are studied by chemical engineers because the molecular level properties are key to designing man-made materials. Although natural polymers exist, such as sugar and protein, this article will focus on synthetic polymers, especially those in the group called “plastics.”

Polymers are macromolecules (large long-chain molecules) with unique structural properties that are useful in manufacturing. Historically, during World War II, a shortage in natural polymers, like silk and rubber, led to the development of man-made polymers, like nylon, synthetic rubbers, and plastics. Plastic manufacturing produced items that were economical, lightweight, waterproof, durable, and either rigid or flexible.

Plastics are made from synthetic chemicals and extracted petroleum products that are made into chains of hydrogen and carbon atoms. These long molecules are composed of repeated basic molecules called monomers. For example, about 50,000 ethylene molecules of ethylene monomers (two carbon atoms bonded to four hydrogen atoms) can be joined into one long polymer called polyethylene or polythene. This process of combining monomers is called polymerization. Another way to make polymers is to remove some atoms from each monomer so they can join up in a process called polycondensation.

The two major types of plastics have very different properties: •

Thermoset plastics, like polyurethanes, polyesters, and epoxy resins, keep their shape once formed and cannot be reformed. Because their structure is so hard and lasting, thermoset plastics are often used in automobiles, aircraft parts, and tires.

•

Thermoplastics, like polyethylene, polypropylene, and polyvinyl chloride, are not as rigid as thermosets. Unlike thermoset plastics, thermoplastics can be reshaped if reheated. Due to this property, they can be easily put into molds or extruded as film, fibers such as Dacron, or thin packaging material like plastic baggies.

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Reading Science 5

Despite the popular benefits of plastics, several varieties can be harmful to both the environment and humans. Plastic containers are coded from 1 to 7, and these codes appear inside an arrow-enclosed triangle. Relatively safe plastics are coded 2, 4, and 5: •

Plastic #2 High-Density Polyethylene (HDPE) is used to hold milk, bottled water, cosmetics, cleaners, and cereal liners. However, the EPA has indicated that HDPE is a potential risk for infants and children. Group #2 plastics are the most commonly used plastics.

•

Plastic #4 Low-Density Polyethylene (LDPE) is used for toys, food containers, and as thin bags for dry cleaning, newspapers, food wrappers, and garbage bags. LDPE is easy to manufacture as thin plastic packaging materials because it has simple repeating units of one carbon and two hydrogen molecules.

•

Plastic #5 Polypropylene (PP) is used for food tubs and medicine containers. PP is a stiff, versatile plastic.

Unsafe plastics are coded 1, 3, 6, and 7: •

Plastic #1 Polyethylene Terephthalate (PET) is used to hold beverages and condiments, and can release toxic antimony and phthalates.

•

Plastic #3 Polyvinyl Chloride (PVC) is used in packaging (blister packs), meat wraps, lids, artificial leather, floating toys, etc. PVC leaches toxic chemicals, especially bisphenol A (BPA), which is linked to cancer. Group #3 plastics are the most hazardous plastics.

•

Plastic #6 Polystyrene (PS) is used as “Styrofoam” that is found in food cups and containers. Styrofoam can leach styrene, which has been connected to cancer.

•

Plastic #7 Polycarbonates (PC) and other plastics: This category has the highest risk due to the leaching of bisphenol A (BPA), which interferes with the body’s hormones and is also linked to cancer. PC is used in a wide variety of hard plastics applications, such as auto and tool parts, computers, and custom packaging. BPA is now banned in infant beverage containers or drink holders.

Lastly, a major benefit of plastics, its durability, has led to the greatest hazard on the environment: plastic is not biodegradable. Enormous amounts of plastic ends up in landfills, and the rest simply gets thrown by roadsides or left as property waste. Plastic pollution has become a worldwide issue. Environmentalists are hoping public education, stricter laws, and grassroots recycling efforts will help to alleviate pollution from plastics.

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Reading Science 1.

Which of the following lists include three man-made polymers? A.

Silk, natural rubber, plastic

Man-made rubber, sugar, silk

Plastic, nylon, man-made rubber

Sugar, silk, nylon

Plastics are made from which of the following? A.

Glass

Petroleum

Natural plastics

Wool and other natural fibers

What is true of thermoset plastics? A.

They are often used in automobile parts.

They do not hold their shape well.

They may be reformed over and over again.

They are often used to make plastic baggies.

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Reading Science 4.

What do plastics coded as #1, #3, #6, and #7 have in common? A.

They are all biodegradable.

They are all choking hazards.

They are classified as the safest of the coded plastics.

They can leach toxic chemicals.

What are some detrimental consequences of widespread plastic usage? A.

Less recycling

Links to cancer and pollution

Public awareness and cooperation

Higher crime rates

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C2C Molecular Structure in Designed Materials

Math Connections Name:

Date:

Group:

As electrons travel through wires, they meet resistance. A formula has been created to measure how strongly certain materials resist the flow of electrons, ρ. The formula to calculate for ρ is based on R (the electrical resistance), A (the area of the cross sectional) and l (the length of the material). 1.

Start your formula. The relationship between ρ and the electrical resistance (R) is directly proportional. (Write what the equation would look like so far.)

Add to your formula. The relationship between ρ and the length of the material (I) is inversely proportional.

Finish your formula. The relationship between ρ and the area of the cross-sectional (A) is directly proportional.

Rewrite the equation to solve for R. This is known as Pouillet’s law.

Rewrite the equation to solve for A.

Rewrite the equation to solve for l.

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Math Connections Use the equations below, and your final equation from number 3, to answer the following questions. σ=1/ρ

ρ=E/J

In the equations, σ represents the measurement of conductivity, ρ represents the electrical resistivity, E represents the magnitude of the electric field, and J represents the magnitude of the current’s density. 7.

Calculate for ρ, if E is 127 volts per meter, and J is 13 amps per square meter.

Calculate for σ, if E is 734 volts per meter, and J is 154 amps per square meter.

Write an equation to calculate for σ, given the magnitude of the electric field and the magnitude of the current’s density.

10. Calculate for J, if E is 14 volts per meter, R is 28 ohms, A is 4 square meters, and l is 9 meters.

11. Write an equation to calculate for σ, given R, A, and l.

12. Explain in detail why the following is true or not true.

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A I

E J

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C2C Molecular Structure in Designed Materials

WRITING SCIENCE Name:

Date:

Group:

LOOK

THINK Copper wire is found in homes all across the United States. It is used to transmit, generate, and distribute electricity. Silver is considered a better conductor than copper; however, its cost makes it an unrealistic material to choose when running electricity in homes. If you have ever accidentally heated something in the microwave with aluminum foil, you have seen its abilities as a conductor based on the sparks seen. Think about the molecular structure of metals when used in electrical designs. WRITE Describe the importance of the molecular structure of metals when utilized as conductors Be sure to clearly state your central idea; organize your thoughts; develop your essay in detail; choose your words carefully; and use correct spelling, capitalization, punctuation, and grammar.

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WRITING SCIENCE

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High School Chemistry

C2D

Bonding Models

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C2D Bonding Models

Student Handout Name:

Date:

Rules for Drawing Electron Dot Structures The octet rule states that atoms are more stable when they contain eight valence electrons in their outer electron shells. Atoms with less than eight electrons will either transfer or share electrons to meet this need. Whether electrons are transferred or shared depends on the type of bond created. In ionic bonds, electrons are transferred, so atoms will either gain or lose electrons. In covalent bonds, electrons are shared between atoms. Hydrogen and helium are the exceptions to the octet rule, as they are most stable with two electrons in their electron shell. When drawing an electron dot formula, draw a line between atoms that share a pair of electrons. Start by drawing the electron dot structure, and then connect the connecting pair of electrons with a line. Each connecting line represents two electrons. One example has been completed below. Complete the table by marking whether it is an ionic or covalent bond. Compound

Electron Dot Structure

Electron Dot Structure Diagram

Ionic (I) or Covalent (C)

Hydrogen chloride

Carbon dioxide

Sodium oxide

Water

Lithium sulfide

Nitrogen trichloride

Calcium bromide

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C2D Bonding Models

Student Journal Name:

Date:

Group:

Background: Electron Dot Formula Basics 1.

What do the dots in an electron dot formula represent?

Describe the pattern of electron dot formulas as you move from left to right in a period of the periodic table.

How are the electron dot formulas of ions different than those of neutral atoms? What do you need to add for cations and for anions?

Draw the correct electron dot structures for the atoms or ions. Neutral Atom

Ion

Neutral Atom

Ion

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C2D Bonding Models

Student Journal Background: Electron Dot Formula Basics, Continued 5.

Bag 1: Draw the electron dot formulas for both the neutral atom and for the appropriate ion of the four randomly drawn elements. Neutral Atom

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Appropriate Ion

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C2D Bonding Models

Student Journal Part I: Electron Dot Formulas for Covalent Compounds 1. What are covalent compounds, and how are they formed?

2. What is meant by “a full octet,” and why do atoms prefer to have this?

3. When atoms combine covalently, creating bonds, how are the shared pairs of electrons represented in the electron dot formula? In other words, what is placed between the atoms of a molecule to represent a shared pair of electrons?

4. Draw the correct electron dot formula for each atom, and then for each molecule. Include the bonding pair and all lone pairs of electrons. Make sure to show the correct structural formula of each molecule. Atom

Atom

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Molecule

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Student Journal Part I: Electron Dot Formulas for Covalent Compounds, Continued 5.

Bag 2: Draw the electron dot formulas for the covalent compounds that can be formed. Draw the correct electron dot formula for each element individually in the first column. Then draw the correct structural formula of the molecule formed in the second column. Electron Dot Formula of Individual Atoms

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Correct Electron Dot Formula of the Covalent Compound Formed

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Student Journal Part II: Electron Dot Formulas for Ionic Compounds 1. Illustrate the process that occurs during the bonding of the ions below.

F Mg

F K +

K 2. Describe how the electron dot formulas of ionic compounds illustrate that electrons are either gained or lost during the bonding process.

3. Why does the electron dot formula for the anion have a charge, and what does this represent? Why does the cation have a charge and a full octet?

4. Why is it important to include a charge when creating electron dot formulas for ions?

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Student Journal Part II: Electron Dot Formulas for Ionic Compounds, Continued 5.

Bags 3A and 3B: Draw the electron dot formulas for the ionic compounds that can be formed. Electron Dot Formula of Individual Atoms

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Correct Electron Dot Formula of the Ionic Compound Formed

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Student Journal Part III: Drawing Electron Dot Formulas for Any Compound 1. Create an electron dot formula for formaldehyde. The chemical formula for this compound is H2CO. Describe (in words) the steps you need to take to draw the correct structure. Support your description with illustrations.

2. What possible errors did you encounter while drawing this structure? Did it have to do with the bonds between the atoms? If so, what did you have to do to correct this?

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Student Journal Part III: Drawing Electron Dot Formulas for Any Compound, Continued 3.

Bag 4: Draw the correct electron dot formula for the compounds. Use additional paper, as needed. Chemical Formula of Compound

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Correct Electron Dot Formula of the Compound

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Student Journal Reflections and Conclusions 1. To understand electron dot formulas, you must first organize your data. Briefly explain how you organized the data.

2. How do the ball and stick models compare to your predictions? In other words, what does the ball and stick model show that an electron dot formula cannot?

3. Describe how the molecular models made with the balloons helped you better understand molecular structure.

4. Drawing correct electron dot formulas takes practice. Mistakes are common when learning how to do this. When drawing your diagrams, which of your mistakes led to a better understanding of how to draw correct electron dot formulas?

5. Sometimes a drawing of an electron dot formula has more valence electrons than the initial total number of valence electrons the person counted. What needs to be adjusted in the drawing to account for this, and how does that adjustment “fix” the problem?

6. On a separate sheet of paper, use all of the following terms to develop a graphic organizer based on what you’ve learned in this Explore. Terms: Ions, ionic compound, electron dot formula, valence electrons, electron pair, bonding pair, lone pair, single bond, double bond, triple bond, covalent compound, VSEPR, molecular structure, molecular geometry, central atom, electron pair geometry, bond angle © Accelerate Learning Inc. - All Rights Reserved

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Reflect Many scientists ran experiments in order to formulate the atomic model. Their experiments defined the atom as we know it today. Niels Bohr refined the structure of the atom by developing a model with a central nucleus surrounded by electrons moving in orbits and in different energy levels. Moving from the nucleus outward, each orbit represented higher energy than the previous one. Although the Bohr model is no longer the current model of the atom (it is now known as the Quantum model), the idea of specific energy levels for electrons became an important one that endured even after the model itself was found to be inadequate. Scientists today accept and continue to use the idea that each electron in an atom has a specific energy. They also accept the idea that electrons absorb or lose discrete quantities of energy called quanta. However, they reject Bohr’s idea that electrons follow specific orbital paths around the nucleus. Instead, scientists now think that the exact location of an electron in an atom cannot be known at any given time. Instead, we can know only a range of probable locations. These locations are best described as electron clouds because they are three-dimensional spaces with fuzzy borders. Scientists use the term orbital when they refer to the location of an electron in this model. You can use electron dot structures (also called Lewis structures) to show the valence electrons around an atom. For example, the element sodium has one valence electron, so it is written with one electron dot surrounding its elemental symbol (Na), as shown to the right. In contrast, the element chlorine has seven valence electrons, so it is written with seven electron dots surrounding its elemental symbol (Cl).

valence electrons: the negatively charged particles in the outer energy shell of an atom

Electron dot structures provide useful information about compounds— substances that consist of two or more types of atoms chemically bonded together. How do you think an electron dot structure is helpful when describing the chemical bonds that can form between elements? Molecular Bonding The strength of chemical bonds depends on how the valence electrons are behaving in compound. Electrons can be shared equally, unequally, given away or gained. This behavior will determine the type of bond. Lewis Dot Structures can be used to model molecular bonds because electron dot models show the number of valence electrons in each element. We can use these models to describe the electrons involved in a chemical bond between elements. Thus, electron dot formulas can show how bonds form between two atoms. We write electron dots slightly differently depending on the type of chemical bond.

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STEMscopedia: BONDING MODELS Ionic Bonding In ionic bonds, one or more valence electrons move from one atom (usually a metal) to another atom (usually a nonmetal). The atom that loses electrons becomes a positively charged ion, or cation; the atom that gains electrons becomes a negatively charged ion, or anion. The attraction between the newly formed cation and anion results in the formation of an ionic bond. When writing the electron dot structure of an ionic bond, we place the anion and the cation in brackets beneath their respective charges. For example, sodium chloride (NaCl) is an ionic compound of sodium (Na) and chlorine (Cl). As you saw above, the electron dot structure of sodium contains one electron, and the electron dot structure of chlorine contains seven electrons. To form an ionic bond, one electron transfers from sodium to chlorine. This gives the anion, chloride, a stable nucleus surrounded by eight valence electrons. It also gives the cation, sodium, a stable nucleus surrounded by eight valence electrons. An ionic bond forms when an electron moves from an atom of sodium (Na) to an atom of chlorine (left). The resulting ionic compound, sodium chloride (NaCl), consists of a positively charged cation (Na+) and a negatively charged anion (Cl–). Note how the cation and anion are placed within brackets on the right side of the chemical equation. Covalent Bonding In covalent bonds, atoms share valence electrons. A covalent compound, or molecule, may contain two or more atoms. Covalent bonds will usually occur between two non-metals, including most organic molecules. In molecules with more than two atoms, one atom is the central atom. We place electron pairs around all atoms to fulfill the octet rule, which is the tendency of most atoms to have eight valence electrons. For example, carbon tetrafluoride (CF4) is a covalent compound containing one central carbon atom (C) surrounded by four fluorine atoms (F). A carbon atom has four valence This electron dot structure electrons, and a fluorine atom has seven valence electrons. Therefore, shows carbon tetrafluoride. in carbon tetrafluoride the carbon atom shares one valence fluorine The line between the atom. Due to this sharing of valence electrons, the carbon and fluorine central carbon atom (C) atoms all fulfill the octet rule. and each fluorine atom (F) represents a single *An important exception to the octet rule is hydrogen; a hydrogen atom chemical bond formed by has at most two valence electrons. two shared electrons. The remaining electrons are placed in pairs as dots around each fluorine atom.

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STEMscopedia: BONDING MODELS Molecules can also form between atoms that share more than one electron pair. For example, in a molecule of oxygen gas (O2), each oxygen atom (O) has six valence electrons and needs two electrons to complete its valence orbital. Thus, the two oxygen atoms share two electron pairs.

Look Out The sharing of two electron pairs between atoms is a double bond. Similarly, the nitrogen atoms (N) in a molecule of nitrogen gas (N2) share three electron pairs. This type of covalent bond is a triple bond. Double and triple bonds can also occur between different elements. The following diagrams show electron dot structures for double and triple bonds. Molecules containing atoms of carbon and oxygen almost always have complete valence shells of electrons. Sometimes, however, covalent compounds form with incomplete valence shells. For example, an atom of boron (B) typically contains three valence electrons. In the molecule boron trifluoride (BF3), the boron atom shares an electron pair with each of three fluorine atoms. Each fluorine atom ends up with eight valence electrons, but the boron atom ends up with only six. Some atoms can have more than eight valence electrons. For example, in a molecule of phosphorus pentachloride (PCl5), the central phosphorus atom (P) has five valence electrons. These diagrams show the electron dot structures Each chloride atom has seven valence electrons and forms a single for each of these unusual bond with the phosphorus atom, which ends up with ten valence covalent compounds electrons. Covalent boding is when valence electrons are shared between atoms, causing them to form a relatively weak bond. Not all electrons are shared equally in covalent bonds. In non-polar covalent bonding, the electrons split their time orbiting the different atoms. This true sharing of valence electrons leaves the newly formed molecule as non-polar, or with no net charge. Sometimes though, electrons spend more time orbiting a particular atom in the molecule. This will cause the atom to have a slightly negative charge, thus creating a polar molecule.

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Water is an example of a polar covalent bond. The electrons are unevenly distributed, giving the molecule an overall charge.

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STEMscopedia: BONDING MODELS Application Draw the Lewis Dot structure and determine the type of bond in the following compounds. H2S CH2O NaCl MgO Metallic Bonding In addition to ionic and covalent bonds, there is a third type of chemical bond: metallic bonds. The strongest type of chemical bonds, these are formed between closely grouped atoms of metals. Examples of substances that contain metallic bonds are gold bars, sheets of aluminum foil, cooking pans, and copper wires. In metallic bonds, atoms do not share electrons, nor does one atom give up its electrons to another atom. Instead, electrons travel from one nucleus to another within the metallic structure. Because they do not hold on to their valence electrons, the metal atoms are actually positively charged ions (cations). Together, the electrons make up an electron sea. Metallic bonds form because metal cations are attracted to electrons in the electron sea.

This diagram shows metallic bonds in a sample of copper (Cu). Freely moving valence electrons (red) form an electron sea around the positively charged copper ions (blue).

The structure of the electron sea in a metallic bond provides many of the properties observed in metals. The abilities of the electrons to move throughout the electron sea and of the cations to slide past each other make metallic bonds more ﬂexible than ionic or covalent bonds. This ﬂexibility makes metals malleable, meaning able to be stretched and bent into shapes, and ductile, or able to be shaped into long, thin wires.

Metals like copper are commonly used to make wires that provide power to electrical appliances such as televisions and computers. In addition to their ductility, most metals are good conductors of electricity and heat. When an electric current is applied to a metal, the electrons in the electron sea begin to ﬂow in the direction of the current. When heat is applied to a metal, the electrons in the electron sea vibrate and collide with each other. As a result, energy—whether electrical or thermal— transfers easily through the metal.

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STEMscopedia: BONDING MODELS What Do You Think? Take a look at these photographs. The picture on the left shows an egg in a metal frying pan. The picture on the right shows metal wires in a power cable. Why are metals good materials from which to make frying pans and electrical wires?

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STEMscopedia: BONDING MODELS Connecting With Your Child Molecular Modeling To help your child learn more about molecular bonding and take a look at molecular geometry, create three-dimensional structures of molecules using household items. Here is a typical materials list: • 15–20 toothpicks • 15–20 marshmallows (or another small, soft food) The marshmallows represent atoms and the toothpicks are the chemical bonds between them. Have your child create models of the following molecules, using toothpicks to join marshmallows. Each toothpick represents two valence electrons. (On each marshmallow, write the chemical symbol for the element represented by the marshmallow. Alternatively, you may use a different colored item to represent the different atoms of each element.) • •

Carbon monoxide (CO): A linear molecule with a carbon atom that is triple bonded to an oxygen atom. Use three toothpicks to connect a carbon atom to an oxygen atom. Borane (BH3): A planar molecule with a central boron atom bonded to three hydrogen atoms through single bonds.

Each hydrogen atom should lie in the same plane at 120˚ angles from each other. • • •

Methane (CH4): A tetrahedral molecule with a central carbon atom bonded to four hydrogen atoms through single bonds. Each hydrogen atom should be at an angle of approximately 109.5˚ from the others. Water (H2O): A bent molecule with a central oxygen atom bonded to two hydrogen atoms through single bonds. The two hydrogen atoms should be at an angle of about 104˚ from each other. Ammonia (NH3): A trigonal pyramidal molecule with a central nitrogen atom bonded to three hydrogen atoms through single bonds. Each hydrogen atom should be at an angle of approximately 107˚ from the others.

Here are some questions to discuss with your child: 1. What is the electron dot structure for each molecule? 2.

What information can you get from a molecular structure that cannot be obtained from the electron dot structure of a molecule?

Why are molecular models valuable tools for understanding molecules?

What is the difference between ionic and covalent bonds? Between metallic and covalent bonds?

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Reading Science Name:

Date:

Group:

Bonds and Electrons 1

A chemical compound is a pure chemical substance consisting of two or more different elements that can be chemically separated into their simpler forms by chemical reactions. An example of a chemical compound is table salt, NaCl. This compound is made up of the simpler forms of sodium (Na) and chloride (Cl).

Valence electrons participate in the formation of a chemical bond, which creates chemical compounds. Valence electrons are electrons associated with an atom. The presence of valence electrons can determine the element’s chemical properties and whether it bonds with other elements. These electrons can be found in electron shells, or the orbit followed by electrons around an atom’s nucleus. The closest shell to the nucleus is called the “1 shell,” or “K shell”; followed by the “2 shell,” or the “L shell”; then the “3 shell,” or the “M shell”; and so on.

Lewis dot diagrams are just one way to visually demonstrate the valence electron locations. These diagrams show the bonding between atoms of a molecule and the lone pairs of electrons that may exist in the molecule that can be used in covalent bonds.

Through several forces, elements and compounds can be added together through chemical reactions and are held together by bonds. There are three types of chemical bonds: ionic, covalent, and metallic. Ionic bonds involve attractions between oppositely charged ions. Ions are atoms or molecules where the total number of electrons is not equal to the total number of protons, giving the atom a net positive or net negative electrical charge. An example of this would be sodium and fluorine undergoing a chemical reaction to form sodium fluoride. In this example, sodium loses its outer electron to give it stability; this then, in turn, enters the fluorine atom. The oppositely charged ions here—typically many of them—are then attracted to each other to form a solid.

Covalent bonds involve the sharing of electrons between atoms. For many molecules, the sharing of these electrons allows them to be stable together. An example of this type of bond would be H2 where the two hydrogen molecules share their electron to create a stable environment.

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Reading Science Continued 6

Metallic bonds occur as a result of electromagnetism and describe the attractive force that occurs between electrons and positively charged metal ions. An example of this type of bond would be brass or steel. In the example of brass, chemists can take the individual parts—zinc and copper—through a chemical reaction, and through this reaction, the two molecules will bond and create brass.

Ions are not just involved in bonding, but they can include different forms that can be used in chemical reactions or in nature. One type of ion is a cation. A cation is an ion with more protons than electrons giving it a net positive charge. Another type of ion is an anion. An anion is an ion with more electrons than protons giving it a net negative charge.

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Reading Science 1 A chemical compound is – A a pure chemical substance consisting of two or more different elements that can be chemically separated into their simpler forms by chemical reactions. B the formation of a chemical bond, which creates chemical compounds. C individual elements on the periodic table. D a combination of a metal and gas that creates a new element.

2 Valence electrons are housed in what we call – A rings. B forces. C orbits. D bonds.

3 Lewis dot diagrams are made for only which type of bond? A Metallic B Covalent C Ionic D All of the above

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Reading Science 4 A chemist wants to combine carbon and iron to form steel. What type of bond will this chemist create between these two elements? A Metallic B Covalent C Ionic D All of the above

5 An ion that has too few electrons and is therefore positively charged is an example of what type of ion? A Cation B Anion C Ion D Bond

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Math Connections Name:

Date:

Group:

Bonding Configurations Nonpolar Covalent Bonding: Bonding between two nonmetals; bonds are shared between the elements. Ionic Bonding: Bonding between an anion and a cation; bonds are not shared equally between the elements; anions and cations are shown in brackets with their respective charges. Electron dot structures are models that illustrate how bonds are formed between the valence electrons in elements. They are used to show the bonds between elements in a chemical formula. Create an electron dot structure: Step 1: Write the electron configuration for each element. Example: CCl4 Carbon – 1s22s22p2 Chlorine – 1s22s22p63s23p5 Step 2: Look at the s and p orbitals in the highest shell; this tells you the number of valence electrons. Carbon – 2s22p2 Chlorine – 3s23p5 Step 3: Write the symbol for the central element in the center of the model. Place an electron on each side of the symbol. Once there are four dots, you can start pairing them up until you run out of valence electrons.

Draw the electron dot formulas for the following compounds and describe the bonds between the atoms. Ionic bonds usually form when one of the atoms is a metal. Covalent bonds usually form when neither element is a metal. Bond Covalent

Ionic

Definition Formed when there is an equal sharing of electrons between two atoms Formed when an electron is transferred from one atom to another, creating ions

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Example

Model

H2 NaBr

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Math Connections 1.

CsCl

Li3N

CH4

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WRITING SCIENCE Name:

Date:

Group:

LOOK

THINK about the types of chemical bonds and how they are formed. A compound is a pure substance made up of two or more elements chemically joined together in definite proportions. Chemical bonds are formed from an attraction that holds atoms or ions together. The strength of chemical bonds can vary widely. Three specific types of chemical bonds are ionic bonds, covalent bonds, and metallic bonds. These bonds can be described by the particles being used to make the bonds, particle interaction within the bonds, and what types of atoms participate in each. WRITE the similarities and differences between ionic, covalent, and metallic bonds.

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WRITING SCIENCE

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High School Chemistry

C2E

Naming Compounds

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Student Handout Name:

Date:

Household Item Chart Directions: Write down six of the chemical compounds, formulas, and common names found on the labels of the common household items.

Chemical Compound Name

Formula

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Common Name (if applicable)

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Student Journal Name:

Date:

Group:

Background: Properties of Ionic and Covalent Bonds 1.

What is chemical nomenclature? What is IUPAC?

What are valence electrons? What do they have to do with bonding?

Use the diagram on this page to answer the remaining questions below. 3. What group number contains oxygen on the periodic table? 4.

How many valence electrons does oxygen have before bonding?

What group number contains hydrogen on the periodic table?

How many valence electrons does each hydrogen have before bonding?

How many valence electrons does oxygen have after bonding?

How many valence electrons does each hydrogen have after bonding?

How many hydrogen and oxygen atoms does the water molecule in the diagram contain?

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Student Journal Part I: Naming and Building Ionic Bonds 1. My Question of Inquiry:

2. Explain what ionic bonds are and how they are created.

3. Use your completed ionic bond puzzle pieces to complete the table below. Remember, when writing chemical formulas with polyatomic ions (ions with more than one element), you need to place parentheses around that polyatomic ion to balance charges. Use the first example in the table as a reference. Cation Used

Anion Used

Chemical Formula

Name of Ionic Compound

Fe3+

NO3−

Fe(NO3)3

Iron (III) nitrate

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Student Journal Part I: Naming and Building Ionic Bonds, Continued 4. Describe how ionic charges are used in creating chemical formulas for ionic compounds. Be specific.

5. Describe how ionic compounds are named. What element goes first? What element goes second?

6. When naming ionic compounds, what is critical to include with regard to the charges? In other words, how do you distinguish elements that may have multiple charges when naming ionic compounds? Use iron as an example.

7. Complete the table below. Naming and Writing Ionic Compounds Compound Name

Compound Formula

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Student Journal Part II: Naming Acids and Bases (use Acid-Base Reference Chart) 1. Complete the chart below. The first one has been done for you. Naming Acids Chart Cation

Anion

Chemical Formula

Name of Acid

Acid Type

H+

Cl−

HCl

Hydrochloric Acid

Monoprotic

HNO2

H+

Monoprotic

H+

ClO3−

HClO3

H+

CrO42−

H2CrO4

Diprotic

H2S

Diprotic

H+ H+

Phosphoric Acid

PO43−

Triprotic

2. Complete the chart below. The first one has been done for you. Naming Bases Chart Cation

Anion

Chemical Formula

Name of Base

Na+

OH−

NaOH

Sodium hydroxide

OH−

KOH

Ba2+

OH−

Ba(OH)2

Mg2+

OH− OH−

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Magnesium hydroxide Al(OH)3

Aluminum hydroxide

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Student Journal Part III: Naming Rules for Covalent Bonds 1. Explain what covalent bonds are and how they are created.

2. Complete the chart below. Covalent Compounds Chart 1st Element

2nd Element

Formula

# of 1st Element

# of 2nd Element

Name

CO2

Carbon dioxide

Nitrogen tribromide

P N

P2O5 Br CI

BCl3

Diphosphorous pentoxide 2

Carbon tetrafluoride

(Water)

Tetraphosphorus trisulfide

SO3

SiO2

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Student Journal Part III: Naming Rules for Covalent Bonds, Continued 3. Where are the elements used to create covalent compounds found on the periodic table? What are these elements known as?

4. Describe the naming system used to create covalent compounds. How is the first element named? How is the second element named?

5. What are diatomic molecules? Name all seven of them and include their molecular formulas. Why are they diatomic?

6. Complete the table below. Compound Name

Compound Formula

Molecular bromine

Br2

Molecular fluorine Carbon dioxide SO2 Carbon tetrachloride N2O5 Tetraphosphorous decoxide

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Student Journal Reflection and Conclusions 1. Describe the ways in which ionic compound nomenclature differs from covalent compound nomenclature.

2. To understand chemical nomenclature, you must first organize your data. Briefly explain how you organized the data to understand and use chemical nomenclature. (Include ionic compounds, oxyacids, nonoxyacids, and covalent compounds.)

3. Explain how you would determine whether parentheses are necessary in the formula of an ionic compound containing a polyatomic ion. Use a specific example.

4. Using all of the following terms, develop a graphic organizer based on what you’ve learned in this Explore. Use another sheet of paper, if needed. Terms: Chemical nomenclature, IUPAC, monoatomic ion, polyatomic ion, ions, ionic compound, covalent compound, molecule, diatomic molecule, chemical formula, anion, cation, acid, base, oxyacid, non-oxyacid, salt, electrolyte

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Student Journal Additional Practice: Naming and Writing Compounds (and identifying type of bond) I.

Identify each of the following compounds as either an ionic compound or a covalent compound. Then name the compound. Use the Student Reference Sheets named Summary of Naming Compounds and the Periodic Table. Compound Formula

Ionic or Covalent Bond

IUPAC Name

LiBr N2O3 MgS P4O10 CCl4 ZnSO4 CO SnCl2 AgI FeO BaCl2 CS2 CuBr2

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Student Journal Additional Practice: Understanding Naming and Writing Compounds II. Identify each of the following compounds with the correct IUPAC name. Common Name

Formula

Water

H2O

Natural Gas

CH4

Deicing Salt

CaCl2

Dry Ice

CO2

Muriatic Acid

HCl

Rust

Fe2O3

Milk of Magnesia (a base)

Mg(OH)2

Battery Acid

H2SO4

Carbonated Drinks (an acid)

H2CO3

Table Salt

NaCl

Lye (a base)

NaOH

Baking Soda

NaHCO3

Soda Ash

Na2CO3

Quicklime

CaO

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IUPAC Name

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Student Journal Additional Practice: Understanding Naming and Writing Compounds, Continued III. Complete the acid/base table below with the correct formula or name. Formula

IUPAC

Oxyacid, Non-oxyacid, or Base

HNO3

Nitric acid

oxyacid

Nitrous acid H2SO4

Sulfuric acid Sulfurous acid

HClO3

Chloric acid Chlorous acid

HClO

Hypochlorous acid

HC2H3O2

Acetic acid

HCl

Hydrochloric acid Sodium hydroxide Aluminum hydroxide Ammonium hydroxide

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Student Journal Additional Practice: Understanding Naming and Writing Compounds, Continued IV. Complete the table below with the correct formula or name. Then identify the type of bond. Formula

IUPAC Name

Nitrogen monoxide

Ionic or Covalent Bond

N2O NO2 N2O3 N2O4 N2O5 Cl2

Molecular chlorine Molecular iodine Molecular hydrogen Beryllium sulfide Aluminum sulfide Sodium fluoride Barium hydroxide Ammonium chloride Cesium chloride Potassium permanganate

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STEMscopedia: NAMING COMPOUNDS C2E

Reflect Have you ever listened to people speak a language that you can’t speak? Did you know what they were talking about? Just as people from different cultures have trouble communicating if they don’t share a language, scientists need a common language to communicate with each other. In the early 1900s, scientists were rapidly discovering new chemical substances. Chemists needed a method for naming these substances, so they developed a standard system, or nomenclature. Nomenclature refers to a system for naming things. The nomenclature used by all scientists to name and talk about chemicals was developed by the International Union of Pure and Applied Chemistry (IUPAC). In this companion, we will refer to the IUPAC rules for naming chemicals. Why is it important that all scientists use the same nomenclature for naming and talking about chemical substances? Chemical Formulas of Compounds The name of each chemical element can be abbreviated as a unique chemical symbol. For example, the chemical symbol for the simplest element, hydrogen, is H. The chemical symbol for helium is He, and the chemical symbol for hafnium is Hf. During chemical reactions, elements come together to form different compounds. Scientists represent these compounds by writing chemical formulas. In a chemical formula, the chemical symbols indicate the types of elements that make up a compound, and subscripts represent the quantity of each type of element in the compound. For example, the chemical formula for water is H2O. This formula means that a water molecule contains both hydrogen (H) and oxygen (O). The subscript “2” means that each molecule of water contains two hydrogen atoms. There is no subscript beside the letter “O,” so each water molecule contains only one oxygen atom.

subscript: a small number placed slightly below and to the right of a symbol that it describes

A water molecule (H2O) contains two hydrogen atoms (H) and one oxygen atom (O). ion: an atom or a molecule with an electric charge

Chemical formulas can be written for both ionic and covalent compounds. Ionic Compounds: An ionic compound forms when atoms bond by gaining or losing electrons. This type of compound usually forms between a metal and a nonmetal. A metal atom donates one or more electrons to a nonmetal. The result is a positively charged ion, called a cation, and a negatively charged ion, called an anion. The cation and anion are joined by an ionic bond. © Accelerate Learning Inc. - All Rights Reserved

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STEMscopedia: NAMING COMPOUNDS The total number of electrons lost by the cation (or cations) must equal the total number of electrons gained by the anion (or anions) to form a neutral ionic bond. The smallest unit of an ionic compound is called a formula unit. A familiar ionic compound is sodium chloride, which is commonly called table salt. Each formula unit of sodium chloride is made from one sodium cation and one chlorine anion. A neutral sodium atom (Na) has one valence electron, and a neutral chlorine atom (Cl) has seven valence electrons. The sodium atom donates one electron to the chlorine atom so that each ion now has a complete outer valence shell of electrons. The chemical formula is written NaCl. Notice that neither symbol has a subscript. This means each formula unit of sodium chloride is composed of one atom of each element.

A sodium atom (Na) donates an electron to a chlorine atom (Cl). The result is an ionic bond between a positively charged sodium ion, or cation (Na+), and a negatively charged chlorine ion, or anion (Cl–).

An ionic bond can also form between atoms of magnesium (Mg) and chlorine (Cl). A magnesium atom has two valence electrons to donate, but chlorine needs only one electron to complete its outer valence shell. To create a neutral ionic compound, one magnesium atom must donate one electron each to two chlorine atoms. The chemical formula for this compound, magnesium chloride, is MgCl2. In other words, each formula unit of magnesium chloride has one magnesium ion and two chlorine ions. No matter how large the sample of magnesium chloride, the ratio of magnesium to chlorine ions will always be 1:2.

Balancing the charges in an ionic compound is not always simple. For example, an ionic compound can form between aluminum (Al) and oxygen (O). The aluminum cation has a 3+ charge, and the oxygen anion has a 2– charge. In order for the overall charge to be neutral, two aluminum cations and three oxygen anions must be present in each formula unit: • For aluminum: 2 × (3+) = 6+ • For oxygen: 3 × (2–) = 6– valence electron: an electron found in an atom’s The resulting compound, aluminum oxide, has the chemical formula outer electron shell Al2O3. Covalent Compounds: A covalent compound forms when atoms bond by sharing electrons. This type of compound usually forms between two or more nonmetal elements. Most nonmetals are found on the upper-right side of the periodic table. (The exception is hydrogen, which is also a nonmetal and is found in many covalent compounds.) Each atom in a covalent compound shares one or more electrons with a neighboring atom, forming a covalent bond. The smallest unit of a covalent compound is called a molecule.

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STEMscopedia: NAMING COMPOUNDS In covalent compounds, atoms share electrons to achieve a full valence shell. For example, a molecule of methane (shown at right) is composed of one atom of carbon (C) and four atoms of hydrogen (H). A carbon atom has four valence electrons, and a hydrogen atom has one valence electron. The carbon atom shares one of its electrons with each hydrogen atom, and each hydrogen atom shares its one electron with the carbon atom. This completes the valence orbitals for the carbon atom and each hydrogen atom. The chemical formula of methane is CH4.

Look Out In addition to covalent and ionic bonds, a third type of chemical bond exists called a metallic bond. This type of bond exists between metals. In metallic bonds, electrons are not held tightly by one atom’s nucleus. Instead, the electrons are free to move from one nucleus to another. Because the atom does not hold on to its valence electrons, the metal atoms are actually positively charged ions. Do not confuse metallic bonds with ionic bonds. Although ions form in each bond type, metallic bonds form only among metal atoms, while ionic bonds form between metal and nonmetal atoms. Naming Ionic Compounds The IUPAC rules for naming all chemical compounds have been established so that a particular compound can be clearly identiﬁed. When naming ionic compounds, the cation is named ﬁrst, followed by the anion. The following rules can be applied for naming the ions.

In this metallic bond among copper (Cu) atoms, valence electrons can move from one nucleus to another.

