Summary Problems to accompany
principles of
chemistry
THE MOLECULAR SCIENCE
Moore • Stanitski • Jurs
Principles of Chemistry: The Molecular Science
John W. Moore, Conrad L. Stanitski, Peter C. Jurs
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About the Cover
As you look carefully at the photo on the front cover of this book, you will notice that
each spherical water droplet hanging from a blade of grass acts as a lens and shows
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happening at the molecular scale (nanoscale). In the case of water droplets, what we
observe has to do with the shape and orientation of water molecules. Water droplets
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interconnected puckered hexagonal rings. This gives water other interesting
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range of natural phenomena—from the cosmos to a cell's nucleus to atoms and
molecules—in which chemistry plays a major role.
Printed in the United States of America
1 2 3 4 5 6 7 12 11 10 09
1
CHAPTER 2 SUMMARY PROBLEM
The atoms of one of the elements contain 47 protons and 62 neutrons.
(a) Identify the element and give its symbol.
(b) What is this atom’s atomic number? Mass number?
(c) This element has two naturally occurring isotopes. Calculate the atomic weight of
the element.
Isotope
1
2
Mass
Number
Percent
Abundance
Isotopic Mass
(amu)
107
109
51.84
48.16
106.905
108.905
(d) This element is a member of which group in the periodic table? Is this element a
metal, nonmetal, or metalloid? Explain your answer.
(e) Consider a piece of jewelry that contains 1.00 g of the element. (i) How many moles
of the element are in this mass? (ii) How many atoms of the element are in this
mass? (iii) Atoms of this element have an atomic diameter of 304 pm. If all the atoms
of this element in the sample were put into a row, how many meters long would the
chain of atoms be?
© 2010 Brooks/Cole, Cengage Learning
2
CHAPTER 3 SUMMARY PROBLEM
Part I
During each launch of the Space Shuttle, the booster rocket uses about 1.5 ⫻ 106 lb of ammonium perchlorate as fuel.
1. Write the chemical formulas for (a) ammonium perchlorate, (b) ammonium chlorate,
and (c) ammonium chlorite.
2. (a) When ammonium perchlorate dissociates in water, what ions are dispersed in the
solution?
(b) Would this aqueous solution conduct an electric current? Explain your answer.
3. How many moles of ammonium perchlorate are used in Space Shuttle booster rockets
during a launch?
Part II
Chemical analysis of ibuprofen (Advil) indicates that it contains 75.69% carbon, 8.80% hydrogen, and the remainder oxygen. The empirical formula is also the molecular formula.
1. Determine the molecular formula of ibuprofen.
2. Two 200.-mg ibuprofen tablets were taken by a patient to relieve pain. Calculate the
number of moles of ibuprofen contained in the two tablets.
© 2010 Brooks/Cole, Cengage Learning
3
CHAPTER 4 SUMMARY PROBLEM
Iron can be smelted from iron(III) oxide in ore via this high-temperature reaction in a blast
furnace:
Fe2O3 (s) ⫹ 3 CO(g) 9: 2 Fe( ᐉ) ⫹ 3 CO2 (g)
The liquid iron produced is cooled and weighed.
(a) For 19.0 g Fe2O3, what mass of CO is required to react completely?
(b) What mass of CO2 is produced when the reaction runs to completion with 10.0 g
Fe2O3 as starting material?
Mass Fe (g)
When the reaction was run repeatedly with the same mass of iron oxide, 19.0 g Fe2O3, but differing masses of carbon monoxide, this graph was obtained.
20
18
16
14
12
10
8
6
4
2
0
0
2
4
6
8 10 12 14 16 18 20
Mass CO (g)
(c) Which reactant is limiting in the part of the graph where there is less than 10.0 g
CO available to react with 19.0 g Fe2O3?
(d) Which reactant is limiting when more than 10.0 g CO is available to react with
19.0 g Fe2O3?
(e) If 24.0 g Fe2O3 reacted with 20.0 g CO and 15.9 g Fe was produced, what was the
percent yield of the reaction?
