6.22 A-quantity of 2.00 X 102 mL of O.862 M HC? is

(i,4 Describe the energy conversions that occur in the following processes: (a) You throw a softball up into the
air and catch it. (b) You switch on a flashlight. (c) You
ride the ski lift to the top of a hill and then ski down.
(d) You strike a match and let it burn down.
HBview Questions
0,5 What is heat? How does heat differ from thermal en-
ergyt Under what condition is heat transferred from
one system to another?
6.6 Explain the following terms: thermochemistry, system, surroundings, open system, closed system, iso-
lated system, exothermic process, endotherrnic
6.7 Stoichiometry is based on the law of conservation of
mass. On what law is thermochemistry based?
6.8 Describe two exotherrnic processes and two endothermic processes.
heating conditions, which metal would take longer to
reach a temperature of 21oC?
6.15 Define calorimetry and describe two cormnonly used
6.16 In a calorimetric measurement, why is it important
that we know the heat capacity of the calorimeter?
6.17 A piece of silver of mass 362 g has a heat capacity
of 85.7 J/oC. What is the specific heat of silver?
6.18 A 6.22-kg piece of copper metal is heated from
20.5oC to 324.3oC. Calculate the heat absorbed (in
kJ) by the metal.
6.19 Calculate the amount of heat liberated (in kJ) from
366 g of mercury when it cools from 77.OoC to
6.20 A sheet of gold weighing 10.0 g and at a temperature of 1 8.OoC is placed flat on a sheet of iron weighing 20.0 g and at a temperature of 55.6oC. What is
the final temperature of the combined metals?
Assume that no heat is lost to the surroundings. (Hint:
The heat gained by the gold must be equal to the heat
lost by the iron.)
Review Questions
6.9 Write an expression for the enthalpy of a reaction in
terms of the enthalpies of products and reactants.
Under what condition is the heat of a reaction equal
to the enthalpy change of the same reaction?
6.10 In writing thermochemical equations, why is it important to indicate the physical state (that is, gaseous,
liquid, solid, or aqueous) of each substance?
6.11 Explain the following thermochemical equation:
4NH3(g) + 50;i(g) ? 4NO(g) + 6H20(g)
AH = -905 J
6.21 A O.1375-g sample of solid magnesium is burned in
a constant-volume bomb calorimeter that has a heat
capacity of 1769 J/oC. The calorimeter contains exactly 300 g of water, and the temperature increases
by l.l26oC. Calculate the heat given off by the burning Mg, in kJ/g and in kJ/mol.
6.22 A-quantity of 2.00 X 102 mL of O.862 M HC? is
mixed with 2.00 X 102 mL of O.431 M Ba(OH)2 in
a constant-pressure calorimeter that has a heat capacity of 453 J/oC. The initial temperature of the HC?
and Ba(OH)2 solutions is the same at 20.48oC. For
the process
6.12 Consider the following reaction:
2CH30H(}) + 302(g) -> 4H20(!) + 2CO,(J)
,?p =-1452.8 kJ
What is the value of bH if (a) the equation is multiplied throughout by 2, (b) the direction of the reaction is reversed so that the products become the reactants and vice versa, (c) water vapor instead of
liquid water is formed as the product?
Review Questions
6.13 What is the difference between specific heat and heat
capacity? What are the units for these two quantities?
Which is the intensive property and which is the extensive property?
6.14 Consider two metals A and B, each having a mass of
100 g and an initial temperature of 20oC. The specific heat of A is larger than that of B. Under the same
H+(aq) + OH (aq) -+ H20(I)
the heat of neutralization is -56.2 kJ. What is the fi-
nal temperature of the mixed solution?
Review Queslions
6.23 What is meant by the standard-state condition?
6.24 How are the standard enthalpies of an element and of
a compound deterrnined?
6.25 What is meant by the standard enthalpy of a reaction?
6.26 Write the equation for calculating the enthalpy of a
reaction. Define all the terms.
6.27 State Hess's law. Explain, with one example, the usefulness of Hess's law in thermochemistry.
6.28 Describe how chemists use Hess's law to detern'iine
the b,H( of a compound by measuring its heat (enthalpy) of combustion.
