Chemistry Structure and Dynamics, 5th Edition Test Bank

Type:

Test Bank

Resource:

Chemistry Structure and Dynamics

Edition:

5th Edition

Author(s):

James N. Spencer George M. Bodner Lyman H. Rickard

Chapter 1: Elements & Compounds Introduction 1.

Which of the following correctly describes one of the differences between compounds and mixtures? (a) Compounds contain more than one element but mixtures do not. (b) Compounds have a constant composition but mixtures do not. (c) It is easier to separate the elements in a compound than it is to separate the elements in a mixture. (d) All of these statements are true. (e) None of these statements are true.

Answer: (b)

2.

Which of the following postulates of Dalton's atomic theory is/are still valid today?

(I) Matter consists of particles called atoms.

(II) Atoms are indestructible and indivisible.

(III) All atoms of an element are identical.

(IV) Atoms of different elements differ in mass.

(V) When the atoms of different elements combine to form compounds they combine in simple whole number ratios. (a) All of these statements are still valid today. (b) I and II are still valid, but not III, IV and V. (c) I, II and III are still valid, but not IV and V. (d) I, IV and V are still valid, but not II and III. (e) II and III are still valid, but not I, IV and V.

Answer: (d)

3.

Calculate the number of protons in the average potassium atom if the atomic weight of potassium is 39 amu and the atomic number of this element is 19. (a) 19 (b) 20 (c) 39 (d) 58 (e) The number of protons is impossible to determine from this information.

Answer: (a)

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Isotopes 4.

The earth is constantly bombarded by cosmic rays emitted by the sun. The total energy received in the form of cosmic rays is small — no more than the energy received by the planet from starlight. But the energy of a single cosmic ray is very large, on the order of 200 million kJ/mol. These highly energetic rays react with atoms in the atmosphere to produce neutrons that then react with nitrogen atoms in the atmosphere to produce 14C, which decays back to 14N with a half-life of 5730 years by emitting an electron.

What is the difference between 14C and 14N atoms? (a) They have the same number of protons but different numbers of neutrons. (b) They have the same number of neutrons but different numbers of protons. (c) They have the same number of neutrons and protons but different numbers of electrons. (d) 14N has one more proton and one less neutron than 14C. (e) 14N has one more neutron and one less proton than 14C.

Answer: (d)

5.

Which of the following atoms have the same number of neutrons?

(I) 233Th (II) 235U (III) 238U (IV) 238Np (a) I and II (b) I, II and IV (c) II and III (d) III and IV (e) All of these atoms have the same number of neutrons.

Answer: (a)

6.

Which statement correctly describes the difference between the 12C, 13C and 14C isotopes of carbon? (a) They have the same number of protons but different numbers of electrons. (b) They have the same number of neutrons but different numbers of electrons. (c) They have the same number of electrons but different numbers of protons. (d) They have the same number of protons but different numbers of neutrons. (e) They have the same number of protons, electrons, and neutrons but different atomic masses.

Answer: (d)

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Atoms and Ions 7.

What is the symbol for the atom or ion that has 34 neutrons, 29 protons, and 27 electrons? (a) 34Se2− (b) 29Cu2+ (c) 63Se2− (d) 63Cu2+ (e) 27Co2+

Answer: (d)

8.

How many electrons are in 52Cr3+? (a) 3 (b) 21 (c) 24 (d) 27 (e) 28

Answer: (b)

9.

How many protons are in 32S2–? (a) 2 (b) 14 (c) 16 (d) 18 (e) 32

Answer: (c) 10. Complete the following table.

Symbol

Atomic Number

Mass Number

Number of Protons

Number of Electrons

Number of Neutrons

37Cl

______

_______

_______

_______

_______

_____

10

_______

_______

10

12

18O2−

______

_______

_______

_______

_______

Answer: 17, 37, 17, 17, 20; 22Ne, 22, 10; 8, 18, 8, 10, 10

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11. Complete the following table Symbol

Mass Number

Protons

Neutrons

Electrons

19F

________

______

______

_______

_______

________

12

13

10

Answer: 19, 9, 10, 9; 25Mg2+, 25

12. Complete the table Symbol

Protons

Neutrons

Electrons

35Cl–

_________

_________

__________

_______

12

13

10

Answer: 17, 18, 18, 25Mg2+

13. Describe the difference on the atomic scale between the following three forms of oxygen: O, O2, O2−.

Answer: O represents an atom of oxygen containing 8 protons and 8 electrons.

O2 represents a molecule of oxygen made up of 2 atoms of oxygen bonded together.

O2- represents an anion of oxygen containing 8 protons and 10 electrons.

14. Which choice gives the diagrams in the order that best represents 2O 2, 4O, and X3, respectively?

A

B

C

(a) A B C (b) C B A (c) B C A (d) B A C 1-4

(e) None of the above gives the cartoons in the correct order.

Answer: (b) 15. Atoms are isobaric if they have the same mass. Which of the following atoms or ions are isobaric?

(I) 51Cr3+ (II) 52Cr3+ (III) 55Mn3+ (IV) 59Ni4+ (V) 59Co3+ (a) I and II (b) II and III (c) III and IV (d) IV and V (e) III, IV and V Answer: (d) 16. Which of the following correctly describes a Zn2+ ion that has a mass number of 65? (a) It contains 28 electrons, 30 protons, and 35 neutrons. (b) It contains 30 electrons, 30 protons, and 35 neutrons. (c) It contains 32 electrons, 30 protons, and 35 neutrons. (d) It contains 35 electrons, 35 protons, and 30 neutrons. (e) It contains 37 electrons, 35 protons, and 30 neutrons.

Answer: (a) 17. Calculate the number of electrons in a P3– ion if the mass number of phosphorus is 31 and the atomic number of this element is 15. (a) 12 (b) 15 (c) 18 (d) 31 (e) 34 Answer: (c)

18. Which of the following symbols indicates an atom or ion with 12 protons and 10 electrons? (a) Mg2+ (b) C4+ (c) O2– 1-5

(d) Ne (e) none of these Answer: (a)

19. All of the following have the same number of electrons except: (a) K+ (b) Ca2+ (c) Al3+ (d) S2– (e) Cl– Answer: (c)

20. What is the symbol for the atom or ion that has 26 protons, 30 neutrons, and 24 electrons? (a) 56Cr2+ (b) 54Cr2+ (c) 56Fe2+ (d) 54Fe3+ (e) 54Cr Answer: (c) 21. How many electrons are present in the 77Se2– ion? (a) 32 (b) 34 (c) 43 (d) 77 (e) none of these Answer: (e) 22. How many electrons are present in the 136Ba2+ ion? (a) 56 (b) 58 (c) 78 (d) 134 (e) none of these Answer: (e) 1-6

23. Calculate the number of electrons on a Ce4+ ion if the atomic number of this element is 58 and the mass number of the ion is 140. (a) 54 (b) 58 (c) 62 (d) 136 (e) 140 Answer: (a)

24. Which of the following compounds is composed only of polyatomic ions? (a) NH4NO3 (b) CaO (c) H2SO4 (d) NO2 (e) CO2 Answer: (a)

Metals, Nonmetals, and Semimetals 25. Which of the following is a nonmetal in the same group as silicon? (a) C (b) P (c) Sn (d) Zn (e) none of these Answer: (a)

26. Which of the following is a nonmetal in period 4? (a) C (b) Ca (c) Se (d) Sr (e) none of these Answer: (c)

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27. Which of the following elements is neither a metal nor a nonmetal? (a) Na (b) Mg (c) Ga (d) Si (e) P Answer: (d)

28. Classify the following elements as metals, nonmetals or semi-metals: S, Sb, Sc, Se, Sg, Si, Sm, Sn, and Sr. Answer: Nonmetals: S, Se Semi-metals: Sb, Si Metals: Sc, Sg, Sm, Sn, Sr

The Mass of an Atom The following two questions use the data listed below.

Isotope 16O 17O 18O

Abundance 99.76% 0.04% 0.20%

Mass (amu) 15.9949 16.9991 17.9992

29. If you were able to select a single atom of oxygen from a naturally occurring sample, what would its mass most likely be? (a) 15.9949 amu (b) 16.9991 amu (c) 17.9992 amu (d) 15.9993 amu (e) none of these Answer: (a)

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30. If you were able to select 1000 individual atoms of oxygen from a naturally occurring sample, what would be the average mass of the atoms? (a) 15.9949 amu (b) 16.9991 amu (c) 17.9992 amu (d) 15.9993 amu (e) none of these Answer (d)

The following two questions use the data listed below.

Isotope 20Ne 21Ne 22Ne

% Natural Abundance 90.51 0.27 9.22

Mass (amu) 19.99244 20.99384 21.99138

31. If you select a single random atom of Ne from a naturally occurring sample, its mass is most likely to be (a) 19.99244 amu (b) 20.99384 amu (c) 21.99138 amu (d) 20.179 amu (e) 20.000 amu Answer: (a) 32. If you select 100 random atoms of Ne from a naturally occurring sample, their total mass is most likely to be (a) 1999.244 amu (b) 2099.384 amu (c) 2199.138 amu (d) 2017.9 amu (e) 2000.0 amu Answer: (d)

Balancing Chemical Equations 33. What will be the coefficient for HNO3 when the following equation is balanced?

Ba(OH)2(aq) + HNO3(aq) → Ba(NO3)2(aq) + H2O(l) (a) 1 (b) 2 1-9

(c) 3 (d) 5 (e) none of these Answer: (b)

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34.

What will be the coefficient for oxygen when the following equation is balanced?

C3H8(g) + O2 (g) → CO2(g) + H2O(g) (a) 1 (b) 2 (c) 3 (d) 5 (e) none of these

Answer: (d)

35. What will be the coefficient for NH3 when the following reaction is balanced?

NH3(g) + O2(g) → N2(g) + H2O(l) (a) 1 (b) 3 (c) 4 (d) 5 (e) none of these Answer: (c)

36. Balance the following equations: (I) Na2SO4(s) + C(s) → Na2S(aq) + CO2(g) (II) Mg3N2(s) + H2O(l) → Mg(OH)2(aq) + NH3(g) Answer: (I) Na2SO4(s) + 2 C(s) → Na2S(s) + 2 CO2(g) (II) Mg3N2(s) + 6 H2O(l) → 3 Mg(OH)2(aq) + 2 NH3(g)

37. The balanced equation for the decomposition of ammonium dichromate has which of the following sets of coefficients?

A (NH4)2Cr2O7(s) → B Cr2O3(s) + C N2(g) + D H2O(g) (a) A = 2, B = 2, C = 2, D = 4 (b) A = 1, B = 1, C = 1, D = 4 (c) A = 1, B = 1, C = 1, D = 2 (d) A = 1, B = 1, C = 2, D = 2 (e) A = B = C = D = 1 Answer: (b)

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38. Calculate the sum of the coefficients in the balanced chemical equation for the following reaction: A H2S(g) + B O2(g) → C SO2(g) + D H2O(g) (a) 6 (b) 8 (c) 9 (d) 11 (e) 17 Answer: (c) 39. What is the sum of the coefficients when the following chemical equation is balanced?

A Ca3(PO4)2(s) + B C(s) → C Ca3P2(s) + D CO(g) (a) 6 (b) 12 (c) 18 (d) 20 (e) none of the above Answer: (c) 40. Which illustration best represents the following equation?

2H2 + O2

(a)

2H2O

+

(b)

+

+

(c)

+

+

(d)

+

+ +

(e)

+

+

+

+

+

+ +

+ +

Answer: (c)

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+

+

Chapter 2: The Mole The Mole 1.

Which of the following contains the largest number of hydrogen atoms? (a) 1 mol of H2O (b) 0.5 mol of NH3 (c) 0.20 mol of CH4 (d) 0.20 mol of C6H12O6 (e) 0.01 mol of H3PO4

Answer: (d) 2.

How many platinum atoms does 1.00 g of pure platinum contain? (a) 195 Pt atoms (b) 3.09 x 1021 Pt atoms (c) 6.17 x 1021 Pt atoms (d) 1.95 x 1023 Pt atoms (e) 6.02 x 1023 Pt atoms

Answer: (b)

3.

There are 16 ounces in a lb and 2.20 lbs in a kg. What would be the value of Avogadro's number if this constant was defined as the number of 12C atoms in 12 ounces of 12C? (a) 0.171 x 1023 (b) 6.02 x 1023 (c) 28.4 x 1023 (d) 171 x 1023 (e) 2050 x 1023

Answer: (d)

4.

If the atomic weight of platinum is 195.08 g/mol, what is the mass in grams of a single platinum atom? (a) 1.62 x 10–22 g (b) 3.24 x 10–22 g (c) 5.13 x 10–3 g (d) 195.08 g (e) none of the above

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Answer: (b) 5.

What is the molar mass of an element for which a single atom weighs 3.95 x 10–22 g? (a) 6.56 x 10–46 (b) 1.31 x 10–45 (c) 3.95 x 10–22 (d) 238 (e) 476

Answer: (d) 6.

What is the mass in grams of 0.250 moles of magnesium? (a) 0.0103 g (b) 6.08 g (c) 13.8 g (d) 97.2 g (e) none of these

Answer: (b)

7.

Which of the following represents the largest number of moles of molecules? (a) 10 g NH3 (b) 10 g N2 (c) 10 g Cu (d) 10 g Br2 (e) all of these contain the same number of moles

Answer (a)

8.

How many moles of oxygen atoms are in 0.500 grams of CO2? (a) 0.500 x 10−2 (b) 1.00 (c) 0.250 (d) 1.14 x 10−2 (e) 2.27 x 10−2

Answer: (e)

9.

Arrange the following in order of increasing mass:

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(1) one F atom (2) 1 x 10−20 mol of F (3) 1 x 10−20 g of F (4) one F2 molecule (a) 1 < 2 < 3 < 4 (b) 1 < 4 < 3 < 2 (c) 3 < 1 < 2 < 4 (d) 2 < 1 < 3 < 4 (e) none of these Answer: (b)

10. What is the molar mass of barium nitrate, Ba(NO3)2? (a) 261 g/mole (b) 199 g/mole (c) 213 g/mole (d) 167 g/mole (e) none of these Answer: (a)

11. What is the mass in grams of 0.125 moles of calcium chloride, CaCI2. (a) 111 g (b) 13.9 g (c) 8.88 x 102 g (d) 1.13 x 10–3 g (e) 41.7 g Answer: (b)

12. How many moles of phosphorus atoms are in 100 grams of P4S10? (a) 4.00 (b) 4.00 x 102 (c) 0.225 (d) 0.487 (e) 0.900 Answer: (e)

13. Determine the number of moles in 12.5 g of carbon tetrachloride, CCl4. (a) 8.13 x 10–2 mol

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(b) 12.3 mol (c) 1.92 x103 mol (d) 7.53 x 1024 mol (e) none of these Answer: (a)

14. What is the mass in grams of of 0.158 moles of Cu(NO3)2? (a) 2.96 x 101g (b) 8.42 x 10–4 g (c) 1.19 x 103 g (d) 12.3 g (e) none of these Answer: (a)

15. Calculate the number of moles in 15.0 grams of NaOH. (a) 40.0 mol (b) 6.00 x 102 mol (c) 2.67 mol (d) 0.375 mol (e) none of these Answer: (d) 16. Which of the following contains more atoms of nitrogen, 100 moles of NH3 or 100 moles of N2O5? Explain your reasoning without showing calculations.

Answer: 100 moles of N2O5 because there are two moles of nitrogen per mole of N2O5 molecules while in NH3 there is only one mole of nitrogen atoms per mole of molecules.

17. How many moles of Na2S are found in 20.0 grams of sodium sulfide, Na2S? (a) 0.256 moles (b) 0.364 moles (c) 2.75 moles (d) 3.90 moles (e) none of these Answer: (a)

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18. Calculate the number of nitrogen atoms in 0.125 moles of barium nitrate, Ba(NO3)2. (a) 2.00 atoms (b) 0.125 atoms (c) 0.250 atoms (d) 1.51 x 1023 atoms (e) 7.53 x 1022 atoms Answer: (d)

19. What is the molecular weight of aspirin, C9H8O4? (a) less than 50 amu (b) between 50 and 80 amu (c) between 80 and 120 amu (d) between 120 and 160 amu (e) more than 160 amu Answer: (e)

20. Cocaine is extracted from the leaves of plants of the species E. coca, whereas chocolate is extracted from the seeds of the species E. cacao.

(Never confuse one with the other!) Calculate the molar mass of cocaine, assuming a molecular formula of C17H21NO4. (a) between 0 and 100 g/mol (b) between 100 and 250 g/mol (c) between 250 and 300 g/mol (d) between 300 and 350 g/mol (e) more than 350 g/mol Answer: (d)

21. What is the mass in amu of one molecule of sucrose (cane sugar) that has the formula: C12H22O11? (a) less than 100 amu (b) between 100 and 150 amu (c) between 150 and 250 amu (d) between 250 and 350 amu (e) more than 350 amu Answer: (d)

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22. What is the mass in grams of one molecule of CO2? (a) 4.6 x 10–23 g (b) 7.3 x 10–23 g (c) 28 g (d) 44 g (e) none of the above Answer: (b) 23. The drug for which more prescriptions – and more bad jokes – were written than any drug released recently has the scientific name Sildenafil citrate and the trade name Viagra™. The structure of Sildenafil is shown below.

O

H3 CH2 CO

CH3

HN NH N CH2 CH2 CH3

O2S N N CH3

When this structure is translated into a chemical formula we get a molecular formula of C22H32N6O4S. Calculate the molecular weight of this compound. (a) between 100 and 200 g/mol (b) between 200 and 300 g/mol (c) between 300 and 400 g/mol (d) between 400 and 500 g/mol (e) more than 500 g/mol Answer: (d)

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24. Let’s consider another “miracle drug” with the scientific name orlistat, which is marketed by Roche Pharmaceuticals as Xenical®. The formal chemical name for Xenical® is (s)-2-formylamino-4-methyl-pentanoic acid (s)-1-[[(2S, 3S)-3-hexyl-4-oxo-2-oxetanyl] methyl]-dodecyl ester. Its structure is shown below.

H N

O H

O

O O O

The molecular formula for this wonder diet drug that is supposed to block the uptake of fat is C29H35NO5. Calculate the molecular weight of this compound. (a) between 100 and 200 g/mol (b) between 200 and 300 g/mol (c) between 300 and 400 g/mol (d) between 400 and 500 g/mol (e) more than 500 g/mol Answer: (d) 25. If you eat too much fat while taking Xenical®, you’re going to be very, very “uncomfortable.” You can avoid this problem, of course, by taking the other new diet drug, Sibutramine hydrochloride monohydrate, which is marketed by a subsidiary of BASF as Meridia®. Meridia® is a neurotransmitter that inhibits the uptake of other neurotransmitters, such as dopamine, norepinephrine, and serotonin. It therefore tends to depress one’s appetite and is being studied as a possible treatment for depression. Elemental analysis suggests that this compound is 72.96% C, 9.36% H, 12.67% Cl, and 5.01% N by weight. What is the total number of C, H, N and Cl atoms in the molecular formula of this compound if the molecular weight is 279.85 g/mol? (a) 10 or less (b) between 11 and 20 (c) between 21 and 30 (d) between 31 and 40 (e) between 41 and 50 Answer: (e)

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26. ADM (Archer Daniels Midland) the self-proclaimed “Supermarket to the World” has been in trouble for allegedly fixing the price of lysine, an amino acid that is often deficient in vegetarian diets. Lysine is 19.16% N by weight and each molecule contains two nitrogen atoms. Calculate the molecular weight of lysine. (a) less than 100 g/mol (b) between 100 and 150 g/mol (c) between 150 and 250 g/mol (d) between 250 and 350 g/mol (e) more than 350 g/mol Answer: (b) 27. Calculate the number of carbon atoms in a mole of sucrose, C12H22O11. (a) 12 atoms (b) 6.02 x 1023 atoms (c) 6.62 x 1024 atoms (d) 7.22 x 1024 atoms (e) 1.3 x 1025 atoms Answer: (d) 28. Calculate the number of chlorine atoms in 10.4 grams of chloroform, CHCl3. (a) 5.25 x 1022 atoms (b) 1.05 x 1023 atoms (c) 1.57 x 1023 atoms (d) 2.10 x 1023 atoms (e) 2.07 x 1025 atoms Answer: (c)

29. Sodium tripolyphosphate, Na5P3O10, is added to detergents to increase their cleaning power. Calculate the number of phosphorus atoms in 0.325 moles of this compound. (a) 6.52 x 1022 atoms (b) 1.96 x 1023 atoms (c) 5.87 x 1023 atoms (d) 6.02 x 1023 atoms (e) none of the above Answer: (c)

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Percent Mass 30. Acetic acid, C2H4O2, is the substance that makes vinegar taste sour. What is the % mass of carbon in acetic acid? (a) 20% (b) 25% (c) 35% (d) 40% (e) none of these Answer: (d)

31. Calculate the % mass of sodium in sodium sulfate, Na2SO4. (a) 16% (b) 23% (c) 32% (d) 54% (e) 67% Answer: (c)

32. A hydrate is heated to determine the % water by mass in the hydrate.

During heating of the hydrate a small chip is broken off of the lid of the crucible and falls to the lab bench without being noticed. What effect will this have on the calculated value of the % water by mass? (a) It will be high. (b) It will be low. (c) It will not affect the results. (d) It could be either high or low.

Answer: (a)

33. Which of the following nitrogen-containing compounds would supply the most nitrogen per gram of fertilizer? (a) iron azide, Fe(N3)2, MW = 140 g/mol (b) sodium azide, NaN3, MW = 65 g/mol (c) potassium nitrite, KNO2, MW = 85.1 g/mol (d) potassium azide, KN3, MW = 81.1 g/mol (e) There is not enough information to answer this problem.

Answer: (b)

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34. Apatite is a mineral that is found in tooth enamel. When fluoride toothpastes are used, this mineral is converted to fluoroapatite, which is much harder and therefore more resistant to decay. What is the percent by weight of fluorine in fluoroapatite, Ca5(PO4)3F? (a) between 0 and 0.1% (b) between 0.1% and 1% (c) between 1% and 3% (d) between 3% and 5% (e) more than 5% Answer: (d)

35. In March of 2007 a number of pets were sickened or killed by pet food contaminated with the chemical melamine which is used in the manufacture of certain plastics. It is thought that the melamine was added to make a pet food additive appear to have more protein than it really did. Protein content is determined by measuring nitrogen content of food since protein is the main source of nitrogen in food. What is the % by mass of nitrogen in melamine which has the formula C3H6N6? (a) 66.6% (b) 61.2% (c) 11.1% (d) 82.3% (e) Insufficient information is given to find the %N by mass.

Answer: (a)

36. In March of 2007 a number of pets were sickened or killed by pet food contaminated with the chemical melamine (C3H6N6) which is used in the manufacture of certain plastics. It is thought that the melamine was added to make a pet food additive appear to have more protein than it really did.

Protein content is determined by measuring the %N by mass in food because protein is the main source of nitrogen in food. Proteins are composed of amino acids so the %N by mass of a protein reflects the %N of the individual amino acids. Alanine is a typical amino acid and has the formula C3H7NO2. Calculate the %N by mass in melamine and alanine.

Why would adding a small amount of melamine make something appear to have a high protein content?

Answer: %N in melamine = 66.63% %N in alanine = 15.72%

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Melamine has a very high %N by mass so adding a small amount of this compound to food will significantly increase the %N in the food. Since protein content is calculated from the %N, the melamine will make the protein content artificially high.

Determining the Formula of a Compound 37. Determine the empirical formula of a compound containing 59.0% Na and 41.0% O by mass. (a) NaO (b) Na2O (c) NaO2 (d) Na2O3 (e) none of these Answer: (a) 38. Cobalt oxide used in glazing pottery contains 71.06% cobalt and 28.94% oxygen by mass. What is the empirical formula of this cobalt oxide? (a) CoO (b) Co2O (c) CoO3 (d) Co2O3 (e) Co3O4 Answer: (d) 39. Calculate the molecular weight of the oxygen-carrier protein hemoglobin if this protein is 0.335% Fe by weight and each protein molecule contains four iron atoms. (a) less than 1000 grams per mole (b) between 1000 and 10,000 grams per mole (c) between 10,000 and 50,000 grams per mole (d) between 50,000 and 100,000 grams per mole (e) more than 100,000 grams per mole Answer: (d)

40. Calculate the number of carbon atoms in the empirical formula of nicotine if this compound is 74.0% C, 8.7% H, and 17.3% N by weight. (a) 3 atoms (b) 4 atoms (c) 5 atoms

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(d) 7 atoms (e) 10 atoms Answer: (c)

41. Glucose (blood sugar), acetic acid, and formaldehyde are all 40.0% C, 6.76% H, and 53.3% O by weight. Calculate the total number of carbon, hydrogen, and oxygen atoms in the empirical formula for these compounds. (a) 3 (b) 4 (c) 5 (d) 6 (e) 9 Answer: (b)

42. A compound that is 43.64% P and 56.36% O by weight has a molecular weight of 283.88 g/mol. What is the molecular formula of this compound? (a) PO4 (b) P2O3 (c) P2O5 (d) P4O10 (e) none of the above Answer: (d) 43. A mixed oxide of potassium and vanadium is 28.3% by weight potassium and 37.0% by weight vanadium. What is the empirical formula? (a) KV (b) KVO (c) K2V3O (d) K3V3O (e) KVO3 Answer: (e)

44. The chief ore of manganese is an oxide known as pyrolusite, which is 36.8% O and 63.2% Mn by weight. Which of the following is the correct empirical formula for pyrolusite? (a) MnO (b) MnO2

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(c) Mn2O3 (d) MnO3 (e) Mn2O7 Answer: (b) 45. Nitrogen combines with oxygen to form a variety of compounds. One of these compounds is called nitrous oxide or “laughing gas.” What is the formula of nitrous oxide if this compound is 63.65% N and 36.35% O by weight? (a) N2O (b) NO (c) NO2 (d) N2O3 (e) N2O5 Answer: (a) 46. What is the atomic weight of a metal that forms an oxide, M2O3, which is 17.29% oxygen by weight? (a) 38.3 g/mol (b) 57.4 g/mol (c) 76.6 g/mol (d) 115 g/mol (e) 229 g/mol Answer: (d) 47. Cisplatin is a commercial antitumor drug that contains 65.0% Pt, 23.7% Cl, 9.3% N and 2.0% H by weight. What is the ratio of H to Pt atoms in the empirical formula of this compound? (a) 1:1 (b) 2:1 (c) 3:1 (d) 4:1 (e) 6:1 Answer: (e)

48. What is the atomic weight in amu of the element that forms a fluoride, M 2F4, which is 73.1% fluorine by weight? (a) 14.0 amu 2-13

(b) 16.0 amu (c) 28.0 amu (d) 32.0 amu (e) none of these Answer: (a)

49. What is the empirical formula of a compound that contains C, H, and O if the compound is 40.00% C and 6.71% H by weight? (a) CHO (b) CH2O (c) CH3O (d) CH4O (e) none of the above Answer: (b)

50. Calculate the empirical formula of naphthalene if this compound is 93.71% carbon and 6.29% hydrogen by weight. (a) CH (b) CH2 (c) C3H2 (d) C4H5 (e) C5H4 Answer: (e)

51. The molecular weight of naphthalene is 128.17 grams per mole. Use the results of the previous question to calculate the total number of atoms in a naphthalene molecule. (a) 9 atoms (b) 12 atoms (c) 18 atoms (d) 27 atoms (e) 38 atoms Answer: (c) 52. The term “carbohydrate” once meant a compound with the empirical formula CH2O. What is the molecular formula of glucose (blood sugar) if the molecular weight of this carbohydrate is 180 g/mol?

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(a) CH2O (b) C4H8O4 (c) C6H12O6 (d) C12H22O11 (e) none of the above Answer: (c)

53. A 1.854 g sample of D-ribose contains 0.01236 moles of this sugar which has an empirical formula of CH2O. What is the ratio of the molecular formula to the empirical formula of this compound? (a) 1 (b) 2 (c) 3 (d) 5 (e) none of these Answer: (d) 54. The ribonucleic acids (RNA) that are so important in the biosynthesis of proteins all contain the sugar D-ribose, which contains three elements: C, H, and O. Elemental analysis suggests that D-ribose is 40.00% C and 6.71% H by weight. What is the empirical formula of this compound? (a) CHO (b) CH2O (c) CH3O (d) CH3O2 (e) none of these Answer: (b) 55. Halothane is an anesthetic that is 12.17% C, 0.51% H, 40.48% Br, 17.96% Cl and 28.87% F by weight. Calculate the molecular formula of this compound if each molecule contains one hydrogen atom.

Answer: C2HBrClF3

56. A compound frequently used as a high-temperature lubricant has the formula MS2, where M represents some metallic element. If this compound is 40.06% sulfur by mass, what is the identity of the metal M? (a) K (b) Fe

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(c) Cu (d) Mo (e) Pb Answer: (d)

57. In small quantities, the nicotine in tobacco is addictive. In large quantities, it is a deadly poison. Calculate the molecular formula of nicotine, CxHyNz, if the molecular weight is 162.2 g/mol and 0.438 grams of this compound burn to form 1.188 grams of CO2 and 0.341 grams of water.

Answer: C10H14N2

58. In 1763, the Reverend Edmund Stone suggested chewing the bark of the English willow (Salix abla) to treat a fever because it grows in moist regions where one was likely to catch a fever, and “... many natural maladies carry their cures along with them, or their remedies lie not far from their causes.”

It took 50 years to isolate the active ingredient in the bark, which was named salicylic acid. By the end of the 19th century, salicylic acid was used to treat rheumatic fever, gout and arthritis. Many patients complained of chronic stomach irritation, however. Because his father was one of these patients, Felix Hoffman searched for a less acidic derivative of salicylic acid.

In 1898, Hoffman reported that the acetyl ester of salicylic acid was both more effective and easier to tolerate than the parent compound. He named the compound aspirin, taking the prefix a- from the name of the acetyl group and spirin from the German name of the parent compound spirsure.

Aspirin is 60.00% C, 4.48% H, and 35.5% O by weight. What is the total number of C, H and O atoms in the empirical formula of this compound? (a) less than 15 atoms (b) between 15 and 20 atoms (c) between 20 and 25 atoms (d) between 25 and 35 atoms (e) more than 35 atoms Answer: (c)

59. One of the concerns during the war against Iraq several years ago was fear of chemical warfare agents, such as the nerve agent VX. VX is the most dangerous material ever synthesized; a single drop of VX on the skin will kill an individual within hours. More than 2 million pounds of VX were made in Southern Indiana by 1968. They are still there, stored in 1600 canisters, packed three high. The canisters were welded together several years ago because of fear of exposure to VX if a tornado struck the storage facility.

Elemental analysis of VX gives the following result: 49.41% C, 9.80% H, 2-16

5.24% N, 11.97% O, 11.58% P, and 11.99% S. How many carbon atoms does the empirical formula for this compound contain?

(Hint: Carry out the calculation to at least three significant figures.) (a) 3 or less atoms (b) 4 to 7 atoms (c) 8 to 11 atoms (d) 12 to 15 atoms (e) more than 15 atoms Answer: (c) 60. Assume that the empirical formula of VX (described in the previous question) is the same as the molecular formula. What is the molecular weight of VX? (a) less than 150 grams/mole (b) between 150 and 200 grams/mole (c) between 200 and 250 grams/mole (d) between 250 and 300 grams/mole (e) more than 300 grams/mole Answer: (d) 61. On March 20, 1995, terrorists released an organophosphate nerve gas known as “Sarin” at several points in the Tokyo subway system, killing 11 and injuring more than 5,500 people. Sarin was concealed in lunch boxes and soft-drink containers and placed on subway train floors. It was released as terrorists punctured the containers with umbrellas before leaving the trains. Elemental analysis of Sarin gives the following result: 34.29% C, 7.19% H, 22.84% O, 13.56% F, 22.11% P. How many carbon atoms does the empirical formula of this compound contain? (a) 2 or less atoms (b) 3 to 6 atoms (c) 7 to 10 atoms (d) 11 to 15 atoms (e) more than 15 atoms Answer: (b)

62. A sample of copper(II) sulfate, CuSO4, that weighs 2.47 grams picks up water from the atmosphere to form a hydrate with the formula CuSO4 • x H2O. If the sample weighs 3.86 grams after it picks up water, what is the value of x?

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(a) 2 (b) 2.5 (c) 3 (d) 4 (e) 5 Answer: (e)

63. The thermal decomposition of carbonates leads to the loss of CO2. The decomposition of an unknown carbonate leads to a 35.1% weight loss. The unknown was which of the following compounds. (a) Li2CO3 (b) MgCO3 (c) CaCO3 (d) ZnCO3 (e) BaCO3 Answer: (d)

64. In March of 2007 a number of pets were sickened or killed by pet food contaminated with the chemical melamine which is used in the manufacture of certain plastics. It is thought that the melamine was added to make a pet food additive appear to have more protein than it really did. Melamine is 28.57% C, 4.80% H, and 66.63% N by mass. What is melamine’s empirical formula? (a) C9H5N7 (b) CHN2 (c) C2HN7 (d) CH2N2 (e) CHN Answer: (d)

65. In March of 2007 a number of pets were sickened or killed by pet food contaminated with the chemical melamine which is used in the manufacture of certain plastics. It is thought that the melamine was added to make a pet food additive appear to have more protein than it really did. Melamine has the empirical formula CH2N2. If the molecular weight of melamine is 126.15 g/mol, what is its molecular formula? (a) CH2N2 (b) CH3N3 (c) C3H6N6

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(d) C2H4N4 (e) C4H8N8 Answer: (c)

Mole Ratios and Chemical Equations 66. If 0.50 moles of PbS are combined with 0.50 moles of O2, how many moles of PbO can be produced by the following reaction?

2 PbS(s) + 3 O2(g) → 2 PbO(s) + 2 SO2(g) (a) 0.33 (b) 0.50 (c) 0.75 (d) 2.0 (e) none of these Answer: (a) 67. What would happen to the potential yield of sulfur dioxide in the previous question if the amount of oxygen was doubled? (a) It would decrease by a factor of 2. (b) It would decrease by a factor of 1.5. (c) It would remain constant. (d) It would increase by a factor of 1.5. (e) It would increase by a factor of 2.

Answer: (d)

68. Acetylene gas, C2H2, burns in the presence of oxygen by the following reaction: 2 C2H2(g) + 5 O2 (g) → 2 H2O(l) + 4 CO2(g) If 4.5 moles of acetylene gas are to be burned, how many moles of oxygen will be required for the complete reaction of the acetylene? (a) 4.5 moles (b) 23 moles (c) 8.1 moles (d) 5.0 moles (e) none of these Answer: (e)

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69. For the reaction: 4 NH3(g) + 5 O2(g) → 4 NO(g) + 6 H2O(l), how many moles of ammonia, NH3, are required to react completely with 0.250 moles of oxygen? (a) 0.400 (b) 0.500 (c) 0.200 (d) 0.313 (e) 0.178 Answer: (c)

70. The molecular formula of isooctane is C8H18. It burns in the presence of oxygen to form a mixture of CO2 and H2O. What is the ratio of moles of water to moles of carbon dioxide produced in this reaction? (a) 1:1 (b) 2:1 (c) 9:4 (d) 9:8 (e) none of the above Answer: (d)

71. A laboratory method for preparing oxygen (O2) involves the decomposition of KClO3.

KClO3(s) → 2 KCI(s) + 3 O2(g) How many moles of oxygen are produced when 0.135 moles of KCl are produced by this reaction? (a) 0.135 mol (b) 0.090 mol (c) 0.203 mol (d) 0.405 mol (e) none of these Answer: (c)

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Stoichiometry 72. What mass of O2, in grams, is produced in the decomposition of 100 g of KClO3 by the following reaction?

2 KClO3(s) → 2 KCl(s) + 3 O2(g) (a) 0.816 grams (b) 1.22 grams (c) 39.2 grams (d) 1.50 x 102 grams (e) 26.1 grams Answer: (c)

73. Iron ore is impure Fe2O3. When Fe2O3 is heated with an excess of carbon (coke), iron metal is produced. How many moles of Fe can be produced from 250 g of Fe2O3 by the following reaction?

Fe2O3 (s) + 3 C(s) → 2 Fe(l) + 3 CO(g) 2 (a) 5.00 x 10 mol (b) 125 mol (c) 1.57 mol (d) 3.13 mol (e) 0.785 mol Answer: (d)

74. Calculate the number of moles of iron, Fe, produced when 0.500 moles of carbon monoxide, CO, react with an excess of iron oxide.

Fe2O3(s) + 3 CO(g) → 2 Fe(s) + 3 CO2(g) (a) 2.00 moles (b) 1.00 moles (c) 0.750 moles (d) 0.166 moles (e) none of these Answer: (e)

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75. Hydrogen chloride, HCI, can be made by the reaction of phosphorus trichloride, PCl3, with water and then boiling the HCI gas out of the solution.

Calculate the mass of HCI gas that can be prepared from 15.0 g of PCl3 and an excess of water.

PCl3(g) + 3 H2O(l) → 3 HCI(aq) + H3PO3(aq) (a) 0.328 g (b) 3.97 g (c) 11.9 g (d) 35.7 g (e) 3.00 g Answer: (c)

76. Ethyl alcohol, C2H5OH, is also known as drinking alcohol. It can be produced by the fermentation of glucose by the following reaction.

C6H12O6(aq) → 2 C2H5OH(aq) + 2 CO2(g) How many moles of ethanol can be produced from 100 g of glucose? (a) 100 moles (b) 200 moles (c) 0.556 moles (d) 1.11 moles (e) none of these Answer: (d)

77. Nitrogen reacts with red-hot magnesium to form magnesium nitride, 3 Mg(s) + N2(g) → Mg3N2(s) which reacts with water to form magnesium hydroxide and ammonia, Mg3N2(s) + 6 H2O(l) → 3 Mg(OH)2(aq) + 2 NH3(aq) How many grams of magnesium would you have to start with to prepare 15.0 grams of ammonia? (a) 13.3 g (b) 15.0 g (c) 20.0 g (d) 32.2 g (e) none of the above Answer: (d)

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78. The chemical formula of ethanol is C2H6O. It burns in the presence of excess oxygen to form CO2 and H2O. How many grams of H2O are produced from the combustion of 25.0 g of ethanol? (a) 9.78 g (b) 18.0 g (c) 25.0 g (d) 29.3 g (e) 54.1 g Answer: (d) 79. Sucrose, or “cane sugar,” has the formula C12H22O11. What weight of carbon dioxide can be produced from 10.0 grams of sucrose and 10.0 grams of oxygen?

C12H22O11(s) + 12 O2(g) → 12 CO2(g) + 11 H2O(g) (a) less than 5 grams (b) between 5 and 10 grams (c) between 10 and 15 grams (d) between 15 and 20 grams (e) more than 20 grams Answer: (d)

80. PF3 reacts with XeF4 to give PF5. In theory, how many moles of PF5 can be produced from 100.0 g of PF3 and 50.0 g of XeF4?

2 PF3(g) + XeF4(s) → 2 PF5(g) + Xe(g) (a) 0.121 mol (b) 0.241 mol (c) 0.482 mol (d) 1.14 mol (e) 2.28 mol Answer: (c)

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81. The portable stoves campers use for cooking burn propane, C3H8.

A C3H8(g) + B O2(g) → C CO2(g) + D H2O(g) What weight of propane would have to be burned to form 7.26 g of CO 2? (a) 2.42 g (b) 4.84 g (c) 7.26 g (d) 21.8 g (e) none of the above Answer: (a) 82. Calculate the weight of O2 needed to burn 10.0 grams of hydrogen sulfide, H2S. A H2S + B O2 → C SO2 + D H2O (a) 4.70 g (b) 6.26 g (c) 7.04 g (d) 9.39 g (e) 14.1 g Answer: (e) 83. How many grams of CO2 gas are produced from the combustion of 10.0 grams of isooctane?

2 C8H18 + 25 O2 → 16 CO2 + 18 H2O (a) 1.58 g (b) 3.85 g (c) 14.2 g (d) 30.8 g (e) 35.0 g Answer: (d)

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84. A crucible and sample of CaCO3, weighing 42.670 g, was heated until red hot in order to decompose the sample.

CaCO3(s) → CaO(s) + CO2(g) The empty crucible weighed 35.351 g. What is the theoretical weight of the crucible and residue after the decomposition is complete? (a) 3.219 g (b) 4.100 g (c) 38.570 g (d) 39.451 g (e) 49.989 g Answer: (d)

85. Nitrogen reacts with hydrogen to form ammonia N2(g) + 3 H2(g) → 2 NH3(g) which burns in the presence of oxygen to form nitrogen oxide, 4 NH3(g) + 5 O2(g) → 4 NO(g) + 6 H2O(g) which reacts with excess oxygen to form nitrogen dioxide.

2 NO(g) + O2(g) → 2 NO2(g) How many moles of nitrogen would we have to start with to make 10 moles of nitrogen dioxide? (a) 2.5 mol (b) 5 mol (c) 10 mol (d) 15 mol (e) 20 mol Answer: (b) 86. How many grams of hydrogen peroxide, H2O2, must decompose by the following reaction to produce 0.400 moles of O2?

2 H2O2(aq) → 2 H2O(l) + O2(g) (a) less than 25 grams (b) between 25 and 40 grams (c) between 40 and 60 grams (d) between 60 and 80 grams (e) more than 80 grams Answer: (b) 2-25

Limiting Reagents 87. In the following reaction 10.0 grams of lead sulfide, PbS are allowed to react with 5.00 grams of oxygen, O2. Determine the limiting reagent, the number of grams of PbO that can be produced and the number of moles of reactant that remains unreacted.

2 PbS(s) + 3 O2(g) → 2 PbO(s) + 2 SO2(g) Answer: PbS, 9.33 g PbO, 0.0935 moles O2 remaining

88. Iron metal reacts with chlorine gas by the following reaction.

2 Fe(s) + 3 Cl2 → 2 FeCl3 (I) Predict the moles of product formed if 31.67 grams of Fe reacts with an excess of chlorine gas, Cl2.

(II) Predict the moles of product formed if 0.567 moles of Fe reacts with 0.718 moles of chlorine gas, Cl2.

(III) What is fundamentally different about how the above two problems must be solved?

Answer: (I) 0.567 moles (II) 0.479 moles (III) In part (b) you must find the limiting reagent.

89. How many grams of HCl are produced when 10.0 g of Cl2 and 1.00 g of H2 react?

H2(g) + Cl2(g) → 2 HCl(g) (a) 5.14 g (b) 9.04 g (c) 10.3 g (d) 11.0 g (e) 18.1 g Answer: (c) 90. What would happen to the amount of HCl produced in the previous question if the amount of H2 is doubled? (a) It would decrease by more than a factor of two. (b) It would decrease by a factor of two. (c) It would remain the same. (d) It would increase by a factor of two.

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(e) It would increase by more than a factor of two.

Answer: (c)

Questions 91-92 refer to the following reaction: 2 NO(g) + O2(g) → 2 NO2(g) 91. What is the limiting reagent when 10.0 g of NO react with 10.0 g of O 2 to form NO2? (a) NO (b) O2 (c) NO2 (d) none of these (e) There isn’t enough data to answer the question.

Answer: (a)

92. How many moles of the limiting reagent are present in the previous problem? (a) 0.313 mol (b) 0.333 mol (c) 0.435 mol (d) 0.625 mol (e) none of these Answer: (b)

93. How many grams of MgO can be produced by burning 10.0 grams of Mg in the presence of 10.0 grams of O2? (a) 10.0 g (b) 12.6 g (c) 16.6 g (d) 20.0 g (e) 25.2 g Answer: (c) 94. When glucose reacts with oxygen in living systems, carbon dioxide and water are produced, and a great deal of energy is liberated.

C6H12O6(s) + 6 O2(g) → 6 CO2(g) + 6 H2O(g)

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What weight of carbon dioxide can be produced from the reaction of 10.0 grams of glucose with 10.0 grams of oxygen? (a) 2.29 g (b) 2.44 g (c) 13.8 g (d) 14.7 g (e) none of these Answer: (c)

Density 95. The density of sugar is 1.59 g/mL. Calculate the volume of a teaspoon of sugar if this quantity of sugar weighs 4.50 grams. (a) 0.140 mL (b) 0.353 mL (c) 2.83 mL (d) 6.09 mL (e) 7.16 mL Answer: (c) 96. Iron has a density of 7.86 g/cm3. When iron rusts, it forms hydrated Fe2O3 that has a density of 5.12 g/cm3. What happens to the material as iron rusts? (a) It expands. (b) It contracts. (c) The volume must remain the same. (d) The iron evaporates. (e) The water evaporates.

Answer: (a)

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97. The first silver dollar minted in the Americas was 3.95 cm in diameter and contained 27.07 grams of silver. If the density of silver is 10.5 g/cm 3, what was the height of each coin? (a) 0.00791 cm (b) 0.0316 cm (c) 0.0526 cm (d) 0.210 cm (e) none of these Answer: (d) 98. The density of water in pounds per cubic foot is 62.4 lb/ft3. If a 2 ft3 bag of mulch sold at the local lawn care center weighs 30 lbs, which of the following correctly describes what happens when this mulch is poured on water? (a) The mulch will sink because it is less dense than water. (b) The mulch will sink because it is more dense than water. (c) The mulch will float because it is less dense than water. (d) The mulch will float because it is more dense than water. (e) There is no way of predicting what will happen.

Answer: (c)

99. A sample of sucrose (C12H22O11) that weighs 9.35 g is placed in a 15.00 mL flask. The remaining volume of the flask is filled with benzene — a liquid in which sucrose does not dissolve. The density of benzene is 0.879 g/cm 3.

The sucrose and benzene together weigh 17.33 g. What is the density of sucrose? (a) 0.532 g/cm3 (b) 0.623 g/cm3 (c) 1.03 g/cm3 (d) 1.16 g/cm3 (e) 1.58 g/cm3 Answer: (e)

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100. The following liquids don’t mix with water. Which of these liquids would float on top of water?

(I) benzene, C6H6, d = 0.877 g/cm3 (II) carbon disulfide, CS2, d = 1.263 g/cm3 (III) chloroform, CHCl3, d = 1.483 g/cm3 (IV) ether, (C2H5)2O, d = 0.714 g/cm3 (V) gasoline, d = 0.675 g/cm3 (a) I, II and III (b) II and III (c) I, IV and V (d) None of these liquids would float on top of water. (e) All of these liquids would float on top of water.

Answer: (c)

101. If a beaker that contains 250 mL of chloroform (CHCl3) weighs 524 grams, what is the weight of the beaker? (CHCl3: d = 1.483 g/cm3) (a) about 50 g (b) about 100 g (c) about 150 g (d) about 250 g (e) about 500 g Answer: (c)

102. A student was given three pieces of metal that looked like brass. Sample I had a volume of 17.3 mL and weighed 154.3 g, sample II had a volume of 28.4 mL and weighed 547.2 g, and sample III had a volume of 22.2 mL and weighed 186.1 g. Which sample(s) was (were) brass? (Brass: d  8 g/cm3) (a) only I (b) I and II (c) I and III (d) II and III (e) All of these samples were brass.

Answer: (c)

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103. Retailers purchase gasoline by weight so they get a constant amount of gasoline per dollar (as long as the price is constant). We buy gasoline by volume. 1.000 kg of gasoline occupies a volume of 1.356 L at 15.6oC. What volume of gasoline would you have to purchase to get the same weight at 32.2oC when gasoline has a density of 0.72210 g/mL? (a) 0.9792 L (b) 0.7221 L (c) 1.848 L (d) 1.385 L (e) 1.215 L Answer: (d)

Concentration & Molarity 104. The amount of solute in each of the beakers shown below is represented by the number of black dots. The volume is shown on the side. In which beaker does the solute have the highest concentration?

20 mL

20 mL

20 mL

20 mL

15 mL

15 mL

15 mL

15 mL

10 mL

10 mL

10 mL

10 mL

5 mL

5 mL

5 mL

5 mL

A

B

C

(a) Beaker A (b) Beaker B (c) Beaker C (d) Beaker D (e) The solute concentration is the same in each beaker.

Answer: (d)

105. What is the molarity of 0.0250 liters of a solution containing 7.50 grams of a compound whose molar mass is 152.8 g/mole? (a) 0.125 M (b) 0.509 M (c) 0.491 M 2-31

D

(d) 1.96 M (e) 3.92 M Answer: (d)

106. What weight of HNO3 is present in 13.0 mL of 0.0872 M HNO3? (a) 0.0714 g (b) 0.422 g (c) 0.0180 g (d) 9.39 g (e) 0.181 g Answer: (a)

107. How many moles of Na2SO4 are in 250 mL of a 0.150 M Na2SO4 solution? (a) 0.0375 (b) 0.600 (c) 1.67 (d) 6.00 x 10–4 (e) none of these Answer: (a)

108. How many grams of NaNO3 are required to prepare 2.00 L of a 0.250 M NaNO3 solution? (a) 0.500 g (b) 84.9 g (d) 1.70 x 102 g (d) 42.5 g (e) 6.80 x 102 g Answer: (d)

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109. Calculate the molarity of a solution prepared by dissolving 10.0 grams of NaOH and then diluting the solution to 500 mL. (a) 0.251 M (b) 0.501 M (c) 1.99 M (d) 3.99 M (e) 7.98 M Answer: (b)

110. Calculate the molarity of a 20.0% by weight solution of CuCl2 in water if the density of this solution is 1.205 g/cm3.

