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6: Percent Composition of Compounds

6: Percent Composition of Compounds

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Chapter Review

c. The reaction is performed in a metal cylinder fitted with a piston, and the piston is compressed to

decrease the total volume of the system.

d. Additional O2(g) is added to the system from a

cylinder of pure O2.

36. Consider the general reaction

2A(g) ϩ B(s) 4

3 C(g) ϩ 3D(g)



⌬H ϭ ϩ115 kJ/mol



which has already come to equilibrium. Predict

whether the equilibrium will shift to the left, will

shift to the right, or will not be affected if the changes

indicated below are made to the system.

a. Additional B(s) is added to the system.

b. C(g) is removed from the system as it forms.

c. The volume of the system is decreased by a factor

of 2.

d. The temperature is increased.

37. Hydrogen gas and chlorine gas in the presence of

light react explosively to form hydrogen chloride

H2(g) ϩ Cl2(g) 4

3 2HCl( g)

The reaction is strongly exothermic. Would an increase in temperature for the system tend to favor or

disfavor the production of hydrogen chloride?

38. For the general reaction

A(g) ϩ B(g) ϩ heat 4

3 C(g)

would an increase in temperature tend to favor the

forward or the reverse process? Why?

39. The reaction

3 C2H2Br4( g)

C2H2( g) ϩ 2Br2( g) 4

is exothermic in the forward direction. Will an increase in temperature shift the position of the equilibrium toward reactants or products?

40. The reaction

4NO( g) ϩ 6H2O( g) 4

3 4NH3( g) ϩ 5O2( g)

is strongly endothermic. Will an increase in temperature shift the equilibrium position toward products

or toward reactants?

41. Plants synthesize the sugar dextrose according to the

following reaction by absorbing radiant energy from

the sun (photosynthesis).

6CO2( g) ϩ 6H2O( g) 4

3 C6H12O6(s) ϩ 6O2( g)

Will an increase in temperature tend to favor or discourage the production of C6H12O6(s)?

42. Consider the exothermic reaction

3 CH3OH(l)

CO(g) ϩ 2H2(g) 4

Predict three changes that could be made to the system that would increase the yield of product over

that produced by a system in which no change was

made.



575



17.8 Applications Involving the Equilibrium

Constant

QUESTIONS

43. Suppose a reaction has the equilibrium constant K ϭ

1.3 ϫ 108. What does the magnitude of this constant

tell you about the relative concentrations of products

and reactants that will be present once equilibrium is

reached? Is this reaction likely to be a good source of

the products?

44. Suppose a reaction has the equilibrium constant K ϭ

1.7 ϫ 10Ϫ8 at a particular temperature. Will there be a

large or small amount of unreacted starting material

present when this reaction reaches equilibrium? Is

this reaction likely to be a good source of products at

this temperature?

PROBLEMS

45. For the reaction

Br2(g) ϩ 5F2(g) 4

3 2BrF5(g)

the system at equilibrium at a particular temperature is

analyzed, and the following concentrations are found:

[BrF5(g)] ϭ 1.01 ϫ 10Ϫ9 M, [Br2(g)] ϭ 2.41 ϫ 10Ϫ2 M,

and [F2(g)] ϭ 8.15 ϫ 10Ϫ2 M. Calculate the value of K

for the reaction at this temperature.

46. Consider the reaction

SO2(g) ϩ NO2(g) 4

3 SO3(g) ϩ NO(g)

Suppose it is found at a particular temperature that

the concentrations in the system at equilibrium are

as follows: [SO3(g)] ϭ 4.99 ϫ 10Ϫ5 M, [NO( g)] ϭ

6.31 ϫ 10Ϫ7 M, [SO2(g)] ϭ 2.11 ϫ 10Ϫ2 M, and

[NO2(g)] ϭ 1.73 ϫ 10Ϫ3 M. Calculate the value of K for

the reaction at this temperature.

47. For the reaction

2CO( g) ϩ O2( g) 4

3 2CO2( g)

it is found at equilibrium at a certain temperature

that the concentrations are [CO(g)] ϭ 2.7 ϫ 10Ϫ4 M,

[O2(g)] ϭ 1.9 ϫ 10Ϫ3 M, and [CO2(g)] ϭ 1.1 ϫ 10Ϫ1 M.

Calculate K for the reaction at this temperature.