© 2010 Brooks/Cole, Cengage Learning
4
CHAPTER 5 SUMMARY PROBLEM
Gold in its elemental state can be separated from gold-bearing rock by treating the ore with
cyanide, CN⫺, in the presence of oxygen via this reaction:
⫺
4 Au(s) ⫹ 8 CN⫺ (aq) ⫹ O2 ( g ) ⫹ 2 H2O(ᐉ ) 9: 4 Au(CN) ⫺
2 (aq) ⫹ 4 OH (aq)
The CN⫺ is supplied by NaCN, but Na⫹ is a spectator ion and is left out of the net ionic
equation.
(a) Which reactant is oxidized? What are the oxidation numbers of this species as a
reactant and as a product?
(b) Which reactant is reduced? What are the oxidation numbers of this species as a reactant and as a product?
(c) What is the oxidizing agent?
(d) What is the reducing agent?
(e) What mass of NaCN would it take to prepare 1.0 L of 0.075 M NaCN?
(f ) If the ore contains 0.019% gold by weight, what mass of gold is found in one metric
ton (exactly 1000 kg) of the ore?
(g) How many grams of NaCN would you require to extract the gold in one metric ton
of this ore?
(h) How many liters of the 0.075 M NaCN solution would you require to extract the
gold in one metric ton of this ore?
© 2010 Brooks/Cole, Cengage Learning
5
CHAPTER 6 SUMMARY PROBLEM
Sulfur dioxide, SO2, is a major pollutant emitted by coal-fired electric power generating plants.
A large power plant can produce 8.64 ⫻ 1013 kJ electrical energy every day by burning 7000.
ton coal (1 ton ⫽ 9.08 ⫻ 105 g).
(a) When 1.00 g coal is burned in air, 33 kJ is transferred to the surroundings. Calculate
the quantity of energy transferred to the surroundings by the coal combustion reaction in a plant that burns 7000. ton coal.
(b) What is the efficiency of this power plant in converting chemical energy to electrical energy—that is, what percentage of the thermal energy transfer shows up as
electrical energy?
(c) When SO2 is given off by a power plant, it can be trapped by reaction with MgO in
the smokestack to form MgSO4.
MgO(s) ⫹ SO2 (g ) ⫹ 12 O2 (g) 9: MgSO4 (s)
If 140. ton SO2 is given off by a coal-burning power plant each day, how much MgO
would you have to supply to remove all of this SO2? How much MgSO4 would be
produced?
(d) How much heat transfer does the reaction in part (c) add or take away from the heat
transfer of the coal combustion?
(e) Sulfuric acid comes from the oxidation of sulfur, first to SO2 and then to SO3. The
SO3 is then absorbed by water in 98% H2SO4 solution to make H2SO4.
S(s) ⫹ O2 ( g) 9: SO2 ( g)
SO2 (g ) ⫹ 12 O2 (g) 9: SO3 (g)
SO3 ( g ) ⫹ H2O (in 98% H2SO4 ) 9: H2SO4 ( ᐉ)
For which of these reactions can you calculate ⌬H° using data in Table 6.2 or Appendix J? Do the calculation for each case where data are available.
© 2010 Brooks/Cole, Cengage Learning
6
CHAPTER 7 SUMMARY PROBLEM
(a) Without looking in the chapter, draw and label the first five energy levels of a hydrogen atom. Next, indicate the 2 : 1, 3 : 1, 5 : 2, and 4 : 3 transitions in a hydrogen atom. Look in Table 7.2 to get the measured wavelengths and spectral regions
for these transitions. Now calculate the frequencies (␯) for these transitions. Next,
calculate the energies of the photons that are produced in these transitions.
(b) The photoelectric threshold is the minimum wavelength a photon must have to produce a photoelectric effect for a metal. These three metals exhibit photoelectric
effects when photons of sufficient energies strike their surfaces.
Element
Photoelectric
Threshold (nm)
Lithium
Potassium
Cesium
540
550
660
Which photon energies calculated in part (a) would be sufficient to cause a photoelectric effect in lithium, in potassium, and in cesium?
(c) Vanadium, V, is a transition element named after Vanadis, the Scandinavian goddess
of beauty, because of the range of beautiful colors observed among vanadium compounds. Vanadium forms a series of oxides in which vanadium ions are V2⫹ (in VO),
V3⫹ (in V2O3), V4⫹ (in VO2), and V5⫹ (in V2O5).