Answer: 1.79 M

111. What volume of a solution of sulfuric acid that is 60.0% by weight H2SO4 (MW = 98.09 g/mol, d = 1.50 g/cm3) would be required to prepare 500.0 mL of a 1.50 M solution?

Answer: 81.7 mL

112. Calculate the concentration of a solution prepared by dissolving 10.0 grams of NaNO3 in enough water to give 250.0 mL of solution. (a) 0.0294 M (b) 0.118 M (c) 0.471 M (d) 2.13 M (e) 8.50 M Answer: (c)

113. What is the molarity of concentrated phosphoric acid if this solution is 85.5% H3PO4 by weight and it has a density of 1.70 g/mL? (a) less than 5 M (b) 14.8 M (c) 17.3 M (d) 20.3 M (e) none of the above Answer: (b)

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114. How many mL of 15 M NH3 would you need to make 15.0 mL of a 3.6 M NH3 solution? (a) 0.036 mL (b) 0.63 mL (c) 3.6 mL (d) 8.1 mL (e) 360 mL Answer: (c)

The following four questions (115-119) are based on the diagram above which describes the following lab procedure. Ten mL of an initial solution is pipetted into a volumetric flask. The 10.00 mL of initial solution is diluted to 100.0 mL with solvent to form the stock solution. Twenty five mL of the stock solution is then pipetted into another 100.0 mL volumetric flask and diluted to the 100.0 mL mark 2-34

to form the final solution. The glassware in the diagram has been labeled with letters.

115. Which of the following statements correctly describes the relationship between moles of solute in the labeled containers? (a) B > A (b) D > E (c) C = D (d) B = C (e) A = C Answer: (d)

116. Which of the following statements correctly describes the relationship between the concentrations of solute in the labeled containers? (a) B > A (b) D > E (c) C < D (d) B = C (e) A < C Answer: (b)

117. Which of the following statements correctly describes the relationship between moles of solute in the labeled containers? (a) A = C (b) A = B (c) D > C (d) D > E (e) C > E Answer: (e)

118. Which of the following statements correctly describes the relationship between the concentrations of solute in the labeled containers? (a) D = E (b) C = D (c) A > B (d) C > B (e) E > C Answer: (b)

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119. Calculate the concentration of a solution prepared by diluting 10.00 mL of a 0.150 M solution of KCl to 250 mL. (a) 8.05 x 10–5 M (b) 5.36 x 10–4M (c) 6.00 x 10–3M (d) 0.0805 M (e) 3.75 M Answer: (c)

120. How many mL of a 0.250 M K2SO4 solution is required to prepare 500 mL of a 0.080 M K2SO4 solution? (a) 125 mL (b) 160 mL (c) 174 mL (d) 220 mL (e) 1.56 x 103 mL Answer: (b)

Solution Stoichiometry 121. How much HgS forms when 113 mL of a 0.75 M CaS solution is mixed with 52 mL of a 1.21 M Hg(NO3)2 solution? (a) 6.30 g (b) 8.48 g (c) 14.6 g (d) 19.7 g (e) none of the above Answer: (c)

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122. If adding 22.7 mL of 0.6 M HCl to 1.00 L of a solution of AgNO3 precipitated 1.95 g of AgCl, what was the concentration of the original AgNO 3 solution?

AgNO3(aq) + HCl(aq) → AgCl(s) + HNO3(aq) (a) 1.36 x 10–5 M (b) 1.36 x 10–2 M (c) 0.027 M (d) 0.68 M (e) 13.6 M Answer: (b)

123. A 2.50 g sample of a bronze gong was cut out and dissolved in sulfuric acid.

The copper sulfate produced was mixed with KI to form CuI and the triiodide ion, I3–. The I3– ion was then titrated with thiosulfate, S2O32–.

Cu(s) + 2 H2SO4(aq) → CuSO4(aq) + SO2(g)+ 2 H2O(l) 2 CuSO4(aq) + 5 I–(aq) → 2 CuI(s) + I3–(aq) + 2 SO42–(aq) I3–(aq) + 2 S2O32–(aq) → 3 I–(aq) + S4O62–(aq) If it took 31.5 mL of 1.00 M thiosulfate for this titration, how many moles of copper were present in the original 2.50 g sample? (a) 0.0105 (b) 0.0158 (c) 0.0315 (d) 0.0630 (e) 0.0945 Answer: (c)

124. On April 19, 1995 a bomb exploded, killing 168 people, wounding hundreds of others and destroying the Alfred P. Murrah Federal Building in Oklahoma City. The explosive mixture used in this incident contained ammonium nitrate, NH4NO3, which is commonly used as a fertilizer. What is the molarity of concentrated nitric acid, HNO3(aq), if it takes 5.63 mL of the acid to consume 15.00 mL of 6.00 M NH3(aq)?

NH3(aq) + HNO3(aq) (a) less than 1 M (b) between 1 and 5 M (c) between 5 and 10 M (d) between 10 and 15 M (e) more than 15 M 2-37

NH4NO3(aq)

Answer: (e) 125. One way of determining blood alcohol levels is by titrating a sample of blood according to the following net ionic equation.

C2H5OH(aq) + 2 Cr2O72–(aq) + 16 H+(aq) → 2 CO2(g) + 4 Cr3+(aq) + 11 H2O(l) If 8.76 mL of 0.04988 M Cr2O72– is required for the titration of a 10.00 mL sample of blood, what was the molarity of the alcohol present in the blood sample? (a) less than 0.01 M (b) between 0.01 and 0.03 M (c) between 0.03 and 0.06 M (d) between 0.06 and 0.09 M (e) more than 0.1 M Answer: (b) 126. The reaction between fructose and the periodate ion is described by the following equation.

C6H12O6(aq) + 5 IO4–(aq) → 5 IO3–(aq) + 5 HCO2H(aq) + H2CO(aq) Calculate the molarity of a fructose solution if 22.35 mL of 1.050 M IO4– is required to consume 12.56 mL of this solution. (a) less than 0.25 M (b) between 0.25 M and 0.5 M (c) between 0.5 M and 1.5 M (d) between 1.5 M and 2 M (e) more than 2 M Answer: (b)

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Chapter 3: Structure of the Atom Rutherford’s Model of the Atom 1.

What would you describe as the key result of the Rutherford experiment, in which a metal target was bombarded with -particles? (a) The -particles were able to pass through the target. (b) The particles were found to knock electrons out of the target. (c) Some of the -particles were deflected through large angles. (d) Flashes of light were emitted when the -particles hit the ZnS screen after they passed through the target. (e) The experiment showed the need for better -particle detectors.

Answer: (c)

Electromagnetic Radiation 2.

O2 molecules can dissociate to form O atoms by absorbing electromagnetic radiation. If it takes 498 kJ to dissociate one mole of O2 molecules to form two moles of O atoms, what is the frequency of the light that would have just enough energy to decompose a single O2 molecule into O atoms? (a) 1.25 x 1015 s-1 (b) 8.27 x 10-19 s-1 (c) 7.52 x 1038 s-1 (d) 8.0 x 10-16 s-1 (e) 5.72 x 1016 s-1

Answer: (a)

3.

O2 molecules can dissociate to form O atoms by absorbing electromagnetic radiation. If it takes 498 kJ to dissociate one mole of O2 molecules to form two moles of O atoms, in what portion of the electromagnetic spectrum would light have sufficient energy to cause this reaction to occur? (a) radio wave (b) microwave (c) infrared (d) visible (e) ultraviolet

Answer: (e)

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4.

In a 23.490 kilogauss magnetic field, 13C nuclei absorb electromagnetic radiation in the RF (radio) portion of the spectrum at a frequency of 25.147 MHz (25.147 x 106 s–1.) Calculate the wavelength of this radiation. (a) longer than 10 meters (b) between 10 and 0.10 meters (c) between 0.10 and 10–4 meters (d) between 10–4 and 10–6 (e) shorter than 10–6 meters

Answer: (a) 5.

What is the wavelength in centimeters of light that has a frequency of 2.33 x 1015 s–1? (a) 6.99 x 1025 cm (b) 7.77 x 104 cm (c) 1.290 x 103 cm (d) 1.29 x 10–5 cm (e) none of the above

Answer: (d)

6.

A helium-neon laser emits light with a wavelength of 6328 Å. What is the energy of a 6328 Å photon in joules? (a) 2.179 x 10–19 J (b) 3.139 x 10–19 J (c) 5.448 x 10–19 J (d) 6.328 x 10–19 J (e) 2.179 x 10–18 J

Answer: (b)

7.

If green light has a frequency of 5.0 x 1014 s–1, what is the wavelength (in meters) of this light? (a) 6.7 x 10–24 (b) 2.0 x 10–15 (c) 6.0 x 10–7 (d) 3.0 x 108 (e) none of the above

Answer: (c)

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8.

If X-rays have a shorter wavelength than ultraviolet light, which of the following statements is true? (a) X-rays have smaller frequencies than UV light. (b) X-rays travel faster than UV light. (c) X-rays have more energy than UV light. (d) X-rays have larger amplitudes than UV light. (e) None of the above statements are true.

Answer: (c)

9.

There is a fundamental difference between the long wavelength IR radiation given off by the toy ovens sold at Toys ‘R Us and the shorter wavelength UV radiation emitted by the tanning booths at your local strip mall. Which of the following statements is true? (a) IR radiation has a higher frequency than UV radiation. (b) IR radiation carries more energy per photon than UV radiation. (c) IR radiation carries a larger amplitude than UV radiation. (d) IR radiation travels slower than UV radiation. (e) None of the above statements are true.

Answer: (e)

10. A cheap spectrophotometer can be made using a light-emitting diode (LED) as the source of the light. Assume that you build one of these spectrometers using an LED that gives off green light with a wavelength of 520 nm. What is the energy of a photon of this green light? (a) 1.15 x 10–34 J (b) 1.03 x 10–22 J (c) 3.82 x 10–21 J (d) 3.82 x 10–19 J (e) none of the above Answer: (d)

11. Water ionizes according to the following reaction which requires approximately 1200 kJ/mol of energy H2O(l) + h → H2O+ + e– In what portion of the electromagnetic spectrum are you likely to find the radiation that carries just enough energy to be dangerous to life because it can ionize the water that is so important to living organisms?

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(a) Radio/TV waves,   10 – 0.1 m (b) Microwaves,   0.01 – 10–4 m (c) Ultraviolet,   10–7 – 10–9 m (d) X-rays,   10–10 – 10–12 m (e) -rays,   10–12 – 10–14 m Answer: (c)

12. Which of the following wavelengths or frequencies of light will have photons with the highest energy? (a) 450 nm (b) 4.5 x 1014 Hz (c) 600 nm (d) 5.0 x 1014 Hz (e) none of these Answer: (a) 13. Yellow light has a frequency of 5.2 x 1014 cycles per second while red light has a frequency of 4.3 x 1014 cycles per second. Which of the following correctly describes the relationship between the wavelength and energy of yellow versus red light? (a) Red light has a longer wavelength and more energy than yellow light. (b) Yellow light has a longer wavelength and more energy than red light. (c) Red light has a longer wavelength and less energy than yellow light. (d) Yellow light has a longer wavelength and less energy than red light. (e) None of these.

Answer: (c)

14. Photochemistry is when light causes a chemical reaction to occur. A simple example is when light breaks a Cl2 molecule apart into 2 Cl atoms. If the energy required to break apart a mole of chlorine molecules is 243 kJ, a photon of what frequency will break apart a single Cl2 molecule? (a) 6.1 x 1014 s-1 (b) 4.0 x 10-19 s-1 (c) 3.7 x 1035 s-1 (d) 5.2 x 1015 s-1 (e) none of the above.

Answer: (a)

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15. A recent article in Scientific American notes that life would be impossible on any planet close enough to the Sun to boil water, or far enough from the Sun that water froze. It also argues that the Earth is in one of the few parts of the galaxy where life could exist. Too close to the core, and either collisions with other objects would destroy the planet or cosmic radiation from neighboring stars would destroy life. Too far from the sun and there wouldn’t be enough of the elements needed to form a planet. Let’s assume that radiation becomes particularly dangerous to life when it carries enough energy to ionize a water molecule when it is absorbed.

H2O(l) + h → H2O+ + e–

∆H = 1200 kJ/mol

Use Avogadro’s number and Planck’s constant to calculate the frequency of this radiation to one significant figure. (Hint: Pay close attention to units!) (a) 3 x 109 s–1 (b) 3 x 1012 s–1 (c) 3 x 1015 s–1 (d) 3 x 1018 s–1 (e) 3 x 1021 s–1 Answer: (c)

16. In 1814 Fraunhofer observed a series of dark lines in the sun's spectrum, which he labeled A through H. About 50 years later, Gustav Kirchhoff noticed that the wavelength of light given off when sodium salts are added to a flame is the same as the wavelength of the D line in Fraunhofer's spectrum. He concluded that certain substances give off light when heated that has the same wavelength as the light absorbed under other conditions.

The wavelength of the characteristic yellow-orange light emitted by sodium ions in a burner flame is 589.5923 nm. What is the energy of this light, in units of kJ/mol? (a) less than 1 kJ/mol (b) between 1 and 10 kJ/mol (c) between 10 and 100 kJ/mol (d) between 100 and 1000 kJ/mol (e) more than 1000 kJ/mol Answer: (d)

17. What do we mean when we say that "The energy of the electron in an atom is quantized?" (a) The electron has a very small amount of energy.

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(b) The energy of the electron is proportional to the mass of the nucleus. (c) When an electron changes its energy, it emits a quantum of light. (d) The energy of the electron can have certain fixed energies and not others. (e) The electron must be a wave.

Answer: (d)

The Bohr Model of the Atom 18. The line structure of the emission spectrum of the hydrogen atom suggests that: (a) The atom is composed of a small positive nucleus and electrons. (b) There are many electrons in a hydrogen atom. (c) The position and energy of an electron cannot be determined simultaneously. (d) The energies of electrons are quantized. (e) Each electron has multiple energies.

Answer: (d)

19. Which of the following statements about the Bohr model is wrong? (a) The energy of the electron is quantized. (b) All possible wavelengths can occur in the emission spectrum since n runs from 1 to infinity. (c) The angular momentum of the electron is quantized. (d) All quantized energies are negative. (e) The electron travels around the nucleus in an orbit.

Answer: (b)

The Energy States of the Hydrogen Atom 20. Which of the following transitions in the spectrum of the hydrogen atom results in the absorption of a photon with the largest energy? (a) n = 2 to n = 3 (b) n = 2 to n = 4 (c) n = 1 to n = 4 (d) n = 3 to n = 1 (e) n = 7 to n = 1 Answer: (c)

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21. Which transition shown below will require the absorption of the shortest wavelength photon? (a) n = 1 to n = 3 (b) n = 2 to n = 3 (c) n = 1 to n = 5 (d) n = 8 to n = 1 (e) All of the above will require the same wavelength photon.

Answer: (c)

22. Which transition in the spectrum of the hydrogen atom results in the emission of light with the longest wavelength? (a) n = 3 to n = 2 (b) n = 3 to n = 1 (c) n = 5 to n = 4 (d) n = 2 to n = 3 (e) n = 1 to n = 3 Answer: (c)

23. What is the energy of the light emitted when an electron falls from the n = 4 to the n = 2 orbit in the hydrogen atom? (a) 4.09 x 10–19 J (b) 5.45 x 10–18 J (c) 4.36 x 10–18 J (d) 2.62 x 10–17 J (e) none of the above Answer: (a)

24. If a laser operated on the n = 5 to n = 2 transition of a H atom, what wavelength in nm photon would be emitted? (a) 380 nm (b) 434 nm (c) 1950 nm (d) 3910 nm (e) none of the above Answer: (b)

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The First Ionization Energy 25. Arrange the following elements in order of increasing first ionization energy: S, Ar, Ca (a) S < Ar < Ca (b) Ar < S < Ca (c) Ca < Ar < S (d) S < Ca < Ar (e) Ca < S < Ar Answer: (e)

26. Why is the ionization energy of potassium (K) less than the ionization energy of sodium (Na)? (a) The ionized electron in K is in a higher energy shell than the ionized electron in Na. (b) Because the core charge remains the same going down a group (c) Because K has more protons in its nucleus than Na. (d) Because the core charge decreases down a group (e) Because the ionized electron in K is in a higher energy shell AND because the core charge of K and Na are the same.

Answer: (e)

27. The first ionization energy of the first six elements in the periodic table is given below.

H He Li Be B C

1312.0 kJ/mol 2372.3 520.2 899.4 800.6 1086.4

Why is the first ionization energy of B smaller than Be? (a) Because the first IE decreases as we go across a row of the periodic table from left to right. (b) Because the nucleus of a B (Z = 5) atom contains fewer protons than the nucleus of a Be (Z = 4) atom. (c) Because the outermost electron on B is coming from a 2p, not a 2s orbital.

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(d) Because the atomic number of B is odd, whereas the atomic number of Be is even. (e) For the same reason that the first IE of He is larger than that of H. Answer: (c)

28. Arrange the following elements in order of increasing first ionization energy: Ne, Na, Mg (a) Ne < Na < Mg (b) Na < Mg < Ne (c) Ne < Mg < Na (d) Mg < Ne < Na (e) Na < Ne < Mg Answer: (b)

The Shell Model 29. Construct a shell model for potassium, K, like Figure 3.9 in the text.

Answer: Potassium has a nuclear charge of +19. The shell closest to the nucleus contains 2 electrons. The second and third shells contain 8 electrons each. The fourth shell contains 1 electron.

30. Construct a shell model for calcium, Ca, like Figure 3.9 in the text.

Answer: Calcium has a nuclear charge of +20. The shell closest to the nucleus contains 2 electrons. The second and third shells contain 8 electrons each. The fourth shell contains 2 electrons.

31. What is the core charge for the potassium atom, K? (a) +1 (b) +2 (c) +8 (d) +9 (e) +19 Answer: (a)

32. Which of the following provides the best evidence for the presence of the electron shell structure of atoms?

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(a) The decrease in ionization energy going down a group. (b) The increase in ionization energy across a period. (c) That the value of ms can be +1/2 or –1/2 (d) That the ionization energy of O is less than the ionization energy of N. (e) All of the above provide evidence for the shell structure of atoms.

Answer: (a)

33. What is the core charge for the calcium atom, Ca? (a) +1 (b) +2 (c) +8 (d) +10 (e) +20 Answer: (b)

34. What is the core charge for the silicon atom, Si? (a) +1 (b) +2 (c) +4 (d) +6 (e) +8 Answer: (c)

35. How would you expect the first ionization energy of K to compare to that of Ar? (a) K > Ar (b) K = Ar (c) K < Ar (d) First ionization energies for an atom can vary (e) Ar does not ionize Answer: (c)

36. Arrange the following atoms in order of decreasing energy needed to remove the most tightly held electron (the electron closest to the nucleus).

Na, Mg, Al, P

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(a) Na > Mg > Al > P (b) Na > Mg > P > Al (c) Al > P > Na > Mg (d) Al > P > Mg > Na (e) P > Al > Mg > Na Answer: (e) 37. Why is the first ionization energy of Br less than that of Cl? (a) Bromine has more protons than chlorine. (b) Bromine has more electrons than chlorine. (c) The core charge on bromine is less than that on chlorine. (d) The outer shell electrons in bromine are farther from the nucleus than in chlorine. (e) none of these Answer: (d)

38. Predict the order of decreasing first ionization energies for Ar, K, and Ca. (a) Ar > K > Ca (b) Ca > K > Ar (c) K > Ar > Ca (d) K > Ca > Ar (e) Ar > Ca > K Answer: (e)

39. Which atom is represented by a shell model with two electrons in the first shell, and five electrons in the second shell? (a) Be (b) C (c) N (d) K (e) P Answer: (c)

40. Which of the following atoms has a core charge of +6? (a) B (b) C (c) N (d) Si

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(e) S Answer (e)

41. Which of the following atoms has the lowest first ionization energy? (a) Cl (b) Ar (c) K (d) Ca (e) Sc Answer: (c)

42. Which atom requires the largest amount of energy to remove the most tightly held electron (the electron closest to the nucleus)? (a) K (b) Ca (c) Ga (d) As (e) S Answer: (d)

43. Arrange the following atoms in order of increasing first ionization energies. (a) Be < Mg < Ca < Sr (b) Mg < Ca < Sr < Be (c) Sr < Ca < Mg < Be (d) Ca < Sr < Mg < Be (e) none of these are in the correct order Answer: (c)

44. Arrange the following atoms in order of increasing first ionization energy. (a) P < S < Cl (b) Cl < S < P (c) S < P < Cl (d) S < Cl < P (e) Cl < P < S Answer: (a)

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45. In which of the following groups of atoms does each atom have the same number of outer shell electrons? (a) B, C, N (b) K, Ca, P (c) C, Si, Ge (d) O, F, Cl (e) N, P, Cl Answer: (c)

Photoelectron Spectroscopy and the Structure of Atoms 46. A photoelectron spectrum has three peaks. The first two are one third the size of the last one. Which atom corresponds to this spectrum? (a) He (b) Be (c) C (d) O (e) Ne Answer: (e)

47. How do Photoelectron spectroscopy experiments differ from experiments used to obtain first ionization energies?

Answer: Experiments to obtain first ionization energies measure only the energy required to remove the most loosely held electron from the atom.

Photoelectron spectroscopy experiments measure the energy required to remove each of the electrons from an atom.

48. Which element is represented by the following PES spectrum?

(a) N (b) F 3-13

(c) Na (d) Al (e) O Answer: (c)

49. Which element is represented by the following PES spectrum?

(a) F (b) Si (c) S (d) K (e) Ti Answer: (c) 50. Which element is represented by the following PES spectrum?

(a) Li (b) Mg (c) Ca (d) Be (e) Ne Answer: (e) 51. Which of the following atoms would have four peaks in its PES spectrum? (a) Li (b) Mg (c) F 3-14

(d) K (e) O Answer: (b)

Allowed Combinations of Quantum Numbers 52. Which is a legitimate set of n, l, ml and ms quantum numbers? (a) 0,0,0, 12 (b) 8,4,-3,- 12 (c) 3,3,2,+ 12 (d) 2,1,-2,- 12 (e) 5,3,3,-1 Answer: (b)

53. Which set of n, l, ml and ms quantum numbers is allowed? (a) 4,-2,-1, 12 (b) 4,2,3, 12 (c) 4,3,0,1 (d) 4,0,0,- 12 (e) 4,2,2,2 Answer: (d) 54. Which of the following sets of n, l, ml and ms quantum numbers isn’t allowed? (a) 1,1,0,+ 12 (b) 2,0,0,+ 12 (c) 3,2,1,- 12 (d) 4,1,-1,+ 12 (e) 5,3,2,- 12 Answer: (a)

55. Which of the following selection rules for quantum numbers is incorrectly stated? (a) n is any integer greater than or equal to zero.

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(b) l is any integer between zero and n - 1. (c) ml is any integer between -l and +l. (d) ms is either + 12 or - 12 . (e) All of the above selection rules are correctly stated.

Answer: (a)

56. Which of the following quantum numbers is used to describe the orientation of an orbital in space? (a) n (b) l (c) ml (d) ms (e) None of these quantum numbers describe the orientation of an orbital in space.

Answer: (c)

57. Which set of quantum numbers can be used to describe a 2p electron? (a) 2,1,0,- 12 (b) 2,0,0, 12 (c) 2,2,1, 12 (d) 3,2,1,- 12 (e) 3,1,0, 12 Answer: (a)

58. How many electrons can be placed in a 3d orbital? (a) 2 (b) 6 (c) 8 (d) 10 (e) 18 Answer: (a) 59. Which of the following atomic orbitals doesn’t exist? (a) 3f (b) 3p (c) 5f 3-16

(d) 5d (e) 6s Answer: (a) 60. What is the orbital designation for the quantum numbers: n = 4, l = 2, ml = -1? (a) 4s (b) 4p (c) 4d (d) 4f (e) none of these Answer: (c)

61. Orbitals for which l = 1 are described by which of the following symbols? (a) s (b) p (c) d (d) f (e) g Answer: (b)

Orbitals and the Pauli Exclusion Principle 62. What is the maximum number of unpaired electrons that can be accommodated in a 5d subshell? (a) 3 (b) 5 (c) 6 (d) 7 (e) 10 Answer: (b)

63. What is the maximum number of electrons that can be accommodated in the subshell for which n = 3 and l = 2? (a) 2 (b) 6 (c) 10 (d) 14

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(e) 18 Answer: (c)

64. Calculate the maximum number of electrons that can fit into the n = 4 shell of orbitals. (a) 6 (b) 8 (c) 18 (d) 16 (e) 32 Answer: (e) 65. The Pauli exclusion principle states that: (a) No two electrons in an atom can have the same spin. (b) No two electrons in an atom can occupy the same orbital. (c) Two electrons in the same orbital have identical values of the spin quantum number. (d) No two electrons in an atom can have the same set of four quantum numbers. (e) Two electrons in an atom can have the same principal quantum number.

Answer: (d) 66. Which of the following quantum numbers can have a value that is not an integer? (a) n (b) l (c) ml (d) ms (e) None of the quantum numbers can have any value that is not an integer.

Answer: (d)

Predicting Electronic Configurations 67. Which of the following subshells is filled first when electrons are added to the atomic orbitals on a xenon atom? (a) 4d

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(b) 4f (c) 5s (d) 5p (e) 5d Answer: (c)

68. Which of the following subshells would be filled first when electrons are added to a gold atom? (a) 4f (b) 5d (c) 6s (d) 6p (e) 6d Answer: (c)

69. Determine the group of the periodic table in which an element with the following electron configuration belongs: 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p4 (a) VA (b) IIIA (c) IVA (d) VIA (e) none of these Answer: (d)

70. Determine the period and group in the periodic chart that contains the element with the following electron configuration: 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p4 (a) period 4, group VIA (b) period 6, group IVA (c) period 3, group IV B (d) period 4, group IVA (e) none of these Answer: (a)

71. In what group of the periodic table would an element with the following electron configuration belong?

1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p1

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(a) Group IA (b) Group IlIA (c) Group VA (d) Group VIIA (e) none of the above Answer: (b)

72. A single atom of element 109 has been synthesized. Predict the electronic configuration of this element. Which of the following elements would 109 most resemble? (a) Ta (b) Re (c) Ir (d) Au (e) Tl Answer: (c)

73. A single atom of element 114 was recently synthesized. Predict the electronic configuration of this element. Which of the following elements would 114 most resemble? (a) Au (b) Hg (c) Tl (d) Pb (e) Po Answer: (d)

74. Theoreticians predict that the element with atomic number 120 will be more stable than the elements recently discovered with atomic numbers between 103 and 109. Based on the order of filling of atomic orbitals, the chemistry of this element should most closely resemble the chemistry of which of the following? (a) Ra, Group IIA (b) Pb, Group IVA (c) Po, Group VIA (d) Rn, Group VIIIA (e) one of the transition metals between La and Hg Answer: (a)

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75. In period 6 of the periodic table, the series of elements which differ in the number of electrons contained in f orbitals (a) begins with Cs and ends with Rn. (b) begins with La and ends with Hg. (c) begins with Tl and ends with Rn. (d) begins with La and ends with Rn. (e) begins with Ce and ends with Lu. Answer: (e) 76. A possible set of quantum numbers for the last electron added to form a gallium atom (Z = 31) in its lowest energy state is: (a) 3, 1, 0, - 12 (b) 3, 2, 1, 12 (c) 4, 0, 0, 12 (d) 4, 1, 1, 12 (e) 4, 2, 2, 12 Answer: (d) 77. A possible set of quantum numbers for the highest energy electron in an As3+ ion is: (a) 3, 1, -1, 12 (b) 4, 0, 0, - 12 (c) 3, 2, 0, 12 (d) 4, 1, -1, 12 (e) 5, 0, 0, 12 Answer: (c) 78. The outermost or highest energy electron in element 105 could be characterized by which of the following sets of n, l, ml and ms quantum numbers? (a) 7, 3, -3, - 12 (b) 6, 3, -1, - 12 (c) 6, 2, 0, 12 (d) 5, 3, -1, - 12 (e) 5, 2, 0, 12 Answer: (c) 3-21

79. Which of the following describes a possible set of quantum numbers for the last electron added to form an aluminum atom when atomic orbitals are filled? (a) 1, 0, 0, + 12 (b) 2, 0, 0, + 12 (c) 2, l, 1, + 12 (d) 3, 0, 0, + 12 (e) 3, 1, 1, + 12 Answer: (e)

80. What is the lowest energy electronic configuration of a fluorine atom? (a) 1s2 2s2 2p5 (b) 1s2 2s2 2p6 (c) 1s2 2s2 2p7 (d) 1s2 2s2 2p6 3s1 (e) none of the above Answer: (a)

81. What is the value of x in the following electron configuration for silicon?

1s2 2s2 2p6 3s2 3px (a) 1 (b) 2 (c) 3 (d) 4 (e) 6 Answer: (b)

82. What is the correct electron configuration for the Ti atom? (a) 1s2 2s2 2p6 3s2 3p6 3d4 (b) [Ar] 4s2 4d2 (c) [Ar] 4d4 (d) 1s2 2s2 2p6 3s2 3d2 (e) [Ar] 4s2 3d2 Answer: (e)

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83. Which of the following would have the electron configuration: 1s2 2s2 2p6 3s2 3p6 3d4 (a) Ca2+ (b) Cr2+ (c) Fe2+ (d) Ti2+ (e) none of these Answer: (b) 84. What is the correct electron configuration for the P3+ ion? (a) [Ne] (b) [Ne] 3s2 (c) [Ne] 3s2 3p3 (d) [Ne] 3s2 3p5 (e) [Ne] 3s2 3p6 Answer: (b) 85. What is the electron configuration for the bromide ion, Br–? (a) [Ar] 4s2 4p5 (b) [Ar] 4s2 3d10 4p7 (c) [Ar] 4s2 3d10 4p5 (d) [Ar] 4s2 3d10 4p6 (e) [Ar] 4s2 3d10 3p6 Answer: (d) 86. Which of the following describes the electron configuration for the Sn 2+ ion? (a) [Kr] 4d10 (b) [Kr] 5s2 (c) [Kr] 5s2 5p2 (d) [Kr] 5s2 4d10 (e) [Kr] 5s2 4d10 5p2 Answer: (d)

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Electronic Configurations and Hund’s Rule 87. Give the full electron configuration of arsenic, As. Answer: 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p3

88. Give the abbreviated electron configuration of platinum, Pt. Answer: [Xe] 6s2 4f14 5d8

89. Give the orbital diagram (using parentheses and arrows to represent orbitals and electrons) of phosphorus.

Answer: 1s2 

2s2 

2p6   

3s2 

3p3   

90. Write the electron configuration for the following atoms or ions.

(I) Br − (write the full electron configuration) (II) Si (write the full electron configuration with orbital diagram; use arrows for the electrons) (III) Ba (the abbreviated form) Answer: (I) 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 (II) 1s2 2s2 2p6      (III) [Xe] 6s2

3s2 

3p2  

91. Which electron configuration for carbon would satisfy Hund's rules? (a) 1s2 2s2 2px2 2py0 2pz0 (b) 1s2 2s2 2px1 2py1 2pz0 (c) 1s2 2s2 2px1 2py1 2pz1 (d) 1s2 2s2 2px2 2py1 2pz1 (e) None of these configurations satisfy Hund's rules.

Answer: (b)

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92.

Which atom contains the largest number of unpaired electrons? (a) B (b) N (c) F (d) Ti (e) Cu

Answer: (b)

93. Which atom or ion has the largest number of unpaired electrons? (a) Li+ (b) B (c) C4– (d) N (e) O2– Answer: (d)

The Sizes of Atoms and Their Ions 94. Which of the following atoms or ions is the smallest? (a) Na (b) Na+ (c) Mg (d) Mg2+ (e) Ba2+ Answer: (d)

95. Arrange the following ions in order of decreasing ionic radius. (a) F– > Cl– > Br– > Na+ > Li+ (b) Na+ > Li+ > F– > Cl– > Br– (c) Br– > Cl– > F– > Li+ > Na+ (d) Br– > Cl– > F– > Na+ > Li+ (e) Br–> Cl– >Na+ > F– > Li+ Answer: (d)

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96. For each pair of atoms, which has the smallest metallic radius? (a) Na, Be (b) K, Mg (c) Cs, Ra Answer: Be, Mg, Cs 97. Which ion, Na+ or K+, is the smaller? Explain.

Answer: Na+. The potassium ion will have one more shell of electrons than Na+. The potassium ion will have electrons in the third shell.

98. Which ion, O2– or S2–, is larger? Explain.

Answer: S2–. The sulfur ion will have one more shell of electrons than oxygen. The sulfur ion will have electrons in the third shell.

99.

Which of the following atoms or ions is the smallest? (a) Na (b) Mg (c) Na+ (d) Mg2+ (e) O2–

Answer: (d) 100. Arrange the following ions or atoms in order of increasing size: Cl−, Ar, Ca2+. (a) Cl−< Ar < Ca2+ (b) Ca2+ < Cl−< Ar (c) Ar < Ca2+ < Cl− (d) Ca2+ < Ar < Cl− (e) Cl− < Ca2+ < Ar Answer: (d)

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Second, Third, Fourth, and Higher Ionization Energies 101. First ionization energy has been defined as the energy required to remove the most loosely held electron from an electrically neutral atom. Second ionization energy is the minimum energy required to remove an electron from a +1 cation (from an atom that has already lost its first electron). In the same way third ionization energy is the minimum energy to remove an electron from a +2 cation (from an atom that has lost two electrons).

Below is a table of first, second, third and fourth ionization energies for Na, Mg and Al. Explain this data for the magnesium atom in terms of its electron configuration.

Na Mg Al

1st IE 4958 737.7 577.6

2nd IE 4562.4 1450.6 1816.6

3rd IE 6912 7732.6 2744.7

4th IE 9543 10,540 11,577

Answer: The first and second electrons are relatively easy to remove but the third and fourth require much more energy. This indicates that the third and fourth electrons are removed from a shell closer to the nucleus.

102. If a powerful enough source of energy is used in the PES experiment, it is possible to measure the energy required to remove the 2 nd, 3rd, 4th, and so on, electrons from an atom. Which of the following elements would have the largest fourth ionization energy? (a) Al (b) Si (c) P (d) S (e) Cl Answer: (a)

Average Valence Electron Energies (AVEE) 103. Predict which atom is most likely to be a metal. (a) P (b) Br (c) As (d) Co (e) Ga Answer: (d)

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104. The results of photoelectron spectroscopy studies of the last four elements in the 2nd row of the periodic table are given below.

Element N O F Ne

Ionization energies (in units of MJ/mol) 1s 2s 2p 39.6 2.45 1.40 52.6 3.12 1.31 67.2 3.88 1.68 84.0 4.68 2.08

Which of the following statements is correct? (a) The average valence electron energy for N is larger than O, F, or Ne. (b) The average valence electron energy for O is larger than N, F, or Ne. (c) The average valence electron energy for F is larger than N, O, or Ne. (d) The average valence electron energy for Ne is larger than N, O, or F. (e) The average valence electron energy is the same for all four elements.

Answer: (d)

105. If you believe that the average valence electron energy measures the ability of an atom to hold on to its valence electrons, and therefore the ability of the atom to compete for electrons in any bond it might form, which of these atoms should be the most electronegative? (a) N (b) O (c) F (d) Ne (e) These elements all have the same electronegativity.

Answer: (d)

106. Predict which atom is most likely to form a negative ion. (a) Mg (b) Si (c) S (d) Te (e) Ca Answer: (c)

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107. Given the information below, what is the average valence electron energy for the oxygen atom?

Element O

Ionization energies (in units of MJ/mol) 1s 2s 2p 52.6 3.12 1.31

(a) 19.4 MJ/mol (b) 14.6 MJ/mol (c) 116.7 MJ/mol (d) 1.44 MJ/mol (e) 1.91 MJ/mol Answer: (e)

108. Predict which atom is most likely to form a positive ion. (a) Ge (b) Se (c) Br (d) F (e) S Answer: (a)

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Chapter 4: The Covalent Bond Valence Electrons 1.

Which of the following elements has five valence electrons? (a) Sr (b) K (c) Si (d) Se (e) As

Answer: (e)

2.

Which of the following elements has 2 valence electrons? (a) Ca (b) Ga (c) Ge (d) As (e) Se

Answer: (a)

Using Lewis Structures 3.

How many valence electrons are in the IF2– ion? (a) 7 (b) 14 (c) 20 (d) 21 (e) 22

Answer: (e)

4.

Which of the following species has all its electrons paired? (a) Fe3+ (b) Zn2+ (c) ClO2 (d) NO (e) NO2

Answer: (b)

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5.

How many pairs of bonding electrons are in the valence shell of the iodine atom in the IF2– ion? (a) 0 (b) 1 (c) 2 (d) 3 (e) 4

Answer: (c)

6.

How many pairs of nonbonding electrons are in the valence shell of the iodine atom in the IF2– ion? (a) 0 (b) 1 (c) 2 (d) 3 (e) 4

Answer: (d)

7.

How many pairs of nonbonding electrons are in the valence shell of the iodine atom in the ICl4– ion? (a) 0 (b) 1 (c) 2 (d) 3 (e) 4

Answer: (c)

8.

How many nonbonding electrons are on the central atom in the Lewis structure of xenon tetrafluoride? (a) 0 (b) 1 (c) 2 (d) 3 (e) 4

Answer: (e)

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9.

How many pairs of bonding electrons can be found in the valence shell of the nitrogen atom in HNO2? (Hint: the skeleton structure is HONO) (a) 0 (b) 1 (c) 2 (d) 3 (e) 4

Answer: (d)

10. How many pairs of nonbonding electrons can be found on the chlorine atom in chlorous acid, HOClO? (a) 0 (b) 1 (c) 2 (d) 3 (e) 4 Answer: (c) 11. Which of the following elements would form an XF62- ion that has no nonbonding electrons in the valence shell of the central atom? (a) Ca (b) C (c) Si (d) S (e) P Answer: (c)

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12. Nitrous oxide (N2O) is an anesthetic commonly known as “laughing gas.”

Which of the following is a correct Lewis structure for this compound?

(a) N

N

O

(b) N

N

O

(c) N

N

O

(d) N

N

O

(e) N

N

O

Answer: (c)

13. Which element would form XF5 for which there are no nonbonding electrons in the Lewis structure? (a) N (b) Al (c) Si (d) P (e) Br Answer: (d)

14. If X is a 3rd row element and XCl3 has the Lewis structure Cl

X Cl

X must be: (a) Al (b) Si (c) P (d) S (e) Cl Answer: (c)

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Cl

15. If the XF6– ion has the Lewis structure F F

F Xe

F

F F

X could be which of the following elements? (a) N (b) Al (c) Si (d) P (e) Br Answer: (d)

16. Assume that the XF3 molecule has the following Lewis structure.

F F

X F

The molecule could be: (a) boron trifluoride, i.e., X = B (b) nitrogen trifluoride, i.e., X = N (c) chlorine trifluoride, i.e., X = Cl (d) phosphorus trifluoride, i.e., X = P (e) more than one of the above Answer: (a)

17. Which of the following elements would form a compound with the following Lewis structure?

O

X

(a) Al (b) Si (c) P (d) S (e) Cl Answer: (c)

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Cl

18. If there are two pairs of nonbonding electrons in the valence shell of the central atom in the XF4– ion, in which Group of the periodic table does the element X belong? (a) Group IVA (b) Group VA (c) Group VIA (d) Group VIIA (e) Group VIIIA Answer: (d)

19. What is element X in the following Lewis structure?

H

X

X

H

H

H

(a) B (b) C (c) N (d) O (e) F Answer: (c)

20. Which of the following species contains only single bonds? (a) CN– (b) NO+ (c) CO (d) O22– (e) Cl2CO Answer: (d)

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21. Which of the following would be the best Lewis structure for NOCl?

(a) O

N

Cl

(b) O

N

Cl

(c) O

N

Cl

(d) O

N

Cl

(e) O

N

Cl

Answer: (c)

22. Which of the following compounds would have the same Lewis structure as the N2 molecule?

(I) F2 (II) CO (III) NO (IV) CN– (V) NO+ (a) I (b) II and IV (c) II and III (d) II, IV and V (e) IV and V Answer: (d) 23. Which of the following would have a Lewis structure most like that of CO 32ion? (a) CH3+ (b) NH3 (c) NO3– (d) SO32– (e) SO42– Answer: (c) 24. The correct Lewis structure for CH2Cl2 contains: (a) 2 single bonds, 2 double bonds and no nonbonding electrons (b) 4 single bonds and no nonbonding electrons (c) 4 single bonds and four nonbonding electrons (d) 2 single bonds, 2 double bonds, and eight nonbonding electrons (e) 4 single bonds and 12 nonbonding electrons Answer: (e) 4-7

25. The correct Lewis structure for the NO2+ ion has: (a) 2 double bonds and eight nonbonding electrons (b) 1 double bond, 1 single bond, and 12 nonbonding electrons (c) 2 single bonds and 12 nonbonding electrons (d) 2 triple bonds and 4 nonbonding electrons (e) 1 single bond, 1 double bond, and 8 nonbonding electrons Answer: (a)

26. In which Group of the periodic table does element X belong if there are one pair of nonbonding electrons in the valence shell of the central atom in the XF4– ion? (a) Group IIIA (b) Group IVA (c) Group VA (d) Group VIA (e) none of these Answer: (d)

27. Which element forms a compound with the following Lewis structure?

Cl

X

Cl

(a) Al (b) Si (c) P (d) S (e) I Answer (e)

28. Which of the following are exceptions to the Lewis octet rule?

(I) CO2 (II) BeF2 (III) SF4 (IV) SO3 (V) PCl3 (a) I and II (b) II and III (c) III and IV (d) I, II and III (e) none of these are exceptions to the Lewis octet rule Answer: (b)

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29. Which of the following are exceptions to the Lewis octet rule?

(I) BF3 (II) H2CO (III) XeF4

(IV) H3O+ (V) ClF3

(a) I and II (b) I and III (c) I, III and V (d) II and III (e) III and V Answer: (c) 30. Which of the following elements is the most likely to follow the octet rule? (a) H (b) C (c) Si (d) As (e) All of these elements must follow the octet rule.

Answer: (b)

Bond Lengths 31. Which of the bonds in ethanol, CH3CH2OH, would have a bond length of approximately 0.103 nm?

Covalent Radii: H = 0.037 nm, C = 0.077 nm, O = 0.066 nm (a) C−C (b) C−O (c) C−H (d) O−H (e) not enough information is given Answer: (d)

32. Which of the following compounds/ions has the shortest C−O bond length? (a) CH3OH (b) H2CO (c) CO32– (d) CH3CH2OH (e) all C−O bonds are the same length Answer: (b)

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33. Which of the following bond lengths is the largest? (a) H−C (b) H−S (c) H−O (d) H−F (e) H−H Answer: (b) 34. Which bond length is the shortest in the following compound?

Cl H

C

O

H

F

(a) H−C (b) Cl−C (c) F−C (d) C−O (e) O−H Answer: (e)

35. In which of the following is the carbon to oxygen bond the shortest?

(a) H3C

O

(b) H2C

O

(c) H3C

O

(d) C

H

CH3

O

(e) all C−O bonds are the same length Answer: (d)

36. Which of the following compounds contains the strongest carbon-nitrogen bond? (Hint: Look at their Lewis structures) (a) H2CNH (b) HCN (c) H3CNH2 (d) H3CNO (e) all of these molecules contain C–N single bonds Answer: (b)

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Resonance Hybrids 37. Which of the following molecules or ions is not a resonance hybrid? (a) SCN– (b) CO32– (c) C2H4 (d) N2O (e) all are resonance hybrids Answer: (c)

38. What is the difference between an isomer and a resonance structure?

Answer: Isomers are different arrangements of the location of atoms in a molecule. Resonance structures are different arrangements of multiple bonds in a molecule.

39. Which of the following would be a resonance structure for the compound shown below.

X

(a) X

Z

Z

(b) X

Z

Z

(c) Z

X

X

(d) X

Z

X

Z

X

(e) none of these Answer: (d)

40. Which of the following molecules is best described with two resonance structures? (a) NO2– (b) H2O (c) OF2 (d) ClNO (e) BeF2 Answer: (a)

4-11

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41. Which of the following isn’t a resonance hybrid? (a) CO2 (b) NO2– (c) N2O (d) SO3 (e) SCN– Answer: (a)

42. Which of the following are resonance hybrids?

(I) HCO2– (II) PH3 (III) HCN (IV) SO42–

(V) NO2

(a) I (b) II and IV (c) IV and V (d) I and IV (e) I and V Answer: (e)

Electronegativity 43. Which of the following can be used to predict an element’s electronegativity? (a) average valence electron energy (b) Lewis Structures (c) bond angles (d) bond length (e) resonance Answer: (a)

44. Which of the following elements has the greatest electronegativity? (a) Mg (b) Cu (c) S (d) As (e) Sr Answer: (c)

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45. Which of the following lists the elements in the correct order of increasing electronegativity? (a) Si < S < O (b) S < Si < O (c) S < O < Si (d) Si < O < S (e) O < Si < S Answer: (a)

Partial Charge Formal Charge 46. Which of the following is the best true description of the charge on an atom within a molecule? (a) oxidation number (b) partial charge (c) formal charge (d) all of these are equally good (e) none of these Answer: (b)

47. In which of the following compounds is the bond the least polar? (a) NaCl (b) CH4 (c) CO (d) Cl2 (e) HCl Answer: (d)

48. Use formal charge to decide which of the following skeleton structures will give the best Lewis structure of the molecule. Show the values of formal charge and explain your choice of the best structure.

O H

C

C H

H

O

H

Answer: All atoms in the structure on the left have formal charges of zero.

The central oxygen in the structure on the right has a formal charge of +2 and the carbon has a formal charge of –2. The structure on the left is best.

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49. The “zwitterion” form of the amino acid glycine is: H3N+CH2CO2–. The properties of the various amino acids depend heavily on the presence (or absence) of charge on the side chain. Consider the amino acid arginine, for example, which has the following Lewis structure: H H

H N H

C

N

1 2

N

CH2

CH

3

N H

CH2

H

H

C

H

O

O

Determine the Formal Charge on the nitrogens labeled 1, 2 and 3. (a) N(1) = 0, N(2) = +1, and N(3) = 0 (b) N(1) = 0, N(2) = –1, and N(3) = 0 (c) N(1) = +1, N(2) = 0, and N(3) = +1 (d) N(1) = –1, N(2) = 0, and N(3) = –1 (e) N(1) = +1, N(2) = 0, and N(3) = –1 Answer: (a) 50. Our response to the anthrax scare was a 20-year-old antibacterial known as ciprofloxacin or “Cipro.”

H

1

CH2

N

CH2

CH2

N CH2 2

H CH C

C

C F

N

CH2

C CH

3

CH C O

O C OH

Apply the concept of formal charge to the three nitrogen atoms labeled 1, 2 and 3. Are they neutral, negatively charged, or positively charged? (a) N(1) = -1, N(2) = -1, and N(3) = -1 (b) N(1) = -1, N(2) = 0, and N(3) = 0 (c) N(1) = 0, N(2) = 0, and N(3) = 0 (d) N(1) = 0, N(2) = +1, and N(3) = +1 (e) N(1) = +1, N(2) = +1, and N(3) = +1 Answer: (c)

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The Shapes of Molecules 51. Which element will combine with sulfur to form a linear molecule with the formula XS2? (a) Li (b) B (c) C (d) O (e) F Answer: (c)

52. In which compound are the valence electrons on the sulfur arranged towards the corners of a tetrahedron? (a) SF3+ (b) SF4 (c) SF5 (d) SF6 (e) both SF4 and SF6 Answer: (a)

53. What is the distribution of electrons around the S atom in H2SO4?

[Hint: The skeleton structure of this oxyacid is O2S(OH)2. (a) linear (b) trigonal planar (c) tetrahedral (d) trigonal bipyramidal (e) octahedral Answer: (c)

54. What is the distribution of the oxygen atoms around the N atom in HNO3?

[Hint: The skeleton structure of this oxyacid is HONO2] (a) linear (b) trigonal planar (c) tetrahedral (d) trigonal bipyramidal (e) octahedral Answer: (b)

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55. The bond angles in most symmetric, ABx, molecules are the same. Which of the following is an exception to this rule? (a) NH3 (b) SF6 (c) PCl5 (d) SO3 (e) CO2 Answer: (c) 56. Which term best describes the shape of the I3– ion? (a) linear (b) bent (c) trigonal planar (d) trigonal pyramidal (e) trigonal bipyramidal Answer: (a)

57. What is the molecular geometry of the compound formed by the reaction of nitrogen with fluorine? (a) bent (b) trigonal planar (c) trigonal pyramidal (d) trigonal bipyramidal (e) none of the above Answer: (c) 58. What is the molecular geometry of the SF5+ ion? (a) see-saw (b) square planar (c) square pyramidal (d) trigonal bipyramidal (e) octahedral Answer: (d)

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59. What is the best description of the structure of the SF3+ ion? (a) bent or angular (b) trigonal planar (c) trigonal pyramidal (d) tetrahedral (e) trigonal bipyramidal Answer: (c) 60. The molecular structure of the IO3– ion is best described as: (a) bent or angular (b) trigonal planar (c) trigonal pyramidal (d) trigonal bipyramidal (e) T-shaped Answer: (c)

61. Which of the following molecules has a trigonal planar shape? (a) BeCl2 (b) PCl3 (c) ICl2+ (d) BF3 (e) H3O+ Answer: (d)

62. Which of the following molecules are planar?

(I) SO3 (II) SO32– (III) NO3– (IV) PF3 (V) BH3 (a) I and II (b) I and III (c) I, III and V (d) II and IV (e) III and V Answer: (c)

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63. Which of the following molecules are tetrahedral?