Common Monatomic Ions Ion Name s odium Na + 2+ Mg magnes ium 3+ Al aluminum – F ﬂuoride – Cl chloride 2– O oxide 2– S s ulﬁde

Monatomic Ions: Monatomic ions are formed from one atom. The table at right lists some common monatomic ions. In most cases, the name of a monatomic cation is the same as the name of the atom. An anion’s sufﬁx, however, is changed to “–ide.” The ionic compound NaCl is therefore named sodium chloride. The rule for naming ionic compounds that contain more than one of the same type of ion is similar. For example, CaCl2 is formed from one calcium cation and two chloride anions. It is named calcium chloride.

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STEMscopedia: NAMING COMPOUNDS Monatomic Ions of Metals That Form More Than One Type of Ion: Many transition metals and metals on the right of the periodic table can form cations with different charges. When naming these cations, it is important to differentiate between the ions with different charges. Roman numerals are placed in the chemical name to indicate the correct charge per ion of the cation. The table below shows common examples. Common Monatomic Ions for Metals with More Than One Ion Ion C r2+

Name chromium(II)

Ion P b2+

Name lead(II)

C r3+

chromium (III)

P b4+

lead(IV )

C u+

copper(I)

Hg22+

mercury(I)

C u2+

copper(II)

Hg2+

mercury(II)

F e 2+

iron(II)

S n2+

tin(II)

iron(III)

tin(IV )

Polyatomic Ions: An ion formed from multiple atoms is known as a polyatomic ion. In polyatomic ions, nonmetal atoms form covalent bonds that satisfy the octet rule. The group of atoms, however, has a net positive or negative charge because the number of electrons in the overall structure does not equal the number of protons. The following rules are helpful for identifying polyatomic ions:

octet rule: the tendency of elements to form compounds by gaining, losing, or sharing electrons so that each atom has eight electrons in its valence shell

1. If two ions are composed of the same nonmetal and different numbers of oxygen atoms, the ion with more oxygen atoms is named the root of the nonmetal with the suffix -ate. The other ion is named the root of the nonmetal with the suffix -ite. 2. If four ions are composed of the same nonmetal and different numbers of oxygen atoms, they all are named with the root of the nonmetal, but the ion with the greatest number of oxygen atoms has the prefix per- and the suffix -ate. • The ion with one fewer oxygen atoms has no prefix but has the suffix -ate. • The ion with two fewer oxygen atoms has no prefix but has the suffix -ite. • The ion with the least number of oxygen atoms has the prefix hypo- and the suffix -ite. Ion Ch3COONH4+ CNHSO32SO42CO32206

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Common Polyatomic Ions Name Ion acetate H 3O + ammonium NO2cyanied NO3hydroxide ClO4sulfite ClO3sulfate ClO2carbonate ClO-

Name hydronium nitrite nitrate perchlorate chlorate chlorite hypochlorite

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STEMscopedia: NAMING COMPOUNDS One or both of the ions in an ionic compound may be polyatomic. As usual, the cation is written first followed by the name of the anion. Because the ionic compound must be neutral, more than one polyatomic ion may be needed to balance the charges. In this case, parentheses are used around the entire ion, and a subscript indicates the number of ions. For example, the formula for ammonium sulfate is (NH4)2SO4. Each molecule includes two ammonium ions and one sulfate ion. Naming Covalent Compounds IUPAC rules also exist for naming covalent compounds. In formulas for covalent compounds, the chemical symbols usually follow the order of elements—from left to right or top to bottom—on the periodic table. When naming covalent compounds containing two or more elements, the element that appears first in the chemical formula is named first. The suffix -ide is added to the root of the last element’s name, just as it is for ionic compounds A prefix is also added to each elemental name to denote the number of atoms of each element in the molecule. The table at right lists these prefixes. However, if the molecule contains only one atom of an element, the prefix mono- is often omitted. For example, a covalent compound containing five atoms of oxygen and two atoms of nitrogen in each molecule is N2O5. There are two atoms of the first element (nitrogen) and five atoms of the second element (oxygen). The compound’s chemical name is dinitrogen pentoxide. It is important to remember that some compounds have common names that are used more often. For example, H2O—dihydrogen monoxide (or simply dihydrogen oxide)—is almost always called water.

Preﬁ x

Number of atoms

mono-

ditritetrapentahexaheptaoctanonadeca-

2 3 4 5 6 7 8 9 10

Naming Acids and Bases In addition to ionic and covalent compounds, IUPAC has established rules for naming acids and bases. An acid is a substance that donates positive hydrogen ions (H+), and a base is a substance that accepts hydrogen ions. The naming of acids and bases follows a set of rules similar to the nomenclature rules of other types of compounds. First, let’s investigate the nomenclature rules for two different types of acids. • Binary Acids: A binary acid contains a hydrogen atom and a nonmetal. HCl (hydrogen and chlorine) and HI (hydrogen and iodine) are examples of binary acids. When naming these acids, the preﬁx hydro- is added to the root of the nonmetal, followed by the sufﬁx -ic. The word “acid” is then added. So the name of the binary acid HCl is hydrochloric acid, and the name of HI is hydroiodic acid. •

Oxyacids: An oxyacid is an acid that contains at least one hydrogen atom and a polyatomic ion containing oxygen. Common examples of oxyacids are H2SO4 and HNO2. If name of the polyatomic ion contains an -ate suffix, the suffix is changed to -ic. If the name of the polyatomic ion contains an -ite suffix, the suffix is changed to -ous. The word “acid” is then added. For example, H2SO4 contains a sulfate ion (SO42–), so it is named sulfuric acid. HNO2 contains a nitrite ion (NO2–), so its name is nitrous acid. The prefix hydro- is not used in naming oxyacids.

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STEMscopedia: NAMING COMPOUNDS Finally, let’s describe the naming conventions of substances that are bases. Bases are named according to the nature of the base. •

Bases That Are Ionic Compounds: Commonly, bases contain hydroxide ions (OH–). These bases are metal hydroxides. They are named in the same way that ionic compounds are named. The cation is named ﬁrst, followed by the anion, with a sufﬁx -ide. Sodium hydroxide (NaOH) and magnesium hydroxide (Mg(OH)2) are examples of bases that are also ionic compounds.

•

Bases That Are Covalent Compounds: Not all bases include a hydroxide ion. These bases are named in the same way that covalent compounds are named. Preﬁxes that denote the number of atoms are used for each element, and the sufﬁx -ide is added to the second element in the formula. For example, you may be familiar with ammonia, a base with the chemical formula NH3. Using IUPAC nomenclature rules, NH3 is called trihydrogen mononitride, or simply trihydrogen nitride.

What Do You Think? Take a look at the following formulas. Which formula represents carbonic acid? Which represents calcium hydroxide? Which represents carbonous acid? You can check your answers before the What Do You Know? section of this companion.

Ca(OH2) H2CO H2CO3 Scientists in the Spotlight: International Union of Pure and Applied Chemistry In the late 1800s and early 1900s, scientists in both industry and academia were discovering new elements and compounds at a rapid pace. A standardized nomenclature was needed to name these compounds. In 1919, a group of chemists formed the International Union of Pure and Applied Chemistry (IUPAC) to create a method of standardization in chemistry. This organization created not only the chemical nomenclature system to describe elements and compounds, but it also served to standardize methods for weights and measurements. Today, the IUPAC is an organization made of chemists from all parts of the world. More than 85% of the world’s chemical industries and research facilities are afﬁliated with IUPAC. IUPAC publishes a series of books that describe the naming system for different areas of specialty in chemistry: inorganic, organic, and biochemistry. Because newly discovered compounds may fall into more than one area of naming system, IUPAC may propose the naming system or a modiﬁed naming system in which these compounds should be named. IUPAC encourages the exchange of chemical information and knowledge among scientists in different areas of specialization. Did you correctly identify the compounds? Ca(OH)2 is calcium hydroxide, H2CO2 is carbonous acid, and H2CO3 is carbonic acid. 208

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STEMscopedia: NAMING COMPOUNDS What do you know? Take a look at the chemical formula for each of the compounds in the table below. First, determine whether each formula represents an ionic compound or a covalent compound. Write your answer in the middle column. If the formula represents an acid or a base, include that information in the middle column also. Then using the IUPAC naming conventions, determine the name of each compound. Write the name of each compound in the third column. Compound C aB r2

Type of Compound

Name of Compound

IC l F eO C C l4 S O3 HNO 3 Al(OH) 3 F e 2O 3 P 2O 5 K 2O C u(OH) 2 NaNO 3 H 2C O 3

Connecting WIth Your Child Chemical Nomenclature Around the Home To help your child learn more about chemical nomenclature, find at least five household products made from ionic or covalent compounds that can be named from the rules described in this lesson. Examples of possible household chemicals include table salt (sodium chloride: NaCl), Epsom salt (magnesium sulfate: MgSO4), baking soda (sodium bicarbonate: NaHCO3), and bleach (sodium hypochlorite: NaClO). The chemical compounds can also be one or many of the ingredients in household products. Possible sources for chemical compounds include soap, shampoo, conditioner, cleaning supplies, and toothpaste. Encourage your child to research the identified compounds using books or the Internet (the website of the IUPAC is a good place to begin). Record the chemical formula and the chemical name of each compound. Make sure to follow the IUPAC naming system of the compounds.

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STEMscopedia: NAMING COMPOUNDS Here are some questions to discuss with your child: 1. How do you know which elements are present in the compound? 2. How do you know how many of each type of element are present in the compound? 3. How can you identify the name of the compound? 4. Is the compound ionic or covalent? Is it acidic or basic? 5. Do the chemical compounds you have identiﬁed have more common names? (For example, the common name of sodium bicarbonate is baking soda.)

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Reading Science Name:

Date:

Group:

IUPAC 1

Chemical formulas describe how atoms combine to form every compound that composes matter. Just like any language, the way that elements combine follows very specific rules. As you may know, each element has a unique chemical symbol and a characteristic number of valence electrons. Elements combine in very predictable patterns, but with a huge number of possible variations. How does everyone agree about how to name the compounds that can be created with the 118 known elements? Very specific rules have been created by a standardizing organization called IUPAC, the International Union of Pure and Applied Chemistry. But, what is this organization, and how did it decide upon these rules?

The ancients knew about elements such as gold and silver more than 1,000 years ago. However, it was not until the mid-1700s that scientists began to isolate and identify other elements. By the mid-1800s, there were 62 known elements that needed to be named and categorized. These elements were identified by the differences in their properties and masses. They were also characterized by the ways that they combined with other known elements. These elements were given names by the scientists who discovered them. However, there was not a standardized system to categorize and make use of these elements. How could scientists use what they had learned about their discoveries in a meaningful way?

In 1860, German scientist Friedrich August Kekulé von Stradonitz led a committee to discuss these very issues. During this meeting, the committee members decided upon the first international system to name the known elements and compounds. This was especially important because of the ever-growing number of discoveries. The scientific community needed to know that, when discussing a chemical, they were all referring to the same element or compound. Scientists also needed to be given the proper credit for their studies.

However, it was soon found that the system set by the 1860 committee needed to be modified. New elements were being discovered. The number of known compounds was also rapidly growing. As more and more scientists performed experiments with these chemicals, there became a growing confusion about how to name and categorize them. In 1892, a new committee called the International Union of Chemistry (IUC), met to create and reform even more rules. These rules became known as the Geneva rules. The rules set forth by the IUC became the fundamental standardized rules for all organic compounds.

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Reading Science Continued 5

In 1911, there was another meeting. This was called the International Association of Chemical Societies (IACS). This committee met in Paris, France, to discuss exactly what types of processes the society should address. The final decision included authority over several processes. IACS would oversee the naming (nomenclature) of both inorganic and organic chemicals. It would determine procedures for standardizing the atomic weights of newly and previously discovered elements. IACS would create and maintain tables identifying the physical and chemical properties of elements and compounds. They would also create commissions to review scientific work and standardize both scientific publications and discoveries.

Finally, in 1919, the International Union of Pure and Applied Chemistry was created to follow both the IUC and the IACS. IUPAC has become a very influential organization. IUPAC oversees all chemical nomenclature in the world. Currently, the headquarters of IUPAC is located in Zürich, Switzerland. They have administrative offices in North Carolina known as the “IUPAC Secretariat.” IUPAC consists of many committees, all with varying responsibilities. Each committee is responsible for a different project and for the regulation of their nomenclature field. All of IUPAC’s decisions are built upon the work and regulations of the previous unions and societies. This is the current mission statement of IUPAC:

“The International Union of Pure and Applied Chemistry (IUPAC) serves to advance the worldwide aspects of the chemical sciences and to contribute to the application of chemistry in the service of Humankind. As a scientific, international, non-governmental, and objective body, IUPAC can address many global issues involving the chemical sciences.”

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Reading Science 1 The field of chemistry had incredible advancements in a very short time period, from the mid-1800s to the early 1900s. It was in this time period that the scientific community saw the need for the regulation of the growing number of chemical discoveries. What were the early committees trying to regulate? A Who actually discovered each chemical or element B How the newly discovered elements were named C How the chemical discoveries were published D All of the above

2 Paragraph 4 discusses one of the committees that was formed to help with the regulation of chemical discoveries. This committee, known as the International Union of Chemistry (IUC), developed a specific set of rules for chemists to follow. What were these rules called? A The nomenclature rules B The IUC rules C The Geneva rules D The Paris rules

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Reading Science 3 On this page, you will find the mission statement of one of the chemical regulation committees that met within the last 150+ years. “[This committee] serves to advance the worldwide aspects of the chemical sciences and to contribute to the application of chemistry in the service of Humankind. As a scientific, international, non-governmental, and objective body, [this committee] can address many global issues involving the chemical sciences.” Which committee stated this? A IUPAC B IUC C IACS D None of the above

4 The International Association of Chemical Societies was created to make sure that certain chemical processes and standards were followed. Which of the following was not one of those standards? A The standardization of the atomic weights of newly and previously discovered elements B The regulation of the IUPAC Secretariat office in North Carolina C The naming (nomenclature) of both inorganic and organic chemicals D The creation and maintenance of tables identifying the physical and chemical properties of elements and compounds

5 The number of known elements changed very rapidly over a 150-year period. There were 62 known elements in the mid-1800s. How many known elements (total) are there today? A 98 B 111 C 118 D 125

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Math Connections Name:

Date:

Group:

Part I: Naming Ionic Compounds The following table lists most of the ionic formulas you will need. The charge on other positive ions is indicated by a roman numeral. Use your periodic table for elements not shown on the chart. Name

Formula

Aluminum ion

Al3+

Ammonium ion

Name

Formula

Name

Formula

Bromide ion

Br−

Carbonate ion

CO3−2

NH4+

Chloride ion

Cl−

Hypochlorite ion

ClO−

Calcium ion

Ca2+

Fluoride ion

F−

Chromate ion

CrO42−

Lithium ion

Li+

Iodide ion

I−

Hydroxide ion

OH−

Potassium ion

K+

Nitride ion

N3−

Nitrate ion

NO3−

Silver ion

Ag+

Oxide ion

O2−

Phosphate ion

PO43−

Iron ion

Fe+

Sulfide ion

S2−

Sulfate ion

SO42−

Barium ion

Ba2+

Phosphorus ion

P3−

Bisulfate ion

HSO4−

Magnesium ion

Mg2+

Perchlorate ion

ClO4−

Chromium ion

Cr3+

Steps to Naming Chemical Compounds 1) List the symbols for each ion. 2) Write the symbols for the ions, side by side, with the cation first. 3) Find the LCM of the ions charges. 4) Write the chemical formula, using subscripts, to indicate how many of each ion are needed to make a neutral compound.

1. Give systematic names for the following compounds: a. BaCl2 b. CrCl3 c. PbS d. Fe2O3 2. Write the formulas for the following compounds: a. Magnesium fluoride b. Tin (IV) oxide c. Iron (III) sulfide d. Manganese (IV) oxide © Accelerate Learning Inc. - All Rights Reserved

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Math Connections Part II: Naming Molecular Binary Compounds Molecular compounds combine through covalent bonds and are named differently than ionic compounds. The number of atoms of each element is described using a prefix like mono-, di-, tri-, tetra-, penta-, or hexa-. However, if the first element has only one atom, mono- is not used; the element is simply named. (Example: NO is nitrogen monoxide.) 3. Give systematic names for the following compounds: a. PCl3 b. N2O3 c. P4O7 d. BrF3 e. NCl3 4. Write the formulas for the following molecular compounds: a. Disulfur dichloride b. Iodine monochloride c. Nitrogen triiodide

Part III: Naming Compounds with Polyatomic Ions Identify the polyatomic ion from the chart on the first page. To name the compound, list the cation first and then name the polyatomic ion. 5. Give systematic names for the following compounds: a. LiNO3 b. KHSO4 c. CuCO3 d. Fe(ClO4)3

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Math Connections 6. Write formulas for the following compounds: a. Potassium hypochlorite b. Silver (I) chromate c. Iron (III) carbonate

Part IV: Writing Ionic Formulas 7. Thiosulfate (S2O32−), an ion of sulfur, occurs naturally in hot springs and geysers. Ammonium thiosulfate is used to make fungicide; barium thiosulfate is used in the manufacturing of explosives; and potassium thiosulfate is used in the photo development process. Write the chemical formulas for these three compounds.

8. Iron pyrite, or fool’s gold, is the most common sulfide mineral. It is used in the manufacturing of iron (II) sulfate, commonly used to treat iron deficiency in humans. Write the chemical formula for iron sulfate.

9. Mercury, used in thermometers, is obtained by heating cinnabar ore, or mercury (II) sulfide. Write the ionic formula for mercury (II) sulfide.

10. Arsenic (III) chloride is a poisonous, colorless oil. It’s used to make pharmaceutical compounds, pesticides, and ceramics. Write the ionic formula for arsenic trichloride.

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Writing Science Name:

Date:

Group:

LOOK

THINK Naming and writing chemical formulas vary according to the type of compounds involved. To ensure consistency among scientists, the International Union of Pure and Applied Chemistry (IUPAC) devised several rules of nomenclature. While all nations practice science, they don’t always speak the same language. In order to produce a better way to communicate globally in terms of science, it was inevitable that a common language be established. The IUPAC devised several detailed rules of nomenclature that helped to bridge this language divide. The purpose of this nomenclature is also to name each chemical in a way that leaves no ambiguity and clearly describes the chemical structure of the compound while arriving at a single name. Between language differences and common names, it is a necessity to have a common set of rules for nomenclature to correctly identify compounds. For example: Common Name

Chemical Formula

IUPAC Name

Battery acid

H2SO4

Sulfuric acid

Lye

NaOH

Sodium hydroxide

Baking soda

NaHCO3

Sodium hydrogen carbonate

The IUPAC nomenclature varies depending on the type of compound being named. WRITE Compare the IUPAC nomenclature rules for naming acids and bases. Provide common examples of each to support your statements. Be sure to • clearly state your central idea; • organize your thoughts; • develop your essay in detail; • choose your words carefully; and • use correct spelling, capitalization, punctuation, and grammar.

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Writing Science

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High School Chemistry

C3AB

Chemical Reactions

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Student Handout Name:

Date:

Chart for Demonstration 1: Pachyderm Tusk Cleaner Properties of Reactants Hydrogen peroxide Yeast Water + food coloring Dish soap

Properties of Product Elephant toothpaste

Chart for Demonstration 2: Production of a Precipitate Properties of Reactants Magnesium sulfate (Epsom salts) solid Magnesium sulfate (Epsom salts) dissolved in solution Sodium carbonate (washing soda) solid Sodium carbonate (washing soda) dissolved in solution Properties of Product Production of a precipitate

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Student Journal Name:

Date:

Group:

Part I: Chemical Equation Basics O2 + 2H2 → 2H2O 1.

What are the reactants in this chemical equation?

What are the products in this chemical equation?

What is a balanced chemical equation, and what is used to balance it?

In the space below, illustrate what is occurring at the molecular level in this chemical equation. How many oxygen and hydrogen molecules are needed to create the required number of water molecules? (Draw the atoms involved in the equation).

What are the coefficients in the equation above, and how are they used to balance the equation? How does this relate to the number of atoms on each side of the equation?

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Student Journal Part I: Chemical Equation Basics 6.

Define the law of conservation of mass and use the equation above to justify your answer.

All chemical equations must follow the law of conservation of mass. Using a pencil, apply what you have learned to balance the following equations. (Hint: Compounds such as NO3 can often be balanced as a unit on either side of the equation.) A.

KOH + HNO3 → KNO3 + H2O

H2CO3 → H2O + CO2

Na + O2 → Na2O

Pb(NO3)2 + KCl → PbCl2 + KNO3

C2H6 + O2 → CO2 + H2O

Cl2 + KBr → Br2 + KCl

Part II: Types of Chemical Reactions 8.

State the correct type of chemical reaction for each of the equations that you balanced in Part I of the Student Journal. A.

_____________________________________________________________________

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Student Journal Part III: Creating Chemical Equations 1.

Complete Line 1 of the table below, and then proceed to Line 1 of the table on the next page. Repeat with Line 2 and so forth. You will need to reuse your cut squares for each equation.

Equation Description

Copper reacts with silver nitrate to produce copper(II) nitrate and silver.

Sulfuric acid reacts with aluminum hydroxide to produce water and aluminum sulfate.

Sulfur trioxide gas reacts with water to produce sulfuric acid.

Ammonium chloride reacts with silver nitrate to produce ammonium nitrate and silver chloride.

Potassium chlorate breaks down into potassium chloride and oxygen gas.

Methane (CH4) (g) reacts with oxygen gas to produce carbon dioxide and water vapor.

Reactants

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Products

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Student Journal Part III: Creating Chemical Equations, continued 2.

Write the “balanced chemical equation” for each equation on the previous page. Include the “state of matter” symbols and identify the type of reaction

Balanced Chemical Equation

Type of Reaction

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Student Journal Reflections and Conclusions 1.

A chemical equation represents the compounds involved in a chemical reaction, but it does more than this. What else does a chemical equation represent? (Hint: It has to do with the law of conservation of mass.)

Explain why the law of conservation of mass is critical when writing and balancing chemical equations. What is used to follow this law when balancing equations, and what must be equal in the equation (and why)?

This type of chemical reaction mixes two aqueous compounds. What type of chemical reaction is this?

During this type of chemical reaction, a metal is placed in a colored solution, resulting in the solution turning clear as another metal is formed. What type of reaction is this?

During this type of reaction, a carbon-based compound reacts with oxygen gas, producing carbon dioxide and steam. What type of reaction is this?

Using all of the following terms, develop a graphic organizer based on what you’ve learned in this Explore. Use another sheet of paper, if needed. The terms include the following: chemical formula, reactant, product, coefficient, reaction, chemical equation, law of conservation of mass, combination reaction, decomposition reaction, combustion reaction, double replacement reaction, and single replacement reaction.

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C3AB Chemical Reactions

Student Journal Name:

Date:

Group:

Part I: Chemical Reactions and Equations 1.

CaCl2 + Na2CO3 → CaCO3 + 2NaCl a.

Is this equation balanced?

__________________________________________________________________________ b.

Circle the reactants and place a box around the products.

Write the equation as a sentence.

__________________________________________________________________________ 2.

Na2CO3 + 2H2O → 2NaOH + H2CO3 a.

Is this equation balanced?

__________________________________________________________________________ b.

Circle the reactants and place a box around the products.

Write the equation as a sentence.

__________________________________________________________________________ 3.

Na2CO3 + 2CH3COOH → 2NaCH3COO + H2CO3 a.

Is this equation balanced?

__________________________________________________________________________ b.

Circle the reactants and place a box around the products.

Write the equation as a sentence.

__________________________________________________________________________ 4.

What are the four signs of a chemical reaction? 1.

_____________________________________________________________________

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Student Journal Part II: Signs of a Chemical Reaction Using Known Reactants 1.

Fill in your own information first, and then collect the rest of the data from your lab group (and share your own). Physical Properties (Observed With Hand Lens)

Observations During a Chemical Reaction When Mixed With Water

Observations During a Chemical Reaction When Mixed With Vinegar

Calcium Chloride

Magnesium Sulfate

Sodium Carbonate

Citric Acid

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Student Journal Part III: Identifying Unknown Reactants Fill in your own information first, and then collect the rest of the data from your lab group (and share your own).

Physical Properties (Observed With Hand Lens)

Observations During a Chemical Reaction When Mixed With Water

Observations During a Chemical Reaction When Mixed With Vinegar

Which Unknown Reactants?

Mixture A

Mixture B

Mixture C

Mixture D

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Student Journal Reflections and Conclusions 1.

How were the chemical reactions of the white solids similar? How were they different?

Did you produce a new substance with new properties when you combined Mixture A with water or vinegar? Explain.

Did you produce a new substance with new properties when you combined Mixture B with water or vinegar? Explain.

Did you produce a new substance with new properties when you combined Mixture C with water or vinegar? Explain

Did you produce a new substance with new properties when you combined Mixture D with water or vinegar? Explain.

What was the identity of your unknown mixture? What evidence led you to believe this?

Which of the unknown solids together could be used to make a cold pack?

Which of the unknown solids could be used to make a heat pack?

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STEMscopedia: CHEMICAL REACTIONS C3AB

Reflect Have you ever tried to balance a book on the top of your head while walking across a room? How do you compensate if you begin to feel the book sliding to one side? You might lean in the opposite direction or hold out your hands to steady the book. A chemical reaction converts one or more substances into another. Writing a chemical equation can explain a chemical reaction. Chemical equations require balance of a different kind. What does it mean to balance a chemical equation? How does balancing a chemical equation reﬂect what happens during a chemical reaction? What types of things do you observe during a chemical reaction? Perhaps you hear loud sounds, see bright lights, and feel the warmth of the reaction. Do any of these indicate that a chemical reaction has taken place? Writing Chemical Equations All chemical reactions involve reactants and products. • •

Reactants are the starting substances—the substances before the reaction happens. Reactants are normally written on the left side of a chemical equation. Products are the resulting substances—the substances after the reaction happens. Products are normally written on the right side of a chemical equation.

In a chemical equation, the reactants are separated from the products with an arrow showing the direction of the chemical reaction. Therefore, in a typical chemical equation, the arrow points from left (the reactants) to right (the products). The chemical equation below shows what happens when you open a can of soda. Soda contains liquid water (H2O) and carbon dioxide gas (CO2). (This is why soda is also called carbonated water.) The pressure inside a sealed can causes some of the water and carbon dioxide to combine into a substance called carbonic acid (H2CO3). Opening a soda can releases this pressure. As a result, the carbonic acid undergoes a chemical reaction and is converted back into liquid water and carbon dioxide gas. (You see the CO2 escaping into the air as ﬁzz.) H2CO3(aq) (reactants)

H2O(l) + CO2(g) (products)

The letters in parentheses next to the chemical formula for each molecule describe the physical state of each substance; they are called state symbols. For example, (aq) stands for aqueous and refers to a substance in solution; in addition, (s) refers to a solid, (l) refers to a liquid, and (g) refers to a gas.

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STEMscopedia: CHEMICAL REACTIONS Balancing Chemical Equations During a chemical reaction, the atoms in one group of molecules (the reactants) rearrange to form new molecules (the products). According to the law of conservation of mass, the same atoms must be present in both reactants and products. Therefore, in a chemical equation, the same atomic symbols must appear to the left of the arrow and to the right of the arrow. Take another look at the chemical reaction from the previous example, in which calcium carbonate decomposes into calcium oxide and carbon dioxide: CaCO3(s) CaO(s) + CO2 (g) On the reactants side of the equation, there is one atom of calcium (Ca), one atom of carbon (C), and three atoms of oxygen (O). Likewise, on the products side of the equation, there is one calcium atom, one carbon atom, and three oxygen atoms. The equation is balanced. This is a relatively straightforward reaction. One molecule of reactant equals one molecule of one product and one molecule of another product. Most reactions are more complicated, however, involving different amounts of molecules on both sides of the equation. To balance the equation for such a reaction, you will need to use coefﬁcients: numbers placed in front of a chemical formula to indicate the number of speciﬁc molecules present during a reaction. Let’s look at an example involving the combustion (burning) of propane gas. In this reaction, propane (C3H8) reacts with oxygen (O2) to produce water vapor (H2O) and carbon dioxide (CO2): C3H8(g) + O2(g)

H2O(g) + CO2(g)

How can we tell if this equation is balanced? Begin by counting the atoms of each element on both sides of the equation. You can do this by making a table.

The combustion of propane allows the steak to be cooked on the grill.

How many atoms in reactants?

How many atoms in products?

Are the atoms balanced?

Carbon (C)

Hydrogen (H)

Oxygen (O)

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STEMscopedia: CHEMICAL REACTIONS We will need to use coefﬁcients to balance this equation. Where should we begin? Balancing an equation often involves trial and error, but here are some guidelines: 1. First, try to balance elements that appear in only one molecule on each side of the equation. 2. If several elements appear in only one molecule, begin with the element that has the fewest number of atoms. Let’s take another look at the equation for the combustion of propane: C3H8(g) + O2(g)

H2O(g) +CO2(g)

Both carbon (C) and hydrogen (H) appear in one reactant molecule and one product molecule. Because the unbalanced equation contains fewer carbon atoms, let’s begin with this element. The reactants side of the equation contains three carbon atoms, and the products side contains one carbon atom. To balance the carbon atoms, add a coefﬁcient of 3 to the CO2 molecule on the products side of the equation: C3H8(g) + O2(g)

H2O(g) + 3CO2(g)

This coefﬁcient means the products side of the equation contains three molecules of CO2. As a result, both sides of the equation now contain three carbon atoms. In addition, the products side of the equation now contains seven oxygen atoms (H2O contains one oxygen atom, and 3CO2 contains six oxygen atoms). Because it appears in several products, however, let’s wait to balance oxygen until we have balanced hydrogen. The reactants side of the equation contains eight hydrogen atoms, and the products side contains two hydrogen atoms. To balance the hydrogen atoms, add a coefﬁcient of 4 to the H2O molecule on the products side of the equation: C3H8(g) + O2(g)

4H2O(g) + 3CO2(g)

This coefﬁcient means the products side of the equation now contains four molecules of H2O. As a result, both sides of the equation now contain eight hydrogen atoms. In addition, the products side of the equation now contains ten oxygen atoms (4H2O contains four oxygen atoms and 3CO2 contains six oxygen atoms). We may now balance oxygen. The reactants side of the equation contains only two oxygen atoms. Fortunately, they appear together in a molecule with no other atoms (O2). We can add a coefﬁcient of 5 to O2 without affecting the balance of other atoms in the equation. Here is the balanced equation for the combustion of propane: C3H8(g) + 5O2(g)

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4H2O(g) + 3CO2(g)

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STEMscopedia: CHEMICAL REACTIONS To conﬁrm, count the atoms of each element on both sides of the equation. The products side contains three carbon atoms, eight hydrogen atoms, and ten oxygen atoms. The reactants side also contains three carbon atoms, eight hydrogen atoms, and ten oxygen atoms. The equation is balanced. If, after working with each element, you have still not balanced the equation, you will need to change your coefﬁcients. As before, begin with the element that appears in the fewest number of atoms.

Look Out Continue to try new combinations of coefficients until you have balanced the equation. Do not confuse coefficients with subscripts. • When you change a subscript in a chemical formula, you are changing the type of molecule involved in the reaction. Never change subscripts to balance a reaction! • When you change a coefficient in a chemical formula, you are changing the number of molecules involved in the reaction. For example, consider the reaction of oxygen gas (O2) and hydrogen gas (H2) to produce water vapor (H2O). H2O(g) H2(g) + O2(g) The products side of the equation contains two hydrogen atoms and two oxygen atoms. The reactants side of the equation contains two hydrogen atoms and one oxygen atom. What happens if you add a subscript of 2 to the oxygen atom in the product? H2(g) + O2(g)

H2O2(g)

By changing the subscript, you have changed the molecule. Instead of water (H2O), the product is now hydrogen peroxide (H2O2). Instead, add a coefﬁcient of 2 in front of the water molecule: Changing a subscript may not seem like a big deal, but there is a big difference between a molecule of water (left) and a molecule of hydrogen peroxide (right).

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STEMscopedia: CHEMICAL REACTIONS What Do You Think? Equations for Everyday Reactions There are several different categories of chemical reactions. Reviewing the chemical equations for common reactions in each category is an excellent way to practice balancing equations. Here are ﬁve important reaction types: synthesis, decomposition, combustion, single-replacement, and double-replacement. Review the balanced equations below. •

Synthesis Reaction: Two or more reactants combine to form a larger product. This is also called a synthesis reaction. An example of a synthesis reaction is an alkali metal reacting with a halogen. Here is the balanced equation: A+B K(s) + Cl2(g)

•

Decomposition Reaction: A larger reactant breaks down into two or more products. A decomposition reaction happens inside an automobile’s airbag. A substance called sodium azide breaks down into sodium and nitrogen gas, which inﬂates the airbag: AB 2NaN3(s)

•

A+B 2Na(s) + 3N2 (g)

Combustion Reaction: All combustion reactions involve oxygen as a reactant and release heat. (Burning is a type of combustion—the process of burning always involves oxygen and releases heat.) Water is a common product in many combustion reactions. For example, a burning candle produces water vapor. Burning propane is another example of a combustion reaction. Here is the balanced equation for the formation of carbon dioxide and water by the combustion of propane, shown in the photograph at right: C3H8(g) + 3O2(g)

•

AB 2KCl(s)

H2O(l) + CO2(g)

Single-Replacement Reactions: In a typical single-replacement reaction, the reactants include a compound and a single atom. During the reaction, the single atom changes places with one of the ions in the compound. In other words, a single-replacement reaction looks like this: AX + B

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BX + A

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STEMscopedia: CHEMICAL REACTIONS The reaction of aluminum and copper(II) chloride to form copper and aluminum chloride is an example of a single-replacement reaction: 3CuCl2 + 2Al •

2AlCl3 + 3Cu

Double-Replacement Reactions: In a double-replacement reaction, two reactants appear to exchange ions as they form two products. In other words, a double-replacement reaction looks like this: AY + BX AX + BY

The reaction of lead nitrate and potassium iodide, shown in the photograph at right, is an example of a double-replacement reaction. Notice how the nitrate molecule (NO3) and the iodide ion (I–) appear to change places in this equation: Pb(NO3)2(aq) + 2KI(aq)

PbI2 (s) + 2KNO3(aq)

Look Out Let’s examine the different types of reactions above. What makes certain reactants react when mixed together? Reactivity is based on a chemical’s valence electrons and its properties. Elements have been organized on the periodic table based on valence electrons (electrons in the outermost shell). Elements in the same group (or column) have the same number of electrons in their valence shells, which means they tend to behave similarly during chemical reactions. Elements react based on the number of their valence electrons. Atoms want a full valence shell. In order for this to happen (for many elements), they must bond with other atoms. These valence electrons can either be transferred or shared to create a bond. In general, alkali metals are the most reactive metals (because they only need to lose one electron) and halogens are the most reactive nonmetals (because they only need to gain one electron). Also, as the element increases in period number (or rows), the reactivity decreases. Therefore, the periods closest to the top of the periodic table are more reactive than those toward the bottom of the periodic table.

What Do You Think? Now that we’ve discussed balancing chemical reactions, let’s look at how to identify different types of reactions.Take a look at the following photographs. The picture on the left shows rusty nails. The picture in the middle shows a lit matchstick. The picture on the right shows a rotting pumpkin. These images are all examples of chemical reactions taking place. In each image, how do you know a chemical reaction is occurring?

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STEMscopedia: CHEMICAL REACTIONS

We can observe evidence of chemical reactions. Scientists confirm that a chemical reaction occurs by determining if a new substance, with new properties, is formed. Scientists may perform additional chemical reactions and use instrumentation to confirm these new substances. Yet, there are signs or indicators that can suggest that a chemical reaction has occurred. The only way to know for sure that a chemical reaction has occurred is to identify the new substance. If you observe one or more of these signs, this provides evidence that a chemical reaction may have taken place. Remember, not all chemical reactions will produce one of these signs. Let’s look at the signs that provide evidence of a chemical reaction. •

Production of a gas: One very common reaction that involves the production of a gas is mixing sodium bicarbonate, also known as baking soda (NaHCO3), with an acid such as vinegar (CH3COOH). The products of this reaction are water (H2O), carbon dioxide (CO2), and a substance called sodium acetate (CH3COONaNaOOCCH3). During this reaction, you can see bubbles in the solution, as shown in the picture on the next page. These bubbles are caused by the carbon dioxide gas escaping into the air as the reaction takes place. The chemical equation for this reaction is: CH3COOH + NaHCO3 (vinegar) + (baking soda)

•

CH3COONaNaOOCCH3 + H2O + CO2 (sodium acetate) + (water) + (carbon dioxide)

Production of light: The burning of logs in a fireplace is the reaction of the wood and oxygen along with a heat initiation source. Wood is made of cellulose, a combination of different substances that contain carbon, hydrogen, and oxygen. This reaction produces a large amount of energy. This energy is in the form of both heat and light. This type of reaction is a combustion reaction. It is similar to the reaction that produces the bright light and heat in fireworks.

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STEMscopedia: CHEMICAL REACTIONS •

Change in temperature: Chemical reactions can either give off heat or use heat. Perhaps you have had an injury and applied a chemical heat pack to the area. A chemical heat pack is an example of a reaction that produces heat. A common substance in a heat pack is magnesium sulfate (MgSO4). When the heat pack is activated, the magnesium sulfate reacts with water. The result is the production of heat, which you use to soothe your injury. Chemical cold packs work in an opposite way to use heat when they mix with water. They may feel very cool to the touch. These temperature changes are evidence of a chemical reaction.

•

Formation of a precipitate: A precipitate is a solid substance that forms and separates from a solution. A precipitate often settles to the bottom of a liquid reaction. One common chemical reaction that forms a precipitate is the reaction of solutions of lead nitrate (Pb(NO3)2) and potassium iodide (KI). Each of these substances in a solution is clear and colorless. But if you mix a solution of each substance, lead iodide (PbI2) and potassium nitrate (2KNO3) form as products. Lead iodide is insoluble, so it separates from the solution as a yellow precipitate (shown in the image on the right). The potassium nitrate remains in the solution. The chemical equation for this reaction is: Pb(NO3)2+PbKl

Pbl2+2KNO3

Change in color: You may have seen rust form on a steel object, such as a chain or an automobile. In this chemical reaction, iron (Fe) in the steel reacts with oxygen (O2) in the air, as well as water (H2O), to produce rust (Fe(OH)3). The properties of steel are different from the properties of rust. Steel is a shiny silver metal made from iron and other elements. Rust is a flaky reddish substance. The change in color from silver to red provides evidence that a chemical reaction has happened.

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STEMscopedia: CHEMICAL REACTIONS What Do You Know? Certain evidence indicates that a chemical reaction may have occurred. This evidence includes the production of a gas, the production of light, a change in temperature, a color change, and the formation of a precipitate. Look at the types of reactions given on the next page. For each reaction, decide which of the five types of evidence would be observed. Write your answers in the table provided on the next page. More than one type of evidence may be possible for each example.

