Chemistry for Changing Times, 12E Hill, Kolb, McCreary Solution Manual

Chemistry for Changing Times, 12E By Hill, Kolb, McCreary

Email: Richard@qwconsultancy.com

Chapter 1

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Chapter 1. Chemistry Overview Chapter 1 introduces the student to the concepts of science, and its ability to solve problems as well as its limitations. The student is then introduced to the concept of matter and energy and the physical means by which they are measured.

Lecture Outline 1.1

1.2

1.3

1.4

1.5 1.6 1.7 1.8

Science and Technology: The Roots of Knowledge Many students are confused about the differences between science and technology. Technology is the direct application of knowledge to solve problems. Science seeks an understanding of underlying principles.

The Baconian Dream and the Carsonian Nightmare Bacon was one of the first to propose the experimental approach to science. Bacon also felt science could solve all problems. Rachel Carson was one of the first “environmentalists.” She pointed out the dangers of DDT and other chemicals damaging to the environment, indicating that the application of science was perhaps not all good.

Science: Testable, Reproducible, Explanatory, Predictive, and Tentative The scientific method, from hypothesis to law, theory, or model; is best explained by example.

The Limitations of Science The scientific method is dependent on the ability to control all aspects of a problem; this limits the method to “simple” systems.

Science and Technology: Risks and Benefits Chemistry: Its Central Role Solving Society’s Problems: Scientific Research Research can be basic or applied.

Chemistry: A Study of Matter and Its Changes Matter is anything that has mass.

Mass is a measure of the amount of matter present.

Weight is mass times gravitational attraction.

Matter exhibits chemical and physical properties.

Chemical properties tell us how matter will combine to form new and different substances (a chemical change).

Physical properties are directly observable: color, state (solid, liquid, or gas), and texture are examples.

An excellent demonstration is to burn a candle (chemical change) and boil water (physical change).

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1.10

1.11 1.12

1.13

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Classification of Matter A. States of Matter • Solids, liquids, and gases.

B. Substances and Mixtures • Substances have constant composition.

• The composition of a mixture is variable.

▪ Homogeneous mixtures: appear the same throughout (milk, paint, and saltwater are examples).

▪ Heterogeneous mixtures: appear different throughout (pizza, raisin bread, and chocolate chip cookies are examples).

C. Elements and Compounds • Substances are either elements or compounds.

▪ Elements: fundamental building blocks of all matter. (lead, silver, gold, carbon, and oxygen) ▪ Compounds: two or more elements chemically combined in fixed ratios. (water, ammonia and propane) D. Atoms and Molecules The Measurement of Matter Most students have been introduced to the metric system during their K-12 education; the SI system used in science is based on the metric system. The difference lies in the base units.

• Base units: kilograms (kg) for mass, meter (m) for length, and seconds (s) for time. The four other base units are shown in Table 1.4.

Density Density is the mass-to-volume ratio of matter.

Energy: Heat and Temperature The SI unit of temperature is the kelvin (K); however, the temperature scale used in the chemistry laboratory is normally the Celsius scale. The Celsius scale was developed with the freezing and boiling points of pure water at a pressure of 1 atmosphere as the reference frames. (Students find the history of the development of the Fahrenheit scale much more interesting!)

Critical Thinking

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Chapter 1

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Transparencies T-1 T-2 T-3 T-4 T-5 T-6 T-7 T-8 T-9

Inside cover Figure 1.1 Figure 1.4 Figure 1.5 Figure 1.6 Figure 1.7 Figure 1.9

Periodic Table The evaporation of water Risks of death (p. 10) Comparison of copper and sulfur Solids, liquids, and gases A scheme for classifying matter Comparison of metric and customary units Comparison of large and small objects (p. 20) A comparison of the Fahrenheit, Celsius, and Kelvin temperature scales

Demonstrations 1. Place samples of various elements (copper, sulfur, zinc, mercury, aluminum, carbon, etc.) in small stoppered bottles or flasks that can be passed around.

2. Ask students: “Which is heavier, lead or aluminum?” They will say “lead.”

Then bring out two flasks, a 50 mL flask with about 10 mL of lead shot and a 500 mL flask about 3/4 full of aluminum pellets. Point out that the aluminum is actually heavier. But we need to compare samples of the same size (volume).

Weight per unit volume is density. Lead has the greater density.

3. Compare a yardstick and a meter stick. A meter is slightly bigger than a yard.

A sugar cube is approximately 1 mL. An ordinary paper clip weighs about 1 gram.

A liter is a little bit (6%) larger than a quart.

4. Use Crispix and Raisin Bran to illustrate the difference between a compound and a mixture. Raisins and bran flakes are easily separated, and the ratio of raisins to flakes can vary. Crispix has a 1:1 ratio of corn and rice flakes, and they cannot be easily separated.

5. Ice floats on water because the density of ice is less than that of water. But alcohol has a much lower density. If you use alcohol instead of water, the ice will sink. Place a glass of water and a glass of alcohol side by side and add an ice cube to each one.

6. The “cartesian diver” is a popular demonstration about density. Completely fill a 2-liter plastic bottle with water, and put just enough water into a medicine dropper so that it can barely float. Put the dropper into the bottle, and cap the bottle. Squeeze and release the bottle to make the “diver” go up and down.

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Chapter 2

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Chapter 2. Atoms Overview Chapter 2 begins with the Greek concept of matter and moves chronologically forward through the development of the atomistic model of matter.

Lecture Outline 2.1

2.2

2.3

2.4

2.5

2.6

Atoms: The Greek Idea The Greek philosophers had two main ideas regarding matter: The atomistic view (Leucippus and Democritus) The continuous view (Aristotle) Lavoisier: The Law of Conservation of Mass Lavoisier was the first great experimentalist. More than anything, he introduced experimental chemistry to Western civilization. Through his experiments he formulated the law of conservation of mass which states: Matter is neither created nor destroyed during a chemical reaction.

Proust: The Law of Definite Proportions The law of definite proportions is strong evidence for the atomistic nature of matter. If matter were continuous, any ratios by mass would be possible.

John Dalton and the Atomic Theory of Matter John Dalton summarized the work of Lavoisier and Proust to form the famous atomic theory.

Out of Chaos: The Periodic Table John Dalton established relative atomic masses using hydrogen as a base.

Mendeleev arranged a table of elements according to increasing atomic weights placing elements with similar properties in the same column. This was the beginning of the modern periodic table!

Atoms: Real and Relevant

Transparencies T-10 T-11 T-12

Figure 2.1 Figure 2.4 Figure 2.5 Figure 2.6

T-13 T-14 of

Figure 2.7 Figure 2.8

Table 2.1

A sandy beach Mercuric oxide and conservation of mass Law of Definite Proportions and Berzelius’s experiment illustrating the law of definite proportions Electrolysis of water The laws of definite proportions and conservation mass interpreted in terms of Dalton’s atomic theory The law of multiple proportions. A carbon atom can combine with either one or two atoms of oxygen

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Demonstrations 1. Electrolysis of water.

2. Fill 3 flasks with zinc (gray powder), sulfur (yellow powder), and zinc sulfide (white powder), and show to the class to illustrate how the properties of a compound differ from those of the component elements. [Or use mercury (silver liquid), oxygen (colorless gas), and mercuric oxide (red powder).]

3. Use samples of metals and nonmetals to demonstrate the difference in their properties. (Aluminum foil, copper wire, nickel coin, zinc strip; capped bottle of liquid bromine, balloon filled with hydrogen, bottles of carbon, sulfur, etc.)

4. Using the largest nuts and bolts you can find, start with unequal numbers of nuts and bolts (e.g., 5 and 7). Connect the pairs of nuts and bolts, with some remaining unpaired, to illustrate the law of definite proportions.

5. Heat a sample of HgO in a large test tube with a burner. Mercury droplets form on the inside of the tube. Test for oxygen with a glowing splint, which bursts into flame when inserted into the test tube.

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Chapter 3

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Chapter 3. Atomic Structure Overview Chapter 3 examines the experimental evidence leading to modern atomic theory. The chapter opens with the earliest experiments that indicated the electrical nature of matter and concludes with modern quantum theory.

Lecture Outline 3.1 Electricity and the Atom Humphry Davy’s experiments with the voltaic pile led to the electrolysis of molten salts. Michael Faraday refined Davy’s procedures and became the father of electrochemistry. Cathode ray tube experiments conducted by J. J. Thomson allowed him to calculate the mass-to-charge ratio of the electron. Conducting similar experiments, Eugen Goldstein discovered the proton. Millikan’s oil-drop experiment enabled him to calculate the charge of the electron, and thus its mass.

3.2 Serendipity in Science: X-Rays and Radioactivity Roentgen, in 1895, stumbled upon X-rays while conducting cathode ray experiments.

3.3 Three Types of Radioactivity The alpha particle () is a helium nucleus.

The beta particle () is an electron.

The gamma ray () is a high-energy photon.

3.4 Rutherford’s Experiment: The Nuclear Model of the Atom The gold foil experiment was strong and convincing evidence that led to the theory of the nuclear atom.

3.5 The Atomic Nucleus The nucleus contains protons (+1 charge), neutrons (neutral), and nearly all the mass of the atom. Electrons are outside the nucleus, have a -1 charge, and only a fraction of the mass of the proton and neutron. It is the number of protons that defines an element! Isotopes are atoms of the same element having different numbers of neutrons, and thus varying in mass.

3.6 Electron Arrangement: The Bohr Model Spectroscopic evidence led Bohr to conclude that only certain energy levels, or orbits, were allowed for electrons.

3.7 Electron Arrangement: The Quantum Mechanical Model Louis de Broglie showed that the electron had wave properties. Schrodinger used the de Broglie wave theory and the Bohr quantum theory to derive the quantum mechanical theory of the atom. The quantum mechanical theory describes the positioning of electrons about the nucleus and electron energies.

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3.8 Electron Configurations and the Periodic Table Figure 3.23 is an excellent tool for teaching students electronic configurations using shorthand spectroscopic notation based on element locations in the periodic table.

3.9 Which Model to Choose?

Transparencies T-15 T-16 T-17 T-18 T-19

Figure 3.1 Figure 3.2 Figure 3.4 Figure 3.5 Figure 3.6 Figure 3.7

T-20 T-21 T-22 T-23 T-24 T-25 T-26

Figure 3.9 Figure 3.11 Figure 3.12 Figure 3.14 Figure 3.15 Figure 3.16 Figure 3.17

Electrolysis apparatus Thomson’s apparatus showing deflection of cathode rays Millikan oil drop experiment Behavior of radioactive rays in an electric field Rutherford’s gold foil experiment and Model explaining the results of Rutherford’s gold foil experiment Some elements that give colored flames Line spectra of selected elements Possible electron shifts between energy levels Charge cloud representation of atomic orbitals Electron orbitals of the second main shell Order-of-filling chart Valence shell configurations and the periodic table

Demonstrations 1. Demonstrate the use of a cathode ray tube, using a magnet to deflect the beam of electrons.

2. Models that can be used to demonstrate the Rutherford gold foil experiment are described in Tested Demonstrations, Vol. II, by George Gilbert et al., on pages L-2,3,4.

3. An overhead projector version of the Rutherford experiment is described by Hau in J. Chem. Educ. 1982, p. 973.

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Chapter 4

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Chapter 4. Chemical Bonds Overview Chemical bonds determine the properties of matter. Chapter 4 examines ionic and covalent bonds, the symbolism used to represent bonds, and the resulting shape and properties of molecules and compounds that result from different types of bonding.

Lecture Outline 4.1 The Art of Deduction: Stable Electron Configurations The most stable electron configurations are those of the noble gases.

Other elements that can alter their electron structures to become like those of noble gases would become less reactive by doing so Formation of isoelectronic ions 4.2 Lewis Electron-Dot Structures Use the periodic table to determine the number of valence electrons Lewis dot symbols use “dots” to represent valence electrons.

4.3 Sodium Reacts with Chlorine Facts: Sodium reacts with chlorine to form sodium chloride (salt) according to the chemical equation: 2 Na + Cl2 → 2 NaCl (Na+ and Cl- ions) Theory: Chlorine takes an electron from sodium; forming the chloride and sodium ions.

The sodium and chloride ions are pulled together by coulombic forces, forming what is called an ionic bond.

Although the names of atoms and the ions are similar, the atom and the ion are quite different chemically, so it is important that we make distinctions and use precise terminology.

4.4 Using Lewis Symbols: More Ionic Compounds Active metal–active nonmetal compounds form ionic bonds.

Generally speaking, metallic elements in groups 1A and 2A react with nonmetallic elements in groups 6A and 7A to form ionic compounds. These are stable crystalline solids.

Atoms gain or lose electrons to have an octet of electrons, similar to the noble gas configurations.

4.5 Formulas and Names of Binary Ionic Compounds Generally the crossover method can be used to get the correct formula for a binary compound, write each ion with its charge (positive ion to the left), then transpose the charge numbers (but not the plus and minus signs) and write them as subscripts 4.6 Covalent Bonds: Shared Electron Pairs Generally nonmetal-nonmetal compounds form covalent bonds.

The bond formed by a shared pair of electrons is called a covalent bond Formation of double and triple bonds maintain octets

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4.7 Unequal Sharing: Polar Covalent Bonds In H-Cl the chlorine atoms have a greater attraction for a shared pair of electrons than hydrogen atoms do; chorine is said to be more electronegative than hydrogen.

Electronegativity increases as you move towards fluorine.

4.8 Polyatomic Molecules: Water, Ammonia, and Methane In general, many nonmetals often form a number of covalent bonds equal to eight minus the group number.

Hydrogen forms 1 bond, oxygen forms 2 bonds, nitrogen forms 3 bonds, carbon forms 4 bonds.

4.9 Polyatomic Ions Table 4.4 lists the common polyatomic ions to be discussed 4.10 Rules for Writing Lewis Formulas The five steps for writing electron dot formulas appear in this section 4.11 Odd-Electron Molecules: Free Radicals An atom or molecule with unpaired electrons is called a free radical Examples are ClO2, NO, NO2 4.12 Molecular Shapes: The VSEPR Theory Table 4.6 lists bonding and the shapes of molecules 4.13 Shapes and Properties: Polar and Nonpolar Molecules Water, ammonia, and methane are revisited 4.14 A Chemical Vocabulary

Transparencies T-27

page 90

T-28 T-29 T-30 T-31 T-32

Table 4.1 Figure 4.1 Figure 4.2 Figure 4.4 Figure 4.5

T-33 T-34

Figure 4.6 Figure 4.7

T-35 T-36 T-37 T-38

Table 4.5 Figure 4.8 Table 4.6 Figure 4.9 Figure 4.10 Figure 4.11

Shell representations: Na loses an electron; Cl gains an electron Electron dot (Lewis) symbols for selected elements Sodium reacts with chlorine to form sodium chloride Sodium chloride crystal Periodic relationships of some simple ions Pauling electronegativity values of several common elements Representation of polar hydrogen chloride Covalent bonding of representative elements of the periodic table Number of bonds formed by selected elements Shapes of molecules Bonding and the shapes of molecules The methane molecule (and) The ammonia molecule (and) The water molecule

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Demonstrations 1. Use a conductivity apparatus (with light bulb or bell) to test various ionic and covalent compounds.

2. Use balloons to demonstrate the shapes of atomic orbitals. (Fowles, G.W.A., J. Chem.

Educ. 1955, p. 260.)

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Chapter 5

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Chapter 5. Chemical Accounting Overview Chapter 5 explains chemical nomenclature, mass relationships, molar masses, and the gas laws. Nomenclature of metal-nonmetal compounds, nonmetal-nonmetal compounds, polyatomic ions, and binary ionic compounds is discussed. The chapter concludes with balancing chemical equations and mole-mass relationships.

Chapter Outline 5.1 Chemical Sentences: Equations Remember that matter is neither created nor destroyed in chemical reactions and chemical equations and the balancing of chemical equations makes sense.

5.2 Volume Relationships in Chemical Equations 5.3 Avogadro’s Number and the Mole 5.4 Molar Mass: Mole-to-Mass and Mass-to-Mole Conversions 5.5 Solutions Transparencies T-39

Figure 5.2

T-40

Figure 5.3 Figure 5.4 Figure 5.5 Figure 5.6 Figure 5.7

T-41 T-42 T-43

Balancing the chemical equation for the reaction between hydrogen and oxygen Gay-Lussac’s law of combining volumes and Avogadro's explanation of Gay-Lussac’s law One mole of several substances Mole and mass relationships for CO2 Converting from mass of one substance to another

Demonstrations 1. Display 1 mole samples of some common elements and compounds. Stoppered flasks make suitable containers for both display and storage. Use 50 mL flasks for 12 g of carbon, 32 g of sulfur, 207 g of lead, 27 g of aluminum, 18 g of water, 58.5 g of salt (sodium chloride), etc. Use a 250 mL flask for 180 g of glucose, and a 500 mL flask for 342 g of sucrose.

2. Put a little water in an empty soda can and heat it to boiling. Lift the can with tongs, and quickly invert it in a beaker of cold water. Air pressure causes the can to collapse.

3. Fill a gallon milk jug (or a 2-liter soda bottle) with very hot water. Then pour out the water and quickly stopper the bottle. The bottle collapses on cooling.

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Chapter 6

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Chapter 6. Gases, Liquids, Solids, and Intermolecular Forces Overview Chapter 6 explains the characteristics of solids, liquids, and gases. The chapter begins with an examination of the three phases of matter, then discusses the forces between molecules, and finishes with a discussion of the ideal gas law.

Chapter Outline 6.1 Solids, Liquids, and Gases 6.2 Ionic Bonds 6.3 Forces Between Molecules 6.4 Forces in Solutions 6.5 Gases: The Kinetic-Molecular Theory 6.6 The Simple Gas Laws Pressure is inversely proportional to volume.

Volume is directly proportional to temperature.

Volume is directly proportional to the number of moles.

6.7 The Ideal Gas Law PV = nRT is the ideal gas law Transparencies T-44 T-45 T-46 T-47 T-48 T-49 T-50 T-51

Figure 6.1 Figure 6.3 page 000 Figure 6.4 Figure 6.5 Figure 6.9 Figure 6.11 Figure 6.12 Figure 6.13 Figure 6.15

Intermolecular and intramolecular forces Changes in state of matter on heating or cooling Electrolytes in solution Dipole-dipole forces in a solid and in a liquid (and) Dispersion forces Sodium chloride dissolving in water A kinetic theory view of Boyle’s law A graphic representation of Boyle’s law A graphic representation of Charles’ law Molar volume of a gas

Demonstrations 1. Display 1 mole samples of some common elements and compounds. Stoppered flasks make suitable containers for both display and storage. Use 50 mL flasks for 12 g of carbon, 32 g of sulfur, 207 g of lead, 27 g of aluminum, 18 g of water, 58.5 g of salt (sodium chloride), etc. Use a 250 mL flask for 180 g of glucose, and a 500 mL flask for 342 g of sucrose.

2. Put a little water in an empty soda can and heat it to boiling. Lift the can with tongs, and quickly invert it in a beaker of cold water. Air pressure causes the can to collapse.

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3. Fill a gallon milk jug (or a 2-liter soda bottle) with very hot water. Then pour out the water and quickly stopper the bottle. The bottle collapses on cooling.

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Chapter 7

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Chapter 7. Acids and Bases Overview Acids and bases are important to us in many ways. We consume products that are acids or bases on a daily basis. The balance of acids and bases is crucial for the life processes. Chapter 7 defines acids and bases, distinguishes between strong and weak acids and bases, introduces the pH scale, and discusses some acid-base consumer issues.

Chapter Outline 7.1 Acids and Bases: Experimental Definitions 7.2 Acids, Bases, and Salts 7.3 Acidic and Basic Anhydrides 7.4 Strong and Weak Acids and Bases Strong acids and bases undergo 100% dissociation. There are six strong acids: H2SO4, HClO4, HNO3 , HCl, HBr, HI 7.5 Neutralization 7.6 The pH Scale pH of 7 is neutral; pH < 7 is acidic, pH > 7 is basic 7.7 Acid Rain 7.8 Antacids: A Basic Remedy 7.9 Acids and Bases in Industry and in Us Transparencies T-52 T-53 T-54 T-55 T-56 and

T-57

Figure 7.1 Table 7.1 Table 7.2 Figure 7.2 Figure 7.3 Figure 7.4

Some common acids, bases, and salts Some familiar acids Common bases Litmus in acid and in base An acid is a proton donor. A base is a proton acceptor Metal oxides react with water to form hydroxide ions

Figure 7.5

Ammonia is a base because it accepts a proton from water The pH scale

Figure 7.7

Demonstrations 1. Place 6 or 8 small beakers on the overhead projector. Put a little (10 mL) dilute acid in half of the beakers and a little dilute base in the other half. Into each acid-base pair add several drops of some selected indicator (e.g., methyl orange, bromocresol green, methyl violet, phenol red, etc.) to show the variety of color changes exhibited by acidbase indicators.

2. Some indicators exhibit more than one color change. In three small beakers on the overhead projector place small amounts of dilute acid, water, and dilute base. Then add a drop of thymol blue indicator to each sample. The colors will be red, yellow, and

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blue.

3. On the overhead projector place four beakers containing small amounts of water, 0.01 M HCl, 1 M HCl, and 6 M HCl. To each sample add a few drops of methyl violet indicator solution. The colors will be violet, blue, green, and yellow.

4. On the overhead projector place six small beakers containing 5-10 mL of the following: 6 M NaOH, 1 M NaOH, tap water, distilled water, 1 M HCl, and 6 M HCl. To each sample add a few drops of bromothymol blue indicator. The colors will be violet, blue, green, yellow, orange, and red. (If the tap water turns blue, add a little more indicator solution.)

5. Place some flower petals in a beaker with about 1/4 cup of rubbing alcohol. (Almost any kind of flower petals can be used.) Use a stirring rod to mash up the petals in the alcohol. (Sometimes it helps to heat the mixture for a few minutes.) Pour off the clear colored extract, filtering it if necessary, and pour into several small beakers. Add vinegar to one, ammonia to another, and other acids or bases to the rest.

6. Take several red cabbage leaves and place them in a blender along with about a cup of water. Blend the mixture until the cabbage has been chopped into fine pieces.

Pour off the liquid, using a tea strainer to get rid of the very small bits of solid material.

Pour some of the deep-colored liquid into each of five glasses or beakers. Into each container pour one of the following: vinegar, bottled soda water, tap water, baking soda dissolved in water, and clear ammonia. The colors will be red, pink, violet, blue, and green.

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Chapter 8

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Chapter 8. Oxidation and Reduction Overview The process of oxidation and reduction occurs in and around us every second of every day. Chapter 8 begins by defining oxidation-reduction reactions, then discusses some important types of oxidation-reduction reactions. The chapter concludes with a brief overview of the importance of oxidation-reduction reactions to living organisms.

Chapter Outline 8.1 Oxidation and Reduction: Three Views Oxidation is the gaining of oxygen atoms. Reduction is the loss of oxygen atoms.

Oxidation is the loss of hydrogen atoms. Reduction is the gaining of hydrogen atoms.

Oxidation is the loss of electrons. Reduction is the gaining of electrons 8.2 Oxidizing and Reducing Agents The substance oxidized is the reducing agent.

The substance reduced is the oxidizing agent.

8.3 Electrochemistry: Cells and Batteries The passage of the electrons of an oxidation-reduction reaction through a circuit yields energy.

8.4 Corrosion Loss of useful metals through the process of oxidation is a costly but unavoidable process.

8.5 Explosive Reactions An explosion is a very rapid oxidation-reduction reaction.

8.6 Oxygen: An Abundant and Essential Oxidizing Agent Oxygen is the most importance oxidizing agent. Oxygen exists as O2 and constitutes 21% of the earth’s atmosphere.

8.7 Other Common Oxidizing Agents 8.8 Some Reducing Agents of Interest 8.9 A Closer Look at Hydrogen Hydrogen is found in nature as H2. Hydrogen reacts violently with oxygen to form water (O2 + 2 H2  2 H2O). This is the reaction that powers the space shuttle’s main engines.

8.10 Oxidation, Reduction, and Living Things Plants use solar energy and reduce carbon dioxide to form complex carbohydrates. Animals oxidize complex carbohydrates for energy.

Transparencies T-58 T-59 T-60 T-61

Figure 8.2 Figure 8.4 Figure 8.5 Figure 8.6

Three different views of oxidation and reduction Reaction of copper with silver nitrate A simple electrochemical cell Cross section of a zinc-carbon cell ..

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T-62

Figure 8.7 Figure 8.8

T-63

Figure 8.10

A lead-acid cell The corrosion of iron requires water, oxygen, and an electrolyte Hydrogen and helium in lighter-than-air craft

Demonstrations 1. Using tongs, hold a strip of magnesium metal in a flame. The oxidation reaction gives off an intense white light.

2. Place a small pile of iodine crystals mixed with zinc dust in the middle of a Pyrex watch glass. The elements do not react. Then add a little water with a medicine dropper.

The reaction to produce zinc iodide proceeds with the evolution of heat and purple vapor from the sublimation of iodine.

3. Mix a little shampoo into some fresh 3% hydrogen peroxide and then add a little MnO2. The oxygen gas given off from the oxidation of the hydrogen peroxide produces foam.

4. To a small pile of fine crystals of potassium permanganate in a Pyrex watch glass add a little glycerol. The reaction starts slowly but produces a lot of white smoke and a violet flame.

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Chapter 9

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Chapter 9. Organic Chemistry Overview Chapter 9 discusses the chemistry of the carbon atom, the only element that has a major field of chemistry devoted to the study of its compounds. Of the tens of millions of known chemical compounds, more than 95% are compounds of carbon. The types of hydrocarbons are discussed, and various functional groups are introduced.

Lecture Outline 9.1 The Unique Carbon Atom Carbon can form an almost infinite number of molecules of various shapes, sizes, and compositions.

Hydrocarbons: Compounds only containing H and C, often called saturated hydrocarbons.

9.2 Alkanes Alkanes are hydrocarbons with only single carbon-to-carbon bonds.

Formulas and names for the first 10 hydrocarbons and their isomers are discussed 9.3 Cyclic Hydrocarbons: Rings and Things Cyclic hydrocarbons are ring or loop structures.

9.4 Unsaturated Hydrocarbons: Alkenes and Alkynes Alkenes are hydrocarbons with at least one double carbon-to-carbon bond.

Alkynes are hydrocarbons with at least one triple carbon-to-carbon bond.

9.5 Aromatic Hydrocarbons: Benzene and Relatives Aromatic hydrocarbons have delocalized electrons.

9.6 Chlorinated Hydrocarbons: Many Uses, Some Hazards Chlorofluorocarbons (CFCs) are one type of chlorinated hydrocarbon.

9.7 The Functional Group Functional groups determine the chemistry of the organic molecule.

Table 9.4 shows selected organic functional groups.

9.8 The Alcohol Family R-OH, such as ethanol CH3CH2OH OH 9.9 Phenols (an aromatic hydrocarbon and alcohol) 9.10 Ethers R-O-R’ 9.11 Aldehydes and Ketones O

Aldehyde R–CHO R

C

H

O

Ketone – a carbon chain with a carbon - oxygen double bond R 9.12 Carboxylic Acids R-COOH 9.13 Esters: The Sweet Smell of RCOOR’

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9.14 Amines and Amides Amine- NH2 Amide- oxygen double bonded to a carbon with a nitrogen next to the carbon.

9.15 Heterocyclic Compounds: Alkaloids and Others Compounds with atoms other than carbon in the ring are heterocyclic.

Transparencies T-64 T-65 T-66 T-67 T-68 T-69

Figure 9.1 Figure 9.2 Figure 9.3 Figure 9.6 Figure 9.7 Table 9.4 Figure 9.10 Figure 9.11 Figure 9.12

Methane and ethane Propane Butane and isobutane Ethylene Acetylene Selected organic functional groups Methanol Ethanol Formaldehyde, acetaldehyde, acetone

Demonstrations 1. On the overhead projector place small beakers containing about 5-10 mL samples of an alkane and an alkene (e.g., cyclohexane and cyclohexene). Add one dropperful of dilute KMnO4 solution to each sample. The purple permanganate will remain unreacted in the alkane, but the alkene will be oxidized by the permanganate, and the purple color will disappear.

2. Place a small beaker containing a little ethyl alcohol on the stage of an overhead projector. Add some acidic dichromate solution and note the darkening of the mixture.

This is the reaction on which the “breathalyzer” test is based.

3. Place a little alcohol in a gallon milk jug, cap the bottle, and shake vigorously. Uncap the bottle, pour out the alcohol, and then replace the cap on the “empty” bottle until ready for demonstration. Hold a flame close to the mouth of the opened bottle and note the violent combustion.

4. Demonstrate a “carbide cannon” in which calcium carbide powder is reacted with water to form acetylene, which is then ignited by a spark.

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Chapter 10

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Chapter 10. Polymers: Giants Among Molecules Overview Chapter 10 introduces the student to the marvelous variety of materials made from polymers. The chapter starts by discussing the history of polymers, then discusses the various types of polymers and their uses. The chapter concludes with a look into the future developments and uses of polymers.

Lecture Outline 10.1 Polymerization: Making Big Ones Out of Little Ones Polymers are huge molecules made from many small molecules, monomers.

10.2 Natural Polymers Wood, wool, and proteins are all natural polymers.

10.3 Celluloid: Billiard Balls and Collars The synthetic material celluloid is derived from natural cellulose Celluloid is used to make billiard balls and historically movie film, and to make collars stiff.

10.4 Polyethylene: From the Battle of Britain to Bread Bags Polyethylene is the simplest and cheapest polymer made from the monomer ethylene (CH2=CH2).

10.5 Addition Polymerization: One + One + One . . . Gives One! Polypropylene, polystyrene, vinyl polymers, and Teflon are examples of addition polymers, which are made by adding monomers together to form a macromolecule. Table 10.2 shows some addition polymers.

10.6 Rubber and Other Elastomers Natural rubber is isoprene, a naturally occurring polymer.

Vulcanization is the addition of sulfur to isoprene to make the substance harder.

Synthetic rubber, SBR, is a copolymer of styrene and butadiene.

10.7 Condensation Polymers: Splitting Out Water Water is eliminated, and the monomers are added together to form the polymer.

Nylon, polyester, and Bakelite are examples of condensation polymers.

10.8 Properties of Polymers In crystalline polymers the monomers are in line, giving high tensile strength.

In amorphous polymers the monomers are randomly ordered.

The glass transition temperature is an important property of polymers.

10.9 Disposal of Plastics Plastics last almost forever. Landfills, incineration, biodegradation, and recycling are ways of disposing of plastics.

10.10 Plastics and Fire Hazards Fumes given off during house fires are very dangerous owing in large part to the polymers in the modern home.

10.11 Plasticizers and Pollution Plasticizers make some plastics, such as vinyl polymers, more flexible.

10.12 Plastics and the Future

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Artificial lungs and hearts made of plastic will be medical uses in the future.

Transparencies T-70 T-71 T-72

Figure 10.1 Table 10.1 Figure 10.4

Models of a polyethylene molecule Some addition polymers Vulcanization of rubber cross-links the molecular chains

Figure 10.5 Figure 10.6

Vulcanization and elasticity Organization of polymer molecules. (a) Crystalline arrangement (b) Amorphous arrangement

(and) T-73

Demonstrations 1. Onto a Petri dish about half-full of acetone, place a cup made of polystyrene foam. The cup will appear to melt away as the polystyrene dissolves in the acetone.

2. Beads of polystyrene can be mixed with beads of polyethylene and then separated by pouring the mixture into a beaker filled with water. The polyethylene will float, while the polystyrene will sink.

3. A sample of high density polyethylene (HDPE) can be separated from a sample of low density polyethylene (LDPE) using a 1:1 mixture of methanol and water. The LDPE will float, while the HDPE will sink. (The HDPE sample can be cut from a gallon milk jug. The LDPE sample can be cut from a waxy squeeze bottle. The samples might be cut in the shapes of an H and an L, respectively.)

4. Insert a copper wire into a cork, and holding the cork heat the copper wire in a flame.

Touch the hot wire to a piece of plastic and return the wire to the flame. If the flame turns green, the plastic contains a halogen. It is likely to be polyvinyl chloride (PVC). (This is the Beilstein test.)

5. If you hold a plastic foam egg carton in a flame, the flame becomes very sooty because the foamed plastic is polystyrene, which has a very high carbon-to-hydrogen ratio.

6. If you hold a piece of polyethylene in a flame, the odor will be like that of a burning candle.

7. Make “slime” by adding about 5 mL of saturated borax solution to 30 mL of 4% poly(vinyl alcohol) solution in a paper cup. Stir vigorously with a wooden stick until the mixture forms a gel. Remove the gel from the cup and knead it into a ball. When laid on a flat surface, the ball flattens. When pulled apart very slowly, it stretches. When pulled apart rapidly, it breaks. The “slime” can usually be stored in an air-tight plastic bag for at least a week. After a while it will start to mold and should be thrown away.

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Chapter 10

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8. Another polymer gel, similar to “slime,” can be made from a white glue, such as Elmer’s. Mix about 30 mL of the glue with 30 mL of water and mix well. (A few drops of food color can be added, if desired.) Then add about 20 mL of 4%, or saturated, borax solution and stir well until the mixture forms a semi-solid mass on the end of the stirrer.

Remove the gel and form it into a ball. Try bouncing the ball. The polymer can be treated and stored like “slime.” (Adding a few drops of Lysol to the borax solution will slow down the mold growth.)

9. Polystyrene is a glassy polymer, whereas polyethylene is crystalline. When you step on a transparent polystyrene cup, it shatters. When you step on a cup or bottle made of polyethylene, it bends but does not break.

10. Have six or more students stand before the class, an arm’s length apart from each other, with their hands clasped about 6 inches in front of their chests. They represent ethylene molecules. “Initiate the reaction” by unclasping the hands of the student on the far left, placing an object such as an eraser, in her right hand. Her left hand then “attacks” the next “molecule,” grabbing his right hand and leaving his left hand free to “attack” the next one, in a “chain reaction.” At the end of the line, the last student has a free left hand, which you can grab as a “chain stopper.”

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Chapter 11

1

Chapter 11. Nuclear Chemistry Overview Chapter 11 discusses the structure of the nucleus, radioactivity, half-life, artificial transmutation, and induced radioactivity. These concepts are then applied to nuclear medicine, radioisotopic dating, the nuclear bomb, and nuclear power.

Lecture Outline 11.1 Natural Radioactivity Background Radiation The two particles of the nucleus are the proton (+) and the neutron (neutral).

Two atoms with the same number of protons but different number of neutrons are isotopes.

11.2 Nuclear Equations The equations are simple if you follow the rule that both mass and charge are conserved. That is, the sums of the nucleons and nuclear charges must be the same before and after a reaction. Another way to think about the problem is that the sum of the superscripts on the left side of the reaction arrow must equal the sum of the superscripts on the right side of the arrow. The same rule holds for the subscripts.

11.3 Half-Life The time it takes for half of a radioactive substance to disintegrate.

11.4 Radioisotopic Dating Use of the half-lives of radioactive isotopes and assumption of constant decay rates over time makes it possible to determine the age of items.

11.5 Artificial Transmutation Bombardment of various nuclei with protons, neutrons, alpha particles, or other fundamental particles brings about transmutation, each element being changed into another.

11.6 Uses of Radioisotopes The main use of radioisotopes is as tracers. Because radioisotopes are radioactive, their locations can be traced.

Radioactive isotopes are used for diagnostic tests as well as for therapy.

More people have had their lives improved by the use of radioactive materials than have been killed or injured!

11.7 Penetrating Power of Radiation Gamma > Beta > Alpha 11.8 Energy from the Nucleus E = mc2 Binding Energy 11.9 The Building of the Bomb Chain Reactions The highly secret Manhattan Project was set up to do this.

11.10 Radioactive Fallout 11.11 Nuclear Power Plants

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Chapter 11

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Nuclear power plants cannot explode! To explode, the U-235 must be enriched to 90%. Nuclear power plants use uranium that is only 3% U-235.

11.12 Thermonuclear Reactions Thermonuclear reactions are those involving nuclear fusion. These are the reactions that power the sun and all other stars.

11.13 The Nuclear Age Transparencies T-74

Figure 11.1

T-75

Figure 11.2 Figure 11.3

Most of our exposure to ionizing radiation comes from natural sources Nuclear emission of an alpha particle and a beta particle (and) Nuclear change accompanying (a) positron

emission T-76 T-77 T-78 T-79

Table 11.2 Table 11.4 Figure 11.4 Table 11.6 Figure 11.8

T-80 T-81 T-82

Figure 11.10 Figure 11.12 Figure 11.13

and (b) electron capture Nuclear change during radioactive decay (and) Subatomic particles Radioactive decay of iodine and half-life Some radioisotopes used in medicine The relative penetrating powers of alpha, beta, and gamma radiation Variation of nuclear stability with mass number Schematic representation of a nuclear chain reaction The “Little Boy” atomic bomb

Demonstrations 1. Demonstrate the use of a Geiger counter. Using a compound of thorium or uranium, show the large amount of radioactivity in such a compound. (If you can find an old piece of orange dinnerware, it can be compared with a piece of some other color. The old orange pieces were glazed with uranium oxide and are quite radioactive.)

2. Using the alpha, beta, and gamma emitters that came with the Geiger counter (or can be purchased separately), show that a sheet of paper will stop alpha rays, and a sheet of aluminum will stop alpha or beta rays, but neither of these will stop gamma radiation. A thin piece of lead foil will readily block out alpha or beta radiation, but it takes several cm of lead to block out gamma radiation.

3. Show (qualitatively) that as you move the detector away from the source of radiation, the count goes down significantly. (If you move just 3 feet away from a radioactive source, you get only 1/9 as much radiation.)

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Chapter 12

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Chapter 12. Chemistry of the Earth Overview Chapter 12 discusses the components of the earth and the use that modern civilization makes of the Earth’s constituents. The chapter begins by describing the substances in the earth’s crust and how they formed. Next, the chapter discusses uses and abuses modern civilization makes of the Earth’s resources.

Lecture Outline 12.1 Spaceship Earth: The Materials Manifest The earth is mostly oxygen, silicon, and hydrogen.

12.2 The Lithosphere: Organic and Inorganic Organic: coal, petroleum, and natural gas Inorganic: silicates, carbonates, oxides, and sulfides 12.3 Meeting Our Needs: From Sticks to Bricks 12.4 Silicates and the Shapes of Things The silicates are quartz, mica, and asbestos.

12.5 Modified Silicates: Ceramics, Glass, and Cement 12.6 Metals and Ores The most important metals are copper, iron, and aluminum.

12.7 Earth’s Dwindling Resources Transparencies T-83 T-84 T-85 T-86 T-87

Figure 12.2 Figure 12.3 Figure 12.4 Figure 12.5 page 330 page 330 Figure 12.6 Figure 12.9

The silicate tetrahedron (and) The chemical structure of quartz Sheets of mica and the chemical structure of mica Chrysotile asbestos and its chemical structure Maglev train Aerogel A modern blast furnace Composition of municipal wastes in the United States

Demonstrations 1. Make a “super ball” by adding 6-8 mL of ethyl alcohol to 25 mL of sodium silicate in a paper cup. Stir with a wooden stick until the mixture forms a solid material. Wearing disposable plastic gloves, press the solid between the palms of the hands and roll it into a ball.

2. Using a plain or a quartered petri dish, place four pieces of rock (e.g., limestone, marble, sandstone, and granite) in the dish. With a dropper, add a little 3 M HCl to each rock. If bubbles form, the rock is probably calcium carbonate. Limestone and marble are two of the many forms of calcium carbonate.

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Chapter 13

1

Chapter 13. Air Overview Chapter 13 begins by defining the layers and composition of the earth’s atmosphere. The chapter then moves on to discuss the dynamic nature of the atmosphere as seen through the nitrogen and oxygen cycles. The chapter concludes with detailed discussion about various air pollutants.

Lecture Outline 13.1 Earth’s Atmosphere: Divisions and Composition Air is composed of 78% nitrogen.

13.2 Chemistry of the Atmosphere The majority of nitrogen fixation is performed by bacteria.

13.3 Pollution Through the Ages 13.4 Coal + Fire → Industrial Smog The word smog comes from smoke and fog.

13.5 Automobile Emissions In urban areas, automobile contribution to smog can exceed 90%.

13.6 Photochemical Smog: Making Haze While the Sun Shines Photochemical smog is produced when it is sunny and warm.

13.7 Acid Rain: Air Pollution → Water Pollution Acid rain is rain having a pH less than 5.6.

13.8 The Inside Story: Indoor Air Pollution Indoor air pollution poses a major health concern.

13.9 Ozone: The Double-Edged Sword Ozone shields humans from UV light.

13.10 Carbon Dioxide and the Greenhouse Effect The greenhouse effect is trapping of reradiated solar rays by the atmosphere.

13.11 Who Pollutes? Who Pays?

Air pollution causes damage to buildings, clothing, and other objects.

Transparencies T-88 T-89 T-90 T-91 T-92

Figure 13.1 Figure 13.2 Figure 13.3 Figure 13.5 Figure 13.7

T-93 T-94 T-95

Figure 13.10 Figure 13.11 Figure 13.12 Figure 13.13

The atmosphere The nitrogen cycle The oxygen cycle Electrostatic precipitator Some chemical processes in the formation of photochemical smog Formation of acid rain The ozone hole The greenhouse effect (and) Carbon dioxide content of the atmosphere and surface temperature variations

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Chapter 13

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Demonstrations 1. Dip some filter paper (or white paper toweling) into a solution of cobalt(II) chloride, and let it dry. The paper will be blue when it is dry, but when the atmosphere is humid, it will cause the paper to turn pink. Try holding the paper over a steaming cup of coffee or boiling water.

2. (Air pressure) Pour about 10 mL water into an empty soda can and heat it to boiling.

Then with tongs quickly invert the can into a container of cool water. The pressure of the air crushes the can.

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Chapter 14

1

Chapter 14. Water Overview The chapter begins by discussing the special properties of water as compared with other substances. Next, the water cycle, biological and chemical contamination, and water treatment methods are introduced.

Lecture Outline 14.1 Water: Some Unique Properties Water has a high dipole moment, a high heat of vaporization, a high heat capacity, and a low heat of fusion.

14.2 Water in Nature Saltwater constitutes nearly 98% of the water on the earth.

14.3 Chemical and Biological Contamination Fewer than 10% of the people of the world have access to sufficient clean water.

14.4 Industrial Use 14.5 Groundwater Contamination: Tainted Tap Water Chemical dumps, garbage dumps, underground storage tanks, and overuse of fertilizers are all potential sources of groundwater contamination.

14.6 Making Water Fit to Drink 14.7 Wastewater Treatment Plants Drinking water comes from lakes, reservoirs, wells, and aquifers.

14.8 The Newest Soft Drink: Bottled Water 14.9 Water Pollution and the Future There are between 4000 and 40,000 cases of waterborne illnesses in the United States each year.

Transparencies T-96 T-97 T-98 T-99

Figure 14.2 Figure 14.4 Figure 14.5 Figure 14.6 Figure 14.7

The hydrological (water) cycle Groundwater likely to be contaminated by nitrates Municipal water purification plant Primary sewage treatment plant (and) Secondary sewage treatment plant

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Chapter 14

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Demonstrations 1. (High heat of vaporization of water) Mix ordinary rubbing alcohol with an equal volume of water. Dip a piece of cloth (such as a handkerchief) into the alcohol mixture, gently squeeze out most of the excess liquid, and then using tongs, brush the cloth past a flame so that it catches fire. Wave the burning cloth slowly through the air to show that it really is burning. When the flame goes out, hold up the cloth to show that it is still intact.

Let a few students feel the cloth to see how cool it is. The heat from the burning alcohol was used up in vaporizing the water.

2. (High surface tension of water) Fill a container several inches deep with water, and float a plastic strawberry basket on the surface. The basket is actually heavier than water, but the surface tension of the water keeps it afloat. Adding a few drops of detergent lowers the surface tension of the water, and the basket sinks. (This demonstration can be done with a floating needle on the stage of an overhead projector.)

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Chapter 15

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Chapter 15. Energy Overview Chapter 15 begins by discussing the sources of energy, and the laws of thermodynamics. Next, the chapter discusses the energy requirements and costs of a technologically advanced culture.

Lecture Outline 15.1 Heavenly Sunlight Flooding Earth with Energy The sun is the source of energy on planet Earth.

15.2 Energy and Chemical Reactions Chemical reactions depend on several factors and are either endothermic or exothermic.