48. For the reaction

CO2(g) ϩ H2(g) 4

3 CO(g) ϩ H2O(g)

the equilibrium constant, K, has the value 5.21 ϫ 10Ϫ3

at a particular temperature. If the system is analyzed at

equilibrium at this temperature, it is found that

[CO(g)] ϭ 4.73 ϫ 10Ϫ3 M, [H2O(g)] ϭ 5.21 ϫ 10Ϫ3 M,

and [CO2(g)] ϭ 3.99 ϫ 10Ϫ2 M. What is the equilibrium

concentration of H2(g) in the system?

49. The equilibrium constant for the reaction

3 2HF( g)

H2( g) ϩ F2( g) 4

has the value 2.1 ϫ 103 at a particular temperature.

When the system is analyzed at equilibrium at this



All even-numbered Questions and Problems have answers in the back of this book and solutions in the Solutions Guide.



576 Chapter 17 Equilibrium

temperature, the concentrations of both H2(g) and

F2(g) are found to be 0.0021 M. What is the concentration of HF( g) in the equilibrium system under

these conditions?

50. For the reaction

3 2H2( g) ϩ O2( g)

2H2O( g) 4

Ϫ3



K ϭ 2.4 ϫ 10 at a given temperature. At equilibrium

it is found that [H2O(g)] ϭ 1.1 ϫ 10Ϫ1 M and [H2(g)] ϭ

1.9 ϫ 10Ϫ2 M. What is the concentration of O2(g) under these conditions?

51. For the reaction



58. Write the balanced chemical equation describing the

dissolving of each of the following sparingly soluble

salts in water. Write the expression for Ksp for each

process.

a. Bi2S3(s)

b. Ca(OH)2(s)



c. Co(OH)3(s)

d. Cu2S(s)



59. Ksp for copper(II) hydroxide, Cu(OH)2, has a value

2.2 ϫ 10Ϫ20 at 25 °C. Calculate the solubility of copper(II) hydroxide in mol/L and g/L at 25 °C.

60. Ksp for magnesium carbonate, MgCO3, has a value

3.5 ϫ 10Ϫ8 at 25 °C. Calculate the solubility of magnesium carbonate in mol/L and g/L at 25 °C.



3 2O3(g)

3O2(g) 4

The equilibrium constant, K, has the value 1.12 ϫ 10Ϫ54

at a particular temperature.

a. What does the very small equilibrium constant indicate about the extent to which oxygen gas,

O2(g), is converted to ozone gas, O3(g), at this temperature?

b. If the equilibrium mixture is analyzed and [O2(g)]

is found to be 3.04 ϫ 10Ϫ2 M, what is the concentration of O3(g) in the mixture?

52. For the reaction

3 2NO2( g)

N2O4( g) 4

the equilibrium constant K has the value 8.1 ϫ 10Ϫ3

at a particular temperature. If the concentration of

NO2(g) is found to be 0.0021 M in the equilibrium

system, what is the concentration of N2O4(g) under

these conditions?



17.9 Solubility Equilibria

QUESTIONS

53. Explain how the dissolving of an ionic solute in water represents an equilibrium process.

54. What is the special name given to the equilibrium

constant for the dissolving of an ionic solute in water?

55. Why does the amount of excess solid solute present

in a solution not affect the amount of solute that ultimately dissolves in a given amount of solvent?

56. Which of the following will affect the total amount of

solute that can dissolve in a given amount of solvent?

a. The solution is stirred.

b. The solute is ground to fine particles before dissolving.

c. The temperature changes.

PROBLEMS

57. Write the balanced chemical equation describing the

dissolving of each of the following sparingly soluble

salts in water. Write the expression for Ksp for each

process.

a. AgIO3(s)

b. Sn(OH)2(s)



c. Zn3(PO4)2(s)

d. BaF2(s)



61. A saturated solution of nickel(II) sulfide contains approximately 3.6 ϫ 10Ϫ4 g of dissolved NiS per liter at

20 °C. Calculate the solubility product Ksp for NiS at

20 °C.

62. Most hydroxides are not very soluble in water. For

example, Ksp for nickel(II) hydroxide, Ni(OH)2, is

2.0 ϫ 10Ϫ15 at 25 °C. How many grams of nickel(II)

hydroxide dissolve per liter at 25 °C?

63. The solubility product constant, Ksp, for calcium

carbonate at room temperature is approximately

3.0 ϫ 10Ϫ9. Calculate the solubility of CaCO3 in

grams per liter under these conditions.

64. Calcium sulfate, CaSO4, is only soluble in water to the

extent of approximately 2.05 g/L at 25 °C. Calculate

Ksp for calcium sulfate at 25°C.