(I) SiF4 (II) CH4 (III) NF4+ (IV) BF4– (V) TeF4 (a) I and II (b) I, II and III (c) I, II, III and IV (d) II, III and IV (e) V Answer: (c) 64. Which of the following molecules or ions doesn’t have a tetrahedral shape? (a) SiF4 (b) NF4+ (c) BF4– (d) SeF4 (e) CF4 Answer: (d) 65. Which of the following molecules or ions isn’t planar? (a) NH3 (b) BF3 (c) CO32– (d) NO3– (e) none of these molecules or ions are planar Answer: (a) 66. What is the best description of the shape of the TeH2 molecules? (a) linear (b) bent (c) trigonal planar (d) tetrahedral (e) seesaw Answer: (b)

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67. Metformin is the chemical name for a prescription drug sold as “glucophage,” which is an oral hypoglycemic. In other words, a pill that lowers blood sugar levels so that adult on-set diabetics don’t have to inject themselves with insulin. The Lewis structure of metformin is shown below.

H

2 N

H3C N

C

H

4 N

3 N

H

C

N

1

5

H3C

H

H

Use the electron domain model to predict the geometry around the nitrogen atom in the center of the molecule, e.g., N(3). (a) bent or angular (b) trigonal planar (c) trigonal pyramidal (d) trigonal bipyramidal (e) tetrahedral Answer: (c)

68. Terpin hydrate is an expectorant that was used for many years by people with serious coughs, such as those suffering from pneumonia.

H

H CH2

O

O

CH2

C

CH3 CH

C

H3C CH2

CH2

CH3

Use the electron domain model to predict the distribution of electron domains in the valence shell of the oxygen atoms in this compound. (a) linear (b) bent or angular (c) trigonal planar (d) tetrahedral (e) trigonal bipyramidal Answer: (d)

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69. When the rare gas elements (He, Ne, Ar, Kr, Xe and Rn) were first discovered, a sample of Ar was sent to Moissan, who had discovered the most active of the known elements: F2. The failure of Moissan's attempts to react argon with fluorine, coupled with repeated failures by other chemists to get the more abundant of these gases to undergo chemical reaction, eventually led to their being labeled inert gases. We now know this isn’t so.

Xe reacts with F2 to form a variety of compounds, including XeF4. Which of the following best describes the XeF4 molecule? (a) Trigonal bipyramidal distribution of valence electrons on the Xe with a T-shaped molecular geometry. (b) Trigonal bipyramidal distribution of valence electrons on the Xe with a trigonal planar molecular geometry. (c) Trigonal bipyramidal distribution of valence electrons on the Xe with a pyramidal molecular geometry. (d) Octahedral distribution of valence electrons on the Xe with a square planar molecular geometry. (e) Octahedral distribution of valence electrons on the Xe with a seesaw molecular geometry.

Answer: (d)

70. An element Z forms a molecular compound ZBr3 that is found to have trigonal pyramidal molecular geometry. The element Z must be in which group in the periodic table? (a) IIIA (b) IVA (c) VA (d) VIA (e) VIIA Answer: (c)

71. Tellurium hydride is one of the foulest smelling compounds known. How many pairs of nonbonding electrons are in the valence shell of the tellurium atom in TeH2? (a) 0 (b) 1 (c) 2 (d) 3 (e) 4 Answer: (c)

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72. There is a single pair of nonbonding electrons on the central atom in the Lewis structure of ICl5. This molecule has: (a) a trigonal pyramidal distribution of electrons and an octahedral molecular geometry. (b) an octahedral distribution of electrons and a square pyramidal geometry. (c) a trigonal bipyramidal distribution of electrons and a seesaw molecular geometry. (d) an octahedral distribution of electrons and a trigonal bipyramidal geometry. (e) a trigonal bipyramidal distribution of electrons and trigonal bipyramidal geometry.

Answer: (b)

73. Which of the following molecules are linear?

(I) C2H2 (II) CO2 (III) NO2– (IV) NO2+ (V) H2O (a) I and II (b) I, II and III (c) I, II and IV (d) II and III (e) V Answer: (c) 74. Which of the following has a square-planar geometry? (a) SF4 (b) ClO4– (c) CCl4 (d) ICl4– (e) PO43– Answer: (d) 75. Which of the following species isn’t planar? (a) SO3 (b) SO32– (c) SO2 (d) BF3 (e) NO3– Answer: (b) 4-22

76. Which of the following molecules are best described as bent or angular in shape?

(I) H2S (II) CO2 (III) ClNO (IV) NH2– (V) O3 (a) I and II (b) I, III and IV (c) I, III, IV and V (d) II, III and IV (e) all of these molecules are bent or angular in shape Answer: (c)

77. What is the molecular geometry of xenon tetrafluoride? (a) tetrahedral (b) trigonal bipyramidal (c) seesaw (d) square planar (e) octahedral Answer: (d) 78. The shape of the NO2– ion is best described as: (a) linear (b) T-shaped (c) bent (d) square planar (e) pyramidal Answer: (c)

79. How many pairs of nonbonding electrons are in the valence shell of the central iodine atom in the I3– ion? (a) 0 (b) 1 (c) 2 (d) 3 (e) 4 Answer: (d)

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80. The Raschig process involves the reaction between ammonia and the hypochlorite ion to form hydrazine, which has been used as a rocket fuel.

2 NH3(aq) + OCl–(aq) → H2NNH2(g) + Cl-(aq) +H2O(l) There are isolated cases of individuals making hydrazine unexpectedly in their homes when cleaning a commode. They start with a cleaner that contains ammonia, and then, when this doesn’t work, they mix it with a cleaner (such as Chlorox) that contains the hypochlorite ion. How many pairs of nonbonding electrons can be found on each nitrogen atom in the hydrazine molecule? (Note: Assume that there is a bond between the two nitrogen atoms in hydrazine.) (a) 0 (b) 1 (c) 2 (d) 3 (e) The two N atoms contain a different number of nonbonding electrons.

Answer: (b)

Bond Angles 81. Which of the following compounds/ions contains a bond angle of 120? (a) CH3OH (b) CO32– (c) CO2 (d) CH3Cl (e) CH2Cl2 Answer: (b)

82. Estimate the hydrogen−phosphorus−hydrogen bond angle in PH3. (a) approximately 90 (b) approximately 109 (c) approximately 120 (d) approximately 180 (e) approximately 150 Answer: (b)

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83. Which one of the following molecular geometries has bond angles which are not all equal? (a) tetrahedral (b) trigonal planar (c) trigonal bipyramidal (d) octahedral (e) none of these Answer: (c)

84. For the molecule below, which of the bond angles is (are) about 120?

O CH3

C

O

H

(a) C−C=O (b) H−C−C (c) C−O−H (d) O−C=O (e) both a & d Answer: (e) 85. For the molecule below, which of the bond angles is (are) about 180?

H3C

CH3

(a) H−C−C (b) C−C−H (c) H−C−H (d) all are 180 (e) none are 180 Answer: (e)

The Difference Between Polar Bonds and Polar Molecules 86. Which of the following molecules is most polar? (a) CF4 (b) BrCl5 (c) PCl5 (d) SCl6 (e) XeF2

4-25

Answer: (b) 87. Which of the following species does not have a dipole moment? (a) CH2Cl2 (b) PCl3 (c) ICl2+ (d) SF4 (e) ICl4– Answer: (e)

88. Which of the following molecules is (are) polar? (a) CO2 (b) SO2 (c) SF6 (d) O3 (e) both b and d Answer: (e)

89. Which of the following molecules is (are) nonpolar? (a) OIF5 (b) ClF3 (c) XeF2 (d) O2 (e) both c and d Answer: (e)

90. Which of the following factors influence the polarity of a bond between two atoms? (a) The number of non-bonding pairs around the two atoms. (b) The bond distance between the atoms (c) The partial charge on the two atoms. (d) Both the number of non-bonding pairs and the partial charge. (e) Both the bond distance and the partial charge.

Answer: (e)

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91. CO2 is a non-polar molecule but SO2 is slightly polar. Use these two molecules to explain how two molecules with similar formulas can have different polarities.

Answer: The two molecules have different molecular geometries. CO2 is a linear molecule. The C=O bonds have equal polarity and are 180 apart, making the molecule nonpolar. SO2 has a bent molecular geometry with two bonding domains and one nonbonding domain surrounding the sulfur, resulting in a net molecular dipole.

Integrated Problems 92. Bond dissociation energies are defined as the energy per bond required to break apart 1 mole of a gas-phase molecule into its atoms in the gas phase.

These values for bonds between different atoms are listed in tables and represents average energies of bonds over many different molecules. For example the actual strength of a nitrogen-oxygen bond will vary slightly from one molecule to another. Consider NO+, NO3–, HNO. Order these three molecules in terms of increasing NO bond strength (bond energy). Show work to support your answer and explain your reasoning.

Answer: NO3– < HNO < NO+ NO+ contains an NO triple bond. NO3– and HNO both contain an NO double bond. However, the NO double bond in NO3– can occur in three possible positions. Therefore, it has three resonance forms with three equivalent bonds with each bond behaving as if it was a bond and a third. Triple bonds are more difficult to break than double bonds which are more difficult to break than a bond and a third.

93. The best way for the body to get rid of excess nitrogen from the metabolism of proteins is by excreting urea, a planar molecule which has the following Lewis structure.

O X H2N

NH2

Which element does the "X" stand for in this Lewis structure? (a) C (b) N (c) O (d) P (e) S Answer: (a)

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94. Partial charge, formal charge and oxidation number are all used to describe the arrangement of electrons or charge associated with atoms within a molecule. Determine the formal charge, partial charge and oxidation number on both atoms in IBr. Give a one sentence description for each of the three calculations describing what the calculation tells you. Which of these three calculations most accurately describes the distribution of charge between I and Br? Br

I

formal charge

________

_________

partial charge

________

_________

oxidation number

________

_________

Answer: Br: 0, –0.065, –1; I: 0, +0.065, +1 Formal charge describes how the charge in a bond would be distributed if the bond was a pure covalent bond. It is a bookkeeping device that is useful for determining the best Lewis structure. Oxidation numbers are a bookkeeping device that treat bonds as if they were ionic. They are useful for determining the chemical formula of an ionic compound and in determining if an oxidation-reduction reaction has occurred. Partial charge is the best description of the actual distribution of electron density between two atoms.

95. The partial charge on F in the molecule ClF is ____ ENCl = 3.16 ENF = 3.98 (a) – 1 (b) – 0.11 (c) 0 (d) + 0.11 (e) none of the above Answer: (b)

96. If there is a single pair of nonbonding electrons on the "X" atom in the Lewis structure of XF4, in which group of the periodic table would you be the most likely to find element "X"? (a) Group IlIA (b) Group IVA (c) Group VA (d) Group VIA (e) Group VIIA

4-28

Answer: (d)

97. What would be the distribution of electrons in the valence shell of the "X" atom in the previous question? (a) trigonal planar (b) tetrahedral (c) trigonal bipyramidal (d) octahedral (e) none of these Answer: (c) 98. Which element would form a neutral XF3 molecule with a trigonal planar geometry? (a) Be (b) B (c) C (d) N (e) Cl Answer: (b) 99. The compound XCl3 has a trigonal pyramidal geometry. Element Z is a member of: (a) Group IIA (b) Group IlIA (c) Group VA (d) Group VIA (e) Group VIIA Answer: (c)

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Molecular Orbital Theory 100. Use the following molecular orbital diagram to predict which of the following has the largest bond order:  * 2p

x*

y*

2p

2p  2p x

y 2s*

2s

2s 2s

(a) NO+ (b) NO (c) NO– (d) NO2– (e) all of the above have the same bond order Answer: (a) 101. Use the same molecular orbital diagram to predict which of the following molecules should be paramagnetic.

(I) B2 (II) C2 (III) N2 (IV) O2 (V) F2 (a) I (b) I and II (c) I and IV (d) I, II and IV (e) II, III and IV Answer: (c)

Hybrid Atomic Orbitals 102. Which type of hybrid orbitals are used by the carbon atom in CO2? (a) sp (b) sp2 (c) sp3 (d) dsp3 4-30

(e) d2sp3 Answer: (a) 103. In which of the following compounds is the central nitrogen atom not sp2 hybridized? (a) NO2+ (b) NO2 (c) NO2– (d) NO3– (e) NO2 and NO2– Answer: (a)

104. What is the hybridization of the nitrogen atom in the N2 molecule? (a) sp (b) sp2 (c) sp3 (d) sp3d (e) sp3d2 Answer: (a) 105. What is the hybridization of the central iodine atom in the I3– ion? (a) sp (b) sp2 (c) sp3 (d) sp3d (e) sp3d2 Answer: (d) 106. What is the hybridization of the sulfur atom in the SF3+ ion? (a) sp (b) sp2 (c) sp3 (d) sp3d (e) sp3d2 Answer: (c)

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107. What is the hybridization of the beryllium atom in BeF2? (a) sp (b) sp2 (c) sp3 (d) sp3d (e) sp3d2 Answer: (a) 108. What is the hybridization of the xenon atom in the XeF3+ ion? (a) sp (b) sp2 (c) sp3 (d) sp3d (e) sp3d2 Answer: (d) 109. What is the hybridization of the nitrogen atom in the NO2+ ion? (a) sp (b) sp2 (c) sp3 (d) sp3d (e) sp3d2 Answer: (a)

110. The central atom in all of the following compounds has the same hybridization, except: (a) NO2– (b) CO32– (c) BF3 (d) NH3 (e) the hybridization is the same in all of these compounds Answer: (d)

111. Which of the following set of hybrid orbitals would give rise to a trigonal bipyramidal geometry?

4-32

(a) sp (b) sp2 (c) sp3 (d) sp3d (e) sp3d2 Answer: (d)

4-33

Chapter 5: Ionic and Metallic Bonds The Active Metals 1.

Which of the following sets of elements is arranged in order of increasing metallic character? (a) Al < Cs < Sr (b) Ba < Sr < Ca (c) Na < Mg < Al (d) Al < Ca < Ba (e) none of these

Answer: (d)

2.

Which of the following metals is the most “active,” the metal which should react most rapidly with oxygen or water vapor in the atmosphere? (a) Be (b) Ca (c) Ba (d) Al (e) Ga

Answer: (c)

3.

Which of the following metals would be the most reactive towards air and water? (a) Ca (b) Al (c) Ag (d) Sn (e) Pb

Answer: (a)

5-1

Predicting Formulas of Ionic Compounds 4.

Predict the formula of the neutral ionic compound formed from K + and S2–. (a) KS (b) K2S (c) KS3 (d) K2S3 (e) none of these

Answer: (b)

5.

Predict the formula of the neutral ionic compound formed from Mg 2+ and PO43–. (a) MgPO4 (b) Mg3PO4 (c) Mg(PO4)2 (d) Mg3(PO4)2 (e) none of these

Answer: (d) 6.

Predict the formula of the electrically neutral compound formed from the following ions: Ca2+ and PO43–. (a) CaPO4 (b) Ca2PO4 (c) Ca3(PO4)2 (d) Ca2(PO4)3 (e) none of these

Answer: (c)

7.

Predict the formula of the electrically neutral compound formed from the following ions: K+ and CrO42–. (a) KCrO4 (b) K(CrO4)2 (c) K3(CrO4)2 (d) K2(CrO4)3 (e) none of these

Answer: (e)

5-2

8.

Predict the formula of the ionic compound formed from the following ions: Na+ and S2–. (a) NaS (b) NaS2 (c) Na2S (d) Na2S3 (e) none of these

Answer: (c)

9.

Predict the formula of the ionic compound formed between an aluminum cation, Al3+, and a carbonate polyatomic anion. (a) AlCO3 (b) Al3CO3 (c) Al2CO3 (d) Al3(CO3)2 (e) Al2(CO3)3

Answer: (e)

10. A mineral called montmorillonite will catalyze some interesting reactions.

The formula of one type of montmorillonite is NaCa 4(Si2O5)4F. If Na has a charge of +1, Ca has a charge of +2, and F has a charge of -1, what is the charge on each Si2O5 complex ion? (a) +2 (b) +3 (c) -2 (d) +4 (e) -4 Answer: (c)

11. Aluminum chlorhydrate is added to antiperspirants to stop people from sweating. If this compound contains neutral H2O molecules and Al3+, OH–, and Cl– ions, and the formula of this compound is Alx(OH)5Cl • 2 H2O, what is the value of x in this formula? (a) 0.5 (b) 1 (c) 1.5 (d) 2 (e) 3 5-3

Answer: (d)

12. Predict the formulas for neutral compounds containing the following pairs of ions. (a) H+ and O22– (b) Zn2+ and PO43– (c) K+ and [PtCl6]2– Answer: H2O2; Zn3(PO4)2; K2[PtCl6]

13.

What is the formula of the neutral compound which might be formed from P5+ and S2–? (a) PS (b) PS2 (c) PS3 (d) P2S5 (e) P5S2

Answer: (d)

14. If the formula for aluminum oxide is Al2O3, what would be the formulas for aluminum fluoride and aluminum sulfide?

Answer: AlF3; Al2S3 15. What is the charge on the [Co(NO2)6]x ion if this complex ion is formed by combining Co3+ and NO2- ions? (a) +3 (b) 0 (c) -3 (d) -6 (e) -12 Answer: (c)

Metals, Nonmetals and Their Ions

5-4

16. A main-group metal reacts with hydrogen and oxygen to form compounds with the formulas XH4 and XO2 (X represents the metal). In which group of the periodic table does this element belong? (a) IA (b) IIA (c) IIIA (d) IVA (e) none of these Answer: (d) 17. What is the most likely charge for a calcium ion? (a) -2 (b) -1 (c) 0 (d) +1 (e) +2 Answer: (e)

18. Which of the following elements would be the most likely to form an oxide with the formula XO and a hydride with the formula XH2 (X represents the unknown element)? (a) Na (b) Mg (c) Al (d) Si (e) P Answer: (b)

19. An element, X, forms the ionic compound CaX. X is a member of: (a) Group IA (b) Group IIA (c) Group IVA (d) Group VIA (e) Group VIIA Answer: (d)

5-5

20. Magnesium metal reacts with hydrogen gas to form a white solid (A) with a high melting point. Give the formula of the product and the balanced reaction for its formation.

Answer: MgH2

Mg(s) + H2(g) → MgH2(s)

21. Which one of the following elements is most likely to form a stable -3 anion? (a) P (b) C (c) Ar (d) Be (e) B Answer: (a) 22. Which element is hydrogen most like to combine with to form a hydride? (a) C (b) S (c) N (d) Na (e) O Answer: (d) 23. What is the most likely electronic configuration of the Co2+ ion? (a) [Ar]3d7 (b) [Ar]4s23d7 (c) [Ar]4s23d5 (d) [Kr]5s24d7 (e) None of the above Answer: (a) 24. What is the most likely electronic configuration of the Ru2+ ion? (a) [Kr]4d6 (b) [Kr]5s24d6 (c) [Kr]5s24d4 (d) [Ar]4s23d6 (e) None of the above Answer: (a)

5-6

25. An element reacts with hydrogen and oxygen to form ionic compounds with the formulas MH4 and MO2. In which group of the periodic table does this element, M, belong? (a) IIA (b) IVA (e) VA (d) VIA (e) VIIA Answer: (b)

26. What group of metals react with sulfur to form M2S3 sulfides, reacts with fluorine to form MF3 fluorides, and react with acid to form M3+ ions and H2 gas? (a) IA (b) IIA (c) IIIA (d) IVA (e) VA Answer: (c) 27. Which of the following elements would be the most likely to react with hydrogen to form a compound with the formula XH (X represents the metal)? (a) Na (b) Mg (c) Al (d) Si (e) P Answer: (a) 28. The two most common ions of copper are Cu1+ and Cu2+. Give the most likely electron configuration of each of these ions.

Answer: Cu+ [Ar] 3d10 Cu2+ [Ar] 3d9

29. Which ionic compound contains the superoxide ion? (a) CaO2 (b) LiO2 (c) Na2O2

5-7

(d) H2O2 (e) BaO Answer: (b) 30. Which ionic compound contains the peroxide ion? (a) CaO (b) LiO2 (c) Na2O2 (d) Na2O (e) None of the above Answer: (c)

Predicting the Products of Reactions That Produce Ionic Compounds 31. Which of the following is the most likely product of the reaction between magnesium metal and nitrogen? (a) MgN (b) Mg2N (c) MgN2 (d) Mg2N3 (e) Mg3N2 Answer: (e) 32. Use the positions of gallium and oxygen in the periodic table to predict the formula for gallium oxide. (a) GaO (b) GaO2 (c) GaO3 (d) Ga2O (e) Ga2O3 Answer: (e)

33. Predict the product of the following reaction: Sr(s) + P4(s). (a) SrP (b) SrP2 (c) SrP3 (d) Sr2P3 (e) Sr3P2

5-8

Answer: (e)

34. What would be the product of the reaction between aluminum metal and sulfur? (a) AlS (b) Al2S (c) AlS2 (d) AlS3 (e) Al2S3 Answer: (e) 35.

What compound is formed by the reaction of Na with P4? (a) NaP3 (b) Na3P (c) NaP2 (d) Na2P (e) NaP

Answer: (b) 36. Use the positions of silicon and fluorine in the periodic table to predict the most likely product of the reaction between silicon and fluorine. (a) SiF (b) SiF2 (c) SiF4 (d) Si2F (e) Si3F4 Answer: (c) 37. Predict the product of the reaction between barium and nitrogen. (a) BaN (b) BaN2 (c) Ba2N (d) BaN3 (e) none of these Answer: (e) 38. An element X forms a compound with oxygen, X2O3, and with hydrogen, XH3. The element X must be in which group in the periodic table?

5-9

(a) IA (b) IIA (c) IIIA (d) IVA (e) VIA Answer: (c)

Bond-Type Triangles: The Relationship Among Ionic, Covalent, and Metallic Bonds 39. In Chapter 4 we were able to determine the geometric shape and resulting polarity of individual covalent molecules. In Chapter 5 we did not do this for ionic and metallic compounds. Why not? (a) Ionic and metallic compounds exist as three dimensional networks not as individual molecules. (b) The electrons in ionic compounds are transferred rather than being shared as in covalent compounds. (c) The difference in electronegativities of the bonded atoms is usually greater in ionic than in covalent and metallic compounds. (d) Covalent molecules are much larger than ionic compounds. (e) all of these Answer: (a)

40. If the difference in electronegativity between two elements in a compound is very small (less then 0.2), can you safely conclude that the compound formed from the two elements is covalent? (a) Yes, because a small difference in electronegativity always indicates that a compound is covalent. (b) No, because covalent compounds are formed between two elements with a large difference in electronegativity. (c) Electronegativity has nothing to do with whether a bond is covalent or not. (d) No, because a metallic compound could also be formed between two elements with a small difference in electronegativity. (e) None of the above are correct.

Answer: (d)

41. Compare and contrast (what is the same and what is different) the distribution of electrons in covalent, ionic and metallic bonds.

5-10

Answer: All three types of bonds are formed by the electrostatic attraction between positive and negative charges. Covalent and metallic bonding both involve the sharing of electrons and ionic bonding involves the transfer of electron density. In covalent and ionic bonding the electrons are localized and in metallic bonding the electrons are delocalized.

The bond-type triangle can be used for Questions 42,43,46-49 3.6

CsF

3.4 BaF 2 3.2

CaF 2

3.0

Cs 2O

MgF 2 2.8 BaO 2.6

CaO

2.4 MgO

SiF4

CsCl CsBr Cs 2S CsI CsH

CsCd 13

Be 13Ca

Cs Rb

Na Ca Ba Sr Li

1.4 Mg Ti

SO 2

1.0

N 2F2

NP

Be 13 Ba

1.2

1.2

GaI 3 SiC

AlAs AlSb

1.6

HCl

GaAs SiH 4

0.4 CCl 4

SCl 2

Br 2 2.0

2.2

0.6

OF 2

NO

ICl

CH4

1.8

0.8

GaBr 3

Cd 3 Mg 1.0

CO 2

BN

ZnS

BeC 2

CIF

SeO 2

Hg 2Cl2

MgH 2

1.6 1.4

P4O 7

AlCl 3

EN

1.8 SF4

B 2O 3 AlN

MgBr 2

AIH 3

MgZn 2

0.8

SiO 2

Be 3N 2

BaZn 5

CdLi

HF IF5

HgO

BaI 2

BaSi 2

2.0

ZnO

Mg 3N 2

Li 3 As CaH 2

Cs 3Sb

BeO

BaCl 2 BaBr 2 BaS BaC 2

2.2

BF 3

Ca 3N 2

2.4

2.6

Cd Al In Ga Si Ge B As H Se C Zn Cu Ni Ag P I Be Sn Sb Te Mn Hg

S

2.8 Br

N2

3.0 Cl

0.2

NCl 3

CI 2

O2 3.2

N

3.4

3.6 O

3.8

4.0

F2 4.2 F

Average EN Copyright 1999 by John Wiley and Sons Inc. All rights reserved.

42. Based on the bond-type triangle, which of these covalent compounds has the greatest ionic character? (a) NCl3 (b) SCl2 (c) ICl (d) HCl

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(e) SiC Answer: (e) 43. Using a bond-type triangle, predict which of the following compounds would have approximately 50% ionic bonding and 50% covalent bonding. (a) AlCl3 (b) Li3As (c) AlAs (d) BN (e) none of these Answer: (a)

44. Describe what happens to the distribution of electrons in the bond between two elements as you move from left to right on a bond-type triangle.

Describe what happens to the distribution of the electrons in a bond between two elements as you move from the bottom to the top of a bondtype triangle.

Answer: As you move from left to right the electrons involved in bonding become more localized. As you move from the bottom to the top the electron density is transferred more completely from one atom to another rather than being shared.

45. Use a bond-type triangle to predict the correct arrangement of the following compounds in order of increasing covalent character.

Mg Zn2, CCl4, SiH4 (a) Mg Zn2 < CCl4 < SiH4 (b) CCl4 < Mg Zn2 < SiH4 (c) CCl4 < SiH4 < Mg Zn2 (d) MgZn2 < SiH4 < CCl4 (e) none of these Answer: (d)

46. Use a bond-type triangle to describe the bonding in CaH2. (a) pure ionic (b) pure covalent (c) pure metallic (d) significant amount of both ionic and covalent character

5-12

(e) significant amount of both ionic and metallic character Answer: (e) 47. As you move from the bottom to the top on a bond-type triangle, _________. (a) the solids can increasingly be described as an array of cations in a sea of electrons (b) compounds exist mostly as gases (c) the solid compounds become more conductive (d) the electron density in a bond is increasingly shared between two atoms, rather than being transferred (e) the electrons in a bond are transferred from one atom to another, rather than being shared Answer: (e) 48. According to the supplied bond-type triangle, which of the following compounds is most likely to conduct electricity as a solid? (a) SO2 (b) MgBr2 (c) CdLi (d) NCl3 (e) N2F2 Answer: (c)

49. Based on the given bond-type triangle, what is the highest possible average electronegativity a compound can have and still retain some metallic character? (a) Approximately 2.9 (b) Approximately 2.7 (c) Approximately 1.9 (d) Approximately 1.4 (e) A compound can have metallic character at any average electronegativity.

Answer: (c)

Oxidation Numbers 50. Determine the oxidation number of sulfur in potassium sulfate, K 2SO4. (a) +2 5-13

(b) +4 (c) +6 (d) +8 (e) none of these Answer: (c)

51. Determine the oxidation number of the carbon in bold in the following compound: CH3CCl3. (a) 0 (b) –3 (c) +1 (d) +3 (e) none of these Answer: (d)

52. Determine the oxidation number of sulfur in sulfurous acid, H2SO3. (a) +6 (b) +4 (c) +2 (d) –2 (e) none of these Answer: (b)

53. Determine the oxidation number of the central carbon atom in the following compound.

O

H3C

C

Cl

(a) +2 (b) –2 (c) +4 (d) –4 (e) none of these Answer: (e)

54. Determine the oxidation number of chlorine in NaClO3. (a) –1 (b) +1

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(c) +3 (d) +7 (e) none of these Answer: (e)

55.

Determine the oxidation number of the second (bolded) carbon in CH3CH2NH2. (a) –1 (b) +1 (c) +3 (d) –3 (e) none of these

Answer: (a)

56. Which of the following compounds has sulfur in a +6 oxidation state? (a) SO2 (b) H2S (c) H2SO3 (d) H2SO4 (e) H2S2O3 Answer: (d)

57. What is the maximum possible oxidation state of a P atom? (a) +2 (b) +3 (c) +4 (d) +5 (e) none of the above Answer: (d)

58. What is the oxidation state of the Mo atom in Li2MoO4? (a) 0 (b) +2 (c) +4 (d) +5 (e) +6

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Answer: (e)

59. All but one of the following species contains nitrogen in the same oxidation state. Which one is different? (a) HNO2 (b) NH2– (c) NH3 (d) NH4Cl (e) Na3N Answer: (a)

60. Which compounds contain hydrogen in a negative oxidation state?

(I) H2S (II) PH4+ (III) LiAlH4 (IV) CH4 (V) CaH2 (a) I, II and IV (b) II, III and IV (c) III and IV (d) III and V (e) none of these contain hydrogen in a negative oxidation state Answer: (d)

61. Honda and Toyota both sell hybrid cars that have engines that can run on either gasoline or electricity stored in a battery. The advantage of hybrid cars is simple: Totally electric cars need batteries that weigh almost as much as the vehicle! It isn’t surprising that a high priority is the search for better batteries. Barium ferrate (BaFeO4) is being tested for use in batteries.

Use the positions of barium and oxygen in the periodic table to predict the oxidation state of the iron atom in this compound. (a) +2 (b) +3 (c) +4 (d) +6 (e) +8 Answer: (d)

62. Lithium ion batteries are high energy density 3.6 V batteries that use a lithium salt such as LiCoO2 or LiPF6 for the positive electrode, or cathode.

Use the positions of lithium and fluorine in the periodic table to predict the oxidation state of the phosphorus atom in LiPF6.

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(a) +1 (b) +2 (c) +3 (d) +4 (e) +5 Answer: (e)

Oxidation-Reduction Reactions 63. In the following redox reaction which element is being oxidized?

Cu(s) + 4 HNO3(aq) → Cu(NO3)2 (aq) + 2 NO2 (g) + 2 H2O(l) (a) Cu (b) H (c) N (d) O (e) none of these Answer: (a) 64. Which one of the following is an oxidation-reduction reaction? (a) CaCO3 (s) → CaO (s) + CO2 (g) (b) CaO (s) + 2 HCl (aq) → CaCl2 (aq) + H2O (l) (c) 2 KNO3 (s) → 2 KNO2 (s) + O2 (g) (d) HCl (aq) + NaOH (aq) → NaCl (aq) + H2O (l) (e) none of these Answer: (c) 65. In the following oxidation-reduction reaction, which of the following statements is true?

Sr(s) + 2 H2O(l) → Sr2+(aq) + 2 OH–(aq) + H2(g) (a) H2O is oxidized (b) Sr acts as the oxidizing agent (c) H2 is reduced (d) H2O is the oxidizing agent (e) more than one of the above are true Answer: (d)

66. Phosphorus is reduced in which of the following reactions?

5-17

(a) P4(s) + 5 O2(g) → P4O10(s) (b) P4(s) + 3 O2(g) → P4O6(s) (c) P4(s) + 6 F2(g) → 4 PF3(g) (d) P4(s) + 6 Ca(s) → 2 Ca3P2(s) + H2O(l) (e) P4(s) + 6 Cl2(g) → 4 H3PO3(aq) + 12 HCl(aq) Answer: (d) 67. Which isn’t an oxidation-reduction reaction? (a) 2 Na (s) + Cl2 (g) → 2 NaCl (s) (b) 2 HCl (g) → H2 (g) + Cl2 (g) (c) 2 NO2 (g) + H2O (l) → HNO3 (aq) + HNO2 (aq) (d) Na2CO3 (s) + 2 HCl (aq) → 2 NaCl (s) + CO2 (g) + H2O (l) (e) all of these are oxidation-reduction reactions Answer: (d) 68. Identify the reducing agent and oxidizing agent in the reaction: H2O2(aq) + 2 HI(aq) → 2 H2O(l) + I2(s) Answer: Oxidizing Agent: H2O2, Reducing Agent: HI

Nomenclature 69. What is the correct name for Ca(NO3)2? (a) calcium nitride (b) calcium nitrite (c) calcium nitrate (d) calcium dinitrite (e) calcium dinitrogen hexaoxide Answer: (c) 70. The name for Li3N is: (a) lithium nitride (b) lithium(II) nitride (c) lithium(III) nitride (d) lithium trinitride (e) trilithium nitride Answer: (a)

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71. Which of the following formula/name combinations is incorrect? (a) H2SO3/sulfurous acid (b) HCO3–/hydrogen carbonate ion (c) HBrO3/bromic acid (d) ClO4–/hypochlorate ion (e) CuSO4/copper(II) sulfate Answer: (d)

72. Name this compound: CoSO3 (a) copper persulfate (b) cobalt(II) sulfite (c) cupric sulfate (d) cobalt (II) sulfide (e) cupric sulfite Answer: (b)

73. Name this compound: MgCl2 (a) magnesium dichloride (b) magnesium chloride (c) manganese chloride (d) magnesium chlorate (e) manganese dichoride Answer: (b)

74. Name this compound: SrSO4 (a) strontium sulfide (b) sodium sulfite (c) strontium persulfate (d) sodium sulfurtetroxide (e) strontium sulfate Answer: (e)

75. Give the chemical formula for barium chromate. (a) Ba2CrO4 (b) Ba2CrO3

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(c) BaCr (d) BaCrO4 (e) BCrO3 Answer: (d)

76. Give the chemical formula for ammonium bromide. (a) NBr3 (b) NH4Br (c) AmBr2 (d) NH4Br2 (e) NH4Br3 Answer: (b)

77. Give the chemical formula for sodium phosphite. (a) Na3PO3 (b) SoPO (c) Na2PO4 (d) Na3P2O4 (e) SPO4 Answer: (a)

78. Give the name for the compound Fe2O3. (a) diiron trioxide (b) iron (III) oxide (c) iron (II) trioxide (d) iron oxide (e) iron trioxide Answer: (b)

79. Give the chemical formula for perchloric acid. (a) HClO3 (b) HCl (c) NaHCO3 (d) HClO4 (e) H2ClO3 Answer: (d)

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80. Give the name for the compound ZnSO3. (a) zinc persulfate (b) zinc sulfur trioxide (c) zinc (II) sulfate (d) zinc (I) sulfide (e) zinc sulfite Answer: (e)

81. Give the chemical formula for ammonium cyanide. (a) HCN (b) NH4CN (c) NH2C (d) CH3CN (e) AmCN Answer: (b)

82. Give the name for the compound Cl2O. (a) dichoro oxide (b) chlorine oxide (c) dichlorine monoxide (d) chlorite (e) chlorine dioxide Answer: (c)

83. Give the chemical formula for hyposulfurous acid. (a) H2SO2 (b) H2S (c) NaHSO3 (d) HSO3 (e) H2SO4 Answer: (a)

84. Give the name for the compound HNO2. (a) hydrogen nitride

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(b) pernitric acid (c) nitrous acid (d) hydrogen nitrogen dioxide (e) nitric acid Answer: (c)

85. Give the chemical formula for copper(I)chloride. (a) CoCl2 (b) CuCl (c) CuClO2 (d) CuCl2 (e) CoCl Answer: (b)

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Chapter 6: Gases Temperature 1.

Until recently, materials had to be cooled to temperatures on the order of 4 K to become superconducting. A ceramic has been discovered that becomes superconducting at 92 K. Which of the following temperatures corresponds to 92 K? (a) 32 °F (b) –181°C (c) 212°F (d) 181°C (e) 365 °C

Answer: (b)

2.

Normal body temperature is 98.6 °F. What is the equivalent temperature on the Kelvin scale? (a) 236 K (b) 293 K (c) 310 K (d) 393 K (e) 418 K

Answer: (c)

3.

At what temperature are the readings on the Celsius and Fahrenheit scales the same? At what temperature are the magnitudes the same but the signs different?

Answer: –40 °C = –40 °F; –11.4 °C = 11.4 °F

4.

What is the melting point of table salt (NaCl) in degrees Fahrenheit, if salt melts at 801°C? (a) 463 °F (b) 477 °F (c) 1441 °F (d) 1474 °F (e) 1499 °F

Answer: (d)

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5.

Fahrenheit defined 0° as the lowest temperature that could be achieved by adding salt to ice and found that ice by itself freezes at 32° on this scale.

Unfortunately, Fahrenheit's measurements were wrong. The lowest temperature that can be achieved by adding salt to ice is -21.2 °C. Calculate the equivalent temperature on the Fahrenheit scale. (a) -38.2 °F (b) -11.8 °F (c) -6.2 °F (d) 20 °F (e) 32 °F

Answer: (c) 6.

The Kelvin scale is an absolute temperature scale on which one unit, or one Kelvin, is equal to one unit, or one degree, on the Celsius scale. The Rankine scale is an absolute temperature scale on which one unit is equal to one degree on the Fahrenheit scale. Calculate the temperature on the Fahrenheit scale that corresponds to zero on the Rankine scale.

Answer: -459.7° 7.

Which of the following elements or compounds is most likely to be a gas at room temperature? (a) a metal such as sodium (Na) (b) an alloy of two metals such as bronze, which is a mixture of copper (Cu) and tin (Sn) (c) a salt such as MgO (d) a nonmetal such as chlorine (Cl2) (e) none of these elements or compounds will be a gas at room temperature

Answer: (d)

Gas Laws 8.

Which of the following relationships between the two variables listed for an ideal gas does not give a straight line graph with a positive slope? (a) P and T (b) n and V (c) P and V (d) T and V (e) none of these

Answer: (c) 6-2

9.

How many cm3 of liquid SO2 (d = 1.46 g/cm3) can be obtained by compressing 1.00 liter of the gas collected at 25°C and 1 atm?

Answer: 1.79 cm3

10. Three containers of equal volume are filled with three different gases. The first container is filled with Ar, the second with O2 and the third with CO2.

The pressure and temperature of the three gases are the same. Which of the following is true of the gases in the three containers? (a) The number of moles of gas in the containers is the same. (b) The masses of gas in the three containers are the same. (c) The density of gas in the three containers is the same. (d) (a), (b), and (c) are all true. (e) None of these are true.

Answer: (a)

11. Samples of oxygen and helium have been collected in two separate containers of equal volume. Both gases have the same temperature and pressure. Which of the following statements is true? (a) The container of oxygen has the larger number of moles of gas. (b) The container of helium has the larger mass of gas. (c) The container of helium has the larger number of moles of gas. (d) Both containers have the same mass of gas. (e) Both containers have the same number of moles of gas.

Answer: (e)

12. The density of neon gas will be greatest at: (a) 0C and 1 atm (b) 273 C and 1.5 atm (c) 0C and 2 atm (d) 273C and 1 atm (e) The density of neon will be the same for all of the above cases.

Answer: (c)

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13. Calculate the volume in liters occupied by 2.0 g of Ne at 50.0C and 1.5 atm of pressure. (a) 1.1 L (b) 8.8 L (c) 35 L (d) 47 L (e) none of these Answer: (e)

14. Ether was one of the first anesthetics discovered. Assume that several mL of ether are placed in a bulb with a volume of 293 mL, and the bulb is immersed in water at 36°C until the last drop of liquid disappears, leaving the bulb filled with ether vapor. The bulb is then removed from the water bath, and the weight of ether that condenses in the bulb is measured. If 0.841 grams of ether collect in this experiment at a pressure of 746 mmHg, what is the molecular weight of ether?

Answer: 74.2 g/mol

15. Which sample would have the largest volume at 25°C and 750 mmHg? (a) 25 g NH3 (b) 25 g C2H2 (c) 25 g NO2 (d) 25 g Ar (e) All of these samples would have the same volume at 25°C and 750 mmHg. Answer: (a)

16. What is the volume, in liters, of 1.00 mole of an ideal gas at 25.0°C and 1.00 atm pressure? (a) 0.0409 L (b) 2.05 L (c) 22.4 L (d) 24.5 L (e) none of the above Answer: (d)

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A small gas cylinder of helium used in chemistry lecture demonstrations has a volume of 275 mL at a pressure of 1823 kPa at 24.7°C. Use this information to answer questions 17-19.

17. The volume of helium, in mL, needed to fill the cylinder at 772 torr and 24.7°C is: (a) 45.9 mL (b) 303 mL (c) 4.87 x 103 mL (d) 4.95 x 104 mL (e) 4.86 x 105 mL Answer: (c) 18. At what temperature in degrees Celsius would the helium gas in the cylinder exhibit a pressure of 25.0 atm? (a) 4.96°C (b) 34.3°C (c) 58.8°C (d) 140°C (e) 307°C Answer: (d) 19. The number of grams of helium in the cylinder is: (a) 0.202 g (b) 0.810 g (c) 1.24 g (d) 4.96 g (e) none of these Answer: (b) 20. Polypropylene is a plastic formed by polymerizing propylene, which has the empirical formula CH2. If a sample of propylene with a mass of 21.0 grams occupies a volume of 11.2 L at STP (0 °C and 1 atm), what is the molecular formula of this gas? (a) CH2 (b) C2H4 (c) C3H6 (d) C4H8 (e) none of these Answer: (c) 6-5

21. A 2.91 gram sample of a gaseous compound that contains only boron and hydrogen has a volume of 1.22 L at 25°C and 1.09 atm. What is the formula of this compound? (a) B2H6 (b) B4H10 (c) B5H9 (d) B6H10 (e) B6H12 Answer: (b)

22. What volume of NH3(g) at STP is required to prepare 49.8 g of (NH4)2SO4 by the reaction of ammonia with sulfuric acid? (a) 0.377 L (b) 6.41 L (c) 8.4 L (d) 12.84 L (e) 16.9 L Answer: (e)

23. Uranium reacts with fluorine to produce a compound that is a gas at 57°C. The density of this gas is 13.0 g/L at 57°C and 1 atm pressure. What is the molecular formula of this compound? (a) UF2 (b) UF3 (c) UF4 (d) UF5 (e) UF6 Answer: (e)

24. Which of the following gases would have the largest density at 25°C and 1.00 atm pressure? (a) methane, CH4 (b) acetylene, C2H2 (c) ethylene, C2H4 (d) ethane, C2H6 (e) propane, C3H8

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Answer: (e)

25. Which of the following gases has a density of 1.72 g/L at 10°C and 1 atm pressure? (a) He (b) Ne (c) Ar (d) CH4 (e) P4 Answer: (c)

26. What is the density of NO(g) at 25.0°C and 783 mmHg? (a) 1.26 g/L (b) 2.68 g/L (c) 3.12 g/L (d) 3.76 g/L (e) 22.4 g/L Answer: (a)

27. Nitrogen gas reacts with hydrogen gas to produce gaseous ammonia (NH 3).

How many liters of product at STP can be formed if 6 L of hydrogen gas at STP reacts with a stoichiometric amount of nitrogen gas? (a) 1 L (b) 2 L (c) 3 L (d) 4 L (e) 6 L Answer: (d)

28. We want to increase the volume of a fixed amount of gas from 68.2 to 153 mL while holding the pressure constant. To what temperature in degrees Celsius must we heat this gas if the initial temperature is 21.0 °C? (a) 47.1 °C (b) 387 °C (c) 660 °C (d) 89.5 °C (e) not enough information given

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Answer: (b)

29. 0.500 mol of Ar gas occupies a volume of 4.07 L at 25.0 °C and 3.00 atm pressure. What volume will the gas occupy at STP (0.00°C and 1.00 atm pressure)? (a) 1.24 L (b) 1.48 L (c) 4.07 L (d) 11.2 L (e) 13.3 L Answer: (d)

30. A cylindrical container with a movable piston initially holds 1.5 moles of gas at a pressure of 4.0 atmospheres and a volume of 2.5 liters. If the piston is moved to create a volume of 5.0 liters while simultaneously withdrawing 0.75 moles of gas, what is the final pressure if the temperature is constant? (a) 0.50 atm (b) 1.0 atm (c) 2.0 atm (d) 4.0 atm (e) 8.0 atm Answer: (b)

31. What is the final temperature in kelvins if a sample of ammonia gas, initially at a pressure of 3.00 atmospheres, a temperature of 500.0 K, and a volume of 275 L is changed to a volume of 200.0 L and a pressure of 2.50 atm? (a) 303 K (b) 436 K (c) 573 K (d) 825 K (e) none of the above Answer: (a)

Partial Pressures 32. At 25°C and one atmosphere pressure, four fifths of the pressure of the atmosphere is due to N2 and one fifth is due to O2. What fraction of the pressure at 10 atm. and 100°C would be due to N2?

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Answer: four fifths

33. A gas cylinder contains the following mixture of gases: 10 g H2, 15 g CH4, 20 g SO2 and 12 g Ar. Which gas exerts the largest partial pressure in the container? (a) H2 (b) CH4 (c) SO2 (d) Ar (e) all exert the same pressure Answer: (a) 34. A sealed 2.00 L container at 20.0oC holds 150.0 g of ammonia (NH3), 150.0 g of carbon dioxide (CO2), and 150.0 g of nitrogen(N2). What is the total pressure in atmospheres? (a) 155 (b) 0. 0227 (c) 211 atm. (d) 0. 333 (e) The pressure can’t be determined from the information given.

Answer: (c)

35. What is the total pressure in atmospheres when 0.400 g of H2, 2.00 g of N2, and 10.5 g of CO2 are injected into a 10.0 liter flask at 273 K? (a) 0.571 atm (b) 0.877 atm (c) 1.14 atm (d) 1.75 atm (e) it is impossible to determine the answer from the data given Answer: (c)

36. 0.300 L of H2 gas was collected over water at 27.0°C on a day when the atmospheric pressure was 745 torr. What would be the volume of H 2 gas at 760 torr after the water vapor was removed? (PH2O = 27.0 torr at 27°C) (a) 0.283 L (b) 0.294 L (c) 0.300 L (d) 0.306 L 6-9

(e) 0.318 L Answer: (a)

37. Calculate the volume of the flask that would contain 0.40 moles of oxygen, 0.600 moles of nitrogen, and 1.50 moles of hydrogen at 25.0°C and a total pressure of 1.80 atm. Calculate the partial pressure of each gas.

Answer: 34.0 L; 0.288 atm O2, 0.43 atm N2, 1.08 atm H2.

38. Oxygen is bubbled through water before it is given to a patient in a hospital.

What volume of pure oxygen gas at 21.0°C and 750.0 torr pressure would a patient receive if the patient breathed 0.250 L of O2 bubbled through water at this temperature and pressure? (PH2O = 19.0 torr at 21°C) (a) 0.00633 L (b) 0.244 L (c) 0.250 L (d) 0.256L (e) none of the above Answer: (b)

The Kinetic Molecular Theory 39. Use the Kinetic Molecular Theory to explain on the molecular level the following observations for an ideal gas.

I) Decreasing the volume causes an increase in pressure if temperature and number of moles are held constant.

II) Increasing the temperature causes an increase in pressure if volume and number of moles are held constant.

Answer: I) As the volume decreases there will be more collisions with the sides of the container resulting in an increase in pressure.

II) As the temperature increases the molecules will move more rapidly, striking the sides of the container more often and with more force, resulting in an increase in pressure.

40. Use the Kinetic Molecular Theory of gases to explain why pressure of a gas is proportional to its temperature in a fixed volume container.

Answer: An increase in temperature causes an increase in the kinetic energy of gas particles. This increase in kinetic energy is due to an

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increase in the velocity of the gas particles. Because the gas particles are moving faster, they strike the walls of the container more frequently and with more force, resulting in increased pressure on the walls of the container.

41. Which of the following can explain the increase in pressure that results from decreasing the volume of a fixed amount of gas at a fixed temperature? (a) The gas molecules move faster and therefore hit the walls of the container with more force. (b) The gas molecules have a shorter path between the walls of the container and therefore collide with the walls more frequently. (c) The gas molecules increase in their apparent mass therefore exerting more force when they strike the walls of the container. (d) The gas molecules collide less with each other therefore having more force to exert on the walls of the container. (e) None of these.

Answer: (b)

42. Why does a hot air balloon rise when the air in the balloon is heated? (a) The average kinetic energy of the air molecules increases, and the collisions between these molecules and the walls of the balloon make it rise. (b) The pressure of the gas inside the balloon increases, pushing up on the balloon. (c) The gas expands, forcing some of the gas to escape from the bottom of the balloon, and the decrease in the density of the gas lifts the balloon. (d) The balloon expands, causing it to rise. (e) The hot air rises inside the balloon, which produces enough force to lift the balloon.

Answer: (c)

43. Which would always lead to an increase in the average kinetic energy of a gas? (a) Increasing the volume by decreasing the pressure. (b) Increasing the pressure by decreasing the volume. (c) Increasing the pressure by increasing the number of molecules of gas. (d) Increasing the volume by increasing the temperature of the gas. (e) All of the above are equally effective ways of increasing the average kinetic energy of a gas.

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Answer: (d) 44. Which isn’t a basic assumption of the kinetic theory? (a) Gases consist of a large number of tiny particles in constant random motion. (b) The distance between gas particles is large compared with their diameters, and therefore most of the volume of a gas is empty space. (c) Gas particles move in a straight line until they collide with another gas particle or the walls of the container. (d) The average kinetic energy of the particles in a gas is proportional only to the temperature of the gas. (e) All of the above are basic assumptions of the kinetic theory.