Calculate the heats of combustion of these alcohols
6.29 Which of the following standard enthalpy of formation values is not zero at 25oC? Na(s), Ne(g), CH4(g),
Ss(s), Hg(I), H(g).
in kJ/mol.
6.40 The standard enthalpy change for the following reaction is 436.4 kJ:
6.30 The b'[-[7 values of the two allotropes of oxygen, 02
and 03, are O and 142.2 kJ/mol, respectively, at 25oC.
Which is the more stable fom'i at this temperature?
6.31 Which is the more negative quantity at 25oC: LH7 for
H20(}) or fsl[', for H20(g)?
6.32 Predict the value of LF]7 (greater than, less than, or
equal to zero) for these elements at 25oC: (a) Brz(g);
Br2('), (b) Iz(g): I2(S)-
6.33 In general, compounds with negativeveb'H'i
more stable than those with positive
bHl values.
H202(7) has a negative bH7 (see Table 6.-3). Why,
then, does H202(I) have a tendency to decompose to
H2(.!,) ? H(g) + H(g)
Calculate the standard enthalpy of formation of
atomic hydrogen (H).
6.41 From the standard enthalpies of fomiation, calculate
Ag'n for the reaction
C6H12(I) + 902(g) -+ 6CO2(g) + 6H20(})
For C6H12(I), ul, = -151.9 kJ/mol.
6.42 The first step in the industrial recovery of zinc from
the zinc sulfide ore is roasting, that is, the conversion
of ZnS to ZnO by heating:
H20(}) and 02(g)?
6.34 Suggest ways (with appropriate equations) that would
allow you to measure the bHy values of Ag20(s) and
2ZnS(s) + 302(g) -+ 2ZnO(s) + 2SO2(g)
AH',,, = -879 kJ
CaCl2(s) from their elements. No calculations are
Calculate the heat evolved (in kJ) per gram of ZnS
6.35 Calculate the heat of decomposition for this process
at constant pressure and 25oC:
CaCO3(s) -? CaO(s) + CO;i(g)
6.43 Deterrnine the amount of heat (in kJ) given off when
1.26 X 104 g of ammonia are produced according to
the equation
N2(g) + 3H2(g) -+ 2NH3(g) h"':xn = -92.6 kJ
(Look up the standard enthalpy of formation of the
reactant and products in Table 6.3.)
6.36 The standard enthalpies of formation of ions in aqueous solutions are obtained by arbitrarily assigning a
value of zero to H+ ions; that is, Aq[H+(aq)] '0.
(a) For the following reaction
Assume that the reaction takes place under standardstate conditions at 25oC.
6.44 At 850oC, CaCO3 undergoes substantial decomposi-
tion to yield CaO and CO2. Assuming that the LH7
values of the reactant and products are the same at
850oC as they are at 25oC, calculate the enthalpy
HC?(g)-+H+(aq)+Cl(aq) AHo= -74.9kJ
change (in kJ) if 66.8 g of CO2 are produced in one
calculate LH7 for the Cl ions. (b) Given that AF]7
6.45 From these data,
for OH ions is -229.6 kJ/mol, calculate the en-
thalpy of neutralization when 1 mole of a strong
monoprotic acid (such as HC?) is titrated by 1 mole
of a strong base (such as KOH) at 25oC.
6.37 Calculate the heats of combustion for the following
reactions from the standard enthalpies of formation
listed in Appendix 3:
(a) 2Hz(g) + Oz(g) --> 2HzO(l)
(b) 2C2H2(,!') + 50z(g) ? 4CO2(.!,) + 2HzO(!)
6.38 Calculate the heats of combustion for the following
reactions from the standard enthalpies of formation
listed in Appendix 3:
(a) C:iHa(g) + 302(,!') '-+ 2CO2(,!,) + 2HzO(})
(b) 2H2S(,!') + 30z(g) ? 2H:?O(I) + 2SO2(,!')
