Test Bank For World of Chemistry, Student Edition for Alabama, 1st Edition by National Geographic Ch

Name______________________________ Section__________ Date______________

Chapter 1

Chemistry: An Introduction I. Matching Match the description in Column B with the correct term in Column A. Write the letter in the blank provided. Each term matches with only one description, so be sure to choose the best description for each term.

Column A

Column B

_____ 1. chemistry

A. process used to make and explain discoveries and to solve problems

_____ 2. scientific method

B. involves a number and a unit

_____ 3. observation

C. a known limitation to a possible solution

_____ 4. hypothesis

D. summary of observations

_____ 5. engineering

E. science that deals with the matter of the universe and the changes they undergo

_____ 6. theory

F. explanation that has stood up to much testing and investigation

_____ 7. natural law

G. noting a measurement or quality of some substance or process

_____ 8. measurement

H. the application of scientific practices to solve real-world problems

_____ 9. criteria

I. possible explanation for a problem or question

_____ 10. constraint

J. the requirements that a solution must satisfy

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II. Multiple Choice Choose the one best answer and write its letter in the blank. _____ 11.

Which of the following careers involve chemistry? a) engineer b) beautician

_____ 12.

c) farmer d) all of the above

Which of the following is a microscopic property? a) color of a solid b) density of a liquid

_____ 13.

c) number of atoms in a molecule d) texture of a solid

Which of the following is a quantitative observation? a) The object had a mass of 2.3 g. b) The liquid looked yellow. c) Carbon dioxide was produced. d) A yellow solid formed.

_____ 14.

Which of these is a chemical change? a) Wood burns. b) An egg cooks.

_____ 15.

c) Food is digested. d) all of the above

When using the scientific method, the first step is to a) design an experiment. b) complete several experiments. c) make a hypothesis. d) make observations.

_____ 16.

When using your five senses to gather information, you are a) drawing a conclusion. b) developing a theory. c) making observations. d) establishing scientific law.

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_____ 17.

How do the practices of science and engineering differ? a) Scientists make observations; engineers take measurements. b) Scientists communicate information; engineers develop models. c) Scientists construct explanations; engineers design solutions. d) Scientists use mathematical thinking; engineers use computational thinking.

III. Free Response Answer the questions in the space provided. Use complete sentences and provide support for your answer.

18.

Describe the difference between a natural law and a theory.

19.

Briefly define the term trade-off and explain why engineers often must make trade-offs when developing solutions.

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Chapter 2

Matter I. Matching Match the description in Column B with the correct term in Column A. Write the letter in the blank provided. Each term matches with only one description, so be sure to choose the best description for each term.

Column A

Column B

_____ 1. element

A. not the same throughout

_____ 2. physical properties

B. has mass and occupies space

_____ 3. compound

C. cannot be separated into simpler substances

_____ 4. solid

D. the same throughout

_____ 5. physical change

E. atoms of different elements chemically combined

_____ 6. solution

F. has a fixed shape and volume

_____ 7. homogeneous

G. a homogeneous mixture

_____ 8. gas

H. has a fixed volume but assumes the shape of its container

_____ 9. chemical change

I. inherent characteristics of a substance

_____ 10. matter

J. composition of a substance is not affected

_____ 11. liquid

K. no fixed volume or shape

_____ 12. heterogenous

L. a new substance is formed

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II. Multiple Choice Choose the one best answer and write its letter in the blank. _____ 13.

Anything that has mass and occupies space is called a) a solid.

_____ 14.

b) oxygen

c) ice

d) air

c) melts at 83 °C d) floats in water

Which of the following is a physical change? a) rusting

_____ 17.

d) a compound.

Which of the following is a chemical property? a) burns b) conducts electricity

_____ 16.

c) matter.

Which of the following would be classified as a mixture? a) carbon

_____ 15.

b) a liquid.

b) freezing

c) rotting

d) cooking

Which of the following is characteristic of a compound? a) pure substance b) homogeneous c) contains different kinds of atoms d) all of the above

_____ 18.

A liquid is a state of matter that a) flows. b) takes the shape of its container. c) is easily compressed. d) has a definite shape.

_____ 19.

A chemical change must produce which of the following? a) a new substance b) a pure substance

c) a compound d) a mixture

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_____ 20.

Ice, liquid water, and water vapor have a) similar physical and chemical properties. b) similar physical but different chemical properties. c) different physical but similar chemical properties. d) different physical and chemical properties.

_____ 21.

In a laboratory experiment, a sample of sugar is dissolved in water, and the solution is then evaporated to dryness. Upon further heating, the sugar turns brown. In correct sequence, the following separate changes have occurred: a) physical change, chemical change, physical change b) physical change, physical change, chemical change c) chemical change, physical change, chemical change d) one chemical change, one physical change

_____ 22.

The element silicon is blue-gray in color. It is brittle in texture. It melts at 1410 °C and reacts vigorously with fluorine. Identify which of the descriptions is a chemical property. a) It reacts with fluorine. b) It is blue-gray.

c) It melts at 1410 °C. d) It is brittle.

III. Free Response For questions 23–27, read the following: The substance looked pale yellow and had a density of 3.6 g/mL. It burned readily in air and produced bubbles when reacted with acid. When heated, it changed from solid to liquid at 79 °C and from liquid to gas at 143 °C. Identify the following properties as either chemical or physical.

23.

_______________________ pale yellow

24.

_______________________ density of 3.6 g/mL

25.

_______________________ burned readily in air

26.

_______________________ produced bubbles when reacted with acid

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27.

_______________________ changed from solid to liquid at 79 °C

Answer the question in the space provided. Use complete sentences and provide support for your answer.

28.

Is there any difference between the properties of pure water that has been boiled and condensed and the properties of pure water that has been frozen and then melted?

Explain.

29.

Identify four properties that would allow you to distinguish between a cube of copper and a cube of sugar.

30.

When a small amount of red powder is heated, it darkens and then changes into a shiny silvery liquid. Is the red powder an element or a compound? Explain your answer briefly.

Can you classify the shiny liquid with certainty? Again, explain briefly.

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31.

You are given a flask that contains sea water that has been contaminated with oil. Some sand is also present in the flask. Describe how you would separate the sand, oil, sea salt, and water from each other.

32.

Make a drawing that shows why a gas is compressible but a solid is not.

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Chapter 3

Measurements and Calculations I. Multiple Choice Choose the one best answer and write its letter in the blank. _____ 1.

For the measurement 0.817 m, which digit is the estimated digit? a) 8 b) 1 c) 7 d) There is no estimated digit in this measurement.

_____ 2.

How many significant figures are there in the measurement 0.0120 g? a) two b) three

c) four d) five

_____ 3.

How many significant figures are there in 201.12 s? a) two c) four b) three d) five

_____ 4.

How many of the zeros in 0.00052010 are significant? a) one b) two

_____ 5.

c) five d) six

The correct setup for the conversion of 235 mL to L is: a) 235 mL ×

1000 mL 1L

c) 235 mL ×

1L 235 mL

d) 1 L ×

b) 1000 mL ×

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_____ 6.

Which of the following is the correct scientific notation for 0.0000007841 g? a) 7.841 × 107 g b) 7.841 × 108 g

_____ 7.

How many mm are in 3.6 km? a) 3.6 × 106 mm b) 3.6 × 103 mm

_____ 8.

c) 20,100 kg d) 20 kg

Express the product of 0.0120 cm and 5.350 cm to the correct number of significant figures. a) 0.06420 cm2 b) 0.0642 cm2

_____ 12.

c) 16 cm d) 20 cm

Express the measurement 20,141 kg rounded to two significant figures. a) 2.0 × 104 kg b) 20,140 kg

_____ 11.

c) 254 K. d) –248 K.

Express the sum of 5.0 cm and 11.28 cm to the correct number of significant figures. a) 16.28 cm b) 16.3 cm

_____ 10.

c) 3.6 × 10–6 mm d) 36 mm

The temperature of –25 °C corresponds to a Kelvin temperature of a) 248 K. b) 298 K.

_____ 9.

c) 78.41 × 10–8 g d) 7.841 × 10–7 g

c) 0.064 cm2 d) 0.06 cm2

What is the metric system prefix for the quantity 0.000000001? a) millib) centi-

c) microd) nano-

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_____ 13.

An object has a mass of 2.63 g and a volume of 2.25 mL. What is its density, and will it sink or float in water? a) 1.17 g/mL, sink b) 1.17 g/mL, float

_____ 14.

The density of gold is 19.32 g/cm3 . What is the volume of a sample of gold that has a mass of 54.82 g? a) 1059 cm3 b) 35.50 cm3

_____ 15.

c) –277 °C d) 277 °C

If the temperature of a substance changes by 100 K, by how much does it change in °C? a) 0 °C b) 373 °C

_____ 18.

c) cg d) ng

Helium boils at a temperature of 4 K. What is this temperature in °C? a) –269 °C b) 269 °C

_____ 17.

c) 2.837 cm3 d) 0.3524 cm3

Which of the following mass units is the largest? a) mg b) µg

_____ 16.

c) 5.92 g/mL, sink d) 5.92 g/mL, float

c) 273 °C d) 100 °C

If the mass of 1 L of water is 1 kg, what is the mass of 1 mL of water? a) 1000 g b) 100 g

c) 10 g d) 1 g

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II. Free Response Round each number to the indicated number of significant figures and express it in scientific notation.

19.

2501 (2 S.F.)

______________________

20.

0.030490 (3 S.F.)

______________________

21.

172590 (1 S.F.)

______________________

Give the correct units and the correct number of significant figures for the answer. Do not perform the calculation or give the numerical answer for the question.

22.

12.6 m × 2.0 m × 13.84 m =

______________________

23.

11.63 mL – 8.8 mL =

______________________

24.

(12.36 g – 11.25 g) = 10.31 mL

______________________

25.

18.5 m = 0.035 s

______________________

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Answer the following questions in the space provided. Show your work and remember to present your answers with the appropriate units and number of significant figures.

26.

If 4 quarts = 1 gallon, and 1.06 quarts = 1 liter, how many liters are there in a 55.0-gallon container?

27.

A student wants to confirm whether or not a piece of metal is pure gold. He determines the volume of the piece to be 3.18 mL and the mass of the object to be 50.8 g. Is the object pure gold? (The density of gold is 19.3 g/cm3 .) Support your answer with both a calculation and a brief explanation.

28.

A cube of aluminum measures 4.0 cm on each edge. The density of aluminum is 2.7 g/cm3 . What is the mass of the cube?

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29.

Explain briefly why we round numbers off in chemistry. Why is it important?

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Chapter 4

Chemical Foundations: Elements, Atoms, and Ions I. Matching Match the description in Column B with the correct term in Column A. Write the letter in the blank provided. Each term matches with only one description, so be sure to choose the best description for each term.

Column A

Column B

_____ 1. electron

A. sum of the number of protons and the number of neutrons in the nucleus of an atom

_____ 2. nucleus

B. positively charged subatomic particle

_____ 3. isotopes

C. dense, positively charged center of atom

_____ 4. atomic number

D. subatomic particle with zero charge

_____ 5. mass number

E. atoms with the same number of protons but different numbers of neutrons

_____ 6. ion

F. negatively charged ion

_____ 7. neutron

G. number of protons in the nucleus of an atom

_____ 8. cation

H. atom with a charge

_____ 9. proton

I. positively charged ion

_____ 10. anion

J. negatively charged subatomic particle

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II. Multiple Choice Choose the one best answer and write its letter in the blank. _____ 11.

Which of the following is a compound? a) Fe

_____ 12.

c) O2

d) K

Which is the most abundant element (by mass) in Earth’s crust and in the human body? a) carbon

_____ 13.

b) CO

b) silicon

c) neon

d) oxygen

Which of the following is NOT part of Dalton’s atomic theory? a) All elements are composed of atoms. b) Atoms of the same element are identical. c) Atoms are in constant motion. d) When atoms combine, they do so in simple whole number ratios.

_____ 14.

Which pair contains an element that is NOT diatomic? a) hydrogen and helium b) fluorine and chlorine

_____ 15.

c) bromine and iodine d) oxygen and nitrogen

The element chlorine belongs to which of the following groups? a) alkali metals b) noble gases

_____ 16.

c) halogens d) alkaline earth metals

An ion always contains a) equal numbers of protons and electrons. b) equal numbers of protons and neutrons. c) unequal numbers of protons and neutrons. d) unequal numbers of protons and electrons.

_____ 17.

The identity of an element is determined by its a) mass number. b) number of neutrons.

c) atomic number. d) atomic mass.

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_____ 18.

Which of the following statements is true? a) The proton and the neutron have opposite charges. b) The proton and the electron have opposite charges. c) The electron and the neutron are found in the nucleus. d) Atoms gain or lose protons when they form ions.

_____ 19.

_____ 20.

Which of these atoms has the greatest number of neutrons in its nucleus? a)

56 25Mn

c)

57 27Co

b)

56 26Fe

d)

56 28Ni

Which of the following is most likely to form a cation? a) a metal b) a nonmetal

_____ 21.

c) a metalloid d) a noble gas

How many of each type of subatomic particle are found in

60

Co3+ ?

a) 27 protons, 27 electrons, 32 neutrons b) 27 protons, 24 electrons, 32 neutrons c) 27 protons, 24 electrons, 33 neutrons d) 27 protons, 30 electrons, 33 neutrons _____ 22.

Which of these equations correctly shows the formation of Ba2+ from an atom of barium? a) Ba + 2e– → Ba2+

c) Ba → 2e– + Ba2+

b) Ba + 2p+ → Ba2+

d) Ba → e2– + Ba2+

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III. Free Response 23.

Complete the following table:

Element (atom/ion)

sodium atom

24.

Symbol

Atomic number

No. of protons

No. of neutrons

17

18

1

0

Mass number

No. of electrons

Charge

18 +1

23 11Na

Write the formula for the compounds that would be formed from the following ions: Al3+ and Br– ____________________

K+ and S2– ______________________

Answer the following questions in the space provided. Use complete sentences and provide support for your answer.

25.

Compare Rutherford’s model of the atom to Thomson’s model. Explain Rutherford’s reasoning in developing his model.

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26.

How might the results of Rutherford’s experiment have been different if he had a) used aluminum foil (atomic number 13) rather than gold foil (atomic number 79)?

b) fired negatively charged particles at the gold foil rather than the positively charged alpha particles?

27.

Hydrogen is a unique element in that it can form both a +1 ion and a –1 ion. For each case (H+ and H– ), identify how many of each type of subatomic particle you would find.

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Chapter 5

Nomenclature I. Matching Match the description in Column B with the correct term in Column A. Write the letter in the blank provided. Each term matches with only one description, so be sure to choose the best description for each term.

Column A

Column B

_____ 1. polyatomic ion

A. metal and nonmetal bonded together

_____ 2. binary ionic compound

B. produces H+ ions when dissolved in water

_____ 3. oxyanion

C. two nonmetallic elements bonded together

_____ 4. binary molecular compound

D. charged particle composed of several atoms bonded together

_____ 5. acid

E. polyatomic ion containing oxygen

II. Multiple Choice Choose the one best answer and write its letter in the blank. _____ 6.

When an oxygen atom forms its most common ion, it a) gains one electron. b) gains two electrons.

_____ 7.

c) loses one electron. d) loses two electrons.

When a barium atom forms its most common ion, it a) gains one electron. b) gains two electrons.

c) loses one electron. d) loses two electrons.

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_____ 8.

How many hydrogen atoms are contained in the formula for ammonium sulfate? a) two

_____ 9.

b) four

c) six

Which of the following formulas is INCORRECT? a) barium hydroxide, BaOH2 b) potassium sulfate, K2 SO4

_____ 10.

c) calcium chloride, CaCl2 d) sodium nitrate, NaNO3

The correct name for FeBr3 is a) iron bromide. b) iron(III) bromide.

_____ 11.

c) iron tribromide. d) tribromoiron.

The correct name for P2 O5 is a) phosphorus(II) oxide. b) phosphorus(V) oxide.

_____ 12.

c) diphosphorus pentoxide. d) phosphate.