Answer: (e) 45. Which isn’t one of the postulates of the kinetic molecular theory? (a) At a constant temperature, all of the particles have the same speed. (b) Gas particles are in constant motion. (c) Gas particles move in a straight line between collisions. (d) The volume of the particles is negligibly small compared to the volume of the container. (e) There are no forces of attraction between gas particles.

Answer: (a) 46. Which isn’t a postulate of the kinetic molecular theory? (a) The mass of a molecule is negligibly small. (b) The volume of a molecule is a negligibly small fraction of the volume of the container. (c) The number of molecules in a gas is very large. (d) The molecules are in a state of constant, random motion. (e) The attractive forces between molecules are negligibly small.

Answer: (a) 47. If the average speed of O2 molecules at STP is 4.3 x 102 meters per second, what is the average speed of H2 molecules at the same temperature and pressure? (a) 1.1 x 102 (b) 2.2 x 102 (c) 8.6 x 102 (d) 1.7 x 103

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(e) none of the above Answer: (d) 48. The root-mean-square speed of CH4 molecules at 25°C is about 0.56 km/s.

What is the root-mean-square speed of a H2 molecule at 25°C? (a) 0.070 km/s (b) 0.20 km/s (c) 1.1 km/s (d) 1.6 km/s (e) 4.5 km/s Answer: (d) 49. Assume that a container is filled with a mixture of SO3 and Ne. The molecular weight of SO3 is 80 g/mol and the atomic weight of Ne is 20 g/mol. The average velocity of an SO3 molecule is: (a) one-fourth that of a Ne atom. (b) one-half that of a Ne atom. (c) the same as a Ne atom. (d) two times that of a Ne atom. (e) 4 times that of a Ne atom.

Answer: (b)

Graham’s Laws of Diffusion and Effusion 50. Which of the following gases would diffuse fastest at room temperature? (a) NH3 (b) CO (c) H2S (d) F2 (e) CO2 Answer: (a)

51. Nitrogen and oxygen are allowed to effuse through a porous barrier at 295 K. If nitrogen effuses at a rate of 0.0355 moles per minute, what is the rate of effusion of oxygen? (a) 0.0311 mol/min (b) 0.0332 mol/min (c) 0.0380 mol/min (d) 0.0405 mol/min 6-13

(e) 31.1 mol/min Answer: (b)

52.

235U (234.040947 amu) and 238U (235.043925 amu) are separated on the

basis of the relative rates of diffusion of 235UF6 and 238UF6. Calculate the relative rates of diffusion of these two compounds.

Answer: 1.0043:1, 235UF6 being slightly faster

Deviations from Ideal Gas Behavior: The van der Waals Equation 53. The ideal gas equation predicts that a plot of the pressure times the volume of a gas versus the pressure of this gas should be a horizontal straight line at a fixed temperature. When these quantities are graphed for CO 2, however, we find that PV dips far below the theoretical straight line at first.

This can be explained as follows. (a) The force of attraction between the CO2 molecules makes the product of the pressure times the volume of this gas smaller than predicted from the ideal gas equation. (b) The volume of the CO2 molecules makes the product of the pressure times the volume of this gas smaller than predicted from the ideal gas equation. (c) The force of attraction between the CO2 molecules makes the product of the pressure times the volume of this gas larger than predicted from the ideal gas equation. (d) The volume of the CO2 molecules makes the product of the pressure times the volume of this gas larger than predicted from the ideal gas equation. (e) There is no way to explain the difference between a real gas and the predictions of the ideal gas equation.

Answer: (a)

Integrated Problems 54. Calculate the volume of hydrogen bromide that can be prepared by reacting 10.0 grams of phosphorus tribromide with excess water at 21°C and 753 mmHg. PBr3(l) + 3 H2O(l) → 3 HBr(g) + H3PO3(aq) Answer: 2.65 L

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55. Calculate the volume of oxygen at 0°C and 1 atm pressure that can be produced by decomposing 100 mL of hydrogen peroxide if this solution is 27.6% H2O2 by weight and the density is 1.09 g/cm3.

2 H2O2(l) → 2 H2O(l) + O2(g) Answer: 9.91 L 56. 2.11 moles of C3H8 are allowed to burn in the presence of excess O2 to form CO2 and H2O. C3H8(g) + 5 O2(g) → 3 CO2(g) + 4 H2O(g) The carbon dioxide is isolated from all other gases and stored in a 1.25 L container at 0.00°C. What is the pressure in this container if the gas obeys the ideal gas equation? (a) 18 atm (b) 38 atm (c) 113 atm (d) 1,130 atm (e) none of the above Answer: (c)

57. What is the empirical formula of the only product of the following reaction __ N2O3(g) + __ O3(g) → if 3.00 L of N2O3 react completely with 1.00 L of O3 to give 6.00 L of the product, all gases measured at STP? (a) NO (b) NO2 (c) NO3 (d) N2O5 (e) none of the above Answer: (b)

58. Two identical flasks are at the same temperature. One is filled with 2 grams of hydrogen, the other with 28 grams of nitrogen gas. Which property would be significantly different for the two samples? (a) pressure (b) average kinetic energy (c) density (d) the number of molecules in each container (e) the number of moles of gas in each container

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Answer: (c)

59. Cyclopropane, a gas containing only carbon and hydrogen, is an anesthetic.

If 0.450 L of cyclopropane at 120°C and 0.720 atm reacts with excess O 2 to give 1.35 L of CO2(g) and 1.35 L of H2O(g) at the same temperature and pressure, what is the percent by weight of carbon in cyclopropane? (a) 33.3% (b) 41.8% (c) 45.0% (d) 85.7% (e) 90.0% Answer: (d)

60. A metal reacts with acid to form hydrogen gas as shown by the following equation.

M(s) + 2 H+(aq) → M2+(aq) + H2(g) What is the atomic weight of this metal if 3.49 grams of the metal generates enough hydrogen collected over water to fill a bottle with a volume of 2.20 L at 25.0°C and 1.00 atmosphere pressure under conditions where the vapor pressure of the solution is 23.8 torr? (a) 19.4 g/mol (b) 37.6 g/mol (c) 38.8 g/mol (d) 40.0 g/mol (e) none of the above Answer: (d)

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Chapter 7: Making and Breaking of Bonds The following questions often assume that a truncated table of standard-state thermodynamic data is attached to the exam.

Specific Heat and Heat Capacity 1.

The same amount of heat is added to 20 gram each of iron and water. The temperature increases more in the iron than in the water. Which has the greater specific heat? Explain.

Answer: Water has the greater specific heat. Specific heat is the measure of the amount of heat required to raise the temperature by 1oC of 1 g of a substance. A greater specific heat indicates that more heat will be needed to give a particular increase in the temperature of a substance relative to a substance with a lower specific heat.

2.

Metal A at 40oC is brought into contact with Metal B at 20oC. Assuming you have the same masses of both metals and that Metal A has a greater specific heat than Metal B, what will the final temperature be? (a) 30oC (b) between 20oC and 30oC (c) between 30oC and 40oC (d) 40oC (e) insufficient information given to decide Answer: (c)

3.

When 2.50 g of KClO4 decomposes, 908 J of heat is evolved. If this heat is absorbed by 45.0 g of H2O, what will be the temperature change of the water? (a) 0.207oC decrease (b) 4.82oC increase (c) 6.41oC increase (d) 0.207oC increase (e) none of these

Answer: (b)

4.

How much energy in joules is required to heat 10.0 grams of gold from room temperature (20.0°C) to the temperature of boiling water (100.0°C) if the specific heat of gold is 0.129 g JoC ?

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(a) 1.29 J (b) 10.3 J (c) 103 J (d) 6,200 J (e) none of the above Answer: (c) 5.

A piece of copper metal weighing 145 grams was heated to 100°C and then dropped into 250 grams of water at 25°C. The copper metal cooled down and the water became warmer until both were at a temperature of 28.8°C. Calculate the amount of heat absorbed by the water. Assuming that the heat lost by the copper was absorbed by the water, what is the molar heat J capacity of copper metal? (CH2O = 75.376 molK ) J (a) between 0 and 10 molK J (b) between 10 and 20 molK J (c) between 20 and 30 molK J (d) between 30 and 40 molK J (e) more than 40 molK

Answer: (c) 6. How much heat in joules is produced by mixing 50.0 mL of 1.0 M HBr at 25.6°C with 50.0 mL of 1.0 M KOH at 25.6°C if this reaction produces 100 mL of a solution with a temperature of 32.3°C? (Assume the heat capacity of water is 4.18 g JoC and the density of these solutions is 1.00 cmg3 ) (a) 6.7 J (b) 670 J (c) 1400 J (d) 2800 J (e) 5600 J Answer: (d)

Enthalpies of Reaction 7.

The change in enthalpy, H, for the following balanced reaction is 33.2 kJ/molrxn. How much heat will be absorbed by the reaction of 1.00 g of N 2 with an excess of oxygen?

N2(g) + 2O2(g) → 2 NO2(g) (a) 1.19 kJ (b) 2.37 kJ

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(c) 33.2 kJ (d) 66.4 kJ (e) none of these Answer: (a)

8.

Which of the following is a correct statement for a chemical reaction that gives off energy? The reaction is: (a) exothermic with a +H (b) exothermic with a -H (c) endothermic with a +H (d) endothermic with a -H (e) exothermic with a H=0

Answer: (b)

9.

For the reaction MnO2(s) + CO(g) → MnO(s) + CO2(g) H = -151 kJ/ molrxn.

How many moles of MnO2 are required in order for 418 kJ of heat to be released? (a) 1.31 moles (b) 0.246 moles (c) 0.361 moles (d) 2.77 moles (e) none of these

Answer: (d)

10. How much heat is produced when 0.200 mole of H2(g) reacts with 0.300 mole of Cl2(g) if the enthalpy of reaction for the production of one mole of HCl(g) is -92.3 kJ/molrxn? (a) 18.5 kJ (b) 27.7 kJ (c) 36.9 kJ (d) 55.4 kJ (e) 92.3 kJ Answer: (c)

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11. How much heat in kilojoules is given off when 4.80 g of carbon are burned to produce carbon dioxide, if Hrxn° for the combustion of carbon to form CO2 is -394 kJ/molrxn? (a) 82.1 kJ (b) 158 kJ (c) 394 kJ (d) 985 kJ (e) 1.89 x 103 kJ Answer: (b) 12. Calculate Hrxn° in kJ/mol for the following reaction if 1.00 gram of magnesium reacts with excess fluorine to give off 46.22 kJ of heat.

Mg(s) + F2(g) → MgF2(s) (a) less than 300 kJ/molrxn (b) between 300 and 600 kJ/molrxn (c) between 600 and 900 kJ/molrxn (d) between 900 and 1200 kJ/molrxn (e) more than 1200 kJ/molrxn Answer: (d)

13. For which of the following reactions is the most heat evolved?

CH4(g) + 2 O2(g) → CO2(g) + 2 H2O(g)

(methane)

2 CH3CH2CH2CH3(g) + 13 O2(g) → 8 CO2(g) + 10 H2O(g)

(butane)

2 (CH3)2CHCH3(g) + 13 O2(g) → 8 CO2(g) + 10 H2O(g)

(isobutane)

Answer: methane = 802 kJ/molrxn, butane = 2656 kJ/molrxn, isobutane = 2649 kJ/molrxn. The most heat is evolved for butane.

14. Which of the following reactions is the most likely to give off heat? (a) Al(s) → Al(g) (b) Al(g) → Al3+(g) + 3 e– (c) Al–(g) → Al(g) + e– (d) 2 Al(s) + conc. HNO3(aq) → 2 Al3+(aq) + 3 H2(g) (e) all these reactions should give off energy.

Answer: (d)

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15. Which of the following reactions is the most likely to give off heat? (a) Mg(s) → Mg(g) (b) Mg(g) → Mg2+(g) + 2 e– (c) MgCl2(s) → Mg2+(g) + 2 Cl-(g) (d) Mg(g) + 2 Cl(g) → MgCl2 (e) all these reactions should give off energy Answer: (d) 16. Which of the following reactions is the most likely to be exothermic? (a) H2(g) → 2 H(g) (b) 2 F(g) → F2(g) (c) C(s) → C(g) (d) CCl4(g) → C(g) + 4 Cl(g) (e) HCl(g) → H(g) + Cl(g) Answer: (b)

17. Which of the following reactions is the most likely to be exothermic? (a) CaCO3(s) → CaO(s) + CO2(g) (b) Ca(g) → Ca2+(g) + 2 e– (c) Cl2(g) → 2 Cl(g) (d) Cl(g) + e– → Cl–(g) (e) CaCl2(g) → Ca2+(g) + 2 Cl–(g) Answer: (d)

18. Which of the following reactions is the most likely to be endothermic? (a) 2 Na(s) + 2 H2O(l) → 2 Na+(aq) + 2 OH–(aq) + H2(g) (b) 2 Mg(s) + O2(g) → 2 MgO(s) (c) 2 NaCl(s) → 2 Na(s) + Cl2(g) (d) Na+(g) + e– → Na(g) (e) Cl2(g) + 2e– → 2 Cl–(g) Answer: (c)

19. Which of the following reactions is the most likely to be endothermic? (a) Mg(s) → Mg(g) (b) Na+(g) + Cl–(g) → NaCl(s) (c) Na+(g) + e– → Na(g)

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(d) H+(aq) + OH–(aq) → H2O(l) (e) 2 H2(g) + O2(g) → 2 H2O(l) Answer: (a)

20. Use your understanding of the bonding in the reactants and products to explain why the following reaction is exothermic.

2 Mg(s) + O2(g) → 2 MgO(s) Answer: The ionic bond in MgO must be stronger than the mixture of metallic and covalent bonds in Mg metal and O2 gas, respectively.

Enthalpies of Atom Combination 21. At what temperature are standard-state enthalpy of reaction measurements most often made? (a) 0 K (b) 273.15 K (c) 0°C (d) 25°C (e) more than one of the above Answer: (d)

22. The bond strength between C and a halogen decreases down the halogen group in the periodic table. Based on this trend, predict which of the following gases has the largest (most negative) enthalpy of atom combination? (a) CCl4 (b) CBr4 (c) CF4 (d) CI4 (e) They all have the same enthalpy of atom combination Answer: (c)

23. When carbon is burned in air the following reaction takes place and releases heat.

C(s) + O2(g) → CO2(g) Which of the following is responsible for the heat produced? (a) breaking oxygen-oxygen bonds 7-6

(b) making carbon-oxygen bonds (c) breaking carbon-carbon bonds (d) both (a) and (c) (e) (a), (b) and (c) Answer: (b)

24. Which of the following is an exothermic process? (a) Li(g) → Li+ (g) + e– (b) NH3(g)→ N (g) + 3 H (g) (c) C(g) + 2 O(g) → CO2(g) (d) all of these (e) none of these Answer: (c)

25. Determine the change in enthalpy for the following reaction: SnO2(s) + 2 H2(g) → Sn(s) + 2 H2O(l) HacSnO2(s) = -1381.1 kJ/molrxn HacH2(g) = -435.3 kJ/molrxn

HacSn(s) = -302.1 kJ/molrxn HacH2O(l) = -970.3 kJ/molrxn

(a) -3088 kJ/molrxn (b) +544 kJ/molrxn (c) -544 kJ/molrxn (d) +9.0 kJ/molrxn (e) none of these Answer: (d) 26. Use enthalpies of atom combination to calculate Hrxn° for the following reaction.

C(s) + H2O(g) → CO(g) + H2(g) What would happen to the magnitude of Hrxn° if the reaction consumed liquid water instead of gaseous water?

Compound C(s) H2O(g) H2O(l) CO(g) H2(g)

Hac° (kJ/molrxn) -717 -926 -970 -1076 -435

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Answer: 132 kJ/molrxn; it would become more positive 27. Calculate Hrxn° for the following reaction from the data given below.

4 NH3(g) + 5 O2(g) → 4 NO(g) + 6 H2O(g) Compound NH3(g) O2(g) NO(g) H2O(g)

Hac° (kJ/molrxn) -1171.76 -498.34 -631.62 -926.29

(a) -105.5 kJ/molrxn (b) -905.5 kJ/molrxn (c) -1274.2 kJ/molrxn (d) -1996.2 kJ/molrxn (e) none of the above Answer: (b) 28. Calculate Hrxn° for the reaction: 3 NO2(g) + H2O(l) → 2 HNO3(aq) + NO(g) from the following enthalpy of atom combination data.

Compound NO2(g) H2O(l) HNO3(aq) NO(g)

Hac° (kJ/molrxn) -937.86 -970.30 -1645.22 -631.62

(a) less than -1000 kJ/molrxn (b) between -1000 and -750 kJ/molrxn (c) between -750 and -500 kJ/molrxn (d) between -500 and -250 kJ/molrxn (e) between -250 and 0 kJ/molrxn Answer: (e) 29. Both ethanol (CH3CH2OH) and methanol (CH3OH) have been considered as fuels for automobiles. Which is the better fuel, on a per gram basis, when burned with oxygen?

Compound CH3CH2OH(g) CH3OH(g)

Hac° (kJ/molrxn) -3223.53 -2037.11

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O2(g) H2O(g) CO2(g)

-498.34 -926.29 -1608.53

Answer: 21 kJ/g CH3OH versus 28 kJ/g CH3CH2OH 30.

The following reaction occurs when sucrose (cane sugar) is metabolized by the body.

C12H22O11(s) + 12 O2(g) → 12 CO2(g) + 11 H2O(l) Assume that Hrxn° for this reaction is -5645 kJ/molrxn. What is the value of Hac° for sucrose?

Compound O2(g) CO2(g) H2O(l)

Hac° (kJ/molrxn) -498.34 -1608.53 -970.30

Answer: -18,350 kJ/molrxn

31. What is the sign of the enthalpy of reaction for the reaction: P4O6(s) + 2 O2(g) + 6 H2O(g) → 4 H3PO4(s) assuming that all compounds are present in their most stable state at 25°C and 1 atm pressure?

Compound P4O6(s) O2(g) H2O(g) H3PO4(s)

Hac° (kJ/molrxn) -4393.7 -498.34 -926.29 -3243.3

(a) positive (b) negative (c) impossible to determine from the data Answer: (b)

32. What is the absolute value of the enthalpy of reaction in the previous question? (a) less than 1000 kJ/molrxn (b) between 1000 and 1250 kJ/molrxn (c) between 1250 and 1500 kJ/molrxn (d) between 1500 and 1750 kJ/molrxn

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(e) more than 1750 kJ/molrxn Answer: (e)

33. Calculate the average Si—Br bond strength for an individual Si—Br bond in SiBr4. The Hac for SiBr4(g) is -1272 kJ/molrxn. (a) 3.66 kJ/mol (b) 254 kJ/mol (c) 318 kJ/mol (d) 636 kJ/mol (e) 1272 kJ/mol Answer: (c)

34. The disposable lighters that many smokers carry use butane as a fuel.

Butane occurs as two isomers, n-butane H 3C CH3

CH3 CH

CH2

CH3 CH2

and isobutane

Hac n-butane = -5169.38 kJ/molrxn Hac isobutane = -5177.75 kJ/molrxn

CH3

.

The combustion reaction for both isomers of butane is: 2 C4H10(g) + 13 O2(g) → 8 CO2(g) + 10 H2O(g) (I) Which form of butane will release the most heat when it is burned by the above combustion reaction? Explain your reasoning.

(II) Which form of butane will have the strongest bonds? Explain your reasoning.

Answer: (I) Both forms of butane produce the exact same products.

Therefore the difference in the one that will release the most heat will be the form of butane that requires the least energy to break its bonds. Less energy is required to break the bonds in n-butane therefore it releases the most heat when burned.

(II) More energy is required to break the bonds in isobutane than in nbutane. Therefore, isobutane has the strongest bonds.

35. Use enthalpies of atom combination to calculate the change in enthalpy for the reaction: HCl(aq) + NaOH(aq) → NaCl(aq) + H2O(l)

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HacHCl(aq) = -506.49 kJ/molrxn HacNaCl(aq) = -636.27 kJ/molrxn

HacNaOH(aq) = -1044.25 kJ/molrxn HacH2O(l) = -970.3 kJ/molrxn

Is the reaction endothermic or exothermic?

Answer: -55.8 kJ/molrxn, exothermic

36.

The enthalpy of atom combination for CCl4 (g) is −1306.3 kJ/molrxn.

Calculate the average bond strength of the C−Cl bond. (a) 261.3 kJ/mol (b) 326.6 kJ/mol (c) 1306.3 kJ/mol (d) 3484.0 kJ/mol (e) 5225.2 kJ/mol

Answer: (b)

Bond Length and the Enthalpy of Atom Combination 37. Predict which of the following gases has the largest (most negative) enthalpy of atom combination? (a) CCl4 (b) CBr4 (c) CF4 (d) CI4 (e) They all have the same enthalpy of atom combination.

Answer: (c)

38. Which of the following has the shortest carbon to halogen bond length? (a) CCl4 (b) CBr4 (c) CF4 (d) CI4 (e) All the bonds are the same length.

Answer: (c)

39. Arrange the following in order of increasing enthalpy of atom combination and decreasing bond length.

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(a) Cl2(g) > Br2(g) > F2(g) > I2(g) (b) F2(g) > Cl2(g) > Br2(g) > I2(g) (c) I2(g) > Br2(g) > Cl2(g) > F2(g) (d) F2(g) > Cl2(g) > I2(g) > Br2(g) (e) Cl2(g) > Br2(g) > I2(g) > Br2(g) Answer: (c)

40. Arrange the following in order of increasing bond length and decreasing enthalpy of atom combination. (a) TeH4 > SiH4 > GeH4 > CH4 (b) TeH4 > GeH4 > SiH4 > CH4 (c) CH4 > SiH4 > GeH4 > TeH4 (d) GeH4 > CH4 > TeH4 > SiH4 (e) TeH4 > SiH4 > CH4 > GeH4 Answer: (c)

Hess’s Law 41. Use the following data 2 H2(g) + O2(g) → 2 H2O(g) 2 H2(g) + O2(g) → 2 H2O(l)

Hrxn° = -483.6 kJ/molrxn Hrxn° = -571.6 kJ/molrxn

to calculate Hrxn° for the reaction H2O(l) → H2O(g) (a) -527.6 kJ/molrxn (b) -44.0 kJ/molrxn (c) 44.0 kJ/molrxn (d) 527.6 kJ/molrxn (e) none of the above Answer: (c) 42. What is Hrxn° for the reaction: 2 CO(g) + O2(g) → 2 CO2(g) if: C(s) + 1/2 O2(g) → CO(g) C(s) + O2(g) → CO2(g) (a) -564 kJ/molrxn (b) -282 kJ/molrxn (c) 171 kJ/molrxn (d) 282 kJ/molrxn 7-12

Hrxn° = -111 kJ/molrxn Hrxn° = -393 kJ/molrxn

(e) 564 kJ/molrxn Answer: (a) 43. Calculate Hrxn° for the reaction: C(s) + 2 H2(g) → CH4(g) from the following data: C(s) + O2(g) → CO2(g) H2(g) + 1/2 O2(g) → H2O(l) CH4(g) + 2 O2(g) → CO2(g) + 2 H2O(l)

Hrxn° = -393.9 kJ/molrxn Hrxn° = -285.8 kJ/molrxn Hrxn° = -890.4 kJ/molrxn

(a) -1855.7 kJ/molrxn (b) -214.6 kJ/molrxn (c) -75.1 kJ/molrxn (d) 210.9 kJ/molrxn (e) 1569.9 kJ/molrxn Answer: (c)

44. Given the following data 3 H2(g) + N2(g) → 2 NH3(g) 2 H2(g) + O2(g) → 2 H2O(l)

Hrxn° = -92.4 kJ/molrxn Hrxn° = -571.7 kJ/molrxn

calculate Hrxn° for the following reaction: 4 NH3(g) + 3 O2(g) → 2 N2(g) + 6 H2O(l) (a) -1899.9 kJ/molrxn (b) -1715.1 kJ/molrxn (c) -1530.3 kJ/molrxn (d) -479.3 kJ/molrxn (e) 1530.3 kJ/molrxn Answer: (c)

45. Hydrogen peroxide is a good oxidizing agent because it is a good source of molecular oxygen: 2 H2O2(aq) → 2 H2O(l) + O2(g). Calculate Hrxn° for this reaction from the enthalpies of the following reactions.

2 H2(g) + O2(g) → 2 H2O(l) H2(g) + O2(g) → H2O2(aq) (a) -759.5 kJ/molrxn (b) -383.9 kJ/molrxn (c) -196.0 kJ/molrxn

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Hrxn° = -571.66 kJ/molrxn Hrxn° = -187.8 kJ/molrxn

(d) -98.0 kJ/molrxn (e) 196.0 kJ/molrxn Answer: (c) 46. Calculate Hrxn° for the reaction N2(g) + O2(g) → 2 NO(g) from the enthalpies of the following reactions.

N2(g) + 2 O2(g) → 2 NO2(g) 2 NO(g) + O2(g) → 2 NO2(g)

Hrxn° = 66.4 kJ/molrxn Hrxn° = -114.2 kJ/molrxn

(a) -180.6 kJ/molrxn (b) -47.8 kJ/molrxn (c) 47.8 kJ/molrxn (d) 180.6 kJ/molrxn (e) Hrxn° is impossible to calculate from the information given Answer: (d)

47. Calculate the heat of combustion of propane, C3H8, C3H8(g) + 5 O2(g) → 3 CO2(g) + 4 H2O(g) from the enthalpies of the following reactions.

3 C(s) + 4 H2(g) → C3H8(g) C(s) + O2(g) → CO2(g) H2(g) + 1/2 O2(g) → H2O(g)

Hrxn° = -103.85 kJ/molrxn Hrxn° = -393.51 kJ/molrxn Hrxn° = -241.83 kJ/molrxn

Answer: -2044.0 kJ/molrxn

48. Calculate the heat required to transform 18 g of ice at 0°C to water vapor at 100°C using some or all of the data given below.

H2O(s) → H2O(l) H2O(l) → H2O(g) 2 H2(g) + O2(g) → 2 H2O(g) H2O(liquid, 0°C) → H2O(liquid, 100°C) (a) 40.67 kJ (b) 46.70 kJ (c) 54.23 kJ (d) 296 kJ (e) none of these

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Hrxn° = 6.03 kJ/molrxn Hrxn° = 40.67 kJ/molrxn Hrxn° = -484 kJ/molrxn Hrxn° = 7.53 kJ/molrxn

Answer: (c)

Enthalpies of Formation (Note that some of these are the same problems as found in the section headed Enthalpies of Atom Combination. However in this section all problems are worked using Enthalpies of Formation.)

49. For which of the following substances is the enthalpy of formation, Hf, equal to zero? (a) H2(l) (b) H2O(g) (c) O3(g) (d) F2(g) (e) Na(g) Answer: (d)

50. Determine the change in enthalpy for the following reaction: SnO2(s) + 2 H2(g) → Sn(s) + 2 H2O(l) Hf SnO2(s) = -580.7 kJ/molrxn

Hf H2O(l) = -285.83 kJ/molrxn

(a) -3088 kJ/molrxn (b) + 544 kJ/molrxn (c) -544 kJ/molrxn (d) + 9.0 kJ/molrxn (e) none of these Answer: (d) 51. Use enthalpies of formation to calculate H for the following reaction.

C(s) + H2O(g) → CO(g) + H2(g) What would happen to the magnitude of H if the reaction consumed liquid water instead of gaseous water?

Compound Hf (kJ/molrxn) H2O(g) -241.81 H2O(l) -285.83 CO(g) -110.53 Answer: 131.28 kJ/molrxn; it would become more positive 52. Calculate H for the following reaction from the data given below.

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4 NH3(g) + 5 O2(g) → 4 NO(g) + 6 H2O(g) Compound NH3(g) NO(g) H2O(g)

Hf (kJ/molrxn) -46.11 90.25 -241.81

(a) -105.5 kJ/molrxn (b) -905.4 kJ/molrxn (c) -1274.2 kJ/molrxn (d) -1996.2 kJ/molrxn (e) none of the above Answer: (b) 53. Calculate H for the reaction 3 NO2(g) + H2O(l) → 2 HNO3(aq) + NO(g) from the following enthalpy of formation data.

Compound NO2(g) H2O(l) HNO3(aq) NO(g)

Hf (kJ/molrxn) 33.18 -285.83 -207.36 90.25

(a) less than -1000 kJ/molrxn (b) between -1000 and -750 kJ/molrxn (c) between -750 and -500 kJ/molrxn (d) between -500 and -250 kJ/molrxn (e) between -250 and 0 kJ/molrxn Answer: (e) 54. The following reaction occurs when sucrose (cane sugar) is metabolized by the body.

C12H22O11(s) + 12 O2(g) → 12 CO2(g) + 11 H2O(l) Given that H for this reaction is -5645 kJ/molrxn, what is the value of Hf for sucrose?

Compound Hf (kJ/molrxn) CO2(g) -393.51 H2O(l) -285.83 Answer: -2221 kJ/molrxn

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55. Both ethanol (CH3CH2OH) and methanol (CH3OH) have been considered as fuels for automobiles. Which is the better fuel, on a per gram basis, when burned with oxygen?

Compound CH3CH2OH(g) CH3OH(g) H2O(g) CO2(g)

Hf (kJ/molrxn) -235.10 -200.66 -241.82 -393.51

Answer: 21 kJ/g CH3OH versus 28 kJ/g CH3CH2OH

56. What is the sign of the enthalpy of reaction for the reaction P4O6(s) + 2 O2(g) + 6 H2O(g) → 4 H3PO4(s) assuming that all compounds are present in their most stable state at 25°C and 1 atm pressure?

Compound P4O6(s) H2O(g) H3PO4(s)

Hf (kJ/molrxn) -1640.1 -241.82 -1279.0

(a) positive (b) negative (c) impossible to determine from the data Answer: (b) 57. What is the magnitude of the enthalpy of reaction in the previous question? (a) less than 1000 kJ/molrxn (b) between 1000 and 1250 kJ/molrxn (c) between 1250 and 1500 kJ/molrxn (d) between 1500 and 1750 kJ/molrxn (e) more than 1750 kJ/molrxn Answer: (e)

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58. For which of the following substances is the enthalpy of formation, Hf, equal to zero? (a) H2(g) (b) H2O(g) (c) O3(g) (d) H(g) (e) Na(g) Answer: (d) 59. For which of the following substances is the enthalpy of formation, Hf, equal to zero? (a) H2O(l) (b) H2O(s) (c) Cl(g) (d) P4(s) (e) CO2(g) Answer: (d)

Integrated Problems 60. The reaction below shows two isomers of benzene, a compound with the formula C6H6. The structure of benzene was a point of great controversy in the late 1800’s and the two structures shown in the reaction were the chief contenders.

H H

H

C

H

C

C

C

C C

H

H

H

H

C

C

C

C

C

C

H

H

H

H = 240 kJ/molrxn H

2

1

Which of the following statements is correct? (a) Isomer 1 is more stable than isomer 2 (b) Isomer 2 is more stable than isomer 1 (c) Isomer 2 has overall stronger bonds than isomer 1 (d) Isomer 2 has stronger bonds and is more stable that isomer 1 (e) Isomer 1 has weaker bonds and is more stable than isomer 2

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Answer: (a) 61. Retailers purchase gasoline by weight so they get a constant amount of gasoline per dollar (as long as the price is constant). We buy gasoline by volume. The enthalpy of combustion of 1.00 L of gasoline is –3.539 x 104 kJ at 15.6oC when gasoline has a density of 0.73722 g/mL. What would the enthalpy of combustion of 1.00 L of gasoline be at 32.2 oC when gasoline has a density of 0.72210 g/mL? (a) –3.466 x 104 kJ (b) –3.539 x 104 kJ (c) –2.592 x 104 kJ (d) –2.373 x 104 kJ (e) None of the above Answer: (a)

62. Which of the following is not true for an exothermic reaction? (a) Heat flows from the reaction system to the surroundings. (b) The temperature of the surroundings rise. (c) The enthalpy change for the reaction is negative. (d) The products have a higher energy than the reactants. (e) The bonds in the products are stronger than the bonds in the reactants.

Answer: (d)

63. Write a balanced equation for the combustion of carbon disulfide to form carbon dioxide and sulfur dioxide, draw the Lewis structures of all reactants and products, and calculate Hrxn° for this reaction. _____ CS2(g) + _____ O2(g) → _____ CO2(g) + _____ SO2(g) S

Answer: 1, 3, 1, 2;

S

C

S, O

O, O

C

O

, O

O

;-1105 kJ/mol

Use the following standard enthalpies of atom combination to answer questions 64 through 68. Hac° H2(g) = -435.30 kJ/molrxn Hac° N2(g) = -945.41 kJ/molrxn Hac° O2(g) = -498.34 kJ/molrxn

Hac F2(g) = -157.98 kJ/molrxn Hac° C(s) = -716.68 kJ/molrxn Hac° NH3(g) = -1171.8 kJ/molrxn

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Hac° H2O(g) = -926.3 kJ/molrxn

Hac° HF(g) = -567.7 kJ/molrxn

64. What is Hrxn° for the following reaction?

3 N(g) → 3/2 N2(g) (a) +710 kJ/molrxn (b) -710 kJ/molrxn (c) +1418 kJ/molrxn (d) -1418 kJ/molrxn (e) -473 kJ/molrxn Answer: (d)

65. What is the bond energy of NN? (a) 237 kJ/mol (b) 158 kJ/mol (c) 473 kJ/mol (d) 946 kJ/mol (e) 1892 kJ/mol Answer: (d) 66. What is the Hrxn° for the following reaction as written?

4 NH3(g) + 3 O2(g) → 2 N2(g) + 6 H2O(g) (a) -1266 kJ/molrxn (b) -1184 kJ/molrxn (c) -196 kJ/molrxn (d) +767 kJ/molrxn (e) none of the above Answer: (a)

67. What is the bond energy, H2N—H, for one of the N-H bonds in NH3(g)? (a) 391 kJ/mol (b) 645 kJ/mol (c) 737 kJ/mol (d) 846 kJ/mol (e) 1173 kJ/moI Answer: (a)

7-20

68. How much heat is absorbed or evolved when just enough H2(g) reacts with just enough F2(g) to produce 0.200 mol of HF(g)? (a) 54 kJ of heat is absorbed (b) 54 kJ of heat is evolved (c) 108 kJ of heat is absorbed (d) 108 kJ of heat is evolved (e) 216 kJ of heat is absorbed Answer: (b)

69. Consider the following data for heats of atom combination of three substances in the liquid and gaseous states.

Compound

Hac° (liquid) Hac° (gas) H(gas) - H(liq)

(kJ/molrxn) pentane -6373 1-butanol -6060 2-butanone -5793

(kJ/molrxn) -6347 -6007 -5761

(kJ/molrxn) 26 53 32

The structures of these molecules are shown below.

H

H

H

H

H

H

C

C

C

C

C

H

H

H

H

H

pentane

H

H

H

H

H

H

C

C

C

C

H

H

H

H

1-butanol

OH

H

H

H

O

H

C

C

C

C

H

H

H

H

2-butanone

Explain the relative magnitudes of the differences between the enthalpies of atom combination of the liquid and gaseous states of these three compounds.

Answer: Intermolecular forces are largest in 1-butanol and smallest in pentane.

7-21

Chapter 8: Liquids and Solutions Structure of Gases, Liquids, and Solids 1.

Describe the difference between a solid, liquid, and gas on the macroscopic scale and on the atomic scale.

Answer: On the macroscopic level solids have a fixed volume and shape, liquids have a fixed volume but take the shape of their container and gases take the shape and volume of their container. On the atomic scale the particles which form solids are very close together and their motion is restricted to vibration. In liquids the particles are farther apart with some degree of motion allowed. In gases the particles are widely separated and are in constant random motion.

Intermolecular Forces 2.

Distinguish between the three intermolecular forces: dipole-dipole interaction, dispersion forces and hydrogen bonding.

Answer: Dipole/dipole interaction is found between polar molecules where the negative end of one molecule is attracted to the positive end of another.

Dispersion forces are found between all molecules. The motion of a molecule’s electrons causes the formation of temporary dipoles which can induce dipoles in neighboring molecules leading to a force of attraction.

Hydrogen bonding is a particularly strong dipole/dipole interaction that occurs when hydrogen is covalently bonded to O, N or F. This results in a large positive partial charge on hydrogen which is then strongly attracted to electronegative O, N, or F on neighboring atoms.

3.

The melting point of hexane at 1 atm. pressure, C6H14, is 179 K; the melting point of octane, C8H18, at the same pressure is predicted to be: (a) a lot lower than 179 K (b) somewhat lower than 179 K (c) equal to 179 K (d) somewhat higher than 179 K (e) a lot higher than 179 K

Answer: (d)

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4.

Which of the following molecules would exhibit a dipole-dipole intermolecular force of attraction? (a) N2 (b) PCl3 (c) CCl4 (d) I2 (e) none of these

Answer: (b)

5.

In which of the following substances do you expect to find the strongest intermolecular hydrogen bonding forces? (a) NH3 (b) H2S (c) H2O (d) CF4 (e) NaH

Answer: (c)

6.

Rare gases, such as xenon, can be frozen to form solids. The intermolecular forces responsible for maintaining the solid are best described as: (a) hydrogen bonds (b) covalent bonds (c) dipole-dipole forces (d) dipole-induced dipole forces (e) induced dipole-induced dipole forces

Answer: (e)

7.

Argon has a higher melting point (84 K) than neon (24 K) because there is: (a) stronger dipole-dipole interaction in argon (b) stronger hydrogen bonding in argon (c) stronger dispersion forces in argon (d) stronger ionic bonding in argon (e) none of these

Answer: (c)

8-2

The following three questions refer to the table below: Boiling Points of Three Classes of Organic Compounds MW (g/mole)

BP (C)

MW (g/mole)

BP (C)

Carboxylic Acid

MW (g/mole)

BP (C)

butane CH3(CH2)2CH3

58.1

−0.5

butanal CH3(CH2)2COH

72.2

75.7

butanoic acid 88.1 CH3(CH2)2COOH

164

Pentane CH3(CH2)3CH3

72.2

36.1

pentanal CH3(CH2)3COH

86.1

103

pentanoic acid 102.1 CH3(CH2)3COOH

186

Hexane CH3(CH2)4CH3

86.2

69.0

hexanal CH3(CH2)4COH

100.2

128

hexanoic acid CH3(CH2)4COOH

116.2

205

Heptane CH3(CH2)5CH3

100.2

98.4

heptanal CH3(CH2)5COH

114.2

153

heptanoic acid CH3(CH2)5COOH

130.2

223

Octane CH3(CH2)6CH3

114.2

126

octanal CH3(CH2)6COH

128.2

171

octanoic acid CH3(CH2)6COOH

144.2

239

Alkane

8.

Aldehyde

What intermolecular force is most responsible for the difference in boiling point between octane and pentane? (a) dispersion (b) dipole-dipole interaction (c) dipole-induced dipole interaction (d) hydrogen bonding (e) all of these

Answer: (a)

9.

What intermolecular force is most responsible for the difference in boiling point between hexanal and pentanoic acid? (a) dispersion (b) dipole-dipole (c) dipole-induced dipole (d) hydrogen bonding (e) all of these

Answer: (d)

10. Predict the boiling point of propanal, CH3CH2COH. (a) −40C (b) −10C (c) 48C (d) 78C (e) 110C Answer (c)

8-3

11. The following three compounds have similar molecular weights. Place them in order of increasing boiling point. formaldehyde: H2C=O

methanol: CH3OH

ethane: CH3CH3

(a) formaldehyde < ethane < methanol (b) ethane < methanol < formaldehyde (c) ethane < formaldehyde < methanol (d) methanol < formaldehyde < ethane (e) none of these Answer (c)

12. Which force must be overcome to melt solid pentane, C5H12? (a) metallic bonding (b) covalent bonding (c) ionic bonding (d) dispersion forces (e) all of these must be overcome to melt pentane Answer: (d)

Vapor Pressure and Boiling Point 13. Which of the following characteristics is associated with a low boiling point for a molecular substance? (a) high molecular weight (b) high vapor pressure (c) strong intermolecular forces (d) large dipole moment (e) strong cohesion forces Answer: (b)

14. Which of the following molecules has the highest boiling point? (a) CCl4 (b) CBr4 (c) CF4 (d) CH4 (e) CI4 Answer: (e)

8-4

15. Which of the following compounds has the lowest vapor pressure? (a) CCl4 (b) CBr4 (c) CF4 (d) CH4 (e) CI4 Answer: (e)

16. For which of the following is the boiling point incorrectly given? (a) CH3(CH2)2CH3, −0.5C (b) CH3(CH2)2CHO, 75C (c) CH3(CH2)3OH, 100C (d) CH3(CH2)2COOH, 95C (e) CH3(CH2)3CHO, 103C Answer: (d)

17.

Which compound would you expect to have the highest vapor pressure at 25°C? (a) C7H16 (BP = 98.4°C) (b) C8H18 (BP = 125.7°C) (c) C9H20 (BP = 150.8°C) (d) C10H22 (BP = 174.3°C) (e) It is impossible to tell which compound would have the highest vapor pressure because there is no relationship between vapor pressure and the boiling point of a compound.

Answer: (a)

18.

A liquid is at its equilibrium vapor pressure at 20oC. The following changes are then made: (1) The liquid is transferred into a container with a wider opening.

(2) More liquid is added to the container.

(3) The temperature is increased to 30oC. The vapor pressure is then observed to increase. Which of these changes contributed to the increase in the vapor pressure? (a) (1) only

8-5

(b) (2) only (c) (3) only (d) (1) and (2), but not (3) (e) (2) and (3), but not (1) Answer: (c) 19. The boiling point of HF is lower than that of water despite the fact that the hydrogen bonds in HF are at least as strong as those in water because: (a) water is more polar than HF (b) water has the higher melting point (c) water is less dense (d) each water molecule forms more hydrogen bonds (e) the vapor pressure of HF is higher than that of water Answer: (d)

Phase Diagrams 20. In most cases, reading from left to right across a phase diagram, what is the order in which you would find the three states of matter? (a) gas < liquid < solid (b) gas < solid < liquid (c) liquid < gas < solid (d) liquid < solid < gas (e) solid < liquid < gas Answer: (e)

21. Gases are most likely to be found in a phase diagram under conditions of (a) high temperature and high pressure. (b) high temperature and low pressure. (c) low temperature and high pressure. (d) low temperature and low pressure. (e) moderate temperature and pressure.

Answer: (b) 22. Which statement isn’t consistent with the information contained in the phase diagram of a compound? (a) Liquids can be made to boil by raising the pressure at constant temperature.

8-6

(b) Solids can be made to melt by raising the temperature at constant pressure. (c) Gases are most likely to be found at high temperatures and low pressures. (d) The boiling point of a compound is the temperature at which the vapor pressure of the compound is equal to atmospheric pressure. (e) All of these statements are consistent with the information in the phase diagram of a compound.

Answer: (a)

The following three questions refer to the phase diagram shown below.

23. The regions labeled A,B, and C are (a) the gas, liquid and solid phases, respectively. (b) the solid and liquid phases and the triple point, respectively. (c) the gas and liquid phases and the critical point, respectively. (d) the solid, liquid and gas phases, respectively. (e) pressure, temperature and moles, respectively.

Answer: (d)

24. If the temperature and pressure are changed from the values at point B to the values at point C, the substance will change its phase from (a) liquid to gas. (b) liquid to solid (c) gas to solid. (d) solid to liquid. (e) gas to liquid.

8-7

Answer: (a) 25. For this substance, as the pressure increases, the melting point will (a) increase. (b) decrease. (c) stay the same. (d) could either increase or decrease. (e) the phase diagram gives no information about the melting point.

Answer: (a)

Solutions 26. The best explanation for the fact that water and ethyl alcohol are completely miscible is that (a) both compounds are nonpolar. (b) water is polar and ethyl alcohol is nonpolar. (c) water has an unusually large dielectric constant. (d) ethyl alcohol dissociates in water. (e) hydrogen bonding takes place between the two substances.

Answer: (e)

27. Which of the following will be most soluble in a nonpolar solvent such as carbon tetrachloride, CCl4? (a) KI (b) H2O (c) NH3 (d) CBr4 (e) HF Answer: (d)

28. Which alcohol should be most soluble in a nonpolar solvent such as benzene, C6H6? (a) CH3OH (b) CH3CH2OH (c) CH3CH2CH2OH (d) CH3CH2CH2CH2OH (e) CH3CH2CH2CH2CH2OH

8-8

Answer: (e)

29. Carboxylic acids with the general formula CH3(CH2)nCO2H have a nonpolar CH3–CH2— tail and a polar —CO2H head. What effect would increasing the value of "n" have on the solubility of carboxylic acids? (a) The solubility would increase in both water and in nonpolar solvents such as CCl4. (b) The solubility would decrease in both water and in nonpolar solvents such as CCl4. (c) The solubility would increase in water but decrease in CCl4. (d) The solubility would decrease in water but increase in CCl4. (e) The solubility of carboxylic acids in both water and CCl4 would remain the same.

Answer: (d)

30. Which of the following statements is correct for the solubilities of alcohols with the formula CH3(CH2)nOH? (a) The solubility in water is minimal because the compound is organic. (b) The solubility in water increases as n increases. (c) The solubility in water decreases as n increases. (d) The solubility in water is independent of n. (e) The solubility in water depends only on the extent of hydrogen bonding.

Answer: (c)

31. An alcohol is a compound with the formula CH3(CH2)nOH, in which the CH2(CH2)n group is nonpolar and the OH group is polar. As "n" increases for different alcohols, we would expect the solubilities of the alcohol to (a) increase in both polar and nonpolar solvents. (b) decrease in polar solvents and increase in nonpolar solvents. (c) decrease in both polar and nonpolar solvents. (d) stay the same in polar solvents and decrease in nonpolar solvents. (e) increase in polar solvents and stay the same in nonpolar solvents.

Answer: (b)

Solids Dissolve in Water 32. Aqueous 0.2M solutions of K2SO4 and Pb(NO3)2 are mixed together. Use the solubility rules to predict the identity of the precipitate.

8-9

(a) KNO3 (b) K2Pb (c) PbSO4 (d) (NO3)2SO4 (e) No precipitate will form.

Answer: (c) 33. Aqueous 0.2M solutions of Sr(OH)2 and FeCl3 are mixed together. Use the solubility rules to predict the identity of the precipitate. (a) Cl(OH)3 (b) SrCl2 (c) Sr3Fe2 (d) Fe(OH)3 (e) No precipitate will form.

Answer: (d) 34. Determine the spectator ions in the following reaction.

MgS(aq) + CuSO4(aq) → CuS(s) + MgSO4(aq) (a) Cu2+ and Mg2+ (b) S2– and SO42– (c) Cu2+ and S2– (d) Mg2+ and SO42– (e) There are no spectator ions in this reaction.

Answer: (d)

35. Based on the solubility rules, which of the following compounds should be most soluble? (a) AgS (b) PbCrO4 (c) BaCO3 (d) (NH4)2CO3 (e) FeS Answer: (d)

36. Based on the solubility rules, which of the following compounds should be most insoluble? (a) KCl 8-10

(b) NaNO3 (c) Cu(OH)2 (d) Cu(NO3)2 (e) KOH Answer: (c)

Colligative Properties 37. Which of the following is true? (a) Molarity (M) is calculated by dividing the number of moles of solute by the number of liters of solvent. (b) Molality (m) is calculated by dividing the number of moles of solute by the number of liters of solvent. (c) Mole fraction is calculated by dividing the number of moles of solvent by the number of moles of solute. (d) Mole fraction is calculated by dividing the number of moles of solute by the number of moles of solvent. (e) Statements (a) through (d) are all false.

Answer: (e)

38.

Which of the following concentration units would change when the temperature of the solution changes? (Hint: remember that liquids expand and contract as the temperature changes.)

(I) molarity (II) molality (III) mole fraction of solute (IV) mole fraction of solvent (a) I (b) I and II (c) I, II and III (d) all of the above (e) none of the above

Answer: (a)

39.

Which of the following correctly describes the behavior of the vapor pressure of liquids? (a) It remains constant as the temperature of the liquid increases. (b) It remains constant when a solute is dissolved in the liquid. (c) It increases when a solute is dissolved in the liquid.

8-11

(d) If the vapor pressure of CH3OCH3 is much higher than CH3CH2OH at room temperature, CH3OCH3 should boil at a lower temperature than CH3CH2OH. (e) All of the above are correct.

Answer: (d)

40. What is the mole fraction of CCl4 (MW = 154 g/mol) in a solution prepared by dissolving 32 grams of CCl4 in 75 grams of C6H6 (MW = 78.1 g/mol) ? (a) 0.18 (b) 0.22 (c) 0.30 (d) 0.48 (e) 0.82 Answer: (a) 41. For a solution of 2 g of solute (MW = 200 g/mol) dissolved in 200 g of solvent (MW = 20 g/mol), the mole ratio of solute to solvent is: (a) 1 (b) 10–1 (c) 10–2 (d) 10–3 (e) 10–4 Answer: (d)

42. Which of the following liquid systems would be expected to have the lowest vapor pressure at room temperature? (a) a 1.00 m solution of ethanol in water (b) a 1.00 m solution of sucrose in water (c) a 1.00 m solution of acetone (CH3C(O)CH3) in water (d) a 1.00 m solution of NaCl in water (e) pure water Answer: (d)

43. Raoult's law predicts that the vapor pressure of solvent over a solution of a nonvolatile solute (a) increases with increasing mole fraction of solvent (b) increases with increasing mole fraction of solute (c) is independent of the solute concentration

8-12

(d) decreases with increasing temperature. (e) decreases with decreasing room pressure.