15.3 The Laws of Thermodynamics 15.4 People Power: Early Uses of Energy 15.5 Reserves and Consumption Rates of Fossil Fuels Coal, petroleum, and natural gas 15.6 Coal: The Carbon Rock of Ages 15.7 Natural Gas: Mostly Methane 15.8 Petroleum: Liquid Hydrocarbons 15.9 Convenient Energy: Electricity Electricity is a secondary energy form; it must be produced from a primary energy source.

15.10 Nuclear Fission Fission, although producing abundant energy without adding to the greenhouse effect, has several problems; fuel shortages and waste disposal are two major problems.

15.11 Nuclear Fusion: The Sun in a Magnetic Bottle Fusion power is the promise of the future, A clean, efficient, and unlimited fuel Supply.

15.12 Harnessing the Sun: Solar Energy 15.13 Biomass: Photosynthesis for Fuel 15.14 Hydrogen: Light and Powerful 15.15 Other Renewable Energy Sources 15.16 Energy: How Much Is Too Much?

Transparencies T-100 T-101 T-102 T-103 T-104

Figure 15.2 Figure 15.6 Figure 15.7 Figure 15.9 Figure 15.10

T-105

Figure 15.11

Fuel reserves and consumption rates The carbon cycle The fractional distillation of petroleum Coal-burning power plant Percentages of electricity generation in the U.S. from various energy sources Nuclear power plant

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Chapter 15

T-106 T-107 T-108

2

Figure 15.14 Figure 15.15 Figure 15.16 Figure 15.17

Solar collector for heating Silicon crystals used in solar cells (and) Diagram of a solar cell A hydrogen–oxygen fuel cell

Demonstrations 1. Place several tablespoonfuls of calcium oxide (quicklime) in a plastic zip-lock bag.

Add water (and some red food color, if desired). The bag will get hot and can be passed around.

2. Place some ammonium nitrate or ammonium acetate in a zip-lock bag. Add water (and some blue food color, if desired). The bag will get cold and can be passed around.

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Chapter 16

1

Chapter 16. Biochemistry: A Molecular View of Life Overview Chapter 16 discusses the chemistry of life, from the substances required for energy to the molecules that define the structure of life.

Lecture Outline 16.1 The Living Cell An introduction to the major parts of the cell, structure, and energy production are discussed, including the cell membrane, the nucleus, mitochondria, and chloroplasts.

16.2 Carbohydrates: A Storehouse of Energy Carbon and water (“carbo hydrate”) are the basis for the structures of the simple sugars to the complex carbohydrates, starches and cellulose.

16.3 Fats and Other Lipids Fats are esters of the fatty acids. Fats may be saturated, with no carbon-to-carbon double or triple bonds, or unsaturated, with one or more carbon-to-carbon double or triple bond(s).

16.4 Proteins: Polymers of Amino Acids Muscle, skin, and hair are all mostly proteins. Here the structure and properties of amino acids are discussed.

16.5 The Peptide Bond: Peptides and Proteins The amine linkage is a peptide bond. Polymers of amino acids create a protein.

The 20 common names and importance of amino acid sequence are discussed.

16.6 Structure of Proteins The four levels of proteins are primary, secondary, tertiary, and quaternary.

16.7 Enzymes: Exquisite Precision Machines Enzymes are large protein molecules that serve a specific purpose.

16.8 Nucleic Acids: Parts and Structure Four simple base pairs define everything about you!

16.9 DNA: Self-Replication Replication is the copying or duplication of DNA molecules.

16.10 RNA: Protein Synthesis and the Genetic Code In transcription, the DNA is transcribed to mRNA. The mRNA travels to the ribosome and is then translated into a protein.

16.11 The Human Genome Various topics of genetic testing and research are discussed, from genetic testing and DNA fingerprinting to gene therapy and genetically modified organisms.

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Chapter 16

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Transparencies T-109 T-110 T-111 T-112 T-113 T-114 T-115 T-116 T-117 T-118 T-119 T-120 T-121 T-122 T-123

Figure 16.1 Figure 16.2 Figure 16.3 Figure 16.4 Figure 16.5 Figure 16.6 Figure 16.11 Table 16.3 Figure 16.12 Figure 16.13 Figure 16.14 Figure 16.15 Figure 16.18 Figure 16.20 Figure 16.22 Figure 16.23 Figure 16.25 Figure 16.28

An idealized “typical” plant cell An idealized “typical” animal cell Three common monosaccharides (and) Cyclic structures of glucose, galactose, and fructose Two disaccharides: sucrose and lactose (and) Two polysaccharides: starch and cellulose A short segment of a protein molecule The 20 amino acids in the genetic code Pleated sheet conformation of polypeptide chains Two representations of the α-helical structure of a polypeptide chain The tertiary structure of myoglobin (and) The quaternary structure of hemoglobin The induced-fit model for enzyme action The components of nucleic acids The DNA double helix The pairing of complementary bases Replication of DNA Recombinant DNA techniques

Demonstrations 1. (Test for starch) To several mL of a dilute solution of soluble starch add a drop of iodine in KI solution. The starch-iodine complex is dark blue. (This can be shown on the overhead projector.)

2. (Biuret test for proteins) To 5 mL of 1% albumin or gelatin in a small beaker add 5 mL of 2 M NaOH, and then add several drops of 0.5% copper sulfate. (Do this on the stage of an overhead projector.) A lavender color confirms the presence of peptide bonds.

3. (Enzyme action) Dissolve a little urea in about 10 mL of water in a small beaker. Add a little neutralized indicator solution (e.g., a yellow-green solution of bromthymol blue).

When a small amount of urease is added, the solution turns blue because of the formation of ammonia.

CO(NH2)2 + H2O → CO2 + 2 NH3 4. (Steroid test) Into several mL of CH2Cl2 (in a small beaker on the overhead projector) dissolve a few crystals of cholesterol. Then add 1 mL of acetic anhydride and about 5 drops of concentrated sulfuric acid. The solution will turn pink, then violet; then it will slowly turn blue, and finally green.

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Chapter 17

1

Chapter 17 Food: Molecular Gastronomy Overview The chapter discusses food groups, minerals, and vitamins. Next, food additives, good and bad, are examined. The chapter concludes with an examination of modern food production technology.

Lecture Outline 17.1 Carbohydrate in the Diet Dietary carbohydrates are the sugars and starches.

17.2 Fats and Cholesterol The fats contain abundant calories, but they are harmful in excess.

17.3 Proteins: Muscle and Much More Proteins are built from the amino acids.

17.4 Minerals: Inorganic Chemicals and Life The minerals are inorganic substances required for life. Table 17.3, on page 468, lists the required minerals.

17.5 The Vitamins: Vital, But Not All Are Amines The vitamins fall into two major categories; fat soluble and water soluble.

17.6 Other Essentials: Fiber and Water 17.7 Starvation and Fasting 17.8 Additives to Enhance Our Food Substances are added to food to make it taste better, look better, prevent spoilage, and add minerals and vitamins to the diet.

17.9 Additives to Retard Spoilage Toxins and carcinogens are the two major categories of food poisons.

17.10 Color Additives 17.11 Poisons in Our Food 17.12 A World Without Food Additives Transparencies T-124 T-125 T-126

Figure 17.7 Table 17.4 Figure 17.12 Figure 17.13

Models of fatty acids Some of the vitamins Some molecular flavorings (and) Artificial sweeteners

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Chapter 17

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Demonstrations 1. Fill a small beaker (100–150 mL) about 1/3 full of sugar (sucrose). Carefully add 5–10 mL of concentrated sulfuric acid, and step back. A black cylinder of foamed carbon will rise up out of the beaker and there will be smoke and fumes. The demonstration should be done in a hood or in a well-ventilated area.

2. (Corn starch putty) To 1 cup of corn starch slowly add about 1/3 cup water and mix.

The gooey mixture has some unusual properties. It can be cut with a knife or with scissors, but it flows like a viscous liquid. It looks wet, but if you rub it between your fingers, it gets powdery. It can be rolled into a ball, but then the ball will flatten. It is a non-Newtonian fluid.

3. Put several spoonfuls of baking soda into a glass soda bottle. Add about a cupful of vinegar, and then quickly spread the opening of a balloon over the mouth of the bottle.

The balloon will be blown up by the carbon dioxide generated. The balloon can be removed from the bottle at any time.

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Chapter 18

1

Chapter 18 Drugs: Chemical Cures, Comforts, and Cautions Overview Chapter 18 examines the world of pharmaceuticals, from aspirin to opiates. A drug is a chemical substance that affects the functioning of living things and is used to relieve pain, to treat illness, or to improve one’s state of health.

Lecture Outline 18.1 Scientific Drug Design Chemotherapy (chemical therapy) can be used to treat disease. There are several hundred over the counter (OTC) drugs available.

18.2 Pain Relievers: Nonsteroidal Anti-inflammatory Drugs (NSAIDS) The wonder drugs! Pain reliever, anti-inflammatory, anticoagulant, and fever reducer 18.3 Chemistry, allergies, and the Common Cold There is no cure for the common cold.

18.4 Antibacterial Drugs Effective against bacteria, not viruses!

18.5 Viruses and Antiviral Drugs Some, not all, can be prevented by vaccinations.

18.6 Chemicals Against Cancer Aim of cancer research is to kill the cancerous tissue without killing normal cells 18.7 Hormones: The Regulators Regulate growth, metabolism, reproduction 18.8 Chemistry and Social Revolution: The Pill Pill fools the body into thinking it is pregnant.

18.9 Drugs for the Heart High blood pressure is the most common cardiovascular disease 18.10 Drugs and the Mind Three classes of drugs: stimulant, depressant, and hallucinogen 18.11 Some Chemistry of the Nervous System Neurons make up the nervous system.

18.12 Brain Amines: Depression and Mania A chemical PEA could govern emotions of romantic relationships.

18.13 Anesthetics Anesthetics are depressants!

18.14 Depressant Drugs Alcohol is the most used and abused depressant drug.

18.15 Narcotics Relief of pain and narcosis are effects.

18.16 Antianxiety Agents Valium is the best known drug.

18.17 Stimulant Drugs

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Chapter 18

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Amphetamine was once used for weight loss, it is now used in treating mild depression and narcolepsy. Methamphetamine is inexpensive and widely abused.

18.18 Hallucinogenic Drugs These drugs change the way people see things.

Marijuana: This drug has a history in medicine, particularly in India.

18.19 Drug Problems Primary drug problem is illegal drugs Transparencies T-127 T-128 T-129 T-130 T-131 T-132 T-133 T-134

T-135

Figure 18.1 Figure 18.2 Figure 18.3 page 000 Figure 18.4 page 000 Figure 18.5 Figure 18.11 Figure 18.14 Figure 18.15 Figure 18.18 page 000 page 000 page 000 Figure 18.27

NSAIDs (and) COX-2 inhibitors Histamine and two antihistamines (and) Dextromethorphan and codeine Leading causes of death in the U.S. in 1900 and in 2002 Sulfanilamide and para-aminobenzoic acid (and) Some common penicillins Some steroids Diagram of a human nerve cell (and) Diagram of a synapse Models and structures of enantiomers Morphine (and) Hydrocodone (and) Receptor site for opioids β-Phenylethylamine and related compounds

Demonstrations [None]

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Chapter 19

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Chapter 19. Nutrition in Health and Fitness Overview In this chapter we examine nutrition, electrolyte balance, muscle action, drugs, and other matters of concern to those interested in their own well-being and physical fitness.

Lecture Outline 19.1 Calories: Quantity and Quality Most calories should come from carbohydrates. No more than 25% of calories should come from fats.

19.2 Vitamins and Minerals Table 19.1 is a table of recommended daily adult intakes of vitamins.

19.3 Body Fluids: Electrolytes.

Sodium, potassium, and chloride ions are the major essential electrolytes.

Water is the best drink when you are thirsty!

19.4 Weight-Loss Diets To lose a pound, you must burn off 3500 kcal.

19.5 Exercise for Weight Loss The healthy way to lose weight is to exercise. Not only will you lose weight, but you will also look and feel better.

19.6 Measuring Fitness Men should have more than 3% body fat; women need 10–12%.

19.7 Some Chemistry of Muscles Muscles consume ATP (adenosine triphosphate) for energy.

Aerobic muscle contraction occurs in the presence of oxygen.

Anaerobic muscle contraction occurs when sufficient oxygen is not available.

19.8 Drugs and the Athlete Restorative drugs are pain killers and anti-inflammatory drugs.

Stimulant drugs are used in an attempt to improve performance.

Anabolic steroids are used to build muscle mass.

19.9 Exercise and the Brain Some runners get a euphoric high during and after a long, hard run.

19.10 No Smoking Smoking is one of the worst things you can do to your physical well being.

Transparencies T-136 T-137 T-138

Figure 19.1 Table 19.1 Figure 19.4

The food pyramid Dietary reference intakes of vitamins and minerals Skeletal muscle tissue

Demonstrations [None]

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Chapter 20

1

Chapter 20. Chemistry Down on the Farm Overview This chapter discusses modern food production and associated technologies.

Fertilizers, pest control, herbicides and defoliants, and sustainable agriculture are all discussed from a conventional, modern farming techniques perspective to modern organic/alternative techniques.

Lecture Outline 20.1

20.2

20.3 20.4 20.5 20.6 20.7

Farming with Chemicals: Fertilizers Discuss how we use and produce the three primary plant nutrients, nitrogen, phosphorus, and potassium.

The War Against Pests Discuss how we control substances that kill organisms that we consider pests as well as how to handle pest resistance.

Biological Insect Controls Herbicides and Defoliants Discuss how we control weeds as well as how to handle weed resistance.

Sustainable Agriculture Discuss what is needed to produce food and fiber indefinitely Some Malthusian Mathematics Can We Feed a Hungry World?

Transparencies T-139 T-140 T-141 T-142

Figure 20.2 Figure 20.7 Figure 20.11 Figure 20.16

Flow of nutrients in a simple system Potassium and hydronium ions in the soil Concentration of DDT up the food chain Energy and agriculture

Demonstrations [None]

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Chapter 20

2

Answers to Critical Thinking Exercises The Critical Thinking Exercises are mostly open-ended. The answers given here are not necessarily definitive or complete.

20.1 An advertisement claims that a certain cereal from organically grown grains provides a “life energy” that does not exist in ordinary processed cereals.

1. Is the claim falsifiable? No. At present it is not possible to prove whether or not “life energy” exists at all.

2. Is the claim logical? No. Even if there is something called “life energy,” it is not likely that its presence in food would depend only on how the food was grown.

3. Is the claim reproducible? No. It is not at all clear what kind of difference one would be looking for in these two kinds of cereals.

4. Is the claim sufficient? No. Proof of the claim would be up to the claimant, who would need to define clearly what “life energy” is and to describe exactly how it is measured.

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Chapter 20

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20.2 An internet article claims that NO3 ions from decomposition of manure is better for plant growth than NO3 ions from synthetic fertilizers.

1. Is the claim falsifiable? Yes. Both forms of fertilizer could be compared.

2. Is the claim logical? No. All nitrate ions are the same, the manure may provide other nutrients.

3. Is the claim reproducible? Yes. 4. Is the claim sufficient? No. 20.3 A newspaper claims that certain genetically modified crops may themselves be weeds if not properly controlled.

1. Is the claim falsifiable? No. It is a statement of opinion.

2. Is the claim logical? No. 3. Is the claim reproducible? No. 4. Is the claim sufficient? No. 20.4 Vegetables grown in a back yard garden are more nutritious than those grown on commercial farms.

1. Is the claim falsifiable? Yes. Nutritional studies could be conducted.

2. Is the claim logical? Yes. 3. Is the claim reproducible? Yes. 4. Is the claim sufficient? Yes. 20.5 Nitrate ions from synthetic fertilizers such as ammonium nitrate are less effective plant nutrients than nitrate ions from the breakdown of organic fertilizers by soil bacteria, leading to nonorganic food being less nutritious.

1. Is the claim falsifiable? Yes. Both forms of fertilizer could be compared.

2. Is the claim logical? No. All nitrate ions are the same, the manure may provide other nutrients.

3. Is the claim reproducible? Yes. 4. Is the claim sufficient? No.

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Chapter 21

1

Chapter 21 Household Chemicals Overview The modern home is stocked with a variety of chemical products. Chapter 21 introduces the student to the chemistry of soaps, detergents, bleaches, paints, cosmetics, perfumes, and other chemical compounds used in and around the home.

Lecture Outline 21.1 A History of Cleaning 21.2 Fat + Lye  Soap Soap is a substance formed by mixing a base and a fatty acid.

21.3 Synthetic Detergents Modern detergents are made in the laboratory.

21.4 Laundry Detergent Formulations 21.5 Dishwashing Detergents 21.6 Fabric Softeners: Quaternary Ammonium Salts 21.7 Laundry Bleaches: Whiter Whites Chlorine kills bacteria and whitens clothes.

21.8 All-Purpose Cleaning Products 21.9 Special-Purpose Cleaners 21.10 Solvents, Paints, and Waxes 21.11 Cosmetics: Personal Care Chemicals 21.12 Toothpaste: Soap with Grit and Flavor It’s the same thing you wash your clothes with!

21.13 Perfumes, Colognes, and Aftershaves 21.14 Some Hairy Chemistry 21.15 The Well-Informed Consumer Transparencies T-143

Figure 21.2

T-144

Figure 21.3 Figure 21.4 Figure 21.5 Figure 21.6 Figure 21.9 Figure 21.14

Soap can be made by reacting fats with sodium hydroxide Sodium palmitate, a soap (and) The cleaning action of soap visualized Soap and detergent compared Water softening by ion exchange Optical brightener converts ultraviolet light to visible Disruption of hydrogen bonds by water when hair is wet

Figure 21.15

Disruption of salt bridges by a change in pH

T-145 T-146 T-147 T-147 (and)

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Chapter 21

2

Demonstrations 1. Into a tall glass place some baking soda, some liquid detergent, and a few drops of red food color. Adding vinegar to this mixture will cause it to bubble up and look somewhat like a strawberry ice cream soda.

2. Add several drops of food color to each of two beakers about half full of water. Add a little liquid chlorine bleach to one of the beakers, and watch the color disappear.

..

Chapter 22

1

Chapter 22. Poisons: Chemical Toxicology Overview Many compounds are toxic to humans. Chapter 22 discuses many toxic substances that are known, giving priority to those that are more likely to be encountered in everyday life, their effects, detection, identification, and their antidotes.

Lecture Outline 22.1 Natural Poisons Poisons around the home include insecticides, bleaches, and cleansers.

Azaleas, hydrangeas, holly berries, and wisteria seeds are poisons in the garden.

22.2 Corrosive Poisons: A Closer Look Three types of corrosive poisons are strong acids, strong bases, and strong oxidizing agents. Blood agents: compounds that block oxygen transport are blood agents.

22.3 Metabolic Poisons 22.4 Heavy Metal Poisons 22.5 More Chemistry of the Nervous System Nerve poisons: stopping the acetylcholine cycle.

Nerve poisons are the most deadly killing agents developed for military use.

22.6 The Lethal Dose How much does it take to kill?

22.7 The Liver as a Detox Facility 22.8 Chemical Carcinogens: Slow Poisons Substances shown to cause cancer are known as carcinogens.

22.9 Three Ways to Test for Carcinogens: The Ames test uses bacteria.

Animal testing uses test animals.

Epidemiological studies are a statistical approach.

22.10 Birth Defects: Teratogens 22.11 Hazardous Wastes.

Reactive wastes, flammable wastes, toxic wastes, and corrosive wastes 22.12 What Price Poisons?

Transparencies T-149 T-150

Figure 22.3 Figure 22.4 Figure 22.5

T-151 T-152 T-153

Table 22.1 Figure 22.7 Figure 22.10

Mercury poisoning The acetylcholine cycle (and) Acetylcholinesterase hydrolyzes acetylcholine to acetic acid and choline LD50 values for selected substances Ranking cancer risks A schematic diagram of an incinerator for hazardous wastes

Demonstrations

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Chapter 22

2

[None]

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Instructor's Manual To Chemical Investigations for Changing Times Twelfth Edition Alton Hassell Baylor University Paula Marshall Baylor University Lori Morgan Wagner Baylor University Shannon Dove Baylor University Prentice Hall

Table of Contents

General Comments Coordination with Hill and Kolb Text Required Times for the Investigations Common Safety Regulations Laboratory Instructions Equipment Suppliers Basic Equipment List Investigation #1-#59 Lab Notes Reagent List (all investigations) Common Materials (all investigations) Laboratory Equipment (all investigations)

iv v vi viii xii xiii xv 1 159 164 168

iii

General Comments The safety rules and laboratory instructions can be modified, copied, and supplied to students on the first day of lab.

Equipment listed in the "basic equipment list" will be listed again in the individual lab notes, as will the extra lab equipment necessary to an investigation. An alphabetical list of equipment used in each lab is included after the materials list.

Where 500-mL beakers are called for, 600-mL beakers will do.

It is assumed that lab will be equipped with distilled water and scales, or balances accurate to + 0.1 gm. Where more accurate measurements are necessary, the need for more sophisticated equipment will be noted.

Instructions for preparing specific reagents are given in the lab notes where needed.

The investigations are designed for students to work in pairs.

The notes for each lab give specific instructions about disposal methods for toxic substances.

In general, it should be noted that concentrated acids and bases can be harmful to plumbing if not diluted or neutralized. In addition, it is not a good practice to put metal ions down the drain.

The investigations are planned to fit into either a two-hour or a three-hour lab period. The instructor may need to adjust the timing with pre-lab lectures or solutions prepared ahead of time. Some of the investigations may be combined into one lab period, especially for three-hour periods.

We encourage the use of field trips as investigations. Chemistry is so much more meaningful to the student who sees it first hand in a sewage treatment plant, water treatment plant, or industrial plant. There are also computer simulation programs that are useful for certain concepts. The Serephim Project and the ACS are good sources for computer programs.

A source of information for the disposal of chemicals is The Science Instructor's Safer Source from Flinn Scientific Inc., P.O. Box 219, Batavia, IL 60510.

We would be interested in hearing about any problems you find or improvements that you can suggest for the investigations herein.

iv

Coordination with Chemistry for Changing Times, 12th Ed. Hill and Kolb The investigations are arranged to show the principles of the chapters in Chemistry for Changing Times, 12th Ed., Hill and Kolb. Some of the investigations will also illustrate principles of other chapters. These are listed in the last column.

TEXT CHAPTER 1.

2.

3.

4.

5.

6.

7.

8.

9.

10.

11.

12.

13.

14.

15.

16.

17.

18.

19.

20.

21.

20.

Chemistry Atoms Atomic Structure Chemical Bonds Chemical Accounting Gases, Liquids, Solids, and Intermolecular forces Acids and Bases Oxidation and Reduction Organic Chemistry Polymers Nuclear Chemistry Chemistry of the Earth Air Water Energy Biochemistry Food Drugs Fitness and Health Chemistry on the Farm and in the Garden Household Chemicals Poisons

APPLICABLE INVESTIGATION

ALSO APPLICABLE

1, 2, 3 4 5, 6 7, 8, 9, 10, 11, 12, 13

15 11 3, 13, 22, 56 4, 30, 46

14, 15, 16 17, 18, 19, 20 21, 22, 23 24, 25 26, 27 28 29, 30 31, 32 33, 34, 35, 36 37, 38, 39, 40 41, 42, 43, 44 45, 46, 47 48 49, 50, 51

2, 4, 12, 19, 34 57 40 48, 49, 53 33 58 29, 39 3, 8, 36, 57 31, 45, 50, 51, 56 25, 41, 42, 43, 50, 51 49 32, 41, 42, 45

52 53, 54, 55, 56, 57 58, 59

19, 27 -

v

Required Times for the Investigations Note: The noted times were for a single person working in a lab as the sole user of balances and all other equipment, and who is experienced in the laboratory setting. Allow extra time for students in this course.

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42

Alchemy Density Layers Energy in Physical and Chemical Changes Diffusion Atoms and Light Flame Tests and Analysis Ionic versus Covalent Chemical Reactions Double Replacement Reactions Iron(II) and Iron(III) Mole Relationships Molar Volume of a Gas Hydrated Compounds Solubility Viscosity Gas Laws Acid Neutralization by Antacid Nature’s Indicators pH and Flammability Buffers Oxidation and Reduction Single Replacement Reactions Silver Tarnish Removal Organic Qual Scheme Synthesis of Esters Polymers Polymer Properties Nuclear Radiation Salt on Ice A Copper Cycle O2 and CO2 in Breath and Air Solids in Smoke Clarification of Water Freezing Water Freezing Point Depression Heat Capacity Heat of Reaction Rates of Chemical Reactions Canned Fuel and Evaporation Batteries Test for Vitamin C Proteins

vi

45 minutes 40 minutes 1 hour 1 hour 15 minutes 1 hour 45 minutes 1 hour 1 hour 5 minutes 45 minutes 60 minutes 1 hour 1 hour 1 hour 30 minutes 30 minutes 1 hour 45 minutes 2 hours 1 hour 40 minutes 1 hour 55 minutes 45 minutes 30 minutes 2 hours 35 minutes 1 hour 30 minutes 45 minutes 1 hour 30 minutes 2 hours 3 hours 1 hour 30 minutes 3 hours 1 hour 30 minutes 1 hour 30 minutes 45 minutes 45 minutes 1 hour 30 minutes 1 hour 15 minutes 1 hour 1 hour 45 minutes 45 minutes 1 hour 15 minutes 1 hour 15 minutes

43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59

Enzymes Forensic Chemistry Protein in Milk and Other Foods Production of Carbon Dioxide Food Colors by Paper Chromatography Detection of Alcohol Aspirin Preparation and Purity Test Fats and Iodine Number Carbohydrates Analysis of Fertilizer Soap in Hard Water Personal Products Hair Chemistry Saponification Testing of Sunscreens Lead Detection Parts Per Million

1 hour 30 minutes 45 minutes 1 hour 15 minutes 1 hour 30 minutes 1 hour 30 minutes 45 minutes without disposal 1 hour 30 minutes 55 minutes 1 hour without disposal 1 hour 50 minutes 1 hour 45 minutes 40 minutes 1 hour 45 minutes 50 minutes 30 minutes 1 hour 45 minutes without disposal 35 minutes

vii

Common Safety Regulations Under OSHA Hazard Communication regulations, a specific set of safety rules must be developed and communicated clearly to all employees. The Department of Chemistry has developed safety regulations that comply with the University's Hazard Communication Program.

Although non-employee students are not covered by these regulations, each student should attest by signature that he or she has read and understands the safety rules. These rules must be rigorously and impartially enforced. Willful noncompliance should result in dismissal or suspension from the laboratory.

An excellent reference book published by The American Chemical Society is called Safety in Academic Chemistry Laboratories. This publication can be obtained by calling The American Chemical Society Referral Service at 1-800-227-5558, option 6. Another source of information is the Chemical Manufacturers Association Referral Center at 1-800-262-8200 (24 hours).

General Safety Regulations The Chemistry Department strives to provide safe working conditions in all of its laboratories. The responsibility for safe working conditions rests with everyone, including students, faculty, teaching assistants, and staff. In order to do your part in maintaining safe working conditions, you must obey the General Safety Regulations listed below and report any violations of these regulations which you observe. Such reports should be made to the teaching assistant or to the professor in charge of the laboratory.

1. Eye protection is required at all times in the laboratory and where chemicals are stored and handled. For most laboratory workers, safety glasses with clear side shields are adequate as long as safety showers or eye-wash fountains are near at hand.

2. Long hair and loose clothing must be contained while in the laboratory. Sturdy closed-toe shoes that cover the entire foot and clothing that covers the body constitute proper laboratory attire. You will not be permitted to work in the laboratory if you are barefoot or wearing flipflops or sandals, including sandals with socks. You will not be permitted in the laboratory if you are wearing shorts, short skirts (mini-shorts), halter tops, spaghetti strap blouses, or tank tops.

Lab coats or aprons are suggested.

3. Work only with materials after you know their flammability, reactivity, corrosiveness, and toxicity. This information is readily available through your laboratory instructor and Material Safety Data Sheets. Read your laboratory instructions carefully before beginning any laboratory experiment. When in doubt about any part of the procedure, consult your laboratory instructor before initiating the experiment.

4. Unauthorized experiments, including unapproved variations in experiments and changes in quantities of reagents, are forbidden.

5. Mouth suction should never be used to fill pipettes, to start siphons, or for any other purpose.

6. Don't force glass tubing through rubber stoppers. (Protect your hands with a towel when inserting tubing into stoppers and use water or glycerine as a lubricant.)

viii

7. Wash hands with soap or detergent upon leaving the laboratory.

8. Smoking in any part of the laboratory is absolutely forbidden.

9. Experiments in progress are not to be left unattended without consent of the laboratory instructor.

10. Horseplay, pranks or other acts of mischief are especially dangerous in the laboratories and are absolutely forbidden. Sitting on bench tops is not allowed.

11. No eating or drinking is permitted in the laboratory.

12. Students are not permitted to work in the laboratory unless a teaching assistant is present.

13. All accidents, no matter how trivial, must be reported. If you have an accident or see an accident, it must be reported to the teaching assistant immediately. Never allow an injured person to leave the laboratory for assistance without notifying the teaching assistant. A formal written report must be made of all accidents to the Department Safety Officer. The report will include any injuries or damage incurred, the cause, the effect, and any recommendations for prevention of recurrence.

14. In case of an accident, treat all open wounds as potential sources of blood-borne pathogens.

15. If an accident appears to be beyond your control (i.e., a large chemical spill or fire), actuate the fire alarm and vacate the area. If possible, call the University Department of Public Safety and give them the details of the problem, (i.e., location, chemicals involved).

16. Visitors to the laboratories must obey the same safety regulations as the laboratory workers. This includes the wearing of safety glasses.

17. Chemical laboratories and storage areas constitute a special hazard for small children.

For this reason, the Department does not permit the presence of small children in these areas under any circumstances. Children are not to be allowed elsewhere in the Chemistry Department except under close adult supervision. The responsibility for the safety of such children must be assumed by the supervising adult. Under no circumstances should children be left unattended, even for short periods of time, nor should they be brought into areas of risk.

Chemical Labeling 1. All chemical labels must contain the following information: • chemical identity • appropriate hazard warnings • responsible company and address 2. OSHA standard labels must be used. These labels are available in the stockroom.

3. Proper labels must be on or attached to each hazardous material container.

ix

4. Immediate use, small transfer containers are exempted from these regulations.

5. No labels may be removed or defaced without replacements.

6. Labels must be legible, prominent, and in English.

The National Fire Protection Association (NFPA) hazard identification coding system is used in the Chemistry Department to readily identify chemical hazards. A copy of this labeling system follows.

NFPA Hazard Code Ratings The National Fire Protection Association (NFPA) ratings provide information relating to the severity of health (blue), flammability (red), and reactivity (yellow) under emergency conditions such as spills, leaks, and fires. Each section displays a numerical severity rating from 4 (most severe) to 0 (least severe). The fourth section, as stipulated by NFPA, is left blank (white) and is reserved for indicating any unusual reactivity with water or oxidizing properties. The numerical ratings and their associated descriptions are given on the next page.

x

HEALTH (Blue) 4

Material that on very short exposure could cause death or major residual injury.

3

Material that on short exposure could cause serious temporary or residual injury.

2

Material that on intense or continued, but chronic, exposure could cause temporary incapacitation or possible residual injury.

1

Material that on exposure would cause irritation but only minor residual injury.

0

REACTIVITY (Yellow) 4

Material that in itself is readily capable of detonation, or of explosive decomposition or reaction at normal temperatures and pressures.

3

Material that in itself is capable of detonation or explosive reaction but does not require a strong initiating source; or which must be heated under confinement before initiation; or may react explosively with water.

2

Material that is unstable and readily undergoes violent chemical change at elevated temperatures and pressures, or which reacts violently with water or may form explosive mixtures with water.

1

Material that in itself is normally stable, but can become unstable at elevated temperatures and pressures.

0

Material that in itself is normally stable, even under fire exposure conditions, and is not reactive with water.

Material that on exposure under fire conditions would offer no hazard beyond that of ordinary combustible material.

FLAMMABILITY (Red)

4

3 2

Material that will rapidly or completely vaporize at atmospheric pressure and normal ambient temperature, or is readily dispersed in air and will burn readily.

A liquid or solid that can be ignited under almost all ambient temperature conditions.

SPECIAL WARNINGS (White)

Material that must be moderately heated or exposed to relatively high ambient temperatures before ignition can occur.

OX

Oxidizing material.

1

Material that must be preheated before ignition can occur.

Material that is hazardous when in contact with moisture or water.

0

Material that will not burn.

W

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Laboratory Instructions

1.

Avoid wasting gas, water (both tap and deionized), filter paper, and materials of any kind. Do not return excess quantities to reagent bottles. Do not remove reagent bottles from their proper place.

2.

Do not use your medicine dropper to take reagents from the classroom stock bottles. Instead, pour out a small quantity into your own clean beaker and use your medicine dropper in that beaker. If a dropper bottle is provided with the reagent, hold that dropper near the top of your test tube or other container when dispensing that reagent.

3.

Do not remove acids or alkalis from their metal trays; take what you need at the tray. If acids or alkalis are spilled on the desk or floor, notify the instructor. The clean-up procedure is to wash the desk or floor at once with plenty of water, then wash the desk with sufficient boric acid solution in the case of spilled acids. This will serve to neutralize the acids or alkalis that have penetrated into the wood. Then wash the desk once more with plenty of water.

Commercial sorbents and neutralizers are also available, and should be used if provided.

4.

When leaving the laboratory, make certain that the gas and water are turned off and that your desk is clean and neat.

5.

Maintain an orderly arrangement of the apparatus and materials in your desk.

xii

Equipment Suppliers ACE Glass Inc., P.O. Box 688, 1430 Northwest Blvd., Vineland, NJ 08361-0688 800-223-4524, www.aceglass.com, sales@aceglass.com.

Aldrich Chemical Company, Inc., 1001 West St. Paul Ave., Milwaukee, WI 53233 800-325-3010, sigma-aldrich.com/order Allied Electronics, 800-433-5700, www.alliedelec.com Atomeric Chemetals Corp., 222 Sherwood Ave., Farmingdale, NY 11235-1718 516-694-9000, www.nuclearstreet.com, BABL28A@prodigy.com Carolina Biological Supply Co.,,2700 York Road, Burlington, NC 27215-3398 800-334-5551, www.carolina.com Central Scientific Co., 3300 CENCO Parkway, Franklin Park, IL 60131-1364 708-451-0150 Edmund Scientific Co., 60 Pearce Ave., Tonawanda, NY 14150 800-728-6999, www.scientificsonline.com Fisher Scientific, 711 Forbes Avenue, Pittsburgh, PA 15219-9919 800-766-7000, www.fishersci.com Fisher Science Education, 4500 Turnberry Drive, Hanover Park, IL 60133 info@fisheredu.com Flinn Scientific, P.O. Box 219, Batavia, IL 60510 800-452-1261, www.flinnsci.com, flinn@flinnsci.com Frey Scientific Co., 905 Hickory Lane, P.O. Box 8101, Mansfield, OH 44901-8101 Toll Free: 800-225-FREY, www.freyscientific.com Hach Chemical Co., P.O. Box 389, Loveland, CO 80539-0389 800-227-4224, www.hach.com Hilton Davis Company, 2235 Langdon Farm Road, Cincinnatti, OH 45237-4712 513-841-4000 J.T. Baker Chemical Co., 800-JTBAKER,,www.jtbaker.com Lab Safety Supply, P.O. Box 1368, Janesville, WI 53547-1368 800-356-0783, www.labsafety.com LaMotte Chemical Products Co., Box 329, Chesterton, MD 21620 Toll Free: 800-344-3100

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NCI Chemical Carcinogen Reference Standard Repository, 425 Volker Boulevard, Kansas City, MO 64110-2289, 816-753-7600, Fax: 816-753-3664 Pfaltz and Bauer, Inc., 800-CALL-1-PB, www.pfaltzandbauer.com Sargent-Welch, 800-727-4368, www.sargent-welch.com Sigma Aldrich Chemical Company, 800-325-3010, www.sigma-aldrich.com Spectrum Chemical and Laboratory Products, 800-772-8786, www.spectrumchemical.com TCI America Inc., 800-423-8616, www.tciamerica.com Thomas Scientific, 800-345-2100, www.thomassci.com VWR Scientific, 800-932-5000, www.vwr.com Wako Chemicals USA, Inc., 1600 Bellwood Road, Richmond, VA 23237-1326 877-714-1920, www.wakousa.com Warner Jenkinson Company, 2526 Baldwin Street, St. Louis, MO 63178-053806 Toll Free: 800-325-8110, 314-889-7600 A good source of information on companies which supply specific chemicals (listed by chemical name) is CSU: Chem Sources U.S.A, a publication available at many library reference desks.

xiv

Basic Equipment List (per pair of students)

evaporating dish 2 watch glasses glass-stirring rod with policeman 1 suction apparatus (Buchner funnel and filter flask) 2 glass funnels beakers: 2 50 mL 2 150 mL 1 250 mL 1 600 mL 1 1000 mL graduated cylinders: 1 10 mL 1 100 mL Erlenmeyer flasks: 1 125 mL 2 250 mL 1 500 mL test tubes: 6 small (about 10 cm long) 4 large (about 15 cm long) 2 extra large (about 25 mm x 200 mm) stoppers: 1 hole to fit x-large test tube 1 hole to fit large test tube solid to fit small test tube solid to fit large test tube 2 hole to fit 500 mL Erlenmeyer flask 1 hole to fit 10 mL graduated cylinder 1 hole to fit 250 mL Erlenmeyer flask 1 hole to fit 100 mL graduated cylinder spatula eye dropper thermometer

xv

ringstand ring wire and "asbestos" pad Bunsen burner flint striker or matches 2 clamp holders 2 clamps assorted rubber hoses glass tubes for holed stoppers forceps tongs test tube holder test tube brush 2 squeeze (wash) bottles powdered cleanser

Lab Notes

Note: There is a page of notes to the instructor for each lab, followed by a Pre-lab Question page, and a Report Page with sample answers to all of the questions and most of the data. It should be noted that data acquired with another system might differ.

#1

Alchemy

Reagents zinc, powdered 6 M sodium hydroxide [NaOH] 6 M nitric acid [HNO3]

1 g/pair 25 mL/pair 25 mL/lab

Common Materials copper token, wire, or piece of plate rulers, metric spray can of clear acrylic coating (optional)

1/student 1/pair 1/lab

Laboratory Equipment balance caliper ring stand ring wire gauze evaporating dish forceps laboratory burner or hotplate stirring rod Special Equipment Notes Pre-1982 pennies are pure copper, but are difficult to clean thoroughly enough to plate out well and produce a satisfactory result.

The best tokens are new and very shiny pennies. Any material on the surface will hinder the adherence of the zinc to the copper. Even slightly darkened areas of finger prints are enough to alter the quality of the finished product.

The sodium hydroxide solution can be less concentrated, but it will slow the reaction.

Caution the students to refrain from heating the zinc/copper surface too long as the coin will melt. Constantly turning the coin while heating is a safer method. If the students will lower the coin into the flame for a second on each side and repeat continuously until the gold color appears, then heat each side one more time, they will ensure a complete reaction without misshaping the coin. A hot plate instead of a Bunsen burner will turn the token to a new color without melting it.

Disposal Wet zinc dust exposed to air can burst into flames. It is important that students not put wet zinc dust in the trash; it can cause fires in the trash. The wet zinc solutions should be spread on a metal pan to dry. It must be in a metal pan for the chemical reaction to form zinc oxide. Dry zinc oxide will form which can be buried in a landfill. The used hydrochloric acid should be flushed down the drain with plenty of water.

2

ALCHEMY PRE-LAB QUESTIONS 1) Legend holds that Archimedes was responsible for proving that a metallurgist had “cut”, or diluted, the gold for a Ruler’s crown with a lesser metal. Archimedes asked the Ruler for a mass of gold identical to what he had provided to the metallurgist. What measurements and calculations did Archimedes most likely make?

Assuming that the metallurgist was not stupid and the finished crown had the same mass as the originally supplied gold, Archimedes most likely took two volume measurements by liquid displacement – one of the identical mass of gold and one of the crown. Then finding that they were not the same, he deduced that since the smaller volume was the pure gold, the crown was only partially gold and contained a less dense metal. The metallurgist had indeed kept some gold for himself.

2) The alloy produced in this investigation is brass. There are several different types of brass. Use an internet encyclopedia resource to find at least three different types of brass.

Describe the differences in proportion and variety of metals used, and the effect on the resulting properties of the brass.

Answers will vary but will most likely include some of the following: • • • • • • • • • • • • • • • • • • •

Admiralty brass contains 30% zinc and 1% tin that will inhibit dezincification.

Alpha brasses, also known as Prince's metal, with less than 35% zinc, are malleable, can be worked cold, and are used in pressing, forging, or similar processes.

Alpha-beta brass, also known as Muntz metal and called duplex brass, is 35-45 % zinc and is suited for working hot.

Aluminum brass contains aluminum, which improves its corrosion resistance.

Arsenical brass contains an addition of arsenic and frequently aluminum and is used for boiler fireboxes.

Beta brasses with 45-50 % zinc content can only be worked hot, are harder, stronger, and suitable for casting.

Bronze is an alloy of copper with tin and optionally zinc, silicon, nickel, and other metals.

Calamine brass is a brass alloy manufacturing process from the first millennium BC that was not replaced in Europe until the late 18th century.

Cartridge brass is a 30% zinc brass with good cold working properties.

Common brass, or rivet brass, is a 37% zinc brass, is inexpensive, and is the standard for working cold.

Cupronickel is an alloy of copper with nickel.

High brass contains 65% copper and 35% zinc, has a high tensile strength, and is used for springs, screws, and rivets.

Leaded brass contains addition of lead. It has excellent machinability.

Low brass is a copper-zinc alloy containing 20% zinc with a light golden color, excellent ductility and is used for flexible metal hoses and metal bellows.

Naval brass, similar to admiralty brass, is a 40% zinc brass and 1% tin.

Pinchbeck is a brass that closely resembles gold in appearance.

Red brass is an American term for Cu/Zn/Sn alloy known as gunmetal.

White brass contains more than 50 % zinc and is too brittle for general use.

Yellow brass is the American term for 33% zinc brass.

[Information from Wikipedia]

3) Consider two different forms of brass. Brass A is 65% copper and 35% zinc, and Brass B is 75% copper and 25% zinc. Which form will be the densest? Explain your reasoning.

Brass A will have a smaller density since there is more of the less dense zinc in the alloy.

4) If the crown fit comfortably, what would be the difference between wearing a brass crown and wearing a gold crown of identical mass?

The pure gold crown would be much smaller in volume.

3

ALCHEMY SAMPLE REPORT SHEET I. MASS OF ORIGINAL TOKEN: _____2.50g_____________ II. MASS OF DRY SILVER TOKEN: ___2.51g-2.53g_____ Instructor's Initials for silver token: ____________________ How does the silver token feel? (Is it smoother than the copper token? Are there rough spots?)

If the penny is very clean and bright to begin with, the coating will not change the texture.

III. MASS OF COOLED GOLD TOKEN: ___2.51g-2.53g______ Instructor's Initials for gold token: _______________ How does the gold token feel?

If it is heated too long, the copper layer with its brass coating will bubble up and become rough.

Otherwise, it should remain very smooth.

IV. DECISION Do you tell the King that the token is really turned to silver and then to gold, or that the silver-colored token is just zinc-coated copper and that the zinc diffuses into the copper to make brass, a solid solution of copper and zinc?

EXTRA Don't lose your head by a wrong decision.

Consider the densities.

Copper to Silver to Gold X

Copper to Zinc coating to Brass

V. DENSITY EXTRA If you claim the token is not gold, how can you Density = mass/volume. prove it by the density?

The density of gold is 19.3g/cm3, so the mass of a gold token would be greater than the experimental token’s mass. Therefore, the token is not gold.

Many tokens are not flat but an average thickness can be calculated from the measured mass and radius. t = m/(πr2d) where r is radius (1/2 of the diameter), π is 3.14, t is thickness, d is density, and m is the measured (actual) mass.

Measurements of copper token: diameter ____1.9____________ cm

EXTRA densities are: radius ______0.95___________ cm copper 9.0 g/cm3 silver 10.5 g/cm3 mass _______2.50_________ g gold 19.3 g/cm3 zinc 7.1 g/cm3 calculated average thickness _______0.10___________ cm The volume of a token can be calculated from volume (v) = πr2t.

4

volume of token _____0.27_______ cm3 You can calculate the mass of the token using the equation: mass (m) = density (d) × volume (v). The token thickness did not change significantly as it changed from copper to silver to gold, so you can calculate the expected masses of a solid silver and a solid gold token using the volume of the copper token and the densities of silver and of gold. You can then compare those masses with the actual masses.