(i) Write the full electron configuration for a vanadium atom.
(ii) Using the noble gas notation, write the electron configuration for a V2⫹ ion.
(iii) Write the orbital box diagram for a V3⫹ ion.
(iv) How many unpaired electrons are in the V4⫹ ion?
(v) Write the set of four quantum numbers for each valence electron in a V2⫹ ion.
(vi) Is the V5⫹ ion paramagnetic? Explain.
(d) An element is brittle with a steel-gray appearance. It is a relatively poor electrical
conductor and forms a volatile molecular chloride and a molecular hydride that
decomposes at room temperature. Is the element an alkaline earth metal, a transition
metal, a metalloid, or a halogen? Explain your answer.
(e) Give two ways in which p orbitals and d orbitals are (i) alike; (ii) different.
© 2010 Brooks/Cole, Cengage Learning
7
CHAPTER 8 SUMMARY PROBLEM
Glycine, C2H5NO2, is a naturally occurring amino acid with this Lewis structure,
H COC
"
H
N9C9C9O9 H
H
H
that contains the amine functional group, !NH2, and a carboxylic acid functional group,
!COOH.
(a) List the C!H, O!H, C!N, N!H, C!C, C!O, and C"O bonds in glycine in
order of increasing bond length.
(b) List the C!H, O!H, C!N, C!C, C!O, and C"O bonds in glycine in order of
increasing bond strength.
(c) List the C!N, C!O, C!H, O!H, and C!C bonds in glycine in order of increasing
bond polarity.
(d) The ionization of glycine in solution forms H⫹ and glycinate ions. Write the structural formula for glycinate ion.
(e) Glycinate ion is an example of a resonance hybrid. Write two Lewis structures that
contribute to this resonance hybrid.
(f) Glycine also forms a zwitterion. To form the zwitterion, an H⫹ is released from the
carboxylic acid group and attaches to the nitrogen of the amine group. Write the
Lewis structure of the zwitterion form of glycine.
(g) Using the Lewis structure drawn for part (f ), calculate the formal charges on each of
its atoms. What is the net charge of the zwitterion?
© 2010 Brooks/Cole, Cengage Learning
8
CHAPTER 9 SUMMARY PROBLEM
Part I
Write the Lewis structures and give the electron-pair geometry, the shape and bond angles,
and the hybridization of the central atom in these species:
(a)
(b)
(c)
(d)
(e)
OF2
NH2⫺
BI3
ClCN
the hybridization of each carbon atom in diacetylene, C4H2
Part II
Aspartame is a commonly used artificial sweetener (NutraSweet®) that was discovered accidentally by a chemist in the laboratory carelessly licking his fingers after synthesizing a new
compound (aspartame). The sweetener’s popularity stems from the fact that aspartame tastes
over 100 times sweeter than sugar. Aspartame has this structural formula.
H COC
H
H
H COC
N
C
C
N
H
C
H
H
O
C
O
H
H
C
C
C
H
H
O
C
H
H
(a)
(b)
(c)
(d)
(e)
Write the molecular formula of aspartame.
Identify the hybridization of each nitrogen atom.
Identify the hybridization of carbon in the CH2 groups.
Identify the hybridization of the carbon atom in the C"O group.
Identify the hybridization of the oxygen atom to which a methyl (!CH3) group is
attached.
(f ) Label each of the sigma and pi bonds involving carbon and those involving nitrogen.
Part III
The structural formula for the open-chain form of glucose is
H
O
OH H
H
H
C 9 C 9 C 9 C 9 C 9 C 9 OH
H
OH H
OH OH H
Glucose dissolves readily in water. Use molecular structure principles to explain why
glucose is so water-soluble.
© 2010 Brooks/Cole, Cengage Learning
9
CHAPTER 10 SUMMARY PROBLEM
The air that enters an automobile engine contains the oxygen that reacts with the hydrocarbon fuel vapors to provide the energy needed to move the vehicle. Prior to the combustion
process the fuel-air mixture is compressed, and then it is ignited with a spark. Most of the
fuel is completely burned to CO2 and H2O, and some of the nitrogen in air is converted to
nitrogen monoxide (NO). For simplicity, assume the fuel has a formula of C8H18 and a density
of 0.760 g/mL.