Answer: (a)

44. When a nonvolatile solute is dissolved in a volatile solvent, the vapor pressure of the solvent over the solution (a) increases as the mole fraction of solute increases. (b) decreases as the mole fraction of solute increases. (c) is independent of the mole fraction of solute. (d) changes in an unpredictable way with changes in the mole fraction of the solute. (e) decreases as the temperature increases.

Answer: (b)

45. When a nonvolatile solute is dissolved in a volatile solvent, the vapor pressure, freezing point, and boiling point (a) are all higher for the solution than for the pure solvent. (b) are all lower for the solution than for the pure solvent. (c) change differently, with vapor pressure increasing and boiling and freezing points decreasing. (d) change differently, with boiling point increasing and vapor pressure and freezing point decreasing. (e) do not change because the solute is nonvolatile.

Answer: (d)

46. A 10.0% by weight solution of NaCl (78.5 g/mol) in water has a molality of: (a) 0.127 m (b) 1.42 m (c) 1.50 m (d) 7.85 m (e) impossible to determine unless you know the density of the solution Answer: (b)

47. 0.400 mole of CCl4 and 0.600 mole of CHCl3 are mixed at 43°C. The vapor pressure of pure CCl4 at this temperature is 0.354 atm and the vapor pressure of pure CHCl3 at this temperature is 0.526 atm. What is the vapor pressure of the solution?

8-13

(a) 0.285 atm (b) 0.423 atm (c) 0.457 atm (d) 0.526 atm (e) 0.880 atm Answer: (c)

48. What is the molality of a bromine solution made by dissolving 39.95 g of Br2 in 5.00 x 102 g of CCl4? (a) 0.0800 (b) 0.250 (c) 0.500 (d) 1.00 (e) 2.00 Answer: (c) 49. KCl is a strong electrolyte while sucrose (table sugar) is not. Which of the following lists 0.10 m KCl, 0.10 m sucrose, and pure water in order of increasing boiling point? (a) KCl < sucrose < water (b) water < KCl = sucrose (c) KCl = sucrose < water (d) water < sucrose < KCl (e) sucrose < KCl < water Answer: (d)

50. Which of the following solutions has the highest boiling point? (a) 0.1 m NaCl (b) 0.1 m sucrose (c) 0.1 m Na2CO3 (d) 0.1 m NH3 (e) 0.1 m NH4Cl Answer: (c)

51. Pure ethanol has a boiling point of 78.0 °C. What is the boiling point of a solution that contains 12.7 grams of KI (MW = 166 g/mol) dissolved in 125 grams of ethanol (MW = 46.0 g/mol)? Assume that KI is a strong electrolyte in ethanol. (Ethanol: kb = 1.22 °C/m)

8-14

(a) 76.5°C (b) 77.2°C (c) 78.7°C (d) 79.5°C (e) 81.0°C Answer: (d)

52. When a 2.15 gram sample of a nonelectrolyte is dissolved in 105 grams of water, the freezing point of the solution is -0.62°C. If the molal freezing point depression constant for water is 1.86°C/m, what is the molecular weight of the nonelectrolyte?

Answer: 61 g/mol 53. The molal freezing point constant for water is 1.86 °C/m and the freezing point for pure water is 0°C. When 1.50 g of a compound is added to 47.8 g of water, the freezing point of the solution is found to change by 0.174°C. Calculate the molecular weight of the solute.

Answer: 335 g/mol 54. Pure benzene melts at 5.5°C. A solution of 1.25 g of CCl4 in 1.00 x 102 g of benzene would have what freezing point? (Benzene: kb = 5.1°C/m) Answer: 5.09°C

55. Given dilute, equimolar solutions of NaCl and ZnCl2 in water, we would expect the freezing points for the solutions relative to pure water to (a) change by the same amount. (b) change by a ratio of 2:3, with the change for NaCl being 2/3rds that for ZnCl2. (c) change by a ratio of 1:3, with the change for NaCl being 1/3rd that for ZnCl2. (d) change by a ratio of 2:3, with the change for ZnCl2 being 2/3rds that of NaCl. (e) change by a ratio of 1:3, with the change for ZnCl2 being 1/3rd that for NaCl. Answer: (b)

8-15

56. If freezing point depression was used to determine the molecular weight of an acid dissolved in water, and the dissociation of the acid wasn’t taken into account, the determined value of the molecular weight of the acid would be: (a) accurate (b) too high (c) too low (d) inaccurate, but in a direction that cannot be predicted Answer: (c)

57. Which of the following equations would be appropriate for calculating the change in the freezing point of an aqueous NaHSO4 solution? (NaHSO4 dissociates in solution in a two-step process.)

NaHSO4(s) → Na+(aq) + HSO4–(aq) HSO4–(aq) → H+(aq) + SO42-(aq)

(100% dissociation) (10% dissociation)

(a) TFP = (1 + 0.10)kfm (b) TFP = (0.10)kfm (c) TFP = [1 + 2( 0.10)]kfm (d) TFP = (0.20)kfm (e) TFP = (2 + 0.10)kfm Answer: (e)

58. A 0.010 molal solution of ammonia in water is 4.1% ionized.

NH3(aq) + H2O(l) → NH4+(aq) + OH–(aq) Given that pure water freezes at 0°C and the kF for water is 1.86°C/m, which of the following expressions gives the freezing point depression of the aqueous ammonia solution? (a) (b) (c) (d) (e)

TFP = -(1.86 °C/m)(0.01 m) TFP = -(1.86 °C/m)(0.01 m)(1.041) TFP = -(1.86 °C/m)(0.01 m)(2.082) TFP = -(1.86 °C/m)(0.01 m)(4.1) TFP = -(1.86 °C/m)(0.01 m)(8.2)

Answer: (b)

8-16

Chapter 9: Solids Solids 1.

Which of the following types of solids represents the most ordered arrangement of repeating units? (a) polycrystalline (b) crystalline (c) hydrophobic (d) amorphous (e) none of these

Answer: (b)

2.

Diamond is an example of: (a) a molecular solid (b) a network covalent solid (c) an ionic solid (d) a solid solution (e) a metallic solid

Answer: (b)

3.

SiO2 is a network covalent solid. Which of the following is most likely to be true of SiO2? (a) SiO2 would have a low melting point and would conduct electricity when melted. (b) SiO2 would conduct electricity as a solid and its melting point cannot be predicted. (c) SiO2 would have a high melting point and would conduct electricity when melted. (d) SiO2 would have a high melting point and wouldn’t conduct electricity when melted. (e) SiO2 would have a low melting point and wouldn’t conduct electricity when melted.

Answer: (d)

9-1

4.

Which of the following types of solids is most likely to be soft, have a low melting point, and not conduct electricity when melted? (a) ionic (b) molecular (c) network covalent (d) metallic (e) none of the above

Answer: (b)

5.

Which solid is held together by an extended network of covalent bonds? (a) sodium chloride (b) gold (c) calcium carbonate (d) diamond (e) dry ice (solid CO2)

Answer: (d)

6.

Which force makes the most important contribution to the lattice energy of solid CO2? (a) metallic bonding (b) ionic bonding (c) hydrogen bonding (d) dispersion forces (e) all of the above

Answer: (d)

7. Which of the following compounds should have the greatest lattice energy? (a) NaCl (b) NaI (c) KI (d) MgS (e) MgO Answer: (e)

8. Diamond (C(s)) is a network covalent solid. Which of the following is most likely to be true of diamond?

9-2

(a) Diamond would have a low melting point and would conduct electricity when melted. (b) Diamond would conduct electricity as a solid and its melting point cannot be predicted. (c) Diamond would have a high melting point and would conduct electricity when melted. (d) Diamond would have a high melting point and wouldn’t conduct electricity when melted. (e) Diamond would have a low melting point and wouldn’t conduct electricity when melted.

Answer: (d)

9.

A compound that is a poor conductor of electricity when solid, but a very good conductor when molten is most likely to fit into which category? (a) molecular solid (b) covalent solid (c) ionic solid (d) metallic solid (e) any of the above

Answer: (c)

10. Arrange the following ionic compounds in order of increasing lattice energy (energy to break the solid into its gaseous ions).

MgF2, SrF2, BeF2 (a) MgF2 < SrF2 < BeF2 (b) SrF2 < MgF2 < BeF2 (c) MgF2 < BeF2 < SrF2 (d) SrF2 < BeF2 < MgF2 (e) BeF2 < MgF2 < SrF2 Answer: (b) Use the bond-type triangle for questions 11-17.

9-3

3.6

CsF

3.4 BaF 2 3.2

CaF 2

3.0

Cs 2O

MgF 2 2.8 BaO 2.6

CaO

2.4 MgO

SiF4

CsCl CsBr Cs 2S CsI CsH

CsCd 13

Be 13Ca

Cs Rb

Na Ca Ba Sr Li

1.4 Mg Ti

SO 2

1.0

N 2F2

NP

Be 13 Ba

1.2

1.2

GaI 3 SiC

AlAs AlSb

1.6

HCl

GaAs SiH 4

0.4 CCl 4

SCl 2

Br 2 2.0

2.2

0.6

OF 2

NO

ICl

CH4

1.8

0.8

GaBr 3

Cd 3 Mg 1.0

CO 2

BN

ZnS

CIF

SeO 2

Hg 2Cl2

BeC 2

1.6 1.4

P4O 7

AlCl 3

MgH 2

1.8 SF4

B 2O 3 AlN

MgBr 2

AIH 3

MgZn 2

0.8

SiO 2

Be 3N 2

BaZn 5

CdLi

HF IF5

HgO

Li 3 As CaH 2 BaSi 2

2.0

ZnO

Mg 3N 2

BaI 2 Cs 3Sb

BeO

BaCl 2 BaBr 2 BaS BaC 2

2.2

BF 3

Ca 3N 2

2.4

2.6

Cd Al In Ga Si Ge B As H Se C Zn Cu Ni Ag P I Be Sn Sb Te Mn Hg

S

2.8 Br

N2

3.0 Cl

0.2

NCl 3

CI 2

O2 3.2

N

3.4

3.6 O

3.8

4.0

F2 4.2 F

Average EN Copyright 1999 by John Wiley and Sons Inc. All rights reserved.

11. Which of the following compounds will have the most ionic character? (a) AlSb (b) AlN (c) AlLi (d) AlP (e) AlAs Answer: (b)

9-4

EN

12. Which of the following compounds will have the most covalent character? (a) NaCl (b) SrCl2 (c) PCl3 (d) AlCl3 (e) SnCl4 Answer: (c)

13. Which of the following compounds will have the most metallic character? (a) SnBr4 (b) SnO2 (c) SnP (d) SnS (e) SnCl4 Answer: (c)

14. Predict which compound has a high melting point and is a conductor in the molten and aqueous state. (a) AlSb (b) AlN (c) AlLi (d) AlP (e) AlAs Answer: (b)

15. Predict which compound is an electrical insulator and has a relatively low melting point. (a) NaIn (b) SrCl2 (c) PCl3 (d) AlCl3 (e) CaH2 Answer: (c)

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16. Predict which solid is the best electrical conductor. (a) SnBr4 (b) SnO2 (c) SnP (d) SnS (e) SnCl4 Answer: (c)

17. Predict which compound is an insulating ceramic material. (a) MgZn2 (b) Li3P (c) Li2O (d) LiCl (e) GaAs Answer: (e)

Coordination Numbers and the Structure of Metals 18. What is the coordination number of a hexagonal closest packed structure? (a) 2 (b) 4 (c) 6 (d) 8 (e) 12 Answer: (e)

19. Which of the following structures has a coordination number of six? (a) simple cubic (b) body-centered cubic (c) cubic closest packed (d) hexagonal closest packed (e) none of the above Answer: (a)

20. Which of the following correctly describes a body-centered cubic structure? (a) Coordination Number = 6, 1 atom per unit cell

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(b) Coordination Number = 6, 2 atoms per unit cell (c) Coordination Number = 6, 4 atoms per unit cell (d) Coordination Number = 8, 1 atom per unit cell (e) Coordination Number = 8, 2 atoms per unit cell Answer: (e)

Unit Cells 21. What is the ratio of the length of the body diagonal of a cube to the face diagonal of the cube? (a) 0.817 (b) 1.00 (c) 1.22 (d) 1.41 (e) 1.73 Answer: (c)

22. What is the relationship between the unit cell edge length (a) and the radius of a metal atom (r) when the metal crystallizes in a face-centered cubic unit cell? (a) r = 2 a (b) r = ( 2 a)/4 (c) r = ( 2 a)/2 (d) r = ( 3 a)/4 (e) none of the above Answer: (b)

23. Calcium titanate crystallizes in a cubic unit cell with titanium atoms at the corners of the cell, oxygen atoms in the middle of the faces of the cell, and a calcium atom in the center of the cell. What is the formula of calcium titanate? (a) CaTiO (b) CaTiO3 (c) CaTiO4 (d) Ca2TiO4 (e) CaTi8O2 Answer: (b)

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24. Calculate the number of chloride and ammonium ions per unit cell if NH4Cl is a simple cubic unit cell of NH4+ ions with a Cl– ion in the center of the unit cell. (a) 1 NH4+ ion and 1 Cl– ion (b) 2 NH4+ ions and 2 Cl– ions (c) 4 NH4+ ions and 4 Cl– ions (d) 8 NH4+ ions and 8 Cl– ions (e) none of the above Answer: (a)

Determining the Unit Cell of a Crystal 25. What is the density in g/cm3 of sodium if this metal crystallizes in a bodycentered cubic unit cell with a cell edge of 0.429 nm? (a) 0.484 g/cm3 (b) 0.582 g/cm3 (c) 0.967 g/cm3 (d) 1.94 g/cm3 (e) 9.67 x 10-25 g/cm3 Answer: (c) 26. What is the approximate density in g/cm3 of titanium metal if the atomic radius of a titanium atom is 0.145 nm and the atoms form a body-centered cubic unit cell? (a) 0.1 g/cm3 (b) 1 g/cm3 (c) 5 g/cm3 (d) 25 g/cm3 (e) 30 g/cm3 Answer: (c) 27. Tungsten crystallizes in a cubic unit cell with a cell edge of 0.3981 nm. The density of tungsten is 19.35 g/cm3. Calculate the number of atoms per unit cell. (a) 1 (b) 2 (c) 3 (d) 4 (e) 8 9-8

Answer: (d)

28. Iron crystallizes in a cubic unit cell with a cell edge of 0.2866 nm. The density of iron is 7.875 g/cm3. Calculate the number of atoms per unit cell. (a) 1 (b) 2 (c) 3 (d) 4 (e) 8 Answer: (b)

Calculating the Size of an Atom or Ion 29. Calculate the metallic radius in nanometers of the Ag atom if silver crystallizes in a face-centered cubic unit cell with a cell edge of 0.4086 nm.

Answer: 0.1445 nm 30. Cesium iodide consists of a simple cubic lattice of l– ions with Cs+ ions in the center of the cubes. If the cell edge length is 0.445 nm, what is the Cs— I interatomic distance in nanometers? (a) 0.193 nm (b) 0.314 nm (c) 0.385 nm (d) 0.629 nm (e) 0.770 nm Answer: (c) 31. In KF, the K+ and F– ions are almost exactly the same size: 0.134 nm.

Which would be larger: a neutral potassium atom or a neutral fluorine atom?

Answer: neutral K atom

Integrated Problem

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32. Which of the following would be the correct order from lowest melting point to highest for the following substances? (a) NaF < C3H8 < CH3CH2OH (b) C3H8 < NaF < CH3CH2OH (c) C3H8 < CH3CH2OH < NaF (d) CH3CH2OH < C3H8 < NaF (e) none of these Answer: (c)

33. When diamond, NaCl, and water are boiled, the bonds being broken can best be described respectively as: (a) covalent, ionic, covalent (b) ionic, covalent, hydrogen bonds (c) covalent, ionic, hydrogen bonds (d) hydrogen bonds, ionic, covalent (e) none of the above Answer: (c)

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Chapter 10: An Introduction to Kinetics and Equilibrium The Rate of a Chemical Reaction 1.

The following reaction was carried out: Ni(CO)4 + PPh3(g)→ PPh3Ni(CO)3 + CO Over 200 seconds, the concentration of Ni(CO)4 dropped from 10 M to 2.5 M. What is the average rate of the reaction in M/s over this time period? (a) –0.038 M/s (b) 0.050 M/s (c) –0.050 M/s (d) 1.5 x 103 M/s (e) 0.038 M/s

Answer: (e)

2.

Describe the relationship between the rate constants kf and kr for the following one step reaction at equilibrium.

Z(g) + X(g)

kf

Y(g)

Kc= 110−3

kr

(a) kf = kr (b) kf < kr (c) kf > kr (d) kf = kr = 0 (e) none of these Answer: (b)

The Collision Theory of Gas Phase Reactions 3.

For the reaction: CO(g) + Cl2(g)

COCl2(g)

kf is 1x 107 M–1 s–1 and kr = 2 x 102 s–1. What is the equilibrium constant for the reaction? (a) 2 x 10–5 (b) 5 (c) 5 x 104 (d) 2 x 109 (e) none of these

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Answer: (c)

Concentration

Problems 4 through 8 go with the figure below.

Time

4.

Which of the following reactions could be described by the above diagram? (a) 2 N2O2(g) 4 NO2(g) + O2(g) (b) H2(g) + CO2(g) CO(g) + H2O(g) (c) PCl5(g) PCl3(g) + Cl2(g) (d) 2 NO(g) + Br2(g) 2 NOBr(g) (e) 2 O3(g) 3 O2(g)

Answer: (d)

5.

What happens with respect to time to the products of this reaction in the kinetic region of the graph above? (a) the products increase (b) the products decrease (c) the products remain constant (d) cannot be determined from the graph

Answer: (a)

6.

What happens to the rate of the reaction as time passes in the equilibrium region of the graph above? (a) the rate increases (b) the rate decreases (c) the rate remains constant (d) cannot be determined from the graph

Answer: (c)

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7.

From the graph above, which of the following is true when equilibrium has been reached? (a) The concentrations of the reactants are larger than the concentration of the products. (b) The concentrations of the products are larger than the concentration of the reactants. (c) The concentrations of the reactants and products are equal. (d) This cannot be determined from the graph.

Answer: (a)

8.

Which of the following will be true concerning the Kc of the reaction described by the graph above? (a) Kc < 1 (b) Kc > 1 (c) Kc = 1 (d) Kc cannot be determined from the graph

Answer: (a)

9.

For the following reaction: 2 NOCl(g)

2 NO2(g) + Cl2(g)

Kc = 5.6 x 10–6 (at 400 K)

If (NOCl) = 0.0222 M, (Cl2) = 0.0222 M, and (NO2) = 0.989 M at some moment in time, which of the following would be true? (a) The rate of the forward reaction would be exactly equal to the rate of the reverse reaction. (b) Neither the forward nor the reverse reaction would occur because the reaction would be at equilibrium. (c) The rate of the forward reaction would be faster than the reverse reaction. (d) The rate of the reverse reaction would be faster than the forward reaction. (e) There is no relationship between these concentrations and the rate of either the forward or the reverse reaction.

Answer: (d)

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10. Which one of the following graphs best represents the relationship between the concentration of reactants and products with respect to time for the following chemical reaction?

2 NO2(g) (b)

Concentration

NO O2

NO 2

Concentration

(a)

2 NO(g) + O2(g)

NO O2

NO 2

Time

Time

(d)

Concentration

NO, O2

NO 2

NO, O2

Time

Time

(e) NO 2

Concentration

NO 2

Concentration

(c)

O2 NO

Time

Answer (b)

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Use the following graph for answering questions 11 & 12.

Concentration (mol/L)

0.8 0.7 0.6

A

0.5 0.4

B

0.3

C

0.2 0.1

Time

11. Using the above graph determine the balanced chemical equation. (a) 3 A + B 2C (b) 5 A + 3 B 2C (c) 2 C + 3 B 5A (d) A + B C (e) none of these Answer: (a)

12. Which of the following will be true of the equilibrium constant, Kc, for the reaction described in the above graph? (a) Kc  1 (b) Kc  0 (c) Kc < 1 (d) Kc > 1 (e) cannot be determined from graph Answer: (a)

13. At equilibrium, ______________. (a) all chemical processes have ceased. (b) the rate of the forward reaction equals that of the reverse. (c) the rate constant for the forward reaction equals that of the reverse. (d) both the rate of the forward reaction equals that of the reverse and the rate constant for the forward reaction equals that of the reverse. (e) none of the above Answer: (b)

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14. Which of the following graphs best represents the relationship between the concentration of reactants and products with respect to time for the following chemical reaction?

2 H2(g) + O2(g)

2 H2O(g)

(a)

(b)

(c)

(d)

(e)

Answer: (e)

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Equilibrium Constant Expressions 15. Assume that the equilibrium constants for the following reactions are known.

2 SO2(g) + O2(g) 2 CO(g) + O2(g)

2 SO3(g) 2 CO2(g)

What is the equilibrium constant for the following reaction?

2 SO2(g) + 2 CO2(g)

2 SO3(g) + 2 CO(g)

(a) K = K1 x K2 (b) K = K1/K2 (c) K = K2/K1 (d) K = K1 + K2 (e) K = K1 – K2 Answer: (b)

16. Which is the correct equilibrium constant expression for the reaction: 2 SO2(g) + O2(g)

(a) Kc =

(b) Kc = (c) Kc =

(d) Kc =

2 SO3(g)

[SO2]2 [O2] [SO3]2 [SO3]2 [SO2]2 [O2] [SO3]2 [SO2]2 + [O2] [SO3]2 [SO2]2

(e) None of the above Answer: (b)

10-7

K1 K2

17. Assume that the equilibrium constant for the following reaction is known.

2 NO2(g)

K1 = 7.4 x 10–16 (at 25°C)

2 NO(g) + O2(g)

What is the correct value of the equilibrium constant for the opposite reaction?

2 NO(g) + O2(g)

2 NO2(g)

K2 = ?

(a) K2 = K1 = 7.4 x 10–16 (b) K2 = 1/K1 = 1.4 x 1015 (c) K2 = K1(RT) = 1.8 x 10–14 (d) K2 = K1(RT)-1 = 3.0 x 10–17 (e) impossible to determine from this information Answer: (b)

18. For the reaction shown below, at equilibrium I2(g) + Br2(g)

2IBr(g)

Kc = 1.1 x 102

(a) Reactants and products will be present in approximately equal concentrations. (b) Reactants will be favored. (c) Products will be favored. (d) Insufficient information is given to decide. (e) It is impossible for the reaction to reach equilibrium.

Answer: (c) 19. Based on the information given in the previous problem, the equilibrium constant for the following reaction would be: 2IBr(g)

I2(g) + Br2(g)

(a) –1.1 x 102 (b) 9.1 x 10-3 (c) 10.5 (d) 1.1 x 10-2 (e) none of the above Answer: (b)

Reaction Quotients

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20. Under which set of conditions must the following reaction always proceed to the right in order to reach equilibrium?

H2(g) + I2(g)

2 HI(g)

(a) Qc < 1 (b) Qc > 1 (c) Qc = 1 (d) Qc > Kc (e) none of these Answer: (e)

21. Which of the following initial concentrations of reactants and products will cause the reaction to proceed to the right in order to establish equilibrium?

2 NO2(g)

2 NO(g) + O2(g)

initial concentrations(M): (a) (b) (c) (d) (e) none of these

NO2(g) 1x10−3 2x10−5 4x10−2 3x10−3

Kc=3.410−7 (at 200C) NO(g) 1x10−2 1x10−3 1x10−2 2x10−3

O2(g) 2x10−3 4x10−2 2x10−5 4x10−3

Answer: (e)

22. If at any moment in time, we find that Qc is larger than Kc for the reaction Cl2(g) + 3 F2(g)

2 ClF3(g)

we can conclude that: (a) The reaction is at equilibrium. (b) The reaction cannot reach equilibrium. (c) The reaction must shift to the right to reach equilibrium. (d) The reaction must shift to the left to reach equilibrium. (e) None of the above are possible.

Answer: (d)

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23. Assume that the reaction quotient, Qc, for the following reaction at 25°C is 1.0 x 10–8 2 NO2(g)

Kc = 7.4 x 10–16 (at 25°C)

2 NO(g) + O2(g)

From this we can conclude: (a) The reaction is at equilibrium. (b) Equilibrium could be reached by adding enough NO or O2 to the system. (c) The reaction must proceed from left to right to reach equilibrium. (d) The reaction must proceed from right to left to reach equilibrium. (e) The reaction can never reach equilibrium.

Answer: (d) 24. If the equilibrium constant for the following reaction is Kc = 1 x 102 2 C2H6(g) + 7 O2(g)

4 CO2(g) + 6 H2O(g)

and all the concentrations were initially 0.10 M, we can predict that the reaction (a) is at equilibrium initially. (b) must shift from left to right to reach equilibrium. (c) must shift from right to left to reach equilibrium. (d) cannot reach equilibrium. (e) cannot be determined unless we have the information necessary to calculate Qc for the reaction.

Answer: (b)

25. For the following reaction: 2 NOCl(g)

2 NO(g) + Cl2(g)

Kc = 5.6 x 10–6 (at 400 K)

If (NOCl) = 0.0222 M, (Cl2) = 0.0222 M, and (NO) = 0.989 M at some moment in time, we can conclude that (a) the reaction is at equilibrium. (b) some of the NOCl must be converted to NO and Cl2 to reach equilibrium. (c) some of the NO and Cl2 must be converted to NOCl to reach equilibrium. (d) it is impossible to determine which way the reaction must go to reach equilibrium. (e) the reaction can never reach equilibrium.

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Answer: (c)

26. Suppose that the reaction quotient, Qc, for a chemical reaction was found to be close to unity (Qc  1). Given this information we can accurately predict (a) that the reaction will proceed in a direction to generate more product. (b) that the reaction will proceed in a direction to use up product. (c) that the reaction is close to equilibrium. (d) nothing about the direction in which the reaction will proceed without additional information. (e) none of the above.

Answer: (d)

27. For the following reaction 2 NO2(g)

2 NO(g) + O2(g)

Kc = 7.4 x 10–16

(at 25°C)

If (NO2) = 0.10 M, (NO) = 0.010 M, and (O2) = 1.0 x 10-5 M, we can correctly predict (a) that the reaction is at equilibrium. (b) that the reaction must shift to the left to reach equilibrium. (c) that the reaction must shift to the right to reach equilibrium. (d) nothing related to the question as to whether the reaction is at equilibrium. (e) nothing related to the question as to which direction the reaction has to shift to reach equilibrium.

Answer: (b)

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Changes in Concentration That Occur as a Reaction Comes to Equilibrium 28. The following reaction was carried out at 250C with the initial concentration of NO2 being 0.70 M and no NO(g) or O2(g) initially present. At equilibrium the NO2 concentration was found to be 0.28 M. Calculate Kc for the reaction.

2NO2(g)

2NO(g) + O2(g)

(a) 0.47 (b) 0.94 (c) 1.9 (d) 0.14 (e) 2.1 Answer: (a)

29. For the reaction below, 0.0500 moles of ethyl acetate are allowed to react with 0.200 moles of water. At equilibrium it is found that 0.0300 moles of acetic acid have been produced. How many moles of ethyl acetate remain at equilibrium?

CH3CH2OCOCH3 + H2O ethyl acetate water

CH3CH2OH + CH3COOH ethyl alcohol acetic acid

(a) 0.0100 moles (b) 0.0300 moles (c) 0.0500 moles (d) 0.170 moles (e) none of these Answer: (e)

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30. A 0.100 M sample of SO2 reacts with 0.060 M O2 to produce SO3 by the reaction shown below. There was no SO3 present initially. The concentration of O2 measured at equilibrium is found to be 0.020 M. What are the equilibrium concentrations of SO2 and SO3?

2 SO2(g) + O2(g)

2 SO3(g)

(a) [SO2] = 0.040M, [SO3] = 0.060M (b) [SO2] = 0.080M, [SO3] = 0.080M (c) [SO2] = 0.080M, [SO3] = 0.120M (d) [SO2] = 0.020M, [SO3] = 0.080M (e) none of these Answer: (d)

31. Calculate the concentrations of Cl2 and ClF3 at equilibrium if a reaction that initially contained 1.000 M concentrations of both Cl2 and F2 is found to have an F2 concentration of 0.700 M at equilibrium.

Cl2(g) + 3 F2(g) Init 1.000M Equil _____

2 ClF3

1.000M 0.00M 0.700M _____

(a) [Cl2] = 0.700M (b) [Cl2] = 0.700M (c) [Cl2] = 0.500M (d) [Cl2] = 0.900M (e) none of these

[ClF3] = 0.700M [ClF3] = 0.300M [ClF3] = 0.300M [ClF3] = 0.200M

Answer: (d)

32. For the reaction 2 NOCl(g)

Kc = 5.6 x 10–6

2 NO(g) + Cl2

(at 400 K)

what is the relationship between the magnitude of the changes in the concentrations of NOCl and Cl2 as this reaction comes to equilibrium? (a) [NOCl] = [Cl2] (b) [NOCl] = 2 [Cl2] (c) [Cl2] = 2 [NOCl] (d) [NOCl] = 1/[Cl2] (e) 2 [NOCl] = [Cl2] Answer: (b)

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33. What is the relationship between the changes in the concentration of NO2 and O2 as the following reaction comes to equilibrium?

2 NO2(g)

2 NO(g) + O2(g)

Kc = 7.4 x 10–16

(at 25°C)

(a) [NO2] = -[O2] (b) [NO2] = -2 [O2] (c) [O2] = -2 [NO2] (d) There is no relationship between the changes in the concentrations of these substances. (e) none of the above Answer: (b)

34. For the following reaction 2 C2H6(g) + 7 O2(g)

4 CO2(g) + 6 H2O(g)

if the reaction mixture initially contained all species at the same initial concentration, C, and if the concentration of O2 at equilibrium was expressed as [O2] = C + 7 C , then the concentration of H2O at equilibrium would be correctly expressed as [H2O] = (a) C + C (b) C - C (c) C + 6C (d) C – 6C (e) none of the above Answer: (d)

35. Calculate the COCl2, CO, and Cl2 concentrations when the following gasphase reaction reaches equilibrium at 300°C.

initial: equilibrium:

COCl2(g)

CO(g) + Cl2(g)

0 ?

0.10 M 0.50 M ? ?

Answer: [CO] = 1 x 10–5, [Cl2] = 0.40, [COCl2] = 0.10 M

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Kc = 4.0 x 10–5

36. Calculate the NO, NO2, and O2 concentrations in the following gas-phase reaction at equilibrium at 200°C. 2 NO2(g) initial: 0 equilibrium: ?

Kc = 1.6 x 10–8

2 NO(g) + O2(g) 0.10 M ?

0.10 M ?

Answer: [NO] = 5.6 x 10–5, [O2] = 0.050, [NO2] = 0.10 M

37. Calculate the equilibrium concentrations of SO3, SO2, and O2 if 0.100 moles of SO3 in a 10-L flask decompose to form SO2 and O2 at 500 K. 2 SO3(g)

Kc = 1.4 x 10–11

2 SO2(g) + O2(g)

Answer: [O2] = 7.0 x 10–6, [SO2] = 1.4 x 10–5, [SO3] = 0.010 M

38. Calculate the concentrations of all reactants and products when a mixture that is initially 0.10 M in N2, H2, and NH3 comes to equilibrium at 600 K. N2(g) + 3 H2(g)

2 NH3(g)

Kc = 0.045

Answer: [NH3]  1.0 x 10–2, [N2] = 0.14, [H2] = 0.22 M 39. Assume an equilibrium constant Kc = 1.7 x 10–2 for the reaction: N2(g) + 3 H2(g)

2 NH3(g)

Calculate (to a first approximation) the equilibrium concentrations of NH3, H2 and N2 when 0.10 moles per liter each of N2 and H2 are mixed and allowed to react to equilibrium.

Answer: NH3 = 1.3 x 10–3, N2  0.10, H2  0.10 M

40. Calculate the COCl2, CO, and Cl2 concentrations when the following gasphase reaction reaches equilibrium at 300°C. COCl2(g) initial: equilibrium:

0.10 M ?

CO(g) + Cl2(g) 0 ?

0 ?

Answer: [CO] = [Cl2] = 2.0 x 10–3, [COCl2] = 0.10 M

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Kc = 4.0 x 10–5

41.

Calculate the NO, NO2, and O2 concentrations in the following gas-phase reaction at equilibrium at 200°C. 2 NO2(g) initial: 0.10 M equilibrium: ?

2 NO(g) + O2(g) 0 ?

Kc = 1.6 x 10–8

0 ?

Answer: [NO2] = 0.10, [NO] = 6.8 x 10–4, [O2] = 3.4 x 10–4 M

Hidden Assumptions that make Equilibrium Calculations Easier What Do We Do When the Approximation Fails?

The following three questions are based on the following reaction: 2H2O(g) 2H2(g) + O2(g) Kc = 2 x 10-42 (at 25oC)

42. Starting with 0.10 M H2O(g), what would be the best strategy for solving for the equilibrium concentrations of H2O, H2 and O2? (a) Since the change in the concentrations are expected to be large going to equilibrium, the equation will have to be solved exactly. (b) Since the changes in the concentrations are expected to be large going to equilibrium, it is best to assume that all of the H2O(g) is converted to products and then calculate the equilibrium concentration. (c) Since the changes in concentrations are expected to be small going to equilibrium, it is best to assume that the changes will be small relative to the initial concentration of H2O(g). (d) Since the changes in concentrations are expected to be small going to equilibrium, it is best to assume that the changes will be large relative to the initial concentration of H2O(g). (e) There is no reasonable way to even approximate the answer to this problem.

Answer: (c)

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43. Starting with only 0.10 M H2O(g), which equation correctly shows the equilibrium concentrations of the reactants and products? (C represents the change in concentration of O2(g)) (a)

2 0.10 -2C 2 x 10-42 =

2 C

2C

(b)

2 C -42

2 x 10

C

=

2 0.10 -C

(c)

2 0.10 +2C 2 x 10-42 = (d)

2 C

2C

2 C

2C 2 x 10-42 =

2 0.10 -2C

(e) None of the above Answer: (d) 44. The equation below solves for the change in concentration for the given reaction assuming the initial concentration of H2O(g) is 0.10 M. Which equation will result if it is assumed that the change will be much smaller than the initial concentration of H2O(g)?

2 2C -42

2 x 10

C

=

2 0.10 -2C

10-17

(a)

2 C -42

2 x 10

C

= 0.10

2

(b)

2 

1 2 x 10-42 =

2 0.10 -2C

(c)

2 C

2C -42

2 x 10

= 2C

2

(d)

C -42

2 x 10

=

2 0.10 -2C

(e) None of the above

Answer: (a)

45. When solving a problem with the initial conditions outlined below, we rearrange the problem to give an intermediate set of conditions where the concentration of one of the reactants or products is equal to zero.

Fe(CO)5 initial: 0.10 M

Fe(CO)4 + CO 0.10 M 0.005 M

Our choice of whether to push the reaction all the way to the right or all the way to the left is determined by which of the following goals? (a) to make both Qc and Kc large (b) to make both Qc and Kc small 10-18

(c) to bring Qc as close as possible to Kc (d) to make the difference between Qc and Kc as large as possible Answer: (c)

Le Châtelier’s Principle 46. What would happen if F2 was removed from the following reaction while it was at equilibrium?

Cl2(g) + 3 F2(g)

2 ClF3(g)

(a) The concentration of Cl2 would increase as the reaction returned to equilibrium. (b) The concentration of both Cl2 and F2 would increase as the reaction returned to equilibrium. (c) The concentration of Cl2 would decrease as the reaction returned to equilibrium. (d) The concentration of both Cl2 and F2 would decrease as the reaction returned to equilibrium. (e) The concentration of both F2 and Cl2 would remain the same because the reaction would be at equilibrium.

Answer: (b)

47. What is the effect of increasing the temperature on the following exothermic reaction?

2 CO(g) + O2(g)

2 CO2(g)

(a) The equilibrium will shift to the right. (b) The equilibrium will shift to the left. (c) The partial pressure of CO2 will increase. (d) The partial pressure of O2 will decrease. (e) It will have no effect.

Answer: (b)

48. What would happen to the extent of the decomposition of PCl5 at a given temperature if the pressure on the system were decreased by increasing the volume of the container at a constant temperature?

PCl5(g)

PCl3(g) + Cl2(g)

(a) It would increase. (b) It would decrease.

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(c) It would remain constant. (d) Not enough information is given to determine this.

Answer: (a)

49. Which of the following factors will cause this chemical reaction to shift its equilibrium to the right?

Xe(g) + 2 F2(g)

XeF4(g)

H = -218 kJ/molrxn

(a) increasing the temperature (b) increasing the pressure (c) decreasing the pressure (d) increasing the concentration of XeF4 (e) decreasing the concentration of F2 Answer: (b)

50. The following chemical reaction has reached equilibrium. Which of the changes listed below would cause the equilibrium to shift back toward the reactants?

N2(g) + 3 H2(g) 2 NH3(g) H = −92.2 kJ/molrxn (a) increasing the pressure (b) increasing the concentration of N2 (c) increasing the temperature (d) decreasing the concentration of NH3 (e) none of these Answer: (c)

51. The following reaction is at equilibrium. Which of the conditions listed below will cause the reaction to shift toward the products?

2 SO3(g) + 2 Cl2(g)

2 SO2Cl2(g)+ O2(g)

(a) The concentration of O2 is increased. (b) The concentration of SO2Cl2 is decreased. (c) The concentration of SO3 is decreased. (d) The pressure is decreased. (e) none of these Answer: (b)

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52. The following reactions are at equilibrium. Beside each reaction is a change that is made to the system which may cause a shift in equilibrium. Which of the reactions will shift its equilibrium toward the products? (a) N2(g) + O2(g) 2 NO(g) (b) PF5(g) PF3(g) + F2(g) (c) Cl2(g) + 3 F2(g) 2 ClF3(g) (d) 2 SO3(g) 2 SO2(g) + O2(g) (e) none of these

the concentration of NO is increased the concentration of F2 is increased the concentration of Cl2 is decreased the concentration of SO2 is decreased

Answer: (d) 53. The following reactions are at equilibrium. Which of the reactions will shift its equilibrium toward the reactant side if the pressure is increased while holding the temperature constant? (a) 2 NO(g) + 2 H2(g) N2(g) + 2 H2O(g) (b) 2 NOCl(g) 2 NO(g) + Cl2(g) (c) 2 NO(g) + O2(g) 2 NO2(g) (d) N2(g) + O2(g) 2 NO(g) (e) NO(g) + NO2(g) N2O3(g) Answer: (b) 54. What would happen if Cl2 is added to the following system at equilibrium?

H2(g) + Cl2(g)

2 HCl(g)

(a) The concentrations of both H2 and HCl would remain the same. (b) The concentrations of both H2 and HCl would increase. (c) The concentrations of both H2 and HCl would decrease. (d) The concentration of H2 would increase, but HCl would decrease. (e) The concentration of H2 would decrease, but HCl would increase.

Answer: (e)

55. Which of the following statements correctly describes the following reaction?

2 NH3(g)

N2(g) + 3 H2(g)

H = 92.2 kJ

(a) An increase in either pressure or temperature shifts the reaction to the left. (b) If the pressure of N2 is doubled and a new equilibrium established, the concentration of H2 at equilibrium will be reduced. (c) More ammonia will be present in the equilibrium mixture if a catalyst is added. (d) none of the above

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(e) all of the above Answer: (b)

56. What would be the effect of removing some NO2 from the following system after the reaction reaches equilibrium at 25°C? 2 NO2(g)

Kc = 7.4 x 10–16

2 NO(g) + O2(g)

(a) The NO and O2 concentrations would increase. (b) The NO and O2 concentrations would decrease. (c) The NO concentration would increase but O2 would decrease. (d) The O2 concentration would increase but NO would decrease. (e) The concentrations of NO and O2 would remain constant.

Answer: (b)

57. The addition of I2(g) to a fixed volume container with I2(g), H2(g), and HI(g) at equilibrium at a given temperature would cause H2(g) + I2(g)

2 HI(g)

(a) the concentrations of both H2 and HI to decrease. (b) the concentrations of both H2 and HI to increase. (c) the concentration of H2 to increase and HI to decrease. (d) the concentration of H2 to decrease and HI to increase. (e) Not enough information is given to determine how the concentrations will change.

Answer: (d)

58. What would happen if O2 were removed from the following system at equilibrium at 25°C? 2 NO2(g)

2 NO(g) + O2(g)

Kc = 7.4 x 10–1 at 25oC

(a) The NO2 and NO concentrations would increase. (b) The NO2 and NO concentrations would decrease. (c) The NO2 concentration would increase and the NO concentration would decrease. (d) The NO2 concentration would decrease and the NO concentration would increase. (e) none of the above Answer: (d)

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59. The equilibrium constant for the following reaction becomes larger as the temperature at which the reaction is run increases.

2 NOCl(g)

2 NO(g) + Cl2(g)

Kc = 5.6 x 10–6 at 400 K

What does this tell us about the reaction? (a) Heat is given off in the reaction. (b) Heat is absorbed in the reaction. (c) Heat is neither given off nor absorbed in the reaction. (d) Increasing the amount of NOCl would also increase the equilibrium constant for the reaction. (e) Increasing the amount of NO would also increase the equilibrium constant for the reaction.

Answer: (b)

60. What is the effect of increasing the temperature on the following reaction?

2 CO(g) + O2(g)

2 CO2(g) + heat

(a) The equilibrium will shift to the left. (b) The equilibrium will shift to the right. (c) The partial pressure of CO2 will increase. (d) The partial pressure of O2 will decrease. (e) It will have no effect.

Answer: (a)

61. What would happen to the extent of the decomposition of PCl5 at a given temperature if the pressure on the system were decreased?

PCl5(g)

PCl3(g) + Cl2(g)

(a) It would increase. (b) It would decrease. (c) It would remain constant. (d) Not enough information is given to determine this.

Answer: (a)

62. What would be the effect of decreasing the pressure on the following reaction after it reaches equilibrium?

2 NO2(g)

2 NO(g) + O2(g)

(a) The NO and O2 concentrations would increase.

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(b) The NO and O2 concentrations would decrease. (c) The NO concentration would increase but O2 would decrease. (d) The O2 concentration would increase but NO would decrease. (e) The concentrations of NO and O2 would remain constant.

Answer: (a)

63. One of the products of the reaction between NO2 and CO is CO2. By writing the equation for the reaction (Hint: the coefficients for all reactants and products are 1), decide whether an increase in pressure: (a) favors CO2 formation (b) decreases CO2 formation (c) has no effect on CO2 formation (d) cannot tell from information supplied Answer: (c)

64. Which of the following equilibria would not be affected by changes in pressure? (a) 2 NO(g) + O2(g) 2 NO2(g) (b) N2O4(g) 2 NO2(g) (c) 4 NH3(g) + 5 O2(g) 4 NO(g) + 6 H2O(g) (d) CaCO3(s) CaO(s) + CO2(g) (e) CO(g) + H2O(g) CO2(g) + H2(g) Answer: (e)

Solubility Product Expressions 65. What is the concentration in moles per liter of pure PbSO4(s) given that it has a density of 6.29 g/cm3? (a) 6.29 (b) 20.7 (c) 1.00 (d) 303 (e) The concentration of a pure solid cannot be found.

Answer: (b)

66. What is the correct solubility constant expression for PbSO4?

Ksp =

[Pb2+][SO42-] [PbSO4]

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(a) (b) Ksp = [Pb2+][SO22-][O2] (c) Ksp = [Pb2+][SO42-] (d) Ksp =

[PbSO4] [Pb2+][SO42-]

(e) none of the above Answer: (c) 67. Which is the correct reaction for the solubility constant expression (Ksp) of CaF2(s)? (a) 2F-(aq) + Ca2+(aq) CaF2(s) (b) 2F(aq) + Ca (aq) CaF2(s) (c) CaF2(s) CaF2(g) (d) CaF2(s) Ca(g) + 2F(g) (e) CaF2(s) Ca2+(aq) + 2F-(aq) Answer: (e)

68. What is the correct solubility constant expression of Ba3(PO4)2? (a) Ksp =

[Ba2+]3[PO43-]2 [Ba3(PO4)2]

(b) Ksp = [Ba2+]2[PO43−-]3 (c) Ksp = [Ba2+][PO43−-] (d) (d) Ksp = [Ba2+]3[PO43−]2 (e) none of the above (d) Ksp = [Ba2+]3[PO43-]2 Formatted

(e) None of the above Answer: (d)

69. What is the correct solubility constant expression for CaF2(s)? (a) Ksp =

[Ca2+] [F ]2 [CaF 2]

10-25

Formatted

(b) Ksp = [Ca2+]2[F−] (c) Ksp = [Ca2+][F−] (d) Ksp = [Ca2+][F2]2 (e) none of the above Answer: (e)

70. What is the correct solubility constant expression for AgCl(s)? (a) Ksp = [Ag+][Cl−] (b) Ksp = [Ag+] + [Cl−] (d) Ksp = [Ag+]/[Cl−] (d) Ksp = [Ag][Cl] (e) none of the above Answer: (a)

Ksp and Solubility 71. What is the solubility in moles per liter of AgBr in water given that the Ksp is 5.0 x 10–13? (a) 5.0 x 10−13 (b) 2.5 x 10−13 (c) 7.1 x 10−7 (d) 4.0 x 10−6 (e) 2.5 x 10−7

Formatted

Answer: (c)

72. What is the solubility in moles/L of CaF2 in water given that the Ksp is 4.0 x 10−11? (a) 2.2 x 10−4 (b) 4.0 x 10−11 (c) 3.5 x 10−4 (d) 6.0 x 10−5 (e) 2.0 x 10−11 Answer: (a)

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73. If 1.3 x 10−5 moles of AgCl will dissolve in 1.0 liter of water at 25C, what is the Ksp of AgCl? (a) 5.1 x 10−9 (b) 3.6 x 10−3 (c) 6.4 x 10−5 (d) 1.3 x 10−5 (e) 1.7 x 10-10 Answer: (e) 74. If 6.5 x 10−5 moles of Ag2CrO4 will dissolve in 1.0 liter of water at 25C, what is the Ksp of Ag2CrO4? (a) 6.5 x 10−5 (b) 4.2 x 10−9 (c) 2.7 x 10−13 (d) 1.1 x 10−12 (e) 3.3 x 10−5 Answer: (d)

75. If the Ksp of PbBr2 is 8.9 x 10-6, what is the solubility of PbBr2(s) in moles/L? (a) 8.9 x 10−6 (b) 1.3 x 10−2 (c) 3.3 x 10−3 (d) 0.21 (e) none of the above Answer: (b) 76. What is the fluoride ion concentration in a saturated solution of SrF2 if the strontium ion concentration is 5.8 x 10−4 M? (a) 5.8 x 10−4 (b) 2.9 x 10−4 (c) 1.2 x 10−3 (d) 1.8 x 10−3 (e) none of the above Answer: (c)

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77. What is the Ksp of SrF2(s) if 0.073 g of SrF2 dissolves in 1.00 liter of solution? (a) 5.8 x 10−4 (b) 1.7 x 10−3 (c) 1.6 x 10−3 (d) 7.8 x 10−10 (e) none of the above Answer: (d) 78. What is the Ksp of PbBr2 if the concentration of PbBr2(aq) in a saturated solution is 1.3 x 10-2 moles/L? (a) 8.8 x 10−6 (b) 1.3 x 10−2 (c) 3.9 x 10−2 (d) 2.2 x 10−2 (e) none of the above Answer: (a)

79. Which of the following equations correctly describes the relationship between the concentrations of Pb2+ and Br- ions in a saturated solution of PbBr2(s)? (a) [Pb2+] = 2[Br-] (b) 2[Pb2+] = [Br-] (c) [Pb2+] = [Br-] (d) 3[Pb2+] = [Br-] (e) none of the above Answer: (b)

80. Which of the following equations correctly describes the relationship between the concentrations of Sr2+ and PO43- ions in a saturated solution of Sr3(PO4)2(s)? (a) 3[Sr2+] = 2[PO43−] (b) [Sr2+] = [PO43−] (c) 2[Sr2+] = 3[PO43−] (d) [Sr2+] = 2/3[PO43−] (e) none of the above Answer: (c) 10-28

81. Which of the following equations describes the relationship between the solubility product for AgCl and the solubility of this compound? (a) Ksp = 2C2 (b) Ksp = C (c) Ksp = 2C (d) Ksp = C2 (e) none of the above Answer: (d) 82. Which of the following equations describes the relationship between the solubility product for Ag2SO4 and the solubility of this compound? (a) Ksp = 4C3 (b) Ksp = 2C (c) Ksp = 4C2 (d) Ksp = C3 (e) none of the above Answer: (a)

83. Of the compounds in the table below, _______ is the least soluble.

Compound CdS CuS PbS MnCO3

Ksp 8.0 x 10-27 6.3 x 10-37 8.0 x 10-28 1.8 x 10-11

(a) CdS (b) CuS (c) PbS (d) MnCO3 (e) The three sulfides are equally soluble, and are all less soluble than MnCO3.

Answer: (b)

Ion Product (Qsp) and Solubility Calculations

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84. A solution is prepared in which the Ag+ ion concentration is 2.5 x 10-3 M and the Cl- ion concentration is 3.2 x 10-4 M. Which of the following statements is true given that the Ksp for AgCl is 1.8 x 10-10? (a) No AgCl(s) will precipitate because Qsp<Ksp. (b No AgCl(s) will precipitate because Qsp>Ksp. (c) No AgCl(s) will precipitate because the Ksp value indicates that AgCl is infinitely soluble in water. (d) A precipitate of AgCl(s) will form because Qsp>Ksp. (e) Insufficient data is given to make a prediction about whether a precipitate will form or not.