CALCULATED MASS Original Token Silver Token Gold Token

ACTUAL MASS ____2.50g_

___2.84g___

_2.51-2.53g

___5.21g____

_2.51-2.53g_

VI. QUESTIONS 1. If the zinc is adhered to the copper instead of bonded to the copper, is the change to silver color a chemical change?

No, this is not a chemical change.

2. If the heating causes the zinc to bond with the copper, is the change to gold color a chemical change?

Yes, the properties of the metal (color, density of the alloy, etc.) are changed.

3. What happened to the mass of a penny when the U.S. changed from pure copper to a copperclad zinc coin? (Hint: Consider the densities given above.)

The mass of the penny decreased because the zinc has a smaller density than the copper.

4. What happened to the mass of a dime when the U.S. changed from pure silver to a silvercopper-silver sandwich?

The mass of the dime decreased because the copper has a smaller density than the silver.

5. How long would it take a person who received an object “changed into gold” to realize they had been duped? What would most likely be the first indication?

As soon as something rubbed or scraped the thick alloy coating from the surface, the copper below it would become visible.

5

#2

Density Layers

Reagents isopropyl alcohol

10 mL/pair

Common Materials antifreeze brake fluid 30 wt motor oil power steering fluid transmission fluid mineral oil Ajax™ laundry detergent (liquid) Downy™ fabric softener Karo™ syrup molasses vegetable oil small cork paraffin thumb tack plastic paper clip aluminum paper clip a small stopper or pieces of rubber band food coloring - green

1 mL/pair 1 mL/pair 1 mL/pair 1 mL/pair 1 mL/pair 10 mL/pair 10 mL/pair 10 mL/pair 10 mL/pair 10 mL/pair 10 mL/pair 1/pair 1 small piece/pair 1/pair 1/pair 1/pair 1/pair 1/pair

Laboratory Equipment 10-mL graduated cylinder 50-mL graduated cylinder 100-mL graduated cylinder funnel small test tube 250-mL beaker Special Equipment Notes The smallest rubber stopper issued to the students will suffice for the one called for in the experiment, but a smaller one would be better.

Disposal Most automobile fluids are oily and can contaminate water if not disposed of properly.

Ethylene glycol is poisonous. Fortunately, service stations are part of our oil-recycling program.

Pour the auto fluids into a marked waste container. The household products solutions can be flushed down the drain.

6

DENSITY LAYERS PRE- LAB QUESTIONS 1) When you are walking through a parking lot and see a puddle of water with a rainbow effect on the surface, you are observing a thin-film optical phenomenon. Do you suspect the material in the film is miscible or immiscible with the water? Explain your reasoning.

The substance must be immiscible in order for it to form a film on the surface. If it dissolved, there would be no top layer, or film.

2) What is the nature of film on the surface of the water in the first question and what was its likely source?

The particles are not miscible so they are probably long-chain organic molecules. Since it is in a parking lot, it is most likely a petroleum product and could be oil or another auto fluid that has dripped from a car, or gasoline.

3) Assume you had to separate a series of layers of materials that have different densities.

Suggest a possible method for separating each of the following: a) solid phase objects of greater density than the liquid in which they are contained The less dense liquid could be poured from the top leaving the more dense solid behind. b) solid phase objects of less density than the liquid in which they are contained The solids could be filtered or they could be skimmed from the surface. c) two immiscible liquids of greatly varying densities Most of the top layer could be poured from the top, and then the rest could be skimmed. Also, a separatory funnel or gravy cup could be used. d) two miscible liquids of greatly varying densities A buret or separatory funnel would be the best choice so that the interface is moved as little as possible.

4) We should be careful not to allow petroleum products to get into our waterways, and yet each time it rains we see evidence of those products in the environment. Can you suggest a creative method to trap these products as they run into street sewers?

It would be very difficult to separate them as they run into the gutters, but a settling basin that held the liquids long enough for them to separate and then allowed the water to drain out of the bottom, holding the less dense petroleum liquids on the surface would work well.

5) Why is it not a problem for vegetable oil to enter waterways when motor oils are an environmental hazard?

Vegetable oils are biodegradable materials whereas petroleum products remain in the environment for many years and can be poisonous.

7

DENSITY LAYERS SAMPLE REPORT SHEET I. AUTO LIQUIDS List the fluids in order of decreasing density (heaviest first). _ antifreeze________________ __brake fluid______________ __transmission fluid_________ __30 wt motor oil ___________ __manual power steering fluid__

Instructor's initials: __________

II. HOUSEHOLD LIQUIDS (the following is sample data only) Mass of 10 mL graduated cylinder 28.07 g Mineral Oil Cylinder and oil 36.66 g Volume 10 minus Cylinder

- 28.07

g

8.59

g

Oil Isopropyl alcohol Cylinder and alcohol minus Cylinder

-

Alcohol

minus Cylinder

-

Ajax Downy fabric softener Cylinder and Downy minus Cylinder Downy

Density

8.59 10

g = _0.859_g/mL mL

35.64 g

Volume

7.7

mL

28.07 g

Density

7.57 g =__0.983_g/mL 7.7 mL

7.57

Ajax laundry detergent Cylinder and Ajax

-

mL

g

34.15

g

Volume

9.4

28.07

g

Density

6.08 9.4

6.08

g

34.76 g

Volume

6.7

mL

28.07 g

Density

6.69 6.7

g = _0.999_g/mL mL

6.69

g

8

mL g = 0.647g/mL mL

Karo Cylinder and Karo

41.07

g

Volume

9.5

mL

minus Cylinder

- 28.07

g

Density

13.0 9.5

g = __1.37__g/mL mL

13.0

g

Karo Molasses Cylinder and Molasses

39.13

g

Volume

7.8

mL

- 28.07

g

Density

11.06 7.8

g = _1.42_g/mL mL

Molasses

11.06

g

OBJECT

ESTIMATED DENSITY

Thumb tack

1.4

g/mL

Aluminum paper clip

0.75

g/mL

Rubber stopper/band

1.3

g/mL

Plastic paper clip

0.65

g/mL

Paraffin

0.6

g/mL

Cork

0.5

g/mL

minus Cylinder

OBJECT

ESTIMATED DENSITY

III. QUESTIONS 1. Why are the auto liquids not poured down the drain?

The liquids can contaminate water and some are poisonous.

2. Of the auto fluids, which ones mix well together? Which fluids form separate layers?

The antifreeze and brake fluid mix well. The power steering fluid does not mix.

3. After observing the difference in the densities of the auto liquids, why is it important that the liquids used in an automobile be uncontaminated?

The liquids do not mix entirely, and all have different densities. This will cause the improper functioning of an automobile.

4. What does it mean if an object sinks to the bottom of the cylinder?

This means that the object is greater than or equal in density than the bottom liquid.

5. What does it mean if an object sits on top of the liquids?

This means that the object is less dense than the top liquid.

6. Karo and molasses are mostly sugar in water, why do they form separate layers?

They form separate layers because they have different proportions of sugar and water which makes their densities different.

7. Which of the household liquids will float on water?

The Ajax, Downy, alcohol, and mineral oil will float.

9

#3

Energy in Physical and Chemical Changes

Reagents ammonium chloride [NH4Cl] calcium chloride [CaCl2] sodium hydroxide [NaOH] solid 1.0 M sodium hydroxide [NaOH] solution sulfuric acid [H2SO4] concentrated 1.0 M hydrochloric acid [HCl] 0.5 M hydrochloric acid [HCl]

6 g/pair 9 g/pair 4 g/pair 50 mL/pair 5 mL/pair 50 mL/pair 100 mL/pair

Common Materials styrofoam cup cardboard (for lid) wire for stirrer

1/pair 1/pair 1/pair

Laboratory Equipment thermometer 100-mL graduated cylinder Special Equipment weighing paper or weighing boats Notes Any moisture on the metal when it is transferred will be a source of error because of the relatively high specific heat of the water.

Disposal All solutions can be flushed down the drain with plenty of water.

ENERGY IN PHYSICAL AND CHEMICAL CHANGES PRE-LAB QUESTIONS 1) In a certain ionic compound, more energy is required to break the lattice energy holding ions together in the crystal than is released when the ions dissolve in water. Will the overall process appear to be endothermic or exothermic? Explain your logic. Will the beaker feel warm or cool?

The required energy is an endothermic process while the released energy is exothermic.

Since more is required, the overall process will be endothermic. The beaker will feel cool as the energy leaves your hand.

2) Consider two solids that are stirred together in a glass beaker at room temperature and interact to produce a liquid while frost is seen to form on the outside of the beaker. Is the process endothermic or exothermic? Discuss the energy changes that can be assumed in the process.

The frost on the beaker would suggest that the process is endothermic. This means that the energy released to form the new substance is not nearly as great as the amount of energy required to break the bonds in the two original compounds.

3) Is the resulting liquid in question #2 most likely the result of a physical or chemical change?

Use energy changes to support your answer.

10

If the solid to liquid change was a physical process, one would expect the energy change to be endothermic so it would be hard to prove that the liquid was the result of a chemical change. However, the fact that the temperature continued well below room temperature would suggest that there is another force at work and that the process is indeed chemical.

4) Two substances in solution at room temperature are mixed in a test tube containing a thermometer to register temperature changes. There is no change on the thermometer but the combined solution immediately becomes cloudy and a white powder eventually settles to the bottom of the test tube. Why is it incorrect to determine that no energy changes occurred?

What is the term for this type of process?

Energy is required to separate dissolved particles from the solution, but it will be released when the new bonds form in the solid. The energy changes in the two opposite processes were simply equal in magnitude. This is an isothermic process.

5) Ice and hot water are combined in the same container. Discuss the energy changes that will take place and any state changes that will result.

The energy from the hot water will be transferred to the ice to melt it. The ice absorbs heat (an endothermic process) as the hot water loses it (an exothermic process).

6) Considering the container discussed in question #5, if ice remains after the temperature stabilizes, what can be assumed about the final temperature?

The final temperature would be the freezing point of the water.

ENERGY IN PHYSICAL AND CHEMICAL CHANGES SAMPLE REPORT SHEET I. PART A: Heat of Solution and Dilution (Physical Changes) TEMPERATURE INITIAL FINAL Ammonium chloride _______ _____ _________ Calcium chloride _________ ________ Sodium hydroxide _________ ________ Sulfuric acid _________ ________

t (final – initial) ________ ________ ________ ________

II. PART B: Heat of Reaction (Chemical Changes) TEMPERATURE INITIAL FINAL HCl and NaOH solution _________ ________

t (final – initial) ________

HCl and solid NaOH

_________

________

________

III. QUESTIONS 1. Is there a difference in the final solution temperature of the HCl with NaOH solution and with solid NaOH? Can the difference by explained by the heat of solution of NaOH? (Hint – Compare the temperature changes for the HCl with solid NaOH to changes for both the water with solid NaOH and the HCl with NaOH solution.)

There is a difference because the solid must first be dissolved and then react.

2. Explain why acid is poured into water instead of water into acid given the fact that if drops splatter during pouring, the drops are usually of the liquid poured into, not the liquid being poured.

Acid is poured into water so that the splattered drops will be water not acid.

3. Could a solution of one of these compounds serve as a cold pack for an athletic injury?

An exothermic reaction could be used for a cold pack.

11

#4

Diffusion

Reagents hydrochloric acid [HCl] concentrated ammonia [NH3] concentrated

5 mL/pair 5 mL/pair

Common Materials cotton swabs ice cubes food coloring blue or green

10 swabs/pair 2 cubes/pair 2 drops/pair

Special Equipment 1" glass tube one hole stoppers stopwatch grease pencil

1/pair 2/pair 1/pair 1/pair

Laboratory Equipment 4 250-mL beakers thermometer 2 watch glasses or evaporating dishes barometer Notes As the students get more adept at spotting the ammonium chloride ring, their times get shorter.

It might be wise to have them do one or two practice runs before they begin taking data.

Disposal All solutions can be flushed down the drain with plenty of water. The cotton swabs can be thrown into the trash.

DIFFUSION PRE-LAB QUESTIONS 1) Give several evidences of diffusion from your own experience.

Answers will vary but should include things like: knowing someone is painting their nails without being able to see them, walking into a house and knowing what is for dinner, knowing there is a fire somewhere without seeing smoke, looking at last night’s leftover iced tea and seeing the gradual fading of color near the top where the ice was instead of a separate layer of water from the melted ice, etc.

2) If gas molecules are moving at a tremendous rate, why is diffusion not an almost instantaneous process?

Molecules don't travel a straight line but a very random path depending on collisions with other molecules.

3) How would an increase in atmospheric pressure be expected to affect the rate of diffusion in a gas?

The rate of both gases would be slower as there would be more molecules in the pathway.

4) Explain how an increase in temperature would be expected to affect the rate of diffusion in a gas.

The rate of diffusion would increase because molecules are moving faster.

5) How does the molar mass of gas particles affect the relative rates of diffusion?

The smaller the molar mass, the faster the rate of diffusion.

12

DIFFUSION SAMPLE REPORT SHEET I. DATA AND CALCULATIONS PART A: The Effect of Temperature on Diffusion in a Liquid Beginning time: ____________________ Beaker #1

Beaker #2

Beaker #3

Temperature Ending time Elapsed Time

depends on temperature

PART B: The Effect of Particle Mass on Diffusion in Gases Trial #1

Trial #2

Trial #3

Time (in sec)

148

120

103

D (HCl to ring) (cm)

8.0

10.1

12.3

D (NH3 to ring) (cm)

22.0

20.2

17.7

R (HCl) (cm/sec)

0.054

0.084

0.119

R (NH3) (cm/sec)

0.13

0.16

0.17

Molar Mass

Average Rate

HCl

36.46 0.083

0.083

NH3

17.03 0.15

0.15

Rate Ratio: Experimental

1.8:1

Calculated

1.5:1

13

Trial #4

II. QUESTIONS 1. What was the effect of the different temperatures on the rates of diffusion in the beakers of water?

The colder the temperature the slower the mixing occurred.

2. Did you notice any “currents”, or convection, as the color moved through the water? If so, account for the cause of the currents.

The water is not homogeneous. The currents are because of the mixing of the waters of different temperatures.

3. Write the balanced equation for the formation of ammonium chloride.

HCl(g) + NH3(g) → NH4Cl(s) 4. Use the experimental ratio and the accepted ratio to calculate percent error.

|E - A| x 100% A Percent error =

1.8 - 1.5  100% = 20% 1.5

5. List possible sources of error that account for the percent error in your experimental data.

An error in seeing ring, width of grease mark, swabs not put in at the same time or put in differently are examples of types of errors.

14

#5

Atoms and Light

Reagents magnesium ribbon 0.1M chloride salts of Li, Na, Ba, Sr, K, and Cu europium oxide [Eu2O3] concentrated HCl (for cleansing flame test wires)

2 in./lab 5 mL enough for demonstration 10 mL/lab

Common Materials candle incandescent bulb in socket

1/lab 1/lab

Special Equipment flame test wires diffraction gratings hydrogen discharge tube helium discharge tube neon discharge tube mercury discharge tube ultraviolet light discharge apparatus earth light (optional)

1/solution 1/student 1/lab 1/lab 1/lab 1/lab 1/lab 1/lab 1/lab

Notes The room must be totally dark in order for the experiment to work well. Sometimes the students hold the diffraction grating turned 90o from the correct orientation. Check each student's view, but this is best done as a demonstration experiment. There may be a hazard from UV light, when burning magnesium or using the mercury discharge tube. The teacher must have adequate light sources and equipment; maybe Physics can supply some.

The HCl is only necessary for cleaning the flame loops. It is not necessary to have it for each pair of students.

Nitrate salts can also be used for the flame tests.

Diffraction gratings are available from Flinn Scientific, Batavia, IL and from Frey Scientific, 905 Hickory Lane, Box 8101, Mansfield, OH 44901-8101.

A new light source not mentioned is the Earth Light by Phillips Petroleum. It gives white light, rather warm, from three phosphors, red, green, and violet, as shown by the diffraction grating. An Earth Light may be purchased from Phillips Lighting, P.O. Box 156, Orangeburg, NY 10962-0156.

Disposal There should not be any disposal to deal with if the test solutions are capped tightly and saved from semester to semester.

15

ATOMS AND LIGHT PRE-LAB QUESTIONS 1) Name several ways in which atoms can be excited.

Heating, electricity, energy from chemical reactions, and absorbing electromagnetic radiation are all ways in which electrons can be excited.

2) When a fuel such as wood is burning in a fireplace, visible light is one of the products of the reaction. What is the source of this light we call the “flame?” (Give your answer in the form of a sequence of events.)

Heat released in the reaction excites the electrons. It also warms the gaseous products so that they rise from the fuel. As the electrons in the atoms reemit the energy that first excited them, they do so as wavelengths of electromagnetic radiation that we perceive as colors of light.

3) How might the normal color of a flame mask the true color of a salt during a flame test?

If the fuel is a source that gives color to the flame, the color will be added to the spectrum resulting from the substance being tested. It would be best to use a fuel that burns with a colorless flame.

4) When the electric element on a home cook top is turned on, it first becomes light orange and then turns red hot. If more current is supplied, it can become hotter. If there was no restriction on the amount of current, at some point the temperature would be sufficient for the burner to glow white. What has happened to the number of individual colors of light being emitted if the burner appears white to the naked eye?

As the temperature increases, more and more different electrons in the atom are excited.

Each electron gives its own color, or colors, of light as the energy is reemitted. When enough different colors of the visible spectrum are present, the overall effect will appear to be white light 5) Video cameras measure colors in Kelvin temperatures. Why is this appropriate?

Colors of light are the key to the temperatures of stars. Since all electromagnetic radiation originates in stars, this is a true and universally accepted standard to colors of light.

ATOMS AND LIGHT SAMPLE REPORT SHEET I. Describe the view of each of the following through a diffraction grating: 1. Candle - complete spectrum 2. Burning magnesium - complete spectrum 3. Incandescent bulb (tungsten) - complete spectrum 4. Hydrogen - three lines, one blue, one violet, and one red. A green line can also be seen.

5. Helium - 2 green lines, 3 yellow-green, 6 yellow, 7 orange, and 3 red lines 6. Neon - several lines in red, orange, yellow and green 7. Mercury - 2 blue lines, 2 yellow-green lines, 3 yellow, 2 orange, and 1 red line II. Describe the color of salts of these metals in a flame: 1. Lithium - red 2. Sodium - orange 3. Barium - lime green 4. Strontium - scarlet 5. Potassium - violet (requires cobalt glass plate to see) 6. Copper - blue green 16

III. QUESTIONS 1. Are the results the same for candle light, magnesium light, and tungsten?

Yes. 2. Could the colors of the flames be used to identify the different metal ions?

Yes. 3. Why is a yellow shirt yellow?

The pigment in a yellow shirt reflects yellow light, so the shirt appears yellow.

4. Large mercury vapor lamps are often used as night time security lights. What color would you expect these to appear to the naked eye?

The color of the lamp is the same as the color of the mercury discharge tube – light aqua/ blue.

5. What applications can you think of for europium compounds?

Student responses will vary. Possibilities include TV monitors, computer monitors, and red indicator lights.

6. What can we learn from a study of light from different stars?

By studying the light given off by stars, the elements produced can be determined and then from these elements, the temperature can be estimated. We can also determine how fast the stars are receding from the Earth.

17

#6

Flame Tests and Analysis

Reagents 1M chloride salt solutions of Na, K, Li, and Cu 1M iron sulfate [FeSO4] 18M sulfuric acid [H2SO4] 1M calcium nitrate [Ca(NO3)2] 1M silver nitrate [AgNO3] 1M hydrochloric acid [HCl] conc. hydrochloric acid [HCl] unknown salts 6M nitric acid (HNO3) 1M sodium hydroxide (NaOH)

10 mL/pair 0.5 mL/pair 0.5 mL/pair 5 mL/pair 5 mL/pair 5 mL/pair 10 mL/pair (for silver reclamation) (for silver reclamation)

Common Materials notebook paper

1 sheet/pair

Special Equipment flame test wires filter paper

1/pair (for silver reclamation)

Laboratory Equipmentl laboratory burner 2 50-mL beakers 10-mL graduated cylinders small test tube 1 400-mL beaker 1 pair forceps Notes An alternative to having each pair of students test each solution at their own lab area is to set up an area with Bunsen burners and the solutions. In this arrangement, each test solution should be supplied with its own flame loop. This also reduces contamination of the loops and preserves the integrity of the tests for the students.

The iron sulfate should be made fresh the day of use.

Caution: Store nitrates away from acids.

Disposal The calcium sulfate, silver chloride, and silver nitrate precipitates can be placed into the trash. Pour the solutions down the drain with plenty of water.

An alternate method for silver disposal is to reclaim it as metal. Dissolve the silver salt in 6 M nitric acid (HNO3). Neutralize with 1 M sodium hydroxide (NaOH) to pH 7. Place a clean copper strip in the solution. Set for twenty minutes. The silver metal will wipe off the copper easily. Wipe the silver into a filter cone (see Investigation #5). Filter the silver metal. Allow it to dry.

18

FLAME TEST AND ANALYSIS PRELAB QUESTIONS 1) Would a flame test work to identify the specific substances in a solution if it contained multiple cations? Explain your answer.

Not with the naked eye because the different colors would be seen simultaneously and each would mask the other. However, if viewed through a spectroscope, the individual lines in the spectra could be used to make a proper identification of the elements present.

2) What is the advantage of a qual scheme that not only identifies, but simultaneously precipitates, a particular ion?

If the ions are precipitated as they are identified, they can be further analyzed, recycled, or simply discarded as impurities that have been removed from the sample.

3) Why would it be important for a technician at a water treatment plant to check the incoming water often for the presence of barium, lead, mercury, or other toxic metallic cations?

These metals should not be present in any water system, either public municipal water supplies, or waterways in nature. Because they are persistent in the environment and are toxic to most all living things, they should not be allowed to enter the food chain at any level.

4) How might the plant operators deal with the presence of such ions?

When undesirable cations are detected in a water system, the technician could add an anion which is insoluble in combination with the ion in question, thus precipitating it.

5) The chloride ion concentration in ocean water is quite high; nearly 55% of the ions present are chloride. What would you expect to be the fate of any ions of silver that found their way into the ocean? Explain your reasoning.

Silver is quite insoluble as silver chloride and would be precipitated immediately upon entering the salt water.

FLAME TEST AND ANALYSIS SAMPLE REPORT SHEET I. FLAME TEST COLORS Sodium Potassium Lithium Copper Unknown Unknown

__orange___ __violet ___ (requires a cobalt glass to view) __red______ __blue green __________ __________

II.* Nitrate test

POSITIVE __________

NEGATIVE __________

Sulfate Chloride Acetate

__________ __________ __________

__________ __________ __________

III.*

SALT

CATION __________

ANION __________

*These answers will depend on the unknown salt used.

19

IV. QUESTIONS 1. What are some situations where an analysis of material is necessary?

Student responses will vary. Possibilities include determination of contaminants in water, input into manufacturing processes, quality control, environmental studies.

2. Is a scheme necessary in those situations?

Yes 3. Why is the Ca(NO3)2 added before AgNO3?

AgNO3 will cause a precipitate while Ca(NO3)2 will not.

4. What would happen if AgNO3 were added before Ca(NO3)2? (Hint: Ag2SO4 is an insoluble compound.)

The Ag2SO4 and AgCl would precipitate and the test for sulfate would not work.

5. Why is the brown ring test done on the original solution instead of the solution left after the sulfate and chloride ions have been removed?

The nitrate is added to the solution that would cause it to test positive for that ion.

20

#7

Ionic versus Covalent

Reagents ethanol [CHeCH2OH] 18 mL/pair covalent (pick three or choices of your own) 2 g of each/pair table sugar styrofoam salicylic acid thioacetamide paraffin Crisco ionic (pick 3 or choices of your own) 2 g of each/pair sodium chloride [NaCl] potassium iodide [KI] sodium sulfate [Na2SO4] potassium chloride [KCl] calcium chloride [CaCl2] sodium hydrogen carbonate [NaHCO3] (sodium bicarbonate) Epsom salts [MgSO4·7 H2O] Common Materials aluminum foil can lids (optional)

4"×4"/pair

Laboratory Equipment ring stand ring wire gauze 6 small test tubes microscale conductivity tester 10-mL graduated cylinder Evaporating dish Special Equipment Notes Covering tin can lids with the foil forms a stable and flat surface and prevents the substances from mixing together as they are heated.

There are other choices besides the ones that are listed. These were easily available at the time of testing. Pick choices that are easily available.

The students may have difficulty “seeing” if the compounds dissolve if they put too much into the water or ethanol because the solution becomes saturated. Be sure to emphasize the importance of dropping only one or two grains into the solvent.

Heating the compounds too long will result in burning of the sugar and decomposition of some ionic compounds. It might be wise to instruct the students to remove the heat source as soon as the first three have melted. .

21

Disposal All of the listed substances except thioacetamide can be thrown into the trash. All of the solutions can be flushed down the drain with lots of water. Thioacetamide should be mixed with a sodium hypochlorite solution and allowed to cool before being flushed down the drain with plenty of water.

IONIC VS. COVALENT PRELAB QUESTIONS 1) Table sugar melts at a much lower temperature than does table salt. Which of these two would you most expect to form an electrolytic solution?

Table salt because high melting point is a characteristic of ionic compounds and ionic compounds generally dissociate to form electrolytic solutions.

2) The labels have fallen off of two containers on a lab counter. The label from one container reads “potassium iodide”, an ionic compound. The other container label reads “dextrose”, a covalent compound and simple sugar. Describe two tests you could perform to determine which container should receive each label.

• Heat small samples of each compound and compare melting points.

• Dissolve a small amount of each compound in water and check for electrolytic properties.

3) Ocean water contains a lot of different ions, but the majority of them are of only seven varieties. Sodium and chloride ions together make up nearly 82% of the ions present.

What property of ionic compounds is responsible for this presence?

Ionic compounds are highly soluble in polar substances such as water. Sodium and chlorine are common in the Earth’s crust and enter the run-off water from melting glaciers, rain, and other precipitation, eventually entering streams and rivers, and ultimately the oceans.

4) Why would we not find a large amount of silicon dioxide (generally accepted as sand) dissolved in ocean water?

Silicon dioxide is a covalently bonded compound and although the bonds are polar, the molecules are not polar, thus it does not dissolve in water.

IONIC VS. COVALENT SAMPLE REPORT SHEET I. PART A: MELTING POINT Test Tube Description of Substance sugar styrofoam salicylic acid thioacetamide paraffin Crisco sodium chloride potassium iodide sodium sulfate

white, granular large chunks white powder yellow, granular translucent chunks white gel white, granular white, granular white, granular

22

Melting Time (seconds) 5 6 10 5 3 3 30+ 20+ 30+

potassium chloride white, granular calcium chloride white to pink, granular sodium hydrogen carbonate white powder Epsom salts white, granular

30+ 30+ 30+ 20+

II. PART B: SOLUBILITY AND CONDUCTIVITY Test Tube Water Soluble Ethanol Soluble Conductivity Covalent sugar Y SLIGHTLY N styrofoam N N N salicylic acid N Y Y thioacetamide Y Y Y paraffin N SLIGHTLY N Crisco N Y N Ionic sodium chloride Y N Y potassium iodide Y N Y sodium sulfate Y N Y potassium chloride Y N Y calcium chloride Y N Y sodium hydrogen carbonate Y N Y Epsom salts Y N Y III. PART C: CLASSIFYING THE COMPOUNDS Test Tube Ionic or Covalent This depends on which substances are chosen.

Compound Name

IV. QUESTIONS 1. What syllable in the name of some compounds is a very good indicator of the class to which it belongs?

All the”- ide” endings denote that they are ionic compounds.

2. According to the data gathered in this lab, what is the most definitive property for use in classifying ionic and covalent compounds?

The melting times and conductivity gave the most accurate results with the exception of salicylic acid and thioacetamide for the conductivity test. (Organic compounds are covalent, but the acids will conduct because of the hydrogen ions that they lose.)

3. Consider the solubility properties of the compounds and apply this to hair care. When hair is rolled and then dried, the curl leaves the hair as soon as it gets moist again.

However, if the hair is treated with a permanent wave, the curl does not wash out. Which type of bond is involved in a “wet set” and which type is involved in a permanent set?

Ionic bonds are involved with the "wet set" because they dissolve in water. However, covalent bonds are involved with a permanent set, because they do not dissolve in water.

4. Is the statement that pure compounds that are liquid are covalent true? If so, give three examples.

Mercury and bromine are elements and are liquid. Almost no ionic compound is liquid.

Many covalent compounds are liquid including water, several alcohols, hexane, ether, and other organic compounds.

23

#8

Chemical Reactions

Reagents copper carbonate [CuCO3] sodium bisulfate [NaHSO4] potassium ferrocyanide [K4Fe(CN)6] (or sodium ferrocyanide) 1M copper sulfate [CuSO4] 1M potassium iodide [KI] 6M hydrochloric acid [HCl] 6M sodium hydroxide [NaOH] pH test paper copper turnings 6M nitric acid ammonium sulfide [(NH4)2S]

0.1 g/pair 0.4 g/pair 0.2 g/pair 20 mL/pair 1.0 mL/pair 10 mL/pair 27 mL/pair .05 g/pair 10 mL/pair 1.0 g/pair

Common Materials steel wool, fine

0.4 g/pair

Special Equipment filter paper

1 sheet/pair

Laboratory Equipment 2 small test tubes 2 large test tubes 50-mL beaker 10-mL graduated cylinder eye dropper stirring rod Notes Caution: Store nitric acid away from copper and ammonium sulfide.

Disposal The solutions can be flushed down the drain with plenty of water and the solids can be placed into a waste container to be buried in a toxic waste dump.

24

CHEMICAL REACTIONS PRELAB QUESTIONS 1) A chemical reaction is defined as a process in which the atomic and electronic arrangements of substances are changed. If a substance does not change during a process, should it be represented as a reactant or a product in the equation?

No, unchanged substances are not represented in a proper chemical equation.

Sometimes only an ion appears on one side of the equation while it is part of a compound on the other side. The missing part of the compound from which the ion came was obviously left unchanged in the reaction, so it does not appear on either side of the equation.

2) Many times a substance such as a catalyst is required for reaction, but is not changed in the course of the reaction. Where would be an appropriate place to indicate the presence of the substance in the equation?

A catalyst is a necessary “condition” for a reaction and is placed over the yield sign with other conditions such as temperature or pressure.

3) When testing the pH of a substance, a drop of the substance is added to a piece of pH paper. A color change is the indicator of pH. Could the change in the pH paper, itself, be a reaction? Support your answer.

Since the color of the substance changes with a change in pH, it is probably a reaction.

The absorption or reflection of specific wavelengths of light is a property that has changed.

4) Some reactions are very subtle and give almost no indicators. Others are so over-the-top that there is no doubt a reaction has occurred. Consider the chemical processes involved in a 4th of July fireworks display. List all of the indicators of a reaction that you can think of that are evident in the explosion of a firework.

Energy changes are evident as heat, light, and sound. There are color changes in some displays. When the firework is initially launched, the rocket portion produces expanding gases. New particulate matter is seen as smoke.

5) Give an example of each of the following changes that would NOT be a chemical change: a) temperature change heating a cup of coffee in a microwave b) formation of a gas boiling water to produce steam c) formation of a solid freezing water to form ice d) color change adding food coloring to cookie dough

25

CHEMICAL REACTIONS SAMPLE REPORT SHEET I. COLOR CHANGE 1. Complete Cu2+ + Fe → Fe 2+ + Cu (s) 2. Where did the iron ions come from?

The iron ions came from the steel wool.

II. DISAPPEARANCE OF METAL 1. What does the color of the solution indicate?

The blue color of the solution indicates the presence of copper ions.

2. Where did the copper go?

The copper goes into solution as Cu2+ ions.

III. NEUTRALIZATION 1. Initial pH _____2_______ Final pH _____10 ______ IV. PRECIPITATION 1. Write the reaction.

Cu2+ + ZnS → Zn2+ + CuS 2. What is the color of the final liquid?

The final liquid should be colorless. The precipitate will be some form of blue (milky, chalky, dark, etc.).

V. HEAT PRODUCTION 1. pH 6 M HCl _____2______ 6 M HCl + 2 mL NaOH _____2______ 4 mL NaOH _____2______ 6 mL NaOH _____2______ 8 mL NaOH _____2______ 10 mL NaOH _____2______ 12 mL NaOH _____10_____ 14 mL NaOH _____10_____ 6 M NaOH _____10 2. Write the reaction.

HCl + NaOH → H2O + NaCl 3. Why were you instructed to place the beaker on the lab bench?

The beaker was placed on the lab bench because it gets very hot.

VI. GAS PRODUCTION 1. Write the reaction.

2NaHSO4 + Fe → 2Na+ + (SO4)2- + H2 + Fe2+ 2. Where did the hydrogen come from?

The hydrogen came from the sodium hydrogen sulfate reacting with the iron.

This formed H2(g).

3. Where did the iron ions come from?

The iron ions came from the steel wool.

26

#9

Double Replacement Reactions

Reagents 1 M copper nitrate [Cu(NO3)2] 1.5 M sodium hydroxide [NaOH] 0.5 M hydrochloric acid [HCl] sodium bicarbonate [NaHCO3] 1 M calcium nitrate [Ca(NO3)2] 1 M sodium carbonate [Na2CO3] 1 M potassium iodide [KI] 1 M lead nitrate [Pb(NO3)2] 1 M copper (II) sulfate [CuSO4] solid sodium sulfite [Na2SO3]] 1 M sodium chloride [NaCl] phenolphthalein

2 mL/pair 3 mL/pair 27 mL/pair 0.5 g/pair 2 drops/pair 6 drops/pair 10 drops/pair 2 drops/pair 6 drops/pair 0.5 g/pair 2 drops/pair 1 drop/pair

Common Materials clear plastic sleeve letter size a sheet of plain white paper

1 sheet/pair 1 sheet/pair

Laboratory Equipment 10-mL graduated cylinder 100-mL graduated cylinder 3 small test tubes

droppers for solutions 50-mL beaker wash bottle

Notes This laboratory exercise works well in conjunction with the Single Replacement Reactions exercise. It can be started as soon as the “5 minute” check has been made in the other lab.

Disposal Except for the lead solutions, the solutions from this investigation are not extremely toxic and can be poured down the drain with plenty of water. The small amount of lead solution used in this experiment can be absorbed by a paper towel and thrown into the trash. It should not be washed down the drain.

DOUBLE REPLACEMENT REACTIONS PRELAB QUESTIONS 1. When solutions of two dissolved salts are combined and result in the formation of an insoluble salt in the form of a precipitate, where are the remaining two ions?

The remaining two ions are not changed in the reaction and are referred to as spectators.

2. Suppose the precipitate in question #1 is filtered from the solution that originally contained four ions. How could the salt that remains in solution be recovered?

The liquid in the solution, i.e., the solvent that is probably water, could be evaporated. The remaining two ions would then be free to form an ionic solid.

3. If there are ions of several soluble salts in one solution, could you devise a way to separate only one of the ions from the solution?

If there is another ion that forms an insoluble salt with the desired ion but not with any of the others, the ion in question could be precipitated as a solid and filtered to remove it.

4. Why is the volume of reactants used not important to the results in Part A of this lab?

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The procedure is not intended to be quantitative, that is a measurement of how much ion is present, but rather qualitative. The production of a precipitate is not dependent upon how much of the two ions are introduced into a common vessel.

5. From the procedure in parts B and C, why do you think there are generally cautionary statements on any household products that are moderately or strongly acidic or basic?

The heat that can result from diluting or neutralizing some acids and bases can cause splattering and spewing. There are also gases produced from the combination of specific acids and bases that are harmful. (Especially anything containing hydrochloric or muriatic acid, HCl, and the base ammonia as the result is a caustic gas, ammonium chloride.)

DOUBLE REPLACEMENT REACTIONS SAMPLE REPORT SHEET I. PART A: Predictions and Results Based Upon Solubility Rules SOLUTIONS

PREDICTION

OBSERVATION

PRECIPITATE

calcium nitrate & sodium carbonate

Y

cloudy white

CaCO3

potassium iodide & lead nitrate

Y

mustard yellow

PbI2

copper (II) sulfate & sodium hydroxide

Y

cloudy and blue

Ca(OH)2

sodium chloride & potassium iodide

N

nothing

none

sodium carbonate & potassium iodide

N

nothing

none

copper sulfate & sodium carbonate

Y

cloudy and gel like

CuCO3

potassium iodide & sodium hydroxide

N

nothing

none

copper sulfate & potassium iodide

Y

brown

CuI2

II. PART B: Gas Forming Reactions Temperature

Temperature

HCl Solution __________ NaHCO3 and HCl mixture __________ Observations Bubbles form rapidly, temperature drops.

Na2SO3 and HCl mixture

__________

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Observations Bubbles form rapidly, temperature drops.

III. PART C: Acid-Base Reactions Indicator Color HCl

colorless____

HCl and NaOH mixture Observations Slight temperature rise

pink________

IV. QUESTIONS 1. Out of the eight precipitation reactions, how many did you correctly predict? _______ 2. In how many of those were you able to correctly determine the precipitate even before observing the reaction? ________ 3. Describe the similarities and differences of the precipitates that were formed. Did you find any characteristics consistent with compounds containing a particular ion?

The copper carbonate was gel like. The precipitates are various colors. Sodium and potassium are always soluble.

4. Explain why such tools as the Periodic Table, Activity Series, and solubility charts are important to the work of a research chemist.

These tools let chemists know what reactions will occur without having to try them out.

5. Write a balanced equation for each reaction that occurred in Part A. (If no precipitate formed, give the reactants and write “no reaction” after the yield sign.) a. Ca2+ + CO32− → CaCO3 b. Pb2+ + I− → PbI2 c. Cu2+ + OH− → Cu(OH)2 d. Na+ + Cl− + K+ + I− → no reaction e. Na+ + CO32− + K+ + I− → no reaction f. Cu2+ + CO32− → CuCO3 g. Na+ + OH− + K+ + I− → no reaction h. Cu2+ + I− → CuI2

29

#10

Iron(II) and Iron(III) Ions

Reagents 0.4 M cobalt(II) nitrate [Co(NO3)2] 12 M hydrochloric acid [HCl] 0.1 M iron(III) chloride [FeCl3] 10.2 M iron(II) ammonium sulfate [Fe(NH4)2(SO4)2] 0.1 M potassium thiocyanate [KSCN] 0.1 M potassium ferrocyanide [K3Fe(CN)6] 0.1 M potassium ferricyanide [K4Fe(CN)6]

5 mL/pair 15 mL/pair 5 mL/pair 5 mL/pair 5 mL/pair 5 mL/pair 5 mL/pair

Common Materials (for Microscale version) clear plastic sleeve letter size a sheet of plain white paper

1 sheet/pair 1 sheet/pair

Laboratory Equipment large test tube small test tubes test tube rack watch glass 10-mL graduated cylinder stirring rod

1/pair 6/pair 1/pair 1/pair 1/pair 1/pair

Notes This laboratory exercise works well in conjunction with the Double Replacement Reactions investigation.

It is VERY important that all solutions be freshly prepared.

The procedure can be performed in a Microscale fashion with the use of white paper in a clear plastic sleeve as the reaction surface for Parts B & C, and a microtitration plate for Part A. Replace all volumes in the procedure with “drop” in place of “mL.” It is easiest to fill and label beral pipets with the solutions and supply them for each pair of students.

Disposal The small amount of solids generated in this experiment can be absorbed by a paper towel and thrown into the trash. They should not be washed down the drain. Solutions from this investigation are collected in a marked container and then precipitated with sulfuric acid, filtered, and disposed of in a land fill. The remaining filtrate is not toxic and can be poured down the drain with plenty of water.

IRON(II) AND IRON(III) IONS PRELAB QUESTIONS 1) According to the information in this investigation, could a chemist ever be certain of the oxidation state of a cobalt ion from its color alone? Explain.

Since some cobalt ions vary in color depending upon the nature of the surrounding environment, no, color alone is not enough to determine the oxidation number of a cobalt ion.

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2) In your own terms, briefly explain the differences in monatomic ions, polyatomic ions, and complex ions.

Monatomic ions are composed of a single atom with a charge. Polyatomic ions are a group of covalently bonded atoms with a charge that is assigned to the group and not a single atom.

Complex ions are composed of individual ions and /or small molecules that are grouped together with a single assigned charge.

3) How might it be an advantage to a chemist if the color of a substance was a definite indication of its oxidation state?

If color alone was a definitive indication of the oxidation state of an ion, many analytical processes could be simplified.

4) Which of the reagents for this lab contain polyatomic ions?

Co(NO3)2, Fe(NH4)2(SO4)2, KSCN, K3Fe(CN)6, and K4Fe(CN)6.

IRON(II) AND IRON(III) IONS SAMPLE REPORT SHEET I. EXPERIMENTAL DATA PART A: Cobalt(II) Ions mL 12 M HCl

Color

1 2 3 4 5 6 7 8 9 10 11 12

___peachy-pink____ ________________ ____pink________ ___pinkish-violet__ ________________ _____violet_______ ________________ ___blue-violet_____ ________________ ____very blue_____ ________________ ________________

PART B: Iron(II) and Iron(III) Thiocyanide Compounds Compound Fe3+ + SCNFe2+ + SCN-

Color ___teal-green_____ __yellowish-green__

PART C: Tests for Iron(II) and Iron(III) Ions Iron Ion Fe3+

ferricyanide (Fe3+)

ferrocyanide (Fe2+)

K3Fe(CN)6

K4Fe(CN)6

deep blue

blood red

orange red

teal blue

FeCl3

Fe2+

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Fe(NH3)2(SO4)2

II. QUESTIONS 1. At what point in the addition of HCl to the cobalt solution did the concentration of [Co(H2O)6]2+ roughly equal the concentration of [CoCl4]2-? How could you tell?

At about 6 mL the color became half way between pink and blue - it was violet. This would suggest half pink ion and half blue ion.

2. What visual result would you expect to see from an increase in the relative amount of water in a solution made of ethanol (alcohol) and water, plus the [CoCl4]2- and [Co(H2O)6]2+ ion combination?

An increase in the percentage of water molecules surrounding the cobalt complexes would probably encourage more of the water-containing complexes to form and that would make it turn blue.

3. Why is the thiocyanate ion test only useful for a solution that contains a single iron ion?

If both the 3+ and 2+ ions are present, the darker color of the 3+ ion would mask the lighter hue of the 2+ ion.

4. What test(s) could you perform to prove beyond doubt that a solution contained both Iron(II) and Iron(III) ions?

In each of two samples of the solution to be tested, add one of the two (ferroocyanate and ferricyanate) ions. If both ions turn the solution blue, then both ions are present. If one is red, only one ion is present.

5. If a chemist is presented with a solution that is known to contain ionized iron atoms, how many steps are necessary to determine if it contains only one, or both, iron ions? What are they and how would you proceed?

Since the thiocyanate ion is inconclusive for both ions at once, it would be best to use only the ferro- and ferricyanate ion tests. The same procedure as for question #4 will work.

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#11

Mole Relationships

Reagents 1M calcium chloride [CaCl2] 30 mL/pair 1M sodium sulfate [Na2SO4] 40 mL/pair 1M calcium nitrate [Ca(NO3)2] 20 mL/pair 1M copper sulfate [CuSO4] 20 mL/pair 1M magnesium chloride [MgCl2] 10 mL/pair 1M sodium hydroxide [NaOH] 10 mL/pair solid sodium bicarbonate [NaHCO3] 0.84 g/pair 1M acetic acid 10 mL/pair Common Materials large party balloons 1/pair Special Equipment filter paper 4 pieces/pair Buchner funnel Laboratory Equipment centigram balance 10-mL graduated cylinder 500-mL Erlenmeyer flask with stopper to fit 250-mL Erlenmeyer flask 1 large test tube Disposal Pour together all the solutions of calcium and sulfate. Filter the calcium sulfate precipitate. Discard all precipitates in a container to be buried in a landfill. Pour the solutions down the drain with lots of water.

MOLE RELATIONSHIPS PRELAB QUESTIONS 1) Refer to procedural steps 1 and 2. How could the sodium chloride (Na+ and Cl- ions) that remains dissolved in the solution be recovered?

The solvent (water) could be evaporated to leave the ionic salt crystals behind.

2) The dissociation of carbonic acid to form water and carbon dioxide gas in step 4 is a major component of the soft drinks we consume. What is happening to the acid content of a coke when the cap is removed and the carbon dioxide gas able to escape?

The acid content is reduced and the drink loses the component of the sharp and bitter flavor that is contributed by the combined presence of carbon dioxide gas and the acid.

The drink will taste much sweeter when it is “flat.”

3) What remains in solution at the end of step 5?