6.39 Methanol, ethanol, and n-propanol are three common
alcohols. When 1.00 g of each of these alcohols is
burned in air, heat is liberated as follows: (a) methanol
(CH30H), -22.6 kJ; (b) ethanol (C2H50H),
-29.7 kJ; (c) n-propanol (C3H70H), -33.4 kJ.
s(rhombic) + 02(g) -+ so2(g) Affixn = -296.06 kJ
S(monoclinic) + 02(g) -+ SO2(g)
h"rxn = -296.36 kJ
calculate the enthalpy change for the transformation
S(rhombic) ? S(monoclinic)
(Monoclinic and rhombic are different allotropic
forms of elemental sulfur.)
6.46 From the following data,
C(graphite) + O;i(g) ? CO2(g) Agn = -393.5 kJ
H:i(g) + 2'Oz(g)?Hz0(I) AFI',n= 285.8kJ
2C2H6(,!,) + 70z(g) ? 4CO2(,!,) + 6H20(I)
Ag, = -31l9.6kJ
calculate the enthalpy change for the reaction
9.52 Draw two resonance structures for diazomethane,
CH2N2. Show formal charges. The skeletal structure
of the molecule is
Review Questions
9.65 What is bond dissociation energy? Bond energies Oy
polyatornic molecules are average values, wherea3
9.53 Draw three reasonable resonance structures for the
OCN ion. Show formal charges.
9.54 Draw three resonance structures for the molecule
N20 in which the atoms are arranged in the order
NNO. Indicate formal charges.
Review Questions
9.55 Why does the octet rule not hold for many compounds
containing elements in the third period of the periodic table and beyond?
9.56 Give three examples of compounds that do not satisfy the octet rule. Write a Lewis structure for each.
9.57 Because fluorine has seven valence electrons
(2s"2ps), seven covalent bonds in principle could
form around the atom. Such a compound might be
FH7 or FCl7. These compounds have never been prepared. Why?
9.58 What is a coordinate covalent bond? Is it different
from a normal covalent bond?
those of diatomic molecules can be accurately deteymined. Why?
9.66 Explain why the bond energy of a molecule is us[-
ally defined in terms of a gas-phase reaction. Why
are bond-breaking processes always endotherrnic qj
bond-forming processes always exothemiic?
9.67 From the following data, calculate the average bond
energy for the N-H bond:
NH3(g) -+ NH2(,!,) + H(g) LHo = 435 kJ
NH2(g) -+ NH(g) + H(,!,) ?o = 381 kJ
NH(g) ? N(g) + H(g)
AHo = 360 kJ
9.68 For the reaction
O(g)+02(g')?03(g) fsHo= -l07.2kJ
Calculate the average bond energy in 03.
9.69 The bond energy of F:?(g) is 156.9 kJ/mol. Calculate
aHi for F(g).
9.70 For the reaction
2CzHb(g) + 70zC!') -'-> 4COz(g) + 6H20(,!,)
(a) Predict the enthalpy of reaction from the average
bond energies in Table 9.4.
Draw three resonance structures of the molecule in
which the octet rule is satisfied for both the Al and
(b) Calculate the enthalpy of reaction from the standard enthalpies of formation (see Appendix 3) of the
reactant and product molecules, and compare the re-
the I atoms. Show formal charges.
sult with your answer for part (a).
9.59 The Al?3 molecule has an incomplete octet around Al.
9.60 In the vapor phase, beryllium chloride consists of discrete BeCl2 molecules. Is the octet rule satisfied for
Be in this compound? If not, can you form an octet
around Be by drawing another resonance structure?
How plausible is this structure?
9.61 0f the noble gases, only Kr, Xe, and Rn are known
to form a few compounds with O and/or F. Write
Lewis strOctures for the following molecules:
(a) XeF2, (b) XeF4, (c) XeF6, (d) XeOF4, (e) Xe02F2.
In each case Xe is the central atom.
9.62 Write a Lewis stmcture for SbCl,.5- Does this molecule obey the octet rule'!
9.63 Write Lewis structures for SeF4 and SeF6. Is the octet
rule satisfied for Se?