Answer: (d) 85. If a solution has a Mg2+ ion concentration of 5.0 x 10-3 M and an OHconcentration of 0.010 M, will a precipitate form? Ksp Mg(OH)2 = 1.7 x 10-6.

Answer: Qsp = (5.0 x 10-3)(0.010)2 = 5.0 x 10-7 Since Qsp < Ksp, no precipitate will form 86. A solution is prepared in which the Pb2+ ion concentration is 0.020 M and the Br- ion concentration is 0.010 M. Which of the following statements is true given that Ksp for PbBr2 is 7.9 x 10-5? (a) No PbBr2(s) will precipitate because Qsp<Ksp. (b No PbBr2(s) will precipitate because Qsp>Ksp. (c) PbBr2 can only be formed from the reaction of Pb(s) with Br2(l). (d) A precipitate of PbBr2(s) will form because Qsp>Ksp. (e) Insufficient data is given to make a prediction about whether a precipitate will form or not.

Answer: (a)

87. A 50.0 mL solution of 0.015 M of Ca(NO3)2 is added to 50.0 mL of 0.010 M NaF. Write a net ionic equation that shows the formation of a precipitate of CaF2(s). Given that the Ksp of CaF2(s) is 4.0 x 10-11, will a precipitate form from this mixture?

Answer: net ionic equation: Ca2+(aq) + 2F-(aq)

CaF2(s)

Qsp = 1.9 x 10-7 Since Qsp > Ksp, a precipitate will form.

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The Common Ion Effect 88. In which of the following will the least PbBr2(s) dissolve? (a) pure water (b) 0.01 M NaBr (c) 0.01 M KBr (d) 0.02 M NaBr (e) PbBr2(s) will be equally soluble in all of the above solutions.

Answer: (d)

89. Which of the following statements is correct? (a) CaCO3(s) is more soluble in 0.1 M Na2CO3 than in pure water. (b) CaCO3(s) is less soluble in 0.1 M Na2CO3 than in pure water. (c) CaCO3(s) has the same solubility in pure water and 0.1 M Na2CO3. (d) More information is needed to decide whether CaCO3(s) is more soluble in pure water or 0.1 M Na2CO3.

Answer: (b)

90. Ag2SO4(s) is in equilibrium with silver and sulfate ions in solution: Ag2SO4(s)

2Ag+(aq) + SO42-(aq)

Which of the following will not increase the amount of solid Ag2SO4(s) present? (a) boiling off some of the water (b) adding Na2SO4 to the solution (c) adding AgNO3 to the solution (d) adding NaCl to the solution (e) All of the above will increase the amount of Ag2SO4 present.

Answer: (d) 91. What is the concentration of Ag+ (mol/L) ion that results from saturating a 0.10 M NaBr solution with AgBr? The Ksp of AgBr is 5.0 x 10-13. (a) 5.0 x 10-12 M (b) 2.2 x 10-6 M

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(c) 0.10 M (d) 1.2 x 10-11 M (e) none of the above Answer: (a)

92. Approximately how many grams of Ag2CO3 will dissolve in 1.0 liter of a solution that is 0.10 M in Na2CO3? The Ksp of Ag2CO3 is 6.2 x 10-12. (a) 2.2 x 10-3 g (b) 1.1 x 10-3 g (c) 1.7 x 10-8 g (d) 3.9 x 10-6 g (e) none of the above Answer: (b)

93. Which of the following changes will shift the following equilibrium to the right?

AgCl(s)

Ag+(aq) + Cl-(aq)

H = 65.4 kJ/molrxn

(a) adding more AgCl(s) (b) adding NaCl (c) adding AgNO3 (d) increasing the temperature (e) None of the above will shift the equilibrium to the right.

Answer: (d)

94. A solution is 0.10 M in Ca(NO3)2. What would the concentration of NaOH need to be for a precipitate of Ca(OH)2 to form? Ksp Ca(OH)2 = 5.5 x 10-6. (a) 5.5 x 10-6 (b) 5.5 x 10-5 (c) 7.4 x 10-3 (d) 1.2 x 10-3 (e) none of the above Answer: (c)

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10-33

Chapter 11: Acids and Bases The Arrhenius Definition of Acids and Bases 1.

Which of the following compounds is not an Arrhenius acid? (a) HCl (b) H2SO4 (c) H2O (d) FeCl3 (e) H2S

Answer: (d)

2.

Which of the following compounds is not an Arrhenius base? (a) NaOH (b) Al(OH)3 (c) Ca(OH)2 (d) H2O (e) NH3

Answer: (e)

The Brønsted-Lowry Definition of Acids and Bases 3.

Consider the following reaction: H2PO4–(aq) + HCO3–(aq)

H2CO3(aq) + HPO42–(aq)

Brønsted would identify the acidic species as: (a) H2PO4– and H2CO3 (b) H2PO4– and HPO42– (c) HCO3– and H2CO3 (d) HCO3– and HPO42– (e) H2CO3 and HPO42– Answer: (a)

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4.

Which compound cannot act as a Brønsted base? (a) NH3 (b) H2O (c) CH4 (d) F− (e) H2S

Answer: (c)

5.

Which of the following compounds cannot be a base? (a) H2O (b) NH3 (c) CO32– (d) OH– (e) NH4+

Answer: (e)

6.

Which of the following compounds cannot be a Brønsted base? (a) OH– (b) H2O (c) NH3 (d) NH4+ (e) All of these compounds can be Brønsted bases.

Answer: (d)

7.

Which of the following isn’t a Brønsted base? (a) PF5 (b) H2O (c) NH3 (d) S2– (e) CN–

Answer: (a)

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Conjugate Acid-Base Pairs 8.

What is the conjugate acid of hydrogen carbonate ion, HCO3−? (a) H3O+ (b) OH− (c) H2O (d) CO32− (e) H2CO3

Answer: (e)

9.

What is the conjugate base of thiocyanic acid, HSCN? (a) H3O+ (b) OH− (c) SCN− (d) CN− (e) H2SCN+

Answer: (c) 10. The dihydrogen phosphate ion, H2PO4−, can behave as both an acid and as a base. What is the chemical formula of the conjugate acid of dihydrogen phosphate? What is the chemical formula of the conjugate base of dihydrogen phosphate? (a) conjugate acid = H3PO4 (b) conjugate acid = H2PO4+ (c) conjugate acid = HPO42− (d) conjugate acid = H3PO4 (e) none of these

conjugate base = HPO42− conjugate base = H2PO42− conjugate base = H3PO4 conjugate base = PO43−

Answer: (a) 11. Which of the following is the conjugate acid of the HPO42– ion? (a) H3PO4 (b) H2PO4– (c) HPO42– (d) PO43– (e) H3O+ Answer: (b)

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12. Ethanol (CH3CH2OH), like water, can behave both as an acid or as a base.

Which of the following structures corresponds to the Brønsted conjugate acid and conjugate base of ethanol? (a) (b) (c) (d) (e)

Conjugate acid CH3CH2OH2+ CH3CH3 CH3CH2OH2O CH3CH2O− none of these

Conjugate base CH3CH2O− CH3CH2− CH3CH2OHO CH3CH2OH2+

Answer: (a) 13. What is the conjugate base of HSO4–? (a) H2O (b) H2SO4 (c) HSO4– (d) SO3 (e) none of the above Answer: (e) 14. Methanol is described by the following skeleton structure.

H H

C O

H

H

Which substance is formed when CH3OH acts as a Brønsted base? (a) CH3O– (b) CH3OH (c) CH3OH2+ (d) CH3OH (e) H3O+ Answer: (c) 15. The compound HB can be represented as  where  represent B- and  represents H+. Given this, ___ would represent the conjugate acid of HB and ____ would represent the conjugate base.

1. 

2. 

3. 

(a) 1,2 (b) 1,5 (c) 3,4

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4. 

5. 

(d) 1,3 (e) None of the above.

Answer: (d) 16. Which of the following is the conjugate base of a strong acid? (a) OH– (b) HSO4– (c) NH2– (d) S2– (e) H3O+ Answer: (b)

17. Which of the following represents a conjugate acid/base pair? (a) H+, OH– (b) H3O+, OH– (c) NH3, NH4+ (d) NaOH, HCl (e) All of the above are conjugate acid/base pairs.

Answer: (c) 18. Which of the following isn’t a Brønsted conjugate acid/base pair? (a) NH3/NH2– (b) CH3CO2H/CH3CO2– (c) H2O/O2– (d) H3O+/H2O (e) HCN/CN– Answer: (c)

Acids and Bases in Water 19. A solution is known to have a hydronium ion concentration of 1.0 x 10 –6 M. What percentage of the total H3O+ ion concentration comes from the dissociation of water? (a) 1% (b) 3% (c) 5% (d) 10%

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(e) 100% Answer: (d)

pH as a Measure of the Concentration of H3O+ Ion 20. What is the [H3O+] concentration of a solution with a pH of 8.75? (a) 1.8 x 10−9M (b) 9.4 x 10−1M (c) 8.7 x 10–5M (d) 5.6 x 108M (e) This is a basic solution, [H3O+] cannot be determined.

Answer: (a)

21. A solution of weak base is added to a beaker of water. Which of the following will be true as the base is added to the water? (a) [H3O+] increases [OH−] decreases (b) [H3O+] decreases [OH−] increases (c) [H3O+] decreases [OH−] decreases + (d) [H3O ] remains constant [OH−] increases (e) none of these

pH increases pH decreases pH increases pH increases

Answer: (e) 22. Determine the [OH−] for an aqueous solution at 25C with a pH of 4.80. (a) 6.3x10−10M (b) 1.6x10−5M (c) 1.6x10−9M (d) 8.0x10−8M (e) none of these Answer: (a)

23. A strong acid is added to water. Which of the following will be true? (a) [H3O+] increases, (b) [H3O+] decreases, (c) [H3O+] decreases, (d) [H3O+] increases, (e) none of these

[OH−] decreases, [OH−] increases, [OH−] decreases, [OH−] increases,

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pH increases pH decreases pH decreases pH increases

Answer: (e) 24. What is the H3O+ ion concentration in a solution that has a pH of 5.75? (a) 5.8 x 10–5 M (b) 3.2 x 10–3 M (c) 0.75 x 10–5 M (d) 3.6 x 10–7 M (e) 1.8 x 10–6 M Answer: (e) 25. What is the pH of a solution that is 1.7 x 10–4 M in H3O+? (a) 3.77 (b) 4.77 (c) 4.23 (d) 10.23 (e) none of these Answer: (a)

26. What is the hydroxide ion concentration in a pH = 5.14 solution? (a) 1.0 x 10–14 (b) 1.4 x 10–9 (c) 1.0 x 10–7 (d) 7.2 x 10–6 (e) 1.4 x 10–2 Answer: (b)

Relative Strengths of Acids and Bases 27. Which equation best describes a 0.10 M solution of a strong acid such as hydroiodic acid (HI: Ka = 3 x 109)? (a) [H3O+] = initial concentration of HI (b) [H3O+]  [OH–] (c) [H3O+]  final concentration of HI (d) [H3O+]  [H2O] (e) [H3O+]  Ka [OH–] Answer: (a)

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28. Which of the following acids would have the strongest conjugate base? (a) HOCl (Ka = 2.9 x 10–8) (b) HOBr (Ka = 2.4 x 10–9) (c) HOI (Ka = 2.3 x 10–11) (d) H2O2 (Ka = 2.2 x 10–12) (e) H2O (Ka = 1.8 x 10–16) Answer: (e) 29. Using a table of Ka’s, determine which of the following is the strongest Brønsted acid. (a) NH4+ (b) HCl (c) HNO3 (d) HSO4– (e) HClO4 Answer: (e) 30. Using a table of Ka’s, determine which of the following is the strongest base. (a) I– (b) NO3– (c) HS– (d) O2– (e) OH– Answer: (d) 31. Using a table of Ka’s, determine which of the following is the strongest acid. (a) NH2– (b) HS– (c) H2O (d) CH4 (e) NH4+ Answer: (e)

11-8

32. If the strength of the following acids increases from left to right CH3NH2 < CH3NH3+ < C6H5NH3 + < C6H5CO2H < HCl the order of increasing base strength, reading from left to right, must be: (a) CH3NH– < CH3NH2 < C6H5NH2 < C6H5CO2– < Cl– (b) Cl– < C6H5CO2– < C6H5NH2 < CH3NH2 < CH3NH– (c) C6H5NH2 < CH3NH2 < Cl– < C6H5CO2– < CH3NH– (d) CH3NH– < C6H5CO2– < Cl– < CH3NH2 < C6H5NH2 (e) none of the above Answer: (b)

33. Explain the following observation: Water becomes acidic when CO 2 is bubbled into it.

Answer: Because of the formation of carbonic acid, H2CO3 which dissociates to produce hydronium ions.

34. Which of the following is the weakest acid? (a) HClO4 (b) HCl (c) HF (d) HI (e) HBr Answer: (c)

35. Which of the following is a weak base? (a) (CH3)3N (b) HF (c) KOH (d) Na (e) H2 Answer: (a)

36. Which of the following is a strong acid? (a) HNO3 (b) H2S (c) HNO2 (d) HCO3–

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(e) HOCl Answer: (a)

Relative Strengths of Conjugate Acid-Base Pairs 37. Which of the following would have the strongest conjugate base? (a) H2O (b) H2S (c) NH3 (d) PH3 (e) CH4 Answer: (e)

38. Which compound should dissolve in water to form an 0.1 M solution with a pH of approximately 9? (a) HCl (b) NaOAc (CH3CO2Na) (c) HOAc (CH3CO2H) (d) NaCl (e) KOH Answer: (b)

Relative Strengths of Pairs of Acids and Bases 39. Arrange the following bases in order of increasing strength: NH3, PH3, H2O and H2S. Answer: H2S < H2O < PH3 < NH3

40. Which of the following statements is true? (a) H2O is a stronger acid than HF(aq). (b) H2O is a stronger acid than H2S(aq). (c) H2O is a stronger acid than the H3O+(aq) ion. (d) H2O is a stronger acid than NH3(aq). (e) None of the above are true.

Answer: (d)

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41. What can be correctly concluded from the fact that the following acid-base reaction proceeds to the right, as written?

NH2–(aq) + HSO4–(aq)

NH3(aq) + SO42–(aq)

(a) NH3 is a stronger base than NH2–. (b) NH3 is a stronger acid than HSO4–. (c) NH3 is a weaker acid than NH2–. (d) NH3 is a weaker base than HSO4–. (e) NH3 is a weaker acid than HSO4–.

Answer: (e)

42. Which of the following statements is true? (a) H2O is a weaker acid than NH3. (b) PH3 is a weaker acid than NH3. (c) OH– is a weaker base than H2O. (d) NH2– is the conjugate acid of NH3. (e) ClO– is a stronger base than ClO4–.

Answer: (e) 43. The hydride ion, (H-), is a stronger base than the hydroxide ion, (OH-). The products of the reaction between H- (aq) and H2O(l) would be. (a) H2O(aq) (b) OH- (aq) + H2(g) (c) OH- (aq) + 2H+(aq) (d) H2O2(aq) (e) No reaction would occur.

Answer: (b)

Relationship of Structure to Relative Strengths of Acids and Bases 44. Which of the following structural factors is primarily responsible for H2Se being a stronger acid than H2S? (a) bond polarity (b) size of the anion (c) electron withdrawing effect of oxygen (d) electric charge on the acid (e) sulfur has d electrons

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Answer: (b) 45. Which of the following is the best Lewis structure for carbonic acid (H2CO3)?

H O

(a)

O OH C H O

(b)

C OH O O

(c) HC O H O H O

(d)

C OH O

(e) None of the above are an acceptable Lewis structure for carbonic acid.

Answer: (d)

46. Of the following, which is the strongest acid (lone pairs omitted for clarity)?

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O

(a)

C H

O H O

(b)

C

Br

O H O

(c)

C F

O H O

(d)

C Cl

O H

(e) All of the above will have approximately the same acid strength

Answer: (c) 47. Which of the following structural factors is primarily responsible for H2AsO4− being a stronger acid than HAsO42−? (a) bond polarity (b) size of the anion (c) electron withdrawing effect of oxygen (d) electric charge on the acid (e) arsenic has d electrons Answer: (d)

48. Which of the following would you expect to be the strongest acid? (a) NH3 (b) PH3 (c) H2O (d) H2S (e) CH4 Answer: (d)

49. Which of the following would be the weakest Brønsted acid? (a) H2Se (b) H2S 11-13

(c) H2O (d) SH– (e) OH– Answer: (e)

50. Which of the following would be the strongest Brønsted base? (a) HSO3– (b) HSO4– (c) H2SO3 (d) SO32– (e) SO42– Answer: (d)

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51. Which of the following would be the strongest Brønsted base? (a) H2O (b) SH– (c) OH– (d) S2– (e) O2– Answer: (e)

Strong Acid pH Calculations 52. What is the pOH of a 0.025M solution of hydrobromic acid (Ka = 1 x 109)? (a) 0.025 (b) 1.60 (c) 12.40 (d) 9.00 (e) 5.00 Answer: (c)

53.

What is the pH of a 0.0250 M HClO4 solution. (a) 2.50 (b) -2.50 (c) 1.60 (d) 12.40 (e) none of these

Answer: (c) 54. Which of the following statements should be true for a 1.0x 10 −4 M solution of a strong acid HBr? (a) [Br−] >> [H3O+] (b) pH = 4.00 (c) [HBr] = 1.0 x 10−4M (d) [Br −] = [HBr] (e) pH = 1.00 Answer: (b)

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55. Which of the following statements is true for a 1.0 M solution of the strong acid HCl in water? (a) [Cl−] > [H3O+] (b) The pH is 1.00 (c) [H3O+] = 1.0 M (d) [HCl] = 1.0 M (e) [H3O+] = [OH−] Answer: (c)

Weak Acid pH Calculations 56. Assume a solution is prepared by adding 1.0 x 10–3 moles of a weak monoprotic acid (Ka = 2.0 x 10–4) to enough water to give a liter of solution.

If you want to compute the H3O+ ion concentration with an error of 5% or less, you can legitimately ignore: (a) both the contribution of the dissociation of water and the C term. (b) neither the dissociation of water nor the C term. (c) the dissociation of water but not the C term. (d) the C term but not the dissociation of water. (e) You can't compute the H3O+ concentration with an error of less than 5%.

Answer: (c) 57. When 1.0 x 10–5 mole of HOCl (Ka = 3.5 x 10–8) is dissolved in pure water and diluted to 1.00 L, which assumption can't be applied in the calculation of the pH of this solution? (a) that the initial concentration of HOCl is much larger than the total H3O+ ion concentration (b) that the H3O+ ion concentration from the dissociation of water can be ignored. (c) that the total H3O+ ion concentration is the sum of the concentrations from the dissociation of HOCl and water (d) all of these assumptions are valid (e) none of these assumptions are valid Answer: (b)

58. For a particular weak acid, HA, in pure water, as the initial concentration of the weak acid increases, the acid-dissociation equilibrium constant of that acid should

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(a) increase. (b) decrease. (c) remain the same. (d) increase until the ionization of water can be neglected. (e) decrease until the ionization of water can be neglected.

Answer: (c)

59. What would happen if more formic acid (HCO2H) was added to the following solution at equilibrium?

HCO2H(aq) + H2O(l)

H3O+(aq) + HCO2–(aq)

(a) [H2O] should increase. (b) [H3O+] and [HCO2–] should both increase. (c) [H3O+] and [HCO2–] should both decrease. (d) [H3O+] should increase but [HCO2–] should decrease. (e) [H3O+] should decrease but [HCO2–] should increase.

Answer: (b)

60. Which of the following equations is valid for a 0.10 M solution of formic acid, HCO2H? HCO2H(aq) + H2O(l)

HCO2–(aq) + H3O+(aq)

Ka = 1.8 x 10–4

(a) [H3O+] < [OH–] (b) [H3O+] = [HCO2H] (c) [H3O+] = [HCO2–] (d) [H3O+] = [H2O] (e) [OH–] = [H2O] Answer: (c) 61. Which of the following isn’t true for the H3O+ ion concentration in a 0.0100 M solution of acetic acid (Ka = 1.8 x 10–5)? (a) It is less than the H3O+ concentration in an 0.0100 M solution of HCl. (b) It is equal to the OAc– ion concentration. (c) It is equal to the OH– ion concentration. (d) The pH of the solution is less than 7.0. (e) None of the above are true.

Answer: (c)

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62. Calculate the approximate pH of a 0.10 M lactic acid solution.

(Ka = 8.4 x 10–4) (a) 2.0 (b) 3.0 (c) 4.0 (d) 7.0 (e) 12.0 Answer: (a)

63. Calculate the approximate pH of a 0.100 M solution of formic acid (HCO2H) in water. (HCO2H: Ka = 1.8 x 10–4) (a) 2.4 (b) 3.7 (c) 4.7 (d) 7 (e) 11.6 Answer: (a)

64. Hydrogen peroxide has been used as a bleach to change hair color, as a disinfectant to treat wounds and as a rocket fuel. It is also a weak acid.

Calculate the pH of a 0.018 M H2O2 solution. (H2O2: Ka = 2.2 x 10–12) (a) 0.6 (b) 6.7 (c) 7.3 (d) 11.7 (e) 13.4 Answer: (b)

65. A 0.10 M solution of a weak acid, HA, is found to be 1.50% ionized.

Calculate Ka for this acid. (a) 1.5 (b) 2.3 x 10–1 (c) 2.3 x 10–5 (d) 2.3 x 10–6 (e) none of the above Answer: (c)

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66. A solution is prepared by dissolving 1.0 x 10–4 moles of HOBr in water and diluting to 1.00 L. What is the pH of this solution? (HOBr: Ka = 2 x 10–9) (a) 1.3 (b) 5.3 (c) 6.3 (d) 8.7 (e) 12.7 Answer: (c)

Base pH Calculations 67. Which of the following groups contains salts that all form basic solutions in water? (a) NaNO3, NH4CN, NaOAc, NH4Cl (b) Na2CO3, KCl, NaOAc, NH4Cl (c) Na2CO3, NaF, NaOAc, NaCN (d) NaHCO3, NaF, NH4Cl, Na2SO3 (e) none of the above Answer: (c) 68. Which of the following compounds aren’t acids in water?

(I) NH4Cl (II) HNO3 (III) NH3 (IV) HI (V) K2S (a) I and II (b) II and III (c) III and V (d) I, II and IV (e) all of the above are acids in water Answer: (c)

69. Kb for methylamine (CH3NH2) is the equilibrium constant for which reaction? (a) CH3NH2 + H2O (b) CH3NH2 + H2O (c) CH3NH2 + H3O+ (d) CH3NH2 + OH(e) None of the above

CH3NH- + H3O+ CH3NH3+ + OHCH3NH3+ + H2O CH3NH- + H2O

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Answer: (b) 70. What is the Kb for methylamine (CH3NH2) if a 0.10 M solution of this weak base is 6.8% ionized? (a) 5.0 x 10-4 (b) 0.068 (c) 4.6 x 10-5 (d) 2.0 x 103 (e) 1.0 Answer: (a) 71. Which of the following won’t produce an acidic solution when dissolved in water? (a) HF (b) NH4Cl (c) CH3CO2H (d) CH3CO2Na (e) HClO4 Answer: (d)

72. Sodium fluoride, NaF (a soluble salt), is dissolved in water. Which one of the following correctly describes the relationships between the concentrations of the species present at equilibrium? (a) [F−] < [HF] and [Na+]  [F−] (b) [F−]  [Na+] and [H3O+]  [F−] (c) [F−]  [HF] and [Na+] < [F−] (d) [OH−] < [HF] and [Na+]  [F−] (e) none of these Answer: (e)

73. The salt NaF is dissolved in water. Which of the following relationships would be true for the species in solution? (a) [Na+]  [F−] >> [HF]  [H3O+] >> [OH−] (b) [Na+]  [F−] >> [HF]  [OH−] >> [H3O+] (c) [Na+] >> [F−]  [HF]  [OH−] >> [H3O+] (d) [Na+]  [HF] >> [F−]  [OH−] >> [H3O+] (e) none of these

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Answer: (b) 74. When NH4Cl is dissolved in water the following species will be present: NH4+, Cl−, OH−, H3O+, and NH3. Which of the following represents the correct relationship between the species in solution? (a) [NH4+]  [Cl−] > [OH−] > [NH3] > [H3O+] (b) [NH4+]  [Cl−] > [NH3]  [H3O+] > [OH−] (c) [NH4+]  [Cl−] > [NH3]  [OH−] > [H3O+] (d) [NH3]  [OH−]  [Cl−] > [NH4+]  [H3O+] (e) [NH3]  [H3O+]  [Cl−] > [NH4+]  [OH−] Answer: (b) 75. What is the value of Kb for the formate (HCO2-) ion if Ka for formic acid (HCO2H) is 1.8x10–4? (a) Kb = Kw x Ka (b) Kb = Ka/Kw (c) Kb = Kw/Ka (d) Kb = Kw + Ka (e) Kb = Kw - Ka Answer: (c)

76. The base-ionization constants for three bases, A(OH), B(OH), and C(OH), are 1.8 x 10–5, 4.4 x 10–4, and 7.4 x 10–4, respectively. Solutions of the bases having equal molarity have pH values that decrease in the following order: (a) A > B > C (b) B > C > A (c) C > B > A (d) A > C > B (e) B > A > C Answer: (c) 77. There are many ways of “fluoridating” water. One approach involves adding a salt of the fluoride ion, such as NaF. Calculate the pH of a 0.15 M NaF solution. (HF: Ka = 7.2 x 10–4) (a) 2 (b) 5.8 (c) 8.2 (d) 10.9

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(e) 12 Answer: (c) 78. Ammonia and the ammonium ion form a conjugate acid-base pair.

NH3(aq) + H2O(l)

NH4+(aq) + OH–(aq)

Calculate the pH of 0.10 M NH3 if Ka for the NH4+ ion is 5.6 x 10–10. (a) 4.7 (b) 5.1 (c) 5.7 (d) 8.9 (e) 11.1 Answer: (e)

79

Which of the following compounds won’t be a base in solution? (a) NH3 (b) CH3NH2 (c) Ca(OH)2 (d) NaOH (e) NH4NO3

Answer: (e) 80. Calculate the pH of a solution prepared by dissolving 1.00 x 10–2 moles of sodium hypochlorite (NaOCl) in enough water to produce a 1.00 L of solution. Hypochlorous acid is a weak monoprotic acid with Ka = 3.2 x 10–8. (a) 4.3 (b) 4.7 (c) 7.5 (d) 9.3 (e) 9.7 Answer: (e) 81. What is the OH– ion concentration in a 0.200 M NaOAc solution?

(HOAc: Ka = 1.8 x 10–5) (a) 1.1 x 10–10 (b) 5.7 x 10–10 (c) 1.8 x 10–6 (d) 1.1 x 10–5 (e) 1.3 x 10–3 11-22

Answer: (d) 82. What is the H3O+ ion concentration in a 0.10 M NH3 solution?

(Kb = 1.8 x 10–5) (a) 7.5 x 10–10 (b) 3.0 x 10–10 (c) 1.8 x 10–6 (d) 1.3 x 10–3 (e) none of these Answer: (e) 83. HA is a weak monoprotic acid with Ka = 2.4 x 10–6. What is [OH–] in a 0.300 M solution of the salt, NaA? (a) 5.8 x 10–12 (b) 2.4 x 10–10 (c) 3.5 x 10–5 (d) 6.5 x 10–5 (e) 1.6 x 10–3 Answer: (c)

84. Which of the following solutions would have a pH greater than 7.0? (a) NH4Cl (b) NaCl (c) CH3COOH (d) CH3COONa (e) HCl Answer: (d)

85. Which of the following solutions has a pH greater than 7.0? (a) 0.10 M NH4Cl (b) 0.10 M NH4NO3 (c) 0.10 M HCN (d) 0.10 M NaCN (e) none of these Answer: (d)

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86. Predict the products of the following acid-base reaction.

NH4Cl(aq) + NaSH(aq) Answer: H2S, NaCl, and NH3

87. Which of the following solutions has the highest pH? (a) 0.3 M Na2CO3 (b) 1 M HOAc (c) 0.3 M NH4Cl (d) water (e) 10–3 M HCl Answer: (a)

88. The addition of sodium formate (HCO2Na) to a solution containing formic acid (HCO2H: Ka = 1.8 x 10–4) will cause (a) the pH to increase. (b) the pH to decrease. (c) no change in pH. (d) a change in the pH, the direction of which cannot be predicted.

Answer: (a)

89. Which of the following would form the most basic 0.100 M solution? (a) NaHSO4 (b) NaHCO3 (c) NaH2PO4 (d) Na2CO3 (e) NaCl Answer: (d)

90. Which of the following compounds would give a 0.10 M solution with a pH less than 7? (a) Na2S (b) K2CO3 (c) KCN (d) NH4Cl (e) Na3PO4 Answer: (d) 11-25

91.

Which of the following is correct (assume all solutions have the same concentration)? (a) A solution of K2SO3 is more basic than a solution of KHSO3. (b) A solution of K2CO3 is less basic than a solution of KHCO3. (c) A solution of NaHSO3 is more acidic than a solution of NaHSO4. (d) A solution of Na2HPO4 is less basic than a solution of NaH2PO4. (e) All of these are correct.

Answer: (a) 92. NaHCO3 can be used to neutralize strong bases, such as NaOH. What conclusion can be drawn from the fact that the following acid-base reaction proceeds to the right as written?

HCO3–(aq) + OH–(aq)

CO32–(aq) + H2O(l)

(a) HCO3– is a stronger acid than H2O. (b) HCO3– is a stronger base than CO32–. (c) HCO3– is a stronger base than OH–. (d) CO32– is a stronger base than OH–. (e) H2O is a stronger acid than HCO3–.

Answer: (a) 93. What can we conclude from the fact that the following reaction proceeds as written?

NaNH2(s) + H2O(l) NH3(aq) + OH–(aq) (a) OH– is a stronger base than NH2–. (b) NH3 is a stronger acid than H2O. (c) NH2– is a stronger acid than H2O. (d) OH– is a weaker base than NH2–. (e) More than one of the above are true.

Answer: (d)

Buffers 94. Which of the following solutions would have the lowest pH? (HCO2H: Ka = 1.8 x 10–4; NH4+: Ka = 5.6 x 10–10) (a) 0.10 M HCO2H (b) 0.10 M NaHCO2 (c) 0.10 M NH4+ (d) a solution that is 0.10 M in both HCO2H and NaHCO2 (e) a solution that is 0.10 M in both NH4+ and NH3

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Answer: (a) 95. Which one of the following mixtures would make the best buffer? (a) HI and NaOH (b) NaCl and NaOH (c) NH3 and CH3COONa (d) HF and NaF (e) NaOH and HCl Answer: (d)

96. Which of the following solutions will show the largest change in pH when 10 mL of 1 M NaOH is added to it? (a) 10 mL of 1 M NaOH (b) 100 mL of water (c) 10 mL of 1 M NaNO3 (d) 10 mL of 0.01 M HOAc (e) 100 mL of 0.1 M HOAc and 0.1 M NaOAc Answer: (d)

97. What is the pH of a solution that is simultaneously 0.10 M in both NH3 and the NH4+ ion? (NH4+: Ka = 5.6 x 10–10) NH3(aq) + H2O(l)

NH4+(aq) + OH–(aq)

(a) 4.75 (b) 7.00 (c) 9.25 (d) The pH of this solution would be impossible to predict. (e) none of the above Answer: (c)

98. What would be the effect of doubling the NH3 concentration in the solution described in the previous question? (a) The pH would increase. (b) The buffering capacity would increase. (c) The NH4+ concentration would increase. (d) All of the above statements are true. (e) None of these statements are true.

Answer: (d)

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99. What is the [HOAc]/[OAc–] ratio in an acetic acid/sodium acetate buffer at pH = 4.9? (HOAc: Ka = 1.8 x 10–5) (a) 0.31 (b) 0.70 (c) 1.4 (d) 2.4 (e) 4.9 Answer: (b)

100. Calculate the pH of a buffer prepared by mixing 0.10 mol of sodium formate and 0.05 mol of formic acid in 1.0 L of solution. (HCO2H: Ka = 1.8 x 10–4) (a) 1.8 x 10–4 (b) 3.44 (c) 4.05 (d) 5.31 (e) none of the above Answer: (c)

101. What is the pH of a solution that is 0.01 M in CH3NH2 and 0.1 M in CH3NH3+?

CH3NH2(aq) + H2O(l)

CH3NH3+(aq) + OH–(aq)

Kb = 4.4 x 10–4

(a) 2.36 (b) 3.36 (c) 9.64 (d) 10.64 (e) 11.64 Answer: (c)

102. If equimolar amounts of NH3 and NH4Cl are dissolved in 1.00 L of water, what is the H3O+ ion concentration in this solution? (NH3: Kb = 1.8 x 10–5) (a) 1.0 x 10–14 (b) 5.5 x 10–10 (c) 1.0 x 10–7 (d) 1.8 x 10–5 (e) It cannot be calculated from the information given.

Answer: (b)

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103. When NH4Cl is added to a 0.10 M NH3 solution, the pH of this solution (a) increases. (b) decreases. (c) remains the same. (d) increases at first and then decreases. (e) decreases at first and then increases.

Answer: (b)

104. When would the pH of a solution prepared by adding sodium formate to formic acid be equal to the pKa of formic acid, HCO2H? (a) when [HCO2H] < [HCO2–] (b) when [HCO2H] = [HCO2–] (c) when [HCO2H] > [HCO2–] (d) The pH of this buffer will never equal the pKa of formic acid. (e) When the temperature is 25°C. Answer: (b)

105. Which of the following solutions would be the best buffer capable of resisting change in pH when a small amount of NaOH is added to the solution? (a) a 1.0 L solution prepared by adding 0.10 mol of carbonic acid and 0.10 mol of sodium bicarbonate (b) a 1.0 L solution prepared by adding 0.20 mol of carbonic acid and 0.20 mol of sodium bicarbonate (c) a 1.0 L solution prepared by adding 0.10 mol of hydrochloric acid and 0.10 mol of sodium hydroxide (d) a 1.0 L solution prepared by adding 0.20 mol of hydrochloric acid and 0.20 mol of sodium hydroxide (e) All of these solutions would have the same buffer capacity.

Answer: (b)

106. What reaction would take place to limit the change in pH when NaOH is added to a mixture of acetic acid (CH3CO2H) and acetate ion (CH3CO2-)? (a) H3O+(aq) + CH3CO2-(aq) → CH3CO2H(aq) + H2O(l) (b) OH-(aq) + CH3CO2H(aq) → CH3CO2-(aq) + H2O(l) (c) H2O(l) + CH3CO2-(aq) → CH3CO2H(aq) + OH-(aq) (d) H2O(l) + CH3CO2H(aq) → CH3CO2-(aq) + H3O+(aq)

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(e) none of the above Answer: (b)

Acid-Base Reactions 107. Predict the products of the following acid-base reaction.

NaOCH3(aq) + NaHCO3(aq) Answer: CH3OH and CO32–

108. Predict the products of the following acid-base reaction.

HOCl(aq) + CH3NH2(aq) Answer: CH3NH3+ and OCl–

109. Which of the following isn’t an acid-base reaction? (a) 2 HgCl2(s) + 2 CH3Li(l) Hg(CH3)2(l) + 2 LiCl2(s) (b) 2 NaCl(s) + H2SO4(aq) Na2SO4(aq) + 2 HCl(g) (c) Na2CO3(aq) + H2SO4(aq) Na2SO4(aq) + 2 H2O(l) + CO2(g) (d) H2C2O4(aq) + 2 NaOH(aq) Na2C2O4(aq) + 2 H2O(l) (e) All of the above are acid-base reactions.

Answer: (a) 110. Which of the following acid-base reactions should go more or less to completion? (a) HClO4(aq) + OCl–(aq) ClO4–(aq) + HOCl(aq) – (b) HF(aq) + I (aq) HI(aq) + F–(aq) (c) H2SO3(aq) + SO42–(aq) H2SO4(aq) + SO32–(aq) – 2– (d) H2S(aq) + 2 Cl (aq) S (aq) + 2 HCl(aq) (e) All of these reactions would go to completion, giving high yields of products.

Answer: (a)

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111. At what point in the following titration curve would the pH of the solution be equal to the pKa of the acid?

pH

(a) 0 mL (b) 50 mL (c) 95 mL (d) 100 mL (e) 120 mL Answer: (b)

Integrated Problems 112. Which of the following compounds can act as either a Brønsted acid or a Brønsted base?

(I) NaHCO3 (II) Na2CO3 (III) H2CO3 (IV) CO2 (V) H2O (a) I (b) I and II (c) II and III (d) I and V (e) V Answer: (d) 113. The dissociation of water into H3O+ and OH- is an endothermic process. As the temperature of pure water is raised, the degree to which water dissociates ________ and the pH ________. (a) increases, increases (b) decreases, increases (c) decreases, stays the same (d) decreases, decreases (e) increases, decreases

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Answer: (e)

114. Rank the following 0.100 M solutions in order of increasing pH (from lowest pH to the highest).

HNO3, KI, HIO3, HIO4 (a) HNO3 < KI < HIO3 < HIO4 (b) HNO3 < HIO3 < HIO4 < KI (c) HNO3 < HIO4 < HIO3 < KI (d) KI < HNO3 < HIO3 < HIO4 (e) none of these Answer: (c) 115. “Cocaine”, as it is shown on television, is actually a salt known as cocaine hydrochloride, [C17H22NO4+][Cl–], which forms white crystals that stupid people snort through straws to make them even more stupid. Recently, a lot of attention has been paid to a frightening substance known as “crack”, which is formed by treating cocaine with sodium hydrogen carbonate, NaHCO3. “Crack” is a neutral compound with the formula C17H21NO4. Which of the following statements is true? (a) “Crack” is the conjugate base of cocaine hydrochloride. (b) “Cocaine” contains the conjugate acid of “crack”. (c) When “cocaine” is treated with NaHCO3, the bicarbonate ion acts as a base, removing an H+ ion. (d) “Crack” can be “smoked” because it is a covalent compound that becomes reasonably volatile when heated, but “cocaine” must be snorted because it is a salt that does not easily boil even when heated. (e) All of the above are true.

Answer: (e)

116. Rank the following in order of increasing pH of 0.1M solutions.

NaI, HI, HIO3, NaIO3 Answer: HI < HIO3 < NaI < NaIO3

Use the following acid-dissociation equilibrium constants for questions 117-119.

HA: Ka = 1.00 x 10–4 HC: Ka = 1.00 x 10–10

HB: Ka = 1.00 x 10–7 HD: Ka = 1.00 x 10–11

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117. Solutions of each acid are prepared in which the initial concentration of the acid is 0.1000 M. Which of the four solutions will be the most acidic? (a) HA (b) HB (c) HC (d) HD (e) Not enough information is provided to answer this question.

Answer: (a)

118. 0.5 M solutions of NaA, NaB, NaC, and NaD are prepared. Which solution is the most basic? (a) NaA (b) NaB (c) NaC (d) NaD (e) All of the solutions have the same pH. Answer: (d)

119. A solution is prepared with initial concentrations of HA and NaA of 0.50 M and 1.00 M respectively. What is the pH of the solution? (a) 3.70 (b) 4.00 (c) 4.30 (d) 9.70 (e) 10.30 Answer: (c)

Diprotic Acids and Bases 120. What is the pH of a 0.10 M succinic acid, H2Sc, solution if we assume that H2Sc is a weak acid that undergoes stepwise dissociation?

H2Sc(aq) + H2O(l) HSc–(aq) + H2O(l)

H3O+(aq) + HSc–(aq) H3O+(aq) + Sc2–(aq)

(a) 2.1 (b) 2.6 (c) 3.3 (d) 4.2 (e) 5.6

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Ka1 = 6.9 x 10–5 Ka2 = 2.5 x 10–6

Answer: (b) Use the following information to answer questions 121-123.

H2Se: Ka1 = 1.0 x 10–4, Ka2 = 1.0 x 10–10 121. In a 0.10 M solution of H2Se, the Se2– concentration is: (a) [Se2–] = 0.10 x Ka2 = 1.0 x 10–11 (b) [Se2–] = Ka2 = 1.0 x 10–10 (c) [Se2–] = (0.10 x Ka2)1/2 = 3.2 x 10–6 (d) [Se2–] = (0.10 x Ka1)1/2 = 3.2 x 10–3 (e) none of the above Answer: (b) 122. The base-ionization constant for the Se2– ion would be: (a) Kb1 = Ka1 = 1.7 x 10–4 (b) Kb1 = Ka2 = 1.0 x 10–10 (c) Kb1 = Kw/Ka1 = 5.9 x 10–11 (d) Kb1 = Kw/Ka2 = 1.0 x 10–4 (e) none of the above Answer: (d) 123. If sodium selenide is added to pure water, we would expect the pH to: (a) increase. (b) decrease. (c) remain the same. (d) change in a direction that can't be predicted from this information.

Answer: (a) 124. Sulfurous acid (H2SO3) is an acid and the sulfite ion (SO32-) is a base.

Which equation correctly describes the relationship between the equilibrium constants for these compounds? (a) Ka1 x Kb2 = Kw (b) Ka1 x Ka2 = Kw (c) Ka2 x Kb2 = Kw (d) Ka1 x Ka2 = 1 (e) Ka1 x Kb2 = 1 Answer: (a)

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125. Succinic acid (HO2CCH2CH2CO2H) is an intermediate in the Kreb's cycle used to “burn” carbohydrates, proteins and lipids. The formula for succinic acid can be abbreviated as H2Sc. H2Sc(aq) + H2O(l) HSc–(aq) + H2O(l)

H3O+(aq) + HSc–(aq) H3O+(aq) + Sc2–(aq)

Ka1 = 6.9 x 10–5 Ka2 = 2.5 x 10–6

Which of the following assumptions may not be valid for solutions of succinic acid because the difference between Ka1 and Ka2 for H2Sc is very small? (a) [H3O+]  Kw/[OH–] (b) [H2Sc] < [Sc2–] (c) [HSc– ]  [H3O+] (d) [H2Sc]  CH2Sc (e) None of these assumptions are legitimate for this solution.

Answer: (c) 126. Using data from the previous problem, what fraction of the HSc – formed in the first step of the dissociation of succinic acid goes on to dissociate to form Sc2– ions in the second step, and what does this lead us to conclude? (a) None of the HSc– ions dissociate further, and the stepwise dissociation assumption is legitimate. (b) Less than 1% of the HSc– ion dissociate to form Sc2– ions, and the stepwise dissociation assumption is legitimate. (c) Between 1 and 5% of the HSc– ions dissociate, and the stepwise dissociation assumption is legitimate. (d) Between 5 and 20% of the HSc– ions dissociate, and the stepwise dissociation assumption is legitimate. (e) Between 5 and 20% of the HSc– ions dissociate, and the stepwise dissociation assumption is not legitimate.

Answer: (b)

127. What is the pH of a 0.10 M H2CO3 solution?

(H2CO3: Ka1 = 4.2 x 10–7, Ka2 = 4.8 x 10–11) (a) 3.1 (b) 3.7 (c) 6.3 (d) 6.8 (e) none of the above Answer: (b) 11-35

128. Use the relationship between Ka1 and Ka2 for H2SO3 and Kb1 and Kb2 for the SO32– ion to predict whether a 0.10 M solution of the HSO3– (bisulfite) ion would be acidic, basic or neutral. (H2SO3: Ka1 = 1.7 x 10–2, Ka2 = 6.4 x 10–8) (a) The solution would be strongly acidic, pH < 2. (b) The solution would be moderately acidic, pH  4. (c) The solution would be very close to neutral, pH  7. (d) The solution would be slightly basic, pH  8. (e) The solution would be strongly basic, pH > 12.

Answer: (b)

129. A diprotic acid, H2A, has the following dissociation constants: Ka1 = 2.1 x 10–7 and Ka2 = 4.3 x 10–13. The H3O+ ion concentration when this acid is dissolved in water is (a) roughly equal to the H2A concentration. (b) much larger than the H2A concentration. (c) much larger than the HA– concentration. (d) approximately equal to the HA– concentration. (e) much less than the HA– concentration.

Answer: (d) 130. For a weak diprotic acid, H2A, for which Ka1 = 2.1 x 10–7 and Ka2 = 4.3 x 10–13 the A2– ion concentration at equilibrium will be: (a) approximately equal to the initial concentration of H2A. (b) roughly equal to Ka2. (c) roughly equal to the HA– concentration. (d) much larger than the HA– concentration. (e) approximately equal to the H3O+ concentration.

Answer: (b)

131. Which of the following equations can be used to calculate Kb1 and Kb2 for Na2Gly from Ka1 and Ka2 for H2Gly? (a) Kb1 = Kw x Ka1 (b) Kb1 = Kw/Ka1 (c) Kb1 = Ka1/Kw (d) Kb1 = Kw x Ka2 (e) Kb1 = Kw/Ka2

and and and and and

Kb2 = Kw x Ka2 Kb2 = Kw/Ka2 Kb2 = Ka1/Kw Kb2 = Kw x Ka1 Kb2 = Kw/Ka1

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Answer: (e) 132. For H2CO3, Ka2 = 4.7 x 10–11. Which of the following numbers is closest to the pH of a 0.200 M Na2CO3 solution? (Hint: Only the first step in the hydrolysis of the CO32– ion need be considered in this calculation.) (a) 5.9 x 10–3 (b) 2.2 (c) 2.48 (d) 11.6 (e) 11.8 Answer: (e)

133. KHS can act as a weak acid or a weak base.

HS–(aq) + H2O(l) HS–(aq) + H2O(l)

S2–(aq) + H3O+(aq) H2S(aq) + OH–(aq)

Ka2 = 1.3 x 10–13 Kb2 = 1.0 x 10–7

Which of the following statements is correct? (a) A solution of 0.1 M K2S and 0.1 M KHS will be acidic. (b) A solution of 0.01 M H2S will have a pH = 2. (c) A solution of 0.1 M KHS will be weakly basic. (d) H2S is over 90% ionized in water solution. (e) The S2– ion concentration in a 0.1 M KHS solution does not depend on pH. Answer: (c)

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Chapter 12: Oxidation-Reduction Reactions E = E° - (0.0257/n) ln Q 1 F = 96,500 C 23 N = 6.02 x 10 particles per mole 1 amp = 1 C/s

Determining Oxidation Numbers 1.

What is the oxidation number of sulfur in the thiosulfate ion, S2O32−? (a) +2 (b) +4 (c) +5 (d) +6 (e) +8

Answer: (a)

2.

What is the oxidation number of the carbon in bold print in the following structure?

O H3C

(a) –1 (b) +2 (c) –3 (d) +4 (e) none of these

C

CH3

Answer: (b)

3.

Determine the oxidation number of the bold face carbon atom in pyruvic acid: H3C

O

O

C

C

(a) −2 (b) +2 (c) +3 (d) +4 (e) none of these Answer: (b)

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O

H

4.

Determine the oxidation number for bromine in HBrO4. (a) +6 (b) +7 (c) +8 (d) −6 (e) −7

Answer: (b)

5.

Determine the oxidation state of the carbon atom in bold print: CH3CH2F (a) 0 (b) +1 (c) +2 (c) +3 (e) −1

Answer: (e)

6.

Determine the oxidation number of arsenic in the anion, HAsO42−. (a) +5 (b) +6 (c) +7 (d) +8 (e) none of these

Answer: (a)

7.

Determine the oxidation state of the carbon atom shown in bold in the following compound: O

(a) −1 (b) 0 (c) +1 (d) +4 (e) none of these

H3C

C

O

Answer: (e)

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CH3

Recognizing Oxidation-Reduction Reactions The information given here is for the next two questions.

One of the most common alcohol breathalyzers used by police for measuring blood alcohol levels in suspected drunk drivers utilizes the oxidation-reduction reaction shown below between ethanol (CH3CH2OH) and the bright orange dichromate ion (Cr2O72-). The chromium is converted to Cr3+ that is green and the ethanol is converted into acetic acid (CH3CO2H). The orange to green color change is used to estimate the blood alcohol content.

2Cr2O72- + 3CH3CH2OH + 16H+ → 4Cr3+ + 3CH3CO2H + 11H2O 8. The carbon in bold in the ethanol has an oxidation number of __ while the carbon in bold in acetic acid has an oxidation number of ___. (a) –1, +3 (b) +1, +2 (c) –2, +3 (d) +2, -3 (e) none of the above Answer: (a)

9. Which of the following statements about this reaction is correct? (a) H+ is the oxidizing agent and dichromate is the reducing agent. (b) Dichromate is oxidized and ethanol is reduced. (c) Ethanol is the reducing agent and dichromate is the oxidizing agent. (d) This is not an oxidation-reduction reaction. (e) None of the above statements is correct.

Answer: (c)

10. How many electrons are transferred in the following oxidation-reduction reaction?

3 Sn2+(aq) + Cr2O72–(aq) + 14 H+ H2O(l) (a) 2 (b) 6 (c) 12 (d) 14 (e) none of these

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3 Sn4+(aq) + 2 Cr3+(aq) + 7

Answer: (b) 11. Which of the following isn’t an example of an oxidation-reduction reaction? (a) Ca3P2(s) + 6 H2O(l) → 3 Ca2+(aq) + 6 OH–(aq) + 2 PH3(g) (b) 2 PH3(g) + 4 O2(g) → H3PO4(s) (c) P4(s) + 5 O2(g) → P4O10(s) (d) 6 Ca(s) + P4(s) → 2 Ca3P2(s) (e) All of the above are oxidation-reduction reactions.