Sodium Chloride 4) If it was determined that the water supply for a city had a minute level of lead ion contamination, the water treatment plant would likely attempt to drop the lead out by precipitating it. Would they want the negative ion they add to the water or the existing lead ion to be the limiting reactant? Explain your reasoning.

No, it would be more efficient to make the lead the limiting reactant by adding more of the negative ion than is needed. This will also function to help ensure that the maximum amount of lead is removed from the system.

33

MOLE RELATIONSHIPS SAMPLE REPORT SHEET I. PRECIPITATION Flask mass: before reaction 116.87 g after reaction 116.84 g I. GAS PRODUCTION Flask mass: before reaction 156.72 g after reaction 156.59 g III. LIMITING REACTANT Record the mass of CaSO4 + Filter paper for each run: Run #1 10 m mol Na2SO4 + 10 m mol CaCl2 9.69 Run #2 20 m mol Na2SO4 + 10 m mol CaCl2 9.37

g g

Run #3 10 m mol Na2SO4 + 20 m mol CaCl2 10.51 g Precipitation in filtered solution Mark each solution that produces a large amount of precipitate.

CuSO4 Ca(NO3)2 Run #1 10 m mol Na2SO4 + 10 m mol CaCl2 _________ Run #2 20 m mol Na2SO4 + 10 m mol CaCl2 ___X____ Run #3 10 m mol Na2SO4 + 20 m mol CaCl2 X _________ IV. QUESTIONS 1. What is the relationship between the masses of reactants and products from the precipitation flask reaction?

The mass of the products is less than the mass of the reactants, but they are still very close because the difference in the fifth digit is negligible.

2. What is the relationship between the masses of reactants and products from the gas production flask reaction?

The product mass is less because gas was produced and can leak from the balloon, but they are still very close.

3. Write the balanced equation for the two flask reactions.

MgCl2 + 2NaOH → Mg(OH)2 + 2NaCl NaHCO3 + CH3COOH → CO2 + H2O + NaCH3COO 4. The three CaSO4 mass should be the same except for a small water mass difference. Explain this statement.

The amount of CaSO4 is the same in all 3 reactions because the limiting reactant in each reaction determines the amount of product which can be formed. In the second the Ca2+ is the limit in run #2 and the SO42- is the limit in t#3.

5. What compound was in excess in run #2?

Na2SO4 6. What compound was the limiting reagent in run #2?

CaCl2 7. What compound was in excess in run #3?

CaCl2 8. What compound was the limiting reagent in run #3?

Na2SO4

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#12

Molar Volume of a Gas

Reagents magnesium ribbon 6.0 M hydrochloric acid [HCl]

about 15 cm/pair about 35 mL/pair

Common Materials sand paper thread

1 piece/pair 3 12” lengths/pair

Special Equipment 50 mL eudiometer tube 1-hole rubber stopper to fit eudiometer 1 liter graduated cylinder wash bottle with distilled water barometer

1/pair 1/pair 1/pair 1/pair

Laboratory Equipment 10-mL graduated cylinder 600-mL beaker ring stand utility (or buret) clamp thermometer centigram balance Notes The thickness of magnesium ribbon varies slightly from supplier to supplier. It would be wise to do a trial run on the lab to ensure that the length of magnesium used will not produce more than about 48 mL of hydrogen gas.

Disposal There are no harmful substances at the end of the lab. The fluid from the eudiometer tube can be flushed down the drain with plenty of water.

MOLAR VOLUME OF A GAS PRELAB QUESTIONS 1) If gases expand to fill any container, what invisible barrier keeps an atmosphere around the Earth?

Gravity holds the atmosphere close to the earth. The atmosphere is denser closest to the Earth’s surface because that is where gravity is the strongest.

2) If the gas collected dissolves in water, as does oxygen or carbon dioxide, what effect will that property have on the final experimental value for the volume of gas produced?

The measured volume will be smaller than the actual volume of gas produced because some will remain dissolved in the water.

3) A eudiometer tube is marked in graduations of 0.1 mL. If a water meniscus sits on the third mark above the 46 mL mark, what is the volume reading?

35

45.7 mL because the eudiometer is marked to show the amount of gas collected with the 0.0 mark at the top.

4) If bubbles of the gas are allowed to remain on the inner surface of the tube and on the thread, what effect will that have on the measured volume of gas in the tube?

The effect will be the same as for dissolved gases; they will not be included in the measured volume.

MOLAR VOLUME OF A GAS REPORT SHEET Data Table Trial 1

Trial 2

Mass of magnesium (g)

0.04343 g

Gas Volume (mL)

38.72 mL

Temperature (oC)

19.2 o

Barometric Pressure (inHg)

28.8 inHg

Trial 3

Calculations 1) Find the moles magnesium consumed by dividing the mass of magnesium by its molar mass (MMMg = 24.31 g/mol). Record this value in the calculations table.

2) Convert your volume in mL to Liters by dividing by 1000mL/L. Record this value in your calculations table.

3) Your thermometer is probably calibrated in degrees Celsius. Convert the experimental temperatures to Kelvin by adding 273.15. Record these values.

4) If your barometer is calibrated in inches of mercury, convert the measurement to centimeters of mercury by multiplying by 2.54 cm/in. Then rewrite the measurement in millimeters of mercury, “mmHg”, by multiplying by 10. Record this value in the calculations table.

5) Find the experimental temperatures on the table of equilibrium vapor pressures for water and subtract the partial pressure of the water to get the “dry” pressure of the hydrogen gas for each trial. Record these values.

6) Convert the dry gas pressures from torr (mmHg) to atmospheres (atm) by dividing the value by 760 mmHg/atm. Record these values.

7) Use the relationship n = (PV)/(RT) where “P” is the corrected pressure in atm, “V” is the experimental volume in liters, “T” is the Kelvin temperature, and “R” is 0.08206 L.atm/mol.K to find the moles of hydrogen gas collected. Record these values.

8) Examine the balanced equation for the reaction between magnesium and hydrochloric acid in the Experimental Plan. There is a 1:1 mole ratio of magnesium consumed to hydrogen gas produced. Using your determination of the moles magnesium consumed in the reaction as the theoretical, or expected, yield “T”, and the moles hydrogen gas actually produced as “A”, find a percent yield for each trial using A/T x100% = % Yield.

36

9) Use the Combined Gas Law to find the equivalent volume of hydrogen gas at standard conditions. Use the relationship V’ = (PVT’)/(TP’) where P, V, and T are your experimental values and T’ is 273 K, while P’ is 1.00 atm. Record this as the volume of a gas at standard conditions (Standard Volume).

10) Using a ratio, find the experimental volume of one mole of hydrogen gas at standard conditions (Exp. SMV), where EV is your experimental standard volume of hydrogen, and Emol is the experimental moles of hydrogen as follows: EV/Emol = SMV/1.00mol or (EV/Emol)/1.00 mol = SMV 11) Using 22.4 L and the accepted value “A,” and your Exp. SMV as the experimental value “E”, find a percent error for your value of the Standard Molar Volume of a gas using the relationship: % Error = |A-E|/A x 100% where |A-E| is the absolute value of the difference in the accepted and experimental values. Record this value in the table.

Calculations Table Trial 1 Moles of magnesium (mol)

0.0017869 mol

Gas Volume (L)

0.03872 L

Temperature (K)

292.4 K

Barometric Pressure (mmHg)

731.4 mmHg

Vapor Pressure of Water (mmHg)

16.7 mm Hg

Dry Pressure of Hydrogen (mmHg)

714.7 mm Hg

Dry Pressure of Hydrogen (atm)

0.940 atm

Moles Hydrogen Collected (mol)

0.0015178 mol

% Yield for Hydrogen Gas

84.9 %

Standard Volume of Hydrogen (L)

0.03402 L

Experimental SMV of a Gas (L)

22.41 L

% Error for SMV of a Gas

0.04464 %

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Trial 2

Trial 3

Post-Lab Questions 1) Why was it necessary to tie the magnesium and pin it with the thread? What does this say about the relative densities of the metal and the water?

The magnesium is less dense than water and will float to the top when the tube is inverted. If this happens, it will not react with the acid.

2) When the eudiometer tube was being raised in the large cylinder to match the water levels inside and outside, what happened to the inner water level? Use gas pressures to explain this observation.

As the eudiometer tube is raised to make the water in the tube match the level of the water outside the tube, it is the higher pressure of the gas on the surface of the water in the open container that supports the water in the tube. As the water levels get closer, less pressure is exerted through the water in the tube and the pressure on the gas is therefore reduced so that it expands, lowering the water level in the tube.

3) Examine the table of equilibrium vapor pressures for water. Assume all measurements for the hydrogen gas were kept constant except for the room temperature. Would an increase of 5 degrees have meant the presence of more or less hydrogen gas?

At a higher temperature the amount of water vapor would increase. If the pressure and volume are constant, this means less of the gas mixture is actually hydrogen.

4) Give an explanation for the fact that some hydrogen bubbles could actually escape down through the hole in the stopper.

As the water is forced out of the bottom of the tube it flows through the little opening in the stopper rather quickly. Some bubbles can get caught in the “current” and exit the tube with the water.

5) List as many sources of error as possible in the data for your experimental value for the moles hydrogen produced and for the determination of the standard molar volume of a gas.

Bubbles of hydrogen gas stuck on the thread, stopper, and sides of the eudiometer; hydrogen gas bubbles escaping through the stopper; impurities remaining on the magnesium ribbon, human error in measurements of the mass of the magnesium ribbon, temperature, barometric pressure, and volume of gas.

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#13

Hydrated Compounds

Reagents hydrated copper sulfate crystals [Cu(SO4).5H2O] hydrated magnesium sulfate crystals [Mg(SO4. 7H2O]

3 g/pair 3 g/pair

Common Materials Equipment evaporating dish watch glass tongs eye dropper ring stand with ring and wire gauze laboratory burner and tubing centigram balance desiccator

2/pair 2/pair 1/pair 1/pair 1/pair 1/pair

Notes Disposal HYDRATED COMPOUNDS PRELAB QUESTIONS 1) If heat is required to remove the water molecules from hydrated crystals, what would you expect to be the result of dropping water onto a dehydrated crystal?

If dehydration is an endothermic process, then re-hydration would be expected to be exothermic.

2) Salt allowed to sit undisturbed in a shaker can “cake” or become stuck together. This happens because most salt contains a bit of calcium chloride and magnesium chloride as minor impurities. Explain what is happening.

The calcium chloride and magnesium chloride are hygroscopic and function as desiccants.

3) Once the heating of a mass of hydrated crystal is completed, it will be placed into a desiccator to cool. From the name of the device, what do you think it is purpose?

A desiccant is a substance that absorbs moisture from the surrounding atmosphere. The dissicator will absorb moisture from the air that is sealed within it as a second material dries and thereby keep humidity of the air low enough to continue removing moisture from the second material.

4) In the bottom of the desiccator is a solid substance. Sometimes this substance is a commercial product called DririteTM, but sometimes it is anhydrous calcium chloride.

What is the function of the substance in the bottom of the desiccator? Name another substance that could be used.

The substance functions to absorb moisture from the surrounding air so that it facilitates the drying of any substance sealed in the desiccator. Magnesium chloride could also be used.

5) Does the substance in the bottom of the desiccator necessarily need to be discarded once it has fully hydrated? What is an alternative?

The desiccant could be heated to drive out the moisture it has absorbed and thus would be “regenerated.”

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6) The manager of a laboratory stockroom might prefer to have the hydrated form of a crystal rather than the anhydrous form. What might be a practical reason for that preference?

A hygroscopic or deliquescent substance will absorb any moisture that comes into contact with it and thereby change its formula weight and properties. It is much more efficient to store a fully hydrated crystal whose constitution and mass will be constant.

HYDRATED COMPOUNDS REPORT SHEET Parts A & B Data Table Magnesium Sulfate

Copper Sulfate

Appearance before heating

white

blue

Appearance after heating

white

light gray

Observations during heating

condensation

condensation

Mass of dish and glass Mass of dish, glass, and salt before heating Mass of dish, glass, and salt after first heating

N/A

Mass of dish, glass, and salt after final heating Parts A & B Calculations 1) For each salt, calculate the mass of the hydrated salt by subtracting the mass of the dish and glass from the total mass before heating. Record these values in the Calculations Table on the next page.

2) For each salt, calculate the mass of the anhydrous salt by subtracting the mass of the dish and glass from the final mass after heating. Record these values.

3) For each salt, calculate the mass of water removed from the salt by finding the difference in the masses of the hydrated and anhydrous forms of the salt. Record these values.

4) For each salt, find the moles of anhydrous salt by dividing the mass of anhydrous salt by its molar mass. Record these values. (The molar mass of CuSO4 is 159.60 g/mol, and the molar mass of MgSO4 is 120.38 g/mol.)

40

5) For each salt, find the moles water that was removed by dividing the mass of the water removed by its molar mass (18.0152 g/mol.) Record these values.

6) For each salt, find the ratio of moles salt to moles water in the hydrate by dividing the moles water by the moles of anhydrous salt. Record these values expressed and “1: __”.

7) Find the percent error for your experimental values by the following formula: % Error = |A-E| / A x 100% where “A” is the accepted value and “E” is the experimental value for the hydrate number. The accepted hydrate values for magnesium chloride and copper sulfate are 7 and 5, respectively. Record the percents error in your calculations table.

8) Round the mole ratio to the nearest whole number to get a hydrate number and report your experimental empirical formula for the hydrated salt in the form of XY.nH20 where “n” is the experimental hydrate number.

9) Gather hydrate numbers from several classmates as comparison to your work.

Part A & B Calculations Table mass of hydrate mass of anhydrous salt mass of water removed moles of anhydrous salt moles of water removed ratio of moles salt to moles water % Error hydrate number exper. hydrate formula

Magnesium Sulfate

Copper Sulfate

<1% 7 MgSO.7H2O

<1% 5 CuSO4.5SO4

Post-Lab Questions 6) Once the mass of the clean and dry evaporating dish and watch glass have been recorded, they are handled only with clean tongs. Why is this important?

Any oils or foreign particles on the dish will alter the mass and change the final results.

7) What is the purpose of the watch glass?

The watch glass prevents loss of material from spattering.

8) If the final mass is taken while the dish is still hot, what would be the likely effect on the experimental ratio of moles water to moles salt? Would it be too high or too low?

Explain the logic of your reasoning.

Taking a mass on a substance while it is hot will cause the balance to read less than the actual mass. This would cause the results to appear as if there was more water removed than was actually removed. Therefore the experimental ratio of water to salt would be too high.

9) Examine data from some classmates. Does there seem to be consistency in the hydrate numbers you found for the two salts? Comment. What law does this data support?

There should be good consistency in all numbers and this is an example of the Law of Definite Proportions.

10) Why was the magnesium sulfate heated until the mass was constant?

41

If there is no indication of dryness by the color of the substance, the only method of determining if it is totally dehydrated is to heat it repeatedly until there is no more water to drive off.

11) A student runs an experiment in which she heats a pink cobalt chloride (CoCl2) crystal and measures the masses before and after heating. During the process she notices the pink color changing to blue and condensation on the watch glass covering the substance.

The masses she records are 3.00 and 1.64 grams. (Hint: The formula masses for cobalt chloride and water are 129.84 g/mol and 18.0152 g/mol respectively.) a) What three evidences are there that the substance was a hydrated crystal?

Condensation, reduced mass, color change b) Calculate the hydrate number for the substance.

5.99 12) From your observations when adding water to the copper sulfate, explain the energy changes taking place.

The substance returned to its original blue color, popped and hissed, and released steam.

From the sound and steam, it can be deduced that heat was released when the water was reabsorbed into the crystal.

13) Cobalt chloride has an interesting property. A CoCl2.4H2O is blue, while CoCl2.6H2O is pink. Can you think of a good use for this substance?

It is often used as an indictor for the moisture content of a desiccant.

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#14

Solubility

Reagents alcohol (ethyl or isopropyl) Common Materials syrup vegetable cooking oil liquid soap chewing gum peanut butter cloth Special Equipment

5 mL/pair 10 mL/pair 10 mL/pair 1 mL/pair 1 stick/pair 1 small scoop/pair 8 in. × 8 in./pair

Laboratory Equipment balance stirring rod 50-mL beaker 10-mL graduated cylinder eye dropper Notes Disposal The solutions in this investigation are not toxic and can be poured down the drain with plenty of water.

SOLUBILITY PRELAB QUESTIONS 1) If the liquid layers in the procedure were added to the cylinder in the reverse order, that is, with the densest first and least dense last, what would you expect to happen at the interface of each liquid?

The substances would naturally move to position the most dense on the bottom and the least dense on top. The movement would cause the water and syrup to mix so there would only be two layers after the substances rearranged themselves.

2) What properties would you expect the molecules of two miscible liquids to have in common?

They would have similar properties of polarity. They would either both be polar, or both be non-polar.

3) Suppose a garment has been stained by a substance that is made of non-polar molecules.

What kind of molecule would you expect to be most effective at removing the stain?

A non-polar molecule would be most efficient in removing a non–polar stain. Dry cleaning agents are non-polar solvents.

4) If a liquid with a density of 1.2 g/mL is mixed in equal proportions with a liquid that has a density of 0.9 g/mL, what would you expect of the combined density if a) the two liquids are completely miscible with additive volumes?

The new density would be an average of the two densities. b) the two liquids are completely miscible without additive volumes?

The new density would be between the two original densities, but not necessarily the average. c) the two liquids are completely immiscible?

Two distinct layers would form with the more dense substance on the bottom. d) the two liquids are miscible but of greatly varying viscosities?

The two substances would tend to form two layers with the more dense substance on the bottom, but the interface between the two layers might not be well defined because of some mixing.

43

SOLUBILITY SAMPLE REPORT SHEET I. DENSITIES OF LIQUIDS SYRUP WATER OIL mass of cylinder and liquid ________g ________g ________g minus mass of cylinder -________g -________g -________g mass of liquid ________g ________g ________g mass of liquid Density of liquids = 10 mL = __1.382__g/mL _1.034__g/mL __0.959_g/mL Which liquid is the most dense? syrup Which liquid is on the bottom of the beaker? syrup Which two layers mix in procedure 4? water and syrup TOP LAYER BOTTOM LAYER mass of cylinder and liquid ________g ________g minus mass of cylinder -________g -________g mass of liquid ________g ________g mass of liquid Density of liquids = 5 mL = __1.106__g/mL __1.094__g/mL Is the density of the mixture an average of the original two layers?

Yes Does soap act as an emulsifying agent?

Yes Why would forming an emulsion be important when washing greasy dishes?

The grease must be held to the water by the soap in order to remove it from the dishes.

II. Volume water 5 mL Volume alcohol 5 mL 10 mL Measured volume __9.5___mL If the volume is different, why?

The molecules will be held tightly together and will fit together closely enough to need less space. It is as if the H2O molecules fit into holes between the alcohol molecules.

What is in peanut butter that allows it to remove the gum?

Oil Why would using peanut butter be better to remove gum from hair than scissors?

Scissors would cut the gum out while peanut butter will remove the gum and leave the hair undamaged.

Why would using peanut butter rather than scissors be better for removing gum from hair?

The water will not dissolve the gum but the peanut butter will dissolve the gum.

Would vegetable oil work in a similar manner to peanut butter to remove gum?

Yes, oil would work, but it is a liquid. The solid peanut butter is easier to work with.

A substance called lecithin allows the vinegar and oil in some dressings to remain mixed.

What qualities would you expect it to have in common with soap?

It is also an emulsifier. In this way, lecithin can attract the oil with one end of the molecule and attract the polar molecules of the vinegar with the other end, allowing them to mix.

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#15

Viscosity

Reagents rubbing alcohol (isopropyl)

30 mL/pair

Common Materials soluble starch vegetable cooking oil honey dishwashing liquid ketchup glass marbles

30 g/pair 30 mL/pair 30 mL/pair 30 mL/pair 30 mL/pair 3/pair

Special Equipment stop watch test tube rack test tubes, large stopper to fit large test tube

1/pair 1/pair 11/pair 1/pair

Notes The soluble starch solution is made by stirring the powdered starch into cold water. Heat the solution with occasional stirring to a slow boil until the solution is clear. When the solution cools and thickens it will appear white and it will be ready to use. It cannot be made up more than one or two days ahead because it will mold easily.

Other common materials that work well are shampoo, conditioner, mustard, and liquid fabric softener. .

Disposal All the solutions in this experiment can be flushed down the drain with plenty of water.

VISCOSITY PRE-LAB QUESTIONS 1) Starch is a linear protein molecule. Knowing this, what might you expect to be the effect on the viscosity of a starch solution if it was allowed to sit for a while?

The linear molecules will settle easily and because of the amount of contact between the molecules, they may be able to form numerous and strong attractions to each other. The result would be that the solution would become more and more viscous as it sits.

2) If two objects of the same material, density, and mass, but of different shape are each allowed to fall through the same viscous fluid, how would shape effect the apparent measurement of the fluid’s viscosity?

It would be expected that the more aerodynamic shape would meet with less resistance and friction from the fluid and would therefore have a slightly increased speed.

3) Considering the cohesive forces between particles in a highly volatile fluid, that is one that evaporates easily, what would you suspect concerning the viscosity of the fluid? Explain your reasoning.

If the fluid evaporates easily, the cohesive forces must be small. It would then be expected that the viscosity is very low because cohesive forces are the major factor in viscosity.

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4) Air is not considered to be a very viscous fluid. However, automobiles and planes are designed with aerodynamics in mind. How does the speed of an object in a fluid affect the force of resistance it experiences from the viscosity?

As an object moves at increasing speeds through any fluid, the number of fluid particles moving over the surface is increased. This increases the total frictional force which is then measured as resistance.

VISCOSITY SAMPLE REPORT SHEET I. PART A: PRODUCING A VISCOSITY SCALE Data Table Mass of one test tube: ____________ g Volume in each tube: 30.0 mL Test Tube

water/starch ratio

Mass

Time for Marble to Reach Bottom

1

0:30

______ g

__>5 hours___

seconds

2

3:27

______ g

_2 hour 32 min

seconds

3

6:24

______ g

__28 min____

seconds

4

9:21

______ g

__3 min____

seconds

5

12:18

______ g

__47 sec___

seconds

6

15:15

______ g

__6.5 sec_ _

seconds

7

18:12

______ g

_2.1 sec_

seconds

8

21:9

______ g

___1.2 sec___

seconds

9

24:6

______ g

___0.8 sec___

seconds

10

27:3

______ g

___0.7 sec___

seconds

11

30:0

______ g

___0.8 sec___

seconds

Calculations Table Density Calculation Test Tube Total mass – test tube mass = mass of solution

Density = mass/volume

1

_______ g -

_______ g

=

________ g

__________ g/mL

2

_______ g -

_______ g

=

________ g

__________ g/mL

3

_______ g -

_______ g

=

________ g

__________ g/mL

4

_______ g -

_______ g

=

________ g

__________ g/mL

5

_______ g -

_______ g

=

________ g

__________ g/mL

6

_______ g -

_______ g

=

________ g

__________ g/mL

7

_______ g -

_______ g

=

________ g

__________ g/mL

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8

_______ g -

_______ g

=

________ g

__________ g/mL

9

_______ g -

_______ g

=

________ g

__________ g/mL

10

_______ g -

_______ g

=

________ g

__________ g/mL

11 _______ g - _______ g = ________ g __________ g/mL Viscosity Scale Divide each time measurement by the smallest one to get a ratio. Give the ratio to three significant digits.

Test Tube Viscosity Scale 1

___________

2

___________

3

___________

4

___________

5

___________

6

___________

7

___________

8

___________

9

___________

10

___________

11

___________

II. Part B: DETERMINATION OF RELATIVE VISCOSITIES OF COMMON SUBSTANCES Data Table Substance Time to Bottom of Tube Mass Measurement Relative Viscosity from Scale Vegetable Oil

_14 s__ seconds

________ g

___________

Rubbing Alcohol _0.6 s _ seconds

________ g

___________

Honey

3 h 12 m seconds

________ g

___________

29 s_ seconds

________ g

___________

Ketchup

> 5 hrs seconds

________ g

___________

Mustard

> 5 hrs seconds

________ g

___________

Molasses

> 5 hrs seconds

________ g

___________

Hair conditioner

47 s

________ g

___________

Dish Liquid

seconds

47

Calculations Tables Substance Mass –

test tube mass = Mass of substance Density = mass/volume

Vegetable Oil

_______ g -

_______ g = ________ g

__________ g/mL

Rubbing Alcohol

______ g -

_______ g = ________ g

__________ g/mL

Honey

_______ g -

_______ g = ________ g

__________ g/mL

Dish Liquid

_______ g -

_______ g = ________ g

__________ g/mL

Ketchup

_______ g -

_______ g = ________ g

__________ g/mL

III. QUESTIONS 1. Do the densities of the starch solutions vary directly with the viscosity ratings?

Yes 2. Do the densities of the other substances vary directly with the viscosity ratings?

Not all. The ketchup has the highest rating , but not the highest density.

3. If you answered differently for the first two questions, give an explanation for the apparent discrepancy.

The molecular structure in ketchup proteins allows them to form a matrix of protein chains that have formed strong attractions to one another.

4. What factors, other than the viscosity of the fluid, effect how fast the marble falls down the test tube? Does this in any way effect the results as they have been gathered? Why or why not?

Density differences. No, because the marble is more dense than all the fluids measured.

5. How would the data be affected by using a pointed object to fall through the fluids rather than a spherical one?

The shape might allow it to move through the fluids more easily because of the streamlined shape.

6. Relate your answer in number five to the general shape of fish and submarines. How is that applicable to automobiles that also travel through a gaseous fluid, air?

The shape of fish and submarines allow them to move through the fluids with less drag caused by the viscosity. Automobiles can also have less drag if they are shaped aerodynamically.

7. Honey and molasses are both sugar solutions. Knowing this, and that molasses is more dense than honey, can you make a prediction as to the viscosity of molasses? Why is this generalization OK to use when comparing these two substances?

Since molasses is more dense it must have a higher percentage of sugar and will then be more viscous. Since both solutions are made of mostly sugar and water, a direct comparison can be made because there are no other major factors affecting the viscosity.

48

#16

Gas Laws

Reagents sodium hydrogen carbonate 1M hydrochloric acid

.20 g/pair 3 mL/pair

Common Materials Special Equipment buret very small (75 mm long) test tube clear tubing

1/pair 1/pair 4 ft./pair

Laboratory Equipment #1 1-hole rubber stopper with short glass tube 10-mL graduated cylinder 100-mL graduated cylinder & 1-hole stopper 250-mL Erlenmeyer flask & 1-hole stopper 1000-mL beaker thermometer barometer Notes In the first procedure it is very important that there are no bubbles in the tubing. This is why clear tubing has been suggested.

In procedure 14, be sure the stopper does not block the pour spout so that the water cannot escape the cylinder during the gas collection.

Disposal The solutions produced in this investigation can be poured down the drain.

GAS LAWS PRELAB QUESTIONS 1) In step 1, air is trapped inside the small cylinder above the water level while the air in the top of the buret is open to the atmosphere. As the buret is raised and lowered the water exerts varying pressures on the trapped gas in the small cylinder. How is the gas pressure at the top of the water in the buret affected? Explain.

When the water levels are even, the pressure of the trapped air is the same as the pressure of the open air at the surface of the buret. When the water level in the closed cylinder is lower than the water in the buret, the water exerts pressure on the trapped air as it attempts to reach the same elevation as the water in the buret. Therefore the air in the cylinder has greater pressure. However, the air pressure on the water in the buret is a direct result of current atmospheric pressure and so it does not change.

2) Gas volumes change with temperature when all other factors are held constant. What would you expect to see happen to a fully inflated mylar balloon if it was taken outside into freezing weather? Why? The cooler temperatures would allow the gas molecules to move with less energy and thereby exert less pressure on the inside of the balloon. The balloon would have a smaller volume of air than the volume of the balloon so it would appear to be wrinkled.

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3) Basketballs inflated at one temperature and then used at another may not behave as expected. Why do balls often seem “flat” in the wintertime?

Most balls are inflated indoors. When they are taken outside to play, the reduced temperatures cause the internal pressure to be lowered. This results in a smaller coefficient of elasticity for the trapped air and the ball does not bounce as well.

GAS LAWS SAMPLE REPORT SHEET I. Boyle's law a. even levels b. first lifted buret c. second lifted buret d. first lowered buret e. second lowered buret

Air volume _5.0 mL__ _4.6 mL__ _________ _5.4 mL__ _________

Pressure 721.4 torr 757.4 torr _________ 643.4 torr_ _________

II. Charles' law

Temperature

Pressure

Room temp Lower temp Higher temp Highest temp

_296 K__ _279 K__ _335 K__ ________

_674 torr _ _417 torr _ _1150 torr _________

III. Gay-Lussac's law

Volume

Temperature

Room temp Lower temp Higher temp Highest temp

65.0 mL_ 79.8 mL_ 91.0 mL_ ________

_320 K___ _280 K __ _348 K __ _________

PV = _3600 ___ _3500____ _________ _3500____ _________ P = T _2.28_____ _1.50 ____ _3.40_____ __________ V = T __.20_____ __.29_____ __.26_____ __________

IV. Ideal gas law Volume Temperature Pressure final volume __68.4___mL ___21.0_ 732.6_mm Hg - initial volume____0____mL + 273_ - vap. -_18.7_mm Hg ______________________ ________ ____________ gas volume __68.4___mL __ 294 _K 713.8_mm Hg pressure

713.8 mm Hg ×

1 atm = 0.939 atm 760 mmHg

10 -3 L = __0.0684____ L mL 1 mol moles = __0.25____g NaHCO3 × = _0.0684__ moles 84 g PV L  atm R= = ___0.0733___ nT K  mol gas volume __68.4____mL ×

50

V. QUESTIONS 1. Are the values for PV all the same?

They should be very close.

V all the same?

T They should be very close.

2. Are the values for

P all the same?

T They should be very close.

3. Are the values for

4. What was the percent error of your value for R? Ractual - Rcalculated × 100% = % error (10.4 %) Ractual 5. If you plotted your temperature volume data, at what temperature would the volume be zero?

From this data it would be at -42 K. It should occur at 0K. 6. Why should you check the tire pressure after the car has been driven long enough for the tires to gain heat?

The air in the tires heats up with the friction of driving and causes more pressure in the tire.

If the pressure were checked when the tire was cool, it would be too high after the tire had been driven a while. However, tires are designed to handle the extra pressure.

51

#17

Acid Neutralization by Antacid

Reagents 6M hydrochloric acid [HCl] sodium hydroxide pellets [NaOH] phenolphthalein

30 mL/pair 20 g/pair 1-2 mL/pair

Common Materials antacid tablets, 3 brands (Alka Seltzer™ suggested)

2 tablets of each/pair

Special Equipment 1 50 mL pipette with safety bulb 1 500 mL volumetric flask with solid stopper to fit mortar and pestle

1/pair 1/pair 1/pair

Laboratory Equipment 100-mL clean squeeze bottle 500-mL Erlenmeyer flask with solid stopper to fit 150-mL Erlenmeyer flask centigram balance Notes Tablets need to be crushed in order to dissolve quickly. White tablets are better. Colored tablets may obscure the end point.

Disposal Flush all solutions down the drain with a lot of water.

ACID NEUTRALIZATION BY ANTACID PRE-LAB QUESTIONS 1) Some antacids neutralize acid very quickly while others neutralize slowly over a longer period of time. The antacids that produce carbon dioxide gas neutralize acid the quickest while antacids containing hydroxides tend to operate over a longer period of time. Name a common antacid that offers immediate relief.

Alka-Seltzer would provide fast relief because of all the carbon dioxide produced as it effervesces.

2) If an antacid tablet requires warming to dissolve, or if it dissolves slowly, would you expect it to work quickly or be designed to operate over an extended period of time?

It takes time for ingested substances to reach body temperature, therefore one would expect the antacid described to act slowly, over an extended period of time.

3) What color would you expect phenolphthalein to be if it was added to the gastric juices of your stomach?

Phenolphthalein would appear clear, or colorless, in gastric juice because digestive secretions are acidic.

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4) Why would it not be a good idea to use antacids as a calcium supplement?

Antacids are intended to alter the pH of the stomach environment. There are other calcium supplements available that will not increase the risk of alkalosis.

5) What happens to the pH of your stomach if you take more antacid than necessary to neutralize the acid?

When more antacid is taken than necessary, the pH of the stomach can be over neutralized making it slightly basic.

ACID NEUTRALIZATION BY ANTACID SAMPLE REPORT SHEET I. DENSITY OF NaOH Mass of cylinder and solution minus Mass of cylinder Mass of 10 mL solution

________________g _ _______________g _________________ __10.0-11.0______g

mass of solution 10.0 - 11.0 =__1.0-1.1____g/mL = volume 10 mL II. STANDARDIZATION OF NaOH Mass of bottle Initial _________________g _ minus Final ________________g _________________ Mass of NaOH added ___18.0-20.0______g Density =

mass added 18.0 - 20.0 g = = _18.0__mL density 1.0 - 1.1 g/mL (0.36 M HCl)(50 mL HCl) (0.36 M HCl)(50 mL HCl) Molarity of NaOH = = = volume NaOH 18.0 mL ____1.0__________Molar *All other answers will depend on the brands used.

III. QUESTIONS 1. Read the antacid label to find the basic compounds and write balanced equations showing the reaction with the hydrogen ion.

Brand #1 The reactions will vary according to brand, but most will contain CO3-2 or Mg(OH)2: Brand #2 CO3-2 + 2H+ → H2CO3 → H2O + CO2 Brand #3 Mg(OH)2 + 2H+ → Mg2+ + 2H2O 2. Which antacid appears to be best at relieving acid stomach?

This varies according to the brands used.

3. Which antacid has more neutralizing power per gram?

This varies according to the brands used.

4. In recent commercials some antacids brag that they contain calcium. If calcium is used by the body to form bones and teeth, why does it matter that it is in an antacid?

When basic compounds such as calcium carbonate react with the hydrogen ion, which is acidic, carbonic acid is produced which quickly dissociates into water and carbon dioxide. This leaves the calcium ion free to be used by the body to build bones and teeth.

Volume of NaOH added

=

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#18

Nature's Indicators

Reagents 0.1M hydrochloric acid [HCl] 0.1M sodium hydroxide [NaOH] pH paper red litmus blue litmus

20 mL/pair 20 mL/pair 1 roll/pair 1 container/pair 1 container/pair

Common Materials red cabbage beets conc. grape juice lemon juice ammonia household cleanser vinegar clear soft drink white or clear shampoo (Clear is better.) liquid drain cleaner knives

1 head/section 1/pair 20 mL/pair 1 mL/pair 1 mL/pair 1 mL/pair 1 mL/pair 1 mL/pair 1 mL/pair 1/pair

Laboratory Equipment 2 250-mL beakers ring stand, ring, and wire gauze 100-mL graduated cylinder laboratory burner 7 100-mL beakers stirring rod 7 test tubes Notes Canned beets and red cabbage will also work.

The skins of black grapes will also work well for indicator preparation, however, this gives essentially the same results as the concentrated grape juice. The grape juice should be 100% and not a grape drink. Other plants which will produce indicators are petals from violets, petunias, hydrangea, and hibiscus; the skins of tomatoes, turnips, red apples, peaches, pears, and red onions; and the juices of pomegranates, cherries, raspberries, blueberries, and blackberries. (Source: WonderScience magazine published by the American Chemical Society.)

The vegetables may also be heated (not boiled) in ethanol to produce an indicator. By this method, less time is required to evaporate the solution down to the desired concentration of color.

There are both acidic and basic drain cleaners. A basic drain cleaner such as Drano™, will provide a second basic solution.

Disposal Dispose of your pH solutions, one at a time by pouring them down the drain with a large amount of water.

54

NATURE'S INDICATORS PRE-LAB QUESTIONS 1) Hydrangea plants (Hydrangea macrophylla) produce very large flower clusters and are common garden shrubs in central and southern states. Many people believe that there are two varieties of this shrub – blue and pink – when, in fact, there is only one. Many homeowners have planted a hydrangea of one ”variety” only to have it slowly transform to the other over a period of years. Give an explanation for this phenomenon in light of the introductory information for this investigation.

The pH of the soil determines the color of the pigment in the petals, because at least one component of that pigment is a natural indicator.

2) Litmus is either red or blue. It is useful only to determine if a substance is classified as an acid or a base. Indicators that are capable of specifying a particular pH exhibit a range of color changes. Would you expect that these “universal indicators” are a single compound, two compounds combined, or several compounds combined? Support your position.

Since most indicators exhibit only a single color change at a narrow pH range, a universal indicator must be a combination of several substances, each covering a separate range of transition.

3) Would an experiment using indicators to determine the pH of common substances work in an already colored substance, such as a cola beverage? Why or why not?

It would not be useful to use a color indicator in a colored substance, especially one that is darkly colored because the existing color would mask the color of the indicator.

4) Nearly all indicators are organic molecules that are sensitive to destruction by strong solutions, especially bleaching cleansers. Could the pH of chlorine bleach be determined with an indicator derived from a plant? Explain.

Probably not. Plant substances are organic and chlorine bleach destroys organic molecules.

If it gave an indication of pH, it would be short-lived.

5) A child goes to the kitchen sink to rinse the glass from which he just drank grape juice. As the water runs into the glass the juice residue turns from purple to light blue. What is happening?

The grape juice contains a natural indicator that appears dark purple at a specific pH range.

As the juice is diluted, the pH moves closer to neutral and the transition range of the indicating substance is crossed so that it changes color.

NATURE'S INDICATORS SAMPLE REPORT SHEET INDICATOR COLOR Cabbage pH = 1 dark pink pH = 3 light pink pH = 5 light purple pH = 7 very lt. purple pH = 9 light blue pH = 11 green pH = 13 lime green

Beet dark pink light pink very lt. pink very lt. pink very lt. pink pink yellow

Grape pink light pink very lt. pink very lt. pink clear clear green

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Blue Litmus red red red red red blue blue

Red Litmus red red blue blue blue dark blue dark blue

INDICATOR USED * __Cabbage_______________ SUBSTANCE pH by indicator Lemon juice __1___ Vinegar __3___ Soft drink __5___ Shampoo __7___ Ammonia cleanser _11___ Drain cleaner _11___ *Beets give the following: 1,3,7,7,11,9; Grapes give the following: 1,3,5,7,11,9

pH by paper __2___ __3___ __7___ __7___ _10___ __8___

QUESTIONS 1. Are beverages usually acidic or basic?

Beverages are usually acidic or neutral: water is neutral while orange juice is acidic. This is true because many beverages contain fruit juices or have acid added for flavor.

2. Why should shampoo not be too acidic or too basic?

Shampoo must be neutral because the acid or base could damage the hair.

3. Which indicator did you like best? Why? Student response here.

4. Which indicator covers the largest range of pH? The cabbage, beet, and grape juice cover the full pH scale, but the cabbage changes color at each pH. 5. Which indicator produced the most different colors? cabbage 6. Why are there two types of drain cleaners?

The bathroom drain cleaner has acid which breaks down protein in hair that is usually found in bathroom drains. The second type is basic and is used to react with the grease in kitchen clogs and helps to free the clog.

7. Comment on how closely your pH determinations using the experimental indicator scale matched the pH determinations with the commercial pH paper.

Answers will vary, but the results should be very close.

56

#19

pH and Flammability

Reagents ethanol [CH3CH2OH] acetone [CH3COCH3] sodium tripolyphosphate [Na5P3O10] sodium pyrophosphate [Na4P2O7] sodium carbonate [Na2CO3] sodium silicate [Na2SiO3] pH paper

2 mL/pair 2 mL/pair 2 mL/pair 2 mL/pair 2 mL/pair 2 mL/pair 8-10 strips/pair

Common Materials vanilla extract nail polish remover cologne aftershave lotion paint thinner oven cleaner hair spray toilet bowl cleaner hand dish washing detergent machine dish washing detergent chlorine bleach vinegar drain cleaner ammonia cleanser other products from home matches

2 mL of each/pair

1 book/pair

Laboratory Equipment evaporating dish 50-mL beaker stirring rod laboratory burner Notes The unused paint thinner should not be flushed down the drain but carefully sealed for use another semester. The small amount of paint thinner given to the students should be consumed by the flames. Be sure it is a flammable paint thinner.

There are both acidic and basic drain cleaners. Drano™ is basic; Works™ is acidic.

Disposal Paint thinner will contaminate water. It should be completely consumed by the flames. It could also be evaporated in a fume hood.

57

pH and FLAMMABILITY PRE-LAB QUESTIONS 1) In many homes, multiple flammable compounds normally reside under the kitchen sink.

Another common storage place for household cleaning supplies and solvents is the “utility” closet that may also house a gas fueled water heater or furnace. Comment on the risks and wisdom of this practice. Can you suggest alternatives?

Any place in which an open flame or electrical sparks exist is a hazardous place to store flammable substances; especially unventilated areas where vapors can build up. A good alternative is a ventilated storeroom where there are no gas fueled appliances or appliances on thermostats that generate small sparks when they cycle on.

2) Assume you are the corporate attorney for a company that produces, packages, and markets multiple household cleaners and solvents. What suggestions would you make to the company CEO concerning consumer warnings?

Correct responses may include but may not be limited to: • Flamables should be stored in ventilated areas and away from open flames and electrical sparks.

• Never mix cleaning products.

• Caustic warnings on products of low or high pH. • Keep all household products away from children.

• Toll Free Poison Control number.

3) It has been suggested that a labeling system for easy identification of pH and flammability ratings for common substances should be instigated. If this system used color-coding for ease of identifying the characteristics of a substance, what suggestions would you make for the design?

Answers will vary.

4) Which of the substances we will use in this lab can you find in your own home, apartment or dorm room?

Answers will vary.

5) How can you reduce the risk of accidents and injuries involving household chemicals in your domicile?

Read and follow all warnings and directions for household substances. Adhere to proper storage techniques. Use all household products in open or ventilated areas. Wear protective clothing and gloves when using caustic substances.

58

pH & FLAMMABILITY SAMPLE REPORT SHEET FLAMMABILITY LIGHTS WITH LIGHTS WITH MATCH STRONGER FLAME Ethyl alcohol __X____ _______ Acetone __X____ _______ Vanilla extract _______ __X____ Nail polish remover __X____ _______ Cologne __X ___ _______ After shave lotion _______ __X____ Paint thinner _______ __X____ Oven cleaner _______ _______ Hair spray __X___ _______ pH Brand pH Toilet bowl cleaner _______Saniflush__________ ____2___ Hand dish washing detergent _______Ajax______________ ____6___ Machine dish washing detergent _______Sunlight___________ ___10___ Bleach _______Vista_____________ ___10___ Na5P3O10 ________________________ __ 8 _ Na2CO3 ________________________ ___10___ Na2SiO3 ________________________ ___10___ Ammonia cleaner ________________________ ___10___ Vinegar ________________________ ___4____ Oven cleaner ________________________ ___14___ QUESTIONS 1. Which products have adequate warning about flammability?

Answers will vary depending on the brands used.

2. Which products need more warning about flammability?

Answers will vary depending on the brands used.

3. Which of the products tested unsafe for skin contact? (Products of pH between 5 and 10 are probably safe for at least short periods of time.)

Toilet bowl cleaner Vinegar 4. Which of the products tested have adequate warnings on the label?

Answers will vary depending on the brands used.

5. Why would you not use machine dish-washing detergent to wash dishes by hand?

The pH of the machine dish washing detergent is basic and is not safe for prolonged use.

6. Why are gloves suggested when cleaning the oven?

Oven cleaner is not safe for skin due to the pH. 7. Because of environmental concerns, phosphates have been replaced with other builders such as sodium carbonate and sodium silicate. Does this change make a difference in safety to the consumer?

Sodium carbonate and sodium silicate are more basic than the phosphates which means that sodium carbonate and sodium silicate can be more dangerous to the consumer than phosphates.

59

#20

Buffers

Reagents 1M acetic acid 1M sodium acetate 1M acetic acid & 1M sodium acetate (buffer solution) wide range pH paper 1M hydrochloric acid 1M sodium hydroxide [NaOH]

10 mL/pair 10 mL/pair 20 mL/pair 1 roll/pair 50 mL/pair 50 mL/pair

Common Materials Special Equipment extra 50 mL beakers 10-mL graduated cylinder

2/pair

Notes Be sure students stir the solutions well after each addition before taking the pH measurement.

Disposal All solutions may be mixed to neutralize and poured down the drain with lots of water.

BUFFERS PRE-LAB QUESTIONS 1) When a person exerts a great deal of physical energy, as in running a race or playing oneon-one basketball, one of the by products of muscular activity is lactic acid. If the blood buffering system cannot handle the lowered pH, what condition could result?