65. Approximately 1.5 ϫ 10Ϫ3 g of iron(II) hydroxide,

Fe(OH)2(s), dissolves per liter of water at 18 °C. Calculate Ksp for Fe(OH)2(s) at this temperature.

66. Chromium(III) hydroxide dissolves in water only to

the extent of 8.21 ϫ 10Ϫ5 M at 25 °C. Calculate Ksp for

Cr(OH)3 at this temperature.

67. Magnesium fluoride dissolves in water to the extent

of 8.0 ϫ 10Ϫ2 g/L at 25 °C. Calculate the solubility of

MgF2(s) in moles per liter, and calculate Ksp for MgF2

at 25 °C.

68. Lead(II) chloride, PbCl2(s), dissolves in water to the

extent of approximately 3.6 ϫ 10Ϫ2 M at 20 °C. Calculate Ksp for PbCl2(s), and calculate its solubility in

grams per liter.

69. Mercury(I) chloride, Hg2Cl2, was formerly administered orally as a purgative. Although we usually think

of mercury compounds as highly toxic, the Ksp of

mercury(I) chloride is small enough (1.3 ϫ 10Ϫ18)

that the amount of mercury that dissolves and enters

the bloodstream is tiny. Calculate the concentration

of mercury(I) ion present in a saturated solution of

Hg2Cl2.

70. The solubility product of iron(III) hydroxide is very

small: Ksp ϭ 4 ϫ 10Ϫ38 at 25 °C. A classical method of

analysis for unknown samples containing iron is to

add NaOH or NH3. This precipitates Fe(OH)3, which



All even-numbered Questions and Problems have answers in the back of this book and solutions in the Solutions Guide.



Chapter Review

can then be filtered and weighed. To demonstrate

that the concentration of iron remaining in solution

in such a sample is very small, calculate the solubility

of Fe(OH)3 in moles per liter and in grams per liter.



Additional Problems

71. Before two molecules can react, chemists envision

that the molecules must first collide with one another.

Is collision among molecules the only consideration

for the molecules to react with one another?

72. Why does an increase in temperature favor an increase in the speed of a reaction?

73. The minimum energy required for molecules to react

with each other is called the

energy.

74. A(n)

consumed.



speeds up a reaction without being



75. Equilibrium may be defined as the

of two

processes, one of which is the opposite of the other.

76. When a chemical system has reached equilibrium,

the concentrations of all reactants and products remain

with time.

77. What does it mean to say that all chemical reactions

are, to one extent or another, reversible?

78. What does it mean to say that chemical equilibrium

is a dynamic process?

79. At the point of chemical equilibrium, the rate of the

forward reaction

the rate of the reverse reaction.



87. Many sugars undergo a process called mutarotation, in

which the sugar molecules interconvert between two

isomeric forms, finally reaching an equilibrium between them. This is true for the simple sugar glucose,

C6H12O6, which exists in solution in isomeric forms

called alpha-glucose and beta-glucose. If a solution of

glucose at a certain temperature is analyzed, and it is

found that the concentration of alpha-glucose is twice

the concentration of beta-glucose, what is the value of

K for the interconversion reaction?

88. Suppose K ϭ 4.5 ϫ 10Ϫ3 at a certain temperature for

the reaction

PCl5( g) 4

3 PCl3( g) ϩ Cl2( g)

If it is found that the concentration of PCl5 is twice

the concentration of PCl3, what must be the concentration of Cl2 under these conditions?

89. For the reaction

CaCO3(s) 4

3 CaO(s) ϩ CO2( g)

the equilibrium constant K has the form K ϭ [CO2].

Using a handbook to find density information about

CaCO3(s) and CaO(s), show that the concentrations of

the two solids (the number of moles contained in 1 L

of volume) are constant.

90. As you know from Chapter 7, most metal carbonate

salts are sparingly soluble in water. Below are listed

several metal carbonates along with their solubility

products, Ksp. For each salt, write the equation showing the ionization of the salt in water, and calculate

the solubility of the salt in mol/L.

Salt



Ksp



80. Equilibria involving reactants or products in more

than one state are said to be

.



BaCO3



5.1 ϫ 10Ϫ9



CdCO3



5.2 ϫ 10Ϫ12



81. According to Le Châtelier’s principle, when a large excess of a gaseous reactant is added to a reaction system

at equilibrium, the amounts of products

.



CaCO3



2.8 ϫ 10Ϫ9



CoCO3



1.5 ϫ 10Ϫ13



82. Addition of an inert substance (one that does not participate in the reaction) does not change the

of an equilibrium.