Type of Reaction

Indicator of a Reaction (production of a gas, light, change in temperature, color change, or formation of a precipitate)

Exploding fireworks Baking a cake Burning paper Mixing an antacid tablet and water Making chalk from two liquids Blue litmus paper turning red in acid A salt solution turning cold as the salt dissolves A copper penny tarnishing Oxygen gas and hydrogen gas producing water and heat Testing for carbon dioxide by bubbling a gas in lime-water to produce a milky white solution

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STEMscopedia: CHEMICAL REACTIONS Connecting With Your Child Investigating Chemical Reactions To help your child learn more about chemical reactions, work together to determine how to identify the evidence that may be observed when a chemical reaction occurs. To do so, gather the following materials: • 3 Glasses of water • 2 Effervescent tablets • 1 Tablespoon of sugar • 1 Tablespoon of Epsom salt • Thermometer While performing the chemical reactions, encourage your child to record all observations. Let the first glass contain the control sample in which no chemical reaction occurs. Add a tablespoon of sugar to the water in the glass and stir until the sugar dissolves completely. Record all observations until the sugar dissolves. Remember that this control sample does not involve a chemical reaction because sugar dissolving in water is only a physical change. Then have your child add both of the effervescent tablets to the second glass of water. Record any observations for at least two minutes while the tablets dissolve. Finally, place the thermometer in the third glass and record the initial water temperature. If a thermometer is not available, feel the outside of the glass and record if it feels hot, warm, or cold. Then add a tablespoon of Epsom salt and gently stir the liquid using the thermometer. Make sure to watch the temperature closely and determine how the temperature changes when the salt is added. After performing the reactions, discuss the following questions with your child: 1. In which of the glasses did a chemical reaction take place? How do you know? 2.

Why can you assume that a chemical reaction did not take place in all three glasses? How could you confirm that a chemical reaction took place in the glasses?

Can you write a balanced chemical equation to describe each chemical reaction that occurred?

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C3AB Chemical Reactions

Reading Science Name:

Date:

Group:

Chemical Reactions Affect Van Gogh’s Paintings 1

From vibrant yellow to dull brown, Vincent would not be happy. Vincent Van Gogh painted some of the most vividly colored paintings in the Impressionist era. For several years, however, art conservationists have been noticing a problem. Some of the brilliant yellow sunflowers have been turning a muddy brown. This is due to a chemical reaction in the paint. They need to find a way to stop it. The first step is to understand the reason for this chemical reaction.

Van Gogh’s use of bold, bright colors, especially yellow, was new to the art world of the late 1800s. There was a new type of pigment called “chrome yellow.” Van Gogh realized that it could create vivid, lively colors and used it often in his paintings. However, the bright yellow pigment tends to turn brown after years of exposure to ultraviolet rays, including sunlight. This is due to a chemical reaction in the outer layers of paint.

To learn why this was happening, researchers used x-rays to analyze the exact chemical composition of the paint as it browned. Different artists mixed the pigment with different chemicals to produce paint. The researchers started with three different historically accurate chrome yellow mixtures. Then they artificially “aged” the pigments by exposing them to UV radiation until one of them turned a rich brown. They then identified the changes in that paint’s chromium compounds. Next, they had to determine if the chemical changes in the sample were the same as in the brown from Van Gogh’s paintings. They compared the results of the artificial aging with a very small sample of one of Van Gogh’s paintings. Their suspicions were confirmed. The two paints behaved similarly in their chemical reactions.

The scientists had shown that the paint pigments Van Gogh used were chemically similar to the ones that had been artificially aged. Barium and sulfur were in the white paint that Van Gogh used to mix with the chrome yellow pigment. Scientists suspect that either barium or sulfur is the reactant in the chemical reaction with the yellow chromium. One of the products is a new chromium compound with different properties and a brown color.

Scientists and art conservationists all agree that Van Gogh’s paintings must be protected from UV radiation, particularly sunlight, to prevent them from turning brown. They are still trying to find a chemical antidote. They must identify a process to prevent the chemical reaction. Until then, Van Gogh’s flowers shouldn’t be exposed to sunlight.

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Reading Science 1.

Based on this passage, which of these is the most accurate description of the job of an art conservationist? A.

Mix different types of paint samples

Preserve paintings and other works of art

Buy and sell paintings

Operate x-ray machines

What would be the most likely explanation if none of the tested paint samples had turned brown? A.

The reaction was not actually caused by UV light.

The reaction only happens in real sunlight, not artificial light.

The paint behaves differently on a canvas than it does in a laboratory.

None of the tested mixtures matched the chemical composition of Van Gogh’s paint.

How did modern scientists experiment with Van Gogh’s paints to determine their chemical reactions? A.

They used x-ray analysis.

They used litmus testing.

They burned the paint to examine what elements were left after the chromium was burned off.

They compared the density of modern yellow paint to that of 100-year-old paint.

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Reading Science 4.

Which words help the reader determine the meaning of the word vividly in the first paragraph? A.

Art conservationists

Impressionist era

Brilliant yellow

Muddy brown

What is the main point of the passage? A.

Ultraviolet rays can cause chemical reactions.

Van Gogh was the most famous painter of the Impressionist era.

Art conservationists can always restore art.

Artists’ paints can eventually damage the canvasses.

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Math Connections Name:

Date:

Group:

Balancing Equations 1.

Write a balanced equation for the combustion reaction of propane with oxygen to yield carbon dioxide and water.

Unbalanced equation: C3H8 + O2 CO2 + H2O

Balanced equation:

Write a balanced equation for the reaction.

Unbalanced equation: KClO3 + C12H22O11 KCl + CO2 + H2O

Balanced equation:

Balance the equation for the fermentation of sugar to yield ethyl alcohol.

C6H12O6

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C2H6O + CO2

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Math Connections 4.

Balance the equation for the photosynthesis reaction in green plants. CO2 + H2O

Balance the equation for the synthesis of hydrazine for rocket fuel. NH3 + Cl2

6. •

N2H4 + NH4Cl

Which of the following equations are balanced? The development reaction in photography 2AgBr + 2NaOH + C6H6O2

•

2 Ag + H2O + 2NaBr + C6H4O2

Preparation of household bleach 2NaOH + Cl2

C6H12O6 + O2

NaOCl + NaCl + H2O

Balance the following equations. Mg + HNO3

H2 + Mg(NO3)2

CaC2 + H2O

Ca(OH)2 + C2H2

S + O2 UO2 + HF 250

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C3AB Chemical Reactions

WRITING SCIENCE Name:

Date:

Group:

LOOK

THINK about how chemical reactions occur and how we know they have occurred. During a chemical reaction, molecules of a substance break apart and join molecules of another substance, creating a new compound. Typically, once a chemical reaction has occurred, it cannot be reversed. Chemical reactions can happen quickly, while others take a long amount of time to occur. Various indicators can tell us whether a chemical reaction has occurred. WRITE Explain the various indicators of chemical reactions. What observable evidence is used to infer a chemical reaction has occurred? Provide examples of any experiments you have done in class, and describe the observable evidence that assured you a chemical reaction occurred.

Be sure to clearly state your central idea; organize your thoughts; develop your essay in detail; choose your words carefully; and use correct spelling, capitalization, punctuation, and grammar.

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C3AB Chemical Reactions

WRITING SCIENCE

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High School Chemistry

C3C

The Mole and Avogadro’s Number

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C3C The Mole and Avogadro's Number

Student Journal Name:

Date:

Group:

Background: Understanding the Mole 1. Why was it important for scientists to know the number of atoms in a sample of matter?

2. What was chosen to use as the standard on which to base the atomic masses of all the elements?

3. How many atoms are in 12 g of the standard?

4. What is the name for this number of atoms?

5. What is the SI unit that scientists created to represent this number of atoms?

6. How many marbles are in one mole of marbles?

7. How many atoms are in one mole of sodium?

8. In the study of chemistry, what connection does the mole unit allow you to make? What types of “worlds” does this unit connect? Be as specific as you can.

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C3C The Mole and Avogadro's Number

Student Journal Background: Understanding the Mole, Continued 9. What does percent composition of a chemical compound mean?

10. A formula unit of sodium chloride contains one sodium atom and one chloride atom. What is the percent composition of NaCl? Step 1: Find the molar mass of each element in the compound:

Na=

Cl=

Step 2: Find the total molar mass of the compound:

NaCl=

Step 3: Divide each individual element’s molar mass by the total molar mass of the compound. Multiply by 100 to make a percent. Molar mass of Na × 100 = Total molar mass of NaCl Molar mass of Cl × 100 = Total molar mass of NaCl 11. If done correctly, what should be the sum of the two percentages?

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C3C The Mole and Avogadro's Number

Student Journal Background: Understanding the Mole, continued 12. A compound is composed of 53.3% oxygen (O) and 46.75% silicon (Si). Find the empirical formula of the compound. Step 1. Assume you have 100 g of the compound. This means:

mass(O) = _____ g and mass(Si) = _____ g

Step 2. Find the number of moles of oxygen and silicon using their molar mass. ___g O_____ = _____________ mol O 15.999 g/mol ___g Si____ = __________ mol Si 28.086 g/mol Step 3. Make ratios with the mole amounts. (Put the smallest number in the denominator.) mol O = mol Si

mol O mol Si

Step 4. Round to the nearest whole number or make a fraction with whole numbers. mol O = the subscript for O in the formula mol Si = the subscript for Si in the formula Step 5: Write the formula for the compound. _________________________________________________________________________

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Student Journal Part I: Plan Your Investigation 1. My Question of Inquiry:

2. The Hypothesis:

3. My Prediction:

4. What is the independent variable (also known as the manipulated variable)?

5. What is the dependent variable (also known as the responding variable)?

6. Is there a control group or control variable for this investigation? Explain.

7. What materials, equipment, and technology will you need for this investigation?

8. List all safety precautions that you must take.

9. Follow the procedures listed in the Student Guide to conduct this investigation.

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Student Journal Part II: Implement Your Investigation Collect, Record, and Organize Data 1. Data Collected Mass of empty crucible & lid: Total mass before heating: Appearance of sample: Total mass after heating: Appearance of sample after heating: 2. Data Calculations Mass of magnesium sulfate hydrate sample before heating: Mass of anhydrous magnesium sulfate after heating: Mass of water driven from sample: Analyze Data 1. Determine number of water molecules in magnesium sulfate hydrate: Mass of water lost by the compound Molar mass of water Moles of water Mass of anhydrous magnesium sulfate Molar mass of anhydrous magnesium sulfate (MgSO4 ) Moles of anhydrous magnesium sulfate Make a mole ratio of water to anhydrous magnesium sulfate: Moles of H2O = Moles of MgSO4 Round to the nearest whole number

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Student Journal Part II: Implement Your Investigation, Continued 2. What is the percent composition of water in the hydrate before heating? Show your work in the space provided below.

3. Write the formula for the magnesium sulfate hydrate found in this investigation. 4. What is the name of the hydrate? 5. Write the chemical equation for the decomposition (dehydration) of the magnesium sulfate hydrate based on your results. 6. What was the appearance of the compound before and after it was heated? Why do you think this change occurred?

7. Why did you have to leave the lid on the crucible while you were waiting for it to cool down before you weighed the anhydrous compound?

8. What do you think will happen to the anhydrous magnesium sulfate if you leave it out on the laboratory table and check it the next day?

9. What would happen to the results of your experiment if you did not drive off all of the water before weighing the magnesium sulfate left in the crucible?

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Student Journal Reflections and Conclusions 1. Does the data support or refute your hypothesis? Explain.

2. Was your prediction correct or incorrect? Explain.

3. What type of data did you collect, qualitative or quantitative? Justify your answer.

4. How did the results reveal a relationship between the independent and dependent variables?

5. Where could you have made errors while collecting or organizing data?

6. What do you conclude about this investigation?

7. What would you do differently if you were to conduct this experiment again?

8. Using all of the following terms, develop a graphic organizer. Use additional paper, if needed. Terms: chemical compound, mole, Avogadro’s number, 6.02×1023, atoms, molecules, atomic mass, molar mass, SI unit, amount of a substance, relative mass, percent composition, empirical formula, molecular formula. © Accelerate Learning Inc. - All Rights Reserved

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Student Journal Additional Practice Show all of your calculations. Use additional paper if necessary. 1. How many atoms are in 7.2 moles of chlorine (Cl)?

2. How many moles are in 45.9 g of sulfur (S)?

3. How many moles are in 1.0×1025 atoms?

4. What is the mass of 2.18 moles of copper (Cu) in grams?

5. If I have 25 g of silver (Ag), how many moles do I have?

6. What is the mass of 3.01×1023 atoms of iron (Fe)?

7. How many atoms are in 15.8 g of neon?

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Student Journal Additional Practice, Continued 8. If I have 3.01×1024 molecules of water, how many moles of water molecules do I have?

9. Find the percent composition of potassium permanganate, KMnO4.

10. An unknown compound contains both sulfur and nitrogen. It is composed of 30.1% nitrogen by mass. The rest is sulfur. The actual molar mass of the compound is 92.15 g/mol. Find the molecular formula of the compound.

11. A student measures 1.95 g of a hydrate of cobalt chloride. After heating, 1.06 g of anhydrous cobalt chloride remains. What is the formula of the hydrate?

12. Which compound contains more silicon by percent mass, H4SiO4 (silicic acid) or HSiCl3 (trichlorosilane)? Show your work.

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High School Chemistry

C3C

The Mole and Avogadro’s Number

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STEMscopedia: THE MOLE AND AVOGADRO’S

NUMBER C3C

Reflect Imagine a carpenter who needs enough nails to lay 240 shingles on the roof of a new house. Because each shingle requires 4 nails, the carpenter multiplies 240 by 4 to find the total number of nails he needs: 960. Do you think he will count out each nail separately when he goes to the hardware store to buy supplies? Or will he buy nails already packaged in boxes? Which is easier and why? How do you think the method a carpenter uses to measure the number of nails is similar to the method a chemist uses to measure atoms or molecules in a chemical reaction?

Having the right amount of supplies on hand helps carpenters complete their work efficiently.

The Mole: An SI Unit for Describing an Amount of a Substance Nails are usually packaged in boxes. Each box contains a large number of nails. That’s because many nails may be needed for one project, and buying them by the box is easier. Instead of worrying about individual nails, a carpenter can choose the correct number of boxes of nails needed to complete a project. Chemists use a similar approach when they carry out chemical reactions. Instead of thinking about the individual atoms or molecules involved, they focus on groups of them. Chemists use a unit known as the mole (abbreviated: mol) to define a specific group of particles. One mole is the SI unit for measuring the quantity of matter. It is defined as the amount of a substance that contains the same number of particles as the number of atoms in exactly 12 grams of carbon-12. More simply, 1 mol equals 6.02×1023 particles. These particles can be atoms, molecules, ions, electrons, or any other chemical unit.

Just as a box of nails contains a specific number of nails, a mole of atoms contains a specific number of atoms.

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Scientists in the Spotlight: Amedeo Avogadro Amedeo Avogadro, an Italian scientist who lived from 1776 to 1856, developed the concept of the mole. Avogadro studied to be a lawyer and practiced law near Turin, a city in the northwestern region of Italy. However, Avogadro was fascinated with science. He engaged private tutors in mathematics and science and later gave up law to become a professor of chemistry at the University of Turin.

In 1811, Avogadro published a landmark paper. He had completed some experiments with different types of gases. Avogadro wrote his results in the paper along with the hypothesis that different gases contain the same number of particles whenever the conditions of volume, temperature, and pressure are the same. This was a novel idea: No one had considered comparing the actual numbers of particles in a substance before.

Avogadro’s hypothesis was mostly ignored for more than 50 years. Eventually, scientists recognized its usefulness, and Avogadro’s work became the foundation for the concept of the mole. Scientists later established that 1 mol of a gas occupies a volume of 22.4 L when the temperature is 0°C and the pressure is 1 atmosphere (atm). A sample like this contains 6.02×1023 particles, the numerical value of 1 mol. Today, 6.02×1023 is known as Avogadro’s number in honor of Avogadro’s work. Using Avogadro’s Number Avogadro’s number is used as a conversion factor whenever the number of particles in a substance must be calculated.

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For example, if you have 0.25 moles of hydrogen gas (H2), you can use a conversion factor with Avogadro’s number (shown in red) to determine how many molecules of hydrogen you have:

Alternatively, you may perform the reverse calculation if you know the number of particles in a substance and must determine the number of moles:

Conversions Between Moles and Mass While the mole is a convenient measuring unit for particles, no one has invented a device to actually measure moles directly in a sample. Instead, chemists use an electronic balance to measure the mass of a sample, and then they perform a calculation to convert mass to moles. The calculation uses molar mass as a conversion factor: The mass of a substance can be accurately measured on an electronic balance and used to calculate the number of moles of that substance.

Molar mass is defined as the mass of one mole of a substance. Molar mass has units of grams/mole, and it has the same numerical value as the formula weight (or molecular weight) of the substance. Formula weight (or molecular weight) is determined by adding the atomic masses of all atoms present in one formula unit (or molecule) of the substance. An element’s atomic mass—measured in atomic mass units (amu)— can be found in most copies of the periodic table. For example, suppose you had 55.4 grams of magnesium bromide (MgBr2). First, use the periodic table to determine the formula weight of MgBr2. • One formula unit of MgBr2 consists of one atom of magnesium (Mg) and two atoms of bromine (Br). • The atomic mass of Mg is 24.3 amu. The atomic mass of bromine is 79.9. • Therefore, the formula weight of MgBr2 equals 24.3 amu+(2×79.9 amu), or 184.1 amu.

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Because a substance’s molar mass has the same numerical value as its formula weight, the molar mass of MgBr2 equals 184.1 g/mol. You can use this value as a conversion factor to determine the number of moles in a 55.4-g sample of MgBr2:

Alternatively, suppose you needed exactly 2.00 mol of MgBr2 to carry out a chemical reaction. You can use molar mass to calculate the number of grams of MgBr2 you would need to weigh to obtain the required number of moles:

Considering both Avogadro’s number and molar mass as conversion factors, you can convert between grams, moles, and number of particles:

What Do You Think?

How would you use molar mass and Avogadro’s number to calculate the number of atoms in 1.0 gram of iron? Calculating the Percent Composition of a Compound Molar mass is useful in other types of calculations, too. For example, chemists often want to know the elemental makeup of a compound, including the identity of all the elements present and how much of each element is present. One means for expressing this distribution is by the compound’s percent composition, which gives the percent by mass of each element in a compound. The general formula for calculating percent composition is shown below. Notice that molar mass is used in the denominator:

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For example, the chemical formula of water is H2O. This formula reveals the identities of the elements in the compound: one molecule of water consists of two hydrogen atoms (H) and one oxygen atom (O). Because water consists of two elements, the percent composition calculation will consist of two percentage values: one for each element. To find the percent composition of each element in water, first determine the mass of each element in one mole of water. You also need to determine the molar mass of water by summing the masses of the individual elements:

Next, apply the formula for percent composition by dividing the mass of each element by the molar mass. Multiply the result by 100 to express the result as a percentage:

A good way to check your work is to add the percent values. They should add up to 100.0%. If they don’t, you made a mistake in your calculations somewhere.

A water molecule is 88.8% oxygen and 11.2% hydrogen. (This molecule is not drawn to scale. An oxygen atom has much more mass than a hydrogen atom.)

Look Out Though not explicitly indicated, percent composition is always based on the masses and not on the moles of elements in a compound. Mole fraction is used to indicate the ratio of moles of each element. It is important to avoid confusing these two types of calculations. It helps to think about the case of water as an example. In water, the mole fraction of hydrogen is very high because there are two moles of hydrogen for every one mole of oxygen. Yet, because a hydrogen atom has such a small mass compared to an oxygen atom, the percent by mass of hydrogen in water is very low.

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Empirical and Molecular Formulas The chemical formula of a compound provides information about the elements present in the compound and their mole ratios. However, there are two types of chemical formulas. When chemists analyze new compounds, they often collect data on the elements making up these compounds. They then use the data to determine the following two formulas. • Empirical formula: An empirical formula expresses the lowest whole number ratio of the elements in a compound. • Molecular formula: A molecular formula contains the actual numbers of atoms of each element making up one molecule or formula unit of a compound.

Consider the compound n-octane. One molecule of n-octane An empirical formula expresses has 8 carbon and 18 hydrogen atoms, giving it a molecular the simplest ratio of elements formula of C8H18. The ratio of carbon to hydrogen in this given in the molecular formula compound is 8:18. This is not the smallest whole number of a compound. ratio because both 8 and 18 can be divided by 2 to give a ratio of 4:9. Because the ratio 4:9 cannot be reduced further, it is the smallest whole number ratio. Therefore, the empirical formula for n-octane is C4H9. Often, when a new compound has been discovered, its empirical formula is determined first. This is because a chemist will begin by analyzing the identities of the elements and their percent composition. For example, suppose that an unknown compound was analyzed and found to contain 38.7% carbon (C), 9.8% hydrogen (H), and 51.5% oxygen (O). What is the empirical formula of the compound? To solve this problem, you must assume you have a 100-g sample of the unknown compound. This gives you 38.7 g C, 9.8 g H, and 51.5 g O. Convert each of these mass measurements into moles using the molar mass of each element as a conversion factor:

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Take the smallest mole quantity and divide all of the mole quantities by this number to generate mole ratios:

Rounding each number to an integer, the simplest whole number ratio for C:H:O is: 1:3:1. The empirical formula for this compound is CH3O. Calculating a Molecular Formula The molecular formula of a new compound is usually determined after its empirical formula has been established. In fact, a chemist can use the empirical formula of the compound along with an experimentally determined molar mass to find the molecular formula. For example, suppose the unknown compound above with the empirical formula CH3O was found to have a molar mass of 62.1 g/mol. What is the molecular formula of the compound? You can solve this problem by comparing the experimentally determined molar mass that represents the molecular formula with a calculated molar mass based on the empirical formula. The two molar masses should be related by a whole number value.

empirical formula = CH3O molar mass calculated from empirical formula = 12.01 + 3(1.01) + 16.00 = 31.04 g/mol molecular formula = C?H?O? experimentally determined molar mass = 62.1 g/mol molar mass representing molecular formula molar mass representing empirical formula

62.1 g/mol 31.04 g/mol

2.00

In this case, the two molar masses are related by a whole number factor of 2. This means the molecular formula is related to the empirical formula by a factor of 2. Because the empirical formula is CH3O, the molecular formula must be C2H6O2.

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What Do You Know? Use your understanding of the mole and molar mass to fill in the empty cells in the following table.

Compound Number

Percent composition

Molar mass (g/mol)

5.9% H, 94.1% O

34.02

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159.69

Empirical formula

Molecular formula

Number of molecules or formula units in 1.0 g of the compound

Fe2O3 C10H20O5

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STEMscopedia: THE MOLE AND AVOGADRO’S Connecting With Your Child

NUMBER

Moles and Everyday Substances To help your child develop a better understanding of the mole concept, pose the following challenge: A penny has a mass of 2.50 g and is 2.5% copper (Cu). (A penny contains mostly zinc.) How many pennies are necessary to make up 1 mol of copper? To solve this problem, you must first know the molar mass of copper. According to the periodic table, the average atomic mass of copper is approximately 63.5. Therefore, copper’s molar mass is 63.5 g/mol. How many pennies are necessary to make up 63.5 g, or 1 mol, of copper? To solve this problem, calculate the mass of copper in 1 penny: 2.50 g×0.025=0.0625 g. In other words, one penny contains approximately 0.0625 g of copper, and 1000 pennies contain approximately 62.5 g of copper. You need approximately 1016 pennies to make up 1 mol of copper. Once your child has solved this problem, he or she can solve similar problems involving the composition of other coins. For example: • A nickel is 5.00 g and is 75% copper (Cu) and 25% nickel (Ni). • A dime is 2.268 g and is 91.67% copper and 8.33% nickel. • A quarter is 5.67 g and is 91.67% copper and 8.33% nickel. You can apply the same principles to mixtures other than coins. For example, air is approximately 78% nitrogen and 21% oxygen. (The remaining 1% of air is made up of various other gases.) How many mol of air is necessary to make up 1 mol of oxygen? Discuss the following questions with your child: • Why do scientists measure quantity of matter in moles? (What are the advantages of working in moles?) • Why might 1 mol of one substance have a different volume than 1 mol of a different substance?

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C3C The Mole and Avogadro’s Number

Reading Science Name:

Date:

Group:

Avogadro: From Forgotten Scientist to Annual Mole Day 1

In 1811, Amedeo Avogadro, an Italian lawyer-turned-physicist, discovered that equal volumes of all gases at constant temperature and pressure have an equal number of particles. However brilliant that premise, he did not know what that number should be! Moreover, his ideas were not fully accepted until years later for several reasons.

First, Avogadro did not have a great deal of scientific data to back his hypothesis. He also worked in isolation with little communication with other scientists and was not very well-known in Italy at the time. The journal that published his ideas was not well-known either.

It was not surprising that Avogadro could not calculate what the number of molecules should be, considering the world of chemistry during 19th-century Europe. At that time, atoms and molecules were not separate concepts, so chemists used them practically as synonyms. Moreover, chemists then used equivalents rather than atomic weights when solving their chemistry equations. In short, there was no set way to measure the “amount” of an element or compound, regardless of volume or weight.

By 1858, Avogadro’s ideas were all but forgotten, until Italian chemist Stanislao Cannizzaro presented work that showed that atomic and molecular weight could be found by using Avogadro’s hypothesis. Subsequently, acceptance of Avogadro’s law grew, and it served as the basis for many other discoveries in chemistry.

So what does this have to do with Avogadro’s number? Even though he earned distinction for his hypothesis about equal volumes of gases, he never stated a specific number of molecules for a specific weight-equivalent of a substance. An estimate of that number was not discovered until 1909 by Jean Baptist Perrin and named the constant in honor of Avogadro for having found the connection between the amount of a substance and the number of particles, or molecules.

Modern chemists used the charge on an electron as 1.60217653 x 10-19 coulombs per electron. When the charge on a mole of electrons is divided by the charge on a single electron, the value of Avogadro’s number is calculated as 6.02214154 x 1023 particles per mole. A mole is the amount of pure substance containing the same number of units as atoms in 12 grams of Carbon-12 (i.e., 6.023 X 1023). Today, Avogadro’s number (6.02 x 1023) is called Avogadro’s constant. That number is celebrated annually as worldwide Mole Day on the 23rd day of the 10th month of year (October 23) from 6:02 a.m. to 6:02 p.m., to build interest in chemistry among students in secondary schools and universities.

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Reading Science 1.

What did Amedeo Avogadro discover? A.

All gases weigh the same for equal volumes.

All equal volumes of gases at constant temperature and pressure have an equal number of particles.

All gases have the same condensation point.

All gases at the same temperature have equal volumes.

Why did Avogadro’s hypothesis not receive the attention it deserved? A.

He didn’t have much scientific data to back it up.

He was isolated from colleagues.

The journal in which he published his findings wasn’t widely read.

All of the above.

What is the value of Avogadro’s number? A.

6.023 x 1023

1.60 x 10-19

6.02 x 10-19

None of the above

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Reading Science 4.

In reference to chemistry, what is a mole? A.

A small burrowing insectivore

The amount of pure substance in 1 mL of water

The amount of a pure substance containing the same number of units as atoms in 12 grams of Carbon-12

The charge in coulombs per electron

Why was Avogadro’s hypothesis significant? A.

It allowed for atomic and molecular weight to be found.

It helped lead to the mapping out of the nucleus of an atom.

It explained the big bang theory mathematically.

It led to the discovery of the gravitational constant for Earth.

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C3C The Mole and Avogadro’s Number

Math Connections Name:

Date:

Group:

Part I: The Mole—Calculating the Number of Atoms A mole is the amount of an element whose mass in grams is equivalent to its atomic mass. The number of atoms in one mole of any element is referred to as Avogadro’s number, 6.022×1023 atoms. Avogadro’s number of atoms is also equal to the element’s atomic mass. Using the information above, you can use atomic mass to calculate the amount of an element or the number of atoms in a sample. Example: How many moles and how many atoms of carbon are in a sample weighing 20.26 grams? The atomic mass of carbon is 12.0107 amu. 20.26 g of C× 1.6868 mol C×

1 mol C =1.6868 mol C 12.0107 g of C

6.022×1023 atoms C =1.015×1024 atoms 1 mol C

1. A nickel coin is actually made of 75% copper and 25% nickel. A nickel coin weighs 5.0 grams. a. The atomic mass of copper is 63.546 amu. How many moles and how many atoms of copper are in one nickel?

b. The atomic mass of nickel is 58.6934 amu. How many moles and how many atoms of nickel are in one nickel?

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Math Connections 2. What is the mass in grams of a 2.56-mole sample of gold?

3. How many moles are in a 2-kg sample of iron?

4. How many moles are in a 5.67-g sample of NaCl?

5. How many moles of ions are in 5.5 moles of NaCl?

Part II: Calculating Molecular Mass and Number of Molecules Molecular mass is the sum of the atomic masses for all atoms in a molecule. One mole of any compound is equivalent to the molecular mass in grams of the compound. 6. Write an equation to calculate the molar mass of 1 mole of N2O (laughing gas).

7. Solve the equation to find the molecular mass of laughing gas.

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Math Connections 8. Calculate the molecular masses of the following: a. Sulfuric acid: H2SO4

b. Ammonium: NH4Cl

c. Rubbing alcohol: (CH3)2CHOH

9. How many moles are in a 347.89-g sample of ammonium?

10. What is the molecular mass of milk of magnesia (Mg(OH)2) if 0.0275 mol weighs 1.603 g?

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Math Connections Part III: Calculating Molar Volume of a Gas Molar volume is the volume that one mole of a substance occupies at a given pressure and temperature. One mole of any gaseous compound fills 22.4 L in a closed system. 11. Write an equation to calculate the molar volume of 5 moles of CO2.

12. Solve the equation to find the molar volume of 5 moles of CO2.

13. How many moles are in a 996.4-L container of mixed gases.

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C3C The Mole and Avogadro’s Number

WRITING SCIENCE Name:

Date:

Group:

LOOK

THINK A carbon-12 atom has a mass of 12 amu. 6.02×1023 carbon-12 atoms, or 1 mole of carbon-12 atoms weighs 12 grams. Amadeo Avogadro proposed a concept in 1811 that related the number of molecules and the volume of gases at the same temperature. Scientists later took this idea and developed it into the concept of mole. The mole concept relates the relative mass of elements as compared to a standard, that standard being carbon-12. Carbon-12 was assigned the mass of 12 grams, and it was determined that 12 grams of carbon-12 contained 6.02×1023 atoms, which is also called Avogadro’s number, or the Avogadro constant. This number is referred to simply as a mole. In other words, 1 mole of carbon-12 contains 6.02×1023 atoms. A mole is a unit of measurement that can be used to describe the number of anything, including cars or pennies. If a person owned 1 mole of pennies, he or she would have 6.02×1023 pennies. If 1 mole of pennies were stacked up, it would reach the sun and back over 500 million times! With the mass of carbon-12 assigned to be 12 grams, other elements can then be compared to determine their mass. If 6.02×1023 atoms of magnesium are found to have twice the amount of mass as 6.02×1023 carbon-12 atoms, then the molar mass of magnesium is determined to be 24 grams. Whereas, if 1 mole of helium atoms is found to have only 1⁄3 of the mass of 1 mole of carbon-12 atoms, then the molar mass of helium is determined to be 4 grams. The mass of each element on the periodic table has been determined in this way by its relative mass to carbon. WRITE Discuss the relationship between the relative masses of the elements and the mole concept. Specifically, explain how the molar mass of hydrogen was determined to be 1 gram after one mole of carbon-12 was assigned to have a mass of 12 grams. Be sure to clearly state your central idea; organize your thoughts; develop your essay in detail; choose your words carefully; and use correct spelling, capitalization, punctuation, and grammar. © Accelerate Learning Inc. - All Rights Reserved

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WRITING SCIENCE

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High School Chemistry

C3DE

Stoichiometry

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C3DE Stoichiometry

Student Handout Name:

Date:

Dimensional Analysis Dimensional analysis is a method of solving chemistry problems that use ratios of values, or “per” expressions, that are directly proportional to each other. Use this process to solve the following question:

How many seconds are there in five years? 1.

Use the table, template, and guidelines below to help you solve the problem. Set up all dimensional analysis problems in the same way. • • • •

Start with what you know or what you are given. Place this value, with units, in the top space on the upper far left. Next, place the unit of the value that you are looking for in the space on the upper far right. List all of the conversion factors (“per” expressions) that you will need to get from your given to your wanted. For instance, how many seconds are there in one minute? Use these conversion factors to get from your given to your wanted. Place the units in the template so that all units cancel. If all of the units cancel from your given to your wanted, then you know that you have set the problem up correctly.

The first one has been completed for you. Use all of the spaces provided. Given

Wanted

Ratio

5 years

5 years 1 year

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C3DE Stoichiometry

Student Handout Pro“per”tional Relationships One ratio of values, or “per” expression, that is often used in chemistry is the mole ratio. This is the ratio of the number of moles of each compound present in a balanced chemical equation. This proportion will allow you to calculate the amount of any compound needed or used in a balanced chemical equation. But that is the most important part. The chemical equation must be correctly balanced for the proportion to be correct. Look at the following two examples to help you better understand how these ratios can be used to help you solve chemical calculations, or stoichiometry problems.

NaHCO3 (s) + HCl (aq) -> NaCl (s) + CO2 (g) + H2O (l) •

This ratio may be used to calculate amounts of reactant to product, and vice versa. There is/are _____ mole(s) of NaHCO3 (baking soda) to ______ mole(s) of NaCl (sodium chloride).

•

This ratio may be used to calculate amounts of each reactant needed or used. There is/are _____ mole(s) of NaHCO3 (baking soda) to ______ mole(s) of HCl (hydrochloric acid).

In this problem, the number of atoms of each reactant equals the number of atoms of each product. Let’s try another problem where the coefficients in the balanced chemical equation are different.

C3H8 + 5 O2 -> 4 H2O + 3 CO2 There is/are _____ mole(s) of C3H8 to _______ mole(s) of H2O. There is/are _____ mole(s) of O2 to _______ mole(s) of C3H8. There is/are _____ mole(s) of H2O to _______ mole(s) of CO2. How many moles of propane (C3H8) are required to create three moles of carbon dioxide?

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Student Journal Name:

Date:

Group:

Part I: Stoichiometry 1. Define stoichiometry and explain the type of information it provides.

2. What is a mole ratio, and how is it used?

3. What is dimensional analysis, and what types of numerical relationships are used in this process? Specifically, what are “per” expressions, and how are they used?

4. What is the theoretical yield of a chemical reaction? Why is it rarely used in real life?

5. How do actual and theoretical yields of a chemical reaction generally compare? What can you determine about a reaction by analyzing these yields together?

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Student Journal Part II: Percent Yield 1. Sodium bicarbonate and hydrochloric acid will react to produce sodium chloride, carbon dioxide gas, and water. Write the balanced chemical equation for this reaction in the space below. Make sure to include the correct state symbols.

2. Fill in the table below as you measure the masses of your reactants. Use dimensional analysis to find the mass of HCl used. The density of 1 M HCl is 1.015 g/mL. One drop of HCl is equivalent to 0.05 milliliters. Show your work below.

Mass of empty evaporating dish: Reactant or Product

Chemical Formula

Molar Mass (g/mol)

Mass Used or Produced (g)

Sodium Bicarbonate

Hydrochloric Acid

Sodium Chloride 3. Using the balanced chemical equation and the mass of sodium bicarbonate used, calculate the theoretical yield of sodium chloride that your reaction may produce. Use the process of dimensional analysis to perform your calculation. Circle your answer.

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Student Journal Part II: Percent Yield, Continued 4. In the space provided below, record the actual yield of sodium chloride created.

5. What errors may have created a difference between the theoretical and actual yield?

6. Using your actual and theoretical yields, calculate the percent yield of your reaction.

7. Compare the percent yields of the other groups in your class. Are they identical or do they vary? Discuss the factors that may have caused variation in the class results.

8. Using the balanced chemical equation and the mass of sodium bicarbonate used, find the theoretical mass of carbon dioxide that this reaction could produce.

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Student Journal Reflections and Conclusions 1. What did you discover about the products and reactants of a chemical reaction?

2. Why was it important to accurately measure your product in this investigation? What could have happened if your measurements were inaccurate?

3. How did the ratios between the chemical values (or “per” expressions) help you solve your percent yield problem? That is, which ratios did you use, and how did those ratios help you convert from compound to compound in the chemical equation?

4. How could you modify your investigation to get a better percent yield? Be as specific as you can.

5. The mass of the products in the evaporating dish was less than the mass of the reactants. What contributed to the difference in the masses?

6. Using all of the following terms, develop a graphic organizer based on what you have learned in this Explore. Use another sheet of paper, if needed. Terms: stoichiometry, calculation, reactant, product, balanced chemical equation, coefficient, ratio, “per” expression, molar mass, mole ratio, direct proportion, dimensional analysis, theoretical yield, actual yield, percent yield

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Student Journal Name:

Date:

Group:

Limit Your Fizz

Part I: Trials

Calculate the molar mass of the substances involved in the reaction. 1.

Citric acid (H3C6H5O7)

Baking soda (NaHCO3)

Mixture

Visual Observations

Taste

Trial 1 Trial 2 Trial 3 Experiment 1 a.

Did the mixture fizz? ____________________________________________________

Was it bitter? __________________________________________________________

Was it sour? ___________________________________________________________

Based on the previous trials and the taste of this experiment, which reactant is the limiting reactant? _______________________________________________________

Based on the previous trials and the taste of this experiment, which reactant is the excess reactant? _______________________________________________________

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Student Journal Part II: Calculations You will be given 0.4 g of H3C6H5O7. Using mole ratios, calculate how many grams of baking soda are needed to make the perfect fizzy Kool-Aid® drink and have all the reactants used up. H3C6H5O7 + 3NaHCO3 → Na3C6H5O7 + 3H2O + 3CO2

Part III: Try It Experiment 2: Kool-Aid® with Citric Acid and Sodium Bicarbonate 1. List the amounts of the reactants (your recipe) you are planning on using, based on your calculations.

List the procedures you will use to test your amounts.

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Student Journal 3.

Record your observations, including the taste.

Questions 1. Explain how your calculated value worked out.

How could you change the recipe you calculated to improve it?

If you are given 15 g of baking soda, how many moles of citric acid are needed to react all of the baking soda?

When you are given 8 g of citric acid and 5 g of baking soda, which one is the limiting reactant?