Acidosis could result as the lactic acid lowers the pH of the blood. The lowered pH could in turn effect proteins in nearby cells. This is the cause of the stiffness and soreness associated with the use of muscles that are not accustomed to a higher level of activity.

2) Many consumer products are buffered to protect the body from dangerous pH changes.

Pharmaceutical compounds are usually either basic or acidic and not as often neutral compounds. What systems in the body need the protection of a buffer if the substance is ingested? Injected?

Ingested substances that are acidic or basic must be shielded from the stomach lining if the pH is very far from neutral, or if the exposure will be extended. Inhaled substances can cause irritation to the lining of the nose, mouth, throat, esophagus, and lungs.

3) Fertilizers are formulated to help maintain specific pH levels for certain applications.

Will making the fertilizer the desired pH be enough to accomplish the result? Explain your reasoning.

The nitrates in fertilizer enter the nitrogen cycle in which several nitrogen compounds take part. These various compounds can alter the pH of the soil as the chemistry develops.

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4) Baking soda (sodium hydrogen carbonate, NaHCO3) acts as a buffer, as does sodium tetraborate, also called borax and used as a water softener in laundry applications. Both raise pH and then tend to buffer it to keep it consistent. Both of these compounds are active ingredients in pool maintenance products. The hypochlorite ion (ClO-) is the ingredient used to “shock” a pool. It is also the active ingredient in chlorine bleach.

Why do people spend large sums of money for pool products that are labeled with cute names like “Alk-UpTM” or “SoftSwimTM” instead of buying the substances off the grocery store shelves?

Most consumers do not have the knowledge necessary to maintain the pool water quality on their own. Testing solutions for pH and other indicators are available with instructions that allow a person with virtually no chemical experience to maintain their pool environment. It is also easier to purchase all the necessary chemicals at one place catering to pool needs that to get some things at a grocery store while others are available only at specific businesses.

BUFFERS SAMPLE REPORT SHEET I. Water

Acetic Acid buffer pH 1 drop HCl _____ 2 drops HCl _____ 20 drops HCl _____ 2 mL HCl _____ 3 mL HCl _____ 4 mL HCl _____ 5 mL HCl _____ 6 mL HCl _____ 7 mL HCl _____ 8 mL HCl _____ 9 mL HCl _____ 10 mL HCl _____ 11 mL HCl _____

II. 1 drop NaOH _____ 2 drops NaOH _____ 20 drops NaOH_____ 2 mL NaOH _____ 3 mL NaOH _____ 4 mL NaOH _____ 5 mL NaOH _____ 6 mL NaOH _____ 7 mL NaOH _____ 8 mL NaOH _____ 9 mL NaOH _____ 10 mL NaOH _____ 11 mL NaOH _____

Acetate ion

Acetic acid/acetate

pH ______ ______ ______ ______ ______ ______ ______ ______ ______ ______ ______ ______ ______

pH _____ _____ _____ _____ _____ _____ _____ _____ _____ _____ _____ _____ _____

pH ______ ______ ______ ______ ______ ______ ______ ______ ______ ______ ______ ______ ______

______ ______ ______ ______ ______ ______ ______ ______ ______ ______ ______ ______ ______

_____ _____ _____ _____ _____ _____ _____ _____ _____ _____ _____ _____ _____

______ ______ ______ ______ ______ ______ ______ ______ ______ ______ ______ ______ ______

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III. Buffer pH

______

add 5 mL water _____

add 10 mL water _____

IV.QUESTIONS 1. Which is better at resisting pH change when acid is added: water, acetic acid, acetate ion, or buffer?

The buffer is the best.

2. Which is the worst at resisting pH change when acid is added: water, acetic acid, acetate ion, or buffer?

The water is the worst.

3. Which is better at resisting pH change when base is added: water, acetic acid, acetate ion, or buffer?

The buffer is the best.

4. Which is the worst at resisting pH change when base is added: water, acetic acid, acetate ion, or buffer?

The water is the worst.

5. Why does the buffer solution have a large pH change after 10 mL is added?

The amount of acid added has overcome the conjugate base's ability to absorb. (i.e., the base has been used up.)

6. Why would it be useful to "buffer" an aspirin tablet? (Aspirin is acetyl-salicylic acid.)

The acid can be uncomfortable in the digestive tract that is already acidic. With the buffer, the tablet will resist lowering the pH of the stomach even more.

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#21

Oxidation - Reduction

Reagents zinc, mossy tincture of iodine 1M copper sulfate [CuSO4] 1M sodium chloride [NaCl] 1M acetic acid [HC2H3O2]

1 g/pair 20 mL/pair 30 mL/pair 10 mL/pair 0.5 mL/pair

Common Materials aluminum foil coins, different metals paper towels chlorine bleach

2 (4 in. x 4 in.) squares/pair 2/pair 1 sheet/pair 0.5 mL/pair

Special Equipment voltmeter

1/lab

Laboratory Equipment 3 50-mL beakers Notes Disposal Wet zinc dust exposed to air can burst into flames. It is important that students not put wet zinc dust in the trash; it can cause fires in the trash. The wet zinc solutions should be spread on a metal pan to dry. It must be in a metal pan for the chemical reaction to form zinc oxide. Dry zinc oxide will form which can be buried in a landfill.

OXIDATION AND REDUCTION PRE-LAB QUESTIONS 1) Have you ever experienced a little jolt while the dental assistant was cleaning your teeth?

It happens when the metal instrument they are holding touches a metallic filling, or more often, the metal base to a crown. It is called galvanic shock. What is happening?

A difference in voltage potential exists between the two metals that are in contact, the instrument and the dental work, and this potential is in a liquid environment with plenty of ions in solution. That is the perfect recipe for a small current.

2) Some people will test a small 9-volt battery to see if it is still usable by touching the terminal end of it to their tongue. (Not a wise practice, by-the-way.) Why doesn’t it work just as well to lay a finger across the two terminals?

The surface of a finger is often very dry and therefore no current can be supported.

3) Why must care be taken when serving food from sterling silver if the food is acidic?

Acidic foods provide the necessary ions in solution to support a current. Depending upon other substances present, especially sulfur, an environment allowing the silver to be oxidized and bond with another element is likely.

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4) Consider a galvanized iron pipe full of dirty water that may contain ions of other metals.

The metal ions might have a reduction potential compared to that of the iron so that a voltaic cell is produced that would corrode the pipe. The presence of zinc may prevent the corrosion. On a table of reduction potentials, would you expect zinc to fall between iron and the other metal, or would zinc be located either above or below both iron and the other metal? Explain your reasoning.

In order for the zinc to react before the iron, it would have to lie between the iron and the other metal on the activity series. If it were to either side of the pair, it would not be able to “interfere” with the reaction.

OXIDATION AND REDUCTION SAMPLE REPORT SHEET I. VOLTAIC CELL VOLTAGE

__________V

(Voltage depends on coins used.)

INSTRUCTOR'S INITIALS __________ 1. Would other coins work?

Yes, any two different metals.

2. Would pieces of metal work?

Yes, as long as the metals are different.

3. Would other ionic compounds like KBr work instead of NaCl? As long as the ionic compound was a salt in the molten state or in solution it would allow the current to pass. Crystalline salts would not.

II. ZINC, IODINE, CHLORINE 1. Is zinc in the first solution? In what form?

Yes. It was a solid which became ions when it dissociated in solution.

2. In what form is the zinc when the purple color is gone from the solution?

Zinc is in a solid form.

3. Does bleach change the zinc back to the metal?

No, it changes into Zn(OH)2, which puts Zn2+ions into solution.

4. In what form is the chlorine when the purple color is gone from the solution?

The chlorine is in the gaseous Cl2 form.

5. What is the color of the solution after adding the bleach?

Purple III. COPPER AND ALUMINUM 1. What does the difference in color in the copper solution indicate?

The copper in solution has been reduced to copper metal which then attached to the aluminum. The solution becomes a lighter blue as the copper ions leave.

2. What happened to the aluminum foil?

The copper attaches to it and replaces the aluminum ions which leave and go into solution. Holes may appear in the foil if it is left too long.

3. In what form is the copper at the end of the investigation?

The copper is solid.

4. Could copper be recycled by collecting it on aluminum foil?

Yes.

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#22

Single Replacement Reactions

Reagents 1 M calcium nitrate [Ca(NO3)2] 1 M copper (II) nitrate [Cu(NO3)2] 1 M iron (II) nitrate [Fe(NO3)2] 1 M magnesium nitrate [Mg(NO3)2] 1 M tin nitrate [Sn(NO3)2] 1 M zinc nitrate [Zn(NO3)2] 1 M nitric acid [HNO3] calcium metal pieces, small copper metal wire, small iron metal pieces, small magnesium metal ribbon, small pieces tin metal pieces, small zinc metal pieces, small

1 mL/pair 1 mL/pair 1 mL/pair 1 mL/pair 1 mL/pair 1 mL/pair 1 mL/pair 1/pair 1/pair 1/pair 1/pair 1/pair 1/pair

Common Materials Special Equipment test tubes, small

7/pair

Notes This can be done with less solution in microtitration plates if they are available.

This investigation works well in conjunction with the Double Replacement Reactions investigation. Begin the double replacement lab after the “5 minute” check.

Disposal The solutions can be flushed down the drain with lots of water.

SINGLE REPLACEMENT REACTIONS PRE-LAB QUESTIONS 1) Platinum is a precious metal used in the finest jewelry. Where on the activity series would you expect to find it? Explain your reasoning.

Coins and jewelry are made from metals that undergo very little corrosion so that they retain their value. Therefore we would expect to find platinum at the lowest portion of the table.

2) Sodium metal is near the top of the activity series and therefore reacts easily with moisture in the air, oxygen, and weak acids. Devise a storage method for elemental sodium.

Any method that would be effective in isolating the metal from air and moisture should be accepted as an appropriate answer. Vacuum packaging and storing under a hydrocarbon such as kerosene are common methods.

3) Use the activity series in your textbook to predict which of the following combinations will result in a reaction. Write “yes” or “no” in the blank before each combination. ___no a) metallic iron in a solution of lithium nitrate __yes b) metallic zinc in a solution of copper sulfate __yes_ c) metallic strontium in a solution of nickel chloride

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4) Which of the two reactions below would you expect to happen the fastest? ______ a) between chromium and cadmium nitrate ___X _ b) between barium and lead nitrate

SINGLE REPLACEMENT REACTIONS SAMPLE REPORT SHEET I. REACTION TIMES AND CHANGES OBSERVED 1st METAL Assigned: _Fe_____ Time

HNO3

Zn(NO3)2

Fe(NO3)3

Ca(NO3)2

Mg(NO3)2

Sn(NO3)2 Cu(NO3)2

within five bubbles minutes thirty minutes 2+ hours

less blue in solution

never

X

X

X

X

X

Mg(NO3)2

Sn(NO3)2 Cu(NO3)2

2nd METAL Assigned: Zn______ Time

HNO3

Zn(NO3)2

Fe(NO3)3

Ca(NO3)2

within five bubbles minutes

zinc turns dark

thirty minutes 2+ hours never

yellow tint X

X

X

66

zinc turns dark

III. ACTIVITY SERIES DETERMINATION Zn/Zn2+

Fe/Zn2+

Time

Ca/Zn2+ 5 min

Mg/Zn2+ 5 min

Points

3

3

Sn/Zn2+

Cu/Zn2+

Zn/Fe3+ 2 hrs 1

Fe/Fe3+

Ca/Fe3+ 5 min 3

Mg/Fe3+ 30 min 2

Sn/Fe3+

Cu/Fe3+

Zn/Ca2+

Fe/Ca2+

Ca/Ca2+ 5 min 3

Mg/Ca2+ 5 min 3

Sn/Ca2+

Cu/Ca2+

Zn/Mg2+

Fe/Mg2+

Ca/Mg2+ 5 min 3

Mg/Mg2+

Sn/Mg2+

Cu/Mg2+

Fe/Sn2+ 2 hrs 1

Ca/Sn2+ 5 min 3

Mg/Sn2+ 5 min 3

Sn/Sn2+

Cu/Sn2+

Time Points

Zn/Sn2+ 5 min 3

Fe/Cu2+ 5 min 3

Ca/Cu2+ 5 min 3

Mg/Cu2+ 5 min 3

Sn/Cu2+

Cu/Cu2+

Time Points

Zn/Cu2+ 5 min 3

Fe/H+ 30 min 2

Ca/H+ 5 min 3

Mg/H+ 30 min 2

Sn/H+ 5 min 3

Cu/H+

Time Points

Zn/H+ 30 min 2 Zn Total

Fe Total

Ca Total

Mg Total

Sn Total

Cu Total

9

6

21

16

3

0

Time Points Time Points Time Points

Total Points

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IV. QUESTIONS 1. How useful would you say the Activity Series is for aiding chemists in knowing the probable outcome of a reaction involving two metals?

It would be very useful to know which metal would react with which ion. .

2. Write a balanced equation below for each of the reactions for the first metal you were assigned. (Give the reactants for each attempted reaction. For those in which nothing was visible, write “no reaction” following the yield sign.)

Fe(s) + HNO3 → Fe2+ + H2(g) Fe(s) + Cu2+ → Fe2+ + Cu(s) 3. List the metals we tested in order of activity with the most active metal first and the least active metal last. (i.e., The metal with the most points would be the most active and be placed at the beginning of the list.) calcium, magnesium, zinc, iron, tin, and copper 4. How well does this agree with the accepted activity series: calcium, magnesium, zinc, iron, tin, and copper?

It does agree.

5. What factors could explain any discrepancies you find?

Student may not notice reaction, metal coated with oil, dirt, etc. that restricted reaction.

6. According to the data for our class, where in the series would you place hydrogen?

Nitric acid reacted with all but copper. It belongs between tin and copper.

7. Magnesium and calcium are not expected to react with themselves, and yet there was a definite reaction. What was happening?

There was a slight reaction with water.

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#23

Silver Tarnish Removal

Reagents sodium bicarbonate [NaHCO3] thioacetamide [CH3CSNH2] sodium chloride [NaCl]

13 g/pair 6.5 g/class 13 g/pair

Common Materials tarnished silver aluminum foil bleach (sodium hypochlorite solution)

1 piece/pair 2 ft length/pair (for disposal only)

Special Equipment 2 liter beaker ring stand, ring, and wire gauze forceps laboratory burner or hotplate Notes Jewelry, silverware, small salt and pepper shakers, etc. generally give better results than silver coins because coins tend to have much more than simple tarnish on the surface. Removing the "goo" which accumulates in every day use is difficult to do without also removing the tarnish by abrasion.

(We have had good luck suggesting that our students bring something from home to clean.) Many times, flatware and jewelry can be found inexpensively at flea markets and used semester after semester.

Disposal The thioacetamide will cause tarnish and should not be put into the water system.

The thioacetamide solution should be added slowly to a solution of sodium hypochlorite (bleach).

When the solution has cooled to room temperature, it can be flushed down the drain with plenty of water. The bleach solution should be in an excess of about 25%.

SILVER TARNISH REMOVAL PRE-LAB QUESTIONS 1) List foods that should not be served from silver service dishes.

Foods that are acidic and foods that contain sulfur will increase the likelihood of tarnishing the silver.

2) Silver stored in some plastics tend to tarnish quickly. What would you suspect is an element in the plastic?

Sulfur 3) How could silver be protected from tarnish in homes with gas fueled furnaces and ranges?

The silver can be placed into sealed storage places or encased in fabrics that are treated to remove sulfur from the air before it reaches the silver. (This fabric is referred to as “silver cloth”.)

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4) Some lotions and cosmetic products will cause silver to tarnish on contact. What is happening?

The products must contain sulfur that will react readily with the silver.

SILVER TARNISH REMOVAL SAMPLE REPORT SHEET I. INSTRUCTOR'S INITIAL FOR CLEAN SILVER______________ II. QUESTIONS 1. Look at the aluminum foil. What happened to the sulfur?

The sulfur bonded to the aluminum instead of the silver. The aluminum becomes darkened and may become thin as it bonds to the sulfur.

2. Would this method be preferable to a paste polish if you wanted to remove the tarnish from the crevices of a very intricate piece of silver? (Removing the tarnish from the crevices of a very old piece devalues the piece in the eyes of antique collectors.)

This method would be preferable to a paste method because pastes use abrasives to remove the tarnish, and thus remove silver along with the sulfur.

3. Why is it necessary for the silver and the aluminum to touch?

In order for there to be an exchange of electrons, the two metals need to touch.

4. Is it necessary to be in a chemical laboratory to clean silver?

No, it is not necessary. This procedure can be done at home.

5. There are silver cleaning kits sold on TV and in stores that use a large plate on which the silver sits to be cleaned. What do you think the plate is made of?

The plate is made of some type of metal, probably aluminum ar zinc.

6. Why should this method not be used for silver with stones in it?

The stones could be loosened, they may react with the solutions, or glue could be dissolved.

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#24

Organic Qual Scheme

Reagents hexane n-octene 2-chloro,2-methyl propane benzaldehyde resorcinol acetic acid (conc.) ethanol acetone isobutylamine pH paper litmus, red and blue 1% iron (III) chloride solution cerric nitrate solution (4g in 10 mL 2M HNO3) acetone (as a test reagent) 1% potassium permanganate copper wire (1 cm lengths) Common Materials food coloring paper towels

unknown unknown unknown unknown unknown unknown unknown unknown unknown 1 roll/pair several strips/pair 2 mL/pair 2 mL/pair 2 mL/pair 2 mL/pair 3/pair 1-2 drops/pair

Laboratory Equipment watch glass small test tubes stirring rod Special Equipment waste container marked "organics and water" Notes Hexane, n-octene, and 2-chloro, 2-methyl propane are chosen because they are large enough molecules as to give a definitively negative solubility test. This makes the determination easier.

Likewise, propionaldehyde, resorcinol, acetic acid, ethanol, acetone, and isobutylamine were chosen because they are small enough to give a very definite positive solubility test. At this level of instruction, it is advisable to keep the tests as easily identifiable as possible. (Note that 2,3 dimethyl phenol does NOT work for an easy phenol identification.) pH paper may lead to errors for neutral substances as students may interpret the value as 6 or 8.

Disposal All solutions should be sealed in a waste container marked "organics and water."

71

ORGANIC QUAL SCHEME PRE-LAB QUESTIONS 1) In finding the identity of an unknown organic compound, determining the class to which it belongs is only the first step. What else would need to be known in order to identify the compound specifically?

The molar mass of the molecule.

2) Only tiny amounts of an unknown are required for each test. Why is this advantageous?

In a real-world situation a sample brought to a lab for testing may be very small.

3) If you knew that a compound of the alcohol class was toxic to living cells and made a good antiseptic, could you suspect this is true of other alcohols, also? Explain.

Yes. It is the functional group in common with all the other substances of a class that imparts the particular physical and chemical properties that characterize the class of organic compounds.

ORGANIC QUAL SCHEME SAMPLE REPORT SHEET OBSERVATIONS Water solubility Two layers form with insoluble substances. pH Answers will vary according to which substance is tested and the type of pH paper used.

FeCl3 A positive result becomes tinted purple while a negative result remains slightly yellow.

Cerric nitrate A positive test changes to red while a negative test remains yellow.

Baeyer A positive test turns brown while a negative test remains purple.

Beilstein A negative test produces only a yellow or orange tint to the flame while a positive test give the flame a greenish tint.

Trace the dotted lines according to your test results to find the classification of the unknown.

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Alkanes, Alkenes, Alkyl Halides, Alcohols, Aldehydes Ketones, Car boxylic A cids, Amines, Phenols Water

Soluble?

Insoluble?

Beilstein

p H pap er Negative?

P ositive?

Neutral?

Acidic?

Alky l halide

Amine

F eC l 3

Baey er Negative?

Negative?

P ositive?

Alkane

Alkene

Basic?

P ositive?

P henol

Carboxylic acid

Cerr ic nitr at e Negative?

P ositive?

Alcohol

Baey er

Negative?

P ositive?

Ketone

Aldehyde

I. Classification of the unknown (student responses) #1 _________________ #2 _________________ #3 _________________ II. QUESTIONS 1. Study the qual scheme and give an example of an incorrect determination that could be made by performing the tests out of order.

If the Baeyer test was performed before the solubility test one could incorrectly identify a ketone as an alkane or an aldehyde as an alkene.

2. List an application in which a qual scheme would be useful.

Responses will vary, but could include identifying contaminates in water, forensics, quality control.

3. Many organic compounds have distinct odors. Why does one not simply sniff the unknown to make the identification?

Many organic compounds are also toxic by inhalation or can cause damage to tissues and the brain.

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#25

Synthesis of Esters

Reagents salicylic acid [C7H6O3] methanol [CH3OH] 6M sulfuric acid [H2SO4] acetic acid, concentrated [CH3COOH] n-octanol [CH3(CH2)7OH] n-pentanol [CH3(CH2)4OH] (n-amyl alcohol) ethanol [CH3CH2OH] butyric acid [CH3(CH2)2COOH] benzyl alcohol [ C6H5CH2OH] propanoic acid [CH3CH2COOH] isobutyl alcohol (2-methyl propanol) [CH3CH(CH3)CH2OH] formic acid [HCOOH]

0.1 g/pair 2 mL/pair 1 mL/pair 2 mL/pair 0.5 mL/pair 1 mL/pair 1 mL/pair 1 mL/pair 2 mL/pair 0.5 mL/pair 1 mL/pair 1 mL/pair

Laboratory Equipment 250 mL beaker 6 small test tubes test tube holder ring stand, ring and wire gauze laboratory burner Notes If the butyric acid stays in the hood and capped, the smell should not become overbearing.

Disposal Esters are organic substances that will contaminate waterways if flushed down the drain.

Each bottle of ester can be left open in a hood to evaporate. Although the idea of putting chemicals into the air is not one that should be given to students, the alternative disposal methods are to incinerate the esters or bury them in a toxic waste site.

SYNTHESIS OF ESTERS PRE-LAB QUESTIONS 1) Since we know that esters are produced from an alcohol and a carboxylic acid, would it be worth the time and expense for chemists to simply try every possible combination in an to attempt to discover new scents and flavors? Support your opinion.

The number of potential organic compounds is considered by many people to be limitless, for all practical purposes. It would probably not be worth the time and expense since all the recognizable scents and flavors are generally found in nature. Humans might not find unnatural flavors or scents pleasurable.

2) Any chemical process involves side reactions and a bit of unreacted reagent. Why should you not taste or use any of the ester you will produce in this lab?

It has not been purified to remove products from side reactions or unreacted substances.

3) Vanilla flavoring is a combination of several compounds isolated in nature from the vanilla bean. However, only two of the compounds make up the largest percentage of the mixture and together those two compounds give the familiar flavor we know. Pure 74

vanilla extract is quite expensive, but imitation vanilla, that contains the two major compounds synthesized in a lab, is not. List advantages and disadvantages for the use of the pure extract.

It is an advantage because it is much less expensive that the vanilla isolated from the vanilla bean, but the flavor is not as “full” because it does not contain the great variety (2-methyl propanol) of compounds found in natural vanilla.

4) Many pharmaceuticals are synthesized copies of compounds found in nature. There is not enough of the natural substance to fill the needs of the world’s population, and even if there were, the cost of isolating it from all the other things that come with it is exorbitant. Why do you think that the widespread and inexpensive availability of things such as aspirin and insulin is a relatively recent phenomenon?

Until scientists had the knowledge and the technology to analyze a compound to determine its composition and then the ability to synthesize that molecule, large amounts of natural compounds were not able to be isolated without great time and expense.

SYNTHESIS OF ESTERS SAMPLE REPORT SHEET I. For each compound write a complete equation in addition to observations of aromas and appearance.

1. methyl salicylate -"oil of wintergreen" - appearance is crystalline, odor is mediciny, minty 2. octyl acetate -"orange peelings" - the liquid is clear, coconut smell O

CH 3 COH Acetic Acid

3.

4.

5.

6.

7.

8.

9.

O +

H HO(CH

+

CH 3 CO( CH

2 ) 7 CH 3

Octy l Alcohol

2 ) 7 CH 3

Octy l Acetat e

+

H 2O Water

n-amyl acetate -"banana oil" - clear liquid, banana extract benzyl acetate - "jasmine" - flowery odor ethyl propionate - "rum" - clear liquid, alcohol smell isobutyl formate - "raspberry" - raspberry odor methyl butyrate -"apple" - apple smell ethyl butyrate -"pineapple" - pineapple extract benzyl butyrate - "cherry" - thick white liquid, artificial cherry smell

II. QUESTIONS: 1. What are the physiological properties (effects on humans) of these esters?

This information can be found in a Merck Index.

Esters are poisonous to humans. They can cause headaches, nausea, and dizziness in large quantities. However, the amount used in food flavorings is very small because of the potency of the molecules.

2. Is it valid to be concerned about "synthetic" flavorings in food?

No, the synthetic molecules are identical to the molecules found in nature.

3. List several substances that contain esters. cosmetics, household products, cleaners, foods, soaps, perfumes, air fresheners 4. How would the use of esters change if they were not volatile?

There would be no smell, therefore they would only be useful in foods as flavorings. (However, even flavor is, to a large extent, due to the sense of smell.)

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#26

Polymers

Reagents hexamethylenediamine [H2N(CH2)6NH2] sodium hydroxide pellets [NaOH] phenolphthalein (or food coloring) sebacoyl chloride [ClCO(CH2)8COCl] hexane [C6H14] phthalic anhydride sodium acetate glycerol (glycerine) ethylene glycol sulfuric acid 18M (for disposal) ParafilmTM

1.5 g/pair 0.5 g/pair 1 mL/pair 1 mL/pair 25 mL/pair 8 g/pair .3 g/pair 2 mL/pair 1 mL/pair 50 mL as needed

Common Materials large metal paper clips aluminum foil (tear 4" strips from roll and halve)

1/pair 1 4"x4" piece/pair

Laboratory Equipment 2 watch glasses ring stand 2 utility clamps tweezers or forceps 150 mL beaker large beaker (400 or 600 mL) Special Equipment disposable test tubes

1/pair

Notes Disposal Hexamethylene diamine, sebacoyl chloride, and hexane are all irritating to humans. Ethylene glycol is toxic. It would not be good to place any of these in the water system.

The solid polymers and foil can be thrown into the trash. When all of the polyester has dissolved, the sulfuric acid in the test tubes can be flushed down the drain with lots of water. Hopefully the nylon reactants will be used up. If not, mix the remaining reactants together to form a ball of nylon that can be thrown into the trash.

POLYMERS PRE-LAB QUESTIONS 1) Study the structural reaction equation for the two polyester polymers. Are they monopolymers or copolymers?

Both polyesters are copolymers.

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2) What structural difference do you see in the molecules of the two polyesters? How might the structure alter the viscosity, mechanically?

The polyester formed with the glyptal is more three-dimensional and the molecules would have more opportunity to “hang up” on each other. This could increase the viscosity.

3) Is the nylon a monopolymer or a copolymer?

The nylon is a coplymer.

4) Which reactions are condensation reactions and which are addition reactions?

The polyester reaction is a condensation polymerization reaction because one H2O molecule is produced for every linkage made in the chain. The nylon reaction is a condensation polymerization reaction because a HCl molecule is produced.

POLYMERS SAMPLE REPORT SHEET I. INSTRUCTOR'S INITIALS LINEAR: _________ GLYPTAL: _________

NYLON: _________

II. ROPE LENGTH Large beaker circumference ___________cm (C=π ×.diameter) Number of turns of beaker ___________ Rope length = circumference × turns = __________cm × ________ Rope length = _______________cm III. QUESTIONS 1. Which of the polyesters is more viscous?

The glyptal is more viscous; it is thicker and flows less than the linear polymer.

2. Did the difference in viscosity fit your prediction based upon the structures of the molecules?

The structure of the glyptal causes it to be more brittle because it is cross-linked.

3. Is either polyester brittle?

They may both be brittle, but the glyptal should be more brittle.

4. Describe the properties of the nylon rope.

It is slimy feeling and fibrous. When is dries it is fragile and crumbly.

5. Did you find that the strength and brittleness of the nylon rope changed as it dried?

Support your answer with your observations and theories.

As the nylon rope dried, it became fibrous and very brittle. This could be caused by the increased amount of linking that occurs as the fluid portion evaporates.

6. Why were the reactants planned to be completely used up?

Hexamethylenediamine and sebacoyl chloride are irritating to the eyes, skin, and respiratory system. These chemicals can also damage the water system; so if they are not used completely, they can cause bodily or environmental harm.

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#27

Polymer Properties

Reagents acetone polyvinyl alcohol

10 mL/pair 4 g/pair

Common Materials 10 oz styrofoam cups paper clips 4% Borax (sodium borate) solution 8% Borax (sodium borate) solution Elmers GlueTM 12" rulers

4/pair 1/pair 50 mL/pair 35 mL/pair 60 mL/pair 1/pair

Laboratory Equipment 100 mL graduated cylinder 150 mL beaker stirring rod Special Equipment acetone resistant gloves

1 pair/student

Notes The Borax solutions required heating with constant stirring in order to get all the solid dissolved. The liquid is clear when all the borate is in solution.

Disposal All acetone should be absorbed/evaporated. Other solutions can be poured down the drain with excess water.

POLYMER PROPERTIES PRE-LAB QUESTIONS 1) Styrofoam is very brittle. What will you expect to find when you examine the hardness of styrene?

It will be brittle and hard also, but much more dense.

2) The slime produced in steps 5-7 will continue to cross -link and exclude water molecules from its matrix until all the water is gone. What will you expect to happen to the elasticity of the slime with time?

As the moisture is excluded and the cross-linking increases, the slime should get denser.

3) As water leaves the glue and borate polymers produced in steps 8-11, they become less and less elastic until, eventually, they are solids. How will you store them if you want to keep them?

Storage in a sealed container should slow the evaporation of water from the polymers.

4) Why might a toy manufacturer be interested in polymer properties?

Many children’s toys are bouncy, elastic, or very pliable. Depending on whether the manufacturer wants a soft, pliable toy for an infant, or a hard, durable toy for an older

78

child, the polymer properties are determined by the chemical makeup. The density and elasticity of a ball will determine what it is good for.

POLYMER PROPERTIES SAMPLE REPORT SHEET I. STYROFOAM TO STYRENE mass initial volume Styrofoam _________g ________mL Styrene _________g ________mL Density Styrofoam ___.048__g/mL

final volume volume ________mL ______mL ________mL ______mL Styrene ___.92___g/mL

II.

BORATE POLYMERS I II III Highest bounce __14 __cm ___20__cm ___22__cm mass initial volume final volume volume I _________g ________mL ________mL ______mL II _________g ________mL ________mL ______mL III _________g ________mL ________mL ______mL Density I __1.13__g/mL II __1.12__g/mL III __1.19__g/mL Elasticity observations (Elasticity decreases as density increases.)

I Stretches easily if pulled slowly.

II Tears if pulled very far.

III Stretches slightly.

Tensile Strength (Tensile strength increases with density of the polymer.)

I Tears when pulled quickly.

II Tears easily when pulled quickly.

III Pops apart when pulled quickly, but one must pull harder than for the other two.

II. QUESTIONS 1. Why is the density so different between styrene and styrofoam?

The styrofoam is full of tiny air pockets which make it light weight and an excellent insulator.

2. Which of the three borate polymers is the most dense?

The one made with 8% Borax and Elmer's Glue.

3. Which of the three borate polymers has the most bounce?

The one made with 8% Borax and Elmer's Glue.

4. Is there a relationship between density and bounce?

Yes, the more dense polymer had the greater bounce properties.

5. Which of the three borate polymers is the most elastic?

The one made with the polyvinyl alcohol.

6. Is there a relationship between resilience and bounce? How about the amount the ball deformed as it hit the surface?

The more resilient the polymer is, the less it deforms, but the greater the bounce.

7. Which of the three borate polymers has the most tensile strength?

The one made with the 8% Borax and Elmer's Glue.

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#28

Nuclear Radiation

Reagents alpha source beta source gamma source lead sheet

1/pair 1/pair 1/pair 1/pair

Common Materials smoke detector plain white paper aluminum foil

1 per lab several sheets/pair 1 sheet/pair

Special Equipment Geiger Counter

1/pair

Notes For demonstration purposes, finding an older style lantern mantle with the pink center or a piece of old Fiestaware pottery in orange will provide the opportunity to see that we have radiation around us in objects we use. (The glaze in the old orange Fiestaware is colored with uranium salts.)

It can be found fairly easily at flea markets and many dealers will give away broken pieces if they know they are to be used in education. Be sure the piece is not a reproduction.

Disposal There are no reagents or products to be disposed of in this investigation.

NUCLEAR RADIATION PRELAB QUESTIONS 1) Which of the three types of radiation is the most penetrating?

The gamma ray is the hardest to stop.

2) A certain percent of all carbon in living systems is radioactive carbon-14. How does this relate to "background" radiation?

It is a contributing factor in all background measurements.

3) Is food served from older pieces of orange FiestawareTM radioactively contaminated?

Would the food, itself, emit anything that would register on a Geiger counter?

No, the food, itself, would have to contain radioactive isotopes to be radioactive. No activity would be measured from the food. In order fort he dishes to be harmful to one’s health, a piece of the glaze would have to be ingested.

4) There are many uses of radiation in medicine. Radioactive sources can be used to tag the activity in specific organs of the body for diagnostic studies. Radiation is often used to reverse certain diseases. Why then are we concerned with adding radiation sources to the environment, or with increasing amounts of radiation that make it through our atmosphere?

Nuclear medicine employs radiation sources with very short half-lives, so that within a few days, there is no detectable increase over the normal level of radioactivity in the body. Environmental sources pose a different risk. Because the radiation sources are

80

constantly present, the exposure is long-term and therefore constitutes a greater risk, overall.

NUCLEAR RADIATION SAMPLE REPORT SHEET I. Background counts ___57 __counts/5 min = __11.4__ counts/min * This will vary depending on geographical location, building construction, and counter sensitivity.

II. Source 1cm 2cm 3cm 4cm 5cm 6cm 7cm 8cm 9cm 10cm * This will vary depending upon the strength of the sources but should decrease exponentially.

Beta 17.6 16.2 15.2 14.6 14.0 13.3 12.5 (background) ____ Gamma 20.6_ 18.6 16.8 15.0 13.9 (background) ____ ____ ____ _ III. Alpha _should be same as background count_ IV Beta no paper 1 sheet 2 sheets Al foil _17.5__ _17.0_ _15.2_ (background) number of sheets required to match background level __4__ V. Gamma no paper 1 sheet Al foil lead _20.5__ _20.2_ _19.2_ _9.6_ VI. Smoke Detector outside case __14.8__ next to source __15.7__ VII.QUESTIONS 1. Would you be safe from an alpha, beta, or gamma radiation source if you were a mile away?

The alpha and beta radiation would pose no problem, but if the gamma ray were energetic enough, it could travel the distance. However, the "density" of the radiation at that distance makes it unlikely to be a problem.

2. Why are people not concerned about radiation sources in our homes such as luminous clock faces, smoke detectors, indicator lights on appliances (krypton-85), electric blankets (promethium-147), and fluorescent lights (thorium -229)?

The amount of radiation is small and we are usually a distance from the sources. These are low energy sources and they would have to be ingested to pose a problem.

3. What would be the problem of using a radioactive isotope which emitted a high energy gamma in addition to the alpha in a smoke detector?

The high energy gamma would escape the case and over time be a health risk to people near it.

4. Why did people stop making luminescent watch hands with radium? (Radium emits a high energy gamma.)

The radium is worn close to the body and so the radiation can be readily absorbed.

5. If the polonium is in the tobacco, would being in the room with a smoker increase your radiation exposure?

Yes, but only a small amount. The more second hand smoke, the more radiation.

5. Describe how the Geiger counter is designed to allow a person to differentiate between beta and gamma rays.

The tube on the Geiger counter has a thick metal covering that will shield beta radiation from entering the tube and being registered. Gamma radiation easily penetrates the metal covering. The covering can be rotated to shield or to allow radiation into the tube.

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#29

Salt on Ice

Reagents isopropyl alcohol [C3H7OH] sodium chloride [NaCl] calcium chloride [CaCl2]

125 mL/pair 160 g/pair 10 g/pair

Common Materials ice

3000 mL/pair

Laboratory Equipment 1 liter beaker large test tube with a 1-hole stopper to fit thermometer Notes Disposal The solutions can be flushed down the drain with excess water.

SALT ON ICE PRE-LAB QUESTIONS 1) Salt produces ions in water and that is the process by which it lowers the freezing point of the water. However, the ions also provide the means for faster rusting of most iron containing substances. What are the implications of this practice for automobile owners in areas that have many days of freezing weather?

Salt placed on roadways is splattered upon the car causing deterioration to paint and protective coatings, and certainly rust on any exposed iron or steel structures.

2) There are salts that would deliver more ions in solution than magnesium or calcium chloride. Aluminum sulfate would deliver 5 ions per formula unit, but the sulfate is an ion that could alter the pH of the water in the environment. What are some other effects that would have to be taken into consideration before choosing a random salt for the icy roads?

Placing metal ions into the environment can result in a variety of reactions. Many of the anions have undesirable properties including the ability to alter pH. Certain ions are toxic when they enter the food chain.

3) Can you think of a reason that people don’t use calcium chloride or magnesium chloride in their ice cream freezers at home?

They are not as readily available and they are more expensive.

4) The salt that is used to melt ice on roads re-crystallizes once the ice is melted and the water has evaporated. The trouble is that most if it is on the underside of cars or on the roadside where all the slush was thrown by cars passing by. Think of an alternative to salt that would increase safety in cold weather, but not spread corrosive ions in the environment? Would it work as well for continual freezing rains?

Some liquids, including many alcohols, are effective in lowering the freezing point of water. The difficulty is in keeping them on the surface of the road. Smaller molecules of alcohol will evaporate into the atmosphere. As liquids are diluted they become less useful so that continued freezing rain will not be affected by their presence.

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SALT ON ICE SAMPLE REPORT SHEET I. The following is sample data only: Salt Added

0 g.

NaCl

10 g.

NaCl

30 g.

NaCl

50 g.

NaCl

70 g.

NaCl

10 g.

CaCl2

Original Temperature

22.3

22.3

21.8

21.5

22.0

20.3

Temperature at:

1 min.

11.0

10.5

10.0

10.0

7.5

9.5

2 min.

9.0

5.0

4.0

0.5

-1.5

3.5

5 min.

2.1

1.8

-0.5

-3.5

-4.0

1.0

10 min.

1.8

0.8

-2.0

-4.0

-5.5

2.0

Difference in temperature

20.5

21.5

23.8

25.5

27.5

78.3

Initial beaker + ice mass

569.3

533.2

594.8

576.4

597.7

588.9

minus beaker mass

289.1

289.1

289.1

289.1

289.1

289.1

Initial ice mass

288.2

244.1

305.7

287.3

308.6

299.8

Final beaker and ice mass

393.4

353.8

450.3

394.9

361.1

391.5

minus beaker mass

289.1

289.1

289.1

289.1

289.1

289.1

Final ice mass

113.3

64.7

161.2

105.8

72.0

102.4

% ice melted=

59.6

73.5

47.3

63.2

76.7

65.8

initial - final x 100% initial II. QUESTIONS 1. Is there a difference in the rate of cooling for the different amounts of salt?

Yes, the rate increased with the amount of salt added.

2. Was there a difference in the amount of frost produced on the outside of the large beaker?

Yes, faster cooling produced more frost.

3. Is it possible for the surroundings to be so cold that salt will not cause the ice to melt?

Yes, the surroundings could be colder that the lowered freezing point of the water.

4. Which amount of salt caused the most ice to melt?

70 g NaCl 5. Which amount of salt caused the lowest temperature?

70 g NaCl 6. Is it possible to add more salt than is needed?

Yes 7. Which melted ice faster, NaCl or CaCl2? Why? CaCl2 because each molecule gives 3 ions in solution while NaCl gives only 2.

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#30

Copper Cycle

Reagents clean copper turnings or wire red & blue litmus paper nitric acid, concentrated [HNO3] 8.0 M sodium hydroxide [NaOH] 6.0 M hydrochloric acid [HCl] zinc, mossy

0.500 g/pair 1 strip/pair 5 mL/pair 10 mL/pair 10 mL/pair 1 gram/pair

Common Materials Special Equipment 10 mL graduated cylinder 9.0 mm Buchner funnel with filter flask and safety flask filter paper circles to fit Buchner funnel analytical balance forceps dropper

1/pair 1/pair 2 pieces/pair 1/4 pair 1/pair 1/pair

Laboratory Equipment 150 mL beaker 250 m beaker ring stand, ring and wire gauze laboratory burner Notes: Aluminum wire also works well in place of the zinc in the last part of the procedure. A wire is sometimes easier to remove once all the color has left the solution.

Disposal Left over zinc, solid copper, and filter papers should be deposited in the trash. All other fluids may be rinsed down the drain with excess water. Unused acids and bases should be neutralized and rinsed down the drain with excess water.

COPPER CYCLE PRE-LAB QUESTIONS 1) In what instances would continued mining of a metal be preferable to recycling processes?

When recycling is more expensive, or more importantly, more unfriendly to the environment, mining would be preferable.

2) List one or two metals that would not be readily available for recycling programs because of the general manner in which they are used. Are these metals occasionally found in their pure form in nature? Are they relatively expensive or inexpensive metals?

Silver, gold, and platinum are metals used in jewelry and currency because they do not easily corrode. Such objects are not considered disposable and therefore these metals are not often available for major recycling operations.

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3) Would you expect metals on the upper or lower end of the activity series to be more easily recyclable? Explain your reasoning.

Metals that are least reactive will be the ones that are easiest to convert back to their pure, elemental state. Such metals are found at the lower end of the activity series.

4) Why must care be taken when disposing of compounds in the environment that are water soluble?

Water soluble compounds travel easily throughout the ecosystem in waterways, ground water, the atmosphere, and in animal and plant tissues.

5) Discuss the advantages and disadvantages of recycling efforts aimed at toxic metals.

Recycling toxic metals keeps them from being disposed of in landfills or in the open environment. However, the by products of the recycling process will contain small amounts of the toxic elements and these can get into the ecosystem.

6) How would you expect percent yield to be affected by increasing the number of reactions in a metal cycle? How would this most likely impact recycling processes?

The more separate reactions in a series, the greater chance of reducing the percent yield because of inefficient reactions, incomplete reactions, and side reactions.

COPPER CYCLE REPORT SHEET I. DATA AND OBSERVTIONS Part A. Preparation of Copper Nitrate Mass of copper metal used 0.485 g Maximum amount of copper that could possibly be recovered 0.485g Observations As the nitric acid begins to react with the copper, the liquid turns from clear to yellow, to light green, and finally to dark blue-green. The container becomes very hot and steam is visible. A brown-red gas is produced that is denser than air as evidenced by its tendency to remain in the bottom of the beaker.

Balanced equation Cu(s) + 4 HNO3 (aq) → 2 NO2 (g) + Cu(NO3)2 (aq) + 2 H2O (l) Part B. Preparation of Copper Hydroxide Observations When the water is added, the solution goes from dark blue-green to a lighter blue color.

As the NaOH is added, a deep blue gel immediately forms. The reaction is slightly exothermic. The gel clumps and falls to the bottom of the beaker. The supernatant liquid gradually loses its blue tint.

Balanced equation Cu(NO3)2 (aq) + 2 NaOH (aq) → Cu(OH)2(s) +2 NaNO3(aq) Part C. Preparation of Copper Oxide Observations As the blue Cu(OH)2 warms, the outside margin begins to turn brown and then black.

With continued heating the black spreads toward the center and eventually the entire precipitate is black. The CuO is a fine powder when stirred.

Balanced equation Cu(OH)2 (s) → (s) + H2O(l) + CuO (s) Part D. Preparation of Copper Chloride Observations

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As the hydrochloric acid is dribble over the black solid, it slowly appears to disintegrate as the liquid turns a yellow-green tint. Repeated applications of the liquid are required to convert all of the black solid to a dissolved compound in solution. As the product solution is diluted, the color changes from yellow-green to a light aqua tint.

Balanced equation CuO (s) + 2 HCl (aq) → CuCl2 (aq) + H2O (l) Part E. Preparation of Copper Metal Observations Very shortly after the metal is introduced to the aqua solution, the surface darkens and bubbles are seen. The reaction accelerates and becomes quite exothermic. A rust colored feathery/flaky solid is seen to form on the metal and can be dislodged by agitation. The color slowly fades from the solution until no more solid is deposited. After filtration, the solid is found to be powdery.