Answer: (a) 12. Which of the following isn’t an example of an oxidation-reduction reaction? (a) H2(g) + Cl2(g) → 2 HCl(g) (b) Ag+(aq) + 2 NH3(g) Ag(NH3)2+(aq) (c) Hg2Cl2(s) + NH3(g) → Hg(s) + HgNH2Cl(s) (d) 2 Mg(s) + O2(g) → 2MgO(s) (e) Cl2(g) + 2 Br–(aq) → 2 Cl–(aq) + Br2(l) Answer: (b)

Voltaic Cells 13. A voltaic cell is constructed with a Pt wire in an aqueous solution of Br2/Br −(1M) serving as the cathode and a silver wire in a 1M solution of Ag+ serving as the anode. The potential of the cell was measured to be +0.287 V. The standard reduction potential for Br2/Br− is 1.087 V. (I) Write the half-reaction for the reduction taking place in the cell.

(II) Write the half-reaction for the oxidation taking place in the cell.

(III) Write the overall reaction for the oxidation-reduction reaction taking place in the cell.

(IV) Calculate the potential for the Ag/Ag+ half-reaction taking place in the cell.

(V) Determine the standard reduction potential, E°, for the half-reaction below.

Ag+(aq) + e− → Ag(s) Answer: (I) Br2 + 2e− → 2 Br − (II) Ag → Ag+ + e− (III) Br2 + 2Ag → 2 Br − + 2 Ag+ (IV) −0.800V

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(V) +0.800V 14. What will be the coefficients of Ce4+ and Cl−, respectively when the following reaction is balanced?

Ce4+(aq) + Cl−(aq) → Cl2(aq) + Ce3+(aq) (a) 1,1 (b) 1,2 (c) 2.1 (d) 2.2 (e) 3,1 Answer: (d)

15. Use a table of standard reduction potentials to determine which of the following statements is true for the electrochemical cell diagrammed below.

NaCl Al( s)

Fe(s)

Al

3+

Fe

3+

a) The Al is the cathode. b) Electrons move from the Al electrode to the Fe electrode. c) The mass of the Al electrode decreases. d) Both (b) and (c) are true. e) Both (a) and (c) are true.

Answer: (d)

Standard Cell Potentials 16. Use the table of electrode potentials to determine which of the following reactions isn’t spontaneous under standard conditions. (a) Zn(s) + 2 H+(aq) → Zn2+(aq) + H2(g) (b) 2 Ag(s) + Zn2+(aq) → 2 Ag+(aq) + Zn(s)

12-5

(c) Cl2(aq) + 2 Fe2+(aq) → 2 Cl–(aq) + 2 Fe3+(aq) (d) 2 Al(s) + 3/2 O2(g) + 6 H+(aq) → 2 Al3+(aq) + 3 H2O(l) (e) Mg(s) + Cl2(g) → MgCl2(s) Answer: (b)

17. Consider the following generalized half-reactions.

Aox + e– Box + e–

Ared Bred

If E° for the half-reaction involving A is positive and if E° for the half-reaction involving B is negative, then at standard conditions: (a) Aox will reduce Box (b) Ared will reduce Box (c) Bred will reduce Aox (d) Bred will reduce Ared (e) no reaction will occur Answer: (c) 18. The potential of a cell at standard state conditions can be best be described as: (a) Eo (b) 0 volts (c) negative (d) more than one volt (e) None of the above Answer: (a)

19. What is the potential for a voltaic cell produced from the following halfreactions, if all ions are present at 1.0 M concentrations?

Oxidation: Al Al3+ + 3 e– Reduction: Zn2+ + 2 e– Zn (a) -3.59 V (b) -0.943 V (c) 0.943 V (d) 1.12 V (e) 3.59 V Answer: (c)

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20. What is the standard cell potential for the following redox reaction?

2 Fe3+(aq) + 2 I–(aq) → 2 Fe2+(aq) + I2(aq) (a) -1.30 V (b) -0.24 V (c) 0.24 V (d) 1.30 V (e) none of the above Answer: (c)

21. What is the magnitude of the standard-state cell potential for the following redox reaction?

2 Al(s) + 3 Pb2+(aq) → 2 Al3+(aq) + 3 Pb(s) (a) 1.58 V (b) 1.83 V (c) 3.03 V (d) 3.790 V (e) 4.866 V Answer: (a)

Oxidizing and Reducing Agents 22. Oxidation-reduction reactions are often written with double arrows indicating that the reaction can go in either direction depending on the experimental conditions. For the following reaction identify the oxidizing agent on both sides of the chemical equation.

S2O82–(aq) + Zn(s)

Zn2+(aq) + 2 SO42–(aq)

(a) S2O82– and SO42– (b) Zn and SO42– (c) Zn and Zn2+ (d) S2O82– and Zn2+ (e) none of these Answer: (d)

23. Which statement correctly describes the following oxidation-reduction reaction?

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5Cr3+(aq) + 3 MnO4−(aq) + 8 H2O (l) → 5 CrO42−(aq) + 3 Mn2+(aq) + 16 H+(aq) (a) MnO4− and H2O are both reduced. (b) Cr3+ is the reducing agent. (c) Cr3+ and H2O are both reducing agents. (d) MnO4− is oxidized. (e) None of these are correct.

Answer: (b)

24. Which of the following is true about the reaction below?

2VO2+ (aq) + Br2(s) + 2 H2O(l) → 2VO2+(aq) + 2 Br −(aq) + 4H+ (a) Br2 is acting as a reducing agent. (b) VO2+ is acting as a weak reducing agent. (c) VO2+ is oxidized. (d) H2O is reduced. (e) H+ is acting as a weak reducing agent.

Answer: (c)

25. Identify each of the requested compounds for the following redox reaction.

2 MnO4−(aq) + 16 H+(aq) + 10 Cl− (aq) → 2 Mn2+(aq) + 8 H2O + 5 Cl2(g) (I) oxidizing agent (II) reducing agent (III) conjugate oxidizing agent (IV) conjugate reducing agent Answer: MnO4−; Cl−; Cl2; Mn2+

26. Which statement correctly describes the following reaction?

O2(g) + 4 H+(aq) + 4 Cl−(aq) → 2 H2O(l) + 2 Cl2(g) (a) O2 is a reducing agent. (b) H+ and Cl− are oxidized. (c) H+ is reduced. (d) H+ is a reducing agent. (e) Cl− is a reducing agent.

Answer: (e)

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27. For the following reaction, identify the oxidizing agents in the forward and reverse directions, respectively..

P4(s) + HNO3(aq) → H3PO4(aq) + NO(g) (a) P4 and H3PO4 (b) P4 and NO (c) P4 and HNO3 (d) HNO3 and H3PO4 (e) HNO3 and NO Answer: (d)

28. For the reaction: 3 Sn2+(aq) + Cr2O72–(aq) + 14 H+(aq) → 3 Sn4+(aq) + 2 Cr3+(aq) + 7 H2O(l) which of the following statements is true? (a) Both Sn2+ and H+ are oxidizing agents. (b) Cr2O72– is the oxidizing agent. (c) Sn2+ is reduced. (d) The acid isn’t important to the reaction. (e) None of the above are true.

Answer: (b)

Relative Strengths of Oxidizing and Reducing Agents 29. Use a table of standard reduction potentials to determine which of the following is the strongest reducing agent. (a) Co (b) Ca2+ (c) Co3+ (d) Cr (e) Au Answer: (d) 30. Given the following half-cell reduction potentials, determine which of the following is the strongest reducing agent.

In3+

3e− → In

+ Pd2+ + 2e− → Pd

E −0.338 V +0.987 V

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(a) In3+ (b) In (c) Pd2+ (d) Pd (e) Cannot be determined from cell potentials alone Answer (b)

31. Use the following half reactions and accompanying standard reduction potentials to determine the best reducing agent.

Ga3+ + 3e− → Ga Be2+ + 2e− → Be Pd2+ + 2e− → Pd

ERed −0.560 V −1.70 V +0.987 V

(a) Ga3+ (b) Be2+ (c) Be (d) Pd2+ (e) Pd Answer: (c)

32. Given the following half-cell reduction potentials, determine which metal is the best reducing agent.

Co2+ + 2e− → Co Pd2+ + 2e− → Pd Cd2+ + 2e− → Cd

E = −0.28 V E = +0.987 V E = −0.403 V

(a) Co (b) Pd (c) Cd (d) None would be a good reducing agent.

Answer: (c) 33. Given the following redox reactions and cell potentials, rank from weakest to strongest the relative strengths of Al, Fe, and Cr as reducing agents.

Al(s) + Cr3+(aq) → Al3+(aq) + Cr(s) Fe(s) + Cr3+(aq) → Fe3+(aq) + Cr(s) (a) Cr < Fe < Al (b) Al < Fe < Cr (c) Al < Cr < Fe (d) Fe < Cr < Al 12-10

Eo = +0.966 V Eo = −0.70 V

(e) Fe < Al < Cr Answer: (d) 34. Use a table of standard reduction potentials to determine which of the following is the strongest oxidizing agent. (a) H2O2 in base (b) H2O2 in acid (c) O2 in acid (d) CrO42– in acid (e) Br2 Answer: (b) 35. Use a table of standard reduction potentials to determine which is the strongest oxidizing agent among the following. (a) Cl2 gas at 1 atm pressure (b) O2 gas at 1 atm pressure in contact with 1 M acid (c) F2 gas at 1 atm pressure (d) Ag metal (e) H2O2 in acid solution Answer: (c) 36. Which of the following pairs of ions can't coexist in aqueous solution under standard-state conditions because a spontaneous redox reaction occurs? (a) Sn4+ and Fe3+ (b) Sn4+ and Fe2+ (c) Sn2+ and Fe3+ (d) Sn2+ and Fe2+ (e) None of these pairs of ions can coexist.

Answer: (c)

For questions 37-38, use the following information.

Mg2+ + 2 e– Cr2+ + 2 e– 2 H+ + 2 e – Pd2+ + 2 e– Ce4+ + e–

Mg Cr H2 Pd Ce3+

E = -2.37 V E = -0.91 V E = 0.00 V E = 0.987 V E = 1.61 V

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37. Which is the strongest oxidizing agent? (a) Ce4+ (b) Ce3+ (c) H+ (d) Cr2+ (e) Mg Answer: (a) 38. Which of the following reagents should react with H+ to produce H2? (a) either Mg2+ and Cr2+ (b) either Pd and Cr2+ (c) either Pd2+ (d) either Mg and Cr (e) Mg, Cr, and Pd Answer: (d)

39. Which of the following pairs of ions cannot coexist in solution because a spontaneous redox reaction occurs? (a) Cu2+ and Zn2+ (b) Fe3+ and Hg22+ (c) Cu2+ and I– (d) Al3+ and S2– (e) All of these pairs of ions can coexist.

Answer: (b)

40. Cobalt(III) oxide reacts with hydrogen gas to form cobalt metal and water.

Co2O3(s) + 3 H2(g) → 2 Co(s) + 3 H2O(g) What does this tell you about the relative strength of the oxidizing and reducing agents in this reaction? (a) The Co3+ ion is a stronger reducing agent than cobalt metal. (b) Cobalt metal is a better reducing agent than water. (c) Cobalt metal is a weaker reducing agent than hydrogen. (d) The Co3+ ion is a weaker oxidizing agent than water. (e) Statements (a) through (d) above are all false.

Answer: (c)

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Electrochemical Cells at Nonstandard Conditions: The Nernst Equation 41. The equilibrium constant at 298 K for the following reaction is 1.2 x 10 5.

H2(g) + Sn4+(aq) → 2 H+(aq) + Sn2+(aq) What is the value of E° for the reaction? (a) -0.15 V (b) 0.15 V (c) 0.30 V (d) 0.35 V (e) 0.45 V Answer: (b) 42. The equilibrium constant at 25 oC for the following reaction is 2 x 10-3.

H2(g) + 2Ru3+(aq) → 2 H+(aq) + 2Ru2+(aq) What is the value of E° for the reaction? (a) –0.16 V (b) 0.16 V (c) 1.00 V (d) 0.00 V (e) –0.08 V Answer: (e)

43. Calculate the value of the Ksp for CdS at 25 °C from the following data.

CdS + 2 e– Cd2+ + 2 e–

Cd + S2– Cd

E° = -1.21 V E° = -0.40 V

(a) 3 x 10–55 (b) 4 x 10–28 (c) 2 x 10–14 (d) 3 x 1027 (e) 6 x 10 Answer: (b)

44. There are two half-reactions for the reduction of O2 to H2O.

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O 2 + 4 H+ + 4 e – O2 + 2 H2O + 4 e–

2 H2O 4 OH–

E° = 1.229 V E° = 0.401V

Use these half-reactions to calculate the ionization constant for water, Kw, at 25 °C. H+(aq) + OH–(aq)

H2O(l)

Kw = ?

Answer: 1 x 10–14

45. The half-reaction reduction potentials for the mercury(I) and silver(I) ions are given below.

Hg22+ + 2 e– Ag+ + e–

2 Hg Ag

E° = 0.7961 V E° = 0.7996V

Calculate the equilibrium constant for the following reaction under standard-state conditions at 298 K: Hg22+(aq) + 2 Ag(s)

2 Hg(l) + 2 Ag+(aq)

Answer: 0.761

46. Calculate the complex dissociation equilibrium constant (Kd) for the Cd(NH3)42+ complex from the following data at 298K. Cd(NH3)42+ + 2 e– Cd2+ + 2 e– Cd

Cd + 4 NH3

E° = -0.597 V E° = -0.403 V

(a) 2.7 x 10–7 (b) 0.19 (c) 6.6 (d) 3.6 x 106 (e) none of the above Answer: (a)

Electrolysis and Faraday’s Law 47. Calculate the weight of sodium metal that would be produced by the electrolysis of molten sodium chloride for 1.00 hour with a 10.0-amp current.

Answer: 8.58 g

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48. What is the ratio of the weight of Cl2 produced at the anode to the weight of Al produced at the cathode when 5 moles of electrons are passed through a molten sample of AlCl3? (a) about 2 g Cl2 per 1 g Al (b) about 4 g Cl2 per 1 g Al (c) about 8 g Cl2 per 1 g Al (d) about 16 g Cl2 per 1 g Al (e) none of the above Answer: (b)

49. Calculate the amount of aluminum produced in 1.00 hour by the electrolysis of molten AlCl3 if the current is 10.0 A. Answer: 3.35 g Al

50. Suppose a beer can weighs 40.0 g. Determine the amount of time in hours that a current of 100.0 amp need to be passed through a molten AlF 3 electrolysis cell to produce enough Al to replace a discarded beer can.

Sketch the electrolysis cell, labeling the electrodes and showing the direction of electron flow in the external circuit as part of your answer.

Answer: 1.2 hours

51. Which of the following electrolysis processes will produce the largest volume of Cl2 gas at STP? (a) passing 35 amps for 20 hours through an aqueous Na2SO4 solution (b) passing 10 amps for 2.69 hours through a 1 M NaCl solution (c) passing 2 moles electrons through a 1 M CaCl2 solution (d) passing 1.5 moles electrons through a 5 M CaCl2 solution (e) passing 1.4 moles electrons through a 5 M AlCl3 solution Answer: (c)

52. A current of 1.5 amps is applied to a 1.00 L solution of 0.100 M hydrochloric acid for 1.0 hr. What is the pH of the solution after electrolysis is complete?

Answer: 1.4

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53. A molten sample of TiCl4 was electrolyzed for 10.0 hours at 12 amps. What is the ratio of the weight of Cl2 produced compared to that of Ti? (a) 0.34 g Cl2/1.00 g Ti (b) 0.68 g Cl2/1.00 g Ti (c) 0.74 g Cl2/1.00 g Ti (d) 1.48 g Cl2/1.00 g Ti (e) 2.96 g Cl2/1.00 g Ti Answer: (e)

54. An electric current is passed through a solution of CuSO4(aq) producing Cu(s) at the cathode and O2(g) at the anode. If 3.48 L of O2(g) measured at STP is produced at the anode, how many grams of Cu(s) must have been deposited on the cathode? (a) 3.48 g (b) 4.93 g (c) 9.87 g (d) 19.7 g (e) none of these Answer: (d)

55. What is the ratio by weight of Br2 to Cr if a molten sample of CrBr2 is electrolyzed for 4.00 hr. at 10.0 amps? (a) 1.05 g Br2/1 g Cr (b) 1.54 g Br2/1 g Cr (c) 2.05 g Br2/1 g Cr (d) 3.07 g Br2/1 g Cr (e) 4.61 g Br2/1 g Cr Answer: (d)

56. How much Cl2 gas would be collected when a 2.00 M NaCl(aq) solution is electrolyzed for 2.00 hours with a current of 15.0 amps? (a) 0.0220 g (b) 39.7 g (c) 79.4 g (d) 159 g (e) none of the above Answer: (b)

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57. A current of 10.0 amps over a period of 3.00 hr is passed through a solution of molten KCl. What weight of K metal is produced? (a) 0.365 g (b) 0.729 g (c) 21.9 g (d) 43.8 g (e) 87.5 g Answer: (d)

58. A current of 2.0 amps passed through a molten solution of the salt MCl2 for a period of 30 minutes produces 1.32 g of Cl2 gas and 0.453 g of M. Which of the following metals is present in this compound? (a) Mg (b) Ca (c) Zn (d) Sr (e) Hg Answer: (a)

59. What is the oxidation state of the osmium atom in an unknown salt if 26.7 grams of osmium plate out when a current of 15.0 amps is passed through a solution of this salt for 1.00 hour? (a) Os+ (b) Os2+ (c) Os3+ (d) Os4+ (e) Os5+ Answer: (d)

60. What is the oxidation state of the cerium atom in CexZy if a current of 1.2 amps for 3.0 hours deposits 4.70 g of Ce at the cathode? (a) 0 (b) +1 (c) +2 (d) +3 (e) +4 Answer: (e)

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61. What is the oxidation state of the tin atom in an unknown salt if 14.8 grams of tin are plated out when a current of 40 amps is passed through a molten solution of this salt for 10 minutes? (a) +1 (b) +2 (c) +3 (d) +4 (e) +6 Answer: (b)

Electrolysis of Molten and Aqueous Salts 62. Which of the following is the correct half-cell reaction for the anode process in the electrolysis of an aqueous solution of potassium sulfate? (a) SO42–(aq) → SO2(g) + O2(g) + 2 e– (b) H2O(l) → 1/2 O2(g) + 2 H+(aq) + 2 e– (c) H2O(l) + 1/2 O2(g) + 2 e– → 2 OH–(aq) (d) SO42–(aq) + 4 H+(aq) + 2 e– → SO2(g) + 2 H2O(l) (e) 3 H2O(l) + e– → 1/2 H2(g) + 3 OH–(aq) + 2 H+(aq) Answer: (b) 63. Which of the following statements about the electrolysis of an aqueous solution of NaCl is true? (a) O2 is liberated at the cathode. (b) H2 is liberated at the cathode. (c) Na is liberated at the cathode. (d) Cl2 is liberated at the cathode. (e) The decomposition potential must not be exceeded.

Answer: (b)

64. In the electrolysis of NaBr(aq), which of the following products is least likely to be found? (a) H2(g) (b) O2(g) (c) Br2(l) (d) Na(s) (e) NaOH(aq) 12-18

Answer: (d)

Balancing Oxidation-Reduction Equations 65. What is the correct coefficient for Na2MnO4 in the balanced equation for the air oxidation of manganese dioxide to sodium manganate in the presence of sodium carbonate?

MnO2(s) + Na2CO3(aq) + O2(g) → CO2(g) + Na2MnO4(aq) (a) 1 (b) 3/2 (c) 2 (d) 5/2 (e) 3 Answer: (c) 66. What is the coefficient for sulfur in the following redox reaction?

S(s) + KClO3(s) → SO2(g) + KCl(s) + bang! (a) 1 (b) 2 (c) 3 (d) 4 (e) 5 Answer: (c)

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Redox Reactions in Acidic Solutions 67. The following equation, when balanced, involves a total of ___ electrons.

ClO3–(aq) + I2(aq) + H2O(l) → IO3–(aq) + Cl–(aq) + H+(aq) (a) 5 (b) 6 (c) 12 (d) 24 (e) 30 Answer: (e) 68. Complete and balance the following equation: PbO2(s) + Cl–(aq) + H+(aq) → Pb2+(aq) + Cl2(aq) Answer: PbO2(s) + 2 Cl–(aq) + 4 H+(aq) → Pb2+(aq) + Cl2(aq) + 2 H2O(l)

Problems 66-68 refer to the following reaction in acid solution.

CuS(s) + NO3–(aq)

H+

Cu2+(aq) + SO42–(aq) + NO(g)

69. In the balanced half-reaction for the NO3– ion, how many electrons are involved? (a) 1 (b) 2 (c) 3 (d) 4 (e) none of these Answer: (c) 70. In the complete balanced overall reaction, what is the coefficient of the NO 3– ion? (a) 3 (b) 4 (c) 5 (d) 6 (e) none of these Answer: (e)

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71. Which of the following isn’t true? (a) The reducing agent is CuS. (b) The CuS is oxidized. (c) The NO3– ion is the oxidizing agent. (d) The NO3– ion is reduced. (e) The oxidation number of the copper changes from 0 to +2.

Answer: (e) 72. Write a balanced chemical equation for the following reaction.

MnO2(s) + PbO2(s) + H+(aq) → MnO4–(aq) + Pb2+(aq) Answer: 2 MnO2(s) + 3 PbO2(s) + 4 H+(aq) → 2 MnO4–(aq) + 3 Pb2+(aq) + 2 H2O(l)

Questions 73 and 74 refer to the following incomplete, unbalanced equation

Cr2O72–(aq) + NO(g)

H+

Cr3+(aq) + NO3–(aq)

73. How many electrons are in the half-reaction involving Cr2O72–? (a) 2 (b) 3 (c) 4 (d) 5 (e) 6 Answer: (e)

74. What is the coefficient of the NO in the balanced equation? (a) 1 (b) 2 (c) 3 (d) 4 (e) 6 Answer: (b)

75. Write a balanced chemical equation for the following reaction, which can be used to standardize aqueous permanganate ion solutions.

H2C2O4(aq) + MnO4–(aq) + H+(aq) → CO2(g) + Mn2+(aq) 12-21

Answer: 5 H2C2O4(aq) + 2 MnO4–(aq) + 6 H+(aq) → 10 CO2(g) + 2 Mn2+(aq) + 8 H2O(l)

76. Balance the following redox reaction. Clearly show the two balanced halfreactions and indicate which represents oxidation and which reduction.

H2O2(aq) + NO(g) + H+(aq) → NO3–(aq) Answer: 3 H2O2(aq) + 2 NO(g) → 2 NO3–(aq) + 2 H2O(l) + 2 H+(aq) 2 H2O(aq) + NO(g) → NO3–(aq) + 4 H+(aq) + 3 e–(aq) : Oxidation H2O2(aq) + 2 H+(aq) + 2 e–(aq) → 2 H2O(l) : Reduction

77. Balance the following oxidation-reduction equation.

HI(aq) + HNO3(aq) → NO(g) + I2(aq) Answer: 6 HI(aq) + 2 HNO3(aq) → 2 NO(g) + 3 I2(aq) + 4 H2O(l)

78. Balance the following oxidation-reduction equation.

CrO42–(aq) + HSnO2–(aq)

H+

HSnO3–(aq) + CrO2–(aq)

Answer: 2CrO42–(aq) + 3HSnO2–(aq) + 2H+(aq) → 3HSnO3–(aq) + 2CrO2–(aq) + H2O(l)

Redox Reactions in Basic Solutions Questions 79-81 refer to the following reaction which occurs in basic solution: OCl–(aq) → ClO2–(aq) + Cl2(aq) 79. How many OCl– ions are consumed in the balanced equation for this reaction? (a) 2 (b) 3 (c) 4 (d) 5 (e) 6 Answer: (b)

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80. How many OH– ions are involved in the simplest balanced equation for this reaction? (a) 0 (b) 1 (c) 2 (d) 3 (e) 4 Answer: (c) 81. What is the net effect on the OH– ion concentration in this reaction? (a) The OH– ion concentration increases. (b) The OH– ion concentration decreases. (c) The OH– ion concentration remains the same. (d) There is no way to predict what happens to the OH– ion concentration in this reaction.

Answer: (a)

Questions 82 and 83 refer to the following reaction which occurs in basic solution.

CrO42– + PH3 → Cr(OH)4– + P4 82. Which of the following are oxidizing agents in the forward and reverse directions, respectively? (a) CrO42– and PH3 (b) CrO42– and P4 (c) Cr(OH)4– and PH3 (d) Cr(OH)4– and P4 (e) CrO42– and OH– Answer: (b) 83. How many CrO42– ions are consumed in the balanced chemical equation? (a) 1 (b) 2 (c) 3 (d) 4 (e) 6 Answer: (d)

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Chapter 13: Chemical Thermodynamics E = E - (0.0257/n) ln Q

1 F = 96,500 C G = - nFE

R = 8.314 J/mol-K G = - RT ln K

Many of the following questions assume that a truncated table of standard-state thermodynamic data is attached to the exam. The table usually contains data for 6-10 compounds, which pertain to the particular questions on that exam.

Entropy and Disorder 1.

For which of the following processes would S be negative? (a) the melting of ice (b) the freezing of water (c) the evaporation of a liquid (d) the breakup of a large molecule into smaller molecules (e) these all have positive ∆S values

Answer: (b)

2.

Which of the following processes would have a negative entropy change? (a) decomposition of hydrogen peroxide (b) conversion of dry ice to gaseous CO2 (c) melting of ice (d) formation of Al2O3 from its elements (e) dissolution of salt in water

Answer: (d)

Entropy and the Second Law of Thermodynamics 3.

In any spontaneous process if the entropy change of the system is negative, the entropy change of the surroundings must be: (a) positive (b) negative (c) may be either positive or negative (d) must be zero

Answer: (a)

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4.

The second law of thermodynamics states that Suniverse is positive for a spontaneous process. However, our bodies are constantly carrying out spontaneous reactions in which simple molecules and compounds are used to make more complex ordered molecules such as proteins as well as more complex structures such as cells and tissues. According to the second law of thermodynamics how can these reactions be taking place spontaneously?

Answer: Suniverse is positive because the reactions are spontaneous and Ssystem must be negative because the reactions produce more complex molecules, then Ssurroundings must be positive and greater than the magnitude of Ssystem. The surroundings must gain entropy.

5.

Which of the following reactions would have a positive value for ∆S? (a) 3 NO(g) → NO2(g) + N2O(g) (b) 2 CO2(g) → 2 CO(g) + O2(g) (c) 2 I(g) → I2(g) (d) NH3(g) → NH3(l) (e) None of these reactions would have a positive ∆S.

Answer: (b)

6.

Which of the following reaction would be more likely to have a positive value for ∆S°? (a) 2 H2O(g) → 2 H2(g) + O2(g) (b) N2(g) + 3 H2(g) → 2 NH3(g) (c) CO2(g) + H2(g) → CO(g) + H2O(g) (d) H2O(l) → H2O(s) (e) None of these reactions would have a positive ∆S.

Answer: (a)

7.

Which of the following reactions would have a negative ∆S°? (a) H2(g) → 2 H(g) (b) H2(l) → H2(g) (c) H2(g) + Cl2(g) → 2 HCl(g) (d) 2 H2(g) + CO(g) → CH3OH(g) (e) none of the above

Answer: (d)

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Standard-State Entropies of Reaction 8.

Without using thermodynamic tables, predict the signs of H and S for the following reaction. CCl4(l) → CCl4(g) (a) H> 0 and S > 0 (b) H< 0 and S > 0 (c) H> 0 and S < 0 (d) H< 0 and S < 0 (e) none of these

Answer: (a)

9.

Which of the following reactions will most likely have a positive S°? (a) CO(g) + 2H2(g) → CH3OH(l) (b) NaCl(s) → Na+(aq) + Cl− (aq) (c) CH4(g) + 2 O2(g) → CO2(g) + 2 H2O(l) (d) 3H2(g) + N2(g) → 2 NH3(g) (e) none of these

Answer: (b)

10. Carbonated beverages can be made by bubbling carbon dioxide gas through the aqueous beverage and allowing the carbon dioxide to dissolve.

Predict the signs of H and S for the following reaction without using tables.

CO2(g) → CO2(aq) (a) H= +, S= + (b) H= −, S= + (c) H= −, S = − (d) H= +, S= 0 (e) H= 0, S= – Answer: (c)

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Calculating Entropy Changes for Chemical Reactions 11. The value of S for this reaction is: 2 Fe3+(aq) + Cu(s) → 2 Fe2+(aq) + Cu2+(aq) Compound Hac (kJ/molrxn)  Sac (J/molrxn-K) Cu(s) -338.32 -133.23 2+ Cu (aq) -273.55 -266.0 Fe2+(aq) -505.4 -318.2 Fe3+(aq) -464.8 -496.4 (a) between -400 and -200 J/molrxn-K (b) between -200 and 0 J/molrxn-K (c) between 0 and 200 J/molrxn-K (d) between 200 and 400 J/molrxn-K (e) more than 400 J/molrxn-K Answer: (d) 12. Use the Sac values for tin metal (Sac = -124.35 J/molrxn-K), chlorine gas (Sac = -107.33 J/molrxn-K), and tin tetrachloride (Sac = -570.7 J/molrxn-K) to calculate S for the following reaction.

Sn(s) + 2 Cl2(g) → SnCl4(l) (a) -339.0 (b) -231.7 (c) 231.7 (d) 339.0 (e) none of the above Answer: (b) 13. Use the following data to calculate S for the reaction: 2 NH3(g) → N2(g) + 3 H2(g) Compound Sac (J/molrxn-K) H2(g) -98.74 N2(g) -114.99 NH3(g) -304.99 (a) -198.8 J/molrxn-K (b) -91.26 J/molrxn-K (c) -106.22 J/molrxn-K (d) 91.26 J/molrxn-K (e) 198.8 J/molrxn-K Answer: (e)

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Gibbs Free Energy 14. If a reaction is spontaneous, what can be said about the reverse reaction under the same conditions? (a) The reverse reaction is nonspontaneous. (b) The reverse reaction is spontaneous also. (c) The reverse reaction is at equilibrium. (d) This cannot be determined.

Answer: (a) 15. Given that H for the reaction shown below is 92.2 kJ/molrxn, calculate G at 25C. 2 NH3(g) → N2(g) + 3 H2(g) Compound H2(g) N2(g) NH3(g)

Sac (J/molrxn-K) -98.74 -114.99 -304.99

(a) -106.5 kJ/molrxn (b) -33.0 kJ/molrxn (c) 33.0 kJ/molrxn (d) 106.5 kJ/molrxn (e) none of these Answer: (c)

16. For which of the following highly exothermic processes would you expect H and G to be about the same? (a) 2 Al(s) + 3/2 O2(g) → Al2O3(s) (b) 2 H2(g) + O2(g) → 2 H2O(g) (c) 2 Na(s) + 2 H2O(l) → 2 NaOH(aq) + H2(g) (d) 2 NO(g) → N2O4(g) (e) 2 Al(s) + Fe2O3(s) → 2 Fe(s) + Al2O3(s) Answer: (e)

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The Effect of Temperature on the Free Energy of a Reaction 17. Which of the following combinations of H and S indicates a nonspontaneous reaction under all conditions of temperature? (a) H> 0 and S > 0 (b) H< 0 and S > 0 (c) H> 0 and S < 0 (d) H< 0 and S < 0 (e) none of these Answer: (c)

18. Under which one of the following conditions will the following reaction be spontaneous as written?

2CO(g) + O2(g) → 2 CO2(g)

H= −565.97 kJ/molrxn, S= −173.0 J/molrxnK

(a) spontaneous at all temperatures (b) spontaneous at high temperatures (c) spontaneous at low temperatures (d) not spontaneous at any temperature Answer: (c) 19. Which of the following combinations of Ho and So can result in a reaction changing from nonspontaneous to spontaneous when the temperature increases? (a) Ho < 0 and So > 0 (b) Ho < 0 and So < 0 (c) Ho = 0 and So < 0 (d) Ho = 0 and So > 0 (e) Ho > 0 and So > 0 Answer: (e)

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20. Which statement correctly describes the following reaction?

PbO(s) + C(s) → Pb(s) + CO(g)

H = 106 kJ/molrxn S = 188 J/molrxnK,

(a) The reaction is spontaneous at all temperatures. (b) The reaction is spontaneous only below about 300 K. (c) The reaction becomes spontaneous at about 300 K. (d) The reaction becomes spontaneous at about 560 K. (e) The reaction is spontaneous only above 2000 K. Answer: (d) 21. The entropy change in the process: 2 H(g) → H2(g) is -98.74 J/molrxnK. The enthalpy of atom combination of the H2 molecule is 435 kJ/molrxn. Estimate the approximate temperature above which the formation of H2 molecules from H atoms will no longer be a spontaneous process. (a) 500 K (b) 1250 K (c) 2300 K (d) 4400 K (e) 6500 K Answer: (d)

Equilibria Expressed in Partial Pressures 22. Which of the following relationships between Kp and Kc is correct for the reaction N2(g) + 3 H2(g) → 2 NH3(g) when the reaction is carried out at 100C? (a) Kp = Kc (b) Kp < Kc (c) Kp > Kc (d) Kp < 0, Kc < 0 Answer: (b)

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23.

The Kp for the following equation is 3.2 x 10−3. Calculate the value of Kc at 25C. COCl2(g) → CO(g) + Cl2(g) (a) 7.8 x 10−2 (b) 6.6 x 10−3 (c) 1.6 x 10−3 (d) 1.3 x 10−4 (e) none of these

Answer: (d)

24. Assume that Kp = 1.75 for the reaction.

CO2(g) + C(s)

2 CO(g)

Calculate Kp for the following reaction.

CO(g) (a) 0.756 (b) 1.32 (c) 1.75 (d) 3.06 (e) none of the above

1/2 CO2(g) + 1/2 C(s)

Answer: (a)

25. The composition of the following system at equilibrium at 1000 K is 0.202 atm CO2, 0.594 atm CO, and 5 kg C. 2 CO(g)

CO2(g) + C(s)

What is the equilibrium constant for this reaction? (a) 2.94 (b) 14.56 (c) 0.572 (d) 0.337 (e) none of the above Answer: (c)

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26. Given the equilibrium constant for the following reaction, H2(g) + Cl2(g)

2 HCl(g)

Kp = 2.5 x 104

if equimolar amounts of H2 and Cl2 are allowed to reach equilibrium in a system in which no HCl was present initially, the ratio of the partial pressures of H2 to Cl2 at equilibrium will be (a) larger than 1. (b) smaller than 1. (c) equal to 1. (d) a value that can't be generalized by any of the above statements.

Answer: (c) 27. Which statement concerning the following reaction isn’t true?

2 HI(g)

H2(g) + I2(g)

(a) At equilibrium, the concentrations of the three gases do not change with time. (b) Kp = Kc (c) The partial pressures of I2 and H2 are equal at equilibrium, regardless of the initial pressures of HI, H2, and I2. (d) HI decomposes as fast as it is formed once equilibrium is established. (e) The percent dissociation of HI does not depend on the total pressure.

Answer: (c)

Interpreting Standard-State Free Energy of Reaction Data 28. The standard free energy change for the evaporation of water is found to be +8.56 kJ/molrxn.

H2O(l) → H2O(g) This says that the evaporation of water is nonspontaneous under standard conditions. If a beaker of water is set out on a table in the room why will it spontaneously evaporate?

Answer: The free energy change given in the problem is for standard conditions. Water evaporating in a room is not at standard conditions.

There is not 1 atm of water vapor present in the room. Therefore, the free energy change for the conditions in the room must be negative if water is to evaporate spontaneously.

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29. Which of the following statements is true? (a) G = G when the reaction is at equilibrium. (b) G is a measure of how far the reaction is from equilibrium. (c) G is positive for reactions that have too much reactant in the standard state. (d) The value of G does not depend on the temperature of the reaction. (e) None of these statements are true.

Answer: (b)

30. Which of the following statements is false? (a) G is equal to G when the system is at the standard state. (b) G is zero when the system is at equilibrium. (c) G measures how far the reaction is from equilibrium. (d) When G is positive, the reaction should proceed forward to form more product. (e) All of the above statements are true.

Answer: (d)

The Relationship between Free Energy and Equilibrium Constants 31. If the following reaction is spontaneous as written, under standard conditions, Zn(s) + Cu2+(aq) → Zn2+(aq) + Cu(s) which of the following statements is true? (a) Kc > 1 and ∆G° > 0 (b) Kc > 1 and ∆G° < 0 (c) Kc < 1 and ∆G° < 0 (d) Kc > 1 and ∆G° = 0 (e) Kc < 1 and ∆G° = 0 Answer: (b)

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32. For the reaction, 2 SO3(g) → 2 SO2(g) + O2(g), the following thermodynamic data are available.

Compound SO3(g) SO2(g) O2(g)

Hac (kJ/molrxn)

Sac (J/molrxn-K)

-1422.04 -1073.95 -498.34

-394.23 -241.21 -116.97

What is the equilibrium constant for this reaction at 25C? (a) between 10–50 and 10–30 (b) between 10–30 and 10–20 (c) roughly 1 (d) between 1020 and 1030 (e) between 1030 and 1050 Answer: (b) 2 Fe3+(aq) + Cu(s) → 2 Fe2+(aq) + Cu2+(aq) Compound Hac (kJ/molrxn) Cu(s) -338.32 2+ Cu (aq) -273.55 Fe2+(aq) -505.4 3+ Fe (aq) -464.8

33.

 Sac (J/molrxn-K) -133.23 -266.0 -318.2 -496.4

The equilibrium constant for this reaction at 25C is: (a) between 10–30 and 10–16 (b) between 10–16 and 10–10 (c) between 10–10 and 106 (d) between 106 and 1016 (e) larger than 1016 Answer: (d) 34. G° = -53.6 kJ/molrxn for the following reaction at 25°C. Zn2+(aq) + 4 NH3(aq) → Zn(NH3)42+(aq) What is the value of the complex formation equilibrium constant for this reaction? (a) 5 x 10–21 (b) 4 x 10–10 (c) 2.5 x 109 (d) 2.0 x 1020 (e) 1.2 x 1023

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Answer: (c) 35. Use the following information to answer this question.

H2(g) + F2(g) → 2 HF(g) H2(g) + Cl2(g) → 2 HCl(g) H2(g) + Br2(g) → 2 HBr(g) H2(g) + I2(g) → 2 HI(g)

G= -190 kJ/molrxn G = -546 kJ/molrxn G = -106 kJ/molrxn G = 4 kJ/molrxn

An excess of H2 is allowed to react with equal amounts of F2, Cl2, Br2, and I2 in a sealed flask. Which of the following products is present in the largest concentration at equilibrium? (a) HF (b) HCl (c) HBr (d) HI (e) All four products are present in equal concentrations.

Answer: (b) 36. What is the equilibrium constant for the following reaction if G° at 298 K is –16.5 kJ/molrxn?

N2(g) + 3 H2(g) → 2 NH3(g) (a) 2.0 x 10–5 (b) 1.3 x 10–3 (c) 780 (d) 4.6 x 106 (e) 220 Answer: (c) 37. Ka for benzoic acid (C6H5CO2H) is 3.0 x 10–5. What is the free energy change for the dissociation of this acid at 25°C? (a) -25.8 kJ/molrxn (b) -2.2 kJ/molrxn (c) 2.2 kJ/molrxn (d) 25.8 kJ/molrxn (e) none of the above Answer: (d)

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38. Which statement is true if the following reaction is spontaneous as written?

Zn(s) + Cu2+(aq) → Zn2+(aq) + Cu(s) (a) Kc is larger than 1, G° is negative, and E° is positive. (b) Kc is larger than 1, G° is positive, and E° is negative. (c) Kc is larger than 1, G° is negative, and E° is negative. (d) Kc is smaller than 1, G° is negative, and E° is positive. (e) Kc is smaller than 1, G° is positive, and E° is negative.

Answer: (a)

The Temperature Dependence of Equilibrium Constants 39. What happens to the magnitude of the equilibrium constant for the following reaction as the temperature of the system increases?

2 HI(g) → H2(g) + I2(g)

H= -9.48 kJ/molrxn ∆S°= -21.8 J/molrxn-K.

(a) Kc first increases and then decreases. (b) Kc first decreases and then increases. (c) Kc increases. (d) Kc decreases. (e) It is impossible to determine what happens to Kc. Answer: (d) 40. Which of the following statements is true about G, and K for the following reaction when the temperature is increased?

2 CO(g) + O2(g) → 2 CO2(g)

H = −565.97 kJ/molrxn S = −173.00 J/molrxnK

(a) As temperature increases, G will increase (become more positive), and K will increase . (b) As temperature increases, G will decrease (become more negative), and K will decrease. (c) As temperature increases, G increases (become more positive) and K decreases. (d) As temperature increases, G decreases (becomes more negative) and K increases. (e) None of these are true.

Answer: (c)

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41. Although the diatomic molecule N2 has an extremely strong nitrogen/nitrogen triple bond, it will break apart into nitrogen atoms (N2(g) → 2 N(g)) if it is heated to a high enough temperature because (a) the value of G becomes more positive as the temperature increases. (b) the unfavorable enthalpy change for the reaction disappears at higher temperatures. (c) the entropy change for the reaction is favorable and this factor becomes more important as the temperature of the system increases. (d) the equilibrium constant for the reaction decreases with temperature. (e) none of the above statements is a correct explanation.

Answer: (c)

Gibbs Free Energies of Formation and Absolute Entropies 42. Calculate the standard free energy change, G, and standard entropy change, S, for the following reaction at 298 K. Cl2(g) + 3 F2(g) → 2ClF3(g) Compound Hf (kJ/molrxn) ClF3(g) −163 Cl2(g) 0 F2(g) 0

S (J/molrxn-K) 123 223 203

(a) −123 kJ/molrxn, −134 J/molrxn-K (b) −246 kJ/molrxn, −268 J/molrxn-K (c) −326 kJ/molrxn, −268 J/molrxn-K (d) −51 kJ/molrxn, −586 J/molrxn-K (e) −148 kJ/molrxn, −134 J/molrxn-K Answer: (d) 43. Calculate the standard free energy, G, in kJ/molrxn for the reaction below.

4 NO2(g) + 6 H2O(g) → 4 NH3(g) + 7 O2(g) Compound Gf (kJ/molrxn) NO2(g) 51.31 H2O(g) −228.572 NH3(g) −16.45 (a) −1232.0 kJ/molrxn (b) 1232.0 kJ/molrxn (c) −478.9 kJ/molrxn (d) 478.9 kJ/molrxn (e) none of these Answer: (e)

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44. Calculate the standard enthalpy change, H, and standard entropy change, S, for the following reaction at 298 K. Compound CCl4(l) CCl4(g)

CCl4(l) → CCl4(g) Hf (kJ/molrxn) −135.44 −102.9

S (J/molrxn-K) 216.40 309.85

(a) 93.45 kJ/molrxn, 32.5 J/molrxn-K (b) 32.5 kJ/molrxn, 93.45 J/molrxn-K (c) −32.5 kJ/molrxn, 93.45 J/molrxn-K (d) −32.5 kJ/molrxn, −93.45 J/molrxn-K (e) –93.45 kJ/molrxn, 32.5 J/molrxn-K Answer: (b)

Integrated Problems 45. For each of the following variables what is the appropriate sign that is favorable for a spontaneous reaction? (a) H= +, S= +, G= +, E= + (b) H= −, S= +, G= +, E= + (c) H= +, S= +, G= −, E= + (d) H= +, S= +, G= +, E= − (e) H= −, S= +, G= −, E= + Answer: (e) 46. In chapter 12 you used Eo to predict when a redox reaction would be spontaneous. In chapter 13 Go was used to predict if a reaction was spontaneous. Which of the following describes the relationship between Go and Eo? (a) when Go is positive, Eo is negative (b) when Go is positive, Eo is positive (c) when Go is negative, Eo is negative (d) when ∆Go is negative, Eo is zero (e) There is no relationship between Go and Eo. Answer: (a)

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47. The equilibrium constant for the following reaction is Kp = 4.1 x 10–3 COCl2(g)

CO(g) + Cl2(g)

Calculate the equilibrium pressure of CO in a closed vessel that initially contained 0.30 atm of COCl2 and no CO or Cl2.

Answer: CO = 3.5 x 10–2, 48. The equilibrium constant for the following reaction is Kp = 1.0 x 104.

2 NO(g) + O2(g)

2 NO2(g)

Calculate the equilibrium pressure of oxygen in a system that initially contains 0.50 atm of NO, 0.10 atm of O2, and 0.20 atm of NO2.

Answer: O2 = 1.8 x 10–4

Questions 49-51 refer to the following reaction.

CO(g) + 2 H2(g) → CH3OH(g) for which the following data are available.

Compound Hac (kJ/molrxn) CO(g) -1076.38 H2(g) -435.3 CH3OH(g) -2037.11

Sac (J/molrxnK) -121.48 -98.74 -538.19

49. What is the value of S per mole of CH3OH produced in this reaction? (a) -219.2 J/molrxnK (b) -164.1 J/molrxnK (c) -88.4 J/molrxnK (d) 88.4 J/molrxnK (e) 219.2 J/molrxnK Answer: (a) 50. What is G per mole of CH3OH produced in this reaction at 25C? (a) -155.5 kJ/molrxn (b) -24.8 kJ/molrxn (c) +24.8 kJ/molrxn (d) 155.5 kJ/molrxn (e) none of the above Answer: (b)

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51. Which of the following statements about this reaction is true? (a) The reaction can never become favorable. (b) The reaction can never become unfavorable. (c) The reaction will become more favorable as the temperature is increased. (d) The reaction will become less favorable as the temperature is increased. (e) Temperature has no effect on whether this reaction is favorable.

Answer: (d) 52. Kp = 2.1 x 106 for the following reaction at a given temperature.

2 NH3(g)

N2(g) + 3 H2(g)

Starting with ammonia at 4.00 atm in a fixed-volume container, calculate the equilibrium pressures of all three gases. Justify any approximations used in solving this problem.

Answer: NH3 = 1.4 x 10–2, N2 = 2, H2 = 6 atm. If you start with all of the ammonia converted to products, then it is reasonable to assume that only a small change in the concentrations of all species occurs in going to equilibrium.

53. The following equilibrium is established in the gas phase at 600 K by heating phosgene, COCl2.

COCl2(g)

CO(g) + Cl2(g)

Kp = 4.10 x 10–3

If the initial pressure of phosgene is 0.120 atm, what is the equilibrium pressure of Cl2? (a) 4.9 x 10–4 atm (b) 2.0 x 10–2 atm (c) 2.2 x 10–2 atm (d) 0.10 atm (e) none of these Answer: (b)

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54. The equilibrium constant for the following gas-phase reaction is 68.0.

2 IBr(g)

I2(g) + Br2(g)

At equilibrium, twice as much Br2 is present as IBr and the partial pressure of I2 is 1 atm. What is the equilibrium partial pressure of IBr? (a) 0.0074 atm (b) 0.0037 atm (c) 0.029 atm (d) 0.059 atm (e) 136 atm Answer: (c)

55. The equilibrium constant for the reaction of NO(g) with Cl2(g) to produce NOCl(g) is 52. If 0.50 atm of NO and 0.25 atm of Cl2 react to equilibrium in a closed system that initially contained no NOCl, calculate the partial pressures of all three components at equilibrium.

Answer: Cl2 = 0.08, NO = 0.16, NOCl = 0.34 M

56. At a fixed temperature and 1.00 atm total pressure, NO2 gas and N2O4 gas are in equilibrium.

2 NO2(g) N2O4(g) What is the effect of increasing the total pressure to 10.0 atm at a constant temperature? (a) Kp increases, but the partial pressure of N2O4 decreases. (b) Kp decreases, but the partial pressure of N2O4 increases. (c) Both Kp and the partial pressure of N2O4 decrease. (d) Both Kp and the partial pressure of N2O4 increase. (e) Kp remains the same and the partial pressure of N2O4 increases.

Answer: (e)

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Chapter 14: Kinetics The Forces That Control a Chemical Reaction 1.

Which of the following experimental conditions describe a reaction which is kinetically controlled? (a) G = +100 kJ/molrxn; there is no observable reaction when reactants are mixed. (b) G = −100 kJ/molrxn; there is no observable reaction when reactants are mixed. (c) G = +100 kJ/molrxn; the reaction proceeds rapidly when reactants are mixed. (d) G = −100 kJ/molrxn; the reaction proceeds rapidly when reactants are mixed. (e) These conditions do not describe kinetic control.

Answer: (b)

Instantaneous Rates of Reaction 2.

Use the following table to determine the average rate during the period 800 to 1200 seconds for the decomposition of H2O2.

2 H2O2(l) → 2 H2O(l) + O2(g) Time (sec) 0 400 800 1200 1600

(H2O2) M 2.32 1.72 1.30 0.98 0.73

(a) 8.0 x 10−4 M/sec (b) 8.2 x 10−4 M/s (c) 1.2 x 10−3 M/s (d) 1.6 x 10−3 M/s (e) 2.4 x 10−3 M/s Answer: (a)

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3.

Determine the units for the rate constant, k, in the following rate expression: rate = k(A)2(B). (a) s−1 (b) M/s (c) M2/s (d) M3/s (e) M−2 s−1

Answer: (e)

4.

What will be the units of the rate constant, k, for the following reaction?