The correct formula for dinitrogen tetroxide is a) N4 O2 . b) N2 O4 .

_____ 13.

c) N2 O3 . d) NO4 .

The correct formula for potassium sulfate is a) K2 SO3 . b) KSO4 .

_____ 14.

d) eight

c) P2 SO4 . d) K2 SO4 .

What is the correct name for CCl4 ? a) carbon chloride b) monocarbon trichloride

c) carbon tetrachloride d) monocarbon tetrachloride

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_____ 15.

If the formula for an ionic compound is represented by X3 Y2 , what is the charge on the X ion? a) 2+

b) 2–

c) 3+

d) 3–

_____ 16. What is the correct name for H2 SO4 ? a) hydrogen sulfate b) hydrogen sulfite _____ 17.

c) sulfuric acid d) sulfurous acid

Which of the following is NOT a molecular compound? a) P2 O5 b) NaF

_____ 18.

Which of the following is a binary compound? a) H2 CO3

_____ 19.

c) H2 O d) PCl3

b) MgSO4

c) CCl4

d) HNO3

When an atom becomes charged, a) electrons are transferred. b) neutrons are transferred. c) protons are transferred. d) protons, neutrons, and electrons can be transferred.

_____ 20.

Which of the following is INCORRECT? a) P4 O10 , phosphorus oxide b) CaSO4 , calcium sulfate c) NaNO2 , sodium nitrite d) NI3 , nitrogen triiodide

III. Free Response Fill in the blank with the corresponding name or formula.

21.

aluminum sulfide

_______________________

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22.

SF2

_______________________

23.

phosphorus trichloride

_______________________

24.

Zn(NO3 )2

_______________________

25.

iron(III) oxide

_______________________

26.

CuI

_______________________

27.

HNO3

_______________________

28.

aluminum hydroxide

_______________________

29.

CaBr2

_______________________

30.

hydrochloric acid

_______________________

31.

Ba3 (PO4 )2

_______________________

Answer the questions in the space provided.

32.

The formula for the monohydrogen phosphate ion is HPO4 2– . Write the formulas that you would expect for the ionic compound that would form between monohydrogen phosphate and a) barium. b) iron(III). c) vanadium(V). d) aluminum.

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33.

Why does the process of ionization not involve the gaining or losing of protons?

34.

Some metals on the periodic table can form ions with different charges. Identify the charge on the metal ion in each of these compounds and write the name for each compound as well. a) Fe2 (SO4 )3

b) Cu2 O

c) Co(C2 H3 O2 )2

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Chapter 6

Chemical Composition I. Matching Match the description in Column B with the correct term in Column A. Write the letter in the blank provided. Each term matches with only one description, so be sure to choose the best description for each term.

Column A

Column B

_____ 1. atomic mass unit

A. 6.02 × 1023 units

_____ 2. average atomic mass

B. mass of one mole of a substance

_____ 3. one mole

C. simplest whole number ratio of atoms in a compound

_____ 4. mass percent

D. percent by mass of a particular component of a compound or mixture

_____ 5. molar mass

E. weighted average of all the isotopes of an element

_____ 6. empirical formula

F. actual number of atoms of each element in a molecule

_____ 7. molecular formula

G. 1.66 × 10–24 g

II. Multiple Choice Choose the one best answer and write its letter in the blank. _____ 8.

What is the mass of 0.200 moles of Fe? a) 280. g b) 112 g

c) 27.9 g d) 11.2 g

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_____ 9.

The molar mass of the compound HNO3 is a) 31.02 g. b) 63.02 g.

_____ 10.

All of the following are empirical formulas EXCEPT a) Fe(OH)3 . b) Na2 SO4 .

_____ 11.

c) average atomic mass. d) mass percent.

How many moles of water are there in 45 g of water? a) 1.0 moles b) 1.5 moles

_____ 16.

c) 6.02 d) 1 × 1023

Molar mass is a whole number multiple of the a) empirical formula mass. b) amu.

_____ 15.

c) 1.5 d) 2.0

6.02 × 1023 atoms of an element have a mass in grams equivalent to the mass of _____ of the atoms in amu. a) 1 × 10–23 b) 1

_____ 14.

c) C2 H2 . d) CH.

85 g of AgNO3 represents _____ moles of AgNO3 . a) 0.50 b) 0.60

_____ 13.

c) N2 O4 . d) Sn3 (PO4 )2 .

The empirical formula for benzene (C6 H6 ) is a) C6 H6 . b) C3 H3 .

_____ 12.

c) 91.04 g. d) 93.06 g.

c) 2.0 moles d) 2.5 moles

How many atoms are there in 2.00 moles of sodium? a) 6.02 × 1023 atoms b) 1.20 × 1024 atoms

c) 6.02 × 1026 atoms d) 2 atoms

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_____ 17.

What is the mass percent of carbon in carbon dioxide? a) 27.3% b) 33.3%

_____ 18.

c) 66.7% d) 72.7%

A compound has the empirical formula of CH2 O and its molar mass is 120 g.

What is its molecular formula? a) CH2 O b) C2 H3 O2

c) C3 H6 O3 d) C4 H8 O4

III. Free Response Solve the following problems in the space provided. Show your work.

19.

Calculate the mass percent of each element in (NH4 )2 S.

20.

Find the mass, in grams, of 2.65 × 1024 molecules of Cl2 .

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21.

How many grams of sulfur are present in 83.2 g of sulfur dioxide?

22.

Determine the empirical formula for a compound that contains 14.7 g of nickel and 40.0 g of bromine.

23.

An oxide of nitrogen is found to be 30.4% nitrogen by mass. The molar mass of this compound is between 90 and 100 g/mol. a) Determine both the empirical formula and the molecular formula for this compound.

b) Name the compound.

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24.

One important hormone in the human system is thyroxine. Its molecular formula is C15 H11 NO4 I4 . a) How many atoms make up one of these molecules?

b) Is this an empirical formula? Why or why not?

c) How much does exactly one mole of thyroxine weigh?

d) If you had 7120 molecules of thyroxine, how many atoms of iodine would you have?

e) How much would two billion molecules of thyroxine weigh?

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25.

Which of the following weighs the most? the least? Which contains the greatest number of atoms? SHOW ALL WORK! 28.4 grams Al(OH)3 5.12 × 1022 atoms of barium 0.604 mole of ozone, O3

26.

What mass of nitrogen contains the same number of atoms as 48 g of oxygen?

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Chapter 7

Chemical Reactions: An Introduction I. Matching Match the description in Column B with the correct term in Column A. Write the letter in the blank provided. Each term matches with only one description, so be sure to choose the best description for each term.

Column A

Column B

_____ 1. chemical reaction

A. numbers used to balance chemical equations

_____ 2. reactant

B. substance that enters into a chemical reaction

_____ 3. product

C. process in which a new substance is formed

_____ 4. coefficients

D. written on the right side of the arrow in a chemical equation

II. Multiple Choice Choose the one best answer and write its letter in the blank. _____ 5.

The formula for chlorine gas is a) Cl. b) Cl2 (g).

_____ 6.

c) Cl(g). d) 2Cl(g).

Which of the following statements is INCORRECT? a) An equation is balanced by changing subscripts as needed. b) Chemical equations are balanced to show that matter is conserved. c) Reactants are substances written to the left of the arrow. d) Coefficients indicate the relative number of molecules of each substance in a chemical reaction.

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_____ 7.

Which set of coefficients balance the following equation?

AlCl3 + NaOH → Al(OH)3 + NaCl a) 1, 3, 1, 3 b) 3, 1, 3, 1

_____ 8.

When the equation Al + O2 → Al2 O3 is balanced using the smallest whole number coefficients, what is the coefficient for Al2 O3 ? a) 1

_____ 9.

c) 1, 1, 1, 3 d) 3, 1, 1, 1

b) 2

c) 3

In the chemical equation 2H2 (g) + O2 (g) → 2H2 O(l), the term to the right of the arrow means a) 2 grams of water. b) 2 liters of water.

_____ 10.

d) 4

c) 2 molecules of water. d) 2 atoms of water.

Which is the best description of NaCl(aq)? a) sodium chloride molecules dissolved in water b) sodium chloride crystallized from water solution c) sodium atoms and chloride atoms dissolved in water d) sodium ions and chloride ions dissolved in water

_____ 11.

How many nitrate ions are represented in the equation below?

Fe2 (SO4 )3 (aq) + 3Ba(NO3 )2 (aq) → 3BaSO4 (s) + 2Fe(NO3 )3 (aq) a) two b) three

_____ 12.

c) six d) eighteen

A student proposed the following equation for the reaction that takes place when a piece of magnesium is burned in air to produce the white powder called magnesium oxide.

Mg(s) + O2 (g) → MgO(s) Which statement is true? a) The equation is unbalanced with respect to charge. b) The equation contains one or more incorrect formulas. c) The equation is unbalanced with respect to number of atoms.

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d) The equation contains a mistake in phase notation.

When the equation NaOH(aq) + CaCl2 (aq) → Ca(OH)2 (s) + NaCl(aq) is balanced using the smallest whole number coefficients, what is the coefficient for NaOH?

_____ 13.

a) 1 _____ 14.

b) 2

c) 3

d) 4

Which observation (a–c) is evidence that a chemical reaction has occurred? a) The color changes. b) Bubbles form. c) Energy is produced or absorbed. d) All of the above are evidence of a chemical change.

III. Free Response Balance the following chemical equations: Pb(NO3 )2

→

15.

PbI2

+

AgNO3

16.

Mg

+

TiCl4

→

MgCl2

17.

C3 H8

O2

→

CO2

18.

P4

+

O2

→

19.

Na2 CO3

+

HCl →

20.

N2

I2

→

+

+

+

+

+

AgI

Ti H2 O

P4 O10 NaCl +

CO2

+

H2 O

NI3

Write balanced equations for the following reactions: 21.

zinc + hydrochloric acid ® zinc chloride + hydrogen gas

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22.

barium chloride + aluminum sulfate ® barium sulfate + aluminum chloride

23.

calcium hydroxide + nitric acid ® calcium nitrate + water

24.

magnesium + iron(III) chloride ® magnesium chloride + iron

25.

Lithium metal reacts with water to form aqueous lithium hydroxide and hydrogen gas.

Write a balanced chemical equation for the reaction, including abbreviations for the physical states.

26.

What is the primary difference between using a number as a subscript in a chemical formula and using one as a coefficient in a chemical equation? Explain briefly.

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27.

Briefly explain why, in the balanced chemical equation C + O2 → CO2 , we know that 1 g of carbon will not react with exactly 1 g of oxygen gas.

28.

When iron(III) oxide reacts with powdered aluminum, a spectacular reaction occurs and produces molten iron and aluminum oxide. Write and balance the equation for the reaction. Identify which type of chemical reaction this is.

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Chapter 8

Reactions in Aqueous Solution I. Matching Match the description in Column B with the correct term in Column A. Write the letter in the blank provided. Each term matches with only one description, so be sure to choose the best description for each term.

Column A

Column B

_____ 1. nonelectrolyte

A. produces water

_____ 2. strong electrolyte

B. electrons are transferred

_____ 3. spectator ions

C. AB(aq) + CD(aq) → AD(s) + CB(aq)

_____ 4. acid

D. A + B → AB

_____ 5. base

E. produces CO2 and H2 O

_____ 6. decomposition reaction

F. A + BC → B + AC

_____ 7. oxidation–reduction reaction

G. does not dissociate in water

_____ 8. double-displacement reaction

H. produces OH– ions in water

_____ 9. acid–base reaction

I. AB → A + B

_____ 10. single-replacement reaction

J. completely dissociates in water

_____ 11. combustion reaction

K. produces H+ ions in water

_____ 12. synthesis reaction

L. do not take part in reaction

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II. Multiple Choice Choose the one best answer and write its letter in the blank.

For questions 13–16, identify the type of reaction from the following choices: a) combination (synthesis) b) single replacement c) combustion

d) decomposition e) precipitation

_____ 13.

2Na(s) + Cl2 (g) → 2NaCl(s)

_____ 14.

2KClO3 (s) → 2KCl(s) + 3O2 (g)

_____ 15.

F2 (g) + 2NaBr(s) → 2NaF(s) + Br2 (l)

_____ 16.

CH4 (g) + 2O2 (g) → CO2 (g) + 2H2 O(g)

_____ 17.

What are the products of the following reaction?

Na3 PO4 + ZnSO4 → a) Na3 SO4 + ZnPO4 b) NaSO4 + ZnPO4 c) Na2 SO4 + ZnPO4 d) Na2 SO4 + Zn3 (PO4 )2

_____ 18.

What are the spectator ions in the following reaction?

ZnCl2 (aq) + LiOH(aq) → Zn(OH)2 (s) + LiCl(aq) a) Zn2+ and OH– b) Zn2+ and Cl–

_____ 19.

c) Li+ and OH– d) Li+ and Cl–

Which of the following is an oxidation–reduction reaction? a) LiOH(aq) + HCl(aq) → H2 O(l) + LiCl(aq) b) AgNO3 (aq) + H2 SO4 (aq) → Ag2 SO4 (s) + HNO3 (aq) c) Fe(s) + O2 (g) → Fe2 O3 (s) d) AlCl3 (aq) + NaOH(aq) → Al(OH)3 (s) + NaCl(aq)

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_____ 20.

Which substance is INSOLUBLE in water? a) NH4 Cl

_____ 21.

c) Ca(NO3 )2

b) BaSO4

Which type of substance would produce an aqueous solution that is a good conductor of electricity? a) strong electrolyte b) weak electrolyte

_____ 22.

d) Na2 SO4

c) nonelectrolyte d) all of these

Aqueous solutions of sodium iodide and lead(II) nitrate are mixed. A yellow precipitate forms. What would be the balanced net ionic equation for this reaction? a) Na+ (aq) + NO3 – (aq) → NaNO3 (s) b) Na+ (aq) + I– (aq) → NaI(s) c) Pb2+ (aq) + 2NO3 – (aq) → Pb(NO3 )2 (s) d) Pb2+ (aq) + 2I– (aq) → PbI2 (s)

_____ 23.

Aqueous solutions of potassium hydroxide and nitric acid are mixed. What is the net ionic equation for this reaction? a) H+ (aq) + OH– (aq) → H2 O(l) b) K+ (aq) + NO3 – (aq) → KNO3 (s) c) K+ (aq) + OH– (aq) → KOH(s) d) H+ (aq) + NO3 – (aq) → HNO3 (s)

III. Free Response Write balanced equations for the following reactions. Include physical states.

24.

Magnesium metal reacts with aqueous iron(III) chloride to produce aqueous magnesium chloride and iron metal.

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25.

Potassium metal reacts with oxygen gas to form solid potassium oxide.

Predict the products and write balanced equations for the following reactions. Include physical states. (In all cases, a reaction occurs.) Also indicate what type of reaction is occurring.

26.

Zinc metal reacts with aqueous hydrochloric acid.

27.

Butane (C4 H10 ) gas burns in oxygen.

28.

Metallic lithium reacts with nitrogen gas.

29.

Aqueous solutions of copper(II) sulfate and sodium sulfide are mixed; a solid forms. a) Write a molecular equation for this reaction. Include physical states.

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b) Write a complete ionic equation for this reaction. Include physical states.

c) Write a net ionic equation for this reaction. Include physical states.

30.

Briefly explain the difference between a) a synthesis reaction and a decomposition reaction.

b) a single replacement reaction and a double replacement reaction.

c) a nonelectrolyte and a strong electrolyte.

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31.

Both of the given equations are “balanced,” but they are balanced incorrectly. Correctly balance each one. a)

NH3

→

b)

H2

+ Br2

N

+ H3

→

H2 Br2

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Chapter 9

Chemical Quantities I. Matching Match the description in Column B with the correct term in Column A. Write the letter in the blank provided. Each term matches with only one description, so be sure to choose the best description for each term.