9.64 Write Lewis structures for the reaction
AIC?':i + Cl --?AlCl4
What kind of bond joins Al and Cl in the product?
9.71 Classify the following substances as tonic compounds
or covalent compounds containing discrete molecules: CH4, KF, CO, SiCl4, BaCl2.
9.72 Which of the following are tonic compounds? Which
are covalent compounds? RbCl, PF5, BrF3, KO2, CI4
9.73 Match each of the following energy changes with one
of the processes given: ionization energy, electron
affinity, bond dissociation energy, and standard enthalpy of formation.
(a) F(g) + e -+F (g)
(b) Fz(g) -? 2F(g)
(c) Na(g) -+ Na'(g) + e
(d) Na(s) + ,'F2(g) ? NaF(s)
9.74 The formulas for the fluorides of the third-period elements are NaF, MgF2, AlF3, SiF4, PF5, SF6, and
ClF3. Classify these compounds as covalent or tonic.
2C(graphite) + 3H2(g) -+ C2H6(g)
0.47 From the following heats of combustion,
CHsOH(}) + 23C%(g) ? CO2(,!,) + 2H;iO(})
AH',,, = -726.4 kJ
C(graphite) + 02(g) -x CO2(g) Agn = -393.5 kJ
Hz(g) + 2' Oz(g) ? HzO(}) A?n = 285.8 kJ
calculate the enthalpy of formation of methanol
(CH30H) from its elements:
C(graphite) + 2H2(g) + 402(g) -+ CH30H(0
6.48 Calculate the standard enthalpy change for the reaction
6.56 Explain what is meant by a state function. Give two
examples of quantities that are state functions and two
that are not.
6.57 The internal energy of an ideal gas depends only on
its temperature. Do a first-law analysis of the following process. A sample of an ideal gas is allowed
to expand at constant temperature against atmospheric pressure. (a) Does the gas do work on its surroundings? (b) Is there heat exchange between the
system and the surroundings? If so, in which direc-
tion? (c) What is AE for the gas for this process?
6.58 At constant pressure, in which of the following reactions is work done by the system on the surroundings? By the surroundings on the system? In which
of them is no work done?
2Al(s) + Fe203(s) -+ 2Fe(s) + Al203(s)
given that
(a) Hg(}) -+ Hg(g)
(b) 30z(g) '? 2C%(g)
2Al(s)+3HO2(g)-+Al203(s) LHo,n= -160lkJ
(c) CuSO4 a 5H20(s) -+ CuSO=i(s) + 5H20(g)
(d) Hz(g) + Fz(g) -+ 2HF(g)
2Fe(s) + y302(g) -? Fe203(s) AP,n = -821 kJ
Review Questions
6.49 Define the following terms: enthalpy of solution, hydration, heat of hydration, lattice energy, heat of dilution.
6.50 Why is the lattice of a solid always a positive quantity? Why is the hydration of ions always a negative
6.51 Consider two tonic compounds A and B. A has a
larger lattice energy than B. Which of the two compounds is more stable?
6.52 Mg2+ is a smaller cation than Na+ and also carries
more positive charge. Which of the two species has
a larger hydration energy (in kJ/mol)? Explain.
6.53 Consider the dissolution of an tonic compound such
as potassium fluoride in water. Break the process into
the following steps: separation of the cations and anions in the vapor phase and the hydration of the ions
in the aqueous medium. Discuss the energy changes
associated with each step. How does the heat of solution of KF depend on the relative magnitudes of
Uhese two quantities? On what law is the relationship
6.54 Why is it dangerous to add water to a concentrated
acid such as sulfuric acid in a dilution process?
Review Questions
6.55 0n what law is the first law of thermodynamics
based? Explain the sign conventions in the equation
aE = q + w.
6.59 A gas expands and does P-V work on the surroundings equal to 325 J. At the same time, it absorbs 127 J
of heat from the surroundings. Calculate the change
in energy of the gas.
6.60 The work done to compress a gas is 74 J. As a result,
26 J of heat is given off to the surroundings. Calculate
the change in energy of the gas.