2 NO(g) + 2 H2(g) → N2(g) + 2 H2O(g)

rate = k (NO)2(H2)

(a) s−1 (b) M/s (c) M−1 s−1 (d) M−2 s−1 (e) M2 s−1 Answer: (d)

5.

The disproportionation of NO to N2O and NO2 is third order in NO. 3 NO(g) → N2O(g) + NO2(g) What are the units for the rate constant, k? (a) M3/s (b) M2/s2 (c) M−3 s−1 (d) M−2 s−2 (e) none of these

Answer: (e)

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rate = k (NO)3

The Rate Law Versus the Stoichiometry of a Reaction 6.

Nitrogen reacts with hydrogen to form ammonia by the following reaction: N2(g) + 3 H2(g) → 2 NH3(g) If the rate of consumption of hydrogen is 0.060 M/s, what is the rate at which N2 is consumed? (a) 0.010 M/s (b) 0.020 M/s (c) 0.030 M/s (d) 0.060 M/s (e) 0.180 M/s

Answer: (b)

7.

Which equation describes the relationship between the rates at which NO is consumed and N2 is produced in the following reaction?

2 NO(g) + 2 H2(g) → N2(g) + 2 H2O(g) (a) rate N2 = rate NO (b) rate N2 = 2(rate NO) (c) 2(rate N2) = rate NO (d) (rate N2)2 = rate NO (e) rate N2 = (rate NO)2

Answer: (c)

8.

The reaction 2A → B is second order with respect to A. If the concentration of A is decreased by half, what will happen to the rate of the reaction? (a) The rate will double. (b) The rate will decrease by half. (c) The rate will remain constant. (d) The rate will increase by ln(A). (e) The rate will decrease by a factor of 4.

Answer: (e)

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9.

For the reaction 3 N2O(g) + C2H2(g) → 3 N2(g) + 2 CO(g) + H2O(g) if water is produced at the rate of 0.10 M/s, what is the rate of production of N2? (a) 0.10 M/s (b) 0.033 M/s (c) 0.30 M/s (d) 0.20 M/s (e) none of these

Answer: (c)

10. Which of the following equations correctly describes the relationship between the instantaneous rate at which NO2 and Cl2 are consumed in the following reaction?

2 NO2(g) + Cl2(g) → 2 NO2Cl(g) (a) -d(NO2)/dt = 1/2 [-d(Cl2)/dt] (b) -d(NO2)/dt = 1/2 [d(Cl2)/dt] (c) -d(NO2)/dt = -d(Cl2)/dt (d) -d(NO2)/dt = 2 [-d(Cl2)/dt] (e) -d(NO2)/dt = 2 [d(Cl2)/dt] Answer: (d)

11. In the following reaction 2 VO43–(aq) + 3 Zn(s) + 16 H+(aq) → 2 V2+(aq) + 3 Zn2+(aq) + 8 H2O(l) the initial rate of disappearance of VO43– was found to be 0.56 M/s. What is the initial rate of appearance of Zn2+? (a) -0.56 M/s (b) 0.37 M/s (c) 0.56 M/s (d) 0.84 M/s (e) 1.12 M/s Answer: (d)

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12. The instantaneous rate of appearance of water from the reaction 4 NH3(g) + 5 O2(g) → 4 NO(g) + 6 H2O(g) at some moment in time is 21.3 mmHg/min. What is the instantaneous rate of disappearance of NH3 at the same moment in time? (a) 10.7 mmHg/min (b) 14.2 mmHg/min (c) 21.3 mmHg/min (d) 32.0 mmHg/min (e) none of these Answer: (b)

13. When NH3 is treated with O2 at elevated temperatures, the rate of disappearance of NH3 is found to be 3.5 x 10–2 M s–1. The equation for the reaction is 4 NH3(g) + 5 O2(g) → 4 NO(g) + 6 H2O(g) Calculate the rate of appearance of NO and H2O. Answer: NO = 0.035 M s–1; H2O = 0.053 M s–1

14. For which reactant or product in the following reaction will the rate of change of concentration with time be the largest?

6 Fe2+ + Cr2O72– + 14 H+ → 2 Cr3+ + 6 Fe3+ + 7 H2O (a) Cr3+ (b) Cr2O72– (c) Fe2+ (d) H+ (e) H2O Answer: (d)

15. What is the correct rate law for the following reaction?

2 NO2(g) + F2(g) → 2 NO2F(g) (a) rate = kf(NO)2 (b) rate = kf(NO2)2(F2) (c) rate = -kf(NO2)2(F2) (d) rate = -kf(F2) (e) Not enough information is available to determine the rate law.

Answer: (e) 14-5

A Collision Theory of Chemical Reactions 16. Which of the following affect whether a collision between molecules will result in a chemical reaction? (a) change in enthalpy for the reaction (b) orientation of the molecules (c) energy of the collision (d) ‘b’ and ‘c’ only (e) all of these Answer: (d)

The Mechanisms of Chemical Reactions 17. Sum the following elementary steps in a mechanism to determine the stoichiometry of the overall reaction.

Cl2(g) →2Cl(g) Cl(g) + CO(g) → COCl(g) COCl(g) + Cl(g) → COCl2(g) (a) Cl2(g) + 2 CO(g) + 2 Cl(g) →2 COCl2(g) (b) 2 Cl2(g) + CO(g) → COCl2(g) + 2 Cl(g) (c) 2 Cl(g) + CO(g) → COCl2(g) (d) Cl2(g) + CO(g) + Cl(g) → COCl2(g) + COCl(g) (e) none of these Answer: (e) 18. The reaction between chloroform (CHCl3) and chlorine gas (Cl2) proceeds by the following mechanism.

Cl2 → 2 Cl CHCl3 + Cl → CCl3 + HCl CCl3 + Cl → CCl4 Which of the following substances are intermediates in the above mechanism? (a) Cl (b) CCl3 (c) CCl4 (d) HCl (e) Cl & CCl3 Answer: (e)

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19. The overall rate of a chemical reaction is determined by (a) the fastest step in the reaction mechanism. (b) the first step in the reaction mechanism. (c) the slowest step in the reaction mechanism. (d) the last step in the reaction mechanism. (e) ∆G for the overall reaction.

Answer: (c)

20. Which of the following rate laws suggests that the reaction probably occurs in a single step? (a) NO(g) + O2(g) → NO2(g) + O(g) rate = k[NO][O2] (b) H2(g) + Br2(g) → 2 HBr(g) rate = k[H2][Br2]1/2 (c) (CH3)3CBr + OH– → (CH3)3COH + Br– rate = k[(CH3)3CBr] (d) H2O2 + 3 I– + 2 H+ → I3– + 2 H2O rate= kl[H2O2][I–]+K2[H2O2][I–][H+] Answer: (a)

21. Use the rate laws given below to determine which of the following reactions most likely occurs in a single step. (a) 2 NO2(g) + F2(g) → 2 NO2F(g) (b) 2 NOCl(g) → 2 NO(g) + 2Cl2g) (c) NO2(g) + CO(g) → NO(g) + CO(g) (d) 2 O3(g) → 3 O2(g) (e) H2(g) + Br2(g) → 2 HBr(g) Answer: (b)

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rate = k(NO2)(F2) rate = k(NOCl)2 rate = k(NO2)2 rate = k(O3) rate = k(Br2)1/2(H2)

Determining the Order of a Reaction from Rates of Reaction Questions 22-24 refer to the reaction: NO(g) + O3(g) → NO2(g) + O2(g) for which the following rate data were obtained.

Trial #1 Trial #2 Trial #3 Trial #4 Trial #5

Initial (NO) 2.1 x 10–6 M 4.2 x 10–6 M 6.3 x 10–6 M 6.3 x 10–6 M 6.3 x 10–6 M

Initial (O3) 2.1 x 10–6 M 2.1 x 10–6 M 2.1 x 10–6 M 4.2 x 10–6 M 6.3 x 10–6 M

Initial Rate of Reaction 1.6 x 10–5 M s-1 3.2 x 10–5 M s-1 4.8 x 10–5 M s-1 9.6 x 10–5 M s-1 14.4 x 10–5 M s-1

22. The rate law for the reaction would be: (a) zero-order in NO (b) first-order in NO (c) second-order in NO (d) third-order in NO (e) none of the above Answer: (b)

23. If these data were obtained by watching the rate at which NO disappears, what would be the initial instantaneous rate of disappearance of O 3 in Trial 3? (a) 1.6 x 10–5 M/s (b) 3.2 x 10–5 M/s (c) 4.8 x 10–5 M/s (d) 9.6 x 10–5 M s (e) 14.4 x 10–5 M/s Answer: (c) 24. If a trial was run in which the initial (NO) was 3.15 x 10–6 M and the initial (O3) was 3.15 x 10–6 M, the initial rate of reaction would be: (a) 1.6 x 10–5 M/s (b) 3.2 x 10–5 M/s (c) 3.6 x 10–5 M/s (d) 4.8 x 10–5 M/s (e) 7.2 x 10–5 M/s Answer: (c)

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This data applies to the next three problems.

The initial rate of disappearance of bromine (Br2) for the reaction shown below was measured for several different concentrations of bromine, CH3COCH3, and H+ ions. (Note that H+ is a catalyst in this reaction – it participates in the reaction but is not itself consumed) CH3COCH3 + Br2 + H+ → CH3COCH2Br + H+ + BrTrial# 1 2 3 4

[CH3COCH3] M 0.30 0.30 0.30 0.40

[Br2] M 0.050 0.10 0.050 0.050

[H+] 0.050 0.050 0.10 0.20

25. The rate law for this reaction would be: (a) zero-order in H+ and first-order in Br2 (b) first-order in H+ and first-order in Br2 (c) first-order in H+ and zero-order in Br2 (d) second-order in H+ and zero-order in Br2 (e) none of the above.

Answer: (c)

26. The rate law for this reaction would be: (a) zero-order in CH3COCH3 (b) half-order in CH3COCH3 (c) first-order in CH3COCH3 (d) second-order in CH3COCH3 (e) none of the above.

Answer: (c)

27. The rate constant for this reaction is: (a) smaller than 1 x 10-5 (b) between 1 x 10-5 and 1x 10-4 (c) between 1 x 10-4 and 1 x 10-3 (d) between 1 x 10-3 and 1 (e) greater than 1 Answer: (d)

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Initial Rate M/s 6.0 x 10−5 6.0 x 10−5 1.2 x 10−4 3.2 x 10-4

28. The reaction shown below is first order in oxygen and second order in NO and has a rate constant of 7.1 x 109 M-2s-1. What would the reaction rate be when the NO concentration is 0.0010 M and the O2 concentration is 0.034 M? 2NO + O2 → 2NO2 (a) 2.4 x 105 M/s (b) 3.4 x 10-8 M/s (c) 3.7 x 103 M/s (d) 8.9 x 10-5 M/s (e) 2.4 x 102 M/s Answer: (e)

29. The reaction following reaction was studied at 25C, and the following experimental results were obtained.

2I−(aq) + S2O82−(aq) → I2(aq) + 2 SO42−(aq) Trial# 1 2 3 4 5

(I−) M 0.080 0.040 0.080 0.032 0.060

(S2O82−) M 0.040 0.040 0.020 0.040 0.030

Initial Rate M/s 12.5 x10−6 6.25 x10−6 6.25 x10−6 5.00 x10−6 7.00 x10−6

Which 2 trials of experimental data should be chosen to find the order of the reaction with respect to S2O82−? (a) trial 1 and 2 (b) trial 2 and 3 (c) trial 2 and 5 (d) trial 1 and 3 (e) none of these Answer: (d)

30. Use the following data on the initial rate of reaction to determine the overall order of the reaction.

BrO3–(aq) + 5 Br–(aq) + 6 H+(aq) → 3 Br2(aq) + 3 H2O(l) Initial rate of reaction

Initial [BrO3–] Initial [Br–]

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Initial [H+]

9.0 x 10–4 1.4 x 10–3 3.5 x 10–3 8.8 x 10–3

0.10 M 0.10 M 0.25 M 0.25 M

0.10 M 0.15 M 0.15 M 0.15 M

0.10 M 0.10 M 0.10 M 0.25 M

(a) first-order overall (b) second-order overall (c) third-order overall (d) fourth-order overall (e) none of these Answer: (c)

Questions 31 through 35 refer to the reaction for which the following data were obtained.

2 NO(g) + Br2(g) → 2 NOBr(g) Trial 1 Trial 2 Trial 3 Trial 4 Trial 5

Initial (NO) 0.025 M 0.025 M 0.025 M 0.050 M 0.075 M

Initial (Br2) 0.040 M 0.080 M 0.120 M 0.040 M 0.040 M

Initial rate of reaction 1.2 M/s 2.4 M/s 3.6 M/s 4.8 M/s 10.8 M/s

31. The rate law for this reaction would be: (a) zero-order in NO (b) half-order in NO (c) first-order in NO (d) second-order in NO (e) none of the above Answer: (d)

32. The rate law for this reaction would be: (a) zero-order in Br2 (b) half-order in Br2 (c) first-order in Br2 (d) second-order in Br2 (e) none of the above Answer: (c)

33. If these data were obtained by watching the rate at which NO disappears, what would be the initial instantaneous rate of disappearance of Br2 in

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Trial 1? (a) 0.30 M/s (b) 0.60 M/s (c) 1.2 M/s (d) 2.4 M/s (e) 4.8 M/s Answer: (b)

34. The magnitude of the rate constant, k, for this reaction is: (a) less than 1 x 10–4 (b) between 10–4 and 10–2 (c) between 0.01 and 100 (d) between 100 and 10,000 (e) more than 10,000 Answer: (e)

35. A plot of (NO) versus time would most closely resemble: (a) a straight line with a positive slope. (b) a straight line with a negative slope. (c) a straight line with a slope of zero. (d) a curve in which the NO concentration increases rapidly at first and then slows down until eventually a maximum concentration is achieved. (e) a curve in which the NO concentration decreases rapidly at first and then slows down until eventually a minimum concentration is achieved.

Answer: (e)

Question 36 through 40 refer to the reaction Cr2O72–(aq) + 6 I–(aq) + 14 H+(aq) → 2 Cr3+(aq) + 3 I2(aq) + 7 H2O(l) for which the following initial instantaneous rates of reaction were obtained. initial (Cr2O72–) Trial 1 0.0040 M Trial 2 0.0040 M Trial 3 0.0040 M Trial 4 0.0040 M Trial 5 0.0080 M

initial (I–) 0.010 M 0.020 M 0.030 M 0.030 M 0.030 M

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initial (H+) 0.020 M 0.020 M 0.020 M 0.040 M 0.040 M

initial rate 5.0 x 10–4 M/s 20 x 10–4 M/s 46 x 10–4 M/s 179 x 10–4 M/s 355 x 10–4 M/s

36. If the initial rate data are for the rate at which I– disappears, -d(I–)/dt, what would be the rate of disappearance of Cr2O72– in Trial 1? (a) 0.83 x 10–4 M/s (b) 2.5 x 10–4 M/s (c) 5 x 10–4 M/s (d) 10 x 10–4 M/s (e) 30 x 10–4 M/s Answer: (a)

37. The rate law for this reaction would be: (a) zero-order in Cr2O72– (b) half-order in Cr2O72– (c) first-order in Cr2O72– (d) second-order in Cr2O72– (e) none of the above Answer: (c)

38. The rate law for this reaction would be: (a) first-order in both I– and H+ (b) first-order in I– and second-order in H+ (c) first-order in H+ and second-order in I– (d) second-order in both H+ and I– (e) none of the above Answer: (d)

39. The rate constant, k, for this reaction is: (a) less than 10–6 (b) between 10–6 and 10–3 (c) between 10–3 and 103 (d) between 103 and 106 (e) more than 106 Answer: (e) 40. A graph of (Cr2O72–) versus time would most closely resemble (a) a straight line with a negative slope. (b) a straight line with a slope of zero.

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(c) a straight line with a positive slope. (d) a curve in which the Cr2O72– concentration decreases rapidly at first, and then the rate at which Cr2O72– disappears slows down with time. (e) a curve in which the Cr2O72– concentration increases rapidly at first, and then the rate at which Cr2O72– appears slows down with time.

Answer: (d)

The Integrated Forms of Zero-, First-, and Second-Order Rate Laws 41. A certain substance, initially at 0.10 M in solution, decomposes by secondorder kinetics. If the rate constant for this process is 0.40 M−1 min−1, how much time in minutes is required for the concentration to reach 0.020 M? (a) 1.6 x 10−4 min (b) 8.0 x 10−3 min (c) 4.02 min (d) 50 min (e) 100 min Answer: (e) 42. The reaction A → C is found to be zero order. Which of the following will give a linear plot? (a) (A) vs. time (b) ln(A) vs. time (c) 1/(A) vs. time (d) None of these will be linear.

Answer: (a)

43. The half-life for the first-order decomposition of nitramide, NH2NO2, into nitrous oxide, N2O, and water is 123 min at 15°C. NH2NO2 → N2O + H2O What is the value of the rate constant (k) for this reaction?

Answer: 5.64 x 10–3 min–1

44. For the reaction in the above question, how long will it take in minutes for 2.0 g of nitramide to decompose until only 0.20 g remains?

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Answer: 410 min

(a)

Rate

45. Which of the following graphs describes the relationship between rate and time for a first-order reaction?

time

(c)

Rate

(b)

(d)

Rate

time

(e)

Rate

time

time Answer: (b)

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46. Which plot best describe the rate of a second-order reaction?

(a)

ln (A)

2A→B

time

(A)

(b)

(c)

1/(A)

time

(d)

ln (A)

time

(e)

(A)

time

time Answer: (c)

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47. A certain second-order reaction is found to have a rate constant of 0.135 M–1 s–1 What is the half-life of the reaction? (a) 5.1 s (b) 6.5 s (c) 7.4 s (d) none of these (e) impossible to determine from this information Answer: (e)

48. Which of the following statements about the half-life of a reaction is true? (a) The half-life doesn’t depend on the order of the reaction. (b) The half-life of a first-order reaction increases with time. (c) The half-life of a second-order reaction is independent of concentration. (d) A zero-order reaction doesn’t have a half-life. (e) None of the above are true.

Answer: (e) 49. For a second-order reaction, it takes 15 seconds for the initial concentration of a reactant to decrease from 0.60 M to 0.53 M. What is the initial half-life of this reaction? (a) 47 s (b) 94 s (c) 114 s (d) 128 s (e) 228 s Answer: (c) 50. The rate constant, k, for the first-order decay of 14C is 1.21 x 10–4 yr. If analysis of a piece of paper, parchment, or papyri suggests that 97.4% of the 14C that was present initially still remains in the sample, it would most likely be associated with which of the following battles? (a) Battle of Actium, 31 BC, Octavian defeating Mark Anthony (b) Battle of Hastings, 1066 AD, William of Normandy defeating King Harold II of England (c) Battle of Stirling Bridge, 1297 AD, William Wallace defeating John de Warrenne, Earl of Surrey (d) Battle of Yorktown, 1781 AD, Washington defeating Lord Cornwallis (e) Battle of Waterloo, 1815 AD, Wellington defeating Napoleon

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Answer: (d)

51. The decay of radioactive nuclei is a first-order kinetic process. Determine the half-life for the decay of 238U to 206Pb if the first-order rate constant is 4.87 x 10–18 s–1 (a) 3.37 x 10–18 s (b) 39.5 s (c) 1.42 x 1017 s (d) 4.11 x 1017 s (e) none of these Answer: (c) 52. The half-life of 228Ra is 1590 years. How long will it take for a sample of this nuclide to decay to 1.0% of its original activity? (a) 16 years (b) 1.6 x 103 years (c) 4.1 x 103 years (d) 1.1 x 104 years (e) none of these Answer: (d)

53. The rate of a zero-order reaction (a) increases as reactant is consumed. (b) depends on the concentration of products. (c) decreases as reactant is consumed. (d) is independent of temperature. (e) is independent of the concentration of reactants and products.

Answer: (e)

Determining the Order of a Reaction with the Integrated Forms of Rate Laws 54. The reaction, N2O5 → NO2 + 1/2 O2, is first-order in N2O5 with a half-life of 19.25 min. How long would it take in minutes for the N2O5 concentration to decrease from 0.050 M to 0.030 M? Answer: 14.2 min

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The following data table should be used to answer questions 55-57.

Reaction:

A → B

(A), M time, s

0.380 0.319 0.268 0.225 0.189 0.158 0.133 0 10 20 30 40 50 60

55. What is the rate constant, k, for this reaction? (a) 0.0175 (b) 0.0560 (c) 0.0815 (d) 0.112 (e) none of these Answer: (a)

56. What is the half-life for this reaction, in seconds? (a) 29.1 (b) 39.6 (c) 47.3 (d) 151 (e) none of these Answer: (b)

57. What will be the concentration of A when t = 120 seconds? (a) 0.047 (b) 0.10 (c) 0.13 (d) 0.21 (e) none of these Answer: (a)

58. If a plot of 1/[A] versus time produces a straight line, which of the following is true? (a) This reaction is first-order in A. (b) The reaction is second-order in A. (c) The reaction is first-order in two reactants. (d) The rate of reaction does not depend on the concentration of A. (e) none of the above Answer: (b) 14-19

Use the following data to answer questions 59-61. These data were obtained by monitoring the rate at which the OCl– ion was consumed in the presence of a large excess of the I– ion.

OCl–(aq) + I–(aq) → OI–(aq) + Cl–(aq) [OCl–], M 0.00250 0.00137 0.00075 0.00041 time, s 0 3 6 9 59. What is the rate constant (k) for this reaction? (a) between 10–2 and 10–1 (b) larger than 10–1 and smaller than 1 (c) between 1 and 10 (d) larger than 10 and smaller than 100 (e) between 100 and 1000 Answer: (b)

60. What is the half-life in seconds? (a) between 0.01 and 0.1 seconds (b) between 0.1 and 1 second (c) between 1 and 10 seconds (d) between 10 and 100 seconds (e) none of the above Answer: (c) 61. What is the concentration of OCl– after 12 seconds? (a) between 1 x 10–4 and 2 x 10–4 (b) between 2 x 10–4 and 3 x 10–4 (c) between 3 x 10–4 and 4 x 10–4 (d) between 4 x 10–4 and 5 x 10–4 (e) between 5 x 10–4 and 6 x 10–4 Answer: (b)

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Questions 62-64 refer to the following reaction for the decomposition of H2O2.

2 H2O2(aq) → 2 H2O(l) + O2(g) The following data were obtained (H2O2), M: time, s:

10.00 8.25 6.80 5.61 4.63 0 500 1000 1500 2000

62. The rate law for this reaction is: (a) zero-order in H2O2 (b) first-order in H2O2 (c) second-order in H2O2 (d) third-order in H2O2 (e) none of these Answer: (b)

63. The half-life for this reaction is: (a) between 0 and 50 s (b) between 50 and 500 s (c) between 500 and 1500 s (d) between 1500 and 2000 s (e) the half-life can't be obtained from these data Answer: (d)

64. What will be the H2O2 concentration after 4000 seconds? (a) less than 1.5 M (b) between 1.5 and 2.5 M (c) between 2.5 and 3.0 M (d) between 3.0 and 3.5 M (e) between 3.5 and 4.5 M Answer: (b)

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Questions 65-67 refer to the reaction 2 N2O5(g) → 4 NO2(g) + O2(g) for which the following data were obtained.

(N2O5), M: time, s:

5.00 3.52 2.48 1.75 1.23 0 500 1000 1500 2000

65. The rate law for this reaction would be: (a) zero-order in N2O5 (b) half-order in N2O5 (c) first-order in N2O5 (d) second-order in N2O5 (e) none of the above Answer: (c)

66. The half-life for this reaction is: (a) between 0 and 12 s (b) between 12 and 120 s (c) between 120 and 1200 s (d) between 1200 and 12,000 s (e) between 12,000 and 120,000 s Answer: (c)

67. What will be the concentration of N2O5 after 5000 seconds? (a) between 0.001 and 0.010 M (b) between 0.010 and 0.10 M (c) between 0.10 and 0.4 M (d) between 0.4 and 0.8 M (e) between 0.8 and 1.2 M Answer: (c)

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Questions 68 and 69 refer to the reaction 2 H2O2(aq)

Pt

2 H2O(l) + O2(g)

for which the following data were obtained at 25%C. (H2O2), M: time, s:

1.000 0.549 0.301 0.165 0.091 0.050 0.027 0 30 60 90 120 150 180

68. The rate law for this reaction is (a) zero-order in hydrogen peroxide (b) half-order in hydrogen peroxide (c) first-order in hydrogen peroxide (d) second-order in hydrogen peroxide (e) none of the above Answer: (c) 69. When will the concentration of H2O2 reach 0.010 M? (a) before 200 s (b) between 200 and 240 s (c) between 240 and 280 s (d) between 280 and 320 s (e) after 320 s Answer: (b) The Activation Energy of Chemical Reactions 70. H2 and O2 can react explosively. However, a mixture of H2 and O2 can exist indefinitely at room temperature with no reaction occurring. Explain why hydrogen and oxygen do not react under these conditions.

Answer: The activation energy must be supplied in order for the reaction to begin.

71. If the rate constant increases from 0.40 M–1 s–1 at 25C to 0.80 M–1 s–1 at 35C, what is the activation energy in kJ/mol for this reaction? (a) between 0 and 40 kJ mol (b) between 41 and 80 kJ/mol (c) between 81 and 120 kJ/mol (d) between 121 and 160 kJ/mol (e) between 161 and 200 kJ/mol Answer: (b) 14-23

72. A chemical reaction is found to be endothermic. Which of the following correctly describes the activation energy associated with the forward and the reverse reactions? (a) Ea forward = Ea reverse (b) Ea forward < Ea reverse at all temperatures (c) Ea forward > Ea reverse at all temperatures (d) Ea forward = Ea reverse at low temperatures (e) none of these Answer: (c)

73. The activation energy was measured for both the forward (Ea=150 kJ/molrxn) and reverse (Ea = 95 kJ/molrxn) directions of a reversible reaction. What would be the activation energy for the reverse reaction in the presence of a catalyst that decreased the activation energy for the forward reaction to 125 kJ/ molrxn? (a) 25 kJ/molrxn (b) 75 kJ/ molrxn (c) 120 kJ/ molrxn (d) 30 kJ/ molrxn (e) none of these Answer: (e)

74. Which of the following correctly describes the variation of rate constant, kf, with temperature? (a) kf = ln Z - Ea/RT (b) ln kf = ln Z + e–(Ea/RT) (c) ln kf = ln Z - Ea/RT (d) ln kf = ln Z + Ea/RT (e) kf = ln Z - Ea/RT Answer: (c)

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75. Which of the following plots represents the largest rate constant for an endothermic reaction?

(a)

(c)

Energy

(b)

reaction coordinate

(d)

(e)

Answer: (d)

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76. For the gas phase decomposition of N2O2, k is 1.4 s−1 at 400 K and 43 s−1 at 450 K. What is the activation energy for this reaction? (a) 102 kJ/molrxn (b) 3.42 kJ/molrxn (c) 50.0 kJ/molrxn (d) 232 kJ/molrxn (e) 78.4 kJ/molrxn Answer: (a)

Integrated Problems 77. Which statement is true? (a) The rate of appearance of products of a chemical reaction is always equal to the rate of disappearance of reactants. (b) If a reaction follows a second-order rate law, it must have two steps in its reaction mechanism. (c) The half-life for a first-order reaction is always larger than the half-life for a second-order reaction. (d) The half-life for a first-order reaction is independent of the initial concentration of the reactant. (e) The half-life for a second-order reaction is independent of the initial concentration of the reactant.

Answer: (d) 78. For the reaction, 2 NOCl(g) → 2 NO(g) + Cl2(g), the observed rate expression is rate = k(NOCl)2. Which of the conditions listed below would cause a change in the value of the rate constant, k? (a) increase the concentration of NOCl (b) decrease the concentration of NOCl (c) decrease the concentration of NO (d) increase the temperature (e) increase the concentration of NOCl and decrease the concentration of NO and Cl2 Answer: (d)

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79. As the reaction below proceeds, what happens (increase, decrease, or remain constant) to the rate of the forward reaction?

CO(g) + H2O(g) → CO2(g) + H2(g) As the reaction proceeds what happens to the rate of the reverse reaction (CO2 + H2 → CO + H2O)?

Explain the above answers using the collision theory for gas-phase reactions.

Answer: decreases; increases; As the concentration of the reactants decreases there are fewer collisions between reactant molecules and the rate of forward reaction slows. As more product molecules are produced there is more probability of a collision between product molecules and the rate of reverse reaction increases.

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Chapter 15: Nuclear Chemistry Radioactivity and the Structure of the Atom 1. Which form of radiation will be deflected by an electric field? (a)  particles (b)  particles (c)  rays (d)  and  particles (e) All three forms of radiation will be deflected by an electric field.

Answer: (d)

2. Which pair represents the same thing? (a)  and  radiation (b)  particles and -10e (c)  rays and a proton (d)  and  particles (e) Each pair represents the same thing.

Answer: (b) 3. 1H, 2H, and 3H are best described as ____ (a) isotopes (b) isobars (c) isotones (d) compounds (e) alpha particle emitters Answer: (a) 4. 9Li, 9Be, and 9C are best described as ____ (a) isotopes (b) isobars (c) isotones (d) compounds (e) covalent Answer: (b)

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5. 44P, 50Sc, and 52V are best described as ____ (a) isotopes (b) isobars (c) isotones (d) compounds (e) p-block elements Answer: (c)

Modes of Radioactive Decay 6. What is the identity of X in the following equation?

238U

→

4He

+ X

(a) 234Pu (b) 242Ra (c) 234U (d) 234Th (e) none of the above Answer: (d)

7. What is the identity of X in the following equation?

X →

3He

+ -10e

(a) 3H (b) 2H (c) 4Li (d) 234Th (e) 3Li Answer: (a)

8. What is the identity of X in the following equation?

65Zn

+ X →

65Cu

(a) 3H (b) -10e (c) +10e (d) 01n 15-2

(e) none of the above Answer: (b) 9. The decay of 40K to 40Ar can be used to determine the age of very old rocks.

What mode of radioactive decay is occurring in the conversion of 40K to 40Ar? (a)  emission (b) electron capture (c) positron emission (d) neutron emission (e) none of the above Answer: (b)

10. What is the identity of X in the following equation?

19F

+ X →

16O

+ 4He

(a) 11p (b) 3H (c) +10e (d) 01n (e) none of the above Answer: (a)

11. Element number 43, technetium, is not naturally occurring but is prepared by nuclear reactions for use in medical imaging. The starting material, 98Mo, captures one neutron to give a new species X. X decays by electron emission to give species Y which is formed in a metastable state. Y then decays to give 99Tc with the emission of a gamma ray. Identify X and Y and write balanced nuclear equations for their formation and decomposition.

Answer: 98Mo + 1n → 99Mo 99Mo

→ 99mTc +

99mTc

→ 99Tc + 

0 -1 e

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12. What is the identity of X in the following equation?

235U

+

1 0n

→

99Sr

+ 135Te + 2X

(a) +11p (b) 3H (c) +10e (d) 01n (e) none of the above Answer: (d)

13. Which of the following decay modes will result in an increase in the atomic number of a nuclide? (a) electron emission (b) positron emission (c)  emission (d)  ray emission (e) All of the above decay modes will result in an increase in atomic number.

Answer: (a)

Neutron-Rich vs. Neutron-Poor Nuclides 14. Which of the following nuclides would be the most likely to decay by electron emission? (a) 17Ne (b) 18Ne (c) 19Ne (d) 24Ne (e) They are all equally likely to decay by electron emission.

Answer: (d)

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15. Which of the following nuclides would be the most likely to decay by positron emission? (a) 19Na (b) 20Na (c) 21Na (d) 22Na (e) They are all equally likely to decay by positron emission.

Answer: (a)

16. A neutron-poor nuclide can decay by: (a) positron emission (b) electron capture (c)  decay (d) A neutron-poor nuclide cannot decay by any of these routes. (e) A neutron-poor nuclide can decay by any one of these three routes, depending on the mass of the nuclide.

Answer: (e)

17. Which of the following nuclides would be most likely to be considered neutron-rich? (a) 26Si (b) 55Cr (c) 61Cu (d) 4Li (e) 37Ar Answer: (b)

18. Which of the following nuclides is most likely to decay by electron capture? (a) 125I (b) 33P (c) 8B (d) 14C (e) 210Pb Answer: (a)

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Binding Energy 19. The sum of the mass of 11 protons, 12 neutrons, and 11 electrons will be _____ the mass of a 23Na atom. (a) greater than (b) smaller than (c) equal to (d) not related to Answer: (a) 20. The exact mass of a 12C atom does not equal the sum of the mass of 6 protons, 6 neutrons, and 6 electrons because (a) a 12C atom has 7 neutrons, not 6. (b) some of the mass of the protons and neutrons is converted to energy when they are combined to form the 12C atom. (c) carbon occurs naturally as a mixture of isotopes with slightly different masses. (d) the enthalpy of atom combination of 12C is negative. (e) none of the above Answer: (b) 21. 223Rn has a binding energy of 1712.325 MeV. What is the binding energy per nucleon? (a) 19.91076 MeV/nucleon (b) 12.49872 MeV/nucleon (c) 34.83745 MeV/nucleon (d) 1.033238 MeV/nucleon (e) 7.678587 MeV/nucleon Answer: (e) The following data is needed for problems 22-24 Particle 1 +1 p 1 0n 0 -1 e

Mass in amu 1.0072765 1.0086650 0.0005486

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22. What is the binding energy in MeV/atpm for a 14C atom which has an exact mass of 14.003242 amu? (a) between 0 and 30 MeV (b) between 30 and 60 MeV (c) between 60 and 90 MeV (d) between 90 and 120 MeV (e) more than 120 MeV Answer: (d) 23. What is the exact mass in amu of 31P if its binding energy is 8.481184 MeV/nucleon? (a) 30.97377 (b) 30.95322 (c) 31.25601 (d) 32.22563 (e) 30.13725 Answer: (a) 24. What is the binding energy per nucleon for 60Co? The exact mass of 60Co is 59.933822 amu. (a) between 6 and 7 MeV/nucleon (b) between 7 and 8 MeV/nucleon (c) between 8 and 9 MeV/nucleon (d) between 9 and 10 MeV/nucleon (e) more than 10 MeV/nucleon Answer: (c)

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25. In aging red giant stars, 12C undergoes several different fusion reactions, one of which is shown below. What is the energy released in this reaction in MeV? 2 12C → 23Na + 1H Nuclide 12C 23Na 1H

Exact Mass in amu 12.000000 22.989770 1.007825

(a) between and 0 and 1 MeV (b) between 1 and 2 MeV (c) between 2 and 3 MeV (d) between 3 and 4 MeV (e) more than 4 MeV Answer: (c)

The Kinetics of Radioactive Decay and Radioactive Dating 26. All forms of radioactive decay obey a ____ rate law. (a) Zero-order (b) First-order (c) Second-order (d) first or second-order (e) zeroth, first, or second-order Answer: (b)

27. A typical rate law for radioactive decay of a nuclide (N) would be: (a) rate = k (b) rate = k(N)1/2 (c) rate = k(N) (d) rate = k(N)2 (e) The rate law depends on the mode of radioactive decay.

Answer: (c)

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28. An important component of smoke detectors is the  emitter americium-241 which has a half-life of 432 years. What is the rate constant for the decay of 241Am? (a) 432 yr-1 (b) 2.31 x 10-3 yr-1 (c) 1.60 x 10-3 yr-1 (d) 0.693 yr-1 (e) none of the above Answer: (c) 29. An important component of smoke detectors is the  emitter 241Am. The rate constant for the decay of this isotope is 5.1 x 10-11 dis/s. What is the activity in Curies of the 241Am if each detector contains approximately 1.8 x 10-4 g of the isotope?

Answer: 6.2 x 10-4 Ci 30. An important component of smoke detectors is the  emitter 241Am. The rate constant for the decay of this isotope is 1.6 x 10-3 yr-1. If a smoke detector starts with1.8 x 10-4 g of the isotope, how long will it take until 1.0 x 10-4 g of americium remains? (a) 3.7 x 102 years (b) 3.2 x 102 year (c) 4.5 x 102 years (d) 32 years (e) none of the above Answer: (a) 31. 90Sr is a particularly dangerous nuclide formed from the fission of 235U in atomic bombs because it can readily enter the food chain by replacing calcium in plants or animals. If the rate constant for the decay of 90Sr is 0.0274 yr-1, approximately how long will it take for 99.0% of a sample of 90Sr to decay? (a) 168 years (b) 52.0 years (c) 452 years (d) 195 years (e) none of the above

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Answer: (a)

32. The cloth wrappings from a mummified bull found in a pyramid in Dashur, Egypt were dated by 14C analysis. The wrappings had 78.0% of the radioactivity of living tissue. Approximately how old are the wrappings? The rate constant for the decay of 14C is 1.21 x 10-4 yr-1. (a) less than 500 years old (b) between 500 and 1500 years old (c) between 1500 and 2500 years old (d) between 2500 and 3500 years old (e) more than 3500 years old Answer: (c) 33. Which of the following is not an important assumption made in 14C dating? (a) 14C is produced in the atmosphere at a more or less constant rate. (b) A carbon atom typically forms four bonds. (c) Carbon atoms circulate among the atmosphere, the oceans, and living organisms faster than they decay. (d) After death, organisms no longer pick up 14C. (e) All of the above are important assumptions in 14C dating.

Answer: (b)

34. Water at various depths in the ocean is recirculated to the surface at different rates. The length of time elapsed since water has been brought up to the surface of the ocean can be estimated by measuring the radioactivity of the 14C in the CO2 dissolved in the water. If a sample of water brought up from a depth of 1000 meters displays 88.5% of the radioactivity of water at the surface, how “old” is the water? The rate constant for the decay of 14C is 1.21 x 10-4 yr-1. (a) between 900 and 1100 years (b) between 700 and 900 years (c) between 500 and 700 years (d) between 300 and 500 years (e) less than 300 years Answer: (a)

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35. The decay of 40K to 40Ar can be used to determine the age of very old rocks.

If moon rocks from the Sea of Tranquility have 12.5% of the estimated original amount of 40K, approximately how old are the rocks? 40K decays with a rate constant of 5.06 x 10-10 yr-1. (a) between 106 and 107 years (b) between 108 and 109 years (c) between 1010 and 1011 years (d) between 1012 and 1013 years (e) more than 1013 years Answer: (b)

Ionizing Radiation and the Biological Effects of Radiation 36. Which of the following is a free radical? (a) H2O (b) H3O+ (c) O2(d) OH(e) none of the above Answer: (c) 37. The energy required to ionize a single water molecule is 2.02 x 10 -18 J. What wavelength (in nm) photon would deliver this amount of energy? (a) 98.4 nm (b) 115 nm (c) 3.04 x 1015 nm (d) 536 nm (e) none of the above Answer: (a)

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Natural vs. Induced Radioactivity 38. The first synthesis of element number 116 (ununhexium) was accomplished in 2000 at the Joint Institute for Nuclear Research in Dubna. The Uuh was prepared in the reaction shown below. What is the mass number of the isotope of Uuh prepared?

248Cm

+ 48Ca

→

116Uuh

+ 4 1n

(a) 116 (b) 296 (c) 200 (d) 292 (e) none of the above Answer: (d)

39. Element 110, darnstatdium, was prepared in the reaction shown below.

After it was synthesized it decomposed by  decay with a half-life of 270 s.

What is the product obtained from the  decay?

208Pb

+ 62Ni

→

269Ds

+ 1n

(a) 265Hs (b) 265Sg (c) 269Hs (d) 108Bh (e) 267Bh Answer: (a) 40. The half-lives for four of the heaviest known elements are given below.

Does this data support the theoretical predictions that element 114 should be an “island of stability” in a group of very unstable nuclides? Explain.

Element 112Uub 114Uuq 115Uup 116Uuh

t1/2 240 milliseconds 30 seconds 46.6 milliseconds 47 milliseconds

Answer: Yes. The longer half-life for the decay of element 114 relative to the neighboring super-heavy elements indicates that it is significantly more stable. Although a half-life of 30 seconds is not

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very long, it is 125 times as long as the next most stable superheavy element (#112).

Nuclear Fission and Fusion 41. Nuclear fission is energetically favorable for ______ elements while nuclear fusion is energetically favorable for ______ elements. (a) the heavier, the lighter (b) the lighter, the heavier (c) metallic, lighter (d) heavier, p-block (e) low electronegativity, f-block Answer: (a)

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Chapter 16: Organic Chemistry Alkanes and Alkenes – Structure and Nomenclature 1.

Which of the following structures represents the same molecule?

CH3

H3C

H2 C H3C

H2 C

CH3

CH2 C H2

CH

CH3

CH3

H3C

1

CH3

CH3 CH

C H2

C CH3

2

3

H3C

CH3

4

(a) 1 and 2 (b) 3 and 4 (c) 2 and 3 (d) 1 and 4 (e) All of the structures show the same molecule Answer: (c)

2.

Draw the structure of 3-bromo-2-methylhexane.

Answer: CH3 H2 C

CH H3C

CH

CH3 C H2

Br

3. The correct systematic name for the compound shown below would be: CH3 CH3

H2C

CH H3C

CH C H2

CH3

(a) 2,4-dimethylhexane (b) 4-ethyl-2-methylpentane 16-1

(c) isopropylpentane (d) 2-methyl-4-ethylpentane (e) 2,3-dimethylhexane Answer: (a)

4. Ignoring stereochemistry, the correct systematic name for the compound shown below would be:

(a) isooctene (b) 3-methyl-2,3-heptene (c) 4-ethyl-2-pentene (d) 3-methylheptane (e) 4-methyl-2-heptene Answer: (e)

5. The compound shown below is geraniol. It has a rose-like scent and occurs naturally in a number of flowers. What does the electron domain model predict the geometry to be around the carbon labeled 1?

CH3

CH3 H2 C

C H3C

1

C H2

H2 C

C C H2

C H

OH

(a) linear (b) trigonal planar (c) tetrahedral (d) trigonal pyramidal (e) bent Answer: (b)

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6. Which of the following best describes the relationship between the two structures shown below?

(a) They represent the same compound. (b) They represent different compounds that are constitutional isomers. (c) They represent different compounds that are stereoisomers. (d) They represent different compounds that are alkenes. (e) They represent different compounds that are alkynes.

Answer: (a) 7.

Draw structures for all of the stable constitutional isomers of the formula C 4H9Cl.

Answer: Cl H2 C

Cl C H2

CH3 C H2

CH H3C

CH3

CH3 CH3 C H2

CH H3C

Cl C H2

H3C

C

Cl

CH3

Unsaturated Hydrocarbons 8. Which of the following formulas represents an unsaturated hydrocarbon? (a) C3H8 (b) C3H4 (c) C2H4 (d) C5H12 (e) Both b and c are unsaturated hydrocarbons.

Answer: (e) 9. What is the correct systematic name for the compound shown below?

CH3 HC

C

CH CH3

(a) 3-methyl-1-butyne

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(b) 3-methyl-1-butene (c) 2-methyl-3-butyne (d) isobutyne (e) 3,3-dimethyl-1-propyne Answer: (a) 10. A structure of the compound benzene (C6H6) is shown below. Are the carboncarbon in benzene all the same length? Explain your answer.

H H

C C C

H

H C C

C

H

H

Answer: Yes, all of the C-C bonds in benzene are the same length. Even though the resonance structure shown suggests that 3 of the bonds are C-C single bonds and 3 are C=C double bonds, benzene is a hybrid of the resonance structures shown below with all of the carbon-carbon bonds being equivalent.

H H

C C

H C

C H

H H

C C H

C C

H

C H

H C C

C

H

H

Functional Groups 11. The molecule shown below is vanillin – one of the principle components of vanilla. Which of the following functional groups is not present in the structure of vanillin?

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(a) aldehyde (b) ether (c) alcohol (d) ketone (e) All of these functional groups are present Answer: (d) 12. The molecule shown below is Tamiflu (®) which is used to treat some types of influenza. Circle and identify each functional group present in Tamiflu.

Answer:

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13. The compound shown below has an extremely strong smell of bananas. Which functional group is present in the molecule?

CH3 H3C

CH

C H2

H2 C

O O

C

CH3

(a) aldehyde (b) ketone (c) alkene (d) ester (e) ether Answer: (d) 14. The compound shown below is the over-the-counter analgesic acetaminophen.

Which of the following functional groups are present in the molecule?

H N

HO

O

(a) aldehyde (b) alcohol (c) amide (d) ester (e) Both alcohol and amide groups are present.

Answer: (e) 15. The compound shown below is the stimulant amphetamine. Which functional group is present in the molecule?

(a) amine (b) alcohol (c) amide (d) ester

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(e) Both alcohol and amide groups are present.

Answer: (a) Oxidation/Reduction Reactions 16. In which of the following reactions is a carbon atom not oxidized or reduced? (a) CH3CH2Br + OH- → CH3CH2OH + Br(b) CH2=CHCH3 + H2 → CH3CH2CH3 (c) CH3CH3 + ½O2(g) → CH3CH2OH (d) CH3CH2CHO + ½O2(g) → CH3CH2CO2H (e) A carbon atom is oxidized or reduced in all of these reactions.

Answer: (a)

17. What is the oxidation number of the carbon shown in bold in the reactant and in the product?

CH3CH2CHO + ½O2(g) → CH3CH2CO2H

(a) +3 in the reactant and +1 in the product. (b) 0 in the reactant and –2 in the product. (c) –2 in the reactant and 0 in the product. (d) +1 in the reactant and +3 in the product. (e) The carbon atom does not undergo a change in its formal oxidation state.

Answer: (d)

Alcohols 18. What is the correct systematic name for the molecule shown below?

H2 C

HO C H2

CH3 C H2

(a) 1-butenol (b) 1-pentanol (c) 1-butanol (d) isobutanol (e) 1-hydroxybutane

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Answer: (c)

19. Which of the following is the conjugate base of ethanol (CH3CH2OH)? (a) CH3CH2OH2+ (b) CH3CH2O(c) CH3CH3 (d) CH4CH2OH+ (e) CH3CH2CO2Answer: (b)

20. Which of the following alcohols would not be oxidized to the corresponding aldehyde or ketone?

a.

H2 C

HO C H2

CH3 C H2

OH

b.

CH

CH3

H3C

C H2

CH3 OH C

c.

H3C

CH3

H2 C

d.

H3C

CH

OH

CH3

e. All of the alcohols could be oxidized to the corresponding aldehyde or ketone.

Answer: (c)

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Carbonyl Compounds 21. What is the most likely product of the reaction shown below?

Answer:

22. Arrange the following compounds in order from most acidic to least acidic. Explain your answer.

H2 C H3C

H2 C

OH H3C

C H2

H2 C

O C

CH3

H3C

O

O

H2 C H3C

Answer:

H2 C

OH C

<

OH C

H3C

H2 C

OH C H2

O

<

H3C

O C

CH3

O

The carboxylic acid is a reasonably strong weak acid because the anion formed by deprotonation is stabilized by resonance between the two oxygens. The alcohol can be deprotonated leaving a moderately stable anion localized on the single oxygen. The ester is an extremely weak acid because the only H’s in the compound are bonded to carbon and have the extremely low acidity associated with alkanes.

23. Predict the products of the following reaction:

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H3C

CH3

OH H3C

C

+

H2C

H2C N C H2

O

H3C

CH3

O C O

Answer:

CH3

H3C

+

H2C

H2C N H2C

H

CH3

24. The “fishy” odor in cooked fish, which is due to the presence of amines, can be reduced by the addition of lemon juice which contains citric acid. Explain.

Answer: The amines (which are basic) react with the citric acid. The protonated amines (ammonium ions) are now cations which are much less volatile then the neutral amines.

Constitutional and Stereoisomers 25. Which of the following best describes the relationship between the two structures shown below?

(a) The represent the same compound. (b) The represent different compounds that are constitutional isomers. (c) The represent different compounds that are stereoisomers. (d) The represent different compounds that are alkanes. (e) They represent different compounds that are alkynes.

Answer: (c)

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26. Which of the following compounds would have a stereoisomer(s)?

(a) 1 only (b) 2 only (c) 3 only (d) 2 and 3 (e) 3 and 4 Answer: (c)

27. The compound labeled 1 is geraniol. It has a rose-like scent and occurs naturally in a number of flowers. Compounds 2 and 3 are isomers of geraniol. Which statement is correct?

(a) 2 is a constitutional isomer of geraniol and 3 is a stereoisomer. (b) 3 is a constitutional isomer of geraniol and 2 is a stereoisomer. (c) Both 2 and 3 are stereoisomers of geraniol. (d) Both 2 and 3 are constitutional isomer of geraniol (e) 2 and 3 are not isomers of geraniol.

Answer: (a) 28. Which of the molecules shown below is chiral?

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a.

H2 C

HO C H2

CH3 C H2

OH

b.

CH

CH3

H3C

C H2

CH3 OH C

c.

H3C

CH3

H2 C

d.

H3C

CH

OH

CH3

e. All of these compounds are chiral.

Answer: (b) 29. The molecule shown below is Tamiflu ® which is used to treat some types of influenza. Circle each carbon in the molecule that is a chiral center.

Answer:

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Misc. Questions 30. One example of a Fisher-Tropsch reaction is shown below. Calculate the Hrxn° using enthalpies of atom combination.

2CO(g) + 5H2(g) → C2H6(g) + 2H2O(g) Is this reaction endothermic or exothermic?

Compound CO(g) H2O(g) C2H6(g) H2(g)

Hac° (kJ/molrxn) -1076 -926 -2824 -435

Answer: -349 kJ/molrxn; it is exothermic.

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