Balanced equation CuCl2 (aq) + Zn (s) → Cu (s) + ZnCl2 (aq) Zn (s) + 2 HCl (aq) → ZnCl2 (aq) + H2 (g) Amount of copper recovered: 0.479 g Percent yield of copper recovered 98.7 % II. QUESTIONS 1) If a student begins with 1.8 grams of copper metal and has a 67% return for the laboratory procedure, what mass of copper is present after drying? x/1.8 = .67 x = 1.2 g 2) Thinking of the implications in this procedure for the recycling of copper metals a. What might be some drawbacks to recycling processes?

Cost of reagents, cost of disposing of by-products, energy costs, environmental pollutants generated in the process, etc. are all things to be considered. b. What advantages are there to recycling copper?

Copper is not a renewable resource and it is wise to keep it in use. c. What environmental concerns come into question with the procedure?

Mining creates temporary aesthetic environmental issues. Mining can be a dangerous operation and there are by-products to the purification process for copper ores. (There will be other responses also.) d. When would recycling not be the best way to deal with waste metals?

If the recycling process is more harmful to the environment than the mining and purification processes, it would not be a prudent practice. The exception is when we run short of a particular metal and recycling becomes the only alternative.

3) If a student begins with 0.85 grams of copper metal and expects 98% return, what is the minimum mass of zinc metal he should place into the final reaction mixture?

0.85g Cu (1 mol Cu/63.55 g Cu) (1 mol Zn/1 mol Cu) (65.41 g Zn/1 mol Zn) (100%/98%) = 0.89 g Zn 4) A student gets a return of 100.5%. Give as many plausible explanations as you can think of for the obvious error.

The copper may still have a moisture content, adding mass. The copper may contain pieces of zinc. The student may have forgotten to subtract the mass of the weighing paper.

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#31

O2 and CO2 in Breath and Air

Reagents calcium hydroxide solution, saturated [Ca(OH)2]

1500 mL/pair

Common Materials small birthday candles

1/pair

Special Equipment 100 mL Erlenmeyer flask 2-hole stopper to fit the Erlenmeyer flask medium 1-hole rubber stopper (#5 or #6) wax pencils hose clamps

1/pair 1/pair 1/pair 1/4 students 2/pair

Laboratory Equipment 3 lengths of tubing 2 hose clamps 2 1000 mL beakers 250 mL beaker 100 mL graduated cylinder funnel Notes Exhaled breath is generally 79.4% nitrogen, 16.7% oxygen, and 3.8% carbon dioxide.

The hose assembly should be assembled beforehand.

Disposal The solutions can be decanted off the settled solid and flushed down the drain with plenty of water. The solids can be thrown in the trash.

O2 AND CO2 IN BREATH AND AIR PRE-LAB QUESTIONS 1) If we exhale mostly carbon dioxide, why does a person doing CPR blow into the mouth of an accident victim?

There is still some oxygen in exhaled breath. Our lungs only remove a portion of the oxygen in the air while it remains in our lungs.

2) In what common food material do we find carbonic acid?

Many beverages contain carbonic acid as a result of the carbon dioxide dissolved in them. Carbonated beverages include sodas, and sparkling juices.

3) Is the fact that fish can live in water proof that oxygen dissolves in water?

Yes, as members of the animal kingdom, fish require oxygen for survival and it must be absorbed from the water.

87

O2 AND CO2 IN BREATH AND AIR SAMPLE REPORT SHEET The following is sample data.

I. O2 IN AIR 250 mL beaker volume ___274.5_______ mL Step 14 minus final volume -___186.5_______ mL Step 14 volume O2 in air ___88.0 _______ mL % O2 in air = O2 volume x 100% = 88.0 mL x 100% = 9.8 % beaker volume 900 mL II. CO2 IN BREATH volume of displaced H2O ___121.0_______ mL Step 11 volume through Ca(OH)2 -___ 115.0_ _____ mL Step 12 volume CO2 in breath _____6.0 _____ mL % CO2 in air = CO2 volume x 100% == 6 mL x 100% volume of displaced H2O 12 mL III. O2 IN BREATH unlighted candle volume __________ mL lighted candle volume -__________ mL volume O2 in breath __________ mL % O2 in breath = O2 volume x 100% = = volume of displaced H2O

= 50 %

Step 15 Step 16 mL x 100% mL

=

IV. QUESTIONS 1. What would be the effect on your results if the CO2 did not come in contact with the Ca(OH)2 solution?

The reaction would not occur.

2. What difference would you find if you pumped air from an air compressor instead of blowing air from your lungs in order to saturate the water?

The compressor air would not saturate the water quickly because the CO2 content in regular air pumped by the compressor is low compared to that from the lungs.

3. What difference would you find if you pumped air from an air compressor instead of blowing air from your lungs in filling the flask?

Air from a compressor is taken from the atmosphere and has different composition than air from lungs. It has much less CO2 and much more O2.

4. If some other compound besides calcium hydroxide were to be used, what is the necessary requirement for that compound?

A compound that could produce a solid with CO2 , acid, and water would be necessary. It would need to form an insoluble salt of carbonate.

5. What would be the difference if calcium hydroxide were not a base?

Since it would be acidic or neutral, the CO2 gas would not be produced.

6. Would CPR work if your breath had no oxygen in it?

No, our lungs do not remove all of the oxygen from the air we breathe before it is exhaled. If they did, we would not be able to deliver oxygen to a victim.

88

%

#32

Solids in Smoke

Reagents Common Materials cigarettes:

non-filter filter little cigar

matches Special Equipment glass tubing student manual filter paper accurate scale (to +0.001 g)

refer to diagram in 4 pieces/pair 1/lab

Laboratory Equipment 2 filter flasks 2 funnels 1-hole stoppers to fit filter flasks 2 glass tubing sections to fit stoppers ring stand 2 utility clamps eyedropper Notes Disposal Throw all solid residue in the trash.

SOLIDS IN SMOKE PRE-LAB QUESTIONS 1) How will the measurement of smoke from the burning end of a cigarette differ from the smoke that would be exhaled into a room by a smoker?

The smoke exhaled may have passed through a filter before entering the lungs of the smoker and contain less of the solids than the smoke from the end of the cigarette. The exhaled smoke will also have lost some of the particulate matter and oxygen that remained in the lungs of the smoker, as well as gaining some carbon dioxide.

2) We will measure all the solids in the smoke that come from the cigarette. How will this be different than the actual solids a smoker gets from a cigarette?

The smoker exhales some of the solids, not all of them remain in the lungs and airways.

3) Who do you predict will profit from the use of a filtered cigarette? (the smoker, others in the room, or both the smoker and others in the room) Only the smoker will benefit from the filter that removes some of the solids and other matter from the smoke. The majority of the smoke in the room has come from the burning end of the cigarette.

4) A person who claims they do not inhale and are thereby protected from the risks of smoking are forgetting about something important. What is it?

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The radiation from the cigarette is present whether or not they are inhaling the actual smoke.

SOLIDS IN SMOKE SAMPLE REPORT SHEET PART I: SOLIDS IN CIGARETTE SMOKE (The following is sample data.)

Puffed Non-filter Non-filter Filter Little Cigarette Cigarette Cigarette Cigar Cigarette Flask Filter Paper Mass: (initial) 0.890 g (after collecting 0.950 g smoke) Mass of Solids Collected 0.06g Cigarette Mass: (initial) 1.03 g (after burning) 0.80 g Mass of Cigarette Consumed 0.23g g solids  1000 261 g cig consumed

smoker solids  100% room solids

Funnel Flask 0.900 g 0.910 g 0.01g

43.5 600%

Cig. Fl.

Fun. Fl.

Cig. Fl.

Fun. Fl.

Cig. Fl.

Fun. Fl.

0.860 g 0.900 g 0.880g 0.920 g 0.880g 0.890g 0.930 g 0.910 g 0.940g 0.920 g 0.910g 0.900g 0.07g 0.01g

0.06g 0.01g

0.03g 0.01g

1.04 g 0.47 g

1.03 g 0.44 g

6.13 g 4.19 g

0.57 g

0.59 g

1.94 g

123

17.5

102

703%

16.9

15.5

602%

5.15 301%

QUESTIONS 1. Which of the cigarettes in Part I produced the most solids in smoker smoke and in room smoke? Can you explain the reason for your results?

The puffed cigarette produced more solids. Since the combustion reaction is happening at a faster rate, it is not as complete and therefore produces more ash.

2. Which cigarette produced the lowest ratio of smoker smoke to room smoke?

The little cigarette had the lowest ratio.

3. What conclusions can be drawn from the data about being in a room with a smoker?

It is not healthy because smoke solids are still breathed by the nonsmoker.

PART II: EFFECTIVENESS OF A CIGARETTE FILTER Smoker Smoke 1. Initial Filter Paper Mass ____0.88_____ g 2. Filter Paper Mass After 1/3 Cigarette Consumed _____0.91____ g 90

Room Smoke ___0.90____ g

___0.91____ g

3. Filter Paper Mass After 2/3 Cigarette Consumed 4. Filter Paper Mass After Last Third Consumed 5. g Solids From First 1/3 6. g Solids From Second 1/3 7. g Solids From Last 1/3 8. First 1/3 Ratio Smoker Smoke/Room Smoke 9. Second 1/3 Ratio Smoker Smoke/Room Smoke 10. Last 1/3 Ratio Smoker Smoke/Room Smoke

_____0.93____ g

___0.92 ___ g

_____0.94____ g

___0.93____ g

_____0.03____ g

___0.01__ g

_____0.02____ g

____0.01___ g

_____0.01____ g

____0.01___ g

____3/1_____ ____2/1_____

EXTRA Units are not necessary on these numbers since they are ratios.

____1/1_____

QUESTIONS (CONT.) 4. Which portion of the cigarette in Part II yielded the most solids in smoker smoke? Explain your result.

The first part had the most solids. (Although this data does not show it, usually the last portion produces the most solids as the smoke from the earlier parts are filtered by traveling through the tobacco in the last portion.

5. Does the portion of the cigarette effect the amount of solids in the room smoke? If so, why?

No, the smoke released to the room comes from the burning of the tobacco at the end of the cigarette and is not filtered by the tobacco.

6. Does the filter protect a smoker from solids? Does a filter protect someone nearby?

Yes, the smoker receives fewer solids. However, the smoke inhaled by someone nearby does not pass through the filter and therefore contains much more particulate matter.

91

#33

Clarification of Water

Reagents 6M sodium hydroxide [NaOH] 3M sulfuric acid [H2SO4] manganese dioxide (1.0 g MnO2/200 mL to make 0.012 M) sodium hydrogen carbonate (sodium bicarbonate) [NaHCO3] pH paper (wide) range ethyl or isopropyl alcohol

15 mL/pair 20 mL/pair 30 mL/pair 1.5 g/pair 1 roll/pair 50 mL/pair

Common Materials aluminum cans, cut into strips 1 cm × 3 cm fabric dye activated charcoal dry ice, broken or crushed food coloring

2 strips/pair 0.5 mL/pair 0.5 g/pair 100 mL/pair 2 drops/pair

Special Equipment fume hood centrifuge spectrophotometer set at 500 nm (optional) metal cutter 125 mL Erlenmeyer flask inorganic waste container filter paper

1/lab 1/lab 1/lab 1/lab 1 more/pair 1/lab 1/pair

Laboratory Equipment ring stand, ring and wire gauze laboratory burner 150 mL beaker 2 Buchner filter systems evaporating dish 100 mL graduated cylinder mL graduated cylinder pH paper 3 small test tubes stirring rod Notes If there is a problem, it can probably be traced to starting with too much aluminum. One strip may be enough. It is not necessary to dissolve all the Al(OH)3 formed by adding H2SO4.

Surplus Al(OH)3 is removed in the first filtration. Un-neutralized H2SO4 is a problem. The spectrophotometers are not necessary to be able to see some clarification. The experiment is a bit long, so the instructor may want to make those measurements for the class.

Disposal

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The solutions can be flushed down the drain with a lot of water. The solids containing aluminum or magnesium should be collected for burial in a toxic waste site. The charcoal should be collected so that it can be reactivated. The charcoal can be regenerated by heating it to 5001000oC in the presence of steam or carbon dioxide.

CLARIFICATION OF WATER PRE-LAB QUESTIONS 1) Why is speed important in the settling of solid particles from water in a municipal water treatment plant?

There is a large amount of water required from municipal water treatment facilities and if an inordinate amount of time is required in the settling tanks, the rate of purification is slowed.

2) In the initial reaction with the aluminum and sodium hydroxide, what is the source of the hydrogen gas?

Most of the hydroxide ions bonded to the aluminum to form the Al(OH)4- come form the sodium hydroxide. However, the aluminum also reacts with the water in a single replacement reaction. The hydroxide ion from the water bonds to the aluminum and the remaining hydrogen is released in the form of hydrogen gas.

3) Aluminum is actually very reactive with water. From the equation representing the first reaction in the process, what substance is allowing the aluminum to react with the water?

Aluminum reacts with the oxygen in the air to form a layer of aluminum oxide that is similar to silicon oxide, or glass. The sodium hydroxide is very basic and this is the key to allowing the aluminum to react with the water.

4) People who live in wilderness areas or rural areas that are quite distant from the nearest city must clean their own water, depending on the purity of the source to which they have access. List the substances and materials that would be necessary for a person to setup their own mini water treatment plant?

In order to clean water, a person would need a settling tank or a flocculant, activated charcoal.

5) What final step would you suggest the person in question #3 take to be certain the water is safe to drink?

Perhaps a bit of chlorine bleach could be added if there is a suspicion of bacteria in the water supply.

CLARIFICATION OF WATER SAMPLE REPORT SHEET I. ALUMINUM HYDROXIDE Compare the volume of aluminum hydroxide obtained to the original volume of aluminum metal.

It would be larger than the original volume.

II. USE OF ALUMINUM HYDROXIDE Absorbance of test tube with dirty water

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1.18

.

Absorbance of test tube with centrifuged dirty water _____0.15____________ Absorbance of test tube with clarified water _____0.06____________ III. USE OF CHARCOAL Absorbance of test tube with dirty water ________0.62__________ Absorbance of test tube with clarified water ________0.095_________ IV. QUESTIONS 1. What were the bubbles made of which formed when aluminum reacted with alkali?

The bubbles are hydrogen gas.

2. Is it easy to see aluminum hydroxide floating in water? Explain.

Yes, the aluminum hydroxide is a dark solid.

3. What were the bubbles made of which formed when sodium hydrogen carbonate was added to the sulfuric acid solution?

CO2 gas 4. What happens to the colorful decorations on the aluminum can when you dissolve it?

The colorful decorations remain on a thin film.

5. Suppose the activated charcoal absorbed a very toxic substance (not heavy metals). What would you do with the spent charcoal?

If the charcoal was not able to be re-activated then it would be disposed of in a safe container and buried in a toxic waste site.

6. Is the clear water you have prepared now safe to drink? Why or why not?

No, the water is not safe to drink because it still contains dissolved substances and perhaps bacteria.

7. Why isn't a centrifuge used to clean the city water?

This is not a practical or logical way to clean water because it would take an enormous amount of energy and time to centrifuge a city’s water supply.

8. Is the food coloring separated from the water as it freezes?

Yes, the water freezes and excludes impurities in the process. This and other state changes can be used as purifying processes.

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#34

Freezing Water

Reagents alcohol (isopropyl or ethyl)

300 mL/pair

Common Materials dry ice pieces

100 mL/pair

Special Equipment grease pencils Laboratory Equipment ring stand, ring and wire gauze 2 250 mL Erlenmeyer flasks 1000 mL beaker 4 large test tubes

100 mL beaker watch glass 100 mL graduated cylinder thermometer

Notes The order in which the solutions freeze may not always be consistent. It may also vary within the class itself.

Disposal The alcohol can be flushed down the drain with plenty of water. The left over dry ice can be left to sublime or evaporate without melting.

FREEZING WATER PRE-LAB QUESTIONS 1) Why do fish prefer cooler water?

Cooler water will have more dissolved gases, most importantly, more oxygen.

2) Use the knowledge from the introductory paragraphs to explain why the hot water pipes in the exterior walls of poorly insulated houses many times freeze before the cold water pipes.

The oxygen and other dissolved gases have been removed from the heated water and therefore it is pure water. Without the “impurities” the freezing point is not lowered and therefore it will freeze at the expected temperature, and not at some lower temperature as will the water with dissolved gasses and other materials.

3) In preparing tap water for fish tanks, some people first boil it to remove the chlorine gas and then allow it to sit to cool. What additional step should be taken before expecting fish to be able to live in the water? Explain.

Oxygen needs to be re-introduced to the water so that the fish have it available for respiration.

4) Why is freezing water a concern in cracks of roadways and bridges?

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As water freezes it expands and places a tremendous amount of pressure on the structures in which the ice crystals are forming. This further damages cracking rocks and concrete as it forces the cracks open.

FREEZING WATER SAMPLE REPORT SHEET I. FREEZING WATER* Initial time _______ _______

Frozen Minutes time difference _______ ________ _______ ________

Hot boiled water Tap water-cool, unboiled water-cool, boiled and covered _______ _______ ________ Tap water-cool, boiled and stirred _______ _______ ________ * Times will vary greatly within lab sections, but should be under 4 minutes.

Tap

II. QUESTIONS 1. Which freezes first? second? third? last?

This will vary depending upon how the investigation is set up. Expected would normally be 1, 2 and 3, 4.

1. tap H2O, boiled, covered, and cooled 2. tap H2O, boiled, stirred, and cooled 3. tap H2O, un-boiled and cooled 4. hot boiled H2O 2. What is the effect of the initial temperature?

The higher the initial temperature, the longer it takes to freeze.

3. What is the effect of the dissolved gases?

The dissolved gases increase the time it takes the solution to freeze.

4. Is there a difference between the tap water and the cool, boiled, stirred water? If so, why?

There can be some difference because the amount of dissolved gases in the cooled, boiled, and stirred H2O would be less than the tap H2O even though it was stirred.

5. Could fish survive in a frozen pond if water did not dissolve gases?

No, fish require O2 dissolved in the H2O in order to survive.

6. Could fish survive in a frozen pond if water froze from the bottom up?

No, the fish would not survive because they would be frozen as the water continued to freeze. The top layer of ice provides an insulation to protect the lower portion of the water from freezing.

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#35

Freezing Point Depression

Reagents isopropyl or ethyl alcohol glycerin

50 mL/pair 12 mL/pair

Common Materials dry ice pieces paper towel (or cloth) to wrap test tube

20 mL/pair 1 roll/pair

Special Equipment 18 mm × 150 mm test tube (to fit in 25 mm × 200 mm) “low temp” thermometers (optional)

1/pair 1/pair

Laboratory Equipment 10 mL graduated cylinder 25 mm x 200 mm test tube 100 or 150 mL beaker dropper thermometer stirring rod centigram balance Notes Using 10 mL graduated cylinders in place of the smaller test tube will shorten the procedure since it is not necessary to continually empty the contents to measure the volume.

This alteration to the procedure also eliminates the error introduced in the volume measurement each time the liquid is transferred between containers.

Place a large test tube in dry ice before adding alcohol as the dry ice alcohol mixture may freeze solid. This preserves a place to submerge the test tube for cooling. It is also advisable to clamp the larger test tube in place.

The liquid in thermometers that are quickly cooled to low temperatures can sometimes separate. Using thermometers that are specifically designed for use in low temperature applications will help prevent this from occurring. They are also easier to read since the scale will be above the lip of the test tube.

The order of freezing will not be predictable because there are so many variables involved.

Expect the students to get differing results.

Disposal The materials used in this investigation can be flushed down the drain with plenty of water.

The alcohol used with the dry ice can be strained through cheesecloth to remove any foreign solids and then be labeled and saved for use in this investigation in another semester or for another section.

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FREEZING POINT DEPRESSION PRE-LAB QUESTIONS 1) Ethylene glycol is a covalently bonded molecule and each molecule remains intact as it dissolves in water. What would be the effect on the radiator and engine of an automobile if one used a salt that dissociates into ions (charged particles) to lower the freezing point of the water even more.

Using a substance that dissociates into more than one particle would increase the “effective” concentration of dissolved particles. The result would be a greater degree of lowering of the freezing point.

2) Could another compound that is not soluble in water be used as antifreeze?

The compound would not be effective because it needs to be soluble in H2O to lower its freezing point.

3) Since glycerin and water are completely miscible, enough of the antifreeze can be added until the antifreeze becomes the solvent and the water is the solute. In this case, which freezing point is being lowered, technically; that of the water or that of the antifreeze?

Explain.

The solvent is defined as the component of a solution that is in the majority, generally accepted as more than 50%. If the amount of antifreeze were increased until it became a larger percentage than the water, it would then be classified as the solvent and the water as the solute.

4) From your explanation in question #3, will there be a limit to how low the freezing point of water in a radiator can be pushed? Why? Yes. The freezing point can only be as low as the lowest point when one of the substances is acting as the solvent at 50% concentration and the other substance is then the solute at its greatest concentration.

FREEZING POINT DEPRESSION SAMPLE REPORT SHEET I. FREEZING POINT: 100% glycerin 67% First Run Second Run Average +17.0ºC -38.4ºC 0% of glycerin in water (pure water) freezes at 0ºC. *calibration and accuracy will vary Change the average freezing points to the other common temperature scale (i.e. ºC → ºF or ºF → ºC) Average F.P. _____62.6ºF___ __-37ºF_

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50%

33%

-23.0ºC

-11.2ºC

EXTRA ºF=9/5ºC+32º ºC=5/9(ºF-32º) __9.4ºF__ __11.8ºF__

II. DENSITY: 100% glycerin 67% 50% 33% Beaker + mixture g g g g minus Beaker – _ g g g g Mixture g g g g Density = mass in grams = 1.26 g/mL 1.17 g/mL 1.13_g/mL _1.08_g/mL 3 mL 0% glycerol in water (pure water) has a density of 1.0 g/mL. III. QUESTIONS 1. From your graphs, what density is required to protect your automobile's cooling system to _10ºC? (Assume glycerin behaves similarly to ethylene glycol.)

About 1.1 g/mL 2. Should you use pure antifreeze in your car radiator? Why or why not?

No, pure antifreeze does not lower the freezing point as low as a solution of antifreeze and H2O. 3. What is the approximate percentage mixture that has the lowest freezing point? What density is that?

The lowest freezing point should be at 67% ethylene glycol. The density should be lower by about 1.2 g/mL. 4. When a check on an automobile radiator's system is made, the amount of antifreeze is tested by measuring the density of the solution in the radiator. Why does the serviceman know to what temperature the car is protected by measuring density?

The density can determine the ratio of antifreeze to H2O and to what temperature the car is protected.

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#36

Heat Capacity

Reagents ethanol [CH3CH2OH]

10 mL/pair

Common Materials styrofoam cup metal pellets

1/pair 10 g/pair

Laboratory Equipment 2 thermometers 2 600 mL beakers ring stand, ring and wire gauze laboratory burner 10 ml graduated cylinder small test tube Notes It is important that no water be on the pellets as the heat capacity of water is so much larger that it could change the results.

Alcohol density is 0.7899/mL Specific heats of some metals are: aluminum 0.216 cal/g·ºC copper 0.092 cal/g·ºC iron 0.107 cal/g·ºC lead 0.031 cal/g·ºC silver 0.056 cal/g·ºC Disposal The ethanol solution can be flushed down the drain with plenty of water.

HEAT CAPACITY PRE-LAB QUESTIONS 1) Suppose equal amounts of heat are added to two equal masses of metals at the same initial temperature. Metal A has a specific heat that is half that of Metal B. Which metal will have the highest final temperature? Explain your reasoning.

Metal A will have the highest final temperature because it will raise two degrees for every degree increase in Metal B. 2) Styrofoam containers are insulators and transmit very little heat from their contents to the environment. The interior of the Styrofoam maintains the same temperature as its contents. The specific heat of Styrofoam is 0.209 J/g Co, a very small value. Will the heat transferred to the Styrofoam cup present a major error in the laboratory data? Why or why not?

The Styrofoam cup will not absorb enough heat to make a difference in the data because it has a relatively low specific heat and it does not pass heat on to the environment. In

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other words, once it has absorbed enough heat to be at the same temperature as the contents, it will not absorb any more.

3) Metals are notorious for being excellent conductors of heat. What precautions should be taken to minimize the heat loss from the heated metal pellets in the experiment?

In order to expose the pellets to the air for as little time as possible, make the transfer between the hot water bath and the cooler water in the cup quickly.

4) If the metal pellets are wet when they are heated and added to the cool water, how will the heat contained in the water on the metal effect the final temperature?

Water has a relatively high specific heat and therefore every drop of water carries a great deal of heat. The final temperature of the mixture would be higher than it should be and the final results would make the metal appear to have a much higher specific heat than actual.

HEAT CAPACITY SAMPLE REPORT SHEET I. CALCULATED FINAL TEMPERATURE OF MIXTURES 10 g H2O and 10 g H2O

Tf =

10 g H2O and 15 g H2O

Tf =

Thot + Tcold = 2

82C + 6C = __44°C 2

(mass  Ti ) cold + (mass  Ti ) hot mass cold + mass hot (10 g  4C) + (15 g  84C) = = __52°C 10 g + 15 g

(mass  sp. ht.  Ti ) alcohol + (mass  sp. ht.  Ti ) water (mass  sp. ht.) alcohol + (mass  sp. ht.) water 0.588 cal 1.00 cal ( ___7.9 g   __5 C) + (10 g   __85C) g C g C = 0.588 cal 1.00 cal ____7.9 g  + 10 g  g C g C = ______60°C

10 mL alcohol and 10 g H2O Tf =

10 g pellets and 10 g H2O Tf =

(mass  sp. ht.  Ti ) pellets + (mass  sp. ht.  Ti ) water (mass  sp. ht.) pellets + (mass  sp. ht.) water

=

__0.449 cal 1.00 cal  _____81C) + (10 g   __4C) g C g C ___0.449 cal 1.00 cal _____ 10 g  + 10 g  g C g C

( ___10 g 

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= ______28°C II. COMPARISON OF CALCULATED AND MEASURED FINAL TEMPERATURES Calculated Measured 10 g H2O and 10 g H2O

___44°C ___

_________

10 g H2O and 15 g H2O

___52°C ___

_________

10 mL alcohol and 10 g H2O

___60°C ___

_________

10 g pellets and 10 g H2O

___28°C ___

_________

III. QUESTIONS 1. Explain why heat capacity is part of the reason that deserts are hot during the day and cold at night.

The heat capacity of soil is much less than water's heat capacity. During the day in the desert the temperature is raised as the soil and air is heated. At night the heat is lost into the air.

2. Considering the specific heats, why is water used instead of alcohol for cooling such things as engines and saw blades?

Water has a higher specific heat so less water is required to carry off the same amount of heat.

3. Why is the final temperature of the alcohol and water mixture, and the pellet and water combination not a numerical average of the two initials temperatures?

The final temperature is not an average because it has to include the specific heat of the materials. Water and alcohol and water and metal have different specific heats.

4. Which will be hot to touch when left in the sun, a piece of metal or a piece of styrofoam?

The metal will be hotter than the styrofoam.

5. List any reasons why the calculated and the measured temperature are not the same.

Heat loss during the mixing, spillage, heat loss to air, heat loss to the styrofoam.

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#37

Heat of Reaction

Reagents 1.5 M solution sodium hypochlorite [NaClO] 15 mL/pair 1.5 M sodium sulfite solution [Na2SO3] 15 mL/pair ammonium nitrate [NH4NO3] 10 g/pair Common Materials 6.4 oz. cup styrofoam 1/pair 8.5 oz. cup styrofoam 1/pair Laboratory Equipment 1 25 x 200 mm test tube solid stopper to fit test tube filter paper thermometer Notes Two styrofoam cups of the same size will work if the rim is trimmed from the top cup.

Disposal Flush the solutions down the drain with plenty of water. Although nitrates are harmful if they are consumed in water, the amount used in this experiment is not large enough to be a problem.

HEAT OF REACTION PRE-LAB QUESTIONS 1) List all the components in the system that will absorb small amounts of heat, besides the water.

The glass of the test tube, the thermometer, and the styrofoam cup will all absorb small amounts of heat.

2) Considering question #1, will the lab results err on the side of there appearing to be more or less heat exchanged than actual in an exothermic reaction? How about in an endothermic reaction? Explain.

In an exothermic reaction, some of the heat generated will be absorbed by the system and the reaction will appear to be less exothermic. In an endothermic reaction, some of the heat will be absorbed from the system (i.e,. not all from the waste) and it will therefore appear to be less endothermic.

3) Many acids and bases have very high heats of solution. Think of some practical precautions a person should take if they wanted to dilute an acid or base?

Anytime a person is working with dangerous or corrosive substances they should wear protective clothing, goggles, and gloves. Diluting acids and bases that may have high heats of solution would require a spatter proof area, such as a fume hood, in addition to the caution of always pouring the acid or base into the water to dilute it, and not the other way around.

4) What will be the effect on the heat measurement in the second reaction if the final temperature is taken before all of the solid is dissolved?

The final temperature will be too high.

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HEAT OF REACTION SAMPLE REPORT SHEET I. SODIUM SULFITE AND SODIUM HYPOCHLORITE 1. Initial Temp. __23.5__ 6 __25.5__ 12 ________ 1 (Minutes) __24.5__ 7 ________ 13 ________ 2 __24.9__ 8 ________ 14 ________ 3 __25.3__ 9 ________ 15 ________ 4 __25.5__ 10 ________ 5 __25.5__ 11 ________ 2. Temperature Difference ∆t = _25.5oC _ _23.5oC = __+2.0oC Final Initial 3. Water Volume ___23.5 __mL Reaction Solution Volume ___45.0__ __mL 4. Heat of Reaction 1 g 1.0 cal   ×_2.0_ oC = 120 calories ∆H = ( __23.5_mL + __45.0_mL) × mL g C Volume: Reaction Cp(H2O) ∆t Water Solution The 1 g/mL is a conversion factor. Cp(H2O) equals 1.0 calorie per gram degree.

II. AMMONIUM NITRATE AND WATER 1. Initial Temp. __24.0 _ 6 __4.7 __ 12 ________ 1 (Minutes) __15.0 __ 7 ________ 13 ________ 2 __6.0 __ 8 ________ 14 ________ 3 __4.0_ _ 9 ________ 15 ________ 4 __4.5 ___ 10 ________ 5 __4.5 __ 11 ________ o 2. Temperature Difference ∆t = __4.7 _ C _ 24.0__oC = __-19.3_oC Final Initial 3. Water Volume ___23.0 ___mL Reaction Solution Volume ___20.0 ___mL 4. Heat of Reaction 1 g 1.0 cal  ∆H = ( _23 mL + _20__mL) × × __-19.3_oC = -830 calories mL g C Volume: Reaction Cp(H2O) ∆T Water Solution III. QUESTIONS: 1. How would ignoring the heat absorbed by the styrofoam cup, the thermometer, and the glass test tube affect the heat of reaction measured?

The amount of heat absorbed by the cups, thermometer, and test tube is very small.

However, if one wanted very accurate results they would need to be taken into consideration else the heat of reaction measured is too low.

2. Do the reaction solutions actually have the same heat capacity as water?

No, the solutions do not have the same heat capacity as water but the difference is small.

3. Which reaction is exothermic?

Sodium hypochlorite and sodium sulfite.

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#38

Rates of Chemical Reactions

Reagents 0.1M sodium iodide solution [NaI] 30% hydrogen peroxide [H2O2] 3% hydrogen peroxide [H2O2] Common Materials rulers Special Equipment timers or watches with second hands digital analytical balance (+ 0.001g) hot air dryers Laboratory Equipment rulers 50 mL graduated cylinder

7 mL/pair 50 mL/pair 50 mL/pair 1/pair 1/pair 1/pair 1/pair

Notes Have the students bring hand-held hair dryers.

Be sure to watch procedure and use dilute (3%) hydrogen peroxide where called for.

Disposal Flush the solutions down the drain with plenty of water.

RATES OF CHEMICAL REACTIONS PRE-LAB QUESTIONS 1) Reactions are often controlled by temperature. How do we slow reactions involving food in our homes?

Food is generally kept in a refrigerator or freezer to retard the natural processes that would cause it to spoil.

2) Some reactions require an input of energy to begin. The added energy can come from several sources – friction, heat, light, an electrical spark, or a current. In some reactions, once the molecules have enough energy to get over the energy barrier, the reaction proceeds at a very fast rate. What method do we use to give the molecules on the tip of a match enough energy to get over the energy barrier?

Friction is used as the source of energy to start the reaction at the tip of a match.

3) Most of the biochemical reactions that are necessary for life occur in solution. Why does our body raise its temperature when we are ill?

Higher temperatures generally speed up reactions. At higher temperatures, the body can better fight infections.

4) What is the name we give to a catalyst in many biochemical and physiological processes?

A catalyst in a physiological or biochemical system is often referred to as an enzyme.

5) How does the concentration of oxygen in your blood affect the rates of respiratory reactions within the cells?

If the level of oxygen in the blood is elevated, the rate of respiration speeds up.

RATES OF CHEMICAL REACTIONS SAMPLE REPORT SHEET I. Observations of H2O2 before adding NaI

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The H2O2 looks very much like water. If it sets very long, tiny bubbles may collect on the sides of the container.

II. Volume of 0.1 M NaI added __2_______mL *The following is sample data only.

3% with NaI 30% Time (minutes) Mass (g) Mass (g) Mass (g) 0 97.6062 99.6477 99.908 1.0 97.6036 99.6464 99.904 2.0 97.6030 99.9445 99.900 3.0 97.6022 99.6432 99.897 4.0 97.6014 99.6411 99.894 5.0 97.6008 99.6391 99.892 6.0 97.5999 99.6357 99.890 7.0 97.5994 99.6316 99.888 8.0 97.5986 99.6266 99.886 9.0 97.5984 99.6231 99.884 10.0 97.5979 99.6190 99.882 11.0 97.5978 99.6148 99.881 12.0 97.5974 99.6108 99.879 13.0 97.5970 99.6070 99.877 14.0 97.5969 99.6020 99.876 15.0 97.5965 99.5988 99.875 3% with NaI 30% -4 -3 III. Rate = W8 - W12 = _3 x 10 g/min 3.95x10 g/min 1.75x10-3g/min T12 - T8 IV. Class Values: 3.0x10-4 g/min 2 mL vol. NaI -3 7.5x10 g/min 4 mL vol. NaI 1.03x10-2 g/min 6 mL vol. NaI _4.2x10-2 g/min 8 mL vol. NaI V. What is the rate per minute per mL of NaI? Average - 2.25x10-3 g/min-mL VI. QUESTIONS 1. Did heating cause a faster reaction? Yes 2. What would be the effect of cooling H2O2 in a refrigerator?

The gas molecules would not be as agitated and would not be given off as fast.

3. What was the effect of increasing the concentration of the catalyst (NaI)?

The reaction rate increased.

4. What was the decomposition rate before adding the catalyst?

3 x 10-4g/min for 3% H2O2 5. Were these chemicals environmentally safe? Yes 6. What was an appropriate way to store H2O2?

It should be stored in an expandable bottle in a cool, dark place.

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#39

Canned Fuel and Evaporation

Reagents calcium acetate [Ca (C2H3O2)2] 3.2 g/pair 6M sodium hydroxide [NaOH] <1 mL/pair ethanol [CH3CH2OH] (denatured) 70 mL/pair phenolphthalein 0.2 mL/pair acetone 5 mL/pair Common Materials mineral oil 5 mL/lab matches 1 box/pair cotton balls several/pair rubber bands several/pair Laboratory Equipment 10 mL graduated cylinder 50 mL beaker 100 mL graduated cylinder 150 mL beaker watch glass funnel with vacuum tubing Notes Students may bring an empty shallow can with cover if they want to save their canned fuel.

Many chip dip cans have plastic lids and are perfect.

Disposal Any left over gel should be placed in an open container in the hood. When the alcohol is evaporated, the residue may be placed in the trash.

CANNED FUEL AND EVAPORATION PRE-LAB QUESTIONS 1) How might the rate at which the alcohol burns be different if it were not held in the gel?

The gel keeps the alcohol from moving around and restricts the combustion reaction to the alcohol held and exposed at the surface. Without the gel, the alcohol would burn much faster.

2) If canned fuel is kept tightly sealed, it will last almost indefinitely. Does that mean that no alcohol is evaporating from the surface of the gel? Explain your answer.

Once enough alcohol has evaporated to saturate the space between the can lid and the surface of the gel, the amount of alcohol in the vapor state will remain constant. However, individual molecules of alcohol will continue to evaporate and condense, just at equal rates.

The system will have established equilibrium.

3) The “heat index” is a measurement used by weather forecasters to give us an indication of how the combined effects of temperature and humidity affect our bodies. Why does it feel hotter in the summer when the humidity is high?

If the amount of water vapor held in the atmosphere is high, the water from the surface of our bodies will not evaporate as readily and the cooling effect of perspiration is reduced.

4) How will our measurement of the energy absorbed by the alcohol as it evaporates be affected by atmospheric humidity? Explain your reasoning.

The humidity in the room should have no effect since only the amount of alcohol vapor in the air would have an effect on evaporating alcohol.

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CANNED FUEL AND EVAPORATION SAMPLE REPORT SHEET I. INSTRUCTOR'S INITIALS FOR ALCOHOL GEL ________ II. EVAPORATION* Sample Time start Finish Difference Alcohol ________ ________ ________ Fanned alcohol ________ ________ ________ Alcohol in a vacuum ________ ________ ________ Heated alcohol ________ ________ ________ Oil covered alcohol ________ ________ ________ * The times will vary between students but the order of evaporation should be: heated alcohol, alcohol in a vacuum, fanned alcohol, alcohol, and oil covered alcohol.

III. COMPARISON OF RATE OF EVAPORATION Sample Time Temperature Start Finish Difference Start Finish Difference Alcohol ______ ______ ______ ______ ______ ______ Acetone ______ ______ ______ ______ ______ ______ IV. QUESTIONS: 1. Why was the NaOH added to the canned fuel? Why add phenolphthalein?

The NaOH was added because it is a base, and phenolphthalein turns pink in the presence of a base. The only purpose is to color the gel.

2. Would this gel work as a fuel source in a car? Why or why not?

No, because the fuel in a car must be able to vaporize quickly for combustion in the cylinder. The alcohol fuel in the gel is not free to go to the gaseous state so quickly.

3. Which sample of alcohol, besides the sample with the oil barrier, evaporated the slowest?

The alcohol that was left to sit evaporated the slowest.

4. Why did each of the others evaporate more quickly?

Energy was added to their systems which caused the evaporation to speed up, or the vapor pressure above the surface was kept low by continually removing the vaporized particles.

5. What do fanning and the vacuum have in common?

They removed the gaseous molecules which allowed room for more liquid molecules to evaporate.

6. Does the material that evaporates the fastest cause the greatest difference in temperature?

Yes 7. If acetone could be held in a gel, would it burn like alcohol?

Yes, it is flammable.

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#40

Batteries

Reagents 1M copper sulfate [CuSO4] conc. sulfuric acid [H2SO4] 1M aluminum sulfate [Al(SO4)3] potassium sulfate [K2SO4] copper metal strips small gauge copper wire

25 mL/pair 2 mL/pair 25 mL/pair 3 g/pair 1/pair 6-8 in./pair

Common Materials aluminum foil 4 in. x 4 in. square oranges, lemons, or grapefruits steel wool, fine cotton balls paper clips clean, shiny coins (penny, nickel, dime) paper towels

1/pair 1/pair 1 small piece/pair 6/pair 2/pair 6 coins/pair 1 sheet/pair

Special Equipment voltmeter 12 cm. flexible clear plastic tubing

1/lab 1/pair

Laboratory Equipment 30-mL beakers 10-mL graduated cylinder 100-mL graduated cylinder

3/pair 1/pair 1/pair

Notes Large aquarium tubing (3/8 in.) is inexpensive and available at large chain supermarket and discount stores.

Disposal Evaporate the solutions to dryness and place the dry salts in containers to be buried.

BATTERIES PRE-LAB QUESTIONS 1) Why must two different metals be used for the two electrodes in a battery?

If the same metal was used for both electrodes, there would be no difference in the reduction potential, hence no potential for the electrons to cross.

2) Would you expect the voltage in a voltaic cell to increase or decrease the further apart the two metals appear on a table of Standard Reduction Potentials? Explain your reasoning.

The larger the difference in reduction potentials, the greater the force on the electrons.

It would then be expected that there would be a greater current.

3) What will happen if the solution inside the salt bridge develops a gap? Why? If the salt bridge develops a gap, the current cannot flow through it to complete the circuit with the battery.

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4) What acid acts as an electrolyte in a citrus battery?

Citric acid, and perhaps to a lesser degree, ascorbic acid.

5) If the two coins touch without the wet paper towel between them, what voltage will you expect to read from the external wire connecting them? Why? None, because the “shorter” path is directly between the touching metals. The arrangement has developed a “short.”

BATTERIES SAMPLE REPORT SHEET I. SALT BRIDGE BATTERY Voltage 2 V Current 0.4 mA II. CITRUS CELL Voltage 6 V Current 0.2 mA III. COIN BATTERY Coins penny and dime Voltage 0.306 V Current _______ mA Coins penny and nickel Voltage __0.570_____ V Current _______ mA Double battery Coins penny and dime Voltage 0.502 V Current _______ mA Double battery Coins penny and nickel Voltage 0.807 V Current mA III. QUESTIONS 1. Did the intensity of the color of the copper solution change with time?

Yes, it decreased in intensity, becoming less blue.

2. Does the copper electrode look like new copper has plated onto it?

Yes, it has a bright copper color to it.

3. If the copper solution touched a metal paper clip, what happened?

The metal paper clip was coated with copper.

4. Why do the different combinations of coins produce a different voltage?

Different metal combinations have different potentials (different pull on electrons).

5. Is the double battery twice the voltage of the single battery? Why? Yes, because it is two like batteries hooked together to add the voltages.

6. What happens if we use two coins of the same metal to make a coin battery?

Nothing, there is no pull on the electrons.

7. Did the change of concentration in the salt bridge battery result in a change of the voltage?

Yes. As the concentration is lowered, so is the voltage.

8. How are the results of halving the concentration in the salt bridge battery and doubling the number of coins in the coin battery related to each other?

Doubling the number of coins was analogous to adding the effects of two batteries together. Halving the concentration made fewer ions available to move. In both cases, the number of electrons available had an effect on the current produced.

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#41

Test for Vitamin C

Reagents 6M acetic acid [CH3COOH] 1% starch solution iodine solution 6M nitric acid [HNO3] copper metal strips or pieces

2 mL/pair 12 mL/pair 60 mL/pair 10 mL/pair 2-3 pieces/pair

Common Materials Vitamin C tablets, 250 mg citrus juices or powdered drink mixes apple juice or orange juice with vitamin C

1/pair 10 mL/pair 10 mL/pair

Special Equipment 10 mL volumetric pipette 250 mL volumetric flask

1/pair 1/pair

Laboratory Equipment filter paper 10-mL volumetric pipette 250-mL beaker 125-mL Erlenmeyer flask 100-mL graduated cylinder 100-mL beaker

ring stand, ring, and wire gauze laboratory burner centigram balance large test tube small squeeze bottle

Notes Iodine solution Dissolve 25 g of iodine in a solution of 50 g of potassium iodide in 200 mL of water.

Note: Red or purple drinks do not work well because the color interferes with the indicator and determination of the titration end point. Use either orange juice or apple juice. Both are not necessary. Tang™ has a very high concentration of vitamin C. If a juice is used that contains pulp, it is probably best if the pulp is filtered out before lab.

CAUTION: Store nitric acid away from copper strips and acetic acid.

Disposal Pour together the solutions containing the iodide ion and vitamin C. Flush all solutions down the drain with plenty of water.

VITAMIN C PRE-LAB QUESTIONS 1) Since vitamin C is soluble in water, any excess is excreted. Is it better to get all of your RDA at one meal or spread it out over the day? Explain your position.

No, because the vitamin is continually removed from the body, it is important that vitamin C is ingested daily, and several times during the day.

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2) Many soft drinks, snack foods, and energy bars have vitamin C in them. Why do you suppose that is so?

Most conscientious food manufacturers attempt to make foods that are as wholesome as possible and able to help supply the vitamins and nutrients the human body needs. Since vitamin C is continually necessary, easily added, inexpensive, and there is no risk of overdose, it is a logical choice of many food preparation companies.

3) If vitamin C is excreted continually, is it possible to get too much?

No. 4) If copper destroys the vitamin C content of foods, would it be prudent to expect that other metals may have the same effect? If so, which metals might you suspect?

Other metal that could be expected to destroy the vitamin C content would be metals in the same group as copper – silver & gold.