6.61 Calculate the work done when 50.0 g of tin are dissolved in excess acid at 1.00 atm and 25oC:
Sn(s) + 2H+(aq) -+ Sn?'+(aq) + H2(g)
Assume ideal gas behavior.
6.62 Calculate the work done in joules when 1.0 mole of
water vaporizes at 1.0 atm and lOOoC. Assume that
the volume of liquid water is negligible compared
with that of steam at lOOoC and ideal gas behavior.
Additional Pmblems
6.63 The convention of arbitrarily assigning a zero enthalpy value for the most stable form 6f each element
in the standard state of 25oC is a convenient way of
dealing with enthalpies of reactions. Explain why this
convention cannot be applied to nuclear reactions.
6.64 Consider the following two reactions:
A ? 2B
AH',, = LH,
A -+ C
hr,, = bH2
Determine the enthalpy change for the process
2B -+ C
6.65 The standard enthalpy change b.Ho for the thermal
decomposition of silver nitrate according to the fol-
8. Most biological reactions are nonspontaneous. They are driven by the hydrolysis of ATP,
for which LGo is negative.
Entropy (S), p. 727
Free energy (Q), p. 737
Gibbs free energy ((7),
Second law of
thermodynamics, p. 731
Standard entropy of reaction,
p. 737
Standard free energy of
reaction (AGo), p. 738
Standard free energy of
Third law of
thermodynamics, p. 734
formation (AG7), p. 738
Review Questions
18.8 State the third law of thermodynamics and explain its
usefulness in calculating entropy values.
18.1Explainwhatismeantbyaspontaneousprocess.Give Problems
two examples each of spontaneous and nonspontaneous processes.
18.2 Which of the following processes are spontaneous
and which are nonspontaneous? (a) dissolving table
salt (NaCl) in hot soup; (b) climbing Mt. Everest; (c)
spreading fragrance in a room by removing the cap
from a perfume bottle; (d) separating helium and neon
from a mixture of the gases
18.3 Which of the following processes are spontaneous
and which are nonspontaneous at a given tempera-
-----'---------- -- - ?'=--- ----'---
(a) NaNO?,(s) ? NaNO,,(aq) saturated soln
(b) NaNO3(s) l? NaNO3(aq) unsaturated soln
(c) NaNO:i(s) -? NaNO,,(aq) supersaturated soln
18,4 Define entropy. What are the units of entropy?
18.5 How does the entropy of a system change for each
of the following processes?
(a) A solid melts.
(b) A liquid freezes.
(c) A liquid boils.
(d) A vapor is converted to a solid.
(e) A vapor condenses to a liquid.
(f) A solid sublimes.
(g) Urea dissolves in water.
18.6 Referring to the setup in Figure 18.1(a), calculate the
probability of all the molecules ending up in the same
flask if the number is (a) 6, (b) 60, (c) 600.
Review Questions
18.7 State the second law of thermodynamics in words and
express it mathematically.
18.9 For each pair of substances listed here, choose the
one having the larger standard entropy value at 25oC.
The same molar amount is used in the comparison.
Explain the basis for your choice. (a) Li(s) or Li(7);
(b) C2H50H(}) or CH30CH3(I) (Hint: Which molecule can hydrogen-bond?); (c) Ar(g) or Xe(g);
(d) CO(g) or COz(g); (e) O:i(g) or C%(g); (f) NC%(g)
or NxC%C!')
18.10 Arrange the following substances (1 mole each) in
order of increasing entropy at 25oC: (a) Ne(g), (b)
SO2(g), (c) Na(s), (d) NaCl(s), (e) NH:i(g)- Give the
reasons for your arrangement.
18.11 Using the data in Appendix 3, calculate the standard
entropy changes for the following reactions at 25oC:
(a) S(S) + O:z(g) -'-> SOz(g)
(b) MgCO3(s) -? MgO(s) + CO2(g)
18.12 Using the data in Appendix 3, calculate the standard
entropy changes for the following reactions at 25oC:
(a) H2(g) + CuO(s) -+ Cu(s) + H20(g)
(b) 2Al(s) + 3ZnO(s) -+ AlzO:i(s) + 3Zn(s)
(C) CHa(g) + 20z(g) -'-> COz(g) + 2HzO(I)
18.13 Without consulting Appendix 3, predict whether the
entropy change is positive or negative for each of the
following reactions. Give reasons for your predictions.