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C3DE

Reflect Suppose your school cafeteria wants to serve grilled cheese sandwiches. The cafeteria manager must calculate the quantity of bread and cheese he needs to order so that enough sandwiches can be made to feed the number of students expected to buy them. One of the things the manager needs to know is the number of students who will buy grilled cheese sandwiches. Beyond that, what other things must the manager know to determine how much bread and cheese to buy? Does it make sense that he decides to buy twice as many bread slices as cheese slices? Why? Stoichiometry refers to the numerical relationships between the participants in a chemical reaction. To make a grilled cheese sandwich, you use two slices of bread and one slice of cheese. The ratio of bread to cheese is 2:1. The cafeteria manager uses this ratio to calculate the number of bread and cheese slices he will need to make a sufficient number of sandwiches to satisfy the students buying them. In doing this kind of calculation, the manager is applying the same concept that chemists use when they plan chemical reactions. This concept—known as stoichiometry—deals with the numerical relationships between reactants and products in chemical reactions. For example, if a chemist reacts hydrogen gas with oxygen gas to form water, the reaction uses two moles of hydrogen for every one mole of oxygen and produces two moles of water. (A mole is an SI unit used to describe an amount of a substance. One mole is equal to 6.02 × 1023 atoms, molecules, or formula units of a substance.) We know this because of information from the balanced chemical equation for the reaction. The coefficients in the balanced equation show that there is a 2:1 mole ratio of reaction between hydrogen and oxygen. Therefore, if the chemist wants to be sure that all of the oxygen reacts to form water, she will have to supply twice as many moles of hydrogen as the number of moles of oxygen available. Stoichiometry is useful when making certain types of calculations. Earlier, you considered the task of a cafeteria manager who must calculate the quantity of bread and cheese he needs to make enough sandwiches to feed the students at his school. To summarize this process and the numerical relationships involved, we can write an equation like the one below: 2 bread slices + 1 cheese slice 1 sandwich stoichiometry: 2 1 1

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STEMscopedia: STOICHIOMETRY The equation above establishes the stoichiometric relationships between the bread, cheese, and sandwiches produced during this process. These relationships become the basis for conversion factors that allow the cafeteria manager to carry out his calculations of the amounts of bread and cheese he needs to buy. Suppose that 150 students will buy grilled cheese sandwiches. The manager can use that number along with ratios from the equation above as conversion factors to calculate the amounts of bread and cheese needed: Slices of bread needed to buy: 150 sandwiches X 2 slices of bread = 300 slices of bread 1 sandwich Slices of cheese needed to buy: 150 sandwiches X 1 slice of cheese= 150 slices of cheese 1 sandwich In a similar way, chemists use stoichiometric relationships between reactants and products in a chemical reaction to make calculations. For example, suppose a chemist has 3.9 grams of acetylene gas, C2H2. He wants to react it with oxygen to produce carbon dioxide and water according to the following balanced chemical equation. How many moles of oxygen are needed to react with 3.9 g C2H2, and how many moles of carbon dioxide will be produced from the reaction? 2 C2H2(g) + 5 O2(g) 4 CO2(g) + 2 H2O(l) stoichiometry: 2 5 4 2 To solve this problem, first use molar mass to convert grams of C2H2 to moles. Then solve the two problems using the mole ratios of reactants and products as conversion factors: Moles of C2H2 calculated from grams: 3.9g C2H2 x 1 mole C2H2 = 0.15 moles C2H2 26.04g C2H2 Moles of O2 needed to react: 0.15 moles C2H2 x 5 moles O2 = 0.38 moles O2 2 moles C2H2

Moles of CO2 produced:

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0.15 moles C2H2 x 4 moles CO2 = 0.30 moles CO2 2 moles C2H2

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STEMscopedia: STOICHIOMETRY The example above demonstrates that it does not matter whether you are interested in the amount of a reactant or a product when it comes to using stoichiometry. As long as you have a balanced chemical equation, you know the mole ratios of every pair of participants in that reaction, whether the pair consists of two reactants, two products, or a reactant and a product. A limiting reagent determines the amount of product produced. Chemists rarely mix reactants in the exact ratio specified by the balanced equation. Usually they mix reactants in ratios that allow one reactant to be the limiting reagent. The limiting reagent is the reactant that limits, and therefore determines, the amount of product produced. Once the limiting reagent is gone, the entire reaction comes to a halt. The other reactants are said to be in excess because unreacted portions of these reactants remain as leftovers after the reaction is over. Take the following reaction as an example. Three moles of hydrogen gas react with one mole of nitrogen gas to produce two moles of ammonia: 2 NH3(g) 3 H2(g) + N2(g) stoichiometry: 3 1 2 In the illustration below, N2 is shown reacting with the available H2 in the 3:1 molar ratio described by the balanced chemical equation for reaction. However, notice that although all of the available hydrogen reacts, the same is not true of the nitrogen. Some of the nitrogen remains unreacted. In this particular instance, hydrogen is the limiting reagent and nitrogen is present in excess. Therefore, all of the hydrogen is consumed and none remains after the reaction has stopped. Some nitrogen does remain after the reaction because more of it was available than there was hydrogen to react with. Use the limiting reagent concept to determine which reactant is limiting if you combined 0.65 moles of H2 with 0.35 moles of N2 and allowed the reaction above to take place. You can do this by calculating the moles of ammonia you would expect from each of the reactants if they reacted completely. Then compare the two values. Whichever produces the smaller quantity of product must be the limiting reagent. Moles of NH3 produced by all of the available N2:

0.35 moles N2 x 2 moles NH3 = 0.70 moles NH3 1 moles N2

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STEMscopedia: STOICHIOMETRY Moles of NH3 produced by all of the available H2:

0.65 moles H2 x 2 moles NH3 = 0.43 moles of NH3 3 moles H2

The calculation shows that the limiting reagent must be hydrogen because it produced the smaller quantity of product. Therefore, when you run this reaction, you can expect to produce 0.43 moles of ammonia and also to have some unreacted nitrogen gas left over from the reaction. No hydrogen gas will remain since it will be completely consumed.

Look Out As the previous problem demonstrates, the reactant present in the smallest quantity is not necessarily the limiting reagent. In the case above, nitrogen was present at a much smaller quantity than hydrogen. However, because each mole of nitrogen reacts with three moles of hydrogen, this requires a threefold higher quantity of nitrogen with respect to hydrogen. Therefore, you must consider both the stoichiometric ratio of reaction as well as the quantity of reactant present in order to determine which reactant is limiting. Also, let’s make sure the answer has the correct significant figures. In any equation involving multiplication or division, the number of significant figures initially given in the problem need to equal the number of significant figures in the final answer. Remember, any non-zero number is considered significant, and zeros are only considered significant if they trail a non-zero number. Therefore, 0.051 has only two significant figures, 1.062 has four significant figures, and 8.010 has four significant figures.

What Do You Think? Take a look at the photograph at right and the balanced chemical equation that describes the reaction taking place. The two reactants are iron (Fe) and oxygen (O2), and the product is iron (III)—or ferric—oxide (Fe2O3). 4Fe + 3O2 2Fe2O3 If you start with one mole of each, one of these reactants serves as the limiting reagent. Can you identify which reactant that is? Explain your reasoning. You can check your answer before the What Do You Know? section of this companion. Chemists analyze the amount of product formed during a reaction.

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STEMscopedia: STOICHIOMETRY When chemists describe the amount of product formed in a chemical reaction, they use the term, yield. There are three different ways this term is used when talking about the results from a chemical synthesis. The three terms are theoretical yield, actual yield, and percent yield. • Theoretical Yield: Theoretical yield is a calculated value. It represents the quantity of product expected based on the quantity of a limiting reagent used in a reaction. This yield assumes that 100 percent of the limiting reagent is converted to products. • Actual Yield: Actual yield is an observed, measured value. It represents the actual quantity of product formed, isolated, and then measured following a chemical reaction. Often the actual yield is lower than the theoretical yield. For example, a reaction may not proceed with 100 percent conversion of limiting reagent to products. In addition, some product may not be recovered during the steps following the reaction intended to isolate the product. Finally, some products are not very stable and may break down after the reaction, which can also decrease the final quantity isolated. • Percent Yield: Percent yield is a calculated value that provides an indication of how successfully product was isolated. Percent yield is found by dividing the actual yield by the theoretical yield and then multiplying the result by 100 to express a percentage. Percent Yield = Actual Yield x 100 Theoretical Yield Suppose that a chemist carries out a chemical reaction and isolates 135 grams of product. If the theoretical yield for this product was 168 grams, what percent yield did the chemist achieve? Percent Yield = 135 grams 168 grams

100 = 80.4%

Notice that the percent yield calculation can be carried out using either mass or moles. You are calculating a ratio, and so the units used are not critical. It is only important that the same units be used for both the numerator and denominator. Everyday Life: Percent Yield and Pharmaceuticals The pharmaceutical industry produces drugs and therapeutic agents using chemical reactions. Many of the compounds synthesized by pharmaceutical companies involve multiple steps. In these cases, even if the percent yield of every step is very high, the overall yield can wind up quite low because of the loss of yield at each step. For example, suppose that six different reactions are required to synthesize a drug, and that each reaction has a percent yield of 95 percent. The overall yield would be as follows: 0.95 × 0.95 × 0.95 × 0.95 × 0.95 × 0.95 × 100 = 74%

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STEMscopedia: STOICHIOMETRY This is a much lower yield than 95 percent and presents a problem for the pharmaceutical company. It costs money to buy the raw materials and go through all of the processes to synthesize and evaluate drugs. In order to stay in business, the company has to charge enough to cover costs and to make a profit. Therefore, the lower the overall percent yield for synthesizing a drug, the more money the pharmaceutical company has to charge for that drug. As a consequence, pharmaceutical companies are constantly researching better and more efficient ways to synthesize their drugs.

What Do You Think? In the reaction described previously, four moles of iron react with three moles of oxygen gas to produce two moles of ferric oxide: 2 Fe2O3 4 Fe(s) + 3 O2(g) stoichiometry: 4 3 2 Now, calculate how much ferric oxide would be produced if you began the reaction with one mole of each reactant: Moles of Fe2O3 produced by all available Fe: 1 mol Fe x 2 moles Fe2O3 = 0.50 moles Fe2O3 4 moles Fe Moles of Fe2O3 produced by all available O2: 1 mol O2 x 2 moles Fe2O3 = 0.67 moles Fe2O3 3 moles O2 The calculation shows that the limiting reagent must be iron because it produced the smaller quantity of product. What Do You Know? Use your knowledge of stoichiometry to answer the questions about a chemical reaction. In the reaction, magnesium metal reacts with hydrochloric acid according to the following equation: Mg(s) + 2 HCl(aq) MgCl2(aq) + H2(g) Suppose six grams Mg is combined with 10 grams HCl to carry out this reaction. • Which reactant will be the limiting reagent? Answer: HCl 302

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STEMscopedia: STOICHIOMETRY •

What is the theoretical yield of MgCl2 in moles? in grams? Remember to use significant figures in your final answer.

Answer needs 3 sig figs: 0.137 moles MgCl2; 13.0 grams MgCl2 • If the reaction is carried out and 10.3 grams of solid MgCl2 is collected after the reaction is over, what is the percent yield of this product? Answer: 79.2 percent

Connecting With Your Child Practice Stoichiometric Concepts To help your child better understand stoichiometric concepts, use physical models to represent a chemical reaction. You will need a box of large paper clips and a box of small paper clips. 1. Begin by having your child make 10 or more S2 molecules by linking 10 pairs of small paper clips together. Place these in a pile on a flat work surface. 2. Next, have your child make a pile of large paper clips representing the reactant, L, on the work surface next to the S2 molecules. 3. Instruct your child to begin reacting the L and S2 molecules to produce the product molecule, L2S. Do this by unlinking paper clips making up the S2 molecules and then linking together two large paper clips with one small paper clip to create product L2S molecules. 4. Have your child continue until the reaction comes to an end. Ask your child why the reaction ended. Ask: Which reactant was the limiting reagent in this case? Then ask whether other starting conditions could have resulted in a different limiting reagent. 5. Challenge your child to repeat the exercise. This time, work out the ratio of reaction between the two reactants. Then have your child write out the chemical equation for the reaction and balance it. Ask whether the ratio they found using the model agrees with the ratio predicted by the balanced chemical equation for the reaction. 6. Repeat the exercise using other reactant pairs. Some possibilities include: a. LS + S3 L2 + S2 b. LS4 + L2 LS2 Here are some additional questions to discuss with your child: 1. In what way are ratios useful in describing the reactions in this model? 2. How could you expand this model to show that mass is conserved in a chemical reaction? 3. Why is this model not a good one for modeling the concept of percent yield?

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Reading Science Name:

Date:

Group:

The Meeting of Particles 1

Chemists study chemical reactions. During a chemical reaction, reactants with specific properties will go through a chemical change to form products with different properties. In the process, the starting reactants are transformed, and brand new products are formed. The processes that occur during the chemical reactions are not perfect, however. There are many variables that can affect the amount of product formed. These can include temperature, catalysts, impurities, or equipment used. It is important for chemists to be able to both anticipate and calculate for any error that will occur during the reaction. Notice that the previous sentence does not say “that might occur.” You can be certain that there will never be an ideal chemical reaction.

Chemists must first calculate the theoretical yield of a reaction. They will find how much product could “ideally” be formed. This ideal reaction would occur if every atom involved in the reaction connected and reacted with all other atoms producing a 100% yield of product. As already stated, this rarely occurs in real life. In real life, the amount of the actual product created or produced varies somewhat with each experiment. This amount is known as the actual yield. These two values can then be used to calculate the percent yield. This is the ratio of the actual yield to the theoretical yield for a chemical reaction, usually expressed as a percentage.

It would be highly unlikely to have an ideal chemical reaction. Furthermore, as you have seen from a balanced chemical equation, each reaction operates with a fixed ratio of molecules. In other words, a certain number of molecules of each reactant are required to produce a certain number of product molecules. As these molecules are at the particulate level, it is impossible to measure out the exact number of molecules that you would need for each reaction. Scientists must therefore measure reactants at the macroscopic level, for example, in grams. This information is then used to calculate how much product is theoretically expected. This is the expected theoretical yield. It reflects the ideal outcome of the reaction, where the actual yield is what is actually produced.

This connection between theoretical yield and actual yield comes into play in modern industry. There are companies that produce the raw chemicals that other companies use in their manufacturing processes. Other companies rely on accurate measurements of chemical reaction yields to make the highest profit possible. It is critical that those in charge of the chemical processes can accurately calculate the amount of reactant that will produce the best yield of product.

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Reading Science Continued 5

For example, imagine that a specific reaction at a chemical factory requires two different chemical reactants. One reactant is easy to obtain and very inexpensive. The other reactant is very rare and very expensive. It seems logical that the company would want to limit the amount of the expensive reactant that was used to maintain a profit for the company. To do this, the chemist in charge must be able to calculate exactly how much of the expensive reactant is required to produce a specific amount of product. It is important to avoid wasting any of this reactant. In other words, it is critical to know exactly how much product could theoretically be produced with the expensive product.

The chemist must then determine how much of each reactant to use for this reaction. The chemist will keep in mind that this chemical reaction can only proceed as long as the two chemicals involved are reacting with each other. Once one reactant is completely used up (consumed), the reaction will stop. No further product will be produced. The chemist will use stoichiometry to find out what the limiting reagent of this particular reaction is. The limiting reagent is the reactant that will be consumed first. Ideally, for this company, the chemist will want the expensive reactant used first so that there is no waste. Therefore, the chemist will perform the necessary calculations to determine the mass of each reactant to use so that the expensive material is the limiting reactant, and the inexpensive material is the excess.

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Reading Science 1 Paragraph 2 discusses the different types of yields that a reaction can produce. If the reactants produce 100% of the product, what is the yield called? A Ideal yield B Theoretical yield C Actual yield D Chemical yield

2 In a chemical reaction involving two different reactants, the reactant that is consumed first is called– A the limiting reactant. B the consumed reactant. C the critical reactant. D the actual reactant.

3 To properly calculate how much reactant needs to be used in a chemical reaction to produce a specific amount of product, what method of chemical calculation must be used? A Balanced chemical equations B Stoichiometry C Electron dot formulas D Mass equations

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Reading Science 4 In chemical calculations, the theoretical mass of the products and the actual mass of the products can be used to determine what? A The mass of the products B The chemical formulas of the products C The percent yield of the products D None of the above

5 A scientist works for a chemical manufacturing company that creates valuable solutions. In the reaction that the scientist is working on, a red solution in one test tube is an inexpensive reactant, and a green solution in another test tube is a very expensive reactant. The chemist must use stoichiometric calculations to determine the best way to combine these solutions to gain the most profit. Which solution would the chemist want to make the limiting reactant? A The red solution B The green solution C Both solutions should be limiting D Neither solution should be limiting

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Math Connections Name:

Date:

Group:

Part I: Stoichiometry A balanced chemical equation describes the ratio of reactants needed to create a certain product. However, when products are created, it’s impossible to count the actual molecules. Instead, the weight of the reactants is used to ensure the correct amount of each substance is used. A mathematical process called stoichiometry is used to convert between moles and mass of a substance in a reaction. Problem: How many moles of hydrogen are needed to react with 35.0 grams of oxygen? 2H2 + 1O2

2H2O

Step 1: Calculate how many moles of O2 are in 35 grams of O2. Molar mass of O2 = 31.9988 35 g O2 ×

1 mol O2 = 1.0937 mol O2 31.9988 g O2

Step 2: The coefficients describe how many moles of each reactant are needed in the balanced equation. For every 1 mole of O2, 2 moles of H2 are needed. Grams of H2 = 1.0937 mol O2 ×

2 mol H2 4.03176 g H2 × = 8.8198 g H2. 1 mol O2 1 mol H2

1. How many moles of salt (NaCl) are in a teaspoon of salt that contains 5.00 g?

2. How many grams are in 2.578 mol of C7H5BiO4, the main ingredient in Pepto-Bismol tablets?

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Math Connections 3. Nitric acid is a chemical used in fertilizers and explosives. It’s created by a mixture of ammonia and oxygen. NH3 + 2O2

HNO3 + H2O

How many grams of O2 are needed to react with 80.0 g of NH3?

4. Acetaminophen, commonly known as Tylenol, has the formula C8H9NO2. How many moles of Tylenol are in a tablet weighing 500 mg? How many molecules?

5. Tylenol is prepared by reaction of p-aminophenol (C6H7NO) with acetic anhydride (C4H6O3) according to the following equation: C6H7NO + C4H6O3 p-aminophenol acetic anhydride

C8H9NO2 + C2H4O2 acetaminophen

acetic acid

a. How many grams of acetic anhydride are needed to react with 5.50 g of p-aminophenol?

b. How many grams of acetaminophen will result?

c. How many grams of acetic acid are formed as a by-product?

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Math Connections Part II: Yields of Chemical Reactions The amount of product actually formed is called the yield of the reaction and is typically less than the amount predicted by calculations. The percent yield is the amount of product actually formed in a reaction divided by the amount theoretically possible and multiplied by 100%. Percent yield =

Actual yield of product × 100% Theoretical yield of product

6. Clamshells are made of calcium carbonate. When they are heated, carbon dioxide is released and a white substance, calcium oxide, is left. What is the percent yield of the reaction if 4.2 g of calcium oxide is obtained from reaction of 10.2 g of calcium carbonate with sufficient heat? CaCO3(s) + heat

CaO(s) + CO2(g)

7. Medical anesthetic uses diethyl ether (C4H10O), synthesized by the treatment of ethyl alcohol (C2H6O) with an acid. How many grams of diethyl ether would you obtain from 60.0 g of ethyl alcohol if the percent yield of the reaction is 82%? 2C2H6O(I ) + acid

C4H10O(I ) + H2O(I )

8. Iron phosphate (III) reacts with sodium sulfate to produce iron (III) sulfate and sodium phosphate. For the balanced equation shown below, if the reaction of 90.0 grams of FePO4 produces a 56% yield, how many grams of Na3PO4 would be produced? 2FePO4 + 3Na2SO4

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Fe2(SO4)3 + 2Na3PO4

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Math Connections Part III: Calculations of Limiting Reagents When chemists use reactions to produce various substances, the molar ratio of the reactants often is not kept in true proportion based on the balanced chemical equation. Many times an excess of water or other inexpensive reactants are used to make sure the intended products are created. When this occurs, the limited reagent (reactant with the least amount) determines the extent of the reaction. The other reactant is called the excess reagent. 9. Aluminum hydroxide is a chemical used as a fire retardant as well as an antacid. Aluminum sulfite reacts with sodium hydroxide to form sodium sulfite and aluminum hydroxide. In this reaction 20 g of aluminum sulfite is allowed to react with 20 g of sodium hydroxide. Al2(SO3)3 + 6 NaOH

3 Na2SO3 + 2 Al(OH)3

a. Which reactant is limiting, and which is in excess?

b. How many grams of the excess reactant are consumed, and how many grams remain?

c. How many grams of aluminum hydroxide are formed?

10. Sulfur dioxide is mixed with phosphorus pentachloride to produce thionyl chloride and phosphoryl chloride. Suppose 1 kg of SO2 and 1 kg of PCl5 is allowed to react. SO2(l ) + PCl5(l )

SOCl2(l ) + POCl3(l ).

a. Which reactant is limiting?

b. How many kilograms of the excess reactant remain?

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WRITING SCIENCE Name:

Date:

Group:

LOOK

THINK Chemical reactions involve reactants combining to produce products at a specific mole ratio. Stoichiometry is the portion of chemistry dealing with numerical relationships in chemical reactions, where conversion factors such as “per expressions,” or ratios, are used to determine the mass relationships between reactants and products. Mole ratios are used to determine the relationships among moles of various compounds in a reaction, and these ratios can be determined from the coefficients of a balanced chemical equation. The molar masses of reactants and products are used as conversion factors to calculate mass relationships. These calculations can also be used to determine the limiting agent of a chemical reaction and to determine the percent yield of a reaction. Consider the chemical equation for burning hydrogen: 2H2 + O2

2H2O

WRITE Explain the stoichiometric calculations used to determine the mass relationship between the reactants and products of the chemical reaction that occurs when burning hydrogen.

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C3DE Stoichiometry

WRITING SCIENCE

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High School Chemistry

C4ABC

Reaction Rates

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C4ABC Reaction Rates

Student Handout Name:

Date:

Procedure 1.

Use a pipette to place 2 mL of 3% H2O2 and 2 mL of 6% H2O2 in individual test tubes.

Add 10 drops of detergent to each test tube and stir.

Be prepared to measure the time for the foam to reach the top of the test tube after adding the NaI.

Place 10 drops of NaI in the test tube with the 3% solution. Stir briefly, and then measure the time for the bubbles to reach the top of the test tube.

Repeat for the 6% concentration.

Observe and record your qualitative data below.

Create a data table below to organize your quantitative data.

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C4ABC Reaction Rates

Student Journal Name:

Date:

Group:

Part I: Clock Reaction Scientific Investigation Step 1. Question of inquiry: Step 2. Relevance: Step 3. Variables: Independent variable (also known as the manipulated variable): Dependent variable (also known as the responding variable): Control variable(s) or group, also known as constants: Step 4. Hypothesis: Do you need a hypothesis? If so, what is it? How will the responding variable change when the manipulated variable changes?

Step 5. Materials:

Step 6. Safety considerations:

Step 7. Procedure Temperature=

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C4ABC Reaction Rates

Student Journal Step 8. Data collection (Part 1: Effect of changing concentration (at constant temperature) on reaction time): Reaction Time for Five Trials (Part 1) Test Tube Trial

mL Solution A

mL H2O

mL Solution B

mL Total volume Time elapsed

Step 9: Procedure: Step 10: Data collection (Part 2: Effect of changing temperature (at constant concentration) on the reaction rate):

Reaction Time Data Table (Part 2) Test Tube

On Ice

Heated

mL Solution A mL H2O mL Solution B mL Total volume Temperature °C Time elapsed

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C4ABC Reaction Rates

Student Journal Scientific Investigation Step 11. Data analysis (what patterns did you observe in the data?):

Part II: Elephant Toothpaste Observations:

Reactants: _______________________________________________________________ Products: ________________________________________________________________ Catalyst: _________________________________________________________________

Temperature before reaction: __________ Temperature after reaction: ____________

Was there a temperature change? If so, was it hotter or colder?

What was the role of the catalyst in this reaction?

How could the experiment be modified to test the catalyst further?

Draw a heating curve below representing the general energy change of the reaction with the catalyst.

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C4ABC Reaction Rates

Student Journal Reflection and Conclusions 1.

In Part I, does the data support or refute the hypothesis? Why or why not?

What did the results reveal about the relationship between the dependent and independent variables?

What errors could have been made in your experiment?

Write a conclusion statement about the experiment.

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STEMscopedia: REACTION RATES

C4ABC

Reflect In order for a glow stick to glow, you have to start a chemical reaction in it. A glow stick is a plastic tube with a glass vial inside it. When you snap the stick, you break the glass vial. This allows the chemicals that are inside the glass to react with the ones surrounding it in the plastic tube. Light is a form of energy, and many chemical reactions release energy. In the case of the reaction in a glow stick, energy is released as light. Temperature affects the rate of reaction in a glow stick. What would happen to the intensity of the chemical reaction if you put the glow stick in the freezer or hot water? Would temperature changes slow or speed up the chemical reaction? Chemical Kinetics A branch of chemistry called chemical kinetics studies the rate of speed of chemical reactions and the factors affecting that rate. Reactions occur at varying speeds ranging from the slow rusting of iron to the extremely fast decomposition of TNT. Reactions take place when particles collide and chemically interact with a certain amount of energy. The minimum amount of energy needed for the particles to react is called the activation energy, and it is different for each reaction. Rates of Chemical Reaction The speed of reaction is called the rate of reaction.

Rusting at top Slow at bottom

Baking Fast

Explosion Very Fast

The speed of different chemical reactions varies greatly. Some reactions are very fast, and others are very slow.

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STEMscopedia: REACTION RATES The rate of a reaction depends on two things: the frequency of collisions between particles and the energy with which particles collide. If particles collide with less energy than the activation energy, they will not react. The particles will just bounce off of each other. Anything that increases the number of successful collisions between reactant particles will speed up a reaction. Reactions do not proceed at a steady rate. They start off at a certain speed, then get slower and slower until they stop. The concentration of reactants decreases as the reaction progresses. This reduces the frequency of collisions between particles, and so the reaction slows down.

Look Out Reactions do not proceed at a steady rate. When reactants combine, the rate of reaction begins slowly. They start off at a certain speed, then get slower and slower until they stop. The concentration of reactants decreases as the reaction progresses. This reduces the frequency of collisions between particles, and so the reaction slows down. The percentage of the reactant decreases as the percentage of product increases. Factors That Affect Reaction Rate It’s useful to be able to predict whether an action will affect the rate at which a chemical reaction proceeds. Several factors can influence the rate of a chemical reaction. In general, a factor that increases the number of collisions between particles will increase the reaction rate, and a factor that decreases the number of collisions between particles will decrease the chemical reaction rate. •

Nature of Reactants: What the reactants are changes the reaction rate. Many factors affect how a substance will react, such as chemical and physical properties, state, molecular structure, type of bonding, strength and number of chemical bonds, and intermolecular forces.

•

Ionic vs. Covalent Compounds: Ionic compounds in a solution (for example, in water) react faster than covalent compounds. Ionic compounds break apart to form free ions. Covalent compounds that break apart form neutral atoms.

Pb2+(aq) + 2I-(aq)

PbI2(s) is fast at 25oC.

2CO (g) + O2(g)

2CO2(g) is slow at 25oC.

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STEMscopedia: REACTION RATES •

Weaker vs. Stronger Bonds: Reactions involved in breaking weaker bonds are faster than reactions that break stronger bonds.

C2H4 has one double bond: 2C2H6(g) + 7O2(g) C2H4 (g) + 4O2(g) •

4CO2(g) + 6H2O (g) is very, very fast.

2CO2(g) + 4H2O (g) is very fast.

Single vs. Double Bonds: Reactions that break single bonds between atoms are faster than reactions that break double bonds. C2H6 has all single bonds:

2C6H14(g) + 19O2(g)

12CO22O (g) is fast.

2C12H (g) + 37O2(g)

24CO2(g) + 26H2O (g) is slow.

•

Fewer vs. Many Bonds: Reactions involved in breaking fewer bonds are faster than reactions that have to break many bonds:

•

Gases vs. Liquids/Solids: Reactions involving gases are faster than those involving liquids or solids.

•

Same Phase vs. Different Phase: Homogeneous reactions (all reactants in the same phase) tend to be faster than heterogeneous reactions that must occur at an interface. In a heterogeneous reaction, one reactant might be a liquid and another a solid.

•

Concentration of Reactants: A higher concentration of reactants leads to more collisions per unit time, which leads to an increasing reaction rate (except for reactions where concentration has no effect). Similarly, a higher concentration of products created by the reaction tends to be associated with a lower reaction rate. Increasing pressure in gases also increases reaction rates. Pressure is related to concentration and volume. By decreasing the volume available to the molecules of gas, you are increasing the concentration of molecules in a specific space. Changing the pressure of a system works well only for gases. Generally, reaction rates for solids and liquids remain unaffected by increases in pressure.

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STEMscopedia: REACTION RATES •

Surface Area: Making pieces of the reactants smaller increases the surface area. Increasing the surface area of the reactants results in a higher number of reaction sites. Reaction sites are specific sites on molecules at which reactions occur. Increasing the number of reaction sites increases the number of total collisions. The greater the frequency of total collisions, the greater the frequency of effective collisions. If the frequency of effective collisions increases, so too does the reaction rate. Reaction depends on collisions. The more surface area on which collisions can occur, the faster the reaction.

What Do You Think? A large piece of coal takes a long time to ignite, as do piles of grain, yet hundreds have died from coal mine dust and grain silo dust explosions. Which factor affecting the rate of chemical reaction is the cause of these deaths? How can science help us avoid these tragic events? •

Temperature: Usually, an increase in temperature is accompanied by an increase in the reaction rate. Temperature is a measure of the kinetic energy of a system, so higher temperature implies higher Coal average kinetic energy of molecules and more collisions per unit time. A general rule of thumb for most (not all) chemical reactions is that the rate at which the reaction proceeds will approximately double for each 10 degrees Celsius increase in temperature. Once the temperature reaches a certain point, some of the chemical species may be altered (e.g., denaturing of proteins) and the chemical reaction will slow or stop.

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STEMscopedia: REACTION RATES At higher temperatures, a larger fraction of molecules have kinetic energies equal to or greater than the activation energy for the reaction than at lower temperatures. In other words, a greater proportion of molecules have enough kinetic energy to participate in the reaction. By heating a mixture, you will raise the average energy levels of the molecules involved in the reaction. Increasing temperature means the molecules move faster and have more energetic collision and therefore more “effective collisions” that can overcome the activation energy required for a chemical reaction.

•

Presence of Catalysts: Activation energy is the energy needed to put a molecule into a state where it can break and rearrange bonds. Catalysts (e.g., enzymes) lower the activation energy of a chemical reaction and increase the rate of a chemical reaction consumed in the process. Catalysts work by increasing the frequency of collisions between reactants, altering the orientation of reactants so that more collisions are effective, reducing intramolecular bonding within reactant molecules, or donating electron density to the reactants. The presence of a catalyst helps a reaction to proceed more quickly to equilibrium. The presence of a catalyst does the following:

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STEMscopedia: REACTION RATES What Do You Think? We use catalysts to our benefit in many everyday procedures such as soap making, the fermentation of wine to vinegar, and the leavening of bread. However, some catalysts are harmful to our environment, such as chlorofluorocarbons (CFCs), which were used in refrigerants and aerosol propellants. CFC’s catalyze the destruction of the molecule ozone, O3, composed of three oxygen atoms. We need the ozone layer to protect us from harmful ultraviolet radiation from the Sun. At what point do humans draw a line between chemicals we need for society to function and limiting their use due to harmful consequences? What Do You Know? 1. Explain how the following factors affect the rate of chemical reactions: • The nature of the reactants • Concentration of reactants • Surface area • Temperature • Presence of a catalyst 2. Explain the role of total bond energy in chemical reactions. 3. Describe the relationship between particle collisions, activation energy, and the rate of reactions 4. Explain the role of activation energy in chemical reactions.

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STEMscopedia: REACTION RATES Connecting With Your Child Making Cheese To help your child understand how catalysts speed up chemical reactions, try making mozzarella cheese together. Cheese is made from the casein (protein) in milk that has curdled or thickened from a chemical reaction with acids. Rennet introduces an enzyme to milk that works as the catalyst in making cheese, starting the acidification or souring of the milk. The rennet enzymes force the milk to coagulate, or clot. This is referred to as curdling in cheese-making terms. In the curdling process, some protein, fat, and other solids bond together to form curds. When this happens, it causes the water and some other proteins to separate from the curd. The resulting liquid is called whey. The rennet causes the cultured milk to set, producing a soft, ripened curd. Materials: • Gallon of milk (the milk can be pasteurized, but not “ultra-pasteurized” because that process affects the milk’s ability to coagulate or thicken) • Citric acid (available at grocery stores) • Rennet tablets—called Junket (rennet is the enzyme used as catalyst to coagulate milk) • Non-iodized salt (e.g., pickling salt) • Large pot • Meat thermometer • Rubber gloves Procedure 1. Fill two small bowls or cups with 1/4 cup of cool water. To one bowl, add 1 1/2 teaspoon citric acid; to the other, add 1/4 teaspoon of liquid rennet. 2. Pour your milk into the large pot and heat over medium heat until the milk measures 55 degrees Fahrenheit. 3. Add your citric acid mixture and gently stir. Continue to heat milk until it reaches 88 degrees Fahrenheit. It will begin to thicken and separate and won’t look very appetizing. 4. Pour in the rennet mixture and stir gently for 30 seconds. Then leave undisturbed until milk heats to about 100 to 105 degrees Fahrenheit. The curds will separate from the whey and pull away from the edge of the pot. The whey will be a non-cloudy, yellow liquid that resembles Mountain Dew soda. When this happens, turn off the heat. 5. Scoop or pour the curds into a fine mesh sieve and drain as much whey as you can. Reserve the whey in a large bowl or pot. 6. Return the whey to the pot and heat it over medium-high heat to 175 degrees Fahrenheit. While the whey is heating, shape your curds into six balls. 7. Put on the rubber gloves. One at a time put the balls of cheese in a ladle and dip into the hot whey for a few seconds. Pull each ball out and knead it a bit in your hands. Dip it again and knead it some more. Work in 1/4 to 1/2 tsp. of salt. Dip and knead some more. Keep dipping and kneading until the cheese is smooth and pliable and doesn’t break when you stretch it. 8. Keep chilled until you are ready to eat the mozzarella.

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C4ABC Reaction Rates

Reading Science Name:

Date:

Group:

Chemical Reaction Rates 1

A chemical reaction is a process that leads to the transformation of one substance into another. For example, add sodium to chloride and you get sodium chloride, or table salt. This reaction takes time—we call this the reaction rate. Reaction rates, or the speed of a reaction, explains how fast or slow a reaction takes place. For example, the oxidation reaction that occurs in iron when water is present is a slow reaction that can take many years, but the combustion of the cellulose in a tree during a wildfire is a reaction that takes place in a fraction of a second. The speed of these reactions depends on collisions between the particles, and this can be affected through kinetic energy (temperature differences), concentration differences, or pressure differences.

When particles of a reaction hit each other, only a certain percentage of the collisions cause any noticeable or significant chemical change; these successful changes are called successful collisions. The successful collisions have enough energy, also known as activation energy, at the moment of impact to break the preexisting bonds and form all new bonds. This results in the products of the reaction. If we look at the making of table salt, it takes successful collisions between the sodium and the chloride particles to create sodium chloride molecules.

When a catalyst is involved in a collision between the reactant molecules, less energy is required for the chemical change to take place, and hence more collisions have sufficient energy for reaction to occur. The reaction rate therefore increases.

Increasing the concentration of the reactant particles or raising the temperature (kinetic energy), will then bring about higher numbers of collisions and therefore many more successful collisions, thus increasing the rate of the reaction.

Adding pressure to a chemical reaction will also speed up the rate of a reaction, and slow down a reaction if decreased. The speeding up or slowing down of the reaction rate is due to pressing the particles closer together or by allowing the particles to spread out, therefore allowing the particles to be farther apart. The farther apart the particles are, the less likely they are to hit another particle or the slower they will be in hitting another particle.

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C4ABC Reaction Rates

Reading Science Continued 6

Particle size is yet another factor that can affect the reaction rate of a reaction. The smaller the particle size, the faster the rate of reaction because as the surface area is increased, there is more chance for the particles to react. For example, imagine adding 10 grams of table salt to a solution; and then imagine adding a 10 gram cube of the same reactant (salt). The powdered salt should react fully before the single lump of salt. Why? The smaller particles are already dispersed or reduced, whereas the lump must first be dispersed or reduced before reacting.

Lastly, all reactions want to reach a state of equilibrium: the state in which both reactants and products are present in concentrations that have no further tendency to change as time passes. Think about adding a tablespoon of salt to a glass of water, over time the particles of salt will dissolve evenly throughout the water, thus being in equal amounts per ounce of water.

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C4ABC Reaction Rates

Reading Science 1 What is a chemical reaction? A The rate at which a reaction takes place. B The transformation of reactants into products. C The ability of particles to reach equilibrium. D None of the above.

2 What affects collisions between particles in a chemical reaction? A Temperature (kinetic energy) B Pressure C Concentration D All of the above

3 Particle size can have a large effect on the rate of a chemical reaction. Which affects the rate of chemical reactions more with regard to speeding up a reaction? A Smaller particle size B Larger particle size C Both small and large particle size speeds up a reaction. D A catalyst must act upon particle size to affect the rate of a chemical reaction; it cannot affect the reaction alone.