Column A

Column B

_____ 1. stoichiometry

A. calculation of quantities in chemical reactions

_____ 2. limiting reactant

B. reactant that remains when reaction ends

_____ 3. theoretical yield

C. reactant that runs out first

_____ 4. excess reactant

D. amount of product that is recovered

_____ 5. actual yield

E. comparison of predicted and actual amounts of product

_____ 6. percent yield

F. predicted amount of product

II. Multiple Choice Choose the one best answer and write its letter in the blank. _____ 7.

Which of the following is/are conserved in a chemical reaction? a) mass b) moles c) atoms d) both mass and atoms

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_____ 8.

The mole ratio of two components in a chemical reaction is determined from the a) coefficients of each component. b) mass of each component. c) volume of each component. d) number of atoms of each component.

_____ 9.

When doing a mass-mass stoichiometric calculation, the mass of the first component is first converted into a) liters. b) moles.

_____ 10.

c) number of atoms. d) number of molecules.

If 3.0 mol of Al(NO3 )3 are used up in the following balanced reaction, how many moles of HNO3 will be produced?

2Al(NO3 )3 (aq) + 3H2 SO4 (aq) → Al2 (SO4 )3 (s) + 6HNO3 (aq) a) 9.0 mol b) 6.0 mol

_____ 11.

c) 3.0 mol d) 2.0 mol

In a chemical reaction, the limiting reactant is the reactant that a) has the smallest mass. b) runs out first. c) has the greatest mass. d) is left over.

_____ 12.

Given the equation Mg(OH)2 (s) + 2HCl(aq) → MgCl2 (aq) + 2H2 O(l) if 33.5 g of Mg(OH)2 react, what mass of H2 O could be produced? a) 20.7 g b) 29.2 g

_____ 13.

c) 33.5 g d) 67.0 g

If 111.70 g of iron react with oxygen gas to produce iron(III) oxide, how many moles of oxygen gas will be needed? a) 0.667 mol b) 1.50 mol

c) 3.00 mol d) 6.00 mol

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_____ 14.

If 3.00 mol of ZnS are combined with 4.00 mol of O2 , according to the equation below, how many moles of ZnO can be produced?

2ZnS(s) + 3O2 (g) → 2ZnO(s) + 2SO2 (g) a) 2.00 mol b) 2.67 mol

c) 3.00 mol d) 5.67 mol

III. Free Response Answer the questions in the space provided. Use complete sentences and provide support for your answer.

15.

Potassium chlorate (KClO3 ) decomposes to form potassium chloride and oxygen gas. If 5.4 mol of potassium chlorate decompose, how many moles of oxygen could be produced?

16.

When ammonia burns in pure oxygen, the reaction is: 4NH3 (g) + 3O2 (g) → 2N2 (g) + 6H2 O(g) What masses of nitrogen and water could be produced from 45.0 g of ammonia?

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17.

Copper metal reacts with a solution of silver nitrate, AgNO3 , to produce copper(II) nitrate and silver metal. In carrying out this reaction, a piece of copper wire was immersed in a solution of silver nitrate until the reaction stopped. The original mass of the copper wire was 2.36 g. After the reaction stopped, the mass of the wire was 1.03 g. What mass of silver was produced? What type of chemical reaction is this?

18.

a) What mass of zinc could be produced from the reaction of 6.50 g of carbon and 10.0 g of zinc oxide if the reaction is the following: C(s) + 2ZnO(s) → 2Zn(s) + CO2 (g)

b) What type of chemical reaction is this?

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19.

Gastric juice (stomach acid) contains about 3.0 g of HCl per liter. If a person produces about 2.5 L of gastric juice per day, how many antacid tablets (each containing 400 mg of Al(OH)3 ) are needed to neutralize all the HCl produced in one day?

3HCl + Al(OH)3 → AlCl3 + 3H2 O

20.

Alkali metal hydroxides are sometimes used to “scrub” excess carbon dioxide from the air in closed spaces, such as submarines and spacecraft. For example, lithium hydroxide reacts with carbon dioxide according to the balanced chemical equation: 2LiOH + CO2 → Li2 CO3 + H2 O Suppose a lithium hydroxide canister contains 454 g (1 pound) of LiOH. a) What mass of carbon dioxide will the canister be able to absorb?

b) If it is found (after 24 hours) the canister has absorbed 308 g of CO2 , what percentage of its capacity has been reached?

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21.

The manufacture of a chemical compound Q requires five separate chemical reactions.

The initial reactant, compound L, is converted to compound M. Then M is converted to compound N. Compound N is then converted to compound O, and so on. Look at the diagram below to see what each step accomplishes. Provide the missing quantities. You can assume that the reactants and products all react in a 1:1 ratio.

10.0 mol Compound L

50.0% yield

5.0 mol Compound M

_____ % yield

25.0% yield

2.0 mol Compound N

_____ % yield

0.015 mol Compound Q

_____ mol Compound O

10.0% yield

_____ mol Compound P

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Chapter 10

Energy I. Matching Match the description in Column B with the correct term in Column A. Write the letter in the blank provided. Each term matches with only one description, so be sure to choose the best description for each term.

Column A

Column B

_____ 1. energy

A. energy needed to raise temperature of 1 g of a substance 1 °C

_____ 2. potential energy

B. a process that absorbs energy

_____ 3. kinetic energy

C. a measure of the dispersal of energy through a system

_____ 4. work

D. energy of motion

_____ 5. entropy

E. energy needed to raise temperature of 1 g of water 1 °C

_____ 6. temperature

F. force acting over a distance

_____ 7. heat

G. a process that releases energy

_____ 8. system

H. stored energy

_____ 9. exothermic

I. measure of the random motions of molecules and atoms of a substance

_____ 10. endothermic

J. ability to do work or produce heat

_____ 11. calorie

K. flow of energy due to temperature differences

_____ 12. specific heat capacity

L. portion of the universe in which you are interested

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II. Multiple Choice Choose the one best answer and write its letter in the blank. _____ 13.

How many calories of heat were added to 5.0 × 102 g of water to raise its temperature from 25 °C to 55 °C? c) 1.5 × 104 calories d) 6.3 × 104 calories

a) 15 calories b) 30 calories _____ 14.

The amount of heat gained or lost by a substance depends on all of the following EXCEPT a) the change in temperature of the substance. b) the change in volume of the substance. c) the mass of the substance. d) the specific heat capacity of the substance.

_____ 15.

Which of the following has the greatest entropy? a) 100 g of ice at 0 °C b) 100 g of water at 25 °C

_____ 16.

One _____ of water requires 4.184 J of heat to cause a change in temperature of 1 °C. a) molecule

_____ 17.

b) gram

c) mole

d) kilogram

A negative sign for the change in enthalpy for a reaction indicates that the reaction is a) fast.

_____ 18.

c) 100 g of water at 50 °C d) 100 g of steam at 100 °C

b) slow.

c) exothermic.

d) endothermic.

How much heat is evolved if 25.00 g of methanol (CH3 OH) is burned in excess O2 ? The reaction is the following: 2CH3 OH(l) + 3O2 (g) → 2CO2 (g) + 4H2 O(g) a) 566.9 kJ b) 1134 kJ

DH = –1453 kJ

c) 2906 kJ d) 3.634 × 104 kJ

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_____ 19.

Consider the equation PCl5 (g) → PCl3 (g) + Cl2 (g). Which of the following is true? a) Energy is released, and a bond is formed. b) Energy is absorbed, and a bond is formed. c) Energy is released, and a bond is broken. d) Energy is absorbed, and a bond is broken.

_____ 20.

Given the following data: C(s) + O2 (g) → CO2 (g)

DH = –393.5 kJ

2CO(g) + O2 (g) → 2CO2 (g)

DH = –566.0 kJ 1

calculate DH for the reaction C(s) + 2 O2 (g) → CO(g). a) –110.5 kJ _____ 21.

b) –676.5 kJ

c) –959.5 kJ

d) +172.5 kJ

Which of the following describes an endothermic reaction taking place in a glass beaker? a) The beaker feels cold to the touch, as the reaction releases energy. b) The beaker feels hot to the touch, as the reaction releases energy. c) The beaker feels cold to the touch, as the reaction absorbs energy. d) The beaker feels hot to the touch, as the reaction absorbs energy.

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III. Free Response Answer the questions in the space provided. Show all work for calculations.

22.

A chemical reaction takes place. As a result of this reaction, the temperature of the reaction container decreases. a) Is this reaction exothermic or endothermic? Explain your choice.

b) Do the reactants or the products of the reaction have the higher potential energy content? Explain your choice.

23.

How much energy as heat is required to raise the temperature of 38.0 g of water from 25.0 °C to 95.0 °C?

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24.

If a piece of aluminum of mass 4.50 g and temperature 99.5 °C is dropped into 12.0 g of water at 21.0 °C, what will be the final temperature of the water-aluminum mixture? The specific heat capacity of aluminum is 0.902 J/(g•°C)

25.

Given the following data: C(s) + O2 (g) → CO2 (g)

DH = –393.5 kJ

C2 H4 (g) + 3O2 (g) → 2CO2 (g) + 2H2 O(l)

DH = –1410.9 kJ

H2 (g) +

1 2

DH = –285.8 kJ

O2 (g) → H2 O(l)

calculate DH for the reaction 2C(s) + 2H2 (g) → C2 H4 (g).

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26.

Consider the following three thermochemical equations: DH = + 97.8 kJ/mol

4A → 2E B +

1

D → 2

1

C + 2

12A → 4D + 8C

1 2

E

DH = –224.5 kJ/mol DH = +101.1 kJ/mol 1

Now suppose that the first equation is reversed and multiplied by 2 , the second equation is doubled, the third equation is divided by four, and the three adjusted equations are added. What is the net reaction and what is the overall heat of reaction?

27.

Explain briefly why energy must always be absorbed by a molecule in order to break a bond.

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Chapter 11

Modern Atomic Theory I. Matching Match the description in Column B with the correct term in Column A. Write the letter in the blank provided. Each term matches with only one description, so be sure to choose the best description for each term.

Column A

Column B

_____ 1. wavelength

A. region where an electron resides in an atom

_____ 2. frequency

B. energy required to remove an electron from an atom in the gas phase

_____ 3. photon

C. an atomic orbital can hold no more than two electrons, and they must be of opposite spin

_____ 4. excited state

D. tiny packet of energy

_____ 5. ground state

E. inner electrons of an atom

_____ 6. orbital

F. distance between two waves

_____ 7. Pauli Exclusion Principle

G. contains an electron with excess energy

_____ 8. valence electrons

H. electrons in outermost (highest) occupied principal energy level

_____ 9. core electrons

I. lowest possible energy state for an atom

_____ 10. ionization energy

J. the number of waves that pass a given point per unit of time

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II. Multiple Choice Choose the one best answer and write its letter in the blank. _____ 11.

What is the maximum number of electrons that can be in the third principal energy level? a) 2

_____ 12.

c) 18

b) 3s2 3p6

c) 3s2 3p1

d) 4s2 4p1

Which of the following groups of atoms have the same outermost electron configurations but with different n values? a) N, O, F, Ne b) S, Cl, Ar, K

_____ 14.

d) 32

The configuration for the outermost energy level in Al is what? a) 2s2 2p1

_____ 13.

b) 8

c) Ca, Ge, Sr, In d) O, S, Se, Te

When an electron in an atom absorbs energy, it a) moves to a higher energy level. b) emits light. c) moves to its ground state. d) does none of the above.

_____ 15.

Which of the following is the smallest atom? a) K

_____ 16.

d) O

b) Ge

c) Si

d) O

The element whose electron configuration is 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p2 is a) Ge.

_____ 18.

c) Si

Which of the following has the highest ionization energy? a) K

_____ 17.

b) Ge

b) Ga.

c) Si.

d) Al.

As the wavelength of light increases, its frequency a) increases. b) decreases.

c) does not change. d) oscillates.

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_____ 19.

Which neutral atom has six 3d electrons? a) Fe

_____ 20.

c) Ru

b) 5

c) 2

_____ 24.

c) from n = 4 to n = 1 d) from n = 1 to n = 5

Which one of the following elements has five valence electrons? a) B

_____ 23.

d) 0

In a hydrogen atom, an electron undergoing which of the following transitions would emit light of the shortest wavelength? a) from n = 5 to n = 1 b) from n = 5 to n = 2

_____ 22.

d) Rh

How many electrons can be in a 3d sublevel? a) 10

_____ 21.

b) Co

b) C

c) N

d) F

What is the electron configuration for bromine?

2 a) [Ar]4s 3d10 4p5

c) [Ar]4s2 4p5

b) [Kr]4s2 3d10 4p5

1 d) [Se]4p

Which of these statements about electron orbitals is not true? a) Orbitals show the path of the electron. b) Orbitals are distributed in space around the nucleus. c) Orbitals may contain more than one electron. d) Orbitals are not all the same shape.

III. Free Response Answer the questions in the space provided.

For questions 25–30, identify the elements described. Choose the one element that best fits the description. _____ 25.

has an electron configuration of 1s2 2s2 2p6 3s2 3p4

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_____ 26.

contains a full second energy level as its valence shell

_____ 27.

contains only two electrons in its fifth energy level

_____ 28.

has a single electron in 7s

_____ 29.

contains three unpaired electrons in its third energy level

_____ 30.

contains one 3d electron

31.

Using noble gas shorthand, write the electron configuration for astatine (At).

32.

Write electron configurations for each of the following. DO NOT use noble gas shorthand. a) Al

b) Fe

33.

Using atomic structure in your explanation, account for the general trend in atomic size as you go from left to right across a period and from top to bottom down a group on the periodic table.

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34.

Give the symbol of the first element on the periodic table which has a) four unpaired electrons.

__________

b) a half-completed set of d orbitals.

__________

c) three valence electrons.

__________

d) an electron in the fourth energy level.

__________

e) a p electron.

__________

35.

Why is the 1s orbital the first orbital to be filled by an electron in any atom in the ground state?

36.

a) How would the color of light given off by an electron jumping from level 6 to level 1 compare to that given off by an electron jumping from level 2 to level 1?

Explain briefly.

b) How would the color of light compare if many electrons jumped from level 2 to level 1 as opposed to just one electron making that same jump?

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Chapter 12

Chemical Bonding I. Matching Match the description in Column B with the correct term in Column A. Write the letter in the blank provided. Each term matches with only one description, so be sure to choose the best description for each term.

Column A

Column B

_____ 1. bond

A. four electron pairs (four bonds) on central atom

_____ 2. polar covalent bond

B. more than one valid Lewis structure may be drawn for a molecule

_____ 3. covalent bond

C. electrons unequally shared by two atoms

_____ 4. ionic compound

D. two electron pairs (two bonds) on central atom

_____ 5. electronegativity

E. three single bonds and a lone pair on central atom

_____ 6. resonance

F. ability of an atom to attract shared electrons to itself

_____ 7. linear structure

G. force that holds atoms together

_____ 8. trigonal planar structure

H. three electron pairs (three bonds) on central atom

_____ 9. tetrahedral structure

I. formed by a metal and a nonmetal

_____ 10. trigonal pyramid structure

J. electrons shared by two atoms

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II. Multiple Choice Choose the one best answer and write its letter in the blank. _____ 11.

Which of the following atoms becomes larger when it forms its most common ion? a) O

_____ 12.

b) Fe

c) Na

To attain a noble gas configuration, an atom of sulfur must a) gain two electrons. b) lose two electrons.

_____ 13.

d) Ca

c) gain one electron. d) lose one electron.

Which of these is NOT a characteristic of an ionic compound? a) is a solid b) has a low melting point c) is composed of metallic and nonmetallic elements d) will conduct electricity when melted

_____ 14.

An ionic bond is formed a) between two nonmetals. b) when electrons are transferred from one atom to another. c) when electrons are shared between atoms. d) when an element changes into a noble gas.

_____ 15.

Which of the following molecules has a single pair of shared electrons? a) O2

_____ 16.

c) Cl2

d) N2

Which of the following molecules has a dipole moment? a) O2

_____ 17.

b) CO

b) N2

c) CO2

d) HCl

Which of the following molecules exhibits resonance? a) NH3

b) SO3

c) H2 S

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d) CCl4

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_____ 18.