VITAMIN C SAMPLE REPORT SHEET 250 mg I. VITAMIN C CONCENTRATION = 1 mg/mL 250 mL II. TITRATE VITAMIN C a. Initial mass g b. Final mass g c. Mass of iodine added g mass of iodine 1 mL vitamin C d. Grams of iodine per mg vitamin C =  10 mL vitamin C mg vitamin C g iodine 1 mL vitamin C = = 0.01-0.1 g iodine/mg vitamin C  10 mL vitamin C mg vitamin C III. TITRATE JUICE* a. Type of juice b. Initial mass g c. Final mass g d. Mass of iodine added g *These answers depend on the type of juice used. mass of iodine (III d) 1 mg vitamin C e. Mass of vitamin C per mL juice =  10 mL juice mass iodine (II d) g iodine 1 mg vitamin C = = mg vitamin C/mL juice  10 mL juice g iodine f. Volume of juice to satisfy Recommended Daily Allowance mL juice 65 mg vitamin C = mL juice mg vitamin C (IV e) IV. TITRATE HEATED JUICE* a. Initial mass b. Final mass c. Mass of iodine added

g g g 112

mass of iodine (IV c) 1 mg vitamin C  10 mL juice mass iodine (II d) g iodine 1 mg vitamin C = = mg vitamin C/mL juice  10 mL juice g iodine *These answers depend on the type of juice used. d. Mass of vitamin C per mL juice =

V. QUESTIONS 1. What effect did heating have on the vitamin C in the juice?

Heating destroys the vitamin C. 2. Would a copper bowl be a wise choice to use when preparing orange juice?

No, it would not be a wise choice because copper destroys the vitamin C. 3. How might the results be affected if the vitamin C tablets had been sitting open for a year?

The vitamin C would be destroyed, so the results would be no vitamin C found.

4. Can you think of a way to determine if removing the pulp affects the vitamin C content of juice?

Two oranges could be prepared, one for which the pulp has been strained and removed and one prepared including the pulp. With the trials conducted holding all other factors constant, it could be determined of the presence of the pulp effects the amount of vitamin C.

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#42

Proteins

Reagents 6M nitric acid [HNO3] 6M sodium hydroxide [NaOH] calcium oxide [CaO] 1% urease solution 0.1% copper sulfate solution [CuSO4] phenol red 2% urea solution [H2NCONH2] conc. nitric acid [HNO3] ethyl alcohol [CH3CH2OH] 3M sodium hydroxide [NaOH] 1M lead(II) nitrate [Pb(NO3)2] 1M acetic acid [CH3COOH] sodium sulfide [Na2S] sodium carbonate [Na2CO3] 1M iron(III) chloride [FeCl3] red litmus pH paper

1 mL/pair 5 mL/pair 1 g/pair 4 mL/pair 1 mL/pair 3 drops/pair 5 mL/pair 5 mL/pair 3 mL/pair 15 mL/pair 3 mL/pair 3 mL/pair (for disposal only) (for disposal only) (for disposal only) 1 piece/pair 1 roll/pair

Common Materials Egg unflavored gelatin milk

1/every 4 students 1 pkg./4 students 12 mL/pair

Special Equipment nail clippers

1 pair/lab

Laboratory Equipment 10-mL graduated cylinder 3 250-mL beakers filter paper several large test tubes

ring stand, ring, and wire gauze laboratory burner watch glass dropper

Notes Remind students to fill pipette by bulb, not by mouth.

Caution: Store nitric acid away from acetic acid and ammonium sulfide.

Disposal Only the lead solution is toxic. The lead needs to be made into an insoluble solid and then buried in a toxic waste site. Collect all used solutions containing lead. Add a three-fold excess of sodium sulfide (Na2S). Let the solution set for an hour. Adjust the pH of the solution to neutral (pH 7 or more). Filter the precipitate. Discard the precipitate and filter in a container to be buried in a secure toxic waste landfill. The filtrate contains sulfide. Remove the sulfide by adding 1M iron(III) chloride (ferric chloride) in a three-fold excess. Neutralize the solution with

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sodium carbonate (Na2CO3). Filter the precipitate. Put the precipitate and filter in a container to be buried in a toxic waste landfill. The filtrate can be flushed with plenty of water.

The metal ion, copper, is not classified as toxic metal but in general it is better to not put any metal ions into the water system. Copper ions could be precipitated and buried in a landfill.

PROTEINS PRE-LAB QUESTIONS 1) Why is it important to keep the pH levels in the body constantly at the “optimum” pH for protein function?

Protein can be denatured by extreme pH levels and then cease to function properly.

2) Why are we very careful to keep lead and other metals out of the water systems following lab investigations?

Lead and other heavy metals damage the cellular tissues of living things. They also build up in the food chain so that all the agents released into the environment have the potential to return to us through the food chain.

3) What is the possible result of alcohol consumption on enzymes and proteins in the digestive system? In the circulatory system?

The immediate effect in the digestive system would be some cellular damage, but more than likely, the lining of the stomach and other paths would be somewhat protected by the mucus layers. Alcohol in the blood can reach all cells in the body and destroy them.

Cells that regenerate will eventually overcome the damage, but nerve, brain, and liver tissues do not regenerate and damage is therefore cumulative.

4) List as many body parts as you can that are made of protein, for example, hair. Which of these are most in danger of being denatured? Which would cause the most serious effect on the physiology of a human if it were denatured?

Nails, hair, DNA, RNA, enzymes, skin, and all solid cell structures are composed of proteins. Any proteins can be denatured. Those on the surface and exposed are at high risk to environmental agents, while the internal organs are at higher risk from ingested agents. The genetic molecules and enzymes would cause the most far reaching damage to human systems. But blindness resulting from contact with a heavy metal in the eyes could have far reaching effects also.

PROTEINS SAMPLE REPORT SHEET I. XANTHOPROTEIC TEST ON NAIL Positive ________ Negative II. HAIR 1. Test for ammonia Positive ________ Negative 2. Biuret Test Positive ________ Negative III. SALIVA 1. Biuret Test Positive ________ Negative IV. ENZYME 1. Biuret Test 115

_______ _______ _______ _______

Positive ________ Negative _______ 2. Test for enzyme action (ammonia) Positive ___ ____ Negative _______ V. PROTEIN TESTS ALBUMIN GELATIN CASEIN Biuret _________ _________ ________ Xanthoproteic _________ _________ ________ VI. DENATURING ALBUMIN GELATIN CASEIN Heat _________ __________ ________ Alcohol _________ _________ ________ Pb(NO3)2 _________ _________ ________ Acid _________ ____ ____ ________ VII.QUESTIONS 1. If a person has swallowed a heavy metal, would it be useful to drink milk and or egg white? Explain.

Both would be denautred by the metal ion so either would work.

2. Why is alcohol used as a disinfectant?

Alcohol kills bacteria by denaturing the protein of which they are made.

3. What causes cooked egg to be white?

The egg turns white when cooked because the proteins become denatured by the heat.

4. What caused the phenol red to change color in step 8?

Ammonia causes the phenol to turn red; the ammonia is caused by the decomposition of urease.

5. Urease is a specific enzyme. What is meant by that?

It only decomposes urea, a specific task, and is unreactive elsewhere.

6. What else might be added to the list of denaturing agents in Section V above?

Salts and bases.

7. Describe the differences in texture and appearance of the denatured solutions in steps 11, 12, 13, and 14.

In steps 11, 13, and 14, the proteins turn white, as if they have been cooked. In step 12, they clump together.

8. Enzymes, DNA, and other vital biochemical molecules are proteins. What would be the result of denaturing one of these proteins?

This would cause the function of that protein to not be performed causing a breakdown in body processes.

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#43

Enzymes

Reagents 3% hydrogen peroxide [H2O2] 3M sodium sulfide [Na2S] 1M lead(II) nitrate [Pb(NO3)2] 6M hydrochloric acid [HCl] 6M sodium hydroxide [NaOH] 1M iron(III) chloride [FeCl3] sodium carbonate pH paper

30 mL/pair <0.1 mL/pair 0.5 mL/pair <0.1 mL/pair <0.1 mL/pair (for disposal only) (for disposal only) 1 roll/pair

Common Materials potato, freshly peeled

30 g/pair

Special Equipment rubber tubing stopwatches 15 cm test tubes solid stopper to fit 15 cm test tubes 1-hole stopper with glass tube to fit test tube

2 ft/pair 1/pair 1/pair 1/pair 1/pair

Laboratory Equipment ring stand, ring, and wire gauze 2 250-mL beakers 1000-mL beaker 100-mL graduated cylinder 10-mL graduated cylinder

2 150-mL beaker 600-mL beaker utility clamp thermometer

Notes In procedure #4 of the student manual, a 100 mL beaker is sufficient.

It is important that the potato be cut into small pieces. Mashing the pieces with a mortar and pestle will also help.

The hydrogen peroxide solution needs to be fresh and no stronger than 3%.

CAUTION: Store acid away from sodium sulfide. Store lead nitrate away from hydrogen peroxide.

Disposal The lead ions are classified as toxic and sulfide ions are also not wanted in the water system. Pour the lead potato extract and the sulfide potato extract together. Add 10 drops of 3 M sodium sulfide. Let the solution set for an hour. Adjust the pH of the solution to neutral (pH 7 or more). Filter the precipitate. Discard the precipitate and filter in a container to be buried in a secure toxic waste landfill. The filtrate contains sulfide. Remove the sulfide by adding 1M iron(III) chloride (ferric chloride) in a three fold excess. Neutralize the solution with sodium carbonate (Na2CO3). Filter the precipitate. Put the precipitate and filter in a container to be buried in a toxic waste landfill. The filtrate can be flushed with plenty of water. The solid

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pieces of potato can be thrown in the trash. The solutions can be flushed down the drain with plenty of water.

ENZYMES PRE-LAB QUESTIONS 1) Why should it be easier to get the extract from the potato if it is in smaller pieces?

The enzyme is extracted from the cells that are open and exposed on the surface. The more cuts there are in the potato, the more enzyme can be removed. Crushing the potato would be most efficient.

2) If cyanide and sulfide ions will inhibit the function of catalase because it contains iron, what other important protein function might be disrupted by the presence of these ions?

Any protein could possibly be effected, but the enzymes that function to digest, such as amylase, would be expected to be most vulnerable.

3) Our bodies have elaborate systems for maintaining optimum temperature and pH levels.

Why? Proteins are pH and temperature specific. If the environment in the cell, or in the body as a whole, becomes too different than the optimal ranges, the proteins cease to function.

4) Why is it important to use the same amount of catalase and hydrogen peroxide in each trial?

Concentration effects the reaction and if the beginning amounts are different, no valid conclusions can be drawn.

ENZYMES SAMPLE REPORT SHEET *Student data will vary.

I. Catalase Final water level minus initial water level O2 produced II.

III.

=

mL mL mL

Catalase + Sulfide Final water level minus initial water level O2 produced

=

mL mL mL

Catalase + Lead Final water level minus initial water level O2 produced

=

mL mL mL

=

mL mL mL

-

mL mL

IV. Catalase at 37_C Final water level minus initial water level O2 produced V.

Catalase at 70_C Final water level minus initial water level

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O2 produced VI. Catalase in acid Final water level minus initial water level O2 produced VII. Catalase in base Final water level minus initial water level O2 produced

=

mL

=

mL mL mL

=

mL mL mL

VIII. QUESTIONS 1. What effect did the sulfide have?

It reduced the action of the enzyme.

2. What effect did the lead have?

It prevents the enzyme from working by denaturing the protein.

3. At which temperature does catalase work best: 25ºC, 37ºC, or 70ºC? 37o - body temperature 4. At which temperature does catalase work the least well?

25o 5. How does the acidity affect the enzyme action?

Too much acid denatures the enzyme and prevents it from working efficiently.

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#44

Forensic Chemistry

Reagents luminol 3% hydrogen peroxide [H2O2] ethyl alcohol [CH3CH2OH] phenolphthalein sodium hydroxide [NaOH] zinc powder copper oxide, fine powder [CuO]

1 g/pair 31 mL/pair 2 drops/pair 0.25 g/pair 2.5 g/pair 2.5 g/pair 1 g/pair

Common Materials clear tape cotton swabs or Q-tips modeling clay or Handi-Tak® or soft putty substance

2 strips/pair 2/pair small ball/pair

Special Equipment ultraviolet light "dark box"

1 or 2 /class 1 or 2 /class

Laboratory Equipment watch glass eye dropper 150-mL beaker Notes A "dark box" can be made by cutting 2 holes in a cardboard box. One hole will allow the UV light into the box, the other will let the student see the sample.

A visible dried bloodstain on a Styrofoam meat tray works well for the phenolphthalein test.

For the luminol test, it is suggested that styrofoam plates wiped with blood and then rinsed off will work well. While the luminol solution is wet, it will luminesce when exposed to the ultraviolet light, producing a light blue glow. The positive test for blood is a yellow or yellowgreen glow. It will continue to luminesce after the solution is dry.

Disposal The zinc solution needs to be evaporated in a metal pan because metal ions should not be put into the water system. The dried zinc oxide should be put into a container to be buried in a toxic waste landfill. The other solutions can be flushed down the drain with plenty of water.

FORENSIC CHEMISTRY PRE-LAB QUESTIONS 1) When might a latent fingerprint be permanent?

If it is made in paint, wet cement, or some other substance that dries and hardens.

2) If iron is the catalyst for tests that detect the presence of blood, what compounds might give a false positive test?

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Anything that contains iron may be capable of giving a false test. Common materials might include rust, certain iron containing foods, dishes on which meat has been prepared or served, etc.

3) Will the tests for blood work only for human blood, or will they detect any animal blood?

Tests for blood that function to pick up iron or hemoglobin will be positive for any animal blood.

4) What characteristics must be possessed by the powder with which fingerprints will be dusted?

The substance used as finger print powder must be very fine so that it will show the most minute detail, it should be able to be picked up for preservation of the print, it must not degrade with time, and it must be able to be carried and used easily.

FORENSIC CHEMISTRY REPORT SHEET I. PART A: FINGERPRINT DETECTION Latent Print Plastic Print II. PART B: BLOOD DETECTION Luminol Test SAMPLE

POSITIVE

NEGATIVE

_____________

_________

__________

Kastle-Meyer Phenolphthalein Test SAMPLE

POSITIVE

NEGATIVE

_____________

_________

___________

III. QUESTIONS 1. Why does a latent print show up better on a smooth surface?

More of the finger makes contact with the surface so more oil lines are left on the surface. Not all of the finger touches against a rough surface and part of the fingerprint is lost.

2. In the luminol test the iron in blood is acting as a catalyst. Other catalysts also work such as iodine, permanganate, or hypochlorite (bleach). What happens if the suspected area had one of the other catalyst spilled on it.

It would luminesce as if it had had blood on it.

3. Why is a dark room better for observing the luminol test?

Normal light will overpower the glow produced by luminol.

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#45

Protein in Milk and Other Foods

Reagents glacial acetic acid [HC2H3O2] 1:1 ethyl alcohol-ethyl ether conc. nitric acid [HNO3] conc. ammonia [NH3] 0.1% copper sulfate [CuSO4] 6 M sodium hydroxide [NaOH]

5 mL/pair 40 mL/pair 1 mL/pair 1 mL/pair 1 mL/pair 3 mL/pair

Common Materials milk rice white bread butter beans Swiss cheese white meat (turkey or chicken) marshmallow radish table sugar sliced almonds

30 g/pair 1 g/pair 1 g/pair 1 g/pair 1 g/pair 1 g/pair 1 g/pair 1 g/pair 1 g/pair 1 g/pair

Special Equipment cheesecloth, 12 in. square rubber bands (to fit 500 mL beaker) 500 mL beaker filter paper

1/pair 1/pair 1/pair 1 piece/pair

Laboratory Equipment 600-mL beaker 100-mL beaker 125-mL Erlenmeyer flask watch glass thermometer

2 filter flasks ring stand, ring, and wire gauze laboratory burner eye droppers

Notes 1:1 by volume ethyl alcohol (CH3CH2OH) - ethyl ether (C2H5OC2H5) Unbleached muslin works well in place of cheesecloth.

Caution: Store nitric acid away from acetic acid, alcohol, and flammable liquids (ether).

Disposal Although normally it is not good procedure to put metal ions such as copper into the water system, in this investigation, the amount is so small that it provides a very small danger to the environment. Two or three drops of a 1% solution diluted with plenty of water quickly becomes a concentration that is small enough to be negligible. The milk protein and foods can be thrown into the wastebasket. The whey and other solutions can be washed down the sink with plenty of water. The acid should be flushed with an excess of water.

122

PROTEIN IN MILK AND OTHER FOODS PRE-LAB QUESTIONS 1) Enzymes, DNA, and other vital biochemical molecules are proteins. What would be the result of eating too little protein?

Too little protein can cause the functions of enzymes, DNA, and other biochemical molecules to slow or cease. This can cause cellular breakdown resulting in illness.

2) What might you expect to be the visible result if you accidentally splattered sodium hydroxide and copper sulfate on your skin at the same time?

There would be a blue gel substance on the flesh. (But if there was left over sodium hydroxide, it would eventually burn.)

3) Why will we test the inside of the butter bean and the white portion of the bread and radish?

The color from the crust and skin portions might make the test results difficult to determine.

PROTEIN IN MILK AND OTHER FOODS SAMPLE REPORT SHEET I. CASEIN IN MILK Mass of watch glass and casein minus Mass of watch glass

g g

_

_____________ Mass of casein

g

mass of casein  100% mass of milk g casein  100% = ___10-15____ % = 30 g milk Brand Name of Milk _____________________ Percentage of casein in the product label _____________________

The percentage of casein in the milk sample is

II. PROTEIN TESTS Indicate the color of the sample after adding: Sample HNO3 Water & CuSO4 & (Conc.)

NH4OH NaOH Casein yellow orange purple Whey no color no color no color White Meat yellow orange purple Radish no color no color no color marshno color no color no color mallow Rice yellow orange purple Butter Beans yellow orange purple Bread yellow orange purple Sliced yellow orange purple Almond

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Conclusion protein no protein protein no protein no protein protein protein protein protein

Swiss Cheese Sugar

yellow

orange

purple

protein

no color

no color

no color

no protein

III. QUESTIONS 1. Did you remove all of the protein from the whey? Use your results to support your answer.

A large portion, but not all. Some small proteins may not have been clumped together and may have escaped through the cheese cloth. Results will vary.

2. List as many foods as possible which contain protein.

Student response here - nearly anything is appropriate except clear liquids.

3. Were the two protein tests consistent in their results? Explain.

Yes, except where the phase of the sample or the color of the sample interfered with a specific test, if the sample contained protein, both tests gave a positive result.

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#46

Production of Carbon Dioxide

Reagents sodium bicarbonate [NaHCO3] 0.5 M acetic acid [CH3COOH] 0.5 M hydrochloric acid [HCl]

0.68 g/pair 42 mL/pair 14 mL/pair

Common Materials Alka-Seltzer™ Rolaids™ baking powder baking soda

0.34 g/pair 0.34 g/pair 0.34 g/pair 0.34 g/pair

Special Equipment Rubber tubing (at least 24") that will fit snugly on glass tubing

1/pair

Laboratory Equipment 1000-mL beaker 600-mL beaker 100-mL graduated cylinder 10-mL graduated cylinder 15-cm test tube

ring stand, ring, and wire gauze utility clamp 1-hole stopper to fit 100-mL graduated cylinder glass tube for stopper

Notes Be sure tubing assembly in test tube is tightly sealed. TumsTM or another carbonate antacid should also work well for this investigation.

Disposal Flush the solutions down the drain with plenty of water as they are not toxic.

PRODUCTION OF CARBON DIOXIDE PRE-LAB QUESTIONS 1) If a recipe calls for lemon juice, would you expect it to require baking soda or baking powder?

Baking soda because baking powder contains its own acid and requires no additional acid in the recipe.

2) If a recipe calls for both baking soda and baking powder, what is a likely reason?

The baking soda will cause the recipetol rise some before it is baked, and then it will rise again during the baking process when the heat activates the baking powder.

3) A certain recipe for bran muffins will keep for a week when refrigerated. As soon as the batter is placed in the oven it rises. What ingredient provides the carbon dioxide gas?

Baking powder because it waits until it is heated to react and rise.

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4) Carbon dioxide gas is soluble in water. How will this affect the measurement of gas produced by each reaction during the investigation?

If the water is not saturated with carbon dioxide gas before the experiment begins, some of the gas produced will be retained in the water and not become a part of the measured volume of gas collected.

PRODUCTION OF CARBON DIOXIDE SAMPLE REPORT SHEET * All procedures will produce between 30 and 65 mL of CO2.

VI. QUESTIONS 1. What would happen if you swallowed a whole Alka-Seltzer tablet?

There would be a large production of CO2 gas.

2. What occurs if dough is left overnight before it is cooked?

The dough would rise, then fall as the bubbles burst. There would be little reaction left to allow it to rise again while cooking.

3. What is the difference between sodium bicarbonate from the laboratory and baking soda from the kitchen?

Essentially they are the same and would produce the same amount of CO2.

4. Would you ever use calcium carbonate (CaCO3) to substitute for sodium bicarbonate (NaHCO3) in cooking? Why or why not?

No, more CaCO3 would need to be used, and it would leave a bad taste because it is like adding lime to your food. Lime is bitter.

5. If the % NaHCO3 is greater than 100%, which of the following is a logical explanation? a. There is more than 100% NaHCO3 in the material. b. Some of the CO2 gas in the sodium bicarbonate determination was not captured. c. Extra gas was produced by a second reaction.

6. What happens if you are slow getting the stopper in the test tube or the stopper does not make a good seal?

This would result in inaccurate readings because much of the gas produced would escape.

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#47

Food Colors By Paper Chromatography

Reagents butanol [CH2(CH3)3OH] ethanol [CH3CH2OH] 2M ammonia [NH3] conc. acetic acid [CH3COOH] 1% ammonia solution [NH3] pH paper

60 mL/pair 20 mL/pair 20 mL/pair 1 mL/pair 5 mL/pair 1 strip/pair for each color of candy

Common Materials FD&C food colors: B1, B2, Y5, Y6, R3, R40, G grocery store food colors (R, G, B, Y) colored candies (M&M's™ or Skittles™)

10 mL/lab 1 bottle each/lab 8 pieces of one or two colors/pair 2/pair 1 2 in. square/pair 2 cm × 5 cm/pair/color (13 cm/pair/color)

paper clips plastic wrap or sealing wax (or ParafilmTM) undyed wool fabric or yarn (or 100% cotton yarn) Special Equipment chromatography paper

3/pair

Laboratory Equipment 10-mL graduated cylinder 150-mL beaker

watchglass

Notes Cut chromatography paper to a size which will fit the beakers with parafilm on top. Be sure the yard or fabric is a natural fiber - wool works best. Synthetics do not work well. "Undye" the yarn by boiling in ammonium hydroxide for 10 minutes.

The FD&C dyes must be purchased by name not by FD&C number. Double check the CI number to be sure you are getting the correct dye. The company names are abbreviated. See below for full names and see the list of Equipment Suppliers at the beginning of this manual for addresses and numbers. The following is a partial list.

FD&C Name CI # AL ATO FIS HLD NCI SPE TCI VWR WAK WAR B #1 Erioglaucine 42090 x x x x x x B #2 Indigo Carmine 73015 x x x x x x G #3 Fast Green 42053 x x x x x R #3 Erythrosine 45430 x x x x x x R #40 Allura Red AC 16035 x x x x x Y #5 Tartrazine 19140 x x x x x x x x Y #6 Sunset Yellow 15985 x x AL

Aldrich Chemical Company, Inc.

SPE

127

Spectrum Chemical Mfg. Corp.

ATO Atomeric Chemetals Corp.

TCI TCI America, Inc. FIS Fischer Scientific VWR VWR Scientific HLD Hilton Davis Company WAK Wako Chemicals USA, Inc..

NCI NCI Chemical Carc. Ref. Stand. Rep. WAR Warner Jenkinson Company Disposal The solids and yarn can be thrown into the trash. The color cannot be completely removed from the yarn. The solutions can be flushed down the drain with plenty of water.

FOOD COLORS BY PAPER CHROMATOGRAPHY PRE-LAB QUESTIONS 1) Why do you suppose the numbers of the FD&C approved food dyes are not consecutive?

The numbers refer to all the dyes that have been isolated in nature or produced in laboratories that have been presented to the Federal Food, Drug, and Cosmetic Commission for approval. Most are not approved.

2) Would you expect stain removers to be acidic or basic?

Basic environments remove dyes, and extremely basic environments can denature them completely. A high pH also raises the outer layer of natural, animal fibers such as wool and allows the dyes to wash out.

3) Most synthetic fibers must be dyed during their production process. Why do natural fibers take up dye, and synthetic ones do not?

Synthetic fibers are solid and do not have the structure to hold the dye molecules.

4) After cloth is dyed at home, it is often boiled in a vinegar solution. Why? A low pH, or acidic environment, helps to “set” the dye.

FOOD COLORS BY PAPER CHROMATOGRAPHY SAMPLE REPORT SHEET I. DYES

Distance traveled *

Rƒ =

Distance dye traveled Distance solvent traveled

(cm) SOLVENT _________ B1 _________ _0.474_______ B2 _________ _0.400_______ Y5 _________ _0.346_______ Y6 _________ _0.753_______ R3 _________ _0.827_______ R40 _________ _0.548_______ G _________ (blue & yellow)_ *Distances will vary. Rf values for the pure dyes are generally larger than for dyes isolated from the candies.

II. STORE FOOD COLORS

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Solvent Components of: red yellow green blue

Distance traveled _________

Rƒ

Identity B1, B2, Y5, Y6, R3, R40, G

_________ _________ _________ _________ _________ _________ _________ _________

0.556 0.722 0.329 0.575 0.329 0.301 0.575 0.301

__R40______ __R3_______ __Y5_______ __R40______ __Y5_______ __B1_______ __R40______ __B1_______

III. CANDY DYES Skittles Solvent Candy Color #1 #2 #3

DISTANCE TRAVELED _________

#4

_________

#5

_________

Rƒ

IDENTITY B1, B2, etc.

0.322 0.508 0.339 0.525 0.322 0.373 0.458 0.508

_________ _________ _________

__R40 & Y5__ __Y5________ __Y6________ __B1 & Y5____ __R40 & B2___

IV. QUESTIONS 1. What is the shade difference between blue #1 and blue #2?

B1 is not as bright as B2 between red #3 and red #40?

R3 is pink in comparison to R40 between yellow #5 and yellow #6?

Y6 is orange in comparison to Y5 2. Are the food colors pure, that is, made of one dye?

No, they are combinations of the FD&C dyes.

3. Are any of the candy dyes pure?

Yes, two of the candy dyes are pure - #2 and #3. Answers will vary depending on the candy used.

4. Why does a colored acidic food make a worse stain on wool than a colored basic food?

Basic solution cannot hold dye, so the dye will not stain.

#48

Detection of Alcohol

Reagents

129

potassium dichromate [K2Cr2O7] conc. sulfuric acid [H2SO4] ethyl alcohol [CH3CH2OH] acetone [CH3COCH3] sodium sulfide [Na2S] or thioacetamide [CH3CSNH] 3M sodium hydroxide [NaOH] 1M ferric chloride [FeCl3] (or iron(III) chloride) sodium carbonate [Na2CO3]

30 g/pair 30 mL/pair 150 mL/pair 5 mL/pair 47 g/pair 45 g/pair (for disposal only) (for disposal only) (for disposal only)

Special Equipment various lengths of glass tubing (polished ends if possible) stop watches

4 pieces/pair* 1/pair

Laboratory Equipment various lengths of glass tubing & rubber hose 500-mL Erlenmeyer flask 2-hole stopper to fit 500-mL flask funnel

filter paper 125-mL Erlenmeyer flask 600-mL beaker

Notes * Refer to the figure in the student manual for glass needs.

CAUTION: Store acid away from sodium sulfide.

Disposal Dichromate and chromium ion are both classified as toxic chemicals. The disposal uses sulfide because the sulfides are very insoluble in water. Dichromate is also a very powerful oxidizing agent. The chromium should be precipitated with a three-fold excess of sodium sulfide (Na2S) or thioacetamide (CH3CSNH2). We began with 30 grams of potassium dichromate (K2Cr2O7), which represents 0.20 moles of Cr. That means complete precipitation would be ensured with 47 grams of Na2S in solution or 45 grams of thioacetamide. Let the solution set for an hour. Check the pH. If it is acidic, adjust it to neutral with 3 M sodium hydroxide. Both of these steps insure complete precipitation. Filter and discard the precipitate and filter paper in the chromium waste container.

The sulfide is hazardous and must be removed. To the filtrate, add 1 M ferric chloride (iron(III) chloride, FeCl3) until no more solid forms. Test the pH. If the solution is acidic add sodium carbonate (Na2CO3) until it is neutral. Filter the solution. Place the precipitate and filter paper in the waste container marked iron (III) sulfide to be buried in a hazardous waste landfill.

Flush the remaining solution down the drain with plenty of water.

DETECTION OF ALCOHOL PRE-LAB QUESTIONS 1) How could blowing through an alcohol solution simulate the alcohol content of a person’s breath?

As the breath moves through the alcohol solution, it draws alcohol with it because the alcohol is quite volatile. As air leaves the flask and fresh breath enters, the alcohol evaporates to reach a new equilibrium.

130

2) How would you expect the varied concentrations of alcohol to effect the time required for the reduction of the dichromate ion?

A lower concentration of alcohol in the flask will vaporize slower, resulting in a longer time required to build up the concentration in the dichromate solution to give a positive test.

3) Why is a chemically based test not as reliable as an electronically based alcohol test?

There are many factors that effect the accuracy of a chemical test including temperature, the purity of the water in which the dichromate ion solution is made, how fast the individual blows into the container, etc. The ability of the test administrator to discriminate between colors in comparisons is another place in which error can be introduced to the procedure.

4) Discuss ways in which the results of a test of the sort used in this investigation might vary between individuals.

There will be several appropriate responses, but the most obvious is that if one individual blows harder than another, the time required to change the color of the solution would be shorter, even if they had the same actual alcohol content on their breath.

DETECTION OF ALCOHOL SAMPLE REPORT SHEET TIME TO COLOR I. TEST SOLUTION OF CONTROL 50 Alcohol 50 Water _________20 sec______ min 40 Alcohol 60 Water _________45 sec______ min 30 Alcohol 70 Water _________<1_________ min 20 Alcohol 80 Water __________2 ________ min 10 Alcohol 90 Water

__________7-8________

min

II. QUESTIONS 1. What is the color of the control sample?

The control sample changes from bright orange to deep orange.

2. What happens to the temperature of the dichromate solution? Why? It gets warm because the chemical reaction is exothermic.

3. Would this test be convenient for a police officer to carry with him if it were calibrated to determine percent alcohol?

No, the disposal is too difficult. It is also a liquid which could spill.

4. Why was proper disposal of the dichromate and chromium solution important?

These are both toxic chemicals and dichromate is a very strong oxidizing agent.

5. Why would a chromium solid that is more soluble be less satisfactory as a disposal method?

The precipitate must be filterable so the chromium must be less soluble.

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#49

Aspirin Preparation and Purity Test

Reagents salicylic acid [C7H6O3] acetic anhydride [(CH3CO)2O] conc. sulfuric acid [H2SO4] 95% ethyl alcohol [CH3CH2OH] eluting solvent (66:33:1 hexane, ethyl acetate, glacial acetic acid)

11 g/pair 12 mL/pair 1 mL/pair 35 mL/pair 30 mL/pair

Common Materials fresh commercial aspirin aged aspirin ice paper towels, plastic wrap, or foil (to fit over beakers) three paper clips Special Equipment TLC chromatography plates ultraviolet light filter paper for Buchner funnel hair dryer (for this eluting solvent probably not needed) Laboratory Equipment 3 paper clips Buchner funnel 2 600-mL beakers 125- or 250-mL Erlenmeyer flask 150-mL beaker 250-mL beaker

1/2 tablet/pair 1/2 tablet/pair 300 mL/pair 1 square/pair 3/pair 2/pair 1/lab 3/pair

filter paper, large size for drying wash bottle of distilled water balance stirring rod 4 small test tubes mortar and pestle

Notes Age aspirin several days before needed.

The Whatman Flexible-backed TLC plates with polyester backing and fluorescent indicator can be purchased in packages of 25 from Fisher Scientific. They come 20 × 20 cm and can be cut to 6.7 × 20 cm for this investigation.

Another brand of TLC plates is POLYGRAM® SIL G/UV254, 40 × 80 mm silica gel on plastic plates from Bodman Chemical, Aston, PA (phone 1-800-241-8774). They come fifty plates per box. The cost is $.50 per plate but can be cut in half for economy.

CAUTION: Store acetic anhydride away from alcohols.

Disposal The eluting solution is toxic enough that it should not be placed into the water system.

However, it is volatile and can be evaporated in a hood. The prepared aspirin and the solutions used to prepare the aspirin are not toxic. The spotting solutions can be poured down the drain with plenty of water.

132

ASPIRIN PREPARATION AND PURITY TEST PRE-LAB QUESTIONS 1) Why should you discard aspirin once its expiration date has been reached or if it has been stored in a moist or warm environment for very long?

With age and especially in moist or warm environments the aspirin begins to degrade and the acid content rises.

2) If a solid impurity fell in to the beaker during the initial formation of the aspirin, in which step would it be removed?

Once the aspirin has been re-dissolved, any solid impurities can be seen and then filtered out in step #10.

3) If two separate samples of the same compound are analyzed by chromatography, why should the Rf values match even if the solvent fronts move different distances on the two chromatograms?

The Rf value is a ratio. The ratio of the distance the solvent traveled to the distance the compound traveled will remain constant if all other factors are the same.

4) Why is ethanol used to dissolve and spot the samples on the chromatogram?

The eluting fluid is made primarily of non–polar substances. It would not be wise to put water, a very polar molecule, on the chromatography surface. Ethanol also evaporates quickly and allows the spot to be made small.

ASPIRIN PREAPRATION AND PURITY TEST SAMPLE REPORT SHEET I. Mass of paper + salicylic acid g _ minus mass of paper ______________________g _____________________________________________________ Mass of salicylic acid ________9.06 ___________g II. Mass of 100 mL beaker + aspirin _______________________g minus mass of 100 mL beaker _______________________g ______________________________________________________ Mass of aspirin _____7.3_______________g

experimental mass of synthetic product (aspirin)  100% Theoretical mass of synthetic product (aspirin) mass of aspirin (II)  100% = 62% = 180.15 g aspirin/mol mass of salicylic acid (I)  138.12 g sal acid / mol Present aspirin to instructor for initial _______________

III. % Yield =

IV. Describe the appearance of the salicylic acid spot.

The salicylic acid spot is blue/brown under the UV light.

Comment on the presence or absence of salicylic acid in the following : 1. Fresh commercial aspirin no salicylic acid 2. Aged commercial aspirin some salicylic acid 133

3. The aspirin prepared in this investigation

Salicylic Acid Distance traveled

Aspirin Distance traveled

some salicylic acid

Fresh

Aged

Today's Preparation

Spot Solvent Rf

___2.4 cm__ ___4.8 cm__ __ 0.50___

___2.3 cm__ ___4.8 cm__ ___ 0.48___

___2.5 cm__ ___4.8 cm__ __ 0.42___

Spot Solvent Rf

___2.5 cm__ ___4.8 cm__ ___ 0.42___

___1.6 cm__ ___4.8 cm__ ___ 0.33 __

___1.4 cm__ ___4.8 cm__ ___ 0.21 __

V. QUESTIONS 1. Why do you not get 100% yield?

If there was still water in the crystals it would account for a greater than 100% yield. If the yield is less that 100% it will be because aspirin is water-soluble and sticks to the surfaces of containers.

1. Is excess water still in the aspirin when you weighed it?

Yes. 3. Could another compound that was also insoluble in cold water be separated from aspirin by re-crystallization?

Probably not, the two compounds would precipitate together.

4. Does your product aspirin smell of vinegar?

Yes. 5. Why does aspirin that has a sour smell have a greater chance of causing an upset stomach?

The sour smelling aspirin contains acetic acid, and perhaps some salicylic acid that upsets the stomach.

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#50

Fats and Iodine Number

Reagents glycerol potassium hydrogen sulfate [KHSO4] hexane [C6H14] 5% iodine solution concentrated sulfuric acid [H2SO4] 50% sodium thiosulfate 3M sulfuric acid [H2SO4] 6M sodium hydroxide solution [NaOH] pH test strips Common Materials vegetable oil solid margarine liquid margarine peanut or olive oil potato chip potato white meat (chicken or turkey) radish Fritos™ Laboratory Equipment small test tubes stirring rod

0.25 mL/pair 0.5 g/pair 30 mL/pair 0.25 mL/pair 3 mL (for disposal only) (for disposal only) (for disposal only) (for disposal only) 4 mL/pair 4 mL/pair 4 mL/pair 3 mL/pair 3 small pieces/pair 3 small pieces/pair 3 small pieces/pair 3 small pieces/pair 3 small pieces/pair watch glass

Disposal The hexane and I2 used in this investigation are classified as toxic substances. Iodine is a strong oxidizing agent. The hexane may be evaporated in a hood. The iodine has to be reduced with a two-fold excess of 50% sodium thiosulfate solution. Use 3 M sulfuric acid (H2SO4) to make the pH 2-3. Set for an hour. Neutralize the solution with sodium hydroxide (NaOH) or sulfuric acid. Flush down the drain with plenty of water. Place all foods in the trash.

FATS AND IODINE NUMBER PRE-LAB QUESTIONS 1) Industry can adjust whether a material is solid or liquid by adding or deleting double bonds. Which is done to make a material more solid?

Double bonds are added to make the liquids solid.

2) Should a person be worried about the type of fat found in solid vegetable margarine?

Explain.

Yes, because the amount of fat per serving can vary greatly depending upon the type of oil used to produce the margarine. It is also advisable to monitor how much trans fat is in the margarine used.

135

3) Many Mediterranean cultures dip their bread in olive oil mixed with roasted garlic and other spices. Compare the health value of this practice to the traditional bread and butter (or margarine) we use in the US. Oils are generally more unsaturated than solids and are therefore healthier. Olive oil, especially, is one of the more healthy choices of the oils.

4) Would you expect the iodine number to be greater for squeeze margarine or for tub margarine. Why? The iodine number for tub margarine would be expected to be lower because it is more solid, meaning that it is more saturated.

FATS AND IODINE NUMBER SAMPLE REPORT SHEET I. FAT TEST* 1. Record your observations of each test tube. #1 #2 #3 2. Record your observations of each food tested for fat. potato chip raw potato radish meat Fritos™ *These are students’ personal observations.

3.

Liquid Margarine Butter Solid Margarine Vegetable Oil Peanut Oil

# drops of I2 12 19 15 4 4

I. QUESTIONS 1. Rank the following in terms of increasing order of dietary health: butter, solid margarine, or liquid margarine. butter, solid margarine, liquid margarine 2. Why are so many products advertised as being fried in peanut oil or olive oil?

These are healthier than other oils or fats.

3. Where have you smelled burned fat before?

At fast food hamburger places and grilling in the backyard.

4. Would you expect a positive test for fat on a piece of celery?

No.

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#51

Carbohydrates

Reagents Benedict's reagent * Barfoed's reagent * 1% glucose solution 1% fructose solution 1% sucrose solution 1% cellulose solution 1% starch solution iodine solution pH 7 buffer solution 50% sodium thiosulfate solution [Na2S2O3] 3M sulfuric acid [H2SO4] 6M sodium hydroxide solution [NaOH]

4 mL/pair 8 mL/pair 20 mL/pair 20 mL/pair 20 mL/pair 20 mL/pair 20 mL/pair 1 mL/pair 10 mL/pair (for disposal only) (for disposal only) (for disposal only)

Common Materials 1% table sugar, solution 1% honey, solution 1% corn syrup, solution 1% powdered sugar, solution

10 mL/pair 10 mL/pair 10 mL/pair 10 mL/pair

Laboratory Equipment 5 large test tubes

50-mL beaker

Notes Benedict's Reagent Dissolve 4.3 g of finely pulverized copper sulfate in 25 mL of hot water, cool and dilute to 40 mL with water. Dissolve separately 43 g of sodium citrate and 25 g of anhydrous sodium carbonate in 150 mL of water. Heat to effect solution; cool, then add the copper sulfate solution. Dilute to 250 mL. Keep in a cork-stoppered bottle.

"Clinitest" tablets used by diabetics to check sugar levels are a solid form of Barfoed's Reagent.

Barfoed's Reagent Dissolve 66 g of cupric acetate and 10 mL glacial acetic acid in water and dilute to 1 liter.

Iodine solution Dissolve 25 g of iodine in a solution of 50 g of potassium iodide in 200 mL of water. pH 7 Buffer Add 9.1 g KH2PO4 and 18.9 g Na2HPO4 to enough water to make a liter of solution.

Disposal Copper is a metal ion which should not be allowed into the water system. Iodine is a strong oxidizing agent. The iodine has to be reduced with a two-fold excess of 50% sodium thiosulfate solution. Use 3 M sulfuric acid (H2SO4) to make the pH 2-3. Set for an hour. Neutralize the solution with sodium hydroxide (NaOH) or sulfuric acid. Flush down the drain with plenty of water.

The copper solutions should be evaporated to a solid which should be placed in a container to be buried in a secure landfill. All foods may be placed into the trash.

CARBOHYDRATES PRE-LAB QUESTIONS

137

1) Which should provide energy more quickly to the body: table sugar or an apple that contains fructose? Explain your reasoning.

The apple with the fructose would provide the quickest energy because fructose is a simple sugar and table sugar is a disaccharide.

2) Which should provide energy more quickly to the body: sucrose or starch? Explain your reasoning.

Sucrose would provide energy to the body more quickly because it is a disaccharide and is more readily digested than the complex sugar in the starch carbohydrate.

3) What must be the active ingredient in the product Lactaid™, an aid for the digestion of dairy products containing lactose?

LactaidTM must contain lactase.

4) Give the results you would expect to see from each of the three chemical analysis tests, Benedict’s, Barfoed’s, and Iodine, on the simple sugar maltose.

Benedict’s test would be positive within a short amount of time, possibly immediately.

Barfoed’s test may eventually turn positive, and the iodine test will be negative.

5) Which of the four carbohydrate solutions in step 7 would you predict to have a negative result to Benedict’s test.

Starch and cellulose should have a negative result to a Benedict’s test because they do not contain mono- or disaccharides.

CARBOHYDRATES SAMPLE REPORT SHEET I. TESTS Benedict's (Color) Sample Glucose ______ Fructose ______ Sucrose ______ Starch ______ Cellulose ______ Table sugar ______ Corn syrup ______ Honey ______ Powdered sugar ______ Sucrose in saliva ______ Sugar in saliva ______ Starch in saliva ______ Cellulose in saliva ______ II. ENZYME ACTION Sample:

Starch

Barfoed's (Color) (Time)

Iodine (Color)

Conclusion (Type of Carbohydrate)

_____ _____ _____ _____ _____ _____ _____ _____ _____ _____ _____ _____ _____

______ ______ ______ ______ ______ ______ ______ ______ ______ ______ ______ ______ ______

Monosaccharide Monosaccharide Disaccharide Disaccharide Disaccharide Disaccharide Disaccharide Disaccharide Monosaccharide Monosaccharide Monosaccharide Monosaccharide Disaccharide

Sucrose

_____ _____ _____ _____ _____ _____ _____ _____ _____ _____ _____ _____ _____ Table Sugar

138

Cellulose

Color:

______

______

______

______

III. QUESTIONS 1. From your data, which of the first 8 substances would you expect to provide the quickest energy to the body? Why? Glucose and fructose would be expected to provide the quickest energy because they are monosaccharides.

2. Why does the water from the cooking of vegetables taste sweet?

The heat breaks the disaccharides into monosaccharides.

3. In the procedure, heat was used to break the bond between monosaccharide units. In what other ways could the bonds be broken?

Enzymes in the digestive system are efficient in breaking the bonds between monosaccharides. Heat will also break down more complex sugars in to simple ones.

4. Give an explanation for the results in the enzyme action test.

The starch and table sugars are slowly broken down into monosaccharides by the salivary amylase. Amylase does not act on cellulose.