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Reading Science 4 A scientist is adding oxygen to iron in a chemical reaction to create rust. At what point during this reaction will the oxygen and iron quit making rust? A Once the reaction runs out of either oxygen or iron B Once there is enough rust produced C Once enough pressure, temperature, or concentration is reached D Once it reaches equilibrium

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C4ABC Reaction Rates

Math Connections Name:

Date:

Group:

Chemical reactions occur when particles collide with a certain amount of energy. The minimum amount of energy needed for the particles to react is called the activation energy. Some factors that can affect the activation rate of a chemical reaction are concentration of reactants, temperature, surface area of reactants, and concentration of reactants. Use the following information to solve Questions 1–3. The data below display the results from an experiment to see how the surface area of catalyst affects the reaction time. Reaction Time (minutes)

Surface Area of Catalyst (mm2)

1. Graph the data and draw a line of best fit on the graph.

2. Create an equation to describe the relationship. What are the units involved in your equation?

3. How does the surface area affect the reaction time?

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Math Connections Use the following information to help answer Questions 4–6. The chart below displays data from an experiment designed to test the effects of temperature on the reaction rate as Alka-Seltzer® is fully dissolved in water. The rate of the reaction was charted at each increase in temperature. Temp (°C)

Reaction Time (seconds)

Reaction Rate (1/Time) or (s–1)

20°

0.016

30°

0.033

40°

0.066

50°

0.125

55°

0.166

60°

0.25

70°

0.5

4. Graph the data and describe the relationship between the temperature and the reaction time.

5. What effect does the temperature seem to have on the reaction time?

6. What is the relationship between the reaction rate and the reaction time? Why is this?

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C4ABC Reaction Rates

Math Connections Use the following information to answer Questions 7–10. In an experiment designed to test the effects of concentration on rate of reaction, the rate of reaction was calculated using the time it took for a 5-cm piece of magnesium (Mg) ribbon to be dissolved in varying concentrations of (HCl) solution. The results from the experiment are shown in the table below. Trial

Concentration of HCl (cm3)

Time (seconds)

100

136

140

7. Plot the data on the graph provided. 8. How does the shape of the data describe the relationship between the concentration level and the reaction time? 9. Balance the formula representing the chemical reaction in the experiment. Mg+HCl

MgCl2+H2

10. The change in enthalpy can be calculated in terms of the bond energies by using the formula: 𝚫H0=H(products) – H(reactants) If the standard of heat formation in kJ for HCl is -92.8, for MgCl2 is -797.1, for Mg is 0 and for H2 is 0, what is the enthalpy of the reaction? Is the reaction endothermic or exothermic?

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C4ABC Reaction Rates

Writing Science Name:

Date:

Group:

LOOK

THINK Chemical reactions occur at many different rates, ranging from the slow rust of metal to the fast burning of wood. During a laboratory experiment, two students are given a choice of substances to combine and observe. They are familiar with all of the chemical substances and notice that a catalyst is one of them. They discuss the possible addition of the catalyst to their experiment and the expected results. They decide to include the catalyst to see how it affects the reaction. WRITE What is a catalyst, and why is it used in chemical reactions?

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C4ABC Reaction Rates

Writing Science

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High School Chemistry

C4D

Equilibrium

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C4D Equilibrium

Student Journal Name:

Part I: Straw Activity

Date:

Group:

Equilibrium

Transfer Number Water Level A: Products

Water Level B: Reactants

Start 1. 2. 3. 4 5. 6. 1.

What determines if a chemical equilibrium exists?

How did you know if a chemical equilibrium has been reached?

How does this activity simulate chemical equilibrium?

Part II: PhET Simulation

Time

Molecule A

Molecule B

20 s 40 s 60 s 1 min, 20 s 1 min, 40 s 2 min

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C4D Equilibrium

Student Journal A.

Equilibrium Reaction

How does the number of balls changing relate to a reversible reaction?

What happened to the reaction?

Where did the final equilibrium point happen?

What do you think would have happened if we had added molecule B first?

Le Chatelier’s Principle

Add heat to the reaction. Record your observations.

Remove heat from the reaction. Record your observations.

Add more of molecule A. Record your observations.

Add more of molecule B. Record your observations.

Write a definition of Le Chatelier’s principle in your own words based on your observations.

Application

Identify how the following reaction will shift when these stresses are applied to it: H2O (l) + heat ←→H2O (g) Change/stress on system

Direction of shift

Reason for shift

Adding heat Adding water vapor Removing water vapor Decreasing temperature

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STEMscopedia: EQUILIBRIUM C4D

Reflect How does your body get oxygen to all 37 trillion cells? Your blood has a protein in it called hemoglobin that is responsible for binding to and taking oxygen all over the body. Equilibrium must be maintained between hemoglobin and the oxygen it is carrying. The interaction between hemoglobin (Hb) and oxygen is as follows: Hb (aq) + 4O2 (g)

Hb(O2)4 (aq)

Your body can maintain equilibrium as long as there is enough oxygen in the air. A person can cause a shift in this equilibrium by changing altitude.

Hemoglobin maintains an equilibrium with oxygen.

Chemical Equilibrium = No Observable Change on the Molecular Level Chemical equilibrium is the state in which both the reactants and the products in a chemical reaction are present in concentrations that will not change over time on the molecular level. Usually chemical equilibrium is attained when both the forward reaction and the reverse reaction are occurring at the same rate. That is, there is no net change in the concentration of either reactant(s) or product(s). In other words, if we were able to observe the molecules in a chemical reaction, in chemical equilibrium we would see the molecules transform in one direction and then the other. Eventually, we will no longer be able to see the concentrations change. If we were able to look at the reaction, it would appear as if nothing were happening. However, the two reactions are still going on! This is chemical equilibrium. The forward and backward reactions are happening at the same speed and end up canceling each other out. This type of equilibrium is known as dynamic equilibrium.

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STEMscopedia: EQUILIBRIUM Forward and Backward Reactions Use Double Arrows: Equilibrium can be shown in an equation for a reaction by the use of special double right-left arrows. The double arrows do not just indicate that the reaction is reversible. They C+D indicate that the reaction is a reversible reaction in a state of dynamic equilibrium. A + B

Forward Reaction and Backward Reaction: The change from left to right in the equation is known as the forward reaction. The change from right to left is the backward reaction, also known as the reverse reaction. In chemical equilibrium, these two reactions will reach a point where both reactions are happening but it will look like the reaction is finished. However, this is not the case, as some of the molecules are turning into products and some are turning back into reactants. The balanced equation is written with two arrows to indicate that the reaction can proceed in both forward and reverse directions, one pointing from the reactants to the products and the other pointing from the products to the reactants . To avoid confusion, substances on the left side of the chemical equation are called “reactants,” while those on the right side are called “products.” Technically speaking, all chemical reactions are reversible. What are sometimes called irreversible reactions are those reactions that proceed nearly to completion. That is, there is a very high concentration of the products and almost no reactants. In these reactions, the reverse reaction is usually too slow to be observed.

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STEMscopedia: EQUILIBRIUM What Do You Think? Do you think the direction in which a chemical reaction is written matters in equations? H2 + Br2 2HBr

2HBr (Synthesis of hydrogen bromide)

H2 + Br2 (Dissociation of hydrogen bromide)

These two equations are exactly the reverse of each other. The roles of the products and reactants are reversed, so the two chemical equations compliment each other.

Endothermic and Exothermic Occur in Opposite Directions: To develop equilibrium, a chemical reaction must be reversible and not go to completion. Reversible reactions are endothermic in one direction and exothermic in the other. Endothermic (heat absorbed) reactions are favored by increasing the temperature. Exothermic (heat released) reactions are favored by decreasing the temperature.

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STEMscopedia: EQUILIBRIUM Equilibrium happens on its own: Equilibrium is a state attained with no need for outside stimulation. If two substances are put into a mixture, they will combine and react by themselves, and eventually they will reach equilibrium. Scientists call this a spontaneous process; that is, it happens on its own. Chemical equilibrium occurs when the rates of forward and reverse reactions are equal, and the concentrations of products and reactants do not change. Equilibrium results in no charge: A system “at equilibrium” has no observable electric charge. It is neutral. Just as the forward and backward reactions cancel each other out, so do all the pluses and minuses, which gives a total charge of 0. The letter “K” is used to add up all of the actions and conditions in a reaction. “K” is referred to as the equilibrium constant.

Look Out quilibrium can be disturbed: The equilibrium of a chemical system can be E disturbed by a change in concentration, temperature, or pressure. This is explained in Le Chatelier’s principle. The French chemist Henri Le Chatelier came up with a principle for systems in equilibrium. The principle says that if you have a system in equilibrium, and you do anything to it that changes the conditions of the equilibrium, the system will try to adjust to a new state of equilibrium. “A system in equilibrium, when subjected to a stress resulting from a change in temperature, pressure, or concentration, will readjust its chemistry and thereby establish a new state of equilibrium.”

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STEMscopedia: EQUILIBRIUM Changes in concentration, temperature, volume/pressure, and temperature Le Chatelier states that if a system at equilibrium experiences a change, the system will shift its equilibrium to compensate for the change. This doesn’t always mean that there are equal numbers of reactant and product molecules. It all depends on the molecules and conditions of the system. When molecules are left alone, they reach equilibrium. However, that position of equilibrium (the point of equilibrium) can change if something happens to the molecules. • Changing concentration (only with gases or aqueous solutions): If the concentration of a species is lowered or removed, the equilibrium will shift to produce more of that species. On the other hand, if the concentration is increased or added to, the equilibrium will shift to produce less of that species. For example, if we remove some of the Br2, the system shifts toward the left (the 2HBr reverse reaction happens) to produce more Br2.: H2+ Br2

• Changing pressure/volume (only gases): Increasing the pressure of a system causes the system to shift so that the least number of gas molecules. This is because when pressure increases, the volume decreases so there is less space for particles to move around. When the pressure decreases, the system shifts to produce a higher number of gas molecules because there is now a greater volume in the container. Observe the equation below: N2(g) + 3H2(g)

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2NH3(g).

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STEMscopedia: EQUILIBRIUM • Change pressure/volume (continued): If the pressure increases, the system shifts to the right because fewer gas molecules are produced in the forward reaction. Increasing the volume has the same effect as decreasing the pressure. As the volume of a system increases, the molecules have more space to move about, resulting in fewer collisions and a decrease in pressure. Similarly, as the volume of a system decreases, the molecules have less space to move about, resulting in more collisions and an increase in pressure. • Change temperature: A reaction is endothermic if it absorbs heat. A reaction is exothermic if it releases heat. Endothermic (heat-absorbing) reactions are favored by increasing the temperature. Exothermic (heat-releasing) reactions are favored by decreasing the temperature. For example, PCl5(g) + energy, if the temperature increases, the system will in the equation PCl3(g) + Cl2(g) shift to the left, and the reverse reaction will be favored because reaction will use some of the extra energy. In any reactions that can go forward and backward, one direction is endothermic, and the other is exothermic.

• Using a catalyst: A catalyst will increase the speed of a reaction but it has no effect on equilibrium.

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STEMscopedia: EQUILIBRIUM Connecting With Your Child Investigating Chemical Equilibrium To help your child learn more about chemical equilibrium, work together on making these acid-base indicators. In this project your child will learn how to make his or her own indicator that can be used to find out if a solution is acidic or basic (alkaline). What does it mean for a solution to be acidic or alkaline? An acid is a substance that donates hydrogen ions. Because of this, when an acid is dissolved in water, the balance changes between hydrogen ions and hydroxyl ions and the equilibrium shifts. Now there are more hydrogen ions than hydroxyl ions in the solution. This kind of solution is acidic and turns the blue-violet cabbage juice to red. A base is a substance that accepts hydrogen ions. When a base is dissolved in water, the balance between hydrogen ions and hydroxyl ions shifts the opposite way. Because the base “soaks up” hydrogen ions, the result is a solution with more hydroxyl ions than hydrogen ions and the equilibrium shifts. This kind of solution is alkaline and turns the blue-violet cabbage juice to green or yellow. Procedure to make acid-base indicator: 1. Buy a red cabbage. 2. Cut it in thin slices. 3. Put them in a pot and pour in enough water to cover them. 4. Boil for half an hour, then turn off the heat and let the temperature come down. 5. Pour the blue-violet liquid you have obtained into a large, low container 6. The boiled cabbage slices are edible and you can use them in a recipe. 7. Pour 1 centimeter of cabbage juice into a clear plastic container for each test. 8. Now try different kitchen liquids and powders to see if they turn red (acid) or green/yellow (base). 9. Look up your substance on the Internet and see how your findings compare to the substance being an acid or base or neutral.

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C4D Equilibrium

Reading Science Name:

Date:

Group:

Chemical Equilibrium 1

Chemical equilibrium is the state in which both reactants and products are present in concentrations that have no further tendency to change as time passes. Think about adding a tablespoon of sugar to a glass of tea: over time the particles of sugar will dissolve evenly throughout the tea, thus being in equal amounts per ounce of tea.

Reversible reactions result in equilibrium of reactants and products. For a reaction involving two reactants and two products this can be expressed symbolically as:

aA + bB

cC + dD

Where A and B can react to form C and D, or in the reverse reaction, C and D can react to form A and B. 3

If we use our tea with sugar added example, when you add tea (A) and sugar (B) together you make sweet tea (C) and water (D); however, if you were to reverse this you could separate the water/tea from the sugar.

When the concentration of reactants decrease, this slows the rate of the reaction, thus the production of the products. If the concentration of products increases, then the speed of the rate of reverse reactions can also increase. Eventually both concentrations will strive to be constant (or the rates will become equal). Using the tea example, if you decrease the rate of adding sugar to your tea, then the production of sweet tea mixed with water will slow; reversely, if you have a large amount of sweet tea mixed with water, you can separate more particles back into sugar and tea only. Eventually this reaction will stabilize, or reach equilibrium.

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Reading Science 1 What is chemical equilibrium? A Both reactants and products are equal over time B When the reactants produce more products C When the reactants are higher than the products D None of the above

2 Reversible reactions allow products to create the reactants after a chemical reaction has gone from reactants to products. Which of the following is an example of a reversible reaction? A Oxygen combining with hemoglobin in red blood cells in our body, and then the red blood cells delivering the oxygen to needed locations. B Ammonium chloride separating into ammonia and hydrogen chloride gas. C Hydrogen and oxygen combining to form water, and then water being broken down by a cell to be used in cellular respiration. D All of the above are examples of reversible reactions.

3 A scientist is running an investigation that requires him to combine reactants to produce specific products. As he completes his chemical reaction, he is increasing his products. How will this affect the rate of reversible reactions at the end of his investigation? A It will increase the rate of reversible reactions. B It will decrease the rate of reversible reactions. C It will have no effect on the rate of reversible reactions. D None of the above.

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C4D Equilibrium

Math Connections Name:

Date:

Group:

In a chemical reaction, chemical equilibrium is the state in which both reactants and products are present in concentrations with no further tendency to change with time on the molecular level. An equilibrium constant (K) is the ratio of concentrations or pressures of the products to the reactants of a reaction when equilibrium is achieved. Use the following information to answer Questions 1–5. A homogeneous reaction is classified as a reaction in which the states of matter of the reactants and products are both the same. The equilibrium constants for concentration and pressure are calculated differently from that of a heterogeneous reaction. For example, in the following reaction nitrogen and hydrogen combine to form ammonia. Both the reactants and products are gases. N2 (g) + 3H2 (g)

2NH3 (g)

1. The equilibrium constant for concentration, Kc, can be found by using the concentration of each molecule in the following equation: [NH3]2 Kc = [N2][H2]3 If the amounts of molecules in equilibrium in this reaction are H2 = 0.0031M, N2 = 0.085M, and NH3 = 0.031M, what is the equilibrium constant?

2. Using what you learned from the formula above, calculate the equilibrium constant, Kc, for the reaction below assuming the concentrations in equilibrium are NO = 0.062M, H2 = 0.012M, N2 = 0.019M, and H2O = 0.138M. 2NO (g) + 2H2 (g)

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N2 (g) + 2H2O (g)

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Math Connections 3. In the following reaction, ethanoic acid and ethanol react together to form ethyl ethanoate and water. Use the following concentrations under equilibrium to calculate the equilibrium constant, Kc. CH3COOH(aq) + C2H5OH(aq)

CH3COOC2H5(aq) + H2O(l)

Species

Concentration (mol dm−3)

C2H5OH

0.30

CH3COOH

0.13

CH3COOC2H5

0.39

4. When K > 1, the reaction favors the product, and if K < 1, the reaction favors the reactant. In Question 3, which is favored in the reaction?

5. What would the equilibrium constant be if the reaction were reversed?

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Math Connections Use the following information to answer Questions 6–12. The equilibrium concentrations can be calculated for various reactions by using an ICE chart. • “I” stands for the initial concentrations or pressures for each substance in the reaction. • “C” stands for the change in the concentrations for each substance as the system moves towards equilibrium. • “E” stands for the equilibrium concentrations for each substance when the reaction reaches equilibrium. Consider the following problem: A reaction consisting initially of 6.00 moles of NH3, 4.00 moles of N2, and 10.00 moles of H2 in a 10-L beaker was heated to 900 K and given time to reach equilibrium. The equilibrium constant, Kc, is 0.0076. 2NH3 (g)

N2 (g) + 3H2 (g)

6. Step 1 is to convert the initial quantities to molarities by dividing the number of moles of each substance by the volume. Fill in the quantities in the chart below.

NH3

Initial Concentration (M) Change in Concentration (M) Equilibrium Concentration (M)

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Math Connections 7. Step 2 is to calculate the reaction quotient, Q, which can be used to determine the direction a reaction will shift to obtain equilibrium. If K > Q, a reaction will proceed forward; if K < Q, a reaction will proceed in reverse. You can calculate Q using the formula below using the initial concentrations from your chart. Qc=

[N2][H2]3 [NH3]2

8. Compare Qc to Kc to determine which direction the reaction will go.

9. Step 3 is to determine the change in all the substances in terms of x. The ratio of the stoichiometry in the equation is 2:1:3. If you move to the left, you will assign x to the substance with the lowest coefficient and determine the change in the other substances in terms of x. Write the changes in your chart in Question 6.

10. Step 4 is to express the equilibrium concentration in terms of x and the initial amounts you determined in Question 6. For instance, the equilibrium concentration for NH3 is 0.6 + 2x. Write the expressions for the other substances in your chart.

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Math Connections 11. Step 5 is to substitute the expressions in your equation for calculating the equilibrium constant, and solve for x. [N2][H2]3 [NH3]2

= 0.0076

12. Step 6 is to substitute your quantity for x into each of the equilibrium expressions and solve for the equilibrium concentrations.

N2 =

H2 =

NH3 =

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C4D Equilibrium

WRITING SCIENCE Name:

Date:

Group:

LOOK

THINK The seal is perfectly balanced on one flipper while balancing the ball on his nose. A see-saw with two people of the same mass at each end will not go up or down but will stay balanced in the middle. There are many everyday examples of balance in our lives. These types of balancing acts are similar to a chemical reaction that is in equilibrium. WRITE How is a chemical reaction that has reached equilibrium similar to the examples of balance above?

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C4D Equilibrium

WRITING SCIENCE

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High School Chemistry

C5AB

Thermal Energy in Chemical and Physical Processes

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C5AB Thermal Energy in Chemical and Physical Processes

Student Handout Name:

Date:

Calorimetry It is not as difficult as you might think to calculate the heat of a chemical reaction. If you know the heat of the reaction, you can also calculate the specific heat of an unknown substance. Watch the teacher demonstration, and then use this information to calculate the specific heat of the unknown metal used. Fill in the table below to find this value. Mass of the water in the calorimeter (g) ______________________g (Remember that the density of water is 1 g/mL. Use this to calculate the mass of the water.) Starting temperature of the water in the calorimeter (T1)

______________________oC

Ending temperature of the water in the calorimeter (T2)

______________________oC

Change in temperature (ΔT = T2 – T1)

______________________oC

Specific heat of water

4.18 J/(g x oC)

Use this information to calculate how many joules of heat were released by the piece of metal using the process of dimensional analysis. The template below is provided to help you. Use all of the spaces provided. Make sure to keep track of your units and place them in the correct part of the template. q = m • cp • ΔT

Answer = _____________________ J

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C5AB Thermal Energy in Chemical and Physical Processes

Student Journal Name:

Date:

Group:

Background 1.

What occurs when a system is heated or cooled? Include the term heat flow.

The amount of heat that flows between substances and systems can be calculated using the equation Q = m x Cp x ΔT. Describe what each of the terms in this equation means in words.

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C5AB Thermal Energy in Chemical and Physical Processes

Student Journal Part I: Phase Change Energy Flow heating start time______________ ice melting time________________ Time Elapsed (minutes:seconds) 0:00

heating end time_______________ water boiling time_____________

Beaker Temperature (°C)

Time Elapsed (minutes:seconds) 13:00

0:30

3:30

1:00

14:00

1:30

14:30

2:00

15:00

2:30

15:30

3:00

16:00

3:30

16:30

4:00

17:00

4:30

17:30

5:00

18:00

5:30

18:30

6:00

19:00

6:30

19:30

7:00

20:00

7:30

20:30

8:00

21:00

8:30

21:30

9:00

22:00

9:30

22:30

10:00

23:00

10:30

23:30

11:00

24:00

11:30

24:30

12:00

25:00

12:30

25:30

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Beaker Temperature (°C)

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C5AB Thermal Energy in Chemical and Physical Processes

Student Journal Part II: Heating Curve

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Student Journal 1.

What was happening to the water when the lines on the graph were mostly flat?

What was happening to the water when the lines on the graph were mostly sloped?

Was the energy flow into the beaker constant? How did you know? Can you make a useful heating curve if the energy flow is not constant? Why or why not?

Based on your heating curve, determine the melting and boiling points of water. Does this match the accepted values (00C for melting and 1000C for boiling)? What sources of error are there in the experiment that may keep you from having the perfect melting and boiling points? Provide three or more.

Based on your heating curve, explain how adding heat allows phase changes to occur on the molecular level.

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C5AB Thermal Energy in Chemical and Physical Processes

Student Journal Name:

Date:

Group:

Background 1.

What happens to heat energy from the beginning of a reaction to the end of the reaction?

What is the difference between temperature and heat?

Why is it important to measure and record the amount of reactants used during an investigation of chemical reactions?

How can we measure the amount of heat released or absorbed during a chemical reaction?

Data Collection Compound

Mass of Calcium Chloride or Citric Acid

Mass of Soduim Bicarbonate

Mass of Water

Initial Temp. °C

Final Temp. °C

calcium chloride/ sodium bicarbonate citric acid/ sodium bicarbonate Data Analysis Compound

Total Mass of Reactants

∆ T=Tf - Ti

q = m x Cp x ∆T in J

Heat in J/g of Reactants

Endothermic or Exothermic

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Student Journal Graph of reaction with calcium chloride and sodium bicarbonate.

Graph of reaction with citric acid and sodium bicarbonate.

Post-Investigation Discussion 1. How would you describe the reactions in terms of energy flow and systems?

How would you describe the reactions in terms of bond energy?

How could you illustrate the release or absorption of energy from a chemical reaction?

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STEMscopedia: THERMAL ENERGY IN

CHEMICAL AND PHYSICAL PROCESSES C5AB

Reflect Most first aid kits contain a small pouch called a cold pack. Cold packs are not cold at first, but they can be activated to become cold. They are useful when someone gets injured and needs an instant ice pack to keep the injured area from swelling. When you activate a cold pack, you set in motion a chemical reaction that causes a drop in temperature. Other chemical reactions cause temperature to increase. Chemists deal with the temperature changes that accompany chemical reactions when they study thermochemistry. When activated, a cold pack What is thermochemistry, and what does it tell us about chemical feels as cold as a bag of ice reactions? water. All chemical reactions involve a change in energy. Thermochemistry is the study of energy changes that occur in chemical systems. All chemical reactions involve some change in energy as the atoms in reactants recombine to form new products. Depending on the reaction, the energy change may be very small or it may be very large. Each reaction is different, and each reaction can be defined by a specific energy change associated with it. To analyze this more closely, consider a set of reactants before a reaction occurs. This set of reactants can be considered a system. Everything else is considered the surroundings. If the reaction is initiated between the reactants, a change in energy of the system occurs. For any given reaction, the energy of the system (Esystem) either increases or decreases as reactants react to form products. If the energy in the system decreases, energy is released into the surroundings. If the energy in the system increases, energy is absorbed from the surroundings.

Reactions that release more energy than they require are called exothermic reactions. Those that absorb more energy than they release are called endothermic reactions. Heat Lost or Heat Gained We can calculate the amount of heat lost or gained by substances in chemical reactions. Heat is also called enthalpy (Q or ∆H) and is measured using the mass of the substance, specific heat of the substance (energy required to raise one gram of the substance one degree Celsius), and the change in initial vs. final temperature. Heat is measured in joules (SI) or calories.

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CHEMICAL AND PHYSICAL PROCESSES Energy and Reactions All chemical reactions involve a change in energy, as the atoms in reactants recombine to form new products. Energy change can be very large (combustion reactions) or very small, depending on the reaction. It starts with the collision of the reactants. Then there are two other factors that must be satisfied before it is determined whether the reactants can indeed react. The reactants must collide with the proper orientation and with the minimum amount of energy required to start the reaction. This is called the activation energy. Without a collision that has both the right orientation and the minimum activation energy, the reactants will simply bounce off each other and move on. This energy is needed to break the bonds that are in the reactants and is usually in the form of heat or light. Bond Energy Since the law of conservation says that energy before and after the reaction must be the same, we can account for the total energy by looking at the bond energies of each substance in the reaction. We can subtract the sum of the energies of the reactants from the sum of the energies of the reactants. The enthalpy (∆Hrxn) of the reaction can be calculated this way. The sign for this helps us determine whether the reaction is endothermic or exothermic. Endothermic Reactions If the products have greater energy than the reactants, the reaction is endothermic. In an endothermic reaction, the enthalpy (∆H) of the products is greater than the enthalpy of the reactants, showing that energy was absorbed (gained) and ∆H is positive. Endothermic reactions absorb heat, making them feel cold. Generally, endothermic reactions take place more slowly than exothermic reactions because their activation energies are higher. In an endothermic reaction, heat flows into the system from the surroundings. Therefore, ∆H is positive from the energy absorbed during the reaction. Examples include the instant cold packs that mix water and ammonium nitrate to cool an area of the body by absorbing heat, and the photosynthetic process of producing sugar.

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CHEMICAL AND PHYSICAL PROCESSES Exothermic Reactions If the products have lower energy than the reactants, the reaction is said to be exothermic. Exothermic reactions give off heat, causing the things that were formed to have less energy than they did before they reacted. In an exothermic reaction, the enthalpy of the products is less than the enthalpy of the reactants, meaning that energy was released (lost) during the reaction and Q or ∆H is negative. Examples include cooking bread, burning a candle, and crystallizing liquid salts (as in sodium acetate in chemical hand warmers).

What Do You Think? If you look on the package of just about any food item, you will find a nutrition label. This label lists the quantities of fats, proteins, carbohydrates, and various minerals and vitamins present in one serving of the food. This information lets you know about the chemical makeup of the food so that you are aware of the types of nutrients the food will supply to your body. The nutrition label goes beyond listing the chemical content of the food, however. It also states the number of calories present in a serving of the food item. The number of calories lets you know how much energy your body will take in if you eat a serving of the food. This number is obtained by food chemicals using lab techniques in calorimetry. What is calorimetry, and how does it enable scientists to measure the energy content of foods? Measuring Energy Changes in Chemical Reactions Calorimetry is a topic within the larger field of chemistry. The word calorimetry comes from the Latin word calor, which means “heat,” and from the Greek word metron, which means “measure.” Basically, calorimetry is the science of measuring changes in energy that accompany a chemical reaction. Scientists use the term enthalpy change (ΔH) to talk about the energy differences between the reactants and products of a chemical reaction. When reactions occur under constant pressure conditions, all of this energy is in the form of heat. ΔH is often expressed in either joules (J) or calories (cal). One calorie is equivalent to 4.184 joules. A food chemist can determine the enthalpy change for a reaction experimentally using a device called a calorimeter. A calorimeter is an insulated container that the chemist can seal after placing reactants inside. The chemist then inserts a thermometer into the calorimeter so that it extends down into the reaction chamber. This allows scientists to monitor the temperature change that occurs when they initiate a reaction between the reactants inside the reaction chamber.

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CHEMICAL AND PHYSICAL PROCESSES Look Out Heat and temperature are not the same thing. Heat is a form of energy; it is measured in calories or joules. Temperature is a measure of the average kinetic energy of the particles in a substance; it is measured in kelvins or degrees. As you will see in the next section, temperature changes measured during a chemical reaction inside a calorimeter cannot be used directly to measure the enthalpy change (ΔH) of the reaction. Instead, the temperature change must be inserted into a mathematical equation along with other factors to calculate the enthalpy change of the reaction. Determining Changes in Enthalpy So far, we have talked about how a calorimeter can measure changes in temperature. How do scientists measure changes in the enthalpy of a reaction? Before we answer this question, let’s review some terms. First, if pressure remains constant during a reaction, change in enthalpy (ΔH) of a chemical system is equivalent to the heat energy in the system (q): ΔHsystem = q Second, ΔH of a chemical system is equal to the difference between the sum of the enthalpies of products and reactants: ΔHsystem = ∑Hproducts − ∑Hreactants = q Now that we have reviewed the meanings of these terms, we can turn to another equation that relates q to other factors that scientists can measure in a lab: q = m × c × ΔT In this equation, m is the mass of the system (measured in grams), c is the specific heat of the system (measured in J/goC), and ΔT (measured in oC) is the change in temperature of the system. This equation allows a scientist to measure the temperature change of a chemical reaction in a calorimeter, and then use that temperature change to calculate the change in enthalpy for that reaction. The scientist must also measure the system’s mass (m) and know the system’s specific heat (c). For a chemical reaction, the system is composed of the substances undergoing reaction and the solvent. For reactions taking place in water, the mass is the sum of the water and other reactants. Only the specific heat of water is used because it is present in such high amounts compared to the other substances. We haven’t yet used specific heat in any calculations, so let’s take a minute to define this term and learn how to use it. The specific heat (c) of a substance is defined as the amount of heat required to raise the temperature of one gram of the substance by one degree Celsius. For a particular substance, specific heat is a constant in the same way that boiling point and melting point are constants.

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CHEMICAL AND PHYSICAL PROCESSES The table below provides specific heat values for several common substances. Note that specific heat is measured in units of joules per gram per degree Celsius (J/g·°C). Substance

Water

Wood

Glass

Copper

Specific Heat (J/goC)

4.184

1.76

0.84

0.385

What Do You Think? Specific heat is a property of a substance. Each substance has a unique specific heat. To think about this, consider what it feels like to stand on a sandy beach on a hot summer day. When it is very hot, the sand becomes scorching—so hot that it hurts to walk barefoot across the beach to the water. Yet, when you step into the water, it feels cool. You might wonder why the sand is so hot and the water is so cool when both are exposed to the same conditions of air temperature and sunlight. How can the concept of specific heat solve this puzzle? (Here is a hint: The specific heats of water and sand are not the same. In the table above, you can use the specific heat of glass—which is made from sand—to approximate the specific heat of sand. How do the two specific heat values differ? Why does this explain your different experiences with hot sand and cool water?) Measuring the Enthalpy Change for a Reaction The procedure for determining ΔH, or change in enthalpy, of a chemical system begins with laboratory work and then proceeds to calculations as outlined below: 1. Measure the masses of reactants: The masses of the reactants will be important later on during the calculation phase of the procedure. 2. Measure the temperature of reactants: Reactants should be allowed to achieve the same initial temperature. Measure and record this value as the initial temperature, Tinitial. 3. Combine reactants in the calorimeter and initiate reaction: Some reactants react spontaneously when combined. Others need to be ignited to begin a reaction. 4. Monitor the temperature until the reaction is complete: The temperature will rise if the reaction is exothermic; the temperature will fall if the reaction is endothermic. Record the temperature every 30 seconds until it levels off. This indicates the reaction has come to a stop. The temperature reading corresponding to this time point can be taken as the ending temperature of the reaction, Tfinal. 5. Calculate ΔT: Subtract Tinitial from Tfinal to determine ΔT: ΔT= Tfinal − Tinitial.

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CHEMICAL AND PHYSICAL PROCESSES 6. Calculate q: Use the following equation to calculate the enthalpy change for the reaction (q). Sum the measured masses of the reactants to obtain the mass of the reaction solution (m). Use ΔT (calculated above). The value for specific heat (c) can be obtained from reference sources. For aqueous solutions, use the specific heat of water (4.184 J/g·°C). q = m × c × ΔT 7. Adjust the sign to indicate either an exothermic or endothermic reaction: If the reaction is exothermic, include a negative sign before the final result. If the reaction is endothermic, include a positive sign. Let’s look at a specific example to see how the steps above work. Suppose you have 100 g 1M HCl and 100 g 1M NaOH. The initial temperature of both solutions is 22.5°C. You combine the solutions in a calorimeter and observe that the temperature increases as an acid-base reaction occurs: HCl(aq) + NaOH(aq) → H2O(l) + NaCl(aq) The final temperature of the solution is 29.2°C. What is the enthalpy change of this reaction, assuming the specific heat of the reaction solution is 4.184 J/g·°C? Solution: To determine q, calculate m and ΔT, and then multiply these values by c. m = 100.0 g HCl(aq) + 100.0 g NaOH(aq) = 200.0 g ΔT = Tfinal − Tinitial = 29.2°C – 22.5°C = 6.7°C q = m × c × ΔT = (200.0 g) (4.184 J/g·°C) (6.7°C) = 5.6 × 103 J = 5.6 kJ Because temperature increased during the reaction, this reaction must be exothermic. Therefore, place a negative sign in front of the ΔH value to give an enthalpy change of −5.6 kJ for this reaction.

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CHEMICAL AND PHYSICAL PROCESSES What Do You Know? A chemist ran three calorimetry experiments. The table below shows data collected during the experiments. Complete the missing parts of the table. Experiment Number

Specific Heat of a System (J/goC)

4.184

3.921

2.553

Mass of System (g)

10.5

5.9

7.4

Temperature of Reactants (oC)

22.1

20.0

25.3

Temperature of Products (oC)

34.8

11.3

20.7

Type of Reaction Δ H (kJ)

Look Out Heating or cooling a substance can cause it to undergo a phase change. As heat is added to a substance, the kinetic energy increases. As kinetic energy increases, motion increases, eventually breaking intermolecular attractions and causing a state of matter change. When a substance is going through a phase change, the temperature is held constant. For instance, the melting point of salt (sodium chloride) is 801 degrees Celsius. We can take a piece of solid NaCl and heat it to its melting point. The temperature will stay the same the entire time as long as there is still solid. This is because the energy is being used to break the intermolecular attractions instead of raising the temperature. Once these intermolecular attractions are broken and solid NaCl has turned into molten NaCl, the temperature can increase again. Heating Curves A heating curve can be plotted out as temperature vs. time. During a heating curve, kinetic energy and particle motion increase as temperature increases. This shows a slope when temperature is rising and a plateau when the substance is in a phase change indicating that the temperature is constant. The phase changes that these heating curves display are melting points and boiling or vaporization points.

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STEMscopedia: THERMAL ENERGY IN

CHEMICAL AND PHYSICAL PROCESSES Connecting With Your Child Your child can calculate enthalpy and draw a heating curve during a common physical process (melting ice into liquid water followed by heating the water to produce steam). 1. Take the initial weight of a small pot. 2. Fill a small pot two-thirds full with ice and add water to cover the ice. 3. Take the final weight of the small pot. Subtract the weights to get the mass of the ice water. Record this mass. 4. Place a thermometer into the pot and allow the thermometer to sit below the water line. Do not let the sensor/thermometer rest on the bottom or side of the pot. 5. Carefully swirl the pot until the temperature stabilizes at or below zero degrees Celsius. Record the temperature at 30-second intervals. 6. After two minutes (record the time), place the pot on the stove top and add turn on the heat. Warm the ice water and continue to record the temperature at 30-second intervals. NOTE: Once the stove top is set, leave it at this setting and make no adjustments to the heat. 7. Record the time at which all the ice has melted. Also record the time at which the water begins to boil. 8. Continue to heat the water for at least three minutes after the water begins to boil, recording the temperature at 30-second intervals. After three minutes, turn off the stove top and record the time and the final temperature. Your child can now construct a graph (time on the x-axis and temperature on the y-axis) of a heating curve using the data they just collected. Guided questions to ask your students: 1. Was this experiment an endothermic or exothermic reaction? 2.

Explain, in terms of bond energy, what is happening in this reaction.

What do the sloped portions of the graph represent?

What do the plateaus on the graph represent?

Explain, in terms of intermolecular forces, what is happening during a phase change.

Calculate the heat of the reaction from 2 degrees Celsius to 60 degrees Celsius (specific heat of water is 4.184 J/goC and use the mass recorded of the ice water at the beginning of the experiment).

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C5AB Thermal Energy in Chemical and Physical Processes

Reading Science Name:

Date:

Group:

Enthalpy of Reaction 1.

Thermochemical equations may be written to express the amount of energy either required or produced during a chemical reaction. The energy term is usually expressed with the units of kilojoules (kJ) and is a representation of the enthalpy of the reaction. The change in enthalpy that occurs during the reaction is represented by the term ΔH. The symbol Δ stands for the term delta and means “change in.” The delta symbol can be attached to any unit and represents the final value minus the initial value. For chemical equations, this means that the value of H of the reactants will be subtracted from the value of H of the products.

The ΔH value may be written directly into the chemical equation, or it may be written after the equation with either a positive or negative sign. Either way, by adding the energy term, a thermochemical equation is created. If the energy term is on the reactant side, it indicates that the reaction required energy. This means that the reaction absorbed energy, so the ΔH value after the equation is positive. This represents an endothermic reaction. If the energy term is on the product side of the reaction, it indicates that the reaction produced energy. This means that the reaction released energy, and the ΔH value is negative. This represents an exothermic reaction.

There are times when the energy unit is known and is provided within the thermochemical equation. However, there are times when the energy term is either not known or not provided. The amount of energy that is either absorbed or produced during a chemical reaction, or the enthalpy of reaction (ΔH), may be calculated by using the individual ΔHf values for each compound. ΔHf stands for the enthalpy of formation. The symbol Σ stands for “the sum of.” This is a fairly straightforward and simple process that any chemistry student can calculate. You start by finding the ΔHf values of each compound involved in the chemical reaction. You then place these values on the appropriate side of the chemical equation. Finally, you subtract the sum of the ΔHf values of the reactants from the sum of the ΔHf values of the products. The formula looks like this: ΔH = Σ (moles of ΔHf) products – Σ (moles of ΔHf) reactants

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C5AB Thermal Energy in Chemical and Physical Processes

Reading Science 4.

You will be provided with the ΔHf values for each equation. These values are state dependent, meaning that liquid water and steam will have different values. You will need to multiply the ΔHf value for each compound by the number of moles of that compound in the balanced chemical equation, as each value of ΔHf is for 1 mole of compound. Find the enthalpy of reaction for the following equation: Equation: 2 C2H6 (g) + 7 O2 (g) 4 CO2 (g) + 6 H2O (l) ΔH = ? ΔHfC2H6 (g) = - 84.7 kJ/mol ΔHf O2 (g) = 0 kJ/mol ΔHf CO2 (g) = - 393.5 kJ/mol ΔHf H2O (g) = -285.8 kJ/mol

When you are learning how to perform this calculation for the first time, it may be helpful to write the equation so that you can visualize the products minus the reactants, as shown below. Products

Reactants

ΔH = 4 CO2 (g) + 6 H2O (g) (minus)

2 C2H6 (g) + 7 O2 (g)

ΔH = (4 (-393.5 kJ/mol) + 6 (-285.8 kJ/mol)) - (2 (-84.7 kJ/mol) + 7 (0)) When you multiply the number of moles by kJ/mol, the mole unit cancels. ΔH = ((-1,574.0 kJ) + (-1,714.8 kJ)) - (-169.4 kJ) ΔH = - 3,119.4 kJ Therefore, 3,119.4 kJ of energy is released in this reaction, producing an exothermic reaction.