Which of the following does NOT have covalent bonds? a) H2 O

_____ 19.

b) SO2

c) K2 O

The shape of the NF3 molecule is a) linear. b) bent.

_____ 20.

c) tetrahedral. d) trigonal pyramid.

An example of a nonpolar molecule would be a) CO.

_____ 21.

d) PCl3

b) H2 O.

c) Br2 .

d) HF.

Which one of the following statements about solid sodium chloride is INCORRECT? a) Within the crystal, each sodium ion is surrounded by chloride ions. b) Each sodium ion is attracted to chloride ions because of their opposite charges. c) Each sodium ion is tightly bonded to only one chloride ion, forming molecules of NaCl, the “building blocks” of the crystal. d) The bonding in sodium chloride crystals is strong.

_____ 22.

Consider the three compounds NH3 , PH3 , and AsH3 . Which of the following statements about these compounds is INCORRECT? a) Each of the three molecules has a lone pair of electrons. b) Each of the three molecules is trigonal planar. c) Each of the molecules is polar. d) Each of the molecules has covalent bonds.

III. Free Response Answer the questions in the space provided.

23.

Write the complete electron configuration for the simple ion that each of the following atoms forms and indicate the element that has the same configuration as the ion. a) Ca: _________________________________

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24.

b) Al: _________________________________

___________

c) P: __________________________________

___________

Draw Lewis structures for each of the following molecules or ions. Predict the molecular structure for each.

Lewis Structure

Molecular Structure

a) NI3

b) CS2

c) COCl2 (C is the central atom)

25.

Draw Lewis structures for the carbonate ion, CO3 2– , showing all possible resonance structures.

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26.

CO2 is an example of a molecule that is nonpolar, even though it is comprised of two polar bonds. Explain briefly how this can be.

27.

Explain briefly why neon (Ne) is monatomic, but chlorine (Cl) exists as diatomic molecules.

28.

Explain briefly why the shape of aluminum hydride, AlH3 , is trigonal planar, but the shape of ammonia, NH3 , is trigonal pyramidal.

29.

Briefly explain the difference between a polar covalent bond and an ionic bond.

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Chapter 13

Gases I. Matching Match the description in Column B with the correct term in Column A. Write the letter in the blank provided. Each term matches with only one description, so be sure to choose the best description for each term.

Column A

Column B

_____ 1. absolute zero

A. at constant temperature, the volume of a given amount of gas is inversely proportional to its pressure

_____ 2. universal gas constant

B. the total pressure of a mixture of gases is the sum of the pressures that each gas would exert if alone in the container

_____ 3. Boyle’s law

C. temperature at which a gas theoretically occupies zero volume

_____ 4. Charles’s law

D. P1 V1 = P2 V2 T1 T2

_____ 5. Avogadro’s law

E. at constant temperature and pressure, the volume of a gas is directly proportional to the number of moles of gas

_____ 6. combined gas law

F.

_____ 7. Dalton’s law of partial pressures

G. at constant pressure, the volume of a given amount of gas is directly proportional to its Kelvin temperature

_____ 8. molar volume of ideal gas at STP

22.4 L

H. 0.08206

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L atm mol K

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II. Multiple Choice Choose the one best answer and write its letter in the blank. _____ 9.

If the temperature is held constant, what will happen to the pressure exerted by a gas if the volume of the container is decreased? a) increase b) decrease

_____ 10.

c) remain the same d) increase or decrease (it depends)

Which of the following is NOT true about pressure? a) Pressure is exerted equally in all directions. b) Pressure is defined as force per unit area. c) Gas pressure can be measured only in units of mm of mercury. d) Atmospheric pressure decreases with increasing altitude.

_____ 11.

The pressure of a sample of gas is 500. mm Hg, and the volume is 30.0 L. If the volume is changed to 50.0 L, what is the new pressure? c) 7.50 × 103 mm Hg d) 3.00 × 103 mm Hg

a) 833 mm Hg b) 300. mm Hg _____ 12.

If the temperature of a sample of gas is decreased at constant pressure, the volume of the gas will a) increase. b) decrease.

_____ 13.

c) remain the same. d) increase or decrease (it depends).

Standard temperature and pressure, STP, refers to a) 100 °C and zero atmosphere. b) absolute zero and one atmosphere. c) 0 °C and a pressure of 760 mm Hg. d) 0 °C and zero atmosphere.

_____ 14.

A sample of gas in a rigid container (constant volume) is at a temperature of 25.0 °C. If the temperature of the gas is increased to 50.0 °C, what will happen to the pressure exerted by the gas? a) It will double. b) It will halve.

c) It will increase by a small amount. d) It will decrease by a small amount.

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_____ 15.

You have 11.2 L of an ideal gas at STP. The number of molecules of gas in this sample is a) 3.01 × 1023 . b) 6.02 × 1023 .

_____ 16.

Consider the following: You have 2 L of neon gas at a pressure of 2 atm, 2 L of carbon dioxide gas at a pressure of 3 atm, and 2 L of nitrogen gas at a pressure of 4 atm. All three samples are at room temperature. If you transfer all three gases to the same rigid 2 L container, what is the pressure exerted by the nitrogen in the final mixture? a) 2 atm

_____ 17.

b) 3 atm

c) 4 atm

b) 3 atm

c) 4 atm

c) 22.5 mol d) 45.0 mol

The volume of a sample of gas at STP is 9.75 L. What volume will this gas occupy at a temperature of 100 °C and a pressure of 2.50 atm? a) 19.5 L b) 16.4 L

_____ 20.

d) 9 atm

Determine the number of moles of gas in 45.0 L of oxygen at STP. a) 1.61 mol b) 2.01 mol

_____ 19.

d) 9 atm

In question 16, what is the total pressure exerted by the gas mixture? a) 2 atm

_____ 18.

c) 11.2. d) 22.4.

c) 13.3 L d) 5.33 L

What volume of carbon dioxide will be produced from the combustion of 4.50 mol of propane (C3 H8 ) at STP? C3 H8 (g) + 5O2 (g) → 3CO2 (g) + 4H2 O(l) a) 101 L

b) 302 L

c) 403 L

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d) 504 L

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III. Free Response Answer the questions in the space provided. Show your work.

21.

What volume does 18.9 g of CO2 occupy at 57.5 °C and 4.35 atm?

22.

Suppose 3.50 g of CO2 and 11.6 g of O2 are placed in a 3.00 L container at 273 K. What will the total pressure in the container be?

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23.

Calcium carbonate decomposes when heated to form calcium oxide and carbon dioxide.

What volume of carbon dioxide can be produced from 37.0 g of calcium carbonate at a temperature of 25 °C and a pressure of 1.15 atm?

24.

A common piece of lab equipment is the 125 mL Erlenmeyer flask. How many moles of air (whose average molecular weight is 29 g/mol) would actually fit into that flask at 20 °C and 750 torr? How many molecules of air does this represent?

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25.

Suppose that you have a mole of oxygen gas trapped in an expandable container. What effect would heating the gas sample have on each of these changes? Explain your answers briefly. a) the number of oxygen molecules inside

b) the average velocity of the molecules

c) the volume of the gas sample

d) the mass of the gas sample

e) the density of the gas sample

f) the average distance between the gas molecules

g) the pressure inside the container

26.

Explain briefly why gases are much less dense than solids.

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Chapter 14

Liquids and Solids I. Matching Match the description in Column B with the correct term in Column A. Write the letter in the blank provided. Each term matches with only one description, so be sure to choose the best description for each term.

Column A

Column B

_____ 1. intermolecular forces

A. transition from liquid to vapor phase

_____ 2. intramolecular forces

B. forces between nonpolar molecules

_____ 3. dipole–dipole attraction

C. pressure of a vapor in equilibrium with a sample of liquid

_____ 4. hydrogen bonding

D. transition from vapor to liquid phase

_____ 5. London dispersion forces

E. general name for forces between molecules

_____ 6. molar heat of fusion

F. solid with regular arrangement of components

_____ 7. condensation

G. chemical bonds between the atoms of a molecule

_____ 8. evaporation

H. attraction between polar molecules

_____ 9. vapor pressure

I. energy required to melt 1 mole of a substance

_____ 10. crystalline solid

J. strong dipole–dipole attraction including a hydrogen atom

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II. Multiple Choice Choose the one best answer and write its letter in the blank. _____ 11.

Which of the following substances would be expected to have the highest melting point? a) NaCl

_____ 12.

c) CH3 OH

b) 16.7 kJ

c) 12.0 kJ

c) remain the same. d) depend on the identity of the liquid.

The escape of vapor molecules with sufficiently high energy from the surface of a liquid is called a) condensation. b) evaporation.

_____ 15.

d) 5.02 kJ

If you increase the pressure above a liquid, its boiling point will a) decrease. b) increase.

_____ 14.

d) Ne

How much heat is required to melt 50.0 g of ice at 0 °C? (The molar heat of fusion for water is 6.02 kJ/mol.) a) 301 kJ

_____ 13.

b) H2 O

c) boiling. d) sublimation.

Which of the following statements is INCORRECT? a) The boiling point is the temperature at which that vapor pressure of the liquid equals the atmospheric pressure. b) Once a substance has begun to boil, the temperature of the liquid will remain constant. c) Evaporation is the same thing as boiling. d) If the pressure is changed, a substance can be made to boil at a different temperature than its normal boiling point.

_____ 16.

Which substance has the highest vapor pressure? a) CH4

b) KCl

c) H2 S

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d) H2 O

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Questions 17–21 refer to the following graph of a heating curve:

E D

Temp (°C)

C B A

Time (min) Which region on the curve contains _____ 17.

solid and liquid together?

_____ 18.

gas only?

_____ 19.

solid only?

_____ 20.

liquid and gas together?

_____ 21.

liquid only?

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III. Free Response Answer the questions in the space provided. Use complete sentences and provide support for your answer. Show your work for any calculations.

22.

Explain, using principles of intermolecular forces and bonding, why an ionic solid will shatter when struck with a hammer while metals will merely flatten out when struck.

23.

How much heat is necessary to completely vaporize 75.0 g of ice at 0 °C? The molar heat of fusion for water is 6.02 kJ/mol. The specific heat capacity of liquid water is 4.18 J/g•°C. The molar heat of vaporization for water is 40.6 kJ/mol.

24.

Briefly identify the difference between a) melting point and freezing point.

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b)

melting point and molar heat of melting (fusion).

25.

What aspect of a water molecule causes its bonds to be polar? What aspect of a water molecule causes it to have molecular polarity? What aspect of water causes it to exhibit hydrogen bonding?

26.

Briefly explain why different liquids have different boiling points.

27.

What are the strongest bonds that must be broken when a) ice melts?

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b) Cl2 (l) vaporizes?

c) H2 O is separated into hydrogen and oxygen gases?

28.

When 25.23 g of methanol (CH3 OH) freezes, 4.01 kJ of heat is released. What is the molar enthalpy of fusion for methanol?

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Chapter 15

Solutions I. Matching Match the description in Column B with the correct term in Column A. Write the letter in the blank provided. Each term matches with only one description, so be sure to choose the best description for each term.

Column A

Column B

_____ 1. solution

A. amount needed to furnish 1 mol of H+ ions

_____ 2. solvent

B. when a relatively large amount of solute is dissolved

_____ 3. solute

C. solution containing more solute than a saturated solution

_____ 4. saturated solution

D. homogeneous mixture

_____ 5. unsaturated solution

E. component of solution present in lesser amount

_____ 6. supersaturated solution

F. number of moles of solute per liter of solution

_____ 7. concentrated solution

G. less than maximum possible amount of solute is dissolved in solution

_____ 8. dilute solution

H. acid–base reaction

_____ 9. molarity

I. amount needed to furnish 1 mol of OH– ions

_____ 10. neutralization reaction

J. component of solution present in largest amount

_____ 11. one equivalent of an acid

K. when the maximum possible amount of solute is dissolved in solution

_____ 12. one equivalent of a base

L. depends on concentration of solute particles

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_____ 13. colligative property

M. when a relatively small amount of solute is dissolved

II. Multiple Choice Choose the one best answer and write its letter in the blank. _____ 14.

Saturated solutions can be a) concentrated. b) dilute.

_____ 15.

Oils will dissolve best in solvents that are a) polar. b) nonpolar.

_____ 16.

c) 3.5 L of solution. d) 3.5 L of solvent.

The number of moles of solute in 200. mL of a 0.500 M solution is a) 400. mol. b) 100. mol.

_____ 18.

c) either polar or nonpolar. d) neither polar nor nonpolar.

A 3.5 M KNO3 solution contains 3.5 mol of KNO3 dissolved in a) 1.0 L of solution. b) 1.0 L of solvent.

_____ 17.

c) either concentrated or dilute. d) neither dilute nor concentrated.

c) 2.50 mol. d) 0.100 mol.

Adding salt to water before cooking pasta means that the pasta will a) cook at a temperature higher than the normal boiling point of water. b) cook at a temperature equal to the normal boiling point of water. c) cook at a temperature lower than the normal boiling point of water. d) stick together after cooking.

_____ 19.

Dissolving a solute in a solution will a) have no effect on the vapor pressure. b) increase the vapor pressure. c) decrease the vapor pressure. d) will increase the vapor pressure if the solute is ionic and will decrease the vapor pressure it the solute is a nonelectrolyte.

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_____ 20.

A 500. mL solution of a 0.15 M NaCl solution contains a) 75 g of NaCl. b) 0.075 g of NaCl.

_____ 21.

25.0 mL of a 12.0 M HCl stock solution is diluted to a new volume of 1.00 L. The concentration of this new solution is a) 11.39 M. b) 3.00 M.

_____ 22.

c) 1.56 M. d) 0.300 M.

200. mL of a 0.50 M NaOH solution requires 100.0 mL of a HCl solution to neutralize it. What is the concentration of the HCl solution? a) 2.0 M b) 1.0 M

_____ 23.

c) 1.5 g of NaCl. d) 4.4 g of NaCl.

c) 0.50 M d) 0.10 M

What volume of 0.100 M Ca(OH)2 is required to neutralize 100.0 mL of a 0.200 M HCl solution? a) 50.0 mL b) 100. mL

c) 200. mL d) 300. mL

III. Free Response Answer the questions in the space provided. Show your work for any calculations.

24.

List three ways to increase the rate that a solid dissolves in water, and describe why each one speeds up the rate of dissolving. Use complete sentences, and provide support for your answer.

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25.

How much potassium iodide (KI), in grams, should be added to water to produce 2.5 L of solution with a concentration of 0.35 M?

26.

What is the molarity of a solution in which 75.0 g of acetic acid (HC2 H3 O2 ) is dissolved in enough water to make 1500 mL of solution?

27.

50.0 mL of 0.15 M AgNO3 solution is poured into 100.0 mL of 0.10 M MgCl2 solution. a) Write a net ionic equation to show the precipitation of solid AgCl.

b) What mass of solid AgCl will be formed?

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28.

Suppose that you prepare two solutions in the laboratory: a 0.5 M solution of the strong acid HCl and a 0.5 M solution of the strong acid H2 SO4 . Briefly explain why their acid strength is not the same even though they have identical molarities.

29.

How much ethyl alcohol (C2 H5 OH), in both grams and moles, are needed to make 2.45 kg of a 12.0% solution?

30.

Sometimes water is called “the universal solvent.” Briefly explain why you do or do not agree with this description.

31.

A sugar syrup solution contains 20.0% fructose (C6 H12 O6 ) by mass. The density of the solution is 1.22 g/mL. a) How much would 1.0 L of this solution weigh?

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b) How many grams of fructose would be in that one liter of solution?

c) How many grams of water would be present in that same volume?

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Chapter 16

Acids and Bases I. Matching Match the description in Column B with the correct term in Column A. Write the letter in the blank provided. Each term matches with only one description, so be sure to choose the best description for each term.