(a) What are the partial pressures of N2 and O2 in the air before it goes into the engine
if the atmospheric pressure is 734 mm Hg?
(b) If no fuel were added to the air and it was compressed to seven times atmospheric
pressure (approximately the compression ratio in a modern engine), what would the
partial pressures of N2 and O2 become?
(c) Assume the volume of each cylinder in the engine is 485 mL and that the temperature is 150 °C. If 0.050 mL of the fuel is added to the air in each cylinder just prior
to compression, and the fuel is completely vaporized, what would be the partial
pressure of the mixture in a cylinder due to the fuel molecules?
(d) How many moles of oxygen would be required to burn the fuel completely to CO2
and H2O?
(e) If 10.% of the nitrogen in the combustion process is converted to NO, calculate how
many grams of NO are produced in a single combustion reaction.
(f) Write an equation that shows the conversion of NO to a more photoreactive compound.
(g) Calculate the mass of the more photoreactive compound that would be formed from
the NO produced in part (e).
(h) What additional information would you need to calculate the NO emissions for an
entire city for a year?
© 2010 Brooks/Cole, Cengage Learning
10
CHAPTER 11 SUMMARY PROBLEM
Part I
Use the vapor pressure curves shown in Figure 11.5 to answer these questions.
(a)
(b)
(c)
(d)
What is the vapor pressure of diethyl ether at 0 °C?
Does diethyl ether have stronger or weaker intermolecular attractions than ethanol?
At what temperature does diethyl ether have a vapor pressure of 600 mm Hg?
If the normal boiling point of diethyl ether is 34.6 °C and it has a vapor pressure of
410 mm Hg at 20 °C, calculate ⌬H°vap for diethyl ether.
Part II
Consider the phase diagram for xenon shown below. Answer these questions.
Normal melting
point
Pressure (atm)
1.0
Solid
Normal
boiling
point
Liquid
0.5
Gas
0.37
Triple point
0
–125
–121 –120
–112
–115
Temperature (°C)
–110 –108
–105
(a) In what phase is xenon found at room temperature and a pressure of 1.0 atm?
(b) If the pressure exerted on a sample of xenon is 0.75 atm and the temperature is
⫺114 °C, in what phase does xenon exist?
(c) If the vapor pressure of a sample of liquid xenon is 375 mm Hg, what is the temperature of the liquid phase?
(d) What is the vapor pressure of solid xenon at ⫺122 °C?
(e) Which is the denser phase, solid or liquid? Explain.
Part III
Consider the CsCl unit cell shown in Figure 11.23.
(a) How many Cs⫹ ions are there per unit cell?
(b) How many Cl⫺ ions are there per unit cell?
(c) How many Cl⫺ ions share each face of the unit cell?
© 2010 Brooks/Cole, Cengage Learning
11
CHAPTER 12 SUMMARY PROBLEM
An excellent way to make highly pure nickel metal for use in specialized steel alloys is to
decompose Ni(CO)4 by heating it in a vacuum to slightly above room temperature.
Ni(CO) 4 (g) 9: Ni(s) ⫹ 4 CO(g)
The reaction is proposed to occur in four steps, the first of which is
Ni(CO) 4 (g) 9: Ni(CO) 3 ( g) ⫹ CO(g)
Kinetic studies of this first-order decomposition reaction have been carried out between
47.3 °C and 66.0 °C to give the results in this table.*
Temperature (°C)
Rate Constant (s⫺1)
47.3
50.9
55.0
60.0
66.0
0.233
0.354
0.606
1.022
1.873
(a) What is the activation energy for this reaction?
(b) Ni(CO)4 is formed by the reaction of nickel metal with carbon monoxide. If you
have 2.05 g CO and you combine it with 0.125 g nickel metal, what is the maximum
mass of Ni(CO)4 (in grams) that can be formed?
The replacement of CO by another molecule in Ni(CO)4 (in the nonaqueous solvents toluene and hexane) was also studied to understand the general principles that
govern the chemistry of such compounds.*
Ni(CO) 4 ⫹ P(CH3 ) 3 9: Ni(CO) 3P(CH3 ) 3 ⫹ CO
A detailed study of the kinetics of the reaction led to the mechanism
Ni(CO) 4 9: Ni(CO) 3 ⫹ CO
Ni(CO) 3 ⫹ P(CH3 ) 3 9: Ni(CO) 3P(CH3 ) 3
slow
fast
(c) Which step in the mechanism is unimolecular? Which is bimolecular?