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C5AB Thermal Energy in Chemical and Physical Processes

Reading Science 1.

The passage discusses two different types of thermochemical reactions. In one type of reaction, the enthalpy term is located on the reactant side of the equation. What type of reaction is this? A.

Chemical reaction

Endothermic reaction

Thermochemical reaction

Exothermic reaction

Many calculations in chemistry will involve certain symbols, such as the symbol Δ, or delta. What does this symbol stand for or mean? A.

The total

The change in

The unit

The formula

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C5AB Thermal Energy in Chemical and Physical Processes

Reading Science 3.

Calculate the enthalpy of reaction for the following equation. Use the table of the reaction of elements per mole of compound. Remember, the values are state dependent. 2 H2O2 (l)

2 H2O (l) + O2 (g)

Element or Compound

ΔHf in kJ/mol

O2 (g)

0 kJ/mol

H2O (l)

- 285.8 kJ/mol

H2O (g)

- 241.8 kJ/mol

H2O2 (l)

- 187.6 kJ/mol

PCl5 (g)

- 398.9 kJ/mol

H3PO4 (aq)

- 1,277 kJ/mol

N2 (g)

0 kJ/mol

HCl (g)

- 386.2 kJ/mol

HCl (aq)

- 700.0 kJ/mol

Use the space provided to show all of your work. A.

- 196.4 kJ

97.9 kJ

- 473.4 kJ

196.4 kJ

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C5AB Thermal Energy in Chemical and Physical Processes

Reading Science 4.

Calculate the enthalpy of reaction for the following equation. Use the table of the reaction of elements per mole of compound. Remember, the values are state dependent. 2 H2O2 (l)

2 H2O (l) + O2 (g)

Element or Compound

ΔHf in kJ/mol

O2 (g)

0 kJ/mol

H2O (l)

- 285.8 kJ/mol

H2O (g)

- 241.8 kJ/mol

H2O2 (l)

- 187.6 kJ/mol

PCl5 (g)

- 398.9 kJ/mol

H3PO4 (aq)

- 1,277 kJ/mol

N2 (g)

0 kJ/mol

HCl (g)

- 386.2 kJ/mol

HCl (aq)

- 700.0 kJ/mol

Use the space provided to show all of your work. A.

- 196.4 kJ

97.9 kJ

- 473.4 kJ

196.4 kJ

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C5AB Thermal Energy in Chemical and Physical Processes

Reading Science 5.

When calculating the ΔH of a reaction, why is it important that you include the correct state symbols in your balanced chemical equation? A.

So that you know which compounds are involved

So that you have the correct molar ratios

So that you use the correct energy value

The state symbols do not matter when calculating ΔH.

A thermochemical equation indicates the absorption or release of heat in addition to the reactants and products involved in the chemical reaction. The amount of energy available for conversion into heat is represented by a ΔH. What does this term stand for? A.

The enthalpy of the reaction

The kinetic energy of the reaction

The potential energy of the reaction

Not enough information is given to answer the question.

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C5AB Thermal Energy in Chemical and Physical Processes

Math Connections Name:

Date:

Group:

Thermochemistry is the study of energy changes that occur in chemical systems. All chemical reactions involve some change in energy as the atoms in reactants recombine to form new products. Chemical reactions not only generate new compounds, but can also transform or transfer energy. Chemical energy can be transformed into thermal energy, and vice versa. A thermochemical equation is a type of balanced chemical equation that includes the amount of energy absorbed or released during a chemical reaction. The total thermal energy of the system is known as the enthalpy (H) of the system. The heat of reaction, or enthalpy change (∆H), is equal to sum of the change in enthalpy of the products minus the sum of the change in enthalpy of the reactants. ∆H = Σ∆Hθ(products) – Σ∆Hθ(reactants) In an exothermic reaction, heat flows from the system to the surroundings. Therefore, the change in enthalpy (ΔH) is negative. In an endothermic reaction, heat flows into the system from the surroundings. Therefore, ΔH is positive. 1.

Look at the following reaction between sodium hydroxide and hydrochloric acid. NaOH + HCl

H2O + NaCl + 57.9 kJ/mol

Is this an endothermic or exothermic reaction? 2.

Look at the following reaction when ammonium nitrate is dissolved in water. NH4NO3 + H2O + 650.9 kJ/mol

NH4 + NO3

Is this an endothermic or exothermic reaction? 3.

The combustion reaction of propane described by the equation below has an enthalpy change of -2044kJ. What will the enthalpy of the reaction be of 0.25 mol of propane with 1.25 mol oxygen? (The enthalpy change is proportional to the the amount of reactants.) C3H8 + 5O2

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3CO2 + 4H2O 387

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C5AB Thermal Energy in Chemical and Physical Processes

Math Connections Use the following table to help calculate the change in enthalpy for the reactions in questions 4-10. Compound

∆HθkJ/mol

Compound

∆HθkJ/mol

CO2(g)

-394

PbO(s)

-218

H2O(l)

-286

NO2(g)

+34

CH4(g)

-74

NaHCO3(s)

-948

ZnCl2(s)

-416

Na2CO3(s)

-1131

Pb(No3)2(s)

-449

NH4Cl(s)

-315

Nh+4(aq)

-133

Cl-

-167

**The ∆Hf for all elements including the seven diatomic elements (H2, N2, O2, F2, Cl2, Br2, I2) is zero. 4.

The combustion of methane produces carbon dioxide and water. Calculate the ΔH for the equation and classify the reaction as either endothermic or exothermic. CH4(g) + 2O2(g)

CO2(g) +2H2O(l)

When heated above 200oC, sodium bicarbonate decomposes quickly to form sodium carbonate, water, and carbon dioxide. Calculate the ΔH for the reaction and classify the reaction as either endothermic or exothermic. 2NaHCO3(s)

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Na2CO3(s) + H2O(l) + CO2(g)

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C5AB Thermal Energy in Chemical and Physical Processes

Math Connections 6.

Zinc chloride is used to form a colored compound called “Ruhemann’s Purple”, which is used to detect fingerprints under fluorescent light. Write a balanced equation to show zinc and chlorine reacting to form zinc chloride. Calculate the ΔH for the reaction and classify the reaction as either endothermic or exothermic.

Write the balanced equation and calculate the enthalpy change for the reaction showing lead (II) nitrate decomposing to lead (II) oxide, nitrogen dioxide, and oxygen. Classify the reaction as eithehr endothermic or exothermic.

Ammonium chloride dissolves in water, forming its component ions. Calculate the ΔH for the reaction and classify the reaction as either endothermic or exothermic.

NH4Cl(s)

NH4+(aq) + Cl-(aq)

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C5AB Thermal Energy in Chemical and Physical Processes

Writing Science Name:

Date:

Group:

LOOK

THINK A calorimeter is used to experimentally measure thermal energy of a chemical reaction. The equation is used to calculate the heat transferred. Calorimetry is used by many types of scientists in the study of heat transfer. Knowing the quantity of the transfer of heat is essential in all areas, including chemical manufacturing and food production. The basic concept is the same for any calorimeter: a known amount of chemical is reacted, and the temperature change of the surroundings is measured. When placed into a formula including specific heat, the amount of heat transferred can be calculated. WRITE Give an explanation for each letter and symbol found in the equation q=m·c·∆T. Be sure to • clearly state your central idea; • organize your thoughts; • develop your essay in detail; • choose your words carefully; and • use correct spelling, capitalization, punctuation, and grammar.

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C5AB Thermal Energy in Chemical and Physical Processes

Writing Science

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High School Chemistry

C5C

Gas Laws

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C5C Gas Laws

Student Handout Name:

Date:

Charles’s Law - Increased Temperature Record your observations of the teacher’s demonstration using hot water. State your reasons (conclusions) for this reaction and use Charles’s law to explain (justify) what occurred.

Observations Record your direct observations. Include the starting and ending circumferences of the balloon.

Conclusions Describe what was occurring inside the balloon on the particulate level. Specifically, state how temperature affected the gas inside the balloon.

Justifications Write the equation for Charles’s law in the space below. How does this equation justify your conclusions?

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C5C Gas Laws

Student Handout Charles’s Law - Decreased Temperature Record your observations of the teacher’s demonstration using cold water. State your reasons (conclusions) for this reaction and use Charles’ law to explain (justify) what occurred.

Observations Record your direct observations. Include the starting and ending circumferences of the balloon.

Conclusions Describe what was occurring inside the balloon on the particulate level. Specifically, state how temperature affected the gas inside the balloon.

Justifications Write the equation for Charles’s law in the space below. How does this equation justify your conclusions?

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C5C Gas Laws

Student Journal Name:

Date:

Group:

Background 1.

What are the main assumptions of kinetic molecular theory? What type of gases does this theory work for?

What is STP? What are the values that are used? Include the correct units that would be used in your answer.

What is the equation for Charles’ law? ____________________________________________

What is the equation for Boyle’s law? _____________________________________________

How may these two laws be combined? What do they create? State the name and the equation for this law.

A 2.5-L closed container holds an ideal gas that is at STP. The container is heated to 25oF. What is the new pressure in the container after heating? Show your work.

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C5C Gas Laws

Student Journal Background, continued 7.

What is the equation for Avogadro’s law? ________________________________________

If 0.6795 mol of methane gas occupy 15.229 L at STP, then what is the volume for 1 mol of methane? Show all your calculations.

What is the equation for the ideal gas law? What does it describe?

10. What does the value of R represent in the ideal gas equation? What two main values (with units) may be used? Be sure to state which pressure unit goes with each value.

11. How many moles of a gas are contained in a 5-L container at 15oF and at 1.2 atmospheres? First, state which equation you will use and why you will use it. Then solve in the space provided below, showing all of your work.

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C5C Gas Laws

Student Journal Part I: Plan Your Investigation 1.

My question of inquiry:

My hypothesis:

My prediction:

What is the independent variable (also known as the manipulated variable)?

What is the dependent variable (also known as the responding variable)?

Is there a control group or control variable for this investigation? Explain.

What materials, equipment, and technology will be needed for this investigation?

List all safety precautions that must be taken.

Follow the procedures listed in the Student Guide to conduct this investigation.

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C5C Gas Laws

Student Journal Part II: Implement Your Investigation Collect, Record, and Organize Data Procedure I: Volume and Temperature Record your data and calculations in the table below. Balloon #1

Balloon #2

Air Temperature (°C) Initial Conditions

Circumference (cm) Radius (cm) Volume (cm3) Air Temperature (°C)

Final Conditions

Circumference (cm) Radius (cm) Volume (cm3)

Graph. Draw a line graph of volume versus temperature for the gas sample. Title the graph and label the axes with the appropriate units.

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C5C Gas Laws

Student Journal Part II: Implement Your Investigation, continued Collect, Record, and Organize Data Procedure II: Volume and Pressure Record your data and calculations in the table below. Fill in the first three columns of the table below with your volume information, then graph your findings. Number of Largest Volume of Number of Moles Balloon Antacid Tablets Circumference Balloon (cm3) of Gas Balloon #1

Balloon #2

Balloon #3

Volume of Balloon (cm3)

Create a bar graph comparing the number of antacid tablets with the volume of each balloon.

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C5C Gas Laws

Student Journal Part II: Implement Your Investigation, continued Analyze Data 1.

Based on the data and graph from Procedure I, describe the relationship between gas volume and temperature.

Write an equation to express the linear relationship between volume and temperature of a quantity of gas. Use the symbols V and T to represent volume and temperature.

Explain in detail why the volume of the balloon increased when it was placed in the hot water bath. Focus your answer on the particles inside the balloon and why the increase in volume eventually stopped.

Explain the observation of the reduced volume of the balloon at low temperature. Use the concepts of molecular velocity and the collision of particles according to the kinetic molecular theory in your answer.

Water vapor may have become part of the gaseous mixture within your balloons. What happens to the water vapor as it is cooled? Do you think the same would happen for all gases when cooled? Explain what would occur and why.

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C5C Gas Laws

Student Journal Part II: Implement Your Investigation, continued Analyze Data 6.

Based on the data and graph from Procedure II, describe the relationship between the pressure and the volume of a gas.

Think about opening a bottle of soda. Describe in terms of pressure and volume what happens to the gas bubbles.

What did you notice about the volume of the balloons in relation to the number of tablets used? What type of proportion was this? How do you know?

Calculate the number of moles of carbon dioxide gas produced in each balloon. You will have three equations. Make sure to label each equation for each balloon. Use the ideal gas equation for this calculation. Hint: First solve for n, then plug in your numbers. Show your work in the space provided below, then add these values to the table on page 5 of your Student Journal.

10. How did the number of moles of gas in each balloon relate to the volume of gas produced on your graph? Be as specific as possible.

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C5C Gas Laws

Student Journal Reflections and Conclusions 1.

Does the data support or refute the hypothesis? Explain.

Was your prediction correct or incorrect? Explain

How did the results reveal a relationship between the independent and the dependent variable?

Where could errors have been made while collecting or organizing the data?

What do you conclude about this investigation?

What would you do differently if you were to conduct this experiment again?

Using all of the following terms, develop a graphic organizer based on what you have learned in this Explore. Use another sheet of paper if needed.

Terms: gas, pressure, volume, temperature, moles, ideal gas constant, Boyle’s law, Charles’ law, Avogadro’s law, ideal gas law, ideal gas equation 404

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STEMscopedia: GAS LAWS

C5C

Reflect Imagine rafting with a group of friends along a winding river on a hot summer day. Occasionally you hit some exciting rapids, but mostly you cruise along at a leisurely pace. At noon, the group decides to take a lunch break. Everyone jumps out and helps to pull the raft up on shore. It’s very hot, and so your friends find a shady spot a few yards away to eat lunch. Later, when you return to the raft, you are dismayed to find that the raft burst while sitting in the sunlight. What caused that to happen? The volume of a gas varies directly with temperature. The raft burst because the volume of air inside increased as it sat in the hot sunlight. Much earlier that day, when it was cooler, the raft had been inflated with air. As the day continued, conditions got hotter. The temperature of the air inside the raft increased quite a bit, reaching a very high temperature when it was left in the sunlight at noon. The volume of the air increased in response to this temperature increase. However, because the canvas walls of the raft were not flexible, the expanding air exerted enough pressure to cause a hole to develop in a weak spot of the canvas. The rafters have discovered an important gas law. As a gas’s temperature increases, the gas’s volume also increases. This is true of all gases, not just air. The graph at the right shows data collected for a gas as its temperature increased. Notice there is a direct, linear relationship between volume and temperature. In other words, as one variable increases, the other variable increases at the same rate. All other variables must be constant in order to observe this relationship.

Rafters must take care not to leave their raft in the hot sunlight.

As the temperature of a gas increases, its volume increases when all other conditions are kept constant.

Mathematically, we can write the following equation to show this direct, linear relationship: V = kT In this equation, V is volume, T is temperature, and k is a constant representing all other variables (such as pressure). This mathematical relationship can be rearranged and expanded to show how a gas’s volume changes when its temperature changes:

In this equation, V1 and T1 are the starting volume and temperature, and V2 and T2 are the ending volume and temperature. This mathematical relationship is known as Charles’ law in honor of the French scientist, Jacques Charles, whose experimental work supported its discovery.

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STEMscopedia: GAS LAWS What Do You Think? Suppose you inflated a balloon with air at room temperature, and then you placed it in a freezer. What do you think you would observe about the balloon if you took it out of the freezer 30 minutes later? Does this fit the relationship expressed by Charles’ law? Explain. The volume of a gas varies inversely with pressure. Volume and temperature are not the only two properties of a gas that vary in a predictable way. The volume and pressure of a gas also show a mathematical relationship. To think about this, consider how it would feel if you tried squeezing a balloon to decrease its volume. It’s difficult to do because you have to exert a great deal of pressure to force the balloon to assume a smaller size. From this thought experiment, you should be able to say that volume decreases as pressure is increased. This means that volume and pressure are inversely related.

The pressure and volume of a gas vary inversely and in a nonlinear fashion. Pressure is measured in units called atmospheres (atm).

If you actually performed an experiment to measure the change in volume as you increased the pressure, you would generate data similar to that shown in the graph above. Notice that the line drawn through the data points shows a downward trend, as you would expect from an inverse relationship. However, the data fit a curved rather than a straight line. All other variables must be constant in order to observe this relationship. Mathematically, the data can be expressed according to the following equation:

In this equation, V is volume, P is pressure, and k is a constant representing all other variables (such as temperature). English scientist Robert Boyle discovered this relationship and published his discovery in 1662. In his honor, the relationship between pressure and volume of a gas is now widely known as Boyle’s law. Mathematically, Boyle’s law can be written to show how the pressure and volume of the same gas taken at two different times are related: V1P1 = V2P2 In this equation, V1 and P1 are the starting volume and pressure, and V2 and P2 are the ending volume and pressure. Temperature and all other conditions must remain constant for this relationship to hold. Try using Boyle’s law to calculate the pressure of a gas that expands to a volume of 20 liters (L), if the gas’s initial pressure is 2.5 atmospheres and initial volume is 5 liters. You can check your answer following the What Do You Know? section of this companion.

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STEMscopedia: GAS LAWS The pressure of a gas varies directly with temperature. When the volume of the container that is holding a gas is constant, the pressure and temperature of the gas have a direct relationship. Consider an aerosol can. There are warning signs on the side of the container that say cans should not be placed in fire. This is because the fixed volume container is already under a certain amount of pressure at room temperature. As the temperature is increased if placed in a fire, the pressure in the container will rise as well because the gas particles are gaining more kinetic energy and hitting the side of the container at a greater force and frequency. The container can only handle so much pressure before the container bursts at the weakest point. This means that pressure and temperature are directly related. Mathematically, we can write the following equation to show this direct, linear relationship: P = kT In this equation, P is pressure, T is temperature, and k is a constant representing all other variables (such as pressure). This mathematical relationship can be rearranged and expanded to show how a gas’s pressure changes when its temperature changes: In this equation, P1 and T1 are the starting pressure and temperature, and P2 and T2 are the ending pressure and temperature. This mathematical relationship is known as Gay-Lussac’s law in honor of French chemist, Joseph Louis Gay-Lussac. The volume of a gas varies directly with moles. The number of gas particles also affects gas volume. Chemists use units called moles to measure the number of particles in a substance. One mole of any gas contains 6.02 × 1023 particles of that gas. If you double the quantity of moles of a gas sample, the gas’s volume also doubles, as long as the pressure and temperature are kept constant. The Italian scientist Amedeo Avogadro discovered this relationship in the early 1800s. It can be expressed as the following mathematical equation: V = kn

Avogadro’s law states that volume is directly proportional to moles. If the number of moles of a gas In this equation, V is volume, n is the number of moles of gas, and k is a doubles, the gas’s volume constant. Other variables must be constant for this relationship to hold. likewise doubles, provided Known as Avogadro’s law, this expression can be rewritten in a way that pressure and temperature allows scientists to calculate a gas’s change in volume as the number of remain constant. moles changes:

In this equation, V1 and n1 are the starting volume and number of moles, and V2 and n2 are the ending volume and number of moles. Temperature and pressure are constant.

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STEMscopedia: GAS LAWS The ideal gas law combines all the variables in Charles’ law, Boyle’s law, Gay-Lussac’s law, and Avogadro’s law. So far in our discussion of the behavior of gases, we have dealt with pairs of variables including volume and temperature, volume and pressure, pressure and temperature, and volume and moles. In every case, we specified that only the particular pair of variables could change. All other variables had to remain constant in order to observe the behavior of interest. It is interesting to consider whether there is a broader way to describe gas behavior without such restrictions. In other words, is there a way to describe the behavior of a gas that takes into account all the variables that could change? The answer is yes. The ideal gas law is a mathematical description that accounts for all of these properties of a gas simultaneously: volume, pressure, temperature, and moles. This law was developed by incorporating the relationships described in the four individual gas laws previously discussed. The ideal gas law is written mathematically as follows: PV = nRT In this equation, P is pressure, V is volume, n is number of moles, T is temperature, and R is a constant called the gas constant. The units used for the gas constant depend on the units used for pressure and volume of the gas being described. However, whenever this equation is used, temperature must be expressed in units of kelvin and volume must be expressed in liters. The ideal gas law can be used whenever a gas is considered close to being ideal. To be ideal, a gas must have the following characteristics:

Look Out 1. There are no attractive forces between gas particles. 2. The collisions between gas particles are perfectly elastic so that no energy is lost during collisions. In reality, no gas meets these criteria, although a gas approaches the conditions of ideality at high temperatures and low pressures. Under these conditions, the gas particles are spread very far apart and move very rapidly so that any attractive forces that do exist between the particles are minimized.

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STEMscopedia: GAS LAWS Kinetic molecular theory provides a model for explaining gas behavior. The gas laws are important descriptions of gas behavior, and each represents a significant achievement in scientific history. However, each of the gas laws is limited to describing a natural phenomenon without explaining why that phenomenon occurs. The kinetic molecular theory provides that explanation. Kinetic molecular theory (abbreviated KMT) was developed over many years and with contributions from many scientists. It is a general theory that describes matter in the gas state, though it also applies to some aspects of liquids and solids. The following are postulates, or statements that describe the theory. Notice that the postulates describe gases that are ideal: • Matter is composed of particles: These particles are in constant motion, move in a straight line, and change direction only when they collide with one another or with the walls of a container. • Particles are extremely small: The particles of a gas are much smaller than the distance between the particles. In fact, gas particles are so small relative to the distances between them that particle size is considered negligible; that is, it has no effect on the gas’s behavior. • No attractive forces exist between particles in a gas: In a gas, the particles are not attracted to one another, nor are they attracted to the walls of the container. • No energy is lost when gas particles collide: The collisions between particles of gas are perfectly elastic; that is, they always retain the same kinetic energy they had before collision. • Particles have an average kinetic energy that varies with temperature: As temperature increases, the average kinetic energy of the particles increases. This causes gas particles to collide with the walls of the container with greater frequency and with greater force.

According to the kinetic molecular theory, Boyle’s law can be explained by the increased frequency of collisions that occur when volume is decreased. Although not explicitly defined by any of these postulates, the pressure exerted by a gas can be explained using KMT. Pressure results from collisions of gas particles with the sides of their container. The greater the frequency and force of these collisions, the higher the pressure of the gas. As stated earlier, we can also use these postulates to explain each of the gas laws. For example, Boyle’s law states that as volume decreases, pressure increases. If we use the model laid out by KMT, we can describe how the decrease in volume increases the frequency of particle collisions with the walls of the container. Because pressure is defined by the frequency and force of these collisions, we can say a gas’s pressure increases as its volume decreases. We have a model for understanding why we observe the pressure increase predicted by Boyle’s law.

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STEMscopedia: GAS LAWS Everyday Life: Less oxygen is present at high altitudes. People who climb very high mountain peaks take air tanks with them, just as scuba divers take air tanks as they dive underwater. This may seem puzzling. After all, air is present on mountain peaks. Why would mountain climbers need air tanks? Though air is present on a mountaintop, the gas particles making up that air are fewer in number. The mixture of gases remains the same because just the density of particles changes. Because of this, less oxygen is available at high altitudes. This can make it difficult to get enough oxygen without assistance. Therefore, mountain climbers often rely on air tanks during the highest range of their ascents. What Do You Know? Apply what you know about the gas laws to complete the following table. Read the description of each example scenario, and then name the gas law that applies to that scenario. Write an equation to describe the relationship between variables described by each gas law. You may also use words to describe the relationship: Example Scenario

Name of Gas

Equation (or Word Description of

Law

Relationship Between Variables)

An empty soda can is heated and then placed into a cold water bath. The can implodes. As a boy blows up a balloon, its volume increases. A student calculates the moles of a sample of gas in a balloon after she measures the balloon’s volume, air pressure, and air temperature. A balloon pops after landing on a radiator. A large sample of air is compressed into a small volume using high pressure.

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STEMscopedia: GAS LAWS What Do You Think? So what is the ending pressure of the gas? Plug the numbers into Boyle’s law and solve for P2:

As the gas’s volume increases, its pressure decreases to 0.625 atmospheres.

Connecting With Your Child Connecting With Your Child: Constructing and Testing a Solar Balloon To help your child learn more about Charles’ law, research hot air balloons to learn how they operate. (A flame heats air inside the balloon, causing the air’s volume to increase. As the air expands, it lifts the balloon off the ground.) Then have your child research solar balloons to learn how radiant energy from the Sun can be used as a source of energy for heating air molecules inside a balloon to achieve lift. Finally, have your child choose a design they find in their research to construct a solar balloon out of black plastic trash bags. Your child can test his or her balloon by attempting to launch it on a sunny day. Here are some questions to discuss with your child: 1. What allows a traditional hot air balloon to achieve lift? 2. What gas law applies to explain the behavior of a solar balloon? 3. How does a solar balloon achieve lift, and how is it similar to and different from a traditional hot air balloon?

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C5C GasLaws

Reading Science Name:

Date:

Group:

The Ideal Gas Law in Real Life 1

Gas laws run the chemistry of our modern world. However, these laws were not developed overnight. First, understand that gases simply do not behave like other states of matter when temperature, pressure, or volume changes. Over the last three centuries, four famous chemists developed gas laws that drive modern chemistry calculations involving one of those factors kept as constant.

Their work also led to the formation of one quintessential equation called the ideal gas law. Such diverse inventions, such as vehicle air bags and combustion engines, work or fail dramatically due to the ideal gas law.

To begin, compression does not change the finite volume of solids or liquids. However, because gases have indefinite shape and volume, they are compressible so that the volume of a gas in a closed container does change when pressure changes. The first to quantify this relationship was Robert Boyle in 1662. He stated that pressure and volume of a gas have an inverse relationship. In Boyle’s Gas Law, if temperature is constant, as pressure increases, gas volume decreases, and vice versa: pV = constant, where p is pressure and V is volume.

In 1787, Jacques Charles derived his Charles’s gas law (also known as the law of volume): V = kc T, where V is volume, K is a constant, and T is temperature in Kelvin. Charles showed that if pressure is kept constant, when the temperature of a gas increases, its volume will increase. Conversely, when it cools, its volume decreases. So with pressure constant, volume and temperature (in Kelvin) are directly related.

In the third law, developed by chemist Joseph Louis Gay-Lussac, volume is the factor kept constant: if you increase the temperature of a container with a fixed volume, then the pressure inside the container will increase. More precisely, the pressure of a gas has a direct relationship to its temperature (in Kelvin).

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C5C GasLaws

Reading Science 6

Amedeo Avogadro formulated the fourth law in 1811, which shows the connection between gas volume and another factor called “moles.” A mole is the unit that shows the amount of reactants and products in a chemical reaction. One mole equals 6.022 x 1023 particles of that substance. Avogadro’s law states that equal volumes of gases at the same temperature and pressure contain the same number of molecules (moles). The ratio between volume and the number of moles is a constant.Enter the ideal gas law. In 1834, Emile Clapeyron combined all the previous gas laws into one simple but profound law that states pV = nRT Where: p is the pressure of the gas V is the volume of the gas n is the amount of substance of gas (also known as number of moles) R is the ideal universal gas constant (8.31441 J K-1 mol-1 combines Avogadro’s constant and Bozeman’s constant) T is the temperature of the gas

If you know three of the variables (pressure, volume, moles, or temperature), you can find the missing variable. In other words, if you have a balloon filled with a gas, the mole or the amount of gas molecules is fixed. The gas exerts a pressure on the inside of the balloon and expands to a certain volume at a certain temperature. If you increase the temperature (the right side of the equation), then the other side increases (the product of pressure and volume). For example, if you increase the temperature, the volume must increase. If you increase the pressure, then the volume decreases. Both sides of the equation must be equal, so each factor compensates for the change in the other factor.

Two real-life applications owe their success to the ideal gas law. First, airbags are based on the combination of physics and chemistry. During a severe impact, a chemical reaction is kicked off that releases nitrogen gas that adds many moles of gas to the container. In order to balance the ideal gas law equation, volume must increase by a huge amount. This causes the airbag to inflate in milliseconds and then instantly deflate to absorb the impact of your head and upper body during a collision. Combustion engines are the second example. A piston pushes down on a mixture of gas and air which increases pressure. A spark plug ignites the mixture, causing an explosion that increases the air volume. This makes the piston go back down into the cylinder, which in turn opens another valve releasing exhaust gases. The piston moves a crankshaft that turns the wheels down the road. The relationships between pressure, temperature, and volume in the ideal gas law make the engine pistons turn a crankshaft, which makes the car move.

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C5C GasLaws

Reading Science 1.

Which of the following statements is true regarding the properties of gases? A.

Gases have definite volume, indefinite shape, and are compressible.

Gases have indefinite volume, definite shape, and are compressible.

Gases have definite volume, definite shape, and are not compressible.

Gases have indefinite volume, indefinite shape, and are compressible.

Which factors does Boyle’s gas law describe as having an inverse relationship? A.

Pressure and volume

Temperature and volume

Pressure and temperature

Mass and volume

The pressure of a gas has a direct relationship to its temperature (in Kelvin) is quantified in __________. A.

Boyle’s gas law

Charles’s gas law

Gay-Lussac’s gas law

Avogadro’s law

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C5C GasLaws

Reading Science 4.

The equation V = kc T describes the following behavior of gases __________. A.

When pressure is constant, volume and temperature (in Kelvin) are inversely related.

When pressure is constant, volume and temperature (in Kelvin) are directly related.

When volume is constant, pressure and temperature (in Kelvin) are directly related

When pressure is constant, mass and temperature (in Kelvin) are directly related.

Why is the ideal gas law most important? A.

It combines the universal law of gravitation with the other gas laws.

If a combustion engine piston compresses fuel and air, the volume of the mixture will increase.

If an airbag deploys, the temperature of the gas will decrease.

If three of the four gas variables (temperature, pressure, volume, or moles) are known, the missing factor can be determined.

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C5C Gas Laws

Math Connections Name:

Date:

Group:

Gas Laws Gases have various chemical and physical properties that help us predict their behavior. Gases can be described in terms of their pressure (P), temperature (T), volume (V), and number of moles (n). The relationships between these characteristics are defined using the gas laws. Part I: Boyle’s Law Boyle’s law describes the relationship between volume and pressure. It states that the volume of a fixed amount of gas held at a constant temperature varies inversely with its pressure. PV = k, where P is pressure, V is volume, and k is a constant 1.

Suppose you have a gas in a container being pressed down by a piston. The initial volume of the gas is 5.0 m3 and the pressure is 7.5 mmHg. If the volume decreases to 4.0 m3, what is the pressure?

Fill in the following chart with the pressures as the volume changes. Volume (cubic meters) 5 4 3 2 1

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Pressure (mmHg) 7.5

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C5C Gas Laws

Math Connections Part I: Boyle’s Law, continued 3.

Plot the data from the chart on page 1 on the graph below.

Pressure (mmHg)

Pressure vs. Volume

Volume (m3)

What is the shape of the data? Is the relationship proportional?

What is the relationship between the volume and the pressure of the gas in the container?

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C5C Gas Laws

Math Connections Part I: Boyle’s Law, continued 6.

Now, let’s look at the data for the relationship between P = 1/V. Use the data from question 2 to calculate the inverse volume. Fill in the chart below with the data for the inverse of pressure. 1/Volume (1/m3) 1/5=0.2

Pressure (mmHg) 7.5

Plot the data above on the graph below.

Pressure (mmHg)

Boyle’s Law

1/Volume (1/m3) 8.

What is the shape of the data? Is the relationship proportional?

Write an equation of the line shown in the graph. What is the slope, y-intercept, and constant in the equation? How does this relate to the formula P=1/V?

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C5C Gas Laws

Math Connections Part II: Charles’ Law Charles’ Law describes the relationship between volume (V) and temperature (T). It states that the volume of a fixed amount of gas held at a constant pressure varies directly with its absolute temperature. When the number of moles (n) and pressure (P) is held constant, the ratio of volume to temperature is held constant. V/T = k at constant n and P 10. Suppose you have a gas in a container being pressed down by a piston. The initial volume of the gas is 0.25 L and the temperature is 120 K. If the volume increases to 0.50 L (while pressure and moles stay the same), what is the temperature?

11. Fill in the following chart with the temperatures as the volume changes. Volume (liters) 0.25 0.50 0.75 1.00 1.25

Temperature (Kelvin) 120

12. Plot the data above on the graph below.

Temperature (Kelvin)

Charles’ Law

Volume (Liters) 13. Write an equation of the line shown in the graph. What is the slope, y-intercept, and constant in the equation? How does this relate to the formula T = kV?

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C5C Gas Laws

Math Connections Part III: The Ideal Gas Law Avogadro’s law states that the amount of gas varies directly with the volume. Boyle’s, Charles’, and Avogadro’s laws can be combined into a single statement called the ideal gas law, which describes how the volume of gas is affected by changes in pressure, temperature, and quantity. Ideal Gas Law: PV = nRT R = gas constant = 0.082058 (L atm) / (K mol) – same for all gases STP or Standard Temperature and Pressure = 0°C at 1 atm Use this information to answer the question below. 14. The average lung capacity of a 14-year-old is about 3.6 L. How many moles of gas (air) do the lungs of an average 14-year-old hold?

Assume that the person is at 1.00 atm pressure and has a normal body temperature of 37°C.

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C5C Gas Laws

WRITING SCIENCE Name:

Date:

Group:

LOOK

THINK The gas laws were discovered by Boyle, Charles, Gay-Lussac, and Avogadro. These laws were developed to explain the relationships between temperature, pressure, volume, and the number of moles in gases. Newer vehicles are equipped with tire pressure gauges that have been installed to alert drivers when more air is needed in their tires. The weather, which can change fairly quickly during the day in some cities, can affect tire pressure. Susan has noticed that in the mornings her tire pressure gauge light is on. However, when she makes her commute home after work, the light goes off. WRITE Explain why the tire pressure gauge is turning on and off depending on the time of day Susan is driving her vehicle, and explain the implications for airing up her tires properly in the future. Then, use a model to show the relationships between temperature, pressure, volume, and the number of moles in relation to the gas laws. Be sure to • clearly state your central idea; • organize your thoughts; • develop your essay in detail; • choose your words carefully; and • use correct spelling, capitalization, punctuation, and grammar. © Accelerate Learning Inc. - All Rights Reserved

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C5C Gas Laws

WRITING SCIENCE

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High School Chemistry

C6ABE

Solutes and the Process of Dissolving

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C6ABE Solutes and the Process of Dissolving

Student Handout Name:

Dissolution Rates

Date:

Sugar–Water Dissolution

Beaker 1 Type of Sugar:

Beaker 2 Type of Sugar:

Beaker 3 Type of Sugar:

Beaker 4 Type of Sugar:

________________

Prediction:

______________

Time to dissolve:

_______________

Observations:

________________

Describe the difference between the substances in Beaker 1 and Beaker 2, and then describe the difference between the substances in Beaker 3 and Beaker 4.

How did these differences affect the rate of dissolution of these substances?

What factors increased the rate of dissolution? Why?

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C6ABE Solutes and the Process of Dissolving

Student Journal Name:

Date:

Group:

Background 1.

Explain what an aqueous solution is and what it contains.

What is solubility? How is it usually expressed for solutions?

Explain how the dissolving process occurs in solutions until equilibrium is reached. Be as specific as possible.

What factors affect the rate of dissolution?

What factors affect the rate of crystallization?

How do temperature and concentration relate to the saturation of a solution?

Define the term ion. State how this term relates to electrolytic solutions.

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C6ABE Solutes and the Process of Dissolving

Student Journal Part I: Plan Your Investigation 1.

My question of inquiry:

What do you need to do to answer this question?

What are the variables that you will observe?

What do you predict will occur?

What materials, equipment, and technology will you need for this investigation?

What safety precautions must be taken in this investigation?

Record your procedure on the back of this worksheet or in your lab journal if needed.

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C6ABE Solutes and the Process of Dissolving

Student Journal Part II: Implement Your Investigation Collect, Record, and Organize Data Compound

Data

_____°C (Initial)

25°C

40°C

60°C

80°C

Solubility (g) Conductivity Solubility (g) Conductivity Solubility (g) Conductivity Solubility (g) Conductivity Solubility (g) Conductivity Solubility (g) Conductivity

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C6ABE Solutes and the Process of Dissolving

Student Journal Part II: Implement Your Investigation, continued Analyze Data 1.

Use this page to create your solubility graph from the data table. Highlight your assigned compound in red and then graph the results from the rest of the class in various colors. Label each line with the corresponding compound and create a legend for your graph. Remember, the independent variable should be on the x-axis and the dependent variable on the y-axis.

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C6ABE Solutes and the Process of Dissolving

Student Journal Part II: Implement Your Investigation, continued Analyze Data 1.

Why did the different compounds have different solubilities? What factors may have affected this?

Describe the processes that occur in the electrolytic solutions that allow for the movement of an electric current. In other words, what is occurring at the particulate level?

How did the electrolytic and nonelectrolytic solutions differ as you heated them? Explain this behavior.

How did you determine whether your solution was an electrolyte or a nonelectrolyte?

Describe what occurred in the supersaturated solution demonstration. Why did the solution behave that way when the crystal was added?

How did you know when your solutions became saturated? What occurred in the beaker?

Explain why ions need to be present in solution for an electric current to occur and if the ion concentration in solution affected conductivity.

How did the solubility graph that you created relate to the temperature at which the various solutions become saturated? What does this tell you about the saturation levels of each compound?

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C6ABE Solutes and the Process of Dissolving

Student Journal Reflections and Conclusions 1.

Was your prediction correct or incorrect? Explain.

Was there a relationship between the variables you observed?

Where could errors have been made while collecting or organizing data?

What do you conclude about this investigation?

What would you do differently if you were to conduct this experiment again?

Using all of the following terms, develop a graphic organizer. Use additional paper if needed.

Terms: solute, solvent, solution, ionic compound, molecular compound, ions, positive charge, negative charge, neutral charge, conductivity, electrolyte, nonelectrolyte, agitation, dissolution rate, surface area, crystallization, solubility, temperature, saturated, unsaturated, supersaturated

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STEMscopedia: SOLUTES AND THE PROCESS OF DISSOLVING C6ABE

Reflect Imagine you are making lemonade by stirring lemonade mix into a glass of water. One spoonful of mix would likely dissolve easily in the water, producing a solution. (A solution is a mixture of two or more substances spread evenly together.) What would happen if you added another spoonful of mix to the solution? Depending on how much water is in the glass, the additional mix may take longer to dissolve. If you add another spoonful or two, you would reach the point where the lemonade mix no longer dissolves in the water. The lemonade mix that does not dissolve settles to the bottom of the glass. Why is there a limit to how much lemonade mix dissolves in water? Solvation vs. Dissociation In the process of solvation, the solvent (the substance another substance dissolves in) attracts the solute (the substance being dissolved) particles toward itself. This is due to the forces of intermolecular attractions between the solute and solvent particles.