#52

Analysis of Fertilizer

Reagents 5% barium chloride [BaCl2] or barium nitrate [Ba(NO3)2] 139

2 mL/pair

sodium sulfate [Na2SO4] potassium chloride [KCl] ammonium phosphate [(NH4)3PO4] sodium nitrate [NaNO3] 2M ammonium hydroxide [NH4OH] 0.2M iron(II) sulfate solution [FeSO4] conc. sulfuric acid [H2SO4] 10% sodium hydroxide [NaOH] 3M hydrochloric acid [HCl] 6M hydrochloric acid [HCl] concentrated hydrochloric acid [HCl] 6M nitric acid[HNO3] 75% isopropyl alcohol 10% magnesium sulfate [MgSO4.7H2O] red litmus

0.1 g/pair 0.1 g/pair 2.0 g/pair 0.1 g/pair 150 mL/pair 2 mL/pair 4 mL/pair 5 mL/pair 25 mL/pair 30 mL/pair 5 mL/pair 15 mL/pair 20 mL/pair 45 mL/pair 4 pieces/pair

Common Materials inorganic fertilizer sample(s)

10 g/pair

Special Equipment platinum or nichrome wire cobalt glass

1/pair 1/pair

Laboratory Equipment 2 50-mL beakers filter paper Buchner funnel evaporating dish

250-mL beaker 100-mL graduated cylinder 10-mL graduated cylinder balance

Notes In the nitrate test, placing the sulfuric acid under the solution by use of a pipette or dropper may work better than pouring down the side. The brown ring may take 30-45 seconds to form.

The iron(II) sulfate should be prepared daily.

CAUTION: Store sulfuric acid away from calcium nitrate and silver nitrate.

Disposal Barium is very toxic. The barium sulfate should be placed in a waste receptacle to go to the toxic waste site. The solutions can be flushed down the drain with plenty of water. The phosphate precipitate can be thrown into the trash.

ANALYSIS OF FERTILIZER PRE-LAB QUESTIONS

1) What would be the content of a fertilizer that is labeled 20:20:10?

20% of the fertilizer is nitrogen, 20% is phosphorous, and 10% potassium

140

2) Some plants such as conifers and other evergreens need a low soil pH. Which number would you want to be the higher number on the fertilizer package for these plants?

In order to lower the pH of soil, more nitrogen is required, so the first number would be higher.

3) Flowering plants use different nutrients from the soil than does a lawn. If a lawn food is 1530-15, but a food for flowering plants is 30-10-10, which nutrient is needed most for each type of plant?

The lawn needs phosphorus, but the flowering plants need more nitrogen.

4) Some fertilizers are labeled “all-purpose”. What would you expect the relative nutrient content to be in such a fertilizer?

If the fertilizer is intended to work for general use and be useful to any plant, it would need to have a fairly even distribution of nutrients. Some fertilizers have a 20:20:20 rating and would qualify for this purpose ANALYSIS OF FERTILIZER SAMPLE REPORT SHEET I . Describe the result of each test on the known solutions Nitrate - there is a brown ring between the two layers Ammonium - ammonia can be smelled Potassium - the potassium burns violet through the cobalt glass Sulfate - there is a precipitate II . FERTILIZER TEST PRESENT _________ _________ _________ _________ _________

Nitrate Ammonium Potassium Phosphate Sulfate

III. Mass of fertilizer phosphate precipitate and paper Mass of paper Mass of fertilizer phosphate precipitate (MgNH4PO4.6H2O)

ABSENT ________ ________ ________ ________ ________ -

2.84 1.60 1.24

g g g

1 mol P2 O 5 1 mol  245.3 g 2 mol MgNH 4 PO 4  6H 2 O 141.9 g P2 O 5 × = _0.359 g P2O5 1 mol P2 O 5

__1.24____g MgNH4PO4.6H2O ×

0.359 g P2O5  100% 3.00 g fertilizer

= __12.0______% P2O5

141

IV. QUESTIONS 1. Arsenic reacts chemically like phosphorus. Why is fertilizer not expected to contain arsenic?

Arsenic is toxic and can contaminate the soil and harm soil life and any animals that eat the plants.

2. Liquid fertilizers often contain the same ions as a solid fertilizer, only they are dissolved in water. Will the tests in this investigation work on liquid fertilizers?

Yes 3. Which fertilizer would be needed if the soil were poor in nitrogen, a 10-10-10 or a 20-5-5 ?

20-5-5 3. How did your measured percent phosphate compare with the amount indicated on the label of the fertilizer container?

Answers will vary.

5. A popular fertilizer for vegetable gardens is listed as 18-18-21. What does this say about the relative importance of each nutrient to vegetable plants?

Vegetable plants need large amounts of nitrogen and phosphorus, but they also need a larger amount of potassium.

#53

Saponification

Reagents 20% sodium hydroxide solution [NaOH] sodium chloride [NaCl] 1M calcium or magnesium nitrate solution pH paper

15 mL/pair 18 g/pair 1 mL/pair 1 strip/pair

Common Materials cottonseed oil

10 mL/pair 142

scented oil

1-2 drops per pair

Special Equipment filter paper

1 piece/pair

Laboratory Equipment 150-mL beaker 10-mL graduated cylinder ring stand, ring, and wire gauze

100-mL graduated cylinder stirring rod laboratory burner

Notes The reaction also works with Crisco™, etc., but the calculation would not be as meaningful.

If desired, the soap can be produced quicker if a stronger (35%) solution of sodium hydroxide is used. This allows the process to be sped up to the point that the four distinct phases in the soap-making process are clearly distinguishable. (boiling, foaming, sizzling, and forming the waxy soap). The fire should be removed as the waxy substance begins to form, but the stirring should continue until all the material in the bottom of the beaker is waxy. If scented oil is to be added, one to two drops should be added after the formation of the waxy soap. Over cooking produces a hard crumbly soap similar in texture to laundry detergent. Under cooking forms a "liquid soap" which will not set up.

Disposal The solids can be thrown into the trash. The solutions can be flushed down the drain with plenty of water.

SAPONIFICATION PRE-LAB QUESTIONS 1) Baths have not always been comfortable. Using the information about “grandma’s” soap, give several reasons why that was so.

The soap was not “pH balanced” and stung. It also had rough abrasive substances in it.

The soap did not smell nice, in fact quite the opposite.

2) Why do the two reactants in step one remain in separate layers unless stirred?

The sodium hydroxide solution is water based and the water will not mix with the oil component of the process.

3) Why should scent be added to the soap after it is complete but before it is pressed into a cake or molded?

The scent could be destroyed or evaporated by the heating process.

4) Soap that is ground into flakes or tiny pellets is used as dishwasher detergent or laundry soap.

Would these types of soap be cooked a longer or shorter time than the soap you have made?

Why do you think so?

They would have been cooked longer to remove more of the water and other moisture.

SAPONIFICATION SAMPLE REPORT SHEET I. SOAP MASS: ____________ Instructor's approval of soap: ____________ II. QUESTIONS 1. Explain why you washed the soap with salt solution rather than water.

143

The salt solution is used to separate the soap from the glycerol. The salt causes the soap to coagulate and water would not do this.

2. What was the pH of soap in water?

The pH of the soap in water was 10.

3. How do you explain this pH? The soap is basic because of the sodium hydroxide which is used to make it.

4. Describe any observations about washing your hands with a small piece of the soap.

Slippery - some foaming. It stings if not rinsed well.

5. In step 11, which aqueous solution allows the soap to lather more freely?

The distilled water lathers more quickly.

6. What is the chemical name of the other two fatty acids that come from cottonseed oil if the soap, sodium linoleate, is made from the fatty acid, linoleic acid, when neutralized by sodium hydroxide?

The other two fatty acids are oleataic acid and palmataic acid.

7. The three most abundant soap molecules produced from cottonseed oil are sodium linoleate, often written as C18H31O2Na; sodium oleate, written as C18H33O2Na; and sodium palmitate, written as C16H31O2Na. The average soap molecule is about 300 g/mol. The cottonseed oil would consist of the three soap molecules as fatty acids combined on a three-carbon backbone. Cottonmseed oil would be about 872 g/mol. Calculate the expected yield of soap using the cottonseed oil as the limiting reagent. The density of cottonseed oil is 0.92 g/mL. mole cottonseed oil 3 mol soap 300 g soap 0.92 g 10 mL oil × × × × = __9.5_g soap 872 g cottonseed oil mole soap mol soap mL 8. How does your actual mass compare to the calculated yield? Does your soap still have water in it?

Student response here. (It should be less as some will dissolve, some will be left in the beaker, etc.)

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#54

Soap in Hard Water

Reagents calcium nitrate [Ca(NO3)2] sodium carbonate [Na2CO3] sodium meta-phosphate [NaPO3] Common Materials Ivory Snow™ or castille soap Tide™ red food coloring borax Laboratory Equipment stirring rod 250-mL beaker 1000-mL beaker eyedropper

0.56 g/pair 0.1 g/pair 0.2 g/pair 2.5 g/pair 1 g/2 pairs 0.5 mL/pair 0.4 g/pair 125-mL Erlenmeyer flask laboratory burner ring stand, ring, and wire gauze filter paper

Notes It will save lab time if the solutions are prepared ahead of time.

True soaps are becoming harder and harder to find. Check bath bars for a true soap. Castille soap is available from many chemical dealers.

Disposal Flush all solutions down the drain with plenty of water as they are not toxic.

SOAP IN HARD WATER PRE-LAB QUESTIONS 1) List other substances that might qualify to be classified as surfactants besides soap and detergent.

Many solvents act on the surface of objects to remove unwanted dirt or contaminants.

Answers could include, dry cleaning agents, petroleum distillates, alcohols, etc.

2) Have you have ever had the experience of washing your hands or hair and had the sensation that you could not get the soap rinsed? What would you suspect about the metallic ion content of the water?

The water is probably very soft, having a small concentration of metal ions.

3) Dry cleaning is a process in which no water is used. The cleaners in the process still attract the dirt, but do not need to be attracted to polar water molecules. Describe the probable structure of a molecule of the cleaner.

The molecules are non-polar with no water-loving end. They are generally a hydrocarbon molecule and are only attracted to the oil that needs to be removed from the fabric.

4) People who rely on septic tanks for their sewage disposal are ultimately putting all of the waste water from their homes into the environment of the immediate area. Which substances would be harmful to use? Is it your assessment that they should rely on detergent or soaps more heavily? Why? Phosphates would be harmful to the environment and should not be used. The fats in soaps are more readily broken down in the environment and would be preferable to detergents.

145

SOAP IN HARD WATER SAMPLE REPORT SHEET This is sample data only.

I. SOAP AND DISTILLED WATER # of drops to lather _____72__________ volume = # drops X 0.05 mL/drop = 3.6 mL II. SOAP AND TAP WATER # of drops to lather _____360__________ volume = # drops X 0.05 mL/drop = 19.0 mL III. SOAP AND "HARD WATER" # of drops to lather _____770__________ volume = # drops X 0.05 mL/drop = 38.5 mL IV. SOAP AND "HARD WATER" WITH Na2CO3 # of drops to lather _____304__________ volume = # drops X 0.05 mL/drop = 15.2 mL V. SOAP AND "HARD WATER" WITH BORAX # of drops to lather ______688_________ volume = # drops X 0.05 mL/drop = 34.4 mL VI. SOAP AND "HARD WATER" WITH NaPO3 # of drops to lather ______743_________ volume = # drops X 0.05 mL/drop = 37.2 mL VII. DETERGENT AND DISTILLED WATER # of drops to lather ____43___________ volume = # drops X 0.05 mL/drop = 2.15 mL VIII. DETERGENT AND TAP WATER # of drops to lather _____65__________ volume = # drops X 0.05 mL/drop = 3.25 mL IX. DETERGENT AND "HARD WATER" # of drops to lather _____105__________ volume = # drops X 0.05 mL/drop = 5.25 mL X. QUESTIONS 1. Why is there a difference between distilled water and tap water in the volume of soap required for a lather? between tap water and hard water?

Distilled water has all ions removed so less soap is required to lather. Hard water requires even more soap due to the larger amount of metal ions.

2. Why does the addition of a builder to soap alter the amount of soap necessary to produce a lather?

The addition of builders causes less soap to be added to hard water because they tie up the hard water ions.

3. Which is the most effective builder?

Sodium carbonate was the most effective builder.

4. Why do clothes washed in hard water get cleaner when a builder is added?

The builder allows the soap to be more efficient. There is less soap scum formed by the metal ions to remain in clothing.

5. How is the detergent different from the soap in its ability to lather?

The detergent lathers better than soap.

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#55

Personal Products

Reagents Common Materials lipstick cleansing cream bar hand soap baby oil astringent cotton balls cologne or perfume hair spray paper towels Special Equipment 25 mL vials with stoppers or caps boiling stones (marble chips) Laboratory Equipment watch glass

1 tube/4 pairs 5 mL/pair small piece/pair 5 mL/pair 5 mL/pair 3 /pair 15 mL/pair 1 mL/pair 1 roll/lab 3/pair 2/pair small test tube

Notes Cologne will probably give better results than perfume and it is cheaper. A cheap cologne may separate better than an expensive cologne as the components are less well matched. If "imitation" colognes, which are copies of original scents, are available, these are more economical and work just as well. Aftershave will also work in place of the cologne.

Disposal The solutions can all be flushed down the drain with plenty of water. The cotton balls and dried hairspray can be placed into the trash.

PERSONAL PRODUCTS PRE-LAB QUESTIONS 1) Heavy make-up called grease paint is used in theater applications, by the military, and athletes. It does not wash off easily, many times not even with soap. Why not?

Soap and water are the usual agents in “washing.” Grease is an oil-based substance and therefore non-polar. It will not mix with water and is many times impervious to the soap molecule.

2) Why is it an advantage to the wearer that water does not remove grease paint?

Most reasons for wearing grease paint are due in some part to the likelihood that the wearer will be sweating or becoming wet by some other means. The grease will not rinse away.

3) Should an astringent be used before or after removing cosmetic substances from the skin?

Why? An astringent closes the pores to restrict dirt and other substances from entering. It is best used after cleansing so that existing impurities on the surface of the skin are not trapped in the pores.

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4) It is often said that hairspray is highly flammable. Would you expect this to be true when it is dry, or only when it is still wet? Explain.

It is the solvent in the hairspray that is flammable. Once the spray is dry, it no longer contains the flammable substance.

PERSONAL PRODUCTS SAMPLE REPORT SHEET I. QUESTIONS 1. Are the stripes made by the lipstick on your arm the same color as the lipstick, itself?

The lipsticks may or may not appear to be the same color depending on each student.

2. Are all of the stripes made by the same lipstick the same color on everyone's arm?

The lipstick stripes should not be the same color on different students’ arms due to differences in body chemistry and temperature.

3. Do some stripes look moister?

Some may look moister, again depending on each student and the amount of oil contained in each brand.

4. Which "cleanser" dissolved the lipstick best?

Soap & water and cold cream clean the best.

5. Were any of the cleansers unable to clean the lipstick off your arm?

The plain water and astringent did not clean well.

6. Did any cleanser leave a residue?

The baby oil left a residue.

7. Are the pores in the astringent area smaller?

Yes 8. Which cleanser would you use to remove lipstick?

Student response here 9. Do the three notes smell alike?

The notes do not smell the same and are definitely distinguishable from each other.

10. Does one of the notes emit an unpleasant smell?

The middle note or bottom note, depending on the cologne used, should be unpleasant.

11. Describe each note: (The following answers depend on the cologne used.)

Top - a sharp odor Middle - a sweet odor Bottom - a musty odor 12. Is the hair spray flexible?

Yes 13. Is the hair spray sticky?

The hairspray was not sticky when completely dry.

14. Is the hair spray stretchable?

Yes

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#56

Hair Chemistry

Reagents Common Materials ammonia cleanser vinegar chlorine bleach food color (blue is best) shampoo clear tape 2 dozen long straight strands of hair (12 light and 12 dark)

90 mL/pair 250 mL/pair 110 mL/pair 1 mL/pair 1 drop/pair 4 strips/pair 2 doz/pair

Special Equipment Extra stirring rods might speed up the investigation.

Laboratory Equipment 100-mL graduated cylinder 3 250-mL beakers

stirring rod watch glass

Notes Students should collect hair from their hair brushes. Combining the hair from the entire class should allow each pair to have six light and six dark strands. A visit to a hair stylist might be a good source of hair for the experiment.

The times required for curling and bleaching the hair will vary due to hair texture and any residues of shampoo, hairspray, and conditioner left on the hair. Freshly washed hair, especially if washed in baking soda, will receive the treatment best. (Baking soda is an excellent stripper and aids in removing any film from the hair.)

Disposal All of the solutions can be flushed down the drain with plenty of water. The bleach should not be poured in the drain at the same time as the other solutions. Bleach and other cleansers can react to produce poisonous chlorine gas.

HAIR CHEMISTRY PRE-LAB QUESTIONS 1) Why is hair without any residue preferred for the investigation?

Conditioners, gels, sprays, and other substances on the surface of the hair can prevent it from responding to treatment.

2) The natural curl in hair is a function of the cross-sectional shape of hair. Rounder fibers are straight and the more oblong the fibers are, the tighter the curl. What would you expect to find if you viewed a microscopic section of a hair from a horse’s tail? Hair from a poodle?

Hair from the long straight tail is thick, coarse, and heavy. It is probably quite round in cross-section. However, the kinky hair of a poodle is probably oblong in cross-section.

3) In what layer of the hair do you suppose this trait is expressed? The medulla, cortex, or cuticle? Support your answer.

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The great number of tiny fibers in the cortex can have a variety of arrangements. It is probably these fibers that give the shape to the cross-section of a hair.

4) People who spend a lot of time in swimming pools tend to have problems with the texture, strength, and color of their hair. To what would you attribute this?

Pools are full of chemicals and are treated with them continually. If the pool is basic, the water will cause the cuticle to swell and open allowing chemicals to enter the cortex, changing color and other properties of the hair. Acidic pH levels tend to harden the cuticle, making hair brittle.

5) Some hair types will not take a perm well. What might be the reason? (Hint: It can be a genetic thing.)

The cuticle probably lies very flat and close so that chemicals cannot enter the cortex.

HAIR CHEMISTRY SAMPLE REPORT SHEET QUESTIONS 2. Examine the hair that was given the “permanent” and compare it to the control strands.

Does the treated sample show more body and a greater tendency to curl on its own?

Comment on the appearance, texture, and strength of the treated hair compared to the control sample.

The hair that was treated does have more body and will curl on it own or if stretched straight and then released, it will curl back up. It appeared to be shinier than previously and it was smoother. The hair was not as strong as before it was treated.

2. Examine the hair that was bleached. Is the color lighter by 10%, 25%, 50%? (estimate) It was lighter by about 10%-25%.

3. How was the texture and strength of the hair changed?

The hair seemed thinner and rougher. The hair was not as strong as it was before it was placed in the bleach solution.

4. Describe the ability of the normal strand of hair to hold the color when rinsed in water and then when washed with shampoo.

Most of the color was rinsed out when the hair was rinsed with water.

5. Describe the ability of the hair that was treated with the base to hold the color when rinsed with water and then when shampooed. Explain what caused any differences you observe in the ability of the hair from the two coloring processes to hold the color.

The color was retained more than in the strand that was not treated with bleach. The hair treated with bleach retained more color because its coating had been broken down by the bleach, and therefore the coloring was able to dye it more easily.

6. Is the texture or strength of the treated and colored hair altered?

The texture and strength are altered some due to the bleaching.

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#57

Testing Sunscreens

Reagents 0.1 M hydrochloric acid [HCl] 0.1 M silver nitrate [AgNO3] potassium ferrioxalate [K3Fe(C2O4)3.3H2O] 0.3 M potassium ferricyanide [K3Fe(CN)6] 6 M nitric acid [HNO3] 1 M sodium hydroxide [NaOH] copper strip

<0.1 mL/pair <0.1 mL/pair 0.5 g/pair 2 mL/pair (for disposal only) (for disposal only) (for disposal only)

Common Materials plain white paper sunscreens, 3 different SPF factors sunglasses

1 sheet/pair 1 mL each/pair 1/pair

Special Equipment UV light petri dishes filter paper grease pencils

1/lab 1/pair 1/pair 1/pair

Laboratory Equipment eye droppers small test tube

stopper to fit test tube

Notes Potassium ferrioxalate is sometimes sold as potassium ferric oxalate.

Disposal Care must be taken to keep metal ions out of the water system. Thus metal ions should go to the toxic waste site. In this case, the silver and iron compounds are not water soluble so would not be dissolved if a land fill developed a leak. The silver chloride solid should be placed in a container to be buried in a secure toxic waste landfill. An alternate method for silver disposal is to reclaim it as metal. Dissolve the silver salt in 6 M nitric acid (HNO3). Neutralize with 1 M sodium hydroxide (NaOH) to pH 7. Place a clean copper strip in the solution. Set for twenty minutes (or overnight). The silver metal will wipe off of the copper easily. Wipe the silver into a filter cone (see Investigation #5). Filter the silver metal. Allow to dry. Place the used white paper into the trash. All other compounds may be flushed down the drain with plenty of water.

TESTING SUNSCREENS PRE-LAB QUESTIONS 1) In areas of the globe where the seasons produce months at a time with no daylight, what might you expect to find as a common deficiency in humans?

Since sunlight is required for the production of Vitamin D, it could be assumed that people with little or no sunlight for extended periods would have a deficiency of Vitamin D.

151

2) Why would it be advantageous to have your eyes protected from UV light rather than simply shielding them from the brightness of the light?

The UV ray are damaging to tissues, especially the delicate tissues of the eyes.

3) Can you think of other uses for the compound that reacts with UV light to turn blue?

The color changing properties would make it suitable as an indicator for UV radiation.

4) Would an SPF rating mean the same thing on the equator as it would in New York City?

Explain your reasoning.

The SPF rating is an average for the various locations and times of the year on Earth. At the equator, the rays of the sun are nearly perpendicular, meaning that the density of the radiation is greater than at either higher or lower latitudes. The sunscreen would not give the same level of protection to a person at the equator.

TESTING SUNSCREENS SAMPLE REPORT SHEET I. SILVER 1. Did all of the white solid turn black? Why or why not.

No, not all of the solid turned black. The key protected part of it from the UV light.

2. Films are made with AgBr in a gel. What is the chemical reaction that occurs when light hits the film?

Ag+ + 1e- → Ag+ II. SUNSCREEN TESTING (The following is sample data only.)

Brand name Banana Boat Vaseline Intensive Care SPF value 30+ 25

Coppertone 15

1. Which sunscreen protected best, that is absorbed the most UV rays?

2. Which protected the least?

30+

15

3. What color was the white paper under the UV light? Why? The color was blue because the iron(II) reacted with the ferricyanide ion.

4. How can you prevent additional reaction with light of the chemicals on the white paper, that is, how can you "fix" the results?

Permanent "fixing" could be accomplished by converting the silver salt to a salt which is not light sensitive.

5. Did the sunglasses block part of the UV rays?

Yes, the sunglasses protected some from the UV light.

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#58

Lead Detection

Reagents 6 M nitric acid [HNO3] 1M or 6M acetic acid [CH3COOH] 1% sodium sulfite [Na2SO3] 0.1 M potassium iodide [KI] lead (II) acetate [Pb(C2H3O2)2] pH 7 buffer solution 1% starch solution iodine solution 1% glucose solution sodium sulfide [Na2S] 1M ferric chloride [FeCl3] iron(III) chloride sodium carbonate [Na2CO3] pH paper

1 mL/pair 10 -20 mL/pair 25 mL/pair 1 mL/pair <1g/pair 10 mL/pair 20 mL/pair 15 mL/pair 10 mL/pair (for disposal only) (for disposal only) (for disposal only) 1 roll/pair

Common Materials egg samples to test

1 for every 4 pairs 1 small piece/pair

Special Equipment 1000 mL volumetric flask hood centrifuge filter paper

1/lab 1/lab 1/lab 2/pair

Laboratory Equipment 50-mL beaker hot water bath

large test tubes balance

Notes pH 7 buffer solution Add 9.1 g KH2PO4 + 18.9 g Na2HPO4 per liter solution.

Iodine solution Dissolve 25 g of iodine in a solution of 50 g of potassium iodide in 200 mL of water.

Caution: Store acids away from sodium sulfide and store nitric acid away from acetic acid.

Lead paint chips should be readily available from any location where an older house (pre 1960) is being torn down or is vacated and peeling. Marine paints are often lead-based paints.

This investigation can be used in conjunction with investigation #59.

Disposal Lead, of course, is toxic and must be kept out of the water system. Collect all used solutions containing lead. Dissolve the lead iodide in 6 M nitric acid. Add a three-fold excess of sodium sulfide (Na2S). Let the solution set for an hour. Adjust the pH of the solution to neutral (pH 7 or more). Use pH paper to check pH. Filter the precipitate. Discard the precipitate and

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filter paper in a container to be buried in a secure toxic waste landfill. The filtrate contains sulfide. Remove the sulfide by adding 1M iron(III) chloride (ferric chloride) in a three fold excess. Neutralize the solution with sodium carbonate (Na2CO3). Filter the precipitate. Put the precipitate and filter paper in a container to be buried in a toxic waste landfill. The filtrate can be flushed with plenty of water. The other solutions can be poured down the drain with lots of water.

LEAD DETECTION PRE-LAB QUESTIONS 1) Lead is commonly used in “shot” to load shotgun shells for such sports as quail hunting.

Evaluate this practice in light of environmental concerns. Suggest an inexpensive alternative.

Lead in the environment persists for a very long time because it is near the bottom of the activity series and does not react easily with other substances in the environment.

Aluminum or steel shot would be inexpensive to produce and would not be toxic in the environment.

2) Why would soils along highways still contain a high level of lead even though it has not been added to gasoline since 1986?

The lead in gasoline is expelled through the exhaust of the automobile. This lead does not break down or combine with other substances in the environment and remains.

3) Recent news stories have revealed a high level of lead in some trinkets and inexpensive jewelry imported from other parts of the world. These items are often found in toydispensing machines in grocery and variety stores. Examine the risk to children.

Small trinkets are often swallowed or chewed on by young children. They can also be lost outdoors where pets and other animals ingest them. Lead is toxic to living systems.

4) What is the relationship between pencil lead and lead metal?

Pencil lead is not made of lead metal, but made of graphite.

LEAD DETECTION SAMPLE REPORT SHEET I. DETECTING LEAD * Student responses depending on sample.

1. Sample: __________________________________________ Lead positive test Yes ________ No ________ 2. Lead Concentration 207 g Pb ___0.05____g Pb(CH3CO2)2 × = ___0.0415 g Pb 325 g Pb(CH 3 CO 2 ) 2 0.0415 g Pb 1 mL solution  × 106 ppm = ____41.5______ppm Pb 1000 mL 1 g solution

II. EFFECT ON A PROTEIN 1. Describe the effect of lead on albumin.

154

The egg white becomes milky, so the protein is being denatured.

2. One emergency treatment for ingesting lead is to swallow egg white or milk.

Why? Lead will denature the protein in the egg white and be absorbed by it instead of the stomach cells.

3. Why is an emetic (vomit inducer) the next step of the treatment?

The denatured material with the lead must be removed from the body or it will make the person sick as the material is digested and the lead is released into the system.

III. EFFECT ON AN ENZYME 1. Describe the color change, or lack of, in each test tube.

Test tube #1 changes, but test tube #2 does not.

2. Did the amylase in saliva break down the starch in test tube #2? Why not?

No. The protein in saliva has been denatured by the lead.

IV. QUESTIONS 1. What is the purpose of the digestion step (step 2)?

This is a heating step similar to what is done when ingested into the body.

2. The lead content of our drinking water is a concern in many areas of the US. List possible methods by which our water supplies are becoming contaminated.

Answers will vary but may include statements concerning, ground water contamination, run-off, irresponsible dumping of toxic substances in sewers and drains, poor clean-up of old dumps and waste sites, etc.

3. What was the purpose of test tube #3 in the enzyme experiment?

The purpose of test tube #3 is a control sample that shows that no color change occurs when the iodine is in the presence of glucose.

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#59

Parts per Million

Reagents copper (II) chloride [CuCl2] 1M silver nitrate [AgNO3] 6 M nitric acid [HNO3] 1 M sodium hydroxide [NaOH] copper strip pH test paper

0.014 g/pair 12 mL/pair (for disposal) (for disposal) (for disposal) (for disposal)

Special Equipment 1000 mL volumetric flask 10 mL pipette with bulb filter paper

1/pair 1/pair (for disposal only)

Laboratory Equipment balance 50-mL beaker 3 1000-mL beakers

10-mL graduated cylinder dropper

Notes This investigation can be used in conjunction with investigation #57.

Disposal Filter the solutions and discard the solids and filter paper in the silver waste container. The copper solutions should be evaporated and the solid buried in a secure landfill. However in this investigation, the concentrations are so small as to not be a threat to the environment.

An alternate method for silver disposal is to reclaim it as metal. Dissolve the silver salt in 6 M nitric acid (HNO3). Neutralize with 1 M sodium hydroxide (NaOH) to pH 7. Place a clean copper strip in the solution. Set for twenty minutes (or overnight). The silver metal will wipe off of the copper easily. Wipe the silver into a filter cone (see Investigation #5). Filter the silver metal. Allow it to dry.

PARTS PER MILLION PRE-LAB QUESTIONS 1) Why is it a concern that lead was an additive to gasoline for over half a century and is now found in the environment in relatively high concentrations, especially along roadways?

Lead persists for a long time without being changed into other substances or becoming non-toxic. It can be ingested by animals and become part of the food chain or get into waterways through ground water and run-of.

2) Explain why it is unreasonable to expect that there will be absolutely no toxins in anything we take from the environment?

Absolutely pure substances simply do not exist.

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3) Evaluate the importance of determining toxicity levels of various toxins for not only primitive life, but all animals in the food chain, and ultimately for humans.

Since we realize that there will always be impurities in the environment, it is important for maintaining the health of the biosphere to understand the levels that pose problems to living systems.

4) Why is it important that scientists be able to detect minute concentrations of substances in our air and water?

We need to be able to test for the presence of toxic substances in order to monitor the health of the ecosystem.

PARTS PER MILLION SAMPLE REPORT SHEET *The following is sample data only.

I. Mass CuCl2 + beaker minus mass beaker mass CuCl2

_

g g _____0.014____g

Solution #1 Molarity

0.014 g CuCl 2 mL mol =  -3  1000 mL solution 10 L 135 g

1.04x10-4 M

Assume that the solution is 1.00 g/mL 0.014 g CuCl 2 mL   106 ppm = ppm 1000 mL solution 1.00 g Color of solution Drops to precipitation

14 ppm

faint blue 2

Solution #2 Molarity

ppm

1.4  10 -4 M solution #1 10 mL = 1000 mL solution

14 ppm  10 mL = 1000 mL solution

Color of solution Drops to precipitation

1.4×10-6 M

0.14 ppm

colorless 20

Solution #3 Molarity

ppm

1.4  10 -6 M solution #2  10 mL = 1000 mL solution

14 ppm  10 mL = 1000 mL solution

1.4×10-3 ppm

157

1.04×10-8

M

ppb

1.4×10-3ppm ×1000 =

Color of solution Drops of precipitation

1.4

ppb

colorless 50+

II. QUESTIONS 1. Where the chloride ion was detected, what can be said about copper ion in solution?

There is not enough copper to bond to the chloride ion so all the copper is out.

2. Could the chloride in solution number #3 be detected by this method?

No 3. Was there copper in solution #3?

Yes, but there was very little.

4. Can the presence of copper be detected by the color of the solution in #1?

Yes, it is faint blue.

In the solution in #2?

No In the solution in #3?

No Explain.

If the concentration is very low, color is not detectable to the human eye.

5. If the copper had been a toxic element like lead, would solution #3 be safe? Explain your answer.

Yes. The concentration is below the level that is understood to be toxic.

6. Why was it important to use distilled water to rinse the lab ware in this lab?

Tap water contains metal ions that would cause the lab results to be incorrect.

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Reagent List (Alphabetical) Chemical Name Acetic acid - concentrated (Glacial) 6M 1M 0.5 M Acetic anhydride Acetone Alcohol (isopropyl or ethyl) Aluminum sulfate 1 M Ammonia (ammonium hydroxide) concentrated 2M 1% solution Ammonium chloride Ammonium nitrate Ammonium sulfide Ammonium phosphate Barfoed's reagent Barium chloride, 0.1 M Barium chloride, 5% solution Barium nitrate, 5% solution Benedict's reagent Benzaldehyde Benzyl alcohol Buffer solution (pH of 7) Butanol Butyric acid Calcium metal Calcium acetate Calcium chloride, solid 1M Calcium hydroxide, sat. soln.

Calcium nitrate: solid 1M Calcium oxide Cellulose, 1% solution Cerric nitrate solution Chromatography paper Cobalt(II) nitrate 0.4 M Copper: strips and wire (sm. gauge) tokens

Formula [CH3COOH]

[(CH3CO)2O] [CH3COCH3] [CH3CH(OH)CH3] or [CH3CH2OH] [Al2(SO4)3] [NH3] or [NH4OH]

[NH4Cl] [NH4NO3] [(NH4)2S] [(NH4)3PO4] [BaCl2] [BaCl2] [Ba(NO3)2] [C6H5C(O)H] [C6H5CH2OH] [CH3(CH2)3OH] [CH3(CH2)2COOH] [Ca] [Ca(C2H3O2)2] [CaCl2] [Ca(OH)2] [Ca(NO3)2] [CaO] [NaNO3] [Co(NO3)2] [Cu]

159

Labs 24,25,45,47,49,51 41, 58 11, 20, 21, 42 46 49 19, 24, 27,39,48,51 15, 34, 33, 35 40 4, 45 52, 47 47 3 37 8 52 51 5 52 52 51 24 25 51, 58 47 25 22 39 3, 7, 29 11 31 54 6, 9, 11, 22, 53 42 51 24 47 10 22, 24, 40,41,57,59 1

turnings Copper(II) acetate Copper(II) carbonate Copper(II) chloride, solid 1M 0.1 M Copper(II) nitrate, 1 M Copper oxide, fine powder Copper(II) sulfate, solid 0.1% solution 1.0 M Copper(II) sulfate pentahydrate

[Cu(C2H3O2)2] [CuCO3] [CuCl2] [Cu(NO3)2] [CuO] [CuSO4] [CuSO4.5H2O]

8, 30 51 8 59 6 5 9, 22 44 51 42, 45 8, 9, 11, 21, 40 13

Ethyl acetate Ethyl alcohol (ethanol)

[CH3CH2OC2H3O2] [CH3CH2OH]

Ethyl ether Ethylene glycol Europium oxide

[C2H5OC2H5] [HO(CH2)2OH] [Eu2O3]

49 7, 19, 24, 25, 36, 39, 42, 44, 45, 47, 48, 49, 51 45 26 5

Formic acid Fructose, 1% solution

[HCOOH] [C6H12O6]

25 51

Glucose, 1% solution Glycerol (or glycerin)

[C6H12O6] [HOCH2CHOH CH2OH]

51, 58 26, 35, 50

Hexamethylenediamine Hexane Hydrochloric acid: conc. 12 M 6M 3M 1M 0.5 M 0.1 M Hydrogen peroxide, 3% solution 30% solution

[H2N(CH2)6NH2] [C6H14] [HCl]

26 24, 26, 49, 50 4, 5, 6, 10, 52 8,12,17,30,43,52 52 3, 6, 16, 20 3, 9, 46 18, 57 38, 43, 44 38

Iodine, solid Iodine solution Iodine tincture Iron metal, small pieces Iron(II) ammonium sulfate, 0.2 M Iron(III) chloride (Ferric chloride),1 M 0.1 M 1% solution Iron(II) nitrate, 1 M

[I2]

[H2O2]

[Fe] [Fe(NH4)2(SO4)2] [FeCl3] [Fe(NO3)2]

160

51, 58 41, 50, 51, 58 21 22 10 42, 43, 48, 58 10 24 22

Iron(II) sulfate solution, 0.2 M 1M Isbutyl alcohol Isobutylamine Isopropyl alcohol 75% solution

[FeSO4]

Lead (sheet) Lead (II) acetate Lead (II) nitrate solution, 1 M Lithium chloride, 1 M 0.1 M Litmus paper: red blue Luminol

[Pb] [Pb(C2H3O2)2] [Pb(NO3)2] [LiCl]

28 58 9, 42, 43 6 5 18, 42, 52, 30 18, 30 44

Magnesium chloride, 1 M Magnesium ribbon Magnesium sulfate heptahydrate (Epsom Salt) 10% solution Magnesium nitrate, 1M Manganese dioxide, 0.012 M Methanol

[MgCl2]

11 5, 12, 22

Nitric Acid:

[HNO3]

conc.

6M

1M n-octanol n-octene n-pentanol (n-amyl alcohol)

[CH3CH(CH3)CH2OH] [(CH3)2CHCH2NH2] [CH3CH(OH)CH3]

[MgSO4.7H2O] [Mg(NO3)2] [MnO2] [CH3OH]

[CH3(CH2)7OH] [CH3(CH2)5CH2] [CH3(CH2)4OH]

ParafilmTM pH test paper

Phenolphthalein Phenol red indicator Phthalic anhydride Polyvinyl alcohol Potassium chloride, solid 0.1 M 1M Potassium dichromate

[C8H4O3] [(CH2CHOH)x] [KCl] [K2Cr2O7]

161

52 6 25 24 2, 15, 29 52

7, 13 52 22, 53 33 25 30, 42, 45 1, 6, 8, 41, 42, 52, 57, 58, 59 22 25 24 25 26, 47 6, 8, 18, 19, 20, 24, 33, 42, 43, 47, 48, 50, 51, 53, 57, 58, 59 9, 17, 26, 39, 44 42 26 27 7, 52 5 6 48

Potassium dihydrogen phosphate Potassium ferricyanide, 0.3M 0.1 M Potassium ferrocyanide, 0.1 M Potassium ferrioxalate Potassium hydrogen sulfate Potassium iodide, solid 0.1 M 1M Potassium permanganate, 1% soln.

Potassium thiocyanate, 0.1 M Potassium sulfate Propane, 2-chloro, 2-methyl Propanoic acid Radiation Sources (alpha, beta, gamma) Resorcinol Salicylic acid Sebacoyl chloride Silver nitrate 1 M 0.1 M Sodium acetate, solid 1M Sodium bisulfate Sodium carbonate, solid 1M Sodium chloride, solid 0.1 M 1M Sodium citrate Sodium ferrocyanide Sodium hydrogen carbonate, solid (Sodium bicarbonate) Sodium hydrogen phosphate Sodium hydroxide pellets Sodium hydroxide solution: 8M 6M

[KH2PO4] [K3Fe(CN)6]

[KMnO4] [KSCN] [K2SO4] [CH3CCH3ClCH3] [CH3CH2COOH]

51, 58 57 10 8, 10 57 50 7, 51, 58 58 8, 9 24 10 40 24 25

[C6H4(OH)2]

28 24

[K4Fe(CN)6] [K3Fe(C2O4)3.3H2O] [KHSO4] [KI]

[C7H6O3] [ClCO(CH2)8COCl] [AgNO3] [NaC2H3O2] [NaHSO4] [Na2CO3] [NaCl] [NaC6H7O6] [Na4FeCN6] [NaHCO3] [Na2HPO4] [NaOH] [NaOH]

7, 25, 49 26 6, 59 57 26 20 8 19, 42, 43, 54, 51, 48, 58 9 7, 23, 29, 53 5 6, 9, 21 51 8 7, 9, 11, 16, 23, 33, 46 51, 58 3, 17, 26, 44 30 1, 8, 33, 39, 42, 43, 45, 50, 51 42, 48 53 52 9 3, 6, 9, 11, 20, 57, 59

3M 20% solution 10% solution 1.5 M 1M

162

0.1 M Sodium hypochlorite solution, 1.5 M Sodium iodide solution, 0.1M Sodium meta-phosphate Sodium nitrate Sodium phosphate Sodium pyrophosphate Sodium silicate Sodium sulfate, solid 1M Sodium sulfide, solid 3M Sodium sulfite, solid 1% solution 1.5 M Sodium thiosulfate, 50% solution Sodium tripolyphosphate Starch, 1% solution Strontium chloride, 0.1 M Sucrose, 1% solution Sulfuric acid: conc.18 M

[NaClO] [NaI] [NaPO3] [NaNO3] [Na3PO4] [Na4P2O7] [Na2SiO3] [Na2SO4] [Na2S] [Na2SO3] [Na2S2O3] [Na5P3O10] [SrCl2] [C12H22O11] [H2SO4]

6M 3M

18 36 38 54 52 30 19 19 7, 52 11 42, 48, 58 43 9 58 37 50, 51 19 41, 51, 58 5 51 3, 6, 26, 40, 52, 49, 50, 48 25 33, 50, 51

Thioacetamide Tin metal Tin (II) nitrate TLC plates

[CH3CSNH2] [Sn] [Sn(NO3)2]

7, 23, 48 22 22 49

Urea, 2% solution Urease, 1% solution

[H2NCONH2]

42 42

Water, distilled

H2O

4, 7, 26, 30, 41, 42, 49, 54, 57, 59

Zinc metal, small pieces powdered mossy

[Zn]

22 1, 44 21, 30

Zinc nitrate

[Zn(NO3)2]

22

163

Common Materials (Alphabetical) Acrylic, clear spray Aftershave Ajax® laundry detergent (liquid) Almonds, sliced Aluminum cans Aluminum foil Ammonia cleaning product Antacid tablets Alka-Seltzer® Tums® Rolaids® Anti-freeze Aspirin; aged commercial fresh commercial Astringent

1 19 2 45 33 28, 7, 21, 23, 26, 40 18, 19, 56 17, 46 17, 46 17, 46 17, 46 2 49 49 55

Baby oil Baking powder Baking soda Balloons, large party Beets Bleach, chlorine Bloodstains Borax Brake fluid Bread, white Butter Butter beans

55 46 46 11 18 19, 21, 23, 56 44 27, 54 2 45 50 45

Cabbage, red Can lids Candle, emergency size small birthday Cardboard for lid to Styrofoam cup Castille soap (or Ivory Snow®) Charcoal, activated Cheese, Swiss Cheesecloth Cigar, little Cigarettes: filtered non-filtered Citrus fruits Clay, modeling Cleansing cream

18 7 5 31 3 54 33 45 45 32 32 32 40 44 55 164

Cloth Coins Cologne Copper tokens Cork, small Corn syrup Cotton balls Cotton swabs Cottonseed oil Crisco®

14 21, 40 19, 55 1 2 51 39, 40, 55 4, 44 53 7

"Dirty" water Dish washing detergent: machine hand Downy® fabric softener Drain cleaner Drink mixes (fruit) Dry ice

33 19 15, 19 2 18, 19 41 33, 34, 35

Egg, raw Elmer's Glue®

42, 58 27

Fabric dyes Fertilizer Food to test for proteins Food colors (common 4) FD&C food dyes Fritos® Fruit juices, apple & orange

33 52 45 2, 4, 24, 33, 47, 54, 56 47 50 41

Gelatin, unflavored Grape juice, concentrated Gum, chewing

42 18 14

Hair Hair spray Handi-Tak® (or other removable gum adhesive) Honey

42, 56 19, 55 44

Ice Incandescent bulb (with socket) Ivory Snow® (or castille soap)

2, 4, 29, 49 5 54

Karo® syrup Ketchup

2 15

15, 51

165

Knives

18, 43

Lead samples (at least 1%) Lemon juice Lipstick

58 18 55

Marbles, glass Margarine; liquid solid Marshmallow Matches Meat, white (chicken or turkey) Metal pellets M&M's® candies Milk Mineral oil Molasses Motor oil (30 wt.)

15 50 50 45 19, 32, 39 45, 50 36 47 42, 45 2, 39 2 2

Nailclippers Nail polish remover Notebook paper

42 19 6

Olive oil Oven cleaner

50 19

Paint thinner Paper clips large metal aluminum plastic Paper towels Paper, white Paraffin Peanut butter Peanut oil Plastic wrap Plastic sleeve, clear, letter size Potato Potato chip Power steering fluid

19 27, 40, 47, 49 26 2, 29 2 21, 23, 24, 35, 40, 49, 55, 57, 9, 10, 28, 57 2, 7 14 50 47, 49 9, 10 43, 50 50 2

Ruler, metric Radish Rice Rubber bands

1, 12, 16, 27,38 45, 50 45 39, 45

166

Rubber, small pieces (stopper, rubber band, eraser)

2

Sandpaper Scented oil Shampoo, clear or white Silver piece (tarnished) Skittles® candies Smoke detector Soap, bar liquid Soft drink, clear Starch, soluble Steel wool, fine Straw, drinking Styrofoam cups 6.4 oz.

8.5 oz.

10 oz.

Styrofoam pieces (peanuts) Sugar: powdered table Sunglasses Sunscreens Syrup

12 53 18, 56 23 47 28 55 14, 23 18 15 8, 40 16

Tape, clear Thumbtacks Thread Tide® Toilet bowl cleaner Transmission fluid

44, 56 2 12 54 19 2

Vanilla extract Vegetable cooking oil Vinegar Vitamin C tablets

19 2, 14, 15, 50 18, 19, 56 41

Wire for stirrer Wool, undyed or natural color

3 47

37 3, 36, 37 27 7 51 7, 45, 51 57 57 14

167