(d) Add the steps of the mechanism to show that the result is the balanced equation for
the observed reaction.
(e) Is there an intermediate in this reaction? If so, what is it?
(f ) It was found that doubling the concentration of Ni(CO)4 increased the reaction rate
by a factor of 2. Doubling the concentration of P(CH3)3 had no effect on the reaction
rate. Based on this information, write the rate equation for the reaction.
(g) Does the experimental rate equation support the proposed mechanism? Why or
why not?
*See Day, J. P., Basolo, F., and Pearson, R. G. Journal of the American Chemical Society, Vol. 90, 1968;
p. 6933.
© 2010 Brooks/Cole, Cengage Learning
12
CHAPTER 13 SUMMARY PROBLEM
One approach to achieving cleaner-burning fuels and more-efficient automobiles is to extract
hydrogen from gasoline and other liquid fossil fuels. The extracted hydrogen could be combined with oxygen in fuel cells (Section 18.9) like those currently used in spacecraft to generate electricity. The electricity could be used in electric motors to power automobiles and to
provide air conditioning and other amenities expected by automobile buyers. Because electric
motors are far more efficient than current automobile engines, such a car might get 80 miles
per gallon of fuel. By answering the following questions, you can explore how the ideas of
chemical equilibrium and chemical kinetics can be applied to motive power for cars.
(a) The hydrogen extracted from hydrocarbon fuels must be free from soot (solid carbon)
and carbon monoxide, which would interfere with the operation of a fuel cell. Consider possible reactions by which hydrogen could be obtained from a hydrocarbon
such as octane (C8H18). Write an equation for a reaction that you think would not be
suitable, and write an equation for one that you think would be suitable. Explain your
choice in each case.
(b) Use data from Appendix J to calculate the change in enthalpy for each of the two reactions you wrote in Question 1. Predict whether entropy increases or decreases when
each reaction occurs. Is either of them ruled out because it is not product-favored? If
not, continue. If either of the reactions is not product-favored, suggest whether conditions could be altered to make it product-favored.
(c) The chemical process by which hydrogen is obtained for use in synthesizing ammonia
(Haber-Bosch process) involves treating methane (from natural gas) with steam. The
first step in this process is
CH4 ( g ) ⫹ H2O(g) EF CO(g) ⫹ 3 H2 ( g)
1.
2.
3.
4.
5.
Write the equilibrium constant expression for this reaction.
What is the relation between Kc and KP for this reaction?
Calculate the enthalpy change for this reaction.
Based on the equation, predict the sign of the entropy change for this reaction.
Is the reaction product-favored at high temperatures but not at lower temperatures? Or the other way around? Explain.
(d) To remove carbon monoxide from the hydrogen destined for the Haber-Bosch
process, this reaction is used:
CO(g) ⫹ H2O(g) EF CO2 (g) ⫹ H2 (g)
1. Use bond enthalpies to estimate the enthalpy change for this reaction.
2. At 450 °C, KP for this reaction is 6.48. Calculate Kc.
3. Suppose that 0.100 mol CO and 0.100 mol H2O were introduced into an empty
10.0-L flask at 450 °C. Determine the concentration of H2 (g) in the flask once
equilibrium has been achieved.
4. What is the concentration of CO(g) remaining in the flask in part (c)? Is it low
enough that we can say that the hydrogen is free of carbon monoxide?
(e) If you were in charge of designing a system for generating hydrogen gas for use in the
Haber-Bosch process, how might you obtain pure hydrogen? Assume that the process
is based on the two reactions given in Questions 3 and 4. Suggest a chemical reagent
that could be used to react with CO2 (g) and thereby remove it from the hydrogen generated as a product in each of the two reactions. Would this same reagent work if you
needed to remove SO2 (g)?
(f) To get the highest-purity H2 (g) and the maximum yield from the hydrogen-generating
process, what reactant concentration(s) would you increase or decrease? How would
you adjust product concentrations?