Dissociation is the process of an ionic compound breaking apart into its individual ions. Then the solvent that is being used can solvate and surround each ion.

When sodium chloride (NaCl) is placed in water, the water molecules break apart the individual ions. Each positively charged sodium ion (Na+) is surrounded by the negatively charged ends of water molecules. Each negatively charged chloride ion (Cl –) is surrounded by the positively charged ends of water molecules. Water is one of the most abundant liquids on Earth. It is also one of the best solvents. Water dissolves many different types of substances. Solutions in which water is used as the solvent are known as aqueous solutions.

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STEMscopedia: SOLUTES AND THE PROCESS OF DISSOLVING

Why is water such an effective solvent? An important reason is that water molecules are polar. In other words, they have opposite, partial charges. A water molecule (H2O) consists of two hydrogen atoms (H), each of which shares its electron with a central oxygen atom (O). The shared electrons are attracted more to the oxygen atom, so it acquires a partial negative charge. Meanwhile, the hydrogen atoms acquire partial positive charges. What does this have to do with solubility? When other polar molecules A water molecule is polar. or ionic compounds are mixed with water, their positively charged components are attracted to the oxygen atoms in water molecules, and The oxygen atom (red) has a partial negative charge their negatively charged components are attracted to the hydrogen (δ−), and each hydrogen atoms in water molecules. (Ions are particles with electric charges.) atom (white) has a partial The chemical bonds joining some compounds are strong enough to positive charge (δ+). withstand the attractive forces from water molecules. Other compounds, however, are broken apart. This is what happens when a substance dissolves. In contrast, nonpolar substances such as vegetable oil do not dissolve in water. Instead, the denser substance sinks beneath the less dense substance.

Look Out Only some ionic compounds are soluble in water. Not all ionic compounds dissolve in water with the same ease as sodium chloride. In fact, some ionic compounds do not dissolve at all in water—they are insoluble. Other ionic compounds are only slightly soluble in water. The dissolution behavior of any particular ionic compound can be thought of as a competition. The competition is based on relative strengths of attraction. The ions making up an ionic compound are attracted to one another and also to water molecules. Whether or not a compound dissolves in water depends on which of these forces of attraction is greater.

Lead iodide is insoluble in water and forms a distinctive yellow precipitate.

In some cases, the ions are more strongly attracted to one another than they are to water molecules. In these cases, the ionic compound forms a precipitate in water. You can see the precipitate clearly as a separate substance in the water. In other cases, the ions are more strongly attracted to water molecules, and the compound dissolves in water.

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STEMscopedia: SOLUTES AND THE PROCESS OF DISSOLVING

By running many experiments, chemists have developed a general list summarizing the water solubility behavior of many compounds. The following is a short list of examples of compounds that are soluble/insoluble in water: Soluble Compound

Formula

Exceptions

all nitrates

NO3-

none

all acetates

C2H3O2-

none

all perchlorates

ClO4 –

none

all chlorates

ClO3 –

none

most chlorides

Cl –

AgCl, Hg2Cl2, PbCl2

most bromides

Br –

AgBr, Hg2Br2, PbBr2

Insoluble Compound

Formula

Exceptions

most sulfides

BaS, SrS, CaS, (NH4)2S

most carbonates

CO32-

(NH4) 2CO3

most phosphates

PO43-

(NH4) 3PO4

Factors That Affect Solubility Rates The speed at which a solute dissolves in a solvent—or the rate of dissolution—depends on three factors: the temperature of the solution, the agitation (or mixing) of the solution, and the surface area of the solute. If any of these three factors change, the rate of dissolution changes. • Surface Area: When the surface area of a solute increases, the dissolution rate also increases. Suppose you had a spoonful of sugar granules and the same amount of sugar pressed into a cube. The spoonful of sugar has a greater total surface area than the cube, because each granule in the spoonful contacts the solvent separately. So the spoonful of sugar will dissolve at a faster rate than the sugar cube.

•

Solution Temperature: Heating a solution causes the particles in the solution to move more quickly. As a result, the solvent and solute particles collide more frequently and with greater force. Therefore, when the temperature of a solution increases, the rate of dissolution increases. Lemonade mix dissolves more quickly in hot water than in cold water.

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STEMscopedia: SOLUTES AND THE PROCESS OF DISSOLVING

•

Agitation: Agitating a solution increases dissolution rate. If you stir lemonade mix into water, the mix will dissolve more quickly. A washing machine spins as it cleans your clothes to help the detergent dissolve more quickly in the water.

Colligative Properties Colligative properties are properties of solutions effected only by the number of particles in the solution, not the type of particles. Examples of colligative properties are vapor pressure, freezing point, and boiling point. The normal freezing point and boiling point of a solution can be changed when you add more solute particles. Adding solute will increase the boiling point and lower the freezing point of a solution. This is because the solute particles cause a stabilization of the solvent in the liquid phase of matter. Ionic solutes have a greater effect on the colligative properties of a solution than covalent. This is because ionic solutes will dissociate or break apart into ions when they are placed into a solution, whereas molecular solutes will stay intact. For example, if we compare dissolved aluminum chloride, an ionic compound, to dissolved glucose, a molecular compound, we see the following:

We see that one unit of AlCl3 breaks into four particles, increasing the property by that many times. One unit of glucose will stay one unit of glucose, when put into a solution. A real-life application of using colligative properties to our benefit is that we salt the roads when winter weather gets bad and there is a chance of the roads getting icy. Salting the roads causes the freezing temperature of ice to drop from zero degrees Celsius to around negative nine degrees Celsius. This is helpful because it has to be colder than normal outside for water to freeze and stay frozen, and therefore less of it will form on the roads. Workers will typically add a salt like magnesium chloride instead of sodium chloride because when each of these compounds break apart, magnesium chloride breaks into three ions, whereas sodium chloride only breaks into two. Therefore, magnesium chloride will have a greater affect and lower the freezing point of water more.

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STEMscopedia: SOLUTES AND THE PROCESS OF DISSOLVING

Getting Technical: Solubility and Wastewater Treatment Wastewater is water that has been used and discarded. Wastewater is generated in homes, schools, businesses, and factories. The majority of contaminants present in wastewater must be removed before the water can be released back into the environment. Wastewater treatment plants have been designed to carry out this task. These plants generally use multiple steps to remove various types of contaminants. One removal technique involves the addition of ions that form insoluble precipitates with metal ions. This precipitation method is particularly important in the processing of wastewater coming from industrial sites. Adding hydroxide ions (OH–) to heavy metals such as copper, nickel, lead, chromium, and cadmium can precipitate the metals. Once precipitated, a filtration process can collect these metal ions.

Sedimentation tanks at a wastewater treatment facility allow precipitates to be separated from waste liquid.

What Do You Know? Part I: Precipitates Use the solubility chart to determine if the following compounds will dissolve in water or not. 1. Ba(NO3)2 2. PbCl2 3. Ag2S 4. (NH4)2CO3 5. Fe(C2H3O2)2 Part II: Colligative Effects Which of the following will have a greater effect on the boiling point of a solution? 1. C6H12O6 or KCl 2. CaBr2 or LiBr 3. AlF3 or C12H22O11

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STEMscopedia: SOLUTES AND THE PROCESS OF DISSOLVING

Connecting With Your Child Experimenting With Solubility To help your child learn more about the factors that affect solubility and the rate of dissolution, gather the following materials: • 6 Sugar cubes • 2 Glasses • Measuring cup • Spoon • Water (both warm and cold) • Thermometer • Stopwatch Your child will observe what happens in three different scenarios and compare the results of each set of experiments. First, your child should pour one cup of warm water into one glass and a cup of cold water into a second glass. After recording the water’s temperature in each glass, add one sugar cube to each glass at the same time and start the stopwatch. Your child should record which sugar cube dissolved faster. Second, pour one cup of water at the same temperature into two separate glasses. Crush one of the sugar cubes into smaller pieces. Then add the crushed pieces to one glass of water and a complete sugar cube to the other glass at the same time and start the stopwatch. Your child should record which sugar sample dissolved faster. Finally, pour one cup of water at the same temperature into two separate glasses. Then add one sugar cube to each glass at the same time and start the stopwatch. Your child should stir the solution in one glass gently with the spoon, but leave the solution in the other glass unstirred. Record which sugar cube dissolved faster. Here are some questions to discuss with your child: 1. In each separate example, which sugar dissolved faster? Explain why. 2. How is the rate of dissolution different than the solubility of a substance? What was being tested in this experiment? How could you test the other term? 3. Combining all of the data, how could you minimize the dissolving time of a sugar cube? Explain.

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C6ABE Solutes and the Process of Dissolving

Reading Science Name:

Date:

Group:

Solubility and Nutrients 1

The human body is an open system with a constant exchange of nutrients and energy between the surroundings and the body. Food is eaten to provide the energy that is stored in the molecular structure of the food molecules. Food is first broken down in the stomach and then passes through the intestinal tract where it is either absorbed or emitted by the body. If the minerals and nutrients are absorbed, they pass into the bloodstream where they are transported to various parts of the body. Once in the proper location, the nutrients are either utilized or reorganized into new molecules. Whether or not the minerals and nutrients are absorbed and where they are absorbed depends on the solubility of those nutrients.

Several factors affect the solubility of the nutrients that we eat. The term solubility means the amount of a substance that dissolves in a given quantity of solvent at specified conditions of temperature and pressure to produce a saturated solution. Remember that a solvent is the dissolving medium in a solution. As the bloodstream is a fluid environment that transports nutrients, the solubility of these nutrients is vital to the well-being of the body.

Some nutrients are water soluble, meaning that they dissolve in aqueous solutions. Other nutrients are fat soluble, meaning that they dissolve in fatty tissues and oils. Some nutrients are also solubilized by the proteins that we eat and transported through the bloodstream by the proteins. It is very important that we understand the solubility of various nutrients so that we can understand both how our body uses those nutrients and how often we should eat those nutrients.

Several factors affect nutrient solubility. One of those factors is the temperature of the solution. Temperature is the average kinetic energy of the particles of a substance. The higher the temperature, the faster the particles move. This means that at a higher temperature, there is an increased chance that the solute particles will make contact with the solvent particles, increasing the rate that the solute dissolves. Most of the vitamins and minerals that we eat are soluble at body temperature, or 37ºC. A second factor has to do with the concentration of the solute in the solvent. There needs to be enough solvent to surround the solute particle so that it can be dissolved. This goes for both aqueous and fatty solutions in our bodies.

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C6ABE Solutes and the Process of Dissolving

Reading Science Continued 5

A third factor, and perhaps the most important factor for the solubility of nutrients, has to do with the molecular structure of the nutrients. There is an old saying with solubility rules: like dissolves like. This means that the structures of “like” substances will dissolve in each other. What does this mean? Molecules can be polar (have a charge) or nonpolar. Polar molecules tend to dissolve in polar solutions, such as water. Nonpolar fatty molecules tend to dissolve in lipids, such as fats.

The vitamins that we eat can be either water soluble or fat soluble. Water soluble vitamins are not stored in the body; therefore, they must be replenished each day as they are washed out with the urine. These vitamins include the B-complex vitamins and vitamin C. On the other hand, fat soluble vitamins become dissolved in fat before entering the bloodstream, and the excess is stored in the liver. These types of vitamins, which include vitamins A, D, E, and K, do not need to be replenished each day as they can be stored in fat tissues within the body.

Why are these important? Vitamin C is a critical antioxidant that aids in the health of cell membranes. The B-complex vitamins aid in metabolism and cell growth and division. Vitamin D aids in the absorption of calcium, an essential mineral. Calcium controls such functions as muscle contraction, nerve impulses, and blood clotting. If the concentration of calcium in our bloodstream (body fluids) gets too low, the calcium must be replenished by food intake or the body will take the necessary calcium from the bones.

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C6ABE Solutes and the Process of Dissolving

Reading Science 1 Paragraphs 4 and 5 list several factors that are important to the solubility of nutrients. Which of the following is not a factor in nutrient solubility? A The size of the molecule B The temperature of the solution C The molecular structure of the nutrient D The concentration of the nutrient in solution

2 Paragraph 1 discusses why the solubility of nutrients is important to the function of the human body. The solubility of nutrients is important because – A whether or not the nutrients are absorbed depends on their solubility. B where the nutrients are absorbed depends on their solubility. C Both A and B. D None of the above.

3 Which vitamin is water soluble, needs to be replenished each day, and is a critical antioxidant? A Vitamin A B Vitamin C C Vitamin D D Vitamin E

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C6ABE Solutes and the Process of Dissolving

Reading Science 4 Which vitamin is fat soluble and aids in the absorption of the essential mineral calcium? A Vitamin A B Vitamin C C Vitamin D D Vitamin E

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C6ABE Solutes and the Process of Dissolving

Math Connections Name:

Date:

Group:

Part I: Concentration of Reactants in Solution (Molarity) Molarity (M) is a ratio of the number of moles of a substance, or solute, dissolved in each liter of solution. Molarity (M)=moles of solute / liters of solution Moles of solute=molarity × final volume of solution Final volume of solution=moles of solute / molarity Use this information to answer the questions below. 1. Sulfuric acid is a very corrosive acid used as an industrial cleaning agent to remove rust from sheet metal. What is the molarity of a solution made by dissolving 6.652 g of sulfuric acid (H2SO4) in water and diluting to a final volume of 70.0 mL?

2. Nitric acid is used as an reagent in liquid-fueled rockets. The molarity of nitric acid is 16.0 M. How many moles of HNO3 are in 450.0 mL of 16.0 M solution?

3. How many moles of solute are present in the following solutions? a. 2.85 L of 1.59 M LiNO3

b. 23.9 mL of 0.156 M HCl

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C6ABE Solutes and the Process of Dissolving

Math Connections 4. How many grams of solute would you use to prepare the following solutions? a. 1 L of 0.250 M sucrose (C12H22O11)

b. 0.0673 L of 12.9 M osmium(III) fluoride (OsF3)

5. How many milliliters of a 3.5 M iron (II) nitrite (Fe(NO2)2) solution are needed to provide a total of 130.0 g of Fe(NO2)2?

6. If a chemist has 2.4 g of silver nitrate and needs to prepare 1.5 L of 0.010 M solution, will there be enough silver nitrate? If yes, how much is left over If not, how much more does he need?

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C6ABE Solutes and the Process of Dissolving

Math Connections Part II: Diluting Concentrated Solutions Similar to drink and juice mixes, some chemicals are bought and stored in concentrated solutions. They are diluted before they are used. When diluting a concentration, the number of moles stay constant, but the volume changes. Mf=Mi × Vi / Vf Where Mf is the final molarity, Vf is the final volume, Mi is the initial molarity, and Vi is the initial volume. 7. How would you prepare 0.10 L of 0.5 M NaOH solution starting from a concentration of 1.0 M?

8. What is the final concentration if 50.0 mL of a 0.20 M NiCl2 solution is diluted to a volume of 100.0 mL?

9. If 50.mL of 2.00 M H2SO4 solution is diluted to a volume of 200.0 mL, what is the final molarity of the solution?

10. How would you prepare 0.25 L of 0.948 M HCl solution starting from a concentration of 4.74 M?

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C6ABE Solutes and the Process of Dissolving

WRITING SCIENCE Name:

Date:

Group:

LOOK

THINK The concentration of a solution is expressed by the unit of molarity. The higher the molarity, the more concentrated the solution. Molar solutions are used in chemistry labs on a regular basis. These solutions must be produced with accuracy and precision for an experiment to be successful. The molarity calculation formula is: M = mole/liter. A solution can be created from either a solid or a liquid solute, usually with water as the solvent. A solid solute must be weighed out on a scale and then converted to a unit of mole. The solvent must always carry the unit of liter, and must equal the desired volume only after the solute has been added. The overall volume of the solute and solvent must equal the desired volume of the solution. WRITE how a 1 molar potassium chloride solution can be made. State the quantities of the solute and solvent needed.

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C6ABE Solutes and the Process of Dissolving

WRITING SCIENCE

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High School Chemistry

C6CD

Concentrations

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C6CD Concentrations

Student Handout Name:

Date:

A “KOOL” Activity A solution is a liquid mixture containing a solute (the material being dissolved) and the solvent (the material doing the dissolving).

Flask A

Flask B

Water mL

_______________

Powedered Drink Mix:

______________g

Which flask has a higher concentration of the tasty drink? __________________________

Can you tell which flask has a higher concentration by sight? _______________________

For both flasks, what is the solute? ____________________________________________

For both flasks, what is the solvent? ___________________________________________

For both flasks, what is the solution? __________________________________________

Number the following steps in order to prepare a solution in the correct sequence by writing a 1, 2, or 3 on the line. ____ Add the remaining amount of solvent carefully to the meniscus of the flask. ____ Measure the amount of solute needed for the solution and add it to the flask. ____ Add half the amount of solvent to the flask and stir to dissolve the solute.

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C 6AC Motion of Particles and Flow of Energy

Writing Science Name:

Date:

Group:

LOOK

THINK matter changes states all the time around us. Ice cubes melt in our drinks. Water turns to steam when heated. The gas in fluorescent light bulbs turns to plasma and back to gas again as we turn on and off the light. These shifts in phases are caused by the difference in movement of the particles that comprise the matter. WRITE what does particle movement look like in solids, liquids, gases, and plasma?

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C 6AC Motion of Particles and Flow of Energy

Writing Science

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C6CD Concentrations

Student Journal Name: Part I: What’s in a Solution?

Date:

Group:

What is a solute? Define the term, and then describe the units in which it is usually measured.

What is a solvent? Describe the role that a solvent plays in a solution.

In your own words, describe what a solution is. What does it contain? How can the contents of a solution be measured?

Explain why it is important to understand the concentration of a solution. How does this relate to both the solute and the solvent?

You will be preparing a solution of known molarity and diluting that solution to make a new solution with a lower concentration. Create a question of inquiry regarding the preparation and dilution of these solutions.

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C6CD Concentrations

Student Journal Part II: Molarity 1.

A solution consists of the grams of solute plus the volume of solvent. Describe how this relates to the concentration of the solution.

Describe how molarity relates to solution concentration. Include the appropriate units in your description.

Explain why molarity is based on the volume of the solvent and the volume of solution.

Could you calculate the molarity of a concentrated solution if you only knew the molarity and volume of the dilution? What would you need to know? What would you need to do?

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C6CD Concentrations

Student Journal Part III: Calculate the Concentration Use the table below to record the mass of the solute in grams and the volume of the solution in milliliters from your investigation. Make the appropriate conversions, and show your calculations in the space below. Then complete the rest of the table. Solute (g)

Solute (mol)

Solution (mL)

Convert the mass of solute to moles of solute.

Convert the milliliters of solution to liters of solution.

Calculate the molarity of the solution created.

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Solution (L)

Molarity of Solution

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C6CD Concentrations

Student Journal Part IV: Create a Diluted Solution 1.

How would you describe a dilution?

When diluting a solution, would you end up with a larger volume of solution or a smaller volume of solution? Explain why and be specific

Describe the process of properly diluting a solution from a known molarity of a concentrated solution. Be sure to include all of the necessary steps. Why is it important to be both accurate and precise in your measurements?

Using the proper formula, describe the steps required to create 100 mL of a 0.50 M NaCl solution from your concentrated solution. Write out your calculations below and show all of your work

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C6CD Concentrations

Student Journal Reflections and Conclusions 1.

You were asked to be both precise and accurate in your measurements. What is the difference between precision and accuracy?

You were asked to calculate the concentration of solutions. Why is this important when creating a solution?

Explain the steps taken when adding solute to solvent to make a solution.

You have been directed to make 500 mL of an MgCl2 solution using 12 g of MgCl2. What would happen if you added 12 g of MgCl2 to the beaker and then added 500 mL of water? How might this affect the solution concentration?

Using all of the following terms, develop a graphic organizer. Use additional paper if necessary.

Terms: solvent, solute, solution, mole, gram, molarity, concentration, volume, mass, milliliter, liter, equation, calculate

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STEMscopedia: CONCENTRATIONS C6CD

Reflect Lemonade is a tasty, refreshing drink to enjoy on a hot summer day. Tasty, that is, unless it’s not prepared properly. One person might make a batch of lemonade that has a bland taste. Another person might make a batch using the same ingredients but produce lemonade that is too tart to enjoy. A third person might make a batch that tastes absolutely delicious. What is it about three different batches of lemonade—all made using the same ingredients— that give rise to such different results that you can taste? What does this have to do with chemistry? The Concentration of a Solution Lemonade is an example of a solution. A substance that is dissolved in a solution is a solute. The substance in which a solute is dissolved is the solvent. In the case of lemonade, the solute is lemon juice and the solvent is water. Sugar is often added to lemonade, and this represents another solute in the lemonade solution. The difference between a sweet glass and a sour glass of lemonade has to do with the amounts of lemon juice and sugar mixed into a solution. The amount of solute present in a solution is the solution’s concentration. Solutions are important in chemistry because they are used in many laboratory procedures. When using solutions, chemists need to know the identities of the solutes and solvents in the solutions as well as the concentrations of the solutes. Molarity (M) is the concentration of a solution in chemistry. It is defined as the number of moles of solute per liter of solution, or mol/L. This unit of mol/L is also called molar. A two molar solution of sodium hydroxide (NaOH), for example, is written as 2 M NaOH. Some solutions are calculated by percent by mass. Percent by mass is mass of the solute per mass of the solution, then multiplied by 100. This unit is a percentage of the entire solution. Labeling Chemical Containers Chemicals in a lab as well as chemicals in everyday household products must be labeled correctly. The labels must include either the concentration (molarity) or the percent by mass of the active ingredients. In a lab, chemists making solutions must label the name, molarity, and preparation date of the solution, and then initial the container. All of this information is needed to identify the strength of the solution (how hazardous the solution is) and how to store the chemical.

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STEMscopedia: CONCENTRATIONS

Look Out Molarity is moles of solute per liter of solution, not per liter of solvent. To calculate molarity, do not first measure the volume of solvent and then add solute to it. The solute particles take up space, and your final solution volume will be a little larger than desired. Instead, begin with the solute and add enough solvent to produce the required amount of solution. Percent by mass is grams of solute per grams of solution, not per grams of solvent. However, when using percent by mass to calculate molarity, assumptions must be made. Assume that the solutions are dissolved in water and that the mass of the solute is so small, that it is considered negligible when taking the mass of the entire solution into account. Therefore, only the mass of water (instead of the entire solution) is needed. In order to convert mass of water into liters of water, the density of water is needed (density of water = 1 gram/1 mL). What Do You Know? Calculating Concentration in Units of Molarity Make sure the units are correct when calculating molarity. Suppose a chemist makes a saline solution that contains 23.4 grams of sodium chloride (NaCl) in enough water to produce exactly 275 milliliters of solution. What is the molarity of the solution and how would the chemist label the container? 1. Determine the amount of solute: The mass of the solute in this problem is given in grams, so it must be converted to moles. To do this, we first need to use the atomic masses of sodium (Na) and chlorine (Cl) from the periodic table to calculate the molar mass of NaCl:

2. We can then use the molar mass to calculate the moles of NaCl for the given mass:

3. Determine the volume of the solution: The volume of the solution must be converted from milliliters (mL) to liters (L):

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STEMscopedia: CONCENTRATIONS

4. Calculate the molarity: We can now use the moles of solute and liters of solution to calculate the molarity of the solution:

The chemist has produced a 1.45 molar saline solution. The chemist’s label for the container will include: 1.45 M Sodium Chloride; date it was prepared; and the chemist’s initials. Calculating Percent by Mass Make sure the units are correct when calculating percent by mass. Suppose a chemist has 35 grams of NaCl in a 500-gram solution. Calculate the percent by mass of the solution. 1. Use the percent by mass equation (grams solute / grams solution) x 100. 35 grams NaCl x 100 = 7% 500 grams solution Converting Percent by Mass to Molarity Sometimes normal household products list active ingredients by percentages. This is percent by mass. Percent by mass means grams of solute per grams of solution. Therefore, if a bleach solution has a percent by mass of 5.25 percent bleach (NaHClO), assume 5.25 grams NaHClO per 100 grams of solution. Percent by mass can be converted to molarity and vice versa. Again, assume that the solutions are dissolved in water and that the mass of the solute is so small that it is considered to be negligible. Therefore, only the mass of water (instead of the entire solution) is needed. In order to convert mass of water into liters of water, the density of water is needed (density of water = 1 gram/1 mL). Again, make sure the units are correct when calculating. 1. Obtain percent by mass of bleach. (5.25 percent Bleach) 2. Assume 5.25 grams NaHClO 100 grams H2O 3. Convert grams of NaHClO into moles of NaHClO. 5.25 grams NaHClO 1 mole NaHClO = .0695 moles NaHClO 1 75.45 g NaHClO

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STEMscopedia: CONCENTRATIONS 4. Convert grams of water into liters of water by using the density of water. 100 grams of H2O 1 mL H2O 1 L H2O = .1 L H20 1 1 gram H2O 1000 mL H2O 5. Calculate molarity of the bleach solution using the molarity equation. .0695 moles NaHClO = .69 M .1 L H2O

What Do You Think? What is the molarity of a solution made by dissolving one mole of solute in enough solvent to produce one liter of solution? How does this solution’s molarity compare to that of another solution made by dissolving 10 moles of solute in enough solvent to produce one liter of solution? Why would knowing the concentrations be helpful? What Do You Know? Fill in the missing quantities to describe the molarity and percent by mass of a NaCl solution.

Amount of NaCl

Moles of NaCl

Volume of the Original Solution

Density of Water

233.76 g

4 mol

1 g/mL

1 mol

1 g/mL

350.64 g 3 mol 116.88 g

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Molarity (moles/L)

Percent by Mass (g NaCl/ g solution) x 100

17.53 %

1 g/mL

0.5 M

17.53 %

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STEMscopedia: CONCENTRATIONS Connecting With Your Child To help students relate molarity with real-life applications, have them calculate the concentration of normal household products. Most household products list the active ingredient, in the form of a percentage, on the back or side of the container. Tell the student that they have a cut that they would like to treat with hydrogen peroxide (3 percent H2O2). However, the only hydrogen peroxide they have in the house in salon grade (10 percent H2O2). The hazards listed on this bottle include burning of skin, redness, and irritation. Using this percentage calculate the molarity of each hydrogen peroxide solution. Discuss the following questions with students: • What is the molarity of each solution? •

What is the difference between the 3 percent and the 10 percent H2O2?

•

Would you use the 10 percent H2O2 solution on your cut? Why or why not?

•

What would happen if the concentration was not listed on the container?

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C6CD Concentrations

Reading Science Name:

Date:

Group:

Concentrations in Laundry Detergent 1

A detergent is an agent that cleanses by being soluble in water and by helping to dissolve dirt, grease, or scum. Laundry detergent is one example of a commonly used detergent.

One of the new trends is to produce and sell highly concentrated detergents. What does it mean to be highly concentrated? In the case of detergents, concentrated means that the amount of water has been reduced, which, in turn, means that surfactants are present in a higher percentage than other substances (such as water). Surfactants are compounds in detergents that break down the barriers between water and dirt, scum, or oil. They hold these unwanted particles in suspension so they can be “washed away” from the clothes or dishes to which they were attached before the cleaning process began.

What are the benefits of this new trend of higher concentrations in detergents? One positive effect is economical. People are now buying a higher percentage of detergent per volume and are saving money. As compared to less concentrated products, for the same number of uses, environmental benefits are numerous. By reformulating their laundry detergent with 25% less water, one company used 45,000,000 fewer gallons of water while making products for the same number of laundry loads. This results in the need to manufacture fewer plastic containers, producing less pollution. In addition, people use fewer plastic containers, creating less trash and reducing recycling costs. Companies can also send out fewer shipments, which saves fossil fuels and reduces greenhouse gas emissions.

What is the difference in concentrations between regular laundry detergent and high-efficiency (i.e., highly concentrated) laundry detergent? It depends on what the concentrations of each type of detergent are. Stock solution is the term used to describe a highly concentrated solution. Stock solutions are created for the purpose of being accurately diluted to create a solution at a lower concentration. Although new detergents are more highly concentrated than before, they are still not as highly concentrated as commercial stock solutions. Companies use stock solutions to more accurately and uniformly dilute their detergents for customer products

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C6CD Concentrations

Reading Science 5

Concentration is usually expressed as molarity. What is molarity? It is how many moles of solute are contained in a liter of solution. The formula to find molarity is Molarity = moles x liters. Moles are expressed as 1 mole = 6.022 x 1023 atoms or particles. A mole is actually the atomic weight in grams of any molecule in a chemical. Every mole of every molecule in every chemical has the same number of molecules (6.022 x 1023), even though the atomic weight in grams is different. It is 1 gram for hydrogen, while it is 197 grams for gold.

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C6CD Concentrations

Reading Science 1. What does removing 25 percent of water in a detergent cause?

It creates a more dilute solution.

It creates a more concentrated solution.

The amount of water does not affect the concentration.

It causes the detergent to be less effective per volume.

If molarity = moles x liters, solve for liters. A.

Liters = molarity x moles

You cannot solve for liters.

Liters = moles/molarity

Liters = molarity/moles

If a concentrated detergent uses only half the water of the original detergent (which was 2.8 L), and the molarity of the original detergent is .080 M, what will be the molarity of the new, concentrated detergent? A.

1.4 L

160 M

.160 M

It does not change (.080 M).

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C6CD Concentrations

Reading Science 4.

Using the formula for molarity (molarity = moles x liters), what is the molarity of .450 moles of detergent dissolved in 100 mL of water? Calculate the molarity of this solution. A.

45 M detergent

.045 M detergent

.0045 M detergent

222 M detergent

Using the following formula (initial concentration x initial volume = final concentration x final volume), figure out the answer to the following question: If you have a 10x stock solution, how much of this stock would be needed to make 100 L of a 1x solution? A.

100 L

100 mL

10 L

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C6CD Concentrations

Math Connections Name:

Date:

Group:

Wholly Mol! Moles are used as a measuring unit for the number of atoms, molecules, or ions in a substance. To calculate moles, we use Avogadro’s number, which is 6.02 x 1023. To calculate molarity, use the following equation: Moles of Solute Molarity= Liters of Solution

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C6CD Concentrations

Math Connections 1.

The vinegar in the bottle consists of an acetic acid, CH3COOH, concentration based on diluting the acidity to 5%. What is the molarity? (The density of acetic acid is 1.02 g/mL.)

You are asked to make 750 mL of a 0.5 M acetic acid solution. How many moles of acetic acid would you need, and how many grams of acetic acid would you need

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C6CD Concentrations

Writing Science Name:

Date:

Group:

LOOK

THINK Consumers are met with many choices when it comes to products in a grocery store. When choosing between cereals, percent by mass is a useful measurement for overall health. When deciding between bottles of hydrogen peroxide, however, molarity becomes the focus based on the buyer’s purpose for buying the hydrogen peroxide. Think about the mathematical processes used to calculate molarity and percent by mass. WRITE Explain the mathematical process in which a commercial product can be evaluated, based on its concentration, in regard to molarity and percent by mass. Be sure to • clearly state your central idea; • organize your thoughts; • develop your essay in detail; • choose your words carefully; and • use correct spelling, capitalization, punctuation, and grammar.

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C6CD Concentrations

Writing Science

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High School Chemistry

C6FGH

Properties and Models of Acids and Bases

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C6FGH Properties and Models of Acids and Bases

Student Handout Name:

Date:

Common Household Acids and Bases Examine the five common household substances and rank them in order from the most acidic to least acidic (basic). Most acidic

Least acidic (basic)

____________________________________________

Now predict the pH of these household substances. 6.

Saline solution (for contact lenses) pH _____

Drain cleaner pH _____

White vinegar pH _____

Mouthwash pH _____

10. Clear soft drink

pH _____

11. Can you identify a common characteristic of an acid substance? If so, list below.

12. Can you identify a common characteristic of a basic substance?

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C6FGH Properties and Models of Acids and Bases

Student Journal Name:

Date:

Group:

Part I: Acid and Base Investigation 1.

Substance comparison

Conductivity

React w/Zinc

Color with PhTH

Guess?

Substance A Substance B Substance C Substance D Substance E 2.

Explain your reasoning for your guesses.

Based on your data, what can be determined about the pH of acids and bases?

Based on your data, what can be determined about strong versus weak acids and bases?

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Student Journal Background 1.

What are some characteristic properties of an acid? Name two everyday substances you can find in a home that are acidic.

What are some characteristic properties of a base? Name two everyday substances you can find in a home that are basic.

What is the Arrhenius definition of an acid?

What is the Arrhenius definition of a base?

Why did the Arrhenius definitions for acids and bases need to be modified?

What is the Brønsted-Lowry definition of an acid and a base?

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Student Journal Background, continued 7.

How is water able to act as both an acid and a base? Be as specific as you can.

Write the equation for the ionization of the water molecule to form the hydrogen ion and the hydroxide ion.

Compare the [H+] and [OH-] concentrations in a neutral, acidic, and basic solution.

10. The pH scale is a logarithmic scale based on the molar concentration of the hydrogen ion in an aqueous solution as seen by the following equation.

pH = -log [H+]

Use your calculator to find the pH of the following hydrogen ion concentrations. a) [H+] = 1.3 x 10-2, pH = b) [H+] = 5.6 x 10-5, pH = c) [H+] = 7.3 x 10-11, pH = d) [H+] = 3.9 x 10-8, pH = 11. What is the pH of a neutral solution? __________________________________________ 12. What is the H+ concentration of a neutral solution? 13. What is the OH- concentration of a neutral solution? 14.

Write the expression for the ionization constant for water and show its numerical value.

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Student Journal Part II: Neutralization Reactions Write and balance the following acid base reactions. Then label the Brønsted-Lowry acid, base, conjugate acid, and conjugate base. 1.

HCl + NaOH → ___________ + ____________

HCl + Ca(OH)2 → ___________ + ____________

HC2H3O2 + NaOH → ___________ + ____________

HC2H3O2 + Ca(OH)2 → ___________ + ____________

Part III: Plan Your Investigation 1. My question of inquiry:

What do you need to do to answer this question?

What are the variables that you will observe?

What do you predict will occur?

What materials, equipment, and technology will you need for this investigation?

What safety precautions must you take in this investigation?

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Student Journal Part IV: Implement Your Investigation Collect, Record, and Organize Data Procedure A. Strong Acid/Strong Base 1. Place 25 drops of the strong acid into a clean test tube. 2. Place five drops of bromothymol blue in the test tube and gently swirl with the acid. Record the color and pH. 3. Add 25 drops of base to a beaker. (You may need more but this is enough to start.) 4. Place five drops of bromothymol blue to the beaker and gently swirl with the base. Record the color and pH. 5. Add base to the test tube one drop at a time, keeping track of how many drops are added until the solution is green after it is gently swirled. 6. If you go over and the solution turns blue, add the acid one drop at a time until it goes back to green. 7. Record all data. B. Strong Acid/Weak Base 1. Place 25 drops of the strong acid into a clean test tube. 2. Place five drops of bromothymol blue in the test tube and gently swirl with the acid. Record the color. 3. Add 25 drops of base to a beaker. (You may need more but this is enough to start.) 4. Place five drops of bromothymol blue to the beaker and gently swirl with the base. Record the color and pH. 5. Add base to the test tube one drop at a time, keeping track of how many drops are added until the solution is green after it is gently swirled. 6. If you go over and the solution turns blue, add the acid one drop at a time until it goes back to green. 7. Record all data. C. Weak Acid/Weak Base 1. Place 25 drops of the weak acid into a clean test tube. 2. Place five drops of bromothymol blue in the test tube and gently swirl with the acid. Record the color. 3. Add 25 drops of base to a beaker. (You may need more but this is enough to start.) 4. Place five drops of bromothymol blue to the beaker and gently swirl with the base. Record the color and pH. 5. Add base to the test tube one drop at a time, keeping track of how many drops are added until the solution is green after it is gently swirled. 6. If you go over and the solution turns blue, add the acid one drop at a time until it goes back to green. 7. Record all data.

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Student Journal Part IV: Implement Your Investigation, continued D. Weak Acid/Strong Base 1. Place 25 drops of the weak acid into a clean test tube. 2. Place five drops of bromothymol blue in the test tube and gently swirl with the acid. Record the color. 3. Add 25 drops of base to a beaker. (You may need more but this is enough to start.) 4. Place five drops of bromothymol blue to the beaker and gently swirl with the base. Record the color and pH. 5. Add base to the test tube one drop at a time, keeping track of how many are added until the solution is green after it is gently swirled. 6. If you go over and the solution turns blue, add the acid one drop at a time until it goes back to green. 7. Record all data.

Initial acid pH and color

Drops of Acid

Initial base pH and color

Drops of base

Final solution pH and color

SA + SB

SA + WB

WA + WB

WA + SB

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Student Journal Part IV: Implement Your Investigation, continued 1.

What does it mean when we use the terms “strong” or “weak” when referring to acids and bases?

What does it mean when an acid and base “neutralize” in a reaction?

How did the strength of the acid or base affect the amount of drops needed to neutralize completely? Why?

Why do you need an indicator to tell how acidic or basic a solution is?

How did you use the indicator to determine how well you neutralized your solutions?

Based on the pH data you collected in the experiment, determine the hydronium ion concentration of the following.

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Student Journal Reflections and Conclusions 1.

Was your prediction correct or incorrect? Explain.

Was there a relationship between the variables you observed?

Where could you have made errors while collecting or organizing data?

What do you conclude about this investigation?

What would you do differently if you were to conduct this experiment again?

Using all of the following terms, develop a graphic organizer based on what you have learned in this Explore. Use another sheet of paper, if needed. Terms: acid, base, Arrhenius definition, Brønsted-Lowry definition, hydrogen ion, hydroxide ion, aqueous solution, pH scale, ionization constant of water 486

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C6FGH

Reflect Some flowering plants show variations in their flower color depending on soil conditions. Consider hydrangeas, shown in the photos below. When grown in acidic soil, hydrangeas tend to produce blue flowers (left). When grown in basic soil, hydrangeas produce pink flowers (right). What is the difference between acidic and basic soil, and why does this influence flower color in hydrangeas?

Chemists classify some chemical compounds as acids or bases. The terms acidic and basic are used when describing solutions or substances. These terms imply the presence of compounds known as acids or bases. Acids and bases are two classifications of chemical compounds. A compound is classified as an acid or a base based on its chemical properties. The first scientist to recognize and describe acids and bases was the Swedish chemist Svante Arrhenius. In the late 1800s, Arrhenius defined acids and bases as follows: • Arrhenius acid: a substance that dissociates to form hydrogen ions (H+) in water

•

Arrhenius base: a substance that dissociates to form hydroxide ions (OH−) in water

An example of an Arrhenius acid is hydrochloric acid (HCl). When placed in water (H2O), HCl dissociates—or breaks apart—to form hydrogen ions (H+) and chloride ions (Cl−), and an example of an Arrhenius base is sodium hydroxide (NaOH). When placed in water, NaOH dissociates to form sodium ions (Na+) and hydroxide ions (OH−).