83. When the volume of a vessel containing a gaseous

equilibrium system is decreased, the

of the

gaseous substances present is initially increased.

84. Why does increasing the temperature for an exothermic process tend to favor the conversion of products

back to reactants?

85. What is meant by the solubility product for a sparingly

soluble salt? Choose a sparingly soluble salt and show

how the salt ionizes when dissolved in water, and

write the expression for its solubility product.

86. For a given reaction at a given temperature, the special ratio of products to reactants defined by the equilibrium constant is always equal to the same number.

Explain why this is true, no matter what initial concentrations of reactants (or products) may have been

taken in setting up an experiment.



577



91. Teeth and bones are composed, to a first approximation, of calcium phosphate, Ca3(PO4)2(s). The Ksp for

this salt is 1.3 ϫ 10Ϫ32 at 25 °C. Calculate the concentration of calcium ion in a saturated solution of

Ca3(PO4)2.

92. Under what circumstances can we compare the solubilities of two salts by directly comparing the values

of their solubility products?

93. How does the collision model account for the fact

that a reaction proceeds faster when the concentrations of the reactants are increased?

94. How does an increase in temperature result in an increase in the number of successful collisions between

reactant molecules? What does an increase in temperature mean on a molecular basis?

95. Explain why the development of a vapor pressure

above a liquid in a closed container represents an

equilibrium. What are the opposing processes? How

do we recognize when the system has reached a state

of equilibrium?



All even-numbered Questions and Problems have answers in the back of this book and solutions in the Solutions Guide.



578 Chapter 17 Equilibrium

96. Write the equilibrium expression for each of the following reactions.

a. H2(g) ϩ Br2(g) 4

3 2HBr(g)

b. 2H2(g) ϩ S2(g) 4

3 2H2S(g)

c. H2(g) ϩ C2N2(g) 4

3 2HCN( g)

97. Write the equilibrium expression for each of the following reactions.

a. 2O3(g) 4

3 3O2(g)

b. CH4(g) ϩ 2O2(g) 4

3 CO2(g) ϩ 2H2O(g)

c. C2H4(g) ϩ Cl2(g) 4

3 C2H4Cl2(g)

98. At high temperatures, elemental bromine, Br2, dissociates into individual bromine atoms.

Br2( g) 4

3 2Br( g)

Suppose that in an experiment at 2000 °C, it is found

that [Br2] ϭ 0.97 M and [Br] ϭ 0.034 M at equilibrium. Calculate the value of K.

99. Gaseous phosphorus pentachloride decomposes according to the reaction

3 PCl3( g) ϩ Cl2( g)

PCl5( g) 4

The equilibrium system was analyzed at a particular

temperature, and the concentrations of the substances

present were determined to be [PCl5] ϭ 1.1 ϫ 10Ϫ2 M,

[PCl3] ϭ 0.325 M, and [Cl2] ϭ 3.9 ϫ 10Ϫ3 M. Calculate

the value of K for the reaction.

100. Write the equilibrium expression for each of the following heterogeneous equilibria.

a. 4Al(s) ϩ 3O2(g) 4

3 2Al2O3(s)

b. NH3(g) ϩ HCl(g) 4

3 NH4Cl(s)

c. 2Mg(s) ϩ O2(g) 4

3 2MgO(s)

101. Write the equilibrium expression for each of the following heterogeneous equilibria.

a. P4(s) ϩ 5O2(g) 4

3 P4O10(s)

b. CO2(g) ϩ 2NaOH(s) 4

3 Na2CO3(s) ϩ H2O(g)

c. NH4NO3(s) 4

3 N2O(g) ϩ 2H2O(g)

102. What is the effect on the position of a reaction system at equilibrium when an exothermic reaction is

performed at a higher temperature? Does the value of

the equilibrium constant change in this situation?

103. Suppose the reaction system

3 2NO2( g)

2NO( g) ϩ O2( g) 4

has already reached equilibrium. Predict the effect of

each of the following changes on the position of the

equilibrium. Tell whether the equilibrium will shift

to the right, will shift to the left, or will not be affected.

a. Additional oxygen is injected into the system.

b. NO2 is removed from the reaction vessel.

c. 1.0 mole of helium is injected into the system.

104. The reaction

PCl3(l) ϩ Cl2(g) 4

3 PCl5(s)



liberates 124 kJ of energy per mole of PCl3 reacted.

Will an increase in temperature shift the equilibrium

position toward products or toward reactants?