Column A

Column B

_____ 1. Arrhenius acid

A. can act as an acid or a base

_____ 2. Arrhenius base

B. [H+ ] = [OH– ]

_____ 3. weak acid

C. produces OH– ions in aqueous solution

_____ 4. Brønsted-Lowry base

D. completely dissociates in water

_____ 5. strong acid

E. [OH– ] > [H+ ]

_____ 6. amphoteric

F. changes color in acidic or basic solution

_____ 7. neutral solution

G. [H+ ] > [OH– ]

_____ 8. acidic solution

H. exactly enough titrant has been added to react with all the unknown present

_____ 9. basic solution

I. partially dissociates in water

_____ 10. indicator

J. a weak acid and its conjugate base are both present in solution

_____ 11. equivalence point

K. proton (H+ ) acceptor

_____ 12. buffered solution

L. produces H+ ions in aqueous solution

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II. Multiple Choice Choose the one best answer and write its letter in the blank. _____ 13.

Which one of the following 1 M solutions would have a pH of 14? a) NaOH

_____ 14.

b) HCl

c) HF

When the pH of a solution decreases, the [H+ ]∕[OH– ] ratio a) decreases. b) increases.

_____ 15.

d) H2 O

c) stays the same. d) can increase or decrease.

Which of the following statements is INCORRECT? a) An acid may be both strong and dilute. b) No substance may be both a strong acid and a strong base. c) Water is not an amphoteric substance. d) The ion product constant for water is 1 × 10–14 at 25 °C.

_____ 16.

pH is a measure of a) the hydrogen ion concentration. b) the hydrogen gas concentration. c) the hydroxide ion concentration. d) hydrolysis.

_____ 17.

Which of the following is NOT a conjugate acid/conjugate base pair? a) HCl∕Cl– b) H2 O∕O2–

_____ 18.

c) HC2 H3 O2 ∕C2 H3 O2 – d) NH4 + ∕NH3

Calculate [H+ ] in a solution if [OH– ] = 4.20 × 10–8 M at 25 °C and state whether the solution is acidic, basic, or neutral. a) [H+ ] = 2.38 × 10–7 M; basic b) [H+ ] = 7.37 × 10–7 M; basic

c) [H+ ] = 2.38 × 10–7 M,;acidic d) [H+ ] = 4.20 × 10–7 M; neutral

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_____ 19.

If the pOH of a solution is 5.25, what is [H+ ]? c) 5.6 × 10–6 M d) 1.8 × 10–9 M

a) 8.75 M b) 5.25 M _____ 20.

What is the pH of a 0.015 M HCl solution? a) 1.50

_____ 21.

b) 1.82

c) 12.18

Calculate the pH of a solution that has [H+ ] = 2.5 × 10–10 M, and state whether the solution is acidic, basic, or neutral. a) 4.40; acidic b) 9.60; acidic

_____ 22.

d) 12.50

c) 4.40; basic d) 9.60; basic

A neutralization reaction between an acid and a base always produces a) salt and acidic anhydride. b) salt and basic anhydride.

c) water and basic anhydride. d) salt and water.

III. Free Response Answer the questions in the space provided. Show your work for any calculations.

23.

Write a chemical equation that shows how each of the following species behaves as an acid when dissolved in water. a) HCN

b) NH4 +

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c) H2 PO4 –

24.

Calculate [H+ ] for each of the following solutions and state whether the solution is acidic, basic, or neutral. a) [OH– ] = 7.26 × 10–2 M

b) [OH– ] = 3.90 × 10–11 M

25.

Calculate the hydrogen ion concentration, in moles per liter, for solutions with each of the following pH or pOH values and state whether the solution is acidic, basic, or neutral. a) pH = 6.44

b) pH = 12.75

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26.

If 50.0 mL of a HCl solution requires 20.0 mL of 0.20 M NaOH to titrate it to the equivalence point, what is the concentration, in moles per liter, of the HCl solution?

27.

What is the difference between a) an Arrhenius base and a Brønsted-Lowry base?

b) pH and [H+ ]?

28.

Complete the following Brønsted-Lowry equations. In each case, circle the conjugate base. a) NH4 + + Cl– →

b) HCO3 – + H2 PO4 – →

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29.

A basic solution is prepared by dissolving 2.55 g of sodium hydroxide pellets (NaOH) in enough water to make 550.0 mL of solution. What is the pH of the resulting solution?

30.

Compare and contrast a dilute solution of a strong acid with a concentrated solution of a weak acid.

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Chapter 17

Equilibrium I. Matching Match the description in Column B with the correct term in Column A. Write the letter in the blank provided. Each term matches with only one description, so be sure to choose the best description for each term.

Column A

Column B

_____ 1. activation energy

A. substances in a system are in more than one state

_____ 2. catalyst

B. dynamic state where concentrations of all reactants and products remain constant

_____ 3. enzyme

C. all substances in a system are in the same state

_____ 4. equilibrium

D. equilibrium expression for a solid dissolving in water

_____ 5. chemical equilibrium

E. minimum amount of energy required for a reaction to occur

_____ 6. equilibrium expression

F. exact balancing of two processes, one of which is the opposite of the other

_____ 7. solubility product

G. biological catalyst

_____ 8. homogeneous equilibrium

H. when a change is imposed on a system at equilibrium, the position of the equilibrium shifts in the direction that reduces the effect of the change

_____ 9. heterogeneous equilibrium

I. substance that speeds up a reaction without being consumed

_____ 10. Le Châtelier’s principle

J. special ratio of concentrations of products to concentrations of reactants

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II. Multiple Choice Choose the one best answer and write its letter in the blank. _____ 11.

Increasing the surface area of reactants generally a) increases the rate of a reaction. b) decreases the rate of a reaction. c) stops a reaction entirely. d) has no effect on the rate of a reaction.

_____ 12.

Adding a catalyst generally speeds up the rate of a reaction by a) raising the activation energy. b) lowering the activation energy. c) raising the heat of the reaction. d) lowering the heat of the reaction.

_____ 13.

Biochemical changes that occur in the human body are assisted by catalysts called a) pharmaceuticals. b) inhibitors.

_____ 14.

c) enzymes. d) cooperators.

What would be the equilibrium expression for the following reaction?

4HCl(g) + O2 (g) ⇌ 2Cl2 (g) + 2H2 O(g)

_____ 15.

a)

K=

[4HCl][O2 ] [2Cl2 ][2H2 O]

c)

K=

[Cl2 ]2 [H2 O]2 [HCl]4 [O2 ]

b)

K=

[HCl]4 [O2 ] [Cl2 ]2 [H2 O]2

d)

K=

[2Cl2 ][2H2 O] [4HCl][O2 ]

From the following statements, choose which is(are) true.

I. At equilibrium, the concentrations of reactants and products are equal.

II. At equilibrium, the concentrations of the reactants and products do not change over time.

III. At equilibrium, the rates of the forward and reverse reactions are equal.

IV. At equilibrium, the chemical reaction has stopped. a) I only b) III only

c) I, II, III, and IV d) II and III

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_____ 16.

If K = 4.5 × 10–11 for a reaction, which of the following would be true? a) Reactants are favored at equilibrium. b) Products are favored at equilibrium. c) Reactants and products are equally favored at equilibrium. d) The reaction proceeds very rapidly.

_____ 17.

For the system SO2 (g) + Cl2 (g) ⇌ SO2 Cl2 (g) at equilibrium, adding Cl2 (g) to the reaction container will have which of the following effects? a) The reaction will shift to the right. b) The concentration of SO2 Cl2 (g) will increase. c) The concentration of SO2 (g) will decrease. d) All of the above will occur.

_____ 18.

If the equilibrium concentrations of A(g) + 2B(g) ⇌ C(g) are found to be [A] = 0.015 M, [B] = 2.0 × 10–4 M, and [C] = 3.0 × 10–9 M, what would the value of the equilibrium constant be? a) 5.0

_____ 19.

c) 1.0 × 10–3

b) 0.20

For the endothermic reaction A(g) + 2B(g) ⇌ C(g), which of the following will drive the reaction towards the products? a) lowering the temperature b) decreasing the volume

_____ 20.

d) 9.0 × 10–15

c) adding more C(g) d) removing some A(g)

What would be the equilibrium expression for the following reaction?

4CuO(s) ⇌ 2Cu2 O(s) + O2 (g) a)

K=

[Cu2 O][O2 ] [CuO]

c)

K=

b)

K=

[Cu2 O]2 [O2 ] [CuO]4

d)

K = [O2 ]

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[CuO]4 [Cu2 O][O2 ]4

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III. Free Response Answer the questions in the space provided. Show your work for any calculations.

21.

Write equilibrium expressions for the following reactions: a) 2SO2 (g) + O2 (g) ⇌ 2SO3 (g)

b) Al(OH)3 (s) ⇌ Al3+ (aq) + 3OH– (aq)

22.

Calculate the equilibrium constant for the following reaction if at equilibrium the concentrations of the reactants and products are [A] = 0.025 M, [B] = 0.450 M, [C] = 1.2 M, and [D] = 0.30 M. 2A(g) + B(g) ⇌ C(g) + 4D(g)

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23.

Ksp for PbI2 is 7.1 × 10–9 at 25 °C. Calculate the solubility of PbI2 in grams/liter at this temperature.

24.

Consider this reaction equation: A2 B(g) ⇌ 2A(g) + B(g). If you start with 1.2 M A2 B, and if at equilibrium, 50% of it has decomposed, what are the equilibrium concentrations of A and B? What is the value of the equilibrium constant?

25.

Consider this reaction: S8 (g) + 12O2 (g) ⇌ 8SO3 (g) + 808.8 kcal. Predict, using arrows, in which direction the equilibrium will shift when the system is disturbed by each of the following: a) adding oxygen gas b) cooling the reaction vessel c) increasing the size of the container d) removing sulfur trioxide as it is formed e) adding a catalyst to make the reaction happen faster

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Chapter 18

Oxidation–Reduction Reactions and Electrochemistry I. Matching Match the description in Column B with the correct term in Column A. Write the letter in the blank provided. Each term matches with only one description, so be sure to choose the best description for each term.

Column A

Column B

_____ 1. oxidation–reduction reaction

A. electrode where reduction occurs

_____ 2. oxidation

B. reactant containing an element that is reduced

_____ 3. reduction

C. battery

_____ 4. oxidizing agent

D. naturally occurring oxidation of metals

_____ 5. reducing agent

E. reaction in which electrons are transferred

_____ 6. electrochemistry

F. loss of electrons

_____ 7. galvanic cell

G. process where electrical energy is used to produce a chemical change

_____ 8. cathode

H. study of the interchange of chemical and electrical energy

_____ 9. electrolysis

I. gain of electrons

_____ 10. corrosion

J. reactant containing an element that is oxidized

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II. Multiple Choice Choose the one best answer and write its letter in the blank. _____ 11.

Which among the following is a reduction half-reaction? a) Cu(s) → Cu2+ (aq) + 2e– b) 2F – (aq) → F2 (g) + 2e– c) Ag+ (aq) + e– → Ag(s) d) V2+ (aq) → V3+ (aq) + e–

_____ 12.

In which of the following does nitrogen have an oxidation state of +4? a) NH3

_____ 13.

c) NO2

c) Na+

b) O2

d) O2–

For the oxidation–reduction reaction 2K(s) + Cl2 (g) → 2KCl(s), the K is a) the oxidizing agent. b) the reducing agent.

_____ 15.

d) HNO3

For the oxidation–reduction reaction 4Na(s) + O2 (g) → 2Na2 O(s), which species is oxidized? a) Na

_____ 14.

b) NH4 Cl

c) the electron acceptor. d) reduced.

Which of the following is NOT an oxidation–reduction reaction? a) CH4 (g) + 2O2 (g) → 2H2 O(g) + CO2 (g) b) AgNO3 (aq) + KCl(aq) → AgCl(s) + KNO3 (aq) c) 2Al(s) + 3CuCl2 (aq) → 3Cu(s) + 2AlCl3 (aq) d) N2 (g) + 3H2 (g) → 2NH3 (g)

_____ 16.

Which of the following is NOT true? a) Oxidation and reduction must accompany each other in a reaction. b) Oxidation and reduction involve the loss and gain of electrons, respectively. c) Oxidation and reduction involve a change in the oxidation states of the elements involved. d) Oxidation and reduction occur in all chemical reactions.

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_____ 17.

When the oxidation–reduction reaction below is balanced using the half-reaction method, what is the coefficient for Fe? Fe3+ (aq) + Br– (aq) → Fe2+ (aq) + Br2 (aq) a) 1

_____ 18.

b) 2

c) 3

d) 4

Which of the following is true for the reaction below?

Zn(s) + 2HNO3 (aq) → Zn(NO3 )2 (aq) + H2 (g) a) Zn is reduced. b) H is reduced.

c) N is reduced. d) O is reduced.

Questions 19–20 refer to the following oxidation–reduction reaction, which takes place in acidic solution: MnO4 – (aq) + Br– (aq) → MnO2 (s) + BrO3 – (aq) _____ 19.

What is the oxidation state of manganese in MnO4 – ? a) +1

_____ 20.

c) +4

b) 3

c) 6

c) chemical to mechanical d) mechanical to chemical

The component in an electrochemical cell that allows for ion migration is the a) wire. b) cathode.

_____ 23.

d) 7

What is the energy conversion that takes place in a galvanic cell? a) chemical to electrical b) electrical to chemical

_____ 22.

d) +7

How many electrons are transferred in the reaction? a) 2

_____ 21.

b) +3

c) anode. d) salt bridge.

The component in an electrochemical cell where loss of electrons takes place is the a) wire. b) cathode.

c) anode. d) salt bridge.

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III. Free Response Answer the questions in the space provided.

24.

Assign the oxidation state for sulfur in each of the following compounds. a) H2 S

_____

c) SO2 _____

b) H2 SO4 _____ 25.

d) S8

_____

For the reaction below, identify what is being oxidized, what is being reduced, the oxidizing agent, and the reducing agent.

CH4 (g) + O2 (g) → CO2 (g) + H2 O(g)

26.

oxidized:

____________

oxidizing agent:

_____________

reduced:

____________

reducing agent:

_____________

Balance the following oxidation–reduction reactions that take place in acidic solution by using the half-reaction method. Show your work. a) I– (aq) + MnO4 – (aq) → Mn2+ (aq) + I2 (s)

b) NO2 → NO3 – + NO

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27.

Mark a “T” next to each statement below that is true. For the statements which are false, change one or more words to make them true. a) _____ The reaction Fe3+ + e– → Fe2+ is a reduction reaction. b) _____ Reduction takes place at the anode in a galvanic cell. c) _____ The oxidation number of an ion is the same as the electrical charge on the ion. d) _____ The reaction Cu0 → Cu2+ is a balanced oxidation half-reaction. e) _____ The oxidation number of molybdenum in Na2 MoO4 is +4.

28.

Why are elements that appear on the left side of the periodic table such good reducing agents? Why are elements that appear on the right side of the periodic table so readily reduced?

29.

What is the difference between a) an anode and a cathode?

b) a galvanic cell and an electrolytic cell?

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Chapter 19

Radioactivity and Nuclear Energy I. Matching Match the description in Column B with the correct term in Column A. Write the letter in the blank provided. Each term matches with only one description, so be sure to choose the best description for each term.

Column A

Column B

_____ 1. nucleons

A. the change of one element into another

_____ 2. radioactive

B. used in medical field to help provide information about how the body functions

_____ 3. alpha particle

C. helium nucleus

_____ 4. beta particle

D. self-sustaining process

_____ 5. gamma ray

E. instrument for measuring radioactivity levels

_____ 6. positron

F. undergoes spontaneous nuclear decay

_____ 7. nuclear transmutation

G. high energy photon

_____ 8. Geiger counter

H. amount of radioactive material necessary for fission to occur

_____ 9. half-life

I. time it takes for half a sample of radioactive nuclei to decay

_____ 10. radiotracers

J. splitting one heavy nucleus into two nuclei with smaller mass numbers

_____ 11. fusion

K. electron

_____ 12. fission

L. particle with the same mass as an electron but positive charge

_____ 13. chain reaction

M. subatomic particles found in the nucleus

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_____ 14. critical mass

N. combining two light nuclei to form a heavier nucleus

II. Multiple Choice Choose the one best answer and write its letter in the blank. _____ 15.

Identify the missing particle in 239 94Pu → _____ + α. a) b)

_____ 16.

c) d)

b)

210 82Pb 206 81Tl

c) d)

c) electron capture. d) positron production.