(g) In the hypothetical fuel cell system for an electric automobile, hydrocarbon fuel is
vaporized and partially oxidized in a limited quantity of air. In a second step the products of the first reaction are treated with steam over copper oxide and zinc oxide catalysts. In a final purification step, more air is introduced and a platinum catalyst helps
© 2010 Brooks/Cole, Cengage Learning
13
convert carbon monoxide to carbon dioxide. Write a balanced chemical equation for
each of these three steps. Assume that the hydrocarbon fuel is octane.
(h) What are the advantages of generating hydrogen gas from hydrocarbon fuel in an automobile rather than storing hydrogen in a fuel tank? What are the disadvantages of storing hydrogen in a car? What advantages does combining hydrogen with oxygen in a
fuel cell have, as opposed to burning a hydrocarbon fuel in an internal combustion
engine?
© 2010 Brooks/Cole, Cengage Learning
14
CHAPTER 14 SUMMARY PROBLEM
You are asked to prepare three mixtures at 25 °C.
Mixture I:
Mixture II:
Mixture III:
25.0 g CCl4 in 100. mL water
15.0 g CaCl2 in 125 mL water
21 g ethylene glycol (HOCH2CH2OH) in 150. mL water
Answer these questions about these mixtures. If one of the solutes fails to dissolve in water,
some of the questions will not be applicable.
(a) What is the weight percent of the mixture?
(b) What is the mass fraction of the mixture?
(c) Is a solution formed? (If a solution is formed, answer the remaining questions. You
may assume a density of the solution of 1.0 g/mL.)
(d) Name the dissolved species in solution and draw a diagram representing how the
solvent (water) molecules interact with these species.
(e) Express the concentration of the solution in ppm.
(f ) Express the concentration of the solution in molality.
(g) Calculate the vapor pressure of water in equilibrium with the solution.
(h) Calculate the boiling point of the solution.
(i) Calculate the freezing point of the solution.
(j) Calculate the osmotic pressure of the solution.
© 2010 Brooks/Cole, Cengage Learning
15
CHAPTER 15 SUMMARY PROBLEM
Lactic acid, CH3CH(OH)COOH, is a weak monoprotic acid with a melting point of 53 °C. The
acid exists in two forms that have slightly different Ka values. One form, call it Form A, has a
Ka of 1.6 ⫻ 10⫺4; the other form, call it Form B, has a Ka of 1.5 ⫻ 10⫺4. Form A is found in
molasses, beer, wines, and souring milk; Form B is produced by muscle cells during anaerobic
metabolism in which glucose is broken down into lactic acid molecules. When lactic acid
builds up too rapidly in muscle tissue, severe pain results.
(a) Which form of lactic acid (A or B) is the stronger acid?
(b) What should be the measured pKa of a 50:50 mixture of the two forms of
lactic acid? pKa ⫽ ⫺log Ka
(c) A solution of Form A of lactic acid is prepared. Use HL as a general formula for lactic
acid and write the chemical equation for the ionization of lactic acid in water.
(d) If separate 0.10 M solutions of each form (A and B) of lactic acid were prepared,
what would be the pH of each solution?
(e) Before any lactic acid dissolves in water, what reaction determines the pH?
(f) Calculate the pH of a solution made by dissolving 4.46 g of the B form of lactic acid
in 500. mL water.
(g) How many milliliters of a 1.15 M NaOH solution would be required to completely
neutralize 4.46 g of pure lactic acid?
(h) What would be the pH of the solution made by the neutralization if lactic acid were
in the A form? The B form? A 50:50 mixture of the two forms?
© 2010 Brooks/Cole, Cengage Learning
16
CHAPTER 16 SUMMARY PROBLEMS
1. (a) Describe how to prepare a pH 3.70 buffer using formic acid (HCOOH) and sodium
formate, NaHCOO.
(b) Calculate the pH of this buffer after the addition of 0.0050 mol HCl.
(c) How many grams of NaOH could be added to the buffer before its buffer capacity is
just exceeded?
2. The Ka of nitrous acid, HNO2, is 4.5 ⫻ 10⫺4. In a titration, 50.0 mL of 1.00 M HNO2 is
titrated with 0.750 M NaOH.