105. For the process

3 CO2( g) ϩ H2( g)

CO( g) ϩ H2O( g) 4

it is found that the equilibrium concentrations at

a particular temperature are [H2] ϭ 1.4 M, [CO2] ϭ

1.3 M, [CO] ϭ 0.71 M, and [H2O] ϭ 0.66 M. Calculate

the equilibrium constant K for the reaction under

these conditions.

106. For the reaction

3 2NH3( g)

N2( g) ϩ 3H2( g) 4

Ϫ2



K ϭ 1.3 ϫ 10 at a given temperature. If the system

at equilibrium is analyzed and the concentrations of

both N2 and H2 are found to be 0.10 M, what is the

concentration of NH3 in the system?

107. The equilibrium constant for the reaction

2NOCl( g) 4

3 2NO( g) ϩ Cl2( g)

has the value 9.2 ϫ 10Ϫ6 at a particular temperature.

The system is analyzed at equilibrium, and it is found

that the concentrations of NOCl(g) and NO( g) are

0.44 M and 1.5 ϫ 10Ϫ3 M, respectively. What is the

concentration of Cl2(g) in the equilibrium system under these conditions?

108. As you learned in Chapter 7, most metal hydroxides

are sparingly soluble in water. Write balanced chemical equations describing the dissolving of the following metal hydroxides in water. Write the expression

for Ksp for each process.

a. Cu(OH)2(s)

b. Cr(OH)3(s)



c. Ba(OH)2(s)

d. Sn(OH)2(s)



109. The three common silver halides (AgCl, AgBr, and

AgI) are all sparingly soluble salts. Given the values

for Ksp for these salts below, calculate the concentration of silver ion, in mol/L, in a saturated solution of

each salt.

Silver Halide



Ksp



AgCl



1.8 ϫ 10Ϫ10



AgBr



5.0 ϫ 10Ϫ13



AgI



8.3 ϫ 10Ϫ17



110. Approximately 9.0 ϫ 10Ϫ4 g of silver chloride, AgCl(s),

dissolves per liter of water at 10 °C. Calculate Ksp for

AgCl(s) at this temperature.

111. Mercuric sulfide, HgS, is one of the least soluble salts

known, with Ksp ϭ 1.6 ϫ 10Ϫ54 at 25 °C. Calculate the

solubility of HgS in moles per liter and in grams per

liter.

112. Approximately 0.14 g of nickel(II) hydroxide,

Ni(OH)2(s), dissolves per liter of water at 20 °C. Calculate Ksp for Ni(OH)2(s) at this temperature.



All even-numbered Questions and Problems have answers in the back of this book and solutions in the Solutions Guide.



Chapter Review

113. For the reaction N2(g) ϩ 3H2(g) S 2NH3(g), list the

types of bonds that must be broken and the type of

bonds that must form for the chemical reaction to

take place.

114. What does the activation energy for a reaction represent? How is the activation energy related to whether

a collision between molecules is successful?

115. What are the catalysts in living cells called? Why are

these biological catalysts necessary?

116. When a reaction system has reached chemical equilibrium, the concentrations of the reactants and

products no longer change with time. Why does the

amount of product no longer increase, even though

large concentrations of the reactants may still be

present?

117. Ammonia, a very important industrial chemical, is

produced by the direct combination of the elements

under carefully controlled conditions.



579



N2(g) ϩ 3H2(g) 4

3 2NH3(g)

Suppose, in an experiment, that the reaction mixture is

analyzed after equilibrium is reached and it is found,

at a particular temperature, that [NH3(g)] ϭ 0.34 M,

[H2(g)] ϭ 2.1 ϫ 10Ϫ3 M, and [N2(g)] ϭ 4.9 ϫ 10Ϫ4 M.

Calculate the value of K at this temperature.

118. Write the equilibrium expression for each of the following heterogeneous equilibria.

3 Li2CO3(s) ϩ H2O(g) ϩ CO2(g)

a. 2LiHCO3(s) 4

3 PbO(s) ϩ CO2(g)

b. PbCO3(s) 4

3 2Al2O3(s)

c. 4Al(s) ϩ 3O2(g) 4

119. Suppose a reaction has the equilibrium constant K ϭ

4.5 ϫ 10Ϫ6 at a particular temperature. If an experiment is set up with this reaction, will there be large

relative concentrations of products present at equilibrium? Is this reaction useful as a means of producing the products? How might the reaction be made

more useful?



All even-numbered Questions and Problems have answers in the back of this book and solutions in the Solutions Guide.



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