Control rods are used in nuclear reactors to a) cool the reactor chamber. b) absorb neutrons.

_____ 19.

c) spin the turbine. d) check for radiation leaks.

Which radioactive process emits the product with the greatest penetrating power? a) alpha emission b) beta emission

_____ 20.

210 84Po 214 85At

The reaction 137N → 136C + 01e is an example of a) alpha particle production. b) beta particle production.

_____ 18.

237 90Th 239 Np 93

0 Identify the missing particle in 210 83Bi → _____ + –1e.

a)

_____ 17.

235 92U 239 95Am

c) positron emission d) gamma emission

Potassium-40 decays by electron capture. Which equation correctly represents this process? a) b)

40 19K 40 19K

+ –10e → 40 18Ar

c)

+ 10n → 41 19K

d)

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40 19K 40 19K

0 → 40 20Ca + –1e 0 + 40 18Ar + 1e

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_____ 21.

Which type of radioactive product has zero mass and zero charge? a) positron

_____ 22.

b) gamma

half-life 12.26 years 5730 years 14.66 hours 14.3 days

Proton bombardment of plutonium-239 (Z = 94) will result in the production of a neutron and which of the following nuclides? a) b)

_____ 24.

234 92U 239 95Am

c) d)

242 96Cm 240 94Pu

The half-life of carbon-14 is 5730 years. A piece of wood that contained 0.200 g of carbon-14 when it was alive now contains 0.0125 g of carbon-14.

How old is the piece of wood? a) 5730 years old b) 11,460 years old

_____ 25.

d) beta

Given the following information, which nuclide would be considered the “hottest”? nuclide a) hydrogen-3 b) carbon-14 c) sodium-24 d) phosphorus-32

_____ 23.

c) alpha

c) 17,190 years old d) 22,920 years old

A 3.60 g sample of a radioactive material has decayed to 0.45 g over the course of 6.75 hours. What is the half-life of this material? a) 2.25 hours b) 3.38 hours

c) 6.75 hours d) 13.5 hours

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III. Free Response Answer the questions in the space provided. Show your work for any calculations.

26.

Write a balanced nuclear equation for each of the following: a) beta decay of 90 37Rb

b) alpha decay of 230 90Th

c) positron production from 30 15P

27.

2400.0 g of a radioactive sample has decayed for 155 years. If the half-life of this material is 31.0 years, how much of the original sample remains?

28.

The half-life of a particular radioactive material is 2.0 minutes. After 12 minutes, 0.25 g remains of a sample of this material. What was the mass of the original sample of this material?

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29.

A typical fission process occurs after a uranium-235 particle absorbs a neutron and becomes the unstable nuclide, uranium-236. Uranium-236 then splits to produce tellurium-137, zirconium-97, and two neutrons. Write the two equations that represent the nuclear reactions that are described here.

30.

Determine the missing component in each of these equations:

31.

a)

222

b)

40

c)

231

Rn →

K → Th →

40

218

Po + _____

Ca + _____ –1e

+ _____

Titanium-51 decays by beta emission with a half-life of 6 minutes. What fraction of titanium would remain after 30 minutes? After 1 hour? After 2 hours? Comment briefly on your answers.

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Chapter 20

Organic Chemistry I. Matching Match the description in Column B with the correct term in Column A. Write the letter in the blank provided. Each term matches with only one description, so be sure to choose the best description for each term.

Column A

Column B

_____ 1. organic chemistry

A. process in which many small molecules are joined to form a large molecule

_____ 2. biomolecules

B. consists of mostly methane with some ethane, propane, and butane

_____ 3. hydrocarbons

C. two molecules have the same atoms but different bonds

_____ 4. saturated

D. hydrogen atoms are removed, and carbon– carbon multiple bonds are formed

_____ 5. unsaturated

E. compounds composed of carbon and hydrogen

_____ 6. structural isomerism

F. hydrogen atoms are replaced by different atoms

_____ 7. natural gas

G. study of carbon-based compounds and their properties

_____ 8. substitution reaction

H. all carbon–carbon bonds are single bonds

_____ 9. dehydrogenation reaction

I. two different types of monomers form the chain

_____ 10. addition reaction

J. new atoms form single bonds to carbons formerly involved in double or triple bonds

_____ 11. polymerization

K. includes carbon–carbon multiple bonds

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L. molecules that make possible the maintenance and reproduction of life

_____ 12. copolymer

II. Multiple Choice Choose the one best answer and write its letter in the blank. _____ 13.

Which of the following are structural isomers of each other?

1.

2. CH3 CH2 CH2 CH2 CH3

3.

4.

a) 1 and 2 b) 1 and 3 _____ 14.

c) 3 and 4 d) All are structural isomers.

The compound shown below is an example of a(n)

a) primary alcohol. b) secondary alcohol. _____ 15.

The compound C2 H2 is an example of a(n) a) alkane.

_____ 16.

c) tertiary alcohol. d) aldehyde.

b) alkene.

c) alkyne.

A straight-chain carboxylic acid with five carbons would be called a) butanoic acid. b) pentanoic acid.

_____ 17.

d) ketone.

c) hexanoic acid. d) carbonic acid.

The compound 1,2-dibromopropane can be formed from the reaction of a) bromine and propene. b) bromine and propane.

c) bromine and propanoic acid. d) bromine and propanal.

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_____ 18.

The name of a substituent that includes six carbons in an aromatic ring is a) hexane.

_____ 19.

c) propyl.

b) –NH2

c) –COOH

c) halogenation. d) addition.

The hydrogenation of 1-pentene will result in which compound? a) pentane b) pentyne

_____ 22.

c) pentanoic acid d) pentanol

What organic compound is typified by the presence of a triple bond? a) carboxylic acid b) aldehyde

_____ 23.

d) –Cl

Unsaturated hydrocarbons undergo each of the following chemical reactions except a) hydrogenation. b) substitution.

_____ 21.

d) phenyl.

What is the functional group for an alcohol? a) –OH

_____ 20.

b) ethyl.

c) alkyne d) benzene

How many structural isomers exist for butane, C4 H10 ? a) 1

b) 2

c) 3

d) 4

III. Free Response Answer the questions in the space provided.

24.

Identify the functional groups in the following molecules: a)

_______________

b)

_______________

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Name______________________________ Section__________ Date______________

25.

c)

_______________

d)

_______________

e)

_______________

Name the following compounds: a)

_______________

b)

_______________

c)

_______________

d)

_______________

e)

_______________

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26.

Draw structural formulas for the following compounds: a) 2,2-dimethyloctane

b) 2,4-dichlorotoluene

c) 2-methylphenol

d) 2-ethyl-3-methylhexanoic acid

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27.

What is the difference between a) a primary alcohol and a secondary alcohol?

b) an aldehyde and a ketone?

c) saturated and unsaturated?

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Chapter 21

Biochemistry I. Matching Match the description in Column B with the correct term in Column A. Write the letter in the blank provided. Each term matches with only one description, so be sure to choose the best description for each term.

Column A

Column B

_____ 1. biochemistry

A. provides structural integrity and strength for many types of living tissues

_____ 2. protein

B. sequence of amino acids in a polypeptide

_____ 3. fibrous protein

C. catalyzes specific biological reactions

_____ 4. polypeptide

D. includes deoxyribose, phosphate, and nitrogenous bases

_____ 5. primary structure

E. the study of the chemistry of living systems

_____ 6. enzyme

F. a class of polymers including C, H, N, and O, which breaks down into amino acids

_____ 7. monosaccharide

G. multiple amino acids connected by peptide linkages

_____ 8. deoxyribonucleic acid

H. a given segment of DNA that contains the code for a specific protein

_____ 9. ribonucleic acid

I. a class of water insoluble substances that can be extracted from cells by organic solvents

_____ 10. gene

J. carbohydrate that cannot be broken down into simpler carbohydrate units

_____ 11. lipid

K. a class of lipids of which testosterone is an example

_____ 12. steroid

L. includes ribose, phosphate, and nitrogenous bases

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II. Multiple Choice Choose the one best answer and write its letter in the blank. _____ 13.

What is the smallest unit in living things that exhibits the properties normally associated with life? a) bio-chemical b) organism

_____ 14.

Which of the following is an essential element? a) radon

_____ 15.

c) enzyme d) cell

b) nitrogen

c) cesium

_____ is an example of a fibrous protein. a) Hair b) Egg proteins

_____ 16.

b) secondary

c) tertiary

b) secondary

c) tertiary

c) enzyme activity d) lipid activity

Which of the following is an example of a monosaccharide? a) sucrose

_____ 20.

d) quaternary

The “lock and key” model applies to which of the following? a) protein synthesis b) DNA and RNA

_____ 19.

d) quaternary

Disulfide linkages found in proteins help maintain their _____ structure. a) primary

_____ 18.

c) Myoglobin d) None of the above

A certain protein, similar to silk, is very flexible but strong, and is resistant to stretching. These properties are most associated with its _____ structure. a) primary

_____ 17.

d) gold

b) fructose

c) starch

d) cellulose

A polysaccharide is formed when glucose molecules bond via a) disulfide linkages. b) peptide bonds.

c) hydrogen linkages. d) glycoside linkages.

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_____ 21.

Which combination of components forms a nucleotide? a) nitrogenous base, six-carbon sugar, phosphate group b) nitrogenous base, five-carbon sugar, phosphate group c) lipid, starch, nucleic acid d) nucleic acid, polypeptide, starch

_____ 22.

_____ is a nitrogenous base found only in DNA. a) Thymine

_____ 23.

c) Quinine

b) tRNA

c) rRNA

d) cRNA

A lipid that is an ester of glycerol has three fatty acid chains. It is most likely a a) wax. b) phospholipid.

_____ 25.

d) Cytosine

This type of RNA is responsible for bringing amino acids to the ribosome during protein synthesis. a) RNA

_____ 24.

b) Uracil

c) fat. d) steroid.

Cholic acid is an example of a _____ and is produced from _____. a) nucleic acid; DNA b) protein; amino acids c) lipid; cholesterol d) sugar; monosaccharies

_____ 26.

Which of these sugars is used to make DNA? a) ribose b) deoxyribose

_____ 27.

c) dextrose d) glucose

What type of reaction occurs to link monosaccharides together into disaccharides? a) condensation b) dehydration c) combustion d) hydrogenation

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III. Free Response For questions 28–31, read the following: A biochemist is responsible for classifying each of the following substances as a protein, nucleic acid, carbohydrate, or lipid. She has made the following observations.

Sample 1: insoluble in water, similar in composition to estrogen Sample 2: structure formed from two strands wound around each other; distinct units include a nitrogen-base, sugar, and a phosphate group Sample 3: globular in structure, polymer formed by peptide bonds and disulfide linkages Sample 4: composed of carbon, hydrogen, oxygen; soluble in water 28.

Sample 1 is most likely a __________________________.

29.

Sample 2 is most likely a __________________________.

30.

Sample 3 is most likely a __________________________.

31.

Sample 4 is most likely a __________________________.

Answer the questions in the space provided. Use complete sentences.

32.

Explain the significance of enzymes in biological reactions and describe how they function using the “lock and key” model.

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33.

Explain how a surfactant removes grease from fabric. Why are detergents more effective in hard water than soaps?

34.

Briefly describe the difference between a) a monosaccharide and a disaccharide.

b) hydrophobic and hydrophilic.

c) a glycoside linkage and a peptide linkage.

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Chapter Test Answers Chapter 1: Chemistry: An Introduction 1. E, 2. A, 3. G, 4. I, 5. H, 6. F, 7. D, 8. B, 9. J, 10. C, 11. D, 12. C, 13. A, 14. D, 15. D, 16. C, 17. C 18. A natural law is merely a summary of observations about the way nature behaves. It makes no attempt to explain why things occur the way they do. A theory, on the other hand, is a hypothesis (explanation) that has stood up to much experimentation and study. It does attempt to give a possible explanation as to how and/or why phenomena occur.

19. Answers should include the definition of a trade-off as the process of giving up one advantage in favor of another when developing a solution to an engineering problem. Engineers must rank the importance of criteria in order to assess how a solution best meets the needs of a problem; this often means the final design will better favor one criterion over another, but will result in an overall more optimal solution.

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Chapter Test Answers Chapter 2: Matter 1. C, 2. I, 3. E, 4. F, 5. J, 6. G, 7. D, 8. K, 9. L, 10. B, 11. H, 12. A, 13. C, 14. D, 15. A, 16. B, 17. D, 18. B, 19. A, 20. C, 21. B, 22. A 23. physical, 24. physical, 25. chemical, 26. chemical, 27. physical 28. There will be no difference between the properties of pure water that has been boiled and condensed and the properties of pure water that has been frozen and then melted. All of these transitions are physical changes (changes of state), so the identity of the substance (water) has not been altered.

29. Answers will vary, but some possibilities include color, texture, solubility in water, behavior when held in a flame, and malleability.

30. The red powder is a compound since it decomposed into more than one substance. No, you cannot classify the shiny liquid as an element or a compound until further experiments are done with it. Sometimes compounds decompose into elements, and sometimes they decompose into simpler compounds.

31. First, scoop the oil off of the surface of the water. Next, filter the salt water and sand mixture.

The sand will remain in the filter paper. Finally, distill (or simply evaporate) the salt water. The water will evaporate, leaving the salt crystals behind.

32. The drawing should indicate the particles of gas being much more widely separated than the particles of solid.

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Chapter Test Answers Chapter 3: Measurements and Calculations 1. C, 2. B, 3. D, 4. B, 5. C, 6. D, 7. A, 8. A, 9. B, 10. A, 11. B, 12. D, 13. A, 14. C, 15. C, 16. A, 17. D, 18. D 19. 2.5 × 103 , 20. 3.05 × 10–2 , 21. 2 × 105 , 22. m3 , 2 S.F., 23. mL, 2 S.F., 24. g/mL, 3 S.F., 25. m/s, 2 S.F. 4 qt 1L 26. 55.0 gal × × = 208 L 1 gal 1.06 qt 50.8 g = 16.0 g/mL 3.18 mL Therefore, the metal sample is not gold since its density is lower than the density of gold. g mcube = ; mcube = 170 g 28. daluminum = 2.7 cm3 4.0 cm × 4.0 cm × 4.0 cm 29. No measurement is perfect; all measurement devices have limitations to their precision. We round numbers to indicate just how certain we are of the answer to a calculation and how exactly we know the quantity we are measuring.

27. dmetal =

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Chapter Test Answers Chapter 4: Chemical Foundations: Elements, Atoms, and Ions 1. J, 2. C, 3. E, 4. G, 5. A, 6. H, 7. D, 8. I, 9. B, 10. F, 11. B, 12. D, 13. C, 14. A, 15. C, 16. D, 17. C, 18. B, 19. A, 20. A, 21. C, 22. C 23.

Element Symbol Atomic No. of No. of Mass No. of Charge (atom/ion) number protons neutrons number electrons 35 – chlorine 17 17 18 35 18 –1 17Cl ion 1 + hydrogen 1 1 0 1 0 +1 1H ion 23 sodium 11 11 12 23 11 0 11Na atom 24. AlBr3 , K2 S 25. Thomson’s model of the atom consisted of tiny negatively charged particles (electrons) floating in a cloud of positive charge. Rutherford’s model of the atom consisted of a dense, positively charged nucleus at the center of the atom, with the majority of the atom being empty space, in which the electrons resided. He developed his model as a result of an experiment in which alpha particles (which are positively charged) were fired at a thin sheet of gold foil. Had Thomson’s model been correct, the alpha particles would have passed straight through the foil with at most showing only a minimal amount of deflection. However, while most did pass straight through, some showed a great deal of deflection, with a few even being reflected back.

Rutherford concluded there must be a dense, positively charged nucleus at the center of the atom.