(a) Calculate the pH of the solution:
(i) Before the titration begins
(ii) When sufficient NaOH has been added to neutralize half of the nitrous acid
originally present
(iii) At the equivalence point
(iv) When 0.05 mL NaOH less than that required to reach the equivalence point has
been added
(v) When 0.05 mL NaOH more than that required to reach the equivalence point
has been added
(b) Can bromthymol blue be used as the indicator for this titration?
(c) Will methyl red be a satisfactory indicator here?
(d) Use data from part (a) to plot a graph of pH ( y-axis) versus volume of titrant.
3. A 0.500-L solution contains 0.025 mol Ag⫹.
(a) Calculate the minimum mass of NaCl that must be added to precipitate AgCl from
the solution.
(b) If excess Cl⫺ is added to the solution, the AgCl precipitate dissolves due to the formation of [Ag(Cl)2]⫺; Kf of [Ag(Cl)2]⫺ ⫽ 2.5 ⫻ 105. Calculate the minimum amount
of Cl⫺ that must be added to dissolve the precipitate.
© 2010 Brooks/Cole, Cengage Learning
In NaHCOO, HCOO⫺ is the formate
ion.
O
H9C
O⫺
17
CHAPTER 17 SUMMARY PROBLEM
In a blast furnace for making iron from iron ore, large quantities of coke (which is mainly
carbon) are dumped into the top of the furnace along with iron ore (which can be assumed to
be Fe2O3 ) and limestone (which is used to help remove impurities from the iron). The overall
process is
2 Fe2O3 (s) ⫹ 3 C(s) 9: 4 Fe(s) ⫹ 3 CO2 (g)
This reaction can be thought of as a combination of several individual steps.
2 Fe2O3 (s) 9: 4 FeO(s) ⫹ O2 (g)
2 FeO(s) 9: 2 Fe(s) ⫹ O2 (g)
2 C(s) ⫹ O2 ( g ) 9: 2 CO(g)
2 CO(g) ⫹ O2 ( g ) 9: 2 CO2 (g)
(a) Calculate the enthalpy change for each step, assuming a temperature of 25 °C.
Which steps are exothermic and which are endothermic?
(b) Based on the equations, predict which of the individual steps would involve an
increase and which a decrease in the entropy of the system.
(c) Based on your results in parts (a) and (b), what can you say about whether each step
is reactant-favored or product-favored at room temperature? At a much higher temperature (⬎1000 K)?
(d) Calculate the entropy change and the Gibbs free energy change for each reaction
step, assuming a temperature of 25 °C. (Obtain data from Appendix J.)
(e) Keeping in mind the equation ⌬G° ⫽ ⌬H° ⫺ T⌬S° and the fact that the enthalpy
change and entropy change for a reaction do not vary much with temperature, what
would be the slope of a graph of ⌬G° versus T for each of the reactions? For which
of the reactions does ⌬G° become more negative as the temperature increases? For
which does it become more positive? Does this agree with what you predicted in
part (c)?
(f) For which of these reactions might the assumption of nearly constant ⌬H° and ⌬S°
not be valid as the temperature increases from 25 °C? For each reaction you choose,
explain why the assumption might not be correct.
(g) Use your results from previous parts of this problem to estimate the Gibbs free
energy change for each of these reactions at a temperature of 1500 K.
(h) Which of the two iron oxides is more easily reduced at 1500 K? Which of the reactions involving carbon compounds is more product-favored at 1500 K? What chemical reactions do you think are taking place in the hottest part of the blast furnace?
(i) In portions of the furnace where the temperature is about 800 K, would
you predict that the same reactions would be occurring as in the highertemperature part of the furnace? Why or why not?
(j) Show that the individual steps can be combined to give the overall reaction. From
the enthalpy, entropy, and Gibbs free energy changes already calculated, calculate
these changes for the overall reaction.
(k) In a typical blast furnace every kilogram of iron produced requires 2.5 kg iron ore,
1 kg coke, and nearly 6 kg air (to provide oxygen for oxidation of the coke to heat
the furnace and to combine with carbon in the coke, forming CO(g)). How much
Gibbs free energy would be destroyed if the coke were simply burned to form carbon dioxide? Given the quantity of iron produced in a typical furnace, how much
Gibbs free energy is stored by coupling the oxidation of coke to the reduction of
iron oxides? What percentage of the Gibbs free energy available from combustion
of coke is wasted per kilogram of iron produced?