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According to Arrhenius’s definition, only substances that release hydrogen ions may be classified as acids. Likewise, only substances that release hydroxide ions may be classified as bases. However, later scientists found many substances that do not fit these definitions, yet still have properties that are either acid-like or base-like. For example, ammonia, which has the formula of NH3, acts as a base, but it does not break apart to form hydroxide ions in solution. In 1923, two scientists each independently proposed new definitions to describe acids and bases. A Danish chemist, Johannes Brønsted, and an English chemist, Thomas Lowry, introduced the following definitions: • Brønsted-Lowry acid: a substance that donates hydrogen ions HA + H2O → A- + H3O+ •

Brønsted-Lowry base: a substance that accepts hydrogen ions B + H2O → OH- + BH+

For example, hydrochloric acid acts as an acid according to the Brønsted-Lowry definition because it donates hydrogen ions to water. Water in this case acts as a Brønsted-Lowry base: It accepts hydrogen ions to form hydronium ions (H3O+): HCl + H2O Cl− + H3O+ (acid) (base) Ammonia (NH3) acts as a Brønsted-Lowry base because it accepts hydrogen ions from water to form ammonium ions (NH4+). Because water donates hydrogen ions in this case, water acts as a Brønsted-Lowry acid: + H2O NH4+ + OH− NH3 (base) (acid) From the Brønsted-Lowry definitions, you can infer that acids and bases behave in pairs such that there is always an acceptor and a donor of hydrogen ions. This is an important point because it highlights the differences between the Arrhenius and Brønsted-Lowry definitions. Water behaves as a Brønsted-Lowry base in the example involving hydrochloric acid, above, and as a Brønsted-Lowry acid in the example involving ammonia. Depending on the situation, water can act as a donor of hydrogen ions in some situations and as an acceptor in other situations. This property of water allows it to behave as both an acid and a base. Water is said to be amphoteric because it can be classified as both an acid and a base. In fact, water demonstrates both acidic and basic properties when it reacts with itself, as shown below: H2O H3O+ + OHH2O + (acid) (base)

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Two molecules of water (H2O) react to form ions of hydronium (H3O+) and hydroxide (OH−). Similarly, a hydronium ion and a hydroxide ion react to form two water molecules. These reversible reactions happen continually in water.

Look Out Hydrogen ions do not exist freely in aqueous solutions. Rather, they associate with water molecules to form hydronium ions, which have the chemical formula H3O+. However, chemists often refer to hydrogen ions, particularly when talking about pH. When chemists talk about hydrogen ions, they are really talking about hydronium ions. An acid’s power can be expressed several different ways. The most common is through pH. It can also be measured through the concentration of hydrogen ions or the percent disassociation of the acid. Bases are measured along similar lines, only they are measured through pOH or the concentration of hydroxide ions. pH is a mathematical conversion of the actual concentration of hydrogen ions in solution. Typically, the concentration of hydrogen ions is much lower (4.5 x 10-5 mol/L is an average concentration, for example) than a normal solution. Because of this, pH is preferred because it makes numbers easier to grasp. To calculate the pH, you take the negative log of the concentration. The mathematical equation looks like this: pH=-log[H+] Logarithms are a method of taking very large or very small numbers and turning them into more easy to handle chunks. This is done by relating the number back to 10. A solution with a pH of 1.0 has a hydrogen ion concentration of 0.01 mol/L. This is because: 10-1.0 = 0.01 mol/L. A log function on the calculator just makes it easier. The greater the acidity of a solution, the greater is the concentration of hydrogen ions. The actual concentration of hydrogen ions in any solution—even one that is very acidic—tends to be low. Concentrations are often measured in ranges of 10−1 to 10−14 moles per liter. As we discussed earlier, logarithms relate numbers back to 10, so a little shortcut can be taken to estimate how much hydrogen is in solution.

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Look at the table below: pH

10-x

10-1

10-2

10-3

10-4

10-5

10-6

10-7

Actual Concentration

0.1 mol/L

0.01 mol/L or 1 x 10-2 mol/L

0.001 mol/L or 1 x 10-3 mol/L

0.0001 mol/L

0.00001 mol/L

0.000001

0.0000001

or 1 x 10-4 mol/L

or 1 x 10-5 mol/L

mol/L or 1 x 10-6 mol/L

mol/L or 1 x 10-7 mol/L

or 1 x 10-1 mol/L

When thinking about pH, just remember that changing the pH by 1 will change the concentration by 10. An acid with a pH of 2 has 10 times as much hydrogen as an acid with a pH of 3. An acid of pH 3 has 100 times as much hydrogen as an acid of pH 5.

As you go from left to right across this pH scale, the concentration of hydrogen decreases by 10 with each increase of pH by 1. At the same time, the concentration of hydroxide increases by 10 with each increase of pH by 1. At pH 7, the concentrations of hydrogen and hydroxide are exactly equal. The mixture is neutral. At pH greater than 7, there is more hydroxide in the solution than hydrogen, and the solution is called “basic.” So far, we have discussed solutions in terms of their hydrogen ion concentration or pH. In cases of solutions made using water as the solvent, it is also possible to talk about the solutions in terms of their hydroxide ion concentration. This is because water undergoes a reversible reaction in which it dissociates to form a hydrogen ion and a hydroxide ion: H+ + OHH 2O For bases, the power is measured by pOH, which is a conversion from the actual concentration of hydroxide ions in solution. For a solution that is 0.1 mol/L of base, it would have a pOH of 1.0, much like an acid would for pH. Because of the disassociation of water, as the hydroxide ion concentration increases, the hydrogen ion concentration will decrease.

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This means that as pOH becomes smaller and smaller, pH will become larger and larger. This relationship is shown by the equation below: pH + pOH = 14 For the base above, with a pOH of 1.0, it would have a pH of 13. Correspondingly, an acid with a pH of 3.0 would have a pOH of 11.0

What Do You Think? Many common items around your house are acidic or basic. Pick out three foods (like bread or coffee) and write them in the table below. Then, using a search engine, try to find out what the pH of the food is. Then using the pH scale above, figure out if the food is acidic or basic. Food Item

Acidic, Basic, or Neutral

Another way to describe the power of an acid is through its strength. The strength of an acid (or base) describes how much of the acid will react with water (or the base) to produce hydrogen in the solution. A strong acid will react 100 percent, producing the maximum amount of hydrogen in the solution. There are seven strong acids. All of the strong bases are Arrhenius bases, and produce hydroxide in solution. Strong bases will react 100 percent and will produce the maximum amount of hydroxide (or accept the maximum amount of hydrogen) in solution. There are eight strong bases.

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The 7 Strong Acids HCl HBr HI HNO3 H2SO4 HClO3 HClO4

Hydrochloric Acid Hydrobromic Acid Hydroiodic Acid Nitric Acid Sulfuric Acid Chloric Acid Perchloric Acid

The 8 Strong Bases LiOH NaOH KOH RbOH CsOH Ca(OH)2 Sr(OH)2 Ba(OH)2

Lithium Hydroxide Sodium Hydroxide Potassium Hydroxide Rubidium Hydroxide Cesium Hydroxide Calcium Hydroxide Strontium Hydroxide Barium Hydroxide

All other acids and bases are weak. This means that when they are dissolved in water, only part of the hydrogen is released. Or only some of the hydrogen is accepted. Citric acid, phosphoric acid, baking soda, and milk of magnesia are all weak. The amount that reacts is called the percent disassociation. Disassociation refers to the acid or base reacting with the water in solution. The following table shows examples of each type of acid or base. Chemical

Classification

Reaction in Water

Acetic Acid (vinegar) Ammonia (glass cleaner) Hydrochloric Acid Sodium Hydroxide

Weak Acid

HC2H3O2 + H2O ↔ H3O+ + C2H3O2-

Percent disassociation in a 0.10 mol/L solution 1.3 percent

Weak Base

NH3 + H2O ↔ OH- + NH4+

1.3 percent

Strong Acid

HCl + H2O → H3O+ + Cl-

100.0 percent

Strong Base

NaOH → OH- + Na+

100.0 percent

To find the percentage dissociation of an acid or a base, divide the concentration of reacted (or disassociated) acid or base by the total concentration and multiply by 100. For the reaction: HA + H2O ↔ H3O+ + A-

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You can see from the very low percentage dissociation values that weak acids and bases dissociate very weakly compared to strong acids and strong bases.

Look Out Acid strength is not measured in pH. Acid strength is an inherent characteristic of any acid regardless of its concentration. A strong acid can have a high pH if it is diluted and a low pH if it is concentrated. Similarly, a weak acid can have a high pH if it is diluted and a low pH if it is concentrated. Therefore, you cannot use pH as a measure of the strength of an acid. Now that we have defined acids, bases, and pH, we can return to the question raised at the beginning of this companion. Recall that hydrangeas produce blue flowers when grown in acidic soil, but they produce pink flowers when grown in basic soil. So why does soil acidity influence hydrangea flower color? A colored pigment present in the cells in hydrangea flower petals provides color. The pigment molecule is a weak acid. In the acid form (abbreviated below as HA), the molecule has a blue color. In the basic form (abbreviated as A−), the molecule has a pink color: HA + H2O ↔ H3O+ + ABecause this pigment is a weak acid, it can shift between its blue and pink forms depending on pH. The presence of aluminum ions also tends to stabilize the blue acid form. When planted in soils of low pH, hydrangea plants are better able to take up aluminum ions from the soil. As these ions accumulate in the petals, they shift the pigment to the acid (or blue) form. Soils of high pH inhibit the plants from taking up aluminum, resulting in a shift toward the pink form of the pigment. What Do You Know? 1. Sort the following compounds into the categories listed in the table below. Some compounds may be listed in more than one category. Compounds: HBr, H2SO4, Mg(OH)2, NH3, H2O Group Arrhenius Acid

Compound

Arrhenius Base Brønsted-Lowry Acid Brønsted-Lowry Base

2. Choose two reactants from the bank to write a balanced equation showing an acid-base reaction. Make sure you are dealing with an Arrhenius type reaction or a Brønsted-Lowry type reaction Reactant bank: HNO3, LiOH, H2O, NH3, Mg(OH)2, OH-, H3O+

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3. The following is a list of solutions located on a bench in a chemistry lab. Place an arrow on the pH scale below to indicate where each solution falls in terms of its pH. Label your arrow with the letter of the solution. Solution A: 0.10M HCl Solution B: 0.0010M NaOH Solution C: 1.0 × 10−5 M HNO3 Solution D: 0.10M KOH

4. The following list provides the degree of dissociation for several compounds when these compounds are placed in water at 1M concentrations. Sort the list into groups in the table provided below. RbOH: 100 percent dissociated HC3H5O2: 0.4 percent dissociated NH3: 0.4 percent dissociated HI: 100 percent dissociated HClO4: 100 percent dissociated NH2CH3: 2 percent dissociated NaOH: 100 percent dissociated HCHO2: 1 percent dissociated Group Strong Acid

Compound

Strong Base Weak Acid Weak Base

Connecting With Your Child Making Homemade Acid-Base Test Paper You can make acid-base test paper at home! Materials needed: • Red cabbage • Blender • Sauce pan • Measuring cups

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• • • • • • • •

Water Timer Coffee filter or fine mesh strainer Blank white paper Tray large enough to hold a sheet of paper Spoon Paper towels Newspaper

1. Cut two leaves of red cabbage into small squares, and puree in a blender with one cup of tap water. 2. Transfer the liquefied cabbage leaves to a small saucepan. 3. Repeat the step above several times until the saucepan is half full. 4. Heat the cabbage leaf puree to boiling on the stove. 5. Reduce the temperature to medium and allow the liquid to simmer 20 minutes uncovered. 6. Cool the liquid to room temperature, then pass it through a coffee filter or fine mesh strainer to remove solid plant material. 7. Pour enough of the filtered cabbage juice into the tray so that a sheet of paper is covered with liquid. 8. Carefully add a sheet of paper to the liquid. Press the paper with the spoon so that it becomes covered with liquid. 9. Soak the paper for at least five minutes, then lift it out and place it on paper towels that have been layered over newspaper. 10. Repeat to make several sheets of acid-base test paper. 11. Allow the papers to dry. 12. Test your acid-base test paper to see how its color reacts when exposed to acids and bases. Acids to test include vinegar, lemon juice, and soda. Bases to test include baking soda, ammonia, and window cleaner. 13. You can make a liquid slurry of baking soda and water by adding some of the powder to a small bowl, adding water, and stirring. Dilute vinegar, lemon juice, ammonia, and window cleaner with water to cut down on fumes and odors. 14. Try using a paintbrush dipped in each solution to transfer some of the acid or base to the acidbase test paper. Here are some questions to discuss with your child after you’ve tested some items around your house: 1. What color is associated with acidic conditions? 2. What color is associated with basic conditions? 3. How could you use your pH paper to test whether an unknown substance is an acid or a base? 4. Is the compound that you extracted from cabbage leaves itself an acid or a base? 5. How is this similar to the hydrangea pigment discussed earlier?

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Reading Science Name:

Date:

Group:

The Strength of Acids and Bases 1

You are most likely familiar with the terms acids and bases. Hydrochloric acid is a solution that many students have worked with in high school. Sodium hydroxide is a common base. You have also studied chemical reactions and equations. No doubt you are familiar with acid-base reactions, in which an acid and a base are combined, producing water and a salt. However, have you ever wondered what an acid and a base really are? Do you know why they react with other compounds in the ways that they do? Protons are the positive hydrogen ions associated with the acids, or the H+ ions. These protons may react with the hydroxide ions (OH−) of the base, producing water as a result. When acids and bases react with other compounds, what you are observing in those reactions is simply a transfer of protons.

Acids and bases have specific properties based on their chemical structures and arrangements. In the beginning, acids and bases were categorized by taste. Acids had a sour taste, and bases had a bitter taste. In 1884, a scientist named Svante Arrhenius discovered that all acids contained H+ ions and all bases contained OH− ions. Therefore, he concluded that an acid was any solution that contained more H+ than OH− ions. A base was defined as any solution that contained more OH− ions than H+ ions. This became the Arrhenius rule for distinguishing acids from bases.

In 1923, this definition was modified when two scientists, Johannes Brønsted and Thomas Lowry, independently made the same discovery. Their tests found that during an acid-base reaction, a proton was actually transferred from the acid to the base. They defined an acid as a proton donor and a base as a proton acceptor. This meant that any compound that donated a proton could be considered an acid. Any compound that accepted a proton could be a base. This became the Brønsted-Lowry rule for distinguishing acids from bases.

Acids and bases may be strong or weak. What does this mean? There are very few true strong acids. A strong acid is an acid that completely releases its hydrogen ions in solution, and a strong base is a base that completely ionizes in solution. Conversely, there are many weak acids. A weak acid is an acid that does not release all of its hydrogen ions in solution.

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Reading Science Continued A weak base is a base that does not completely ionize in solution. How strong or weak these compounds are can be determined, in part, by the pH level of the solution. The pH scale is a scale ranging from 1 to 14 that measures the concentration of hydrogen ions in a solution. A strong acid will have a pH reading of 1. A strong base will have a pH reading of 14. A neutral solution will have a pH reading of 7. 5

If we use the Brønsted-Lowry definition of acids (and bases), a strong acid will more easily lose its protons than a weak acid. Furthermore, a strong base will more likely attract protons than a weak base. Strong acids include compounds such as hydrochloric acid, or HCl. This acid will completely dissociate in solution, leaving H+ ions and Cl− ions. This acid is “happier” when it is dissociated into ions in solution. A common weak acid is acetic acid, the acid found in vinegar. This acid, CH3CO2H, will only partially dissociate in solution, leaving some H+ ions and some CH3CO2− ions. This acid is “happier” when it is an acid in solution. Therefore, in solution, strong acids tend to be ions, and weak acids tend to be acids.

So what does this have to do with strong and weak bases? Each acid has something called a conjugate base. The conjugate base of HCl is Cl−. The conjugate base of CH3CO2H is CH3CO2−. The strong acid HCl has a weak conjugate base of Cl−. Why? Remember that HCl prefers to be dissociated into ions in solution. So each time a Cl− ion is present in solution, it is very unlikely that it will accept an H+ ion. Therefore, Cl− has a weak attraction for protons, making it a weak base. On the other hand, the CH3CO2− ion is a strong base, as its conjugate acid prefers to be an acid in solution. This means that the CH3CO2− ion has a strong attraction for protons, making this ion a strong base. In general, the stronger the acid, the weaker its conjugate base, and vice versa.

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Reading Science 1 Paragraph 2 discusses one of the first scientific ways to distinguish acids from bases, known as the Arrhenius rule. According to this rule, what would be considered an acid? A Any solution with more H+ than OH− B A proton donor C Any solution with more OH− than H+ D A proton acceptor

2 H2SO4, sulfuric acid, is considered a strong acid. How would you categorize its conjugate base, SO42−? A A strong base B A weak base C An Arrhenius base D A Brønsted-Lowry base

3 According to this definition, a substance that can donate a proton (H+) is considered an acid and a substance that accepts a proton is considered a base. What is this definition known as? A Arrhenius definition B Brønsted-Lowry definition C Neutralization definition D Dissociation definition

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Reading Science 4 When a compound dissolves in an aqueous solution, not all the compound may dissociate and form ions in the solution. If a compound dissociates completely to release all of its hydrogen ions in aqueous solutions, what does it form? A A strong acid B A strong base C A weak base D Both A and C

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Math Connections Name:

Date:

Group:

Part I: Acidity with Arrhenius and Brønsted-Lowry Theories Compounds can be defined as acids or bases depending upon their characteristics and based on the how the molecules of the compound act when placed into water. Several scientists have developed their own definitions for acids and bases. In the late 1800s, Arrhenius defined acids and bases in aqueous solutions as follows: Arrhenius acid: An ionic compound that ionizes in an aqueous solution and produces hydrogen ions (H+) HCl

H+(aq) + Cl–(aq)

Arrhenius base: An ionic compound that ionizes in an aqueous solution and forms hydroxide ions (OH–). NaOH

Na+(aq) + OH–(aq)

However, later scientists found many substances that do not fit these definitions, yet still have other properties of either acids or bases. In 1923, a Danish chemist named Johannes Brønsted and an English chemist named Thomas Lowry independently proposed new definitions to describe acids and bases. Brønsted-Lowry acid: a compound that donates a proton (hydrogen ion H+) Brønsted-Lowry base: a compound that accepts a proton The ion or molecule remaining after the acid has lost a proton is known as that acid’s conjugate base, and the species created when the base accepts the proton is known as the conjugate acid. HCl+H2O

H3O++Cl–

HCl donates a proton to water

NH3+H2O

NH4++OH–

NH3 accepts a proton from water

acid

base

1. Write a balanced equation for the dissociation of each of the following Brønsted-Lowry acids in water. Determine the conjugate base of each acid. a. H2SO4 b. HNO3 2. Label the Brønsted-Lowry acids and bases and show the direction of the proton transfer. OH–+H3O+

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H2O+H2O

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C6FGH Properties and Models of Acids and Bases

Math Connections Part II: The pH Scale The pH of a solution or substance is the number used to describe the hydrogen ion concentration of that solution or substance. The pH scale is a range that goes from 0 to 14. pH

Classification

Highly acidic

Neutral

Highly Basic

Description High H+ ion concentration Hydrogen ion and hydroxide ion concentrations are equal. High OH– ion concentration

The pH scale is used to determine the strength of an acid or base and is equal to the negative logarithm (base 10) of the hydrogen ion concentration [H+], as follows: pH=−log [H+] for acidic solutions The letter p has been chosen to mean “negative logarithm of.” Thus, pH means the negative log of the hydrogen ion (H+) concentration, and pOH means the negative log of the hydroxide ion (OH–) concentration. Example: • HCl, a strong acid, pH of 0, [H+]=1×100 M or [OH–]=1×10–14 M • NaOH, a strong base, pH of 14, [H+]=1×10–14 M or [OH–]=1×100 M 3. Calculate the pH of an aqueous ammonia solution that has an OH− concentration of 3.9×10–3 M. Is the solution more basic or acidic? Step 1: Set up equation:

pOH=−log(3.9×10–3)

Step 2: Solve for pH:

pH+pOH=14

4. Calculate the pH of a sample of seawater that has an OH– concentration of 2.54×10–6 M.

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C6FGH Properties and Models of Acids and Bases

Math Connections 5. Calculate the concentrations of H3O+ and OH– in baking soda (pH 8.3). Step 1: Adjust the pH equation to solve for the concentration of [H+] ions.

pH=log[H3O+]

–pH=log[H3O+]

10–pH=10log[H3O+] since 10logx=x

[H+]=10–pH

Step 2: Substitute pH into equation and solve for concentration. [H3O+]=10–8.3= Step 3: Solve for the pOH. Remember pH+pOH=14. Step 4: Substitute pOH into equation and solve for concentration. [OH–]=10–pOH=

6. Calculate the concentrations of H3O+ and OH– in lemon juice (pH 2.0).

7. Calculate the pH of a 0.20 M solution of NaOH. Is the solution basic or acidic? NaOH

Na++OH–

Step 1: Since NaOH dissociates completely, the moles remain proportional, therefore:

pOH=−log[0.20] pOH=

Step 2: Solve for the pH. pH+pOH=14

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C6FGH Properties and Models of Acids and Bases

WRITING SCIENCE Name:

Date:

Group:

LOOK

THINK The use of pH indicators can be used to create certain colors of solutions. A change in pH will also create different observable colors. The hydrogen ion concentration of a substance can be used to calculate pH. pH can be used to determine the acidity or alkalinity of a solution. The pH scale begins with 0, which indicates an acid and continues through to 14, which indicates a base. The center point of the pH scale is neutral. pH indicators are used to visually observe the pH of a solution. These indicators come in various colors and will turn into specific colors when exposed to a specific range in pH. If a solution contains a pH indicator and displays a certain color, the color may be changed by adding an acid or a base. WRITE Explain the chemistry behind the demonstration, often seen in television commercials, where a colored solution changes colors when another colorless solution is added.

Be sure to • address the prompt • provide support • write legibly and concisely • organize your thoughts; • choose your words carefully; and • use correct spelling, capitalization, punctuation, and grammar.

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C6FGH Properties and Models of Acids and Bases

WRITING SCIENCE

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acid-base reaction

GLOSSARY OF TERMS

acid-base reaction – this type of double-replacement reaction occurs when equal amounts of an acid are added to a base so that the acid and the base neutralize each other, forming water and salt. The general equation is: HX (acid) + MOH (base) = H2O (l) + MX (salt) acidic compound – a compound in which the pH is between 0 and 6.9 acidic solution – a solution in which the hydrogen concentration is greater than the hydroxide concentration

atomic mass

actual yield – a measured quantity of the actual amount of product produced during a chemical reaction alpha particle – a positivelycharged particle identical to the helium atom nucleus, with two protons and two neutrons, that is emitted from radioactive decay of a nucleus anion – any atom or group of atoms with a negative charge aqueous solutions – a solution in which the solvent is water

atom – the smallest particle activation energy – the amount of energy input needed of an element maintaining the chemical identity of that to start a reaction element activity series – a list of atomic mass – the mass of metals ranked by their ability or tendency to react with other an atom, approximately equal to the number of protons and substances neutrons in the atom

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atomic number

atomic number – the number of protons found in the nucleus of an element, represented by the letter Z atomic radii – the distance between the nucleus of an atom to the outermost electron orbital, usually measured in nanometers average atomic mass – average mass of all known isotopes of an element based upon the abundance of the isotopes Avogadro’s Law – states that the number of moles of a gas varies directly with the volume of a gas at constant temperature and pressure Avogadro’s Number – expressed as 6.02 × 1023; the number of representative particles contained in one mole of a substance

beta particle

Avogadro’s Principle – equal volumes of the same gas at the same temperature and pressure contain the same number of molecules balanced chemical equation – a chemical equation in which mass is conserved and each side of the equation has the same number of atoms of each element barometer – an instrument that measures the amount of atmospheric pressure basic solution – a solution in which the hydroxide concentration is greater than the hydrogen concentration beta particle – a negativelycharged electron emitted from the radioactive decay of a nucleus

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binary molecule

GLOSSARY OF TERMS

catalyst

binary molecule – molecular compounds formed by two nonmetal elements or by a metalloid and a nonmetal element

bond strength – the degree to which each atom is joined to each other in a chemical bond which contributes to the valency of the other atom(s)

Bohr Model – this model was defined by Niels Bohr in 1913 and shows the negativelycharged electrons of the atom travelling around the positivelycharged nucleus in discrete circular orbits, much like the structure of the solar system

bonding pairs – the valence electrons in an atom that are bonded or shared with other atoms

Boyle’s Law – states that the pressure of a gas varies inversely, with volume at constant temperature bond – any of several forms of expressed as the equation P1V1 electrostatic attraction between = P2V2 atoms that holds the atoms brittleness – the tendency of together a mineral to fracture or break bond energy – a measure of when pressure is applied the strength of chemical bonds; catalyst – a substance that equal to the heat required to speeds up or promotes a break one mole of molecules chemical reaction without being into their individual atoms chemically changed by the reaction

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cation

cation – any atom or group of atoms with a positive charge Celsius – scale on a thermometer based on the water’s boiling (100oC) and freezing (0oC) points; named after Anders Celsius, a Swedish astronomer

coefficient

chemical properties – characteristics that can only be observed or measured when atoms of matter rearrange during a chemical change chemical reaction – the process by which one or more substances change to produce one or more different substances

central atom – the atom around which other atoms bond chemical reaction – a process in a molecule that leads to the transformation Charles’ Law – states that of a set of chemicals by the temperature of a gas rearranging the atoms or bonds varies directly, with volume at between the atoms constant pressure expressed coefficient – a number placed as the equation V1/T1 = V2/T2 in front of a chemical symbol or chemical names – the names formula in order to balance the of elements or compounds equation based on a standardized naming process

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GLOSSARY OF TERMS colligative properties

colligative properties – properties of solutions that depend upon the ratio of the number of solute particles to the number of solvent molecules in a solution, and not on the type of chemical species present color change – a change that alters the identity of a substance, resulting in a new substance or substances with different properties combination reaction – also known as a synthesis reaction; two or more reactants will combine during a chemical change to create one product. The general equation is: A + X = AX combined gas law – can be used to determine the pressure, volume, or temperature of gases and can be expressed as P1V1/T1 = P2V2/T2 © Accelerate Learning Inc. - All Rights Reserved

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covalent bond

combustion reaction – this type of reaction is an oxidation process in which a compound of carbon, hydrogen, and sometimes oxygen reacts with oxygen gas to produce carbon dioxide gas and water. The general equations are: CxHyOz + O2 CO2 + H2O or CxHy + O2 CO2 + H2O concentration – a measurement of the amount of solute that is dissolved in a given quantity of solvent conductive – having the property to transfer energy from one medium to another through direct contact covalent bond – a chemical bond where one or more electrons is shared by two separate atoms; a form of chemical bonding that is characterized by the sharing of pairs of electrons between atoms 511

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covalent compound

covalent compound – a compound created when two (or more) atoms bond covalently, sharing electrons and creating a molecule Dalton’s Law of Partial Pressures – states that the total pressure of a mixture of gases is the sum of the partial pressures of the gases contained within that mixture decomposition reaction – where a single compound on the reactant side breaks down into two or more products during a chemical change; the general equation is: AX A + X dipole – a pair of equally and oppositely charged poles separated by a distance

dissolve

directly proportional – a relationship between variables in which the value of one variable can always be determined by multiplying the value of the other variable by a constant dispersion force – an intermolecular attraction force that exists between all molecules and is the result of movement of electrons causing slight polar moments dissociation – a process in which molecules or compounds split into smaller pieces dissolve – to become incorporated into a liquid to form a solution

dipole-dipole force – attractive forces between the positive end of one polar molecule and the negative end of another 512

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GLOSSARY OF TERMS double-replacement reaction double-replacement reaction – in this type of reaction, the reacting compounds exchange cations during the chemical reaction. Precipitation reactions and acid-base reactions are both examples of this type of reaction. The general equation is AX + BY = AY + BX electric current – the flow of electricity through a circuit electron – a negativelycharged subatomic particle of the electron cloud; involved in the formation of chemical bonds electron affinity – a neutral atoms likelihood of gaining an electron and becoming a negative ion electron configuration – the arrangement of an element’s electrons in various orbitals around the nucleus to achieve the lowest energy state © Accelerate Learning Inc. - All Rights Reserved

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element identification

electron dot formula – a notation that depicts valence electrons as dots around the atomic symbol of the element; also known as a Lewis dot formula electron orbital – the specific regions around the nucleus of an atom in which electrons travel electronegativity – within a molecule, it is the ability of an atom in that molecule to attract electron pairs to itself element – a pure substance that cannot be separated into simpler substances by physical or chemical means element identification – a symbol and name attached to a specific type of atom

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emission

formulas

emission – the act of discharging energy, light, particles, and different types of waves

excess reagent – any reactant that remains after the reactants in a chemical reaction have completely reacted

empirical formula – a chemical formula with the lowest whole-number ratio of elements in a compound

exothermic – a process in which heat is released

endothermic – a process in which heat is absorbed

fluidity – the physical property that enables a substance to flow

formula mass – the sum of energy – the ability of a system the atomic masses of ionic compounds to do work or produce heat energy level – regions around the nucleus of an atom where electrons may be found

formula unit – the compound created from ionic bonds or the empirical formula of molecular compounds; NaCl is a formula unit

enthalpy – the heat content of a system plus the product of the pressure and volume of the formulas – a short way of expressing information using system symbols and numbers to equilibrium – a condition in represent what makes up an which all competing influences element or compound are balanced

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Gay-Lussac’s Law

GLOSSARY OF TERMS

Gay-Lussac’s Law – states that the volumes of gases undergoing a reaction at constant pressure and temperature are in a simple ratio to each other and to that of the product heat – energy transferred between two objects of different temperatures, moving continually in a predictable pattern from warmer site to cooler site until all sites have reached the same temperature heat of formation – the thermal energy gained or released during the formation of one mole of a substance from its constituent elements heat of fusion – the change in enthalpy required to change a given amount of a substance from a solid to a liquid at the melting point of the substance

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hydrogen bond

heat of vaporization – the change in enthalpy required to change a given amount of a substance from a liquid to a vapor at a given pressure heating curve – shows how the temperature changes as a substance is heated up helium – colorless inert gas that has 2 protons Hess’s law – law that states that regardless of the various stages/steps of a reaction, the total enthalpy change for the reaction is the sum of all changes hydrogen bond – a relatively weak attraction between a hydrogen atom covalently bound to a strong electronegative atom and some other strong electronegative atom

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hydronium ion

hydronium ion – a cation consisting of one oxygen atom and three hydrogen atoms ideal gas – a theoretical gas that perfectly follows the ideal gas law ideal gas law – an equation describing the behavior of a gas in terms of temperature, pressure, volume, and number of moles; expressed as PV = nRT insoluble – a substance that cannot be dissolved within a specific medium insulative – material through which heat or electricity does not easily flow intermolecular force – forces of attraction or repulsion which act between neighboring particles

ionic bond

intramolecular force – any force that holds together the atoms making up a molecule or compound (contains all types of chemical bonds) inversely proportional – one variable decreases while the other variable increases, provided their product remains constant ion – atoms (or groups of atoms) that have an electrical charge due to unequal numbers of protons and electrons ionic bond – a form of chemical bonding that is characterized by the electrostatic attraction that binds oppositely-charged ions together

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ionic compound

GLOSSARY OF TERMS

ionic compound – a compound created when two or more oppositely-charged ions combine to form a formula unit, usually in the form of a crystalline structure

light emission

IUPAC nomenclature – standardized set of rules to generate universal names for chemical compounds

kelvin – the SI base unit of temperature; represented as ionic radii – the radius of either the letter K a positive or a negative ion law of conservation of mass found in an ionic crystal – the law that states matter can change states, but the mass ionization – the process of converting an atom or molecule remains the same into an ion by adding or Lewis structure – shows an removing charged particles element’s symbol surrounded by dots representing the ionization constant – a element’s valence electrons measure of the strength of an placed on the four sides of the acid in solution symbol; also known as electron ions – an atom with a net dot structure electrical charge resulting from light emission – giving off light the loss or gain of electrons isotope – any two or more forms of an element with the same number of protons in the nucleus, but a different number of neutrons © Accelerate Learning Inc. - All Rights Reserved

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limiting reactant

limiting reactant – any reactant that is the first to be completely consumed in a chemical reaction, limiting the amount of product that can be produced lone pair of electrons – pairs of valence electrons that are not bonded to other atoms malleability – a property that describes the ability of the material to be rolled or hammered into thin sheets mass – a measure of how much matter is present in a substance metallic bond – bond characteristic of metals in which free (non localized) valence electrons are shared among positively charged ions, usually in a stable crystalline structure

molarity

mm hg – millimeter of mercury; non-SI unit of pressure approximately equal to 1 Torr or 1/760 atm molar concentration – the amount of moles of the substance divided by the volume of the mixture molar mass – a general expression used to refer to the mass (in grams) of a mole of any substance, expressed as grams per mole, or g/mol. molarity – the concentration of solute in a solution expressed as the number of moles of solute dissolved in 1 liter of solution; expressed by the equation mol/L

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mole

GLOSSARY OF TERMS

mole – the SI unit is used to describe an amount of a substance. The mole is equal to 6.02 × 1023 atoms or molecules of a substance. One mole is the amount of any substance that contains the same number of units as the number of atoms in exactly 12 grams of Carbon-12 mole ratio – the ratio between the amounts of moles of any two substances in a chemical reaction molecular formula – a chemical formula of a compound that is the empirical formula of a compound multiplied by a whole-number subscript molecular geometry – the arrangement of atoms in relation to the central atom of the molecule

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nuclear equation

molecular mass – the sum of the atomic masses of molecular compounds neutral – having no electrical charge neutralization reaction – an acid and a base react quantitatively with each other in a chemical reaction neutron – a subatomic particle of the nucleus of an atom that is without charge and contributes to the mass of an atom nonpolar covalent bond – bonding in which two atoms equally share a pair of electrons with each other nuclear equation – an equation that shows how a nucleus gains or loses subatomic particles

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nuclear fusion

nuclear fusion – a reaction in which two light nuclei combine to produce a nucleus of heavier mass and results in the release of a large amount of energy nuclear model – this model, also known as the planetary model, was defined by Ernest Rutherford in 1911 and shows a dense, positively-charged core surrounded by negative charges nuclear reaction – the process in which a nucleus that is stimulated will emit a nucleon without a significant change in its atomic weight

periodicity

oxidation-reduction reaction – a chemical reaction that involves the transfer of electrons between atoms or molecules, changing the oxidation state of the reactants; the atom or molecule that gains electrons is oxidized, and the atom or molecule that loses electrons is reduced particle size – dimension (normally average diameter) of a single grain percent yield – the ratio of the actual yield to the theoretical yield for a chemical reaction expressed as a percentage, shown as actual yield/ theoretical yield x 100 periodicity – the quality or state of being periodic, or occurring at regular intervals

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pH scale

GLOSSARY OF TERMS

pH scale – a scale ranging from 1-14 that measures the concentration of hydrogen ions in a solution phase transition – the transformation of a thermodynamic system from one phase or state of matter to another one by heat transfer physical property – characteristic that can be observed or measured without changing the substance; for example, color, melting point, or conductivity polar – describes a molecule in which one or more atoms is slightly negative and one or more is slightly positive polarity – a separation of electric charge between atoms in a molecule that result from the unequal sharing of electrons in a covalent bond

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radioactive decay

recipitation reaction – this type of double-replacement reaction occurs in aqueous solutions of ionic compounds where one of the reactants formed is a solid that “precipitates” from the solution. The general equation is: AX (aq) + BY (aq) AY (aq) + BX (s) pressure – force of objects pushing on other objects; an amount of force per unit area product – the ending substance(s), written on the right side of the chemical reaction arrow, that is created during a chemical change radiation – the complete process in which energy is emitted as particles or waves radioactive decay – the spontaneous emission of radiation by an unstable nucleus

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radioactivity

GLOSSARY OF TERMS

solute

shielding effect – an effect radioactivity – the processes by which unstable atomic nuclei that occurs as the inner electron shells of an atom achieve stability “shield” the valence electrons reactant – a substance that in the outer electron shells from takes part in and undergoes the positive pull of the nucleus change during a reaction single-replacement reaction reactant – the starting – a redox reaction where one substance, written on the left element or ion moves out side of the chemical reaction of one compound and into arrow, which will be destroyed another during a chemical change solubility – the ability of reverse reaction – a chemical a substance to dissolve in reaction that results in an another substance to form equilibrium mixture of reactants a homogenous mixture; and products measurement of the ability of a solid to dissolve in a liquid sea of electrons model – the valence electrons (in a metallic soluble – able to be dissolved bond) from the s and p orbitals in another substance to form a of the interacting metal atoms homogeneous mixture delocalize; instead of orbiting solute – a substance that their metal atoms, they form a sea of electrons that surrounds is dissolved, or spread the positively charged nuclei of homogeneously, in another aqueous substance the interacting metal ions

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solution

GLOSSARY OF TERMS

theoretical yield

solution – a homogeneous mixture composed of only one phase (solid, liquid, or gas)

subatomic particles – particles that are smaller than the atom

solvent – a substance in which another substance (solute) is dissolved to form a homogeneous mixture

temperature – a measure of the average amount of kinetic energy of the molecules of a substance; average kinetic energy of all the particles in a material, measured by a thermometer in degrees (usually Celsius or Fahrenheit)

specific heat – the amount of heat required to raise the temperature of one gram of an object by one degree speed of reaction – how fast or slow a reaction takes place strong acid – an acid that completely releases its hydrogen ions in solution strong base – a base that completely ionizes in solution structural formula – a chemical formula that shows the arrangement of atoms in a molecule or a polyatomic ion

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temperature change – increase or decrease of heat energy in a substance, which may be evidence of a new substance formed during chemical change theoretical yield – a quantity that is always calculated and shows the theoretical amount of product that could be produced in an ideal chemical reaction in which there is a complete conversion of reactant(s) to product(s)

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titration

GLOSSARY OF TERMS

weak base

titration – the process of determining the concentration of a solution by adding a known concentration and volume of another solution to produce an observable reaction valence electrons – the electrons in the outermost electron orbital of an atom that influence how an element will react with other substances viscosity – a property of liquid that makes it resistant to flowing volume – a measure of the space that matter occupies weak acid – an acid that does not release all of its hydrogen ions in solution weak base – a base that does not completely ionize in solution

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Chemistry STEMscopes Georgia Student Notebook

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Chemistry Notebook

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