26. a) Due to aluminum’s fewer protons in its nucleus (13 as opposed to 79 for gold), the nucleus of an aluminum atom would have made a much smaller target in terms of the field of positive charge that it produced. Consequently, more alpha particles would have simply passed through the foil without showing any deflection or reflection at all. Only those alpha particles that passed particularly close to the aluminum nuclei would have deflected, and these would be fewer in number than those in the gold foil experiment. b) Most of the negatively charged particles would still have passed right through the gold undeflected. However, rather than being repelled by the nucleus in the atom of gold, the negatively charged particles would have been attracted to the nucleus, perhaps even being absorbed into the nucleus.

27. H+ contains only a proton. H– contains a proton and two electrons. Neither ion contains neutrons.

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Chapter Test Answers Chapter 5: Nomenclature 1. D, 2. A, 3. E, 4. C, 5. B, 6. B, 7. D, 8. D, 9. A, 10. B, 11. C, 12. B, 13. D, 14. C, 15. A, 16. C, 17. B, 18. C, 19. A, 20. A 21. Al2 S3 , 22. sulfur difluoride, 23. PCl3 , 24. zinc nitrate, 25. Fe2 O3 , 26. copper(I) iodide, 27. nitric acid, 28. Al(OH)3 , 29. calcium bromide, 30. HCl, 31. barium phosphate, 32. a) BaHPO4 , b) Fe2 (HPO4 )3 , c) V2 (HPO4 )5 , d) Al2 (HPO4 )3 33. Protons are hidden away in the nucleus of the atom. Electrons, on the other hand, are located on the outside of atoms and are involved in collisions or interactions between atoms.

34. a) Fe3+ , iron(III) sulfate, b) Cu1+ , copper(I) oxide, c) Co2+ , cobalt(II) acetate

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Chapter Test Answers Chapter 6: Chemical Composition 1. G, 2. E, 3. A, 4. D, 5. B, 6. C, 7. F, 8. D, 9. B, 10. C, 11. D, 12. A, 13. B, 14. A, 15. D, 16. B, 17. A, 18. D 19. molar mass of (NH4 )2 S = (2)(14.01 g/mol) + (8)(1.008 g/mol) + 32.06 g/mol = 68.14 g/mol (2)(14.01 g/mol) × 100% = 41.12% N 68.14 g/mol (8)(1.008 g/mol) × 100% = 11.83% H 68.14 g/mol 32.06 g/mol × 100% = 47.05% S 68.14 g/mol 20. 2.65 × 1024 molecules Cl2 × 21. 83.2 g SO2 × 22. 14.7 g Ni ×

1 mol Cl2 23

6.02 × 10 molecules Cl2

×

70.90 g Cl2 = 312 g Cl2 1 mol Cl2

1 mol SO2 1 mol S 32.06 g S × × = 41.6 g S 64.06 g SO2 1 mol SO2 1 mol S

1 mol Ni 1 mol Br = 0.250 mol Ni; 40.0 g Br × = 0.500 mol Br 58.69 g Ni 79.90 g Br

0.250 mol Ni ∕ 0.250 = 1 mol Ni; 0.500 mol Br ∕ 0.250 = 2 mol Br The empirical formula is NiBr2 .

23. a)

30.4 g N 69.6 g O = 2.17 mol N; = 4.35 mol O 14.0 g/mol 16.0 g/mol

There is one mole of N for every two moles of O; therefore, the empirical formula is NO2.

The molar mass of NO2 is 46.0 g/mole. Since the known molar mass of the compound is between 90 and 100 g/mol, the true molecular formula must be N2O4. b) The compound is named dinitrogen tetroxide.

24. a) 35, b) Yes, it is empirical because the nitrogen atom has a subscript of 1. c) 777 g/mol, d) 28,480 atoms I, e)

2 × 109 molecules 6.02 × 1023 molecules/mol

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25.

28.4 g Al(OH)3 7 atoms = 0.364 mol Al(OH)3 × = 2.55 mol atoms 78.0 g/mol molecule

5.12 × 1022 atoms Ba 6.02 × 10

23

= 8.50 × 10–2 mol atoms ×

137.3 g = 11.7 g Ba mol

3 atoms = 1.81 mol atoms molecule 48.0 g 0.604 mol O3 × = 29.0 g O3 mol Therefore, O3 weighs the most, Ba weighs the least, and Al(OH)3 contains the greatest number of atoms.

48 g O 26. = 3.0 mol O = 3.0 mol N × 14.0 g/mol = 42 g N 16.0 g/mol 0.604 mol O3 ×

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Chapter Test Answers Chapter 7: Chemical Reactions: An Introduction 1. C, 2. B, 3. D, 4. A, 5. B, 6. A, 7. A, 8. B, 9. C, 10. D, 11. C, 12. C, 13. B, 14. D 15. PbI2 + 2AgNO3 → Pb(NO3 )2 + 2AgI, 16. 2Mg + TiCl4 → 2MgCl2 + Ti, 17. C3 H8 + 5O2 → 3CO2 + 4H2 O, 18. P4 + 5O2 → P4 O10 , 19. Na2 CO3 + 2HCl → 2NaCl + CO2 + H2 O, 20. N2 + 3I2 → 2NI3 , 21. Zn + 2HCl → ZnCl2 + H2 , 22. 3BaCl2 + Al2 (SO4 )3 → 3BaSO4 + 2AlCl3 , 23. Ca(OH)2 + 2HNO3 → Ca(NO3 )2 + 2H2 O, 24. 3Mg + 2FeCl3 → 3MgCl2 + 2Fe, 25. 2Li(s) + 2H2 O(l) → 2LiOH(aq) + H2 (g) 26. A subscript is used to establish the molecular formula of a given compound; it is never changed when an equation is written or balanced. A coefficient is used as a multiplier to balance the atoms on both sides of an equation.

27. Though the equation shows a 1:1 relationship between the C and the O2 , that relationship pertains to moles and not to grams. Since C and O2 has different molar masses, they will not react in a 1:1 relationship with respect to grams.

28. Fe2 O3 + 2Al → 2Fe + Al2 O3 ; this is a single-replacement reaction.

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Chapter Test Answers Chapter 8: Reactions in Aqueous Solutions 1. G, 2. J, 3. L, 4. K, 5. H, 6. I, 7. B, 8. C, 9. A, 10. F, 11. E, 12. D, 13. A, 14. D, 15. B, 16. C, 17. D, 18. D, 19. C, 20. B, 21. A, 22. D, 23. A 24. 3Mg(s) + 2FeCl3 (aq) → 3MgCl2 (aq) + 2Fe(s), 25. 4K(s) + O2 (g) → 2K2 O(s) 26. Zn(s) + 2HCl(aq) → ZnCl2 (aq) + H2 (g), single-replacement reaction 27. 2C4 H10 (g) + 13O2 (g) → 8CO2 (g) + 10H2 O(g), combustion reaction 28. 6Li(s) + N2 (g) → 2Li3 N(s), synthesis reaction 29. a) CuSO4 (aq) + Na2 S(aq) → CuS(s) + Na2 SO4 (aq), b) Cu2+ (aq) + SO4 2– (aq) + 2Na+ (aq) + S2– (aq) → CuS(s) + 2Na+ (aq) + SO4 2– (aq), c) Cu2+ (aq) + S2– (aq) → CuS(s) 30. a) In a synthesis reaction, two elements or simpler molecules come together to form a more complex compound. A decomposition reaction is just the opposite; one compound breaks down into two or more elements or simpler molecules. b) In a double replacement reaction, a precipitate is formed. A single replacement reaction is an oxidation–reduction reaction, or two elements (either two metals or two nonmetals) exchange places. c) A strong electrolyte dissociates into many ions in solution and therefore conducts electricity well. A nonelectrolyte does not form ions and therefore does not conduct electricity.

31. a) 2NH3 → N2 + 3H2 , b) H2 + Br2 →2HBr

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Chapter Test Answers Chapter 9: Chemical Quantities 1. A, 2. C, 3. F, 4. B, 5. D, 6. E, 7. D, 8. A, 9. B, 10. A, 11. B, 12. A, 13. B, 14. B 3 mol O2 15. 2KClO3 → 2KCl + 3O2 ; 5.4 mol KClO3 × = 8.1 mol O2 2 mol KClO3 16. 45.0 g NH3 × 45.0 g NH3 ×

1 mol NH3 2 mol N2 28.02 g N2 × × = 37.0 g N2 17.03 g NH3 4 mol NH3 1 mol N2

1 mol NH3 6 mol H2 O 18.02 g H2 O × × = 71.4 g H2 O 17.03 g NH3 4 mol NH3 1 mol H2 O

17. Cu + 2AgNO3 → Cu(NO3 )2 + 2Ag, single-replacement reaction mass of Cu reacted = 2.36 g – 1.03 g = 1.33 g Cu 1 mol Cu 2 mol Ag 107.9 g Ag 1.33 g Cu × × × = 4.52 g Ag produced 63.55 g Cu 1 mol Cu 1 mol Ag 1 mol C 2 mol ZnO 18. a) 6.50 g C × × = 1.08 mol ZnO needed 12.01 g C 1 mol C 1 mol ZnO = 0.123 mol ZnO present; ZnO is limiting.

81.38 g ZnO 1 mol ZnO 2 mol Zn 65.38 g Zn 10.0 g ZnO × × × = 8.03 g Zn 81.38 g ZnO 2 mol ZnO 1 mol Zn b) single replacement 3.0 g HCl = 7.5 g HCl 19. 2.5 L × 1L 1 mol HCl 1 mol Al(OH)3 78.0 g Al(OH)3 7.5 g HCl × × × = 5.3 g Al(OH)3 needed 36.5 g HCl 3 mol HCl 1 mol Al(OH)3 1 tablet 5.3 g Al(OH)3 × = 13.25 tablets ⇒ 14 antacid tablets needed 0.400 g (assuming you cannot break apart a tablet) 1 mol LiOH 1 mol CO2 44.0 g CO2 20. a) 454 g LiOH × × × = 416 g CO2 24.0 g LiOH 2 mol LiOH 1 mol CO2 10.0 g ZnO ×

b)

308 g CO2 × 100% = 74.0% CO2 416 g CO2

21. In the M → N step: 2.0 mol⁄5.0 mol × 100% = 40.0% yield In the N → O step: 0.25 × 2.0 mol = 0.50 mol O In the O → P step: 0.10 × 0.50 mol = 0.050 mol P In the P → Q step: 0.015 mol⁄0.050 mol × 100% = 30.0% yield

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Chapter Test Answers Chapter 10: Energy 1. J, 2. H, 3. D, 4. F, 5. C, 6. I, 7. K, 8. L, 9. G, 10. B, 11. E, 12. A, 13. C, 14. B, 15. D, 16. B, 17. C, 18. A, 19. D, 20. A, 21. C 22. a) Endothermic; the temperature of the surroundings (the reaction vessel) has decreased, so heat has flowed from the surroundings into the system. b) The products have the higher energy content, as the system has absorbed the heat energy from the surroundings.

23. q = mCDT = (38.0 g)(l cal/(g•°C))(95.0 °C – 25.0 °C) = 2660 cal (or 11,100 J to 3 S.F.) 24. heat lost by aluminum = heat gained by water –mAl CAl DTAl = mwater Cwater DTwater –(4.50 g)(0.902 J/(g•°C))(Tf – 99.5 °C) = (12.0 g)(4.184 J/(g•°C))(Tf – 21.0 °C) –4.059Tf + 403.871 = 50.208Tf – 1054.368 54.267Tf = 1458.239 Tf = 26.9 °C 25. 2[C(s) + O2 (g) → CO2 (g)] DH = 2(–393.5 kJ) = –787.0 kJ 2[H2 (g) + ½O2 (g) → H2 O(l)] DH = 2(–285.8 kJ) = –571.6 kJ (g) (g) (g) 2H2 O(l) + 2CO2 → C2 H4 + 3O2 DH = –(–1410.9 kJ) = +1410.9 kJ (g) (g) 2C(s) + 2H2 → C2 H4 DH = +52.3 kJ 26. E → 2A DH = –48.9 kJ/mol 2B + D → C + E DH = –449.0 kJ/mol 3A → D + 2C DH = +25.28 kJ/mol A + 2B → 3C DH = –472.6 kJ/mol 27. When atoms or ions are bonded together, there are forces of attraction between the atoms or the ions. Energy must be absorbed to overcome these forces. The stronger the force of attraction, the more energy is required.

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Chapter Test Answers Chapter 11: Modern Atomic Theory 1. F, 2. J, 3. D, 4. G, 5. I, 6. A, 7. C, 8. H, 9. E, 10. B, 11. C, 12. C, 13. D, 14. A, 15. D, 16. D, 17. A, 18. B, 19. A, 20. A, 21. A, 22. C, 23. A, 24. A 25. S, 26. Ne, 27. Sr, 28. Fr, 29. P or V, 30. Sc, 31. [Xe]6s2 4f14 5d10 6p5 , 32. a) 1s2 2s2 2p6 3s2 3p1 , b) 1s2 2s2 2p6 3s2 3p6 4s2 3d6 33. Across a period, the number of protons in the nucleus is increasing, while the electrons are in the same principal energy level. Because of this increasing positive nuclear charge, the electrons feel a greater attraction towards the nucleus and so are drawn in closer to it, thereby decreasing the size of the atoms. Down a group, the electrons are going into higher principal energy levels and so are at a greater distance from the nucleus, thereby making the atoms larger down a group.

34. a) Cr, b) Mn, c) B, d) K, e) B 35. The 1s orbital, being closest to the nucleus, represents the orbital with the lowest potential energy in the atom. Electrons can occupy it with relatively little energy.

36. a) The color emitted in the 6 ® 1 jump would be higher energy. If the colors were both visible, the 6 ® 1 jump would produce light more toward the blue or violet end of the spectrum, and the 2 ® 1 jump would produce light toward the red end of the spectrum. Regardless, the color, frequency, wavelength, and energy of the light would be different. b) If many electrons were making a jump, the light (regardless of its color or frequency) would be brighter.

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Chapter Test Answers Chapter 12: Chemical Bonding 1. G, 2. C, 3. J, 4. I, 5. F, 6. B, 7. D, 8. H, 9. A, 10. E, 11. A, 12. A, 13. B, 14. B, 15. C, 16. D, 17. B, 18. C, 19. D, 20. C, 21. C, 22. B 23. a) 1s2 2s2 2p6 3s2 3p6 , Ar, b) 1s2 2s2 2p6 , Ne, c) 1s2 2s2 2p6 3s2 3p6 , Ar 24. Lewis Structures and Molecular Structures:

a)

trigonal pyramidal

b)

linear

c)

trigonal planar

25.

26. The oxygen atoms located at either end of the linear molecule have higher electronegativities than the carbon in the middle. The oxygens pull the electrons away from the center of the CO2 molecule with equal force in opposite direction. Since the dipoles cancel one another out, the molecule is overall nonpolar.

27. Neon has a full set of valence electrons and has no half-filled orbitals with which to bond.

Chlorine, on the other hand, has seven valence electrons, and it does have one half-filled p orbital with which it can bond to another Cl to establish an octet and have greater stability.

28. In the NH3 molecule, there is a lone pair of electrons on the central atom, N. The repulsion from that pair of electrons on the other electrons in the molecule causes the Hs to occupy three corners of a tetrahedron rather than exist in a trigonal planar shape.

29. A polar covalent bond is one in which there is an unequal sharing of electrons between two atoms. As such, the shared electrons lie closer to one of the bonded atoms than the other, and the bond is unbalanced. Both atoms share the electron pair; the sharing is just unequal. In an ionic bond there is such a difference in electronegativity that one of the atoms actually takes an electron from the other. Therefore, one of the atoms becomes a – ion, and the other becomes a + ion. The attraction between the – and + ions is the ionic bond.