© 2010 Brooks/Cole, Cengage Learning
18
CHAPTER 18 SUMMARY PROBLEM
Many kinds of secondary batteries are known. If it were not for the density of lead, the leadacid storage battery would find far greater application. Lead-acid storage batteries continue to
be used widely as a source of electric power, but automotive engineers and others are seeking
other batteries with more desirable energy-to-mass ratios. When you review Table 18.1 and
consider the chemical properties of all of the other oxidizing and reducing agents shown
there, you might be tempted to create a hybrid battery that would combine some of the desirable features of, say, a PbO2 cathode and some other kind of anode rather than the lead anode
found in the lead-acid storage battery. In that way, at least the high reduction potential of the
half-reaction involving PbO2 might still be used.
(a) What would be the E° value of a cell made using the PbO2 reduction reaction and
magnesium metal as the reducing agent? Write the two half-reactions and the net
cell reaction. Would this cell be the basis for a secondary battery? Explain your
answer.
(b) What would be the E° value of a cell made using the PbO2 reduction reaction and
nickel metal as the reducing agent? Write the two half-reactions and the net cell
reaction. Would this cell have a voltage greater than or less than that of a single cell
of a lead-acid storage battery? Would this cell be the basis for a secondary battery?
Explain. What could you do to the chemistry in the anode compartment to make it a
secondary battery?
(c) If your Ni/PbO2 hybrid battery were a success and it was manufactured for use in
electric automobiles, how many amperes could it produce, assuming that 500.0 g Ni
reacted in exactly 30 min? How much PbO2 would be reduced during this same
period of time?
(d) Of course, batteries must be recharged. How much time would be required to
recharge your Ni/PbO2 battery to its original state (the 500.0 g Ni being converted
back to its original form) if a current of 25.5 A is passed through the battery?
(e) Just as you are getting ready to cash in on the success of your new battery, someone
announces that it has some serious environmental problems. What could these be?
Explain.
© 2010 Brooks/Cole, Cengage Learning
19
CHAPTER 19 SUMMARY PROBLEM
(a) One of the species in the uranium-238 decay series (Figure 19.1) is radon-222, an
alpha emitter. Write the nuclear equation for alpha emission by 222Rn.
(b) Uranium-238 can be converted to plutonium-239 through a series of nuclear reactions involving absorption of a neutron followed by two beta emissions. Write these
three nuclear reactions connecting 238U and 239Pu.
(c) Uranium-235 is the main fissionable nucleus used in nuclear reactors. When it fissions, it can produce 132Sb and 101Nb as products. Write the nuclear reaction for this
fission reaction.
(d) Hydrogen bombs that use fusion reactions were developed following World War II.
One reaction used in a hydrogen bomb was
2
1H
⫹ 31H 9: 42He ⫹ 10n
Calculate the energy released, in kilojoules per gram of reactants, for this
fusion reaction. The necessary nuclear masses are 21H ⫽ 2.01355 g/mol;
3
4
1
1H ⫽ 3.01550 g/mol; 2He ⫽ 4.00150 g/mol; 0n ⫽ 1.00867 g/mol.
(e) The goal of recent nuclear arms treaties has been to dismantle the stockpiles of
nuclear weapons built up by the United States and the former Soviet Union since
World War II, including those containing plutonium-239 (t1/2 ⫽ 2.44 ⫻ 104 years).
How long will it take for the activity of plutonium-239 in a nuclear warhead to
decrease (i) to 75% of its initial activity? (ii) to 10% of its initial activity?
(f ) Deep underground burial has been proposed for long-term storage of the 239Pu
waste removed from nuclear weapons. Based on the answers to part (e), comment
on factors that need to be considered for the storage and burial of such nuclear
waste.
(g) Iodine-131 emits both beta and gamma rays and is used in treating thyroid cancer.
Its half-life is 13.2 h. (i) Write the nuclear reaction for beta decay of 131I. (ii) If a 131I
sample had an activity of 5.0 ⫻ 1011 Bq (5.0 ⫻ 1011 s⫺1), what would its activity be
after 48 h?
© 2010 Brooks/Cole, Cengage Learning