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Chapter Test Answers Chapter 13: Gases 1. C, 2. H, 3. A, 4. G, 5. E, 6. D, 7. B, 8. F, 9. A, 10. C, 11. B, 12. B, 13. C, 14. C, 15. A, 16. C, 17. D, 18. B, 19. D, 20. B 1 mol CO2 21. 18.9 g CO2 × = 0.429 mol CO2 44.01 g CO2 PV = nRT (4.35 atm)(V) = (0.429 mol) #0.08206 V = 2.67 L

L atm $ (330.5 K) mol K

1 mol CO2 = 0.07953 mol CO2 44.01 g CO2 1 mol O2 11.6 g O2 × = 0.3625 mol O2 32.00 g O2 L atm (PA )(3.00 L) = (0.07953 mol) #0.08206 $ (273 K) mol K PA = 0.594 atm CO2 L atm (PB )(3.00 L) = (0.3625 mol) #0.08206 $ (273 K) mol K PB = 2.71 atm O2 PTotal = PA + PB = 0.594 atm + 2.71 atm = 3.30 atm 23. CaCO3 → CaO + CO2 1 mol CaCO3 1 mol CO2 37.0 g CaCO3 × × = 0.370 mol CO2 100.09 g CaCO3 1 mol CaCO3 L atm (1.15 atm)(V) = (0.370 mol) #0.08206 $ (298 K) mol K V = 7.86 L CO2 1 atm L atm 24. (750 torr) # $ (0.125 L) = (n) #0.08206 $ (293 K) 760 torr mol K n = 0.0051 mol (0.0051 mol)(6.02 × 1023 molecules/mol) = 3.1 × 1021 molecules 25. a) No change; no molecules are lost. b) Increase; heat adds kinetic energy to the molecules. c) Increase; molecules move faster and collide more forcefully with the expandable container walls. d) No change; again, no molecules are lost. e) Decrease; since the molecules now occupy more volume, the density will be lower. f) Increase; as the molecules move faster and expand in volume, they are more widely separated from one another. g) No change; since the container walls are allowed to expand, pressure remains the same.

26. Gas molecules occupy much larger volumes than do either liquids or solids. Since density is a measurement of mass per volume, the density of gases is very low comparatively.

22. 3.50 g CO2 ×

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Chapter Test Answers Chapter 14: Liquids and Solids 1. E, 2. G, 3. H, 4. J, 5. B, 6. I, 7. D, 8. A, 9. C, 10. F, 11. A, 12. B, 13. B, 14. B, 15. C, 16. A, 17. B, 18. E, 19. A, 20. D, 21. C 22. When an ionic solid is struck, the ions shift position, placing ions with like charges adjacent to each other. The repulsion that results shatters the crystal. The free-flowing electrons present in a metallic bond act as a “lubricant” between the metal atoms, thus allowing them to move past one another when struck.

1 mol H2 O 23. 75.0 g H2 O × = 4.16 mol H2 O 18.02 g H2 O Melting: (4.16 mol)(6.02 kJ/mol) = 25.1 kJ Heating: q = mCDT = (75.0 g)(4.18 J/(g•°C))(100 °C) = 31,350 J = 31.35 kJ Vaporizing: (4.16 mol)(40.6 kJ/mol) = 168.90 kJ Total: 225 kJ 24. a) The melting point is the temperature at which a solid becomes a liquid. A substance’s freezing point is exactly the same temperature as its melting point. It simply describes the opposite process, that is, the temperature at which the liquid solidifies. b) The melting point is the temperature at which a solid becomes a liquid. The molar heat of melting is the energy required for one mole of the substance to go through that transition.

25. The fact that the hydrogen and oxygen atoms have different electronegativities causes the bond between them to be polar. The fact that the water molecule has a bent shape and two polar bonds gives it molecular polarity. The fact that water contains both hydrogen and a small, highly electronegative atom like oxygen is what allows it to demonstrate hydrogen bonding.

26. Boiling points are primarily determined by the strength of the intermolecular forces between the liquid molecules and by the mass of the molecules themselves. Therefore, since substances differ in the strength of their intermolecular forces and in their molar masses, there will be differences among their boiling points as well.

27. a) hydrogen bonds, b) dispersion forces, c) polar covalent bonds between O and H 1 mol CH3 OH = 0.7875 mol CH3 OH 28. 25.23 g CH3 OH × 32.04 g CH3 OH –4.01 kJ DHfusion = = –5.09 kJ/mol 0.7875 mol

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Chapter Test Answers Chapter 15: Solutions 1. D, 2. J, 3. E, 4. K, 5. G, 6. C, 7. B, 8. M, 9. F, 10. H, 11. A, 12. I, 13. L, 14. C, 15. B, 16. A, 17. D, 18. A, 19. C, 20. D, 21. D, 22. B, 23. B 24. #1 Grind up the solid. A smaller particle size increases the surface area, thereby increasing the number of collisions between the solvent and the solute. #2 Increase the temperature. This increases the kinetic energy of the solvent molecules, thereby increasing the number of collisions between the solvent and the solute. #3 Stir the solution. This increases the number of collisions between the solute and the solvent.

0.35 mol KI 166.0 g KI 25. 2.5 L × × = 145 g KI = 150 g KI (2 S.F.) 1L 1 mol KI 1 mol HC2 H3 O2 1 26. 75.0 g HC2 H3 O2 × × = 0.83 M HC2 H3 O2 60.1 g HC2 H3 O2 1.5 L 27. a) 2Ag+ + 2Cl– → 2AgCl or Ag+ + Cl– → AgCl 0.15 mol AgNO3 1 mol MgCl2 b) 0.0500 L × × = 0.00375 mol MgCl2 needed 1L 2 mol AgNO3 0.10 mol MgCl2 = 0.010 mol MgCl2 present 1L AgNO3 is limiting, so 0.15 mol AgNO3 2 mol AgCl 143.35 g AgCl 0.0500 L × × × = 1.1 g AgCl 1L 2 mol AgNO3 1 mol AgCl 0.1000 L ×

28. Even though both solutions have the same molarity, they do not have the same normality.

The H2 SO4 releases two hydrogens when it ionizes and, as a result, has an effective [H+ ] = 1.0 M; or 2H+ (2H3 O+ ) ions come from H2 SO4 , and 1H+ (1H3 O+ ) ion comes from HCl, though H2 SO4 is a stronger acid than HSO4 – .

12 g C2 H5 OH x = ; x = 294 g C2 H5 OH 29.

100 g solution 2450 g solution 1 mol C2 H5 OH 294 g C2 H5 OH × = 6.38 mol C2 H5 OH 46.07 g C2 H5 OH 30. Even though water does effectively dissolve many, many solutes, it does not dissolve nonpolar compounds (e.g., fats and oils). Therefore, the term universal solvent is not completely accurate.

31. a) 1.22 g/mL × 1000 mL = 1220 g solution 20.0 g fructose b) 1220 g solution × = 244 g fructose 100 g solution c)

1220 g solution – 244 g fructose = 976 g water

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Chapter Test Answers Chapter 16: Acids and Bases 1. L, 2. C, 3. I, 4. K, 5. D, 6. A, 7. B, 8. G, 9. E, 10. F, 11. H, 12. J, 13. A, 14. B, 15. C, 16. A, 17. B, 18. C, 19. D, 20. B, 21. D, 22. D 23. a) HCN(aq) → H+ (aq) + CN– (aq), b) NH4 + (aq) → NH3 (aq) + H+ (aq), c) H2 PO4 – (aq) → H+ (aq) + HPO4 2– (aq) 24. a) 1.38 × 10–13 M, basic, b) 2.56 × 10–4 M, acidic 25. a) 3.6 × 10–7 M, acidic, b) 1.8 × 10–13 M, basic 26. moles H+ = moles OH– (MHCl )(0.0500 L) = (0.20 M)(0.0200 L) MHCl = 0.080 M 27. a) A base, according to the Arrhenius model, is a substance that releases OH– in solution.

According to the Brønsted-Lowry model, a base captures or accepts an H+ . b) Both pH and [H+ ] describe the concentration of hydrogen ions in solution. pH is found by taking the negative logarithm of the[H+ ].

28. a) NH4 + + Cl– → NH3 +HCl (NH3 is the conjugate base) b) HCO3 – + H2 PO4 – → H2 CO3 + HPO4 2– (HPO4 2– is the conjugate base) Theoretically, the equation in b) above could have been written as: HCO3 – + H2 PO4 – → CO3 2– + H3 PO4 (CO3 2– is the conjugate base) However, the ionization constant for H2 PO4 – is greater than that for HCO3 – , so the dihydrogen phosphate ion will be the acid.

2.55 g NaOH 29. moles OH– = = 0.0638 mol 40.0 g/mol NaOH 0.0638 mol [OH– ] = = 0.116 M 0.5500 L pOH = –log(0.116) = 0.936; pH = 14.000 – 0.936 = 13.064 30. A dilute solution of a strong acid does not contain a relatively large number of solute molecules. However, the solute molecules that are there are completely ionized. A concentrated solution of a weak acid would contain many solute molecules (i.e., many moles of solute per liter), but only a small portion of them would be ionized. The solutions may in fact, as a result, have similar concentrations of ions.

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Chapter Test Answers Chapter 17: Equilibrium 1. E, 2. I, 3. G, 4. F, 5. B, 6. J, 7. D, 8. C, 9. A, 10. H, 11. A, 12. B, 13. C, 14. C, 15. D, 16. A, 17. D, 18. A, 19. B, 20. D 21. a)

K=

[SO3 ]2 [SO2 ]2 [O2 ]

b)

K = [Al3+ ][OH– ]3

[C][D]4 [1.2][0.30]4 K= = = 34.56 = 35 (2 S.F.) [A]2 [B] [0.025]2 [0.450] 23. PbI2 (s) ⇌ Pb2+ (aq) + 2I– (aq) Ksp = [Pb2+ ][I– ]2 7.1 × 10–9 = [x][2x]2 7.1 × 10–9 = 4x3 x = 1.21 × 10–3 mol/L (1.21 × 10–3 mol/L)(461.0 g/mol) = 0.56 g/L 24.

A2 B ⇌ 2A + B Initial concentrations: 1.2 M 0M 0M Change in concentrations: –0.6 M +1.2 M +0.6 M Equilibrium concentrations: 0.6 M 1.2 M 0.6 M 2 2 [A] [B] (1.2) (0.6) K= = = 1.4 [A2 B] (0.6) 22.

25. a) ® b) ® c) ¬ d) ® e) no change in either direction

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Chapter Test Answers Chapter 18: Oxidation–Reduction Reactions and Electrochemistry 1. E, 2. F, 3. I, 4. B, 5. J, 6. H, 7. C, 8. A, 9. G, 10. D, 11. C, 12. C, 13. A, 14. B, 15. B, 16. D, 17. B, 18. B, 19. D, 20. C, 21. A, 22. D, 23. C 24. a) –2, b) +6, c) +4, d) 0 25. oxidized: C, reduced: O, oxidizing agent: O2 , reducing agent: CH4 5 (2I– → I2 + 2e– )

26. a)

2 (5e– + 8H+ + MnO4 – → Mn2+ + 4H2 O) 10I– + 16H+ + 2MnO4 – → 5I2 + 2Mn2+ + 8H2 O b)

2 (H2 O + NO2 → NO3 – + 2H+ + e– ) 2e– + 2H+ + NO2 → NO + H2 O H2 O + 3NO2 → 2NO3 – + NO + 2H+

27. a) T, b) F; “anode” should be replaced with “cathode,” c) T, d) F; two electrons are needed on the product side: Cu0 → Cu2+ + 2e– , e) F; the oxidation number is +6 28. Elements on the left side of the periodic table have few valence electrons and low electronegativities. They lose their electrons relatively easily. Elements on the right side, on the other hand, have high electronegativity values and tend to gain electrons readily to achieve an octet of valence electrons.

29. a) An anode is the terminal in a battery where electrons are lost and oxidation occurs. A cathode is the terminal where electrons are gained and reduction occurs. b) A galvanic cell is one in which an oxidation–reduction reaction results in an electrical current. An electrolytic cell is one in which the electrical current causes a chemical reaction.

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Chapter Test Answers Chapter 19: Radioactivity and Nuclear Energy 1. M, 2. F, 3. C, 4. K, 5. G, 6. L, 7. A, 8. E, 9. I, 10. B, 11. N, 12. J, 13. D, 14. H, 15. A, 16. C, 17. D, 18. B, 19. D, 20. A, 21. B, 22. C, 23. B, 24. D, 25. A 90 230 30 226 30 0 4 0 26. a) 90 37Rb → 38Sr + –1e, b) 90Th → 88Ra + 2α, c) 15P → 14Si + 1e 27. (155 years)/(31.0 years) = 5 half-lives. Divide 2400.0 g by 2 five times, or (2400.0 g)/(25 ) = 75.0 g remains.

28. (12 minutes)/(2.0 minutes) = 6 half-lives. Multiply 0.25 g by 2 six times, or (0.25 g)(26 ) = 16 g of original sample.

29. 235U + 1n → 236U 236

U → 137Te + 97Zr + 2 1n 30. a) alpha particle, b) beta particle, c) 231Pa 31. (30 minutes)/(6 minutes) = 5 half-lives (60 minutes)/(6 minutes) = 10 half-lives (120 minutes)/(6 minutes) = 20 half-lives (1/2)5 = 1/32 left after 5 half-lives (1/2)10 = 1/1024 left after 10 half-lives (1/2)20 = 1/1,048,576 left after 20 half-lives Exponential decay results in the remaining fraction being very insignificant after two hours.

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Chapter Test Answers Chapter 20: Organic Chemistry 1. G, 2. L, 3. E, 4. H, 5. K, 6. C, 7. B, 8. F, 9. D, 10. J, 11. A, 12. I, 13. A, 14. B, 15. C, 16. B, 17. A, 18. D, 19. A, 20. B, 21. A, 22. C, 23. B 24. a) alcohol, b) carboxylic acid, c) ester, d) ketone, e) aldehyde 25. a) l-propanol (or n-propanol), b) 5-phenyl-l-hexyne, c) l-chloro-2-methyl-2-butene, d) 5,6-dimethyl-2-heptene, e) 1,4-dichlorobenzene (or p-dichlorobenzene) 26. a)

b)

c)

d)

27. a) A primary alcohol has the –OH functional group attached to the first carbon on the chain.

A secondary alcohol has the –OH functional group attached to a carbon that is attached to two other carbons. b) An aldehyde and a ketone have the carbonyl (–C=O) functional group in common. However, an aldehyde has the carbonyl group at the end of a carbon chain and has at least one hydrogen attached to the carbonyl group. A ketone has carbon chains attached to either side of the carbonyl group. c) A saturated compound is one in which all carbons are singly bonded to other atoms. An unsaturated compound contains at least one double or triple bond.

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Chapter Test Answers Chapter 21: Biochemistry 1. E, 2. F, 3. A, 4. G, 5. B, 6. C, 7. J, 8. D, 9. L, 10. H, 11. I, 12. K, 13. D, 14. B, 15. A, 16. B, 17. C, 18. C, 19. B, 20. D, 21. B, 22. A, 23. B, 24. C, 25. C, 26. B, 27. B 28. lipid, 29. nucleic acid, 30. protein, 31. carbohydrate 32. Enzymes are proteins that catalyze specific biological reactions. Without them, life would be impossible. In the “lock and key” model, the shapes of the enzyme and the reacting molecule (substrate) are such that they fit together in a specific manner like the way a key fits in a lock.

The reaction occurs, the products are liberated, and the enzyme can start again with a new molecule.

33. Surfactants have long nonpolar tails with a polar head. They form micelles in water (a polar solvent) wherein the nonpolar tails of the surfactant molecules are directed inward, and the polar heads are directed outward. Micelles can absorb nonpolar grease molecules into their interior, thus allowing the grease to be carried away by the water. Detergents are preferable to soaps in hard water because soaps tend to form precipitates (“soap scum”) with the Mg2+ and Ca2+ ions present in hard water. Detergents do not do this.

34. a) A monosaccharide is a simple sugar (which is an aldehyde or ketone with a number of –OH groups attached). A disaccharide is a more complex sugar made by attaching two monosaccharides together. b) Hydrophobic refers to something that is not water-soluble and usually pertains to nonpolar molecules. Hydrophilic describes a substance that does dissolve in water and, as a result, is a polar molecule. c) A glycoside linkage attaches monosaccharides together to make polysaccharides or carbohydrates. A peptide linkage attaches amino acids together to make polypeptides or proteins.

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