(a) 2KClO,,(s) ? 2KClO,,(s) + 02(g)
(b) H20(.!,) -'-> HzO(l)
(c) 2Na(s) + 2H20(}) -+ 2NaOH(aq) + H2(g)
(d) N2(,!,) -+ 2N(g)
18.14 State whether the sign of the entropy change expected
for each of the following processes will be positive
or negative, and explain your predictions.
(a) PCl3(I) + Cl2(g) -+ PC?s(s)
(b) 2HgO(s) -+ 2Hg(I) + 02(,!,)
(C) H2(,!,) -? 2H(g)
(d) U(s) + 3F2(g) ? UF6(s)
Rev;ew Questions
18.15 Define free energy. What are its units?
18.16 Why is it more convenient to predict the direction of
a reaction in terms of AGsYS instead of 'univ? Under
what conditions can LGsys' be used to predict the
2HzO(g) :? 2H:?(.!') + Oz(g)
18.27 (a) Calculate bGo and KP for the following equilib-
rium reaction at 25oC. The LG7 values are O for
Cl2(g), -286 kJ/mol for PC?:i(g), and -325 kJ/mol
for PCl5(g).
PClb(g) ? PC?:t(!') + Clz(g)
spontaneity of a reaction?
(b) Calculate AG for the reaction if the partial pres-
18.17 Calculate bGo for the following reactions at 25oC:
(a) N2(.!,) + Oz(g) -+ 2NO(g)
(b) HzO(!) -'-> H20(.!,)
(C) 2CzHz(g) + 502(,!,) -'-> 4CO2(,!,) + 2H;iO(I)
(Hint: Look up the standard free energies of formation of the reactants and products in Appendix 3.)
i8.l8 Calculate LGo for the following reactions at 25oC:
(a) 2Mg(s) + 02(g) -+ 2MgO(s)
(b) 2SOz(g) + Oz(g) -+ 2SO:iC!')
(c) 2CzHb(g) + 702(g) -€ 4CO2(,!,) + 6HzO(I)
sures of the initial mixture are Pp,?l5 = 0.0029 atm,
Pp(.13 = 0.27 atm, and P(212 = 0.40 aim.
18.28 The -equilibrium constant ;Kp) for the reaction
Hz(g) + CO2(,!') ':? HzO(g) + CO(,!')
is 4.40 at 2000 K. (a) Calculate bGo for the reaction.
(b) Calculate AG for the reaction when the partial
pressures are PH2 '- 0.25 atm, P([email protected] = 0.78 atm,
PH20 = 0.66 atm, ffid P(20 = 1.20 atm.
18.29 Co;sider the decomposition of calcium carbonate:
CaCO3(s) '? CaO(s) + CO2(g)
See Appendix 3 for thermodynarnic data.
18.19 From the values of AH and LS, predict which of the
following reactions would be spontaneous at 25oC:
Reaction A: AH = 10.5 kJ, AS = 30 J/K; reaction B:
LH = 1.8 kJ, AS = -113 J/K. If either of the reactions is nonspontaneous at 25oC, at what temperature
might it become spontaneous?
!8.20 Find the temperatures at which reactions with the following LH and AS values would become spontaneous: (a) LH = -126 kJ, AS = 84 J/K; (b) LH =
-11.7 kJ, AS = -105 J/K.
Calculate the pressure in atm of CO2 in an equilibrium process (a) at 25oC and (b) at 800oC. Assume
that ?o = ?77.8 kJ and LSo = 160.5 J/K for the
temperature range.
18.30 The equilibrium constant Kp for the reaction
CO(g) + Clz(g) '? COC?;i(g)
is 5.62 X 1035 at 25oC. Calculate AG7 for COCl2 at
18.31 At 25oC, biGo for the process
HzO(I) '=' H20(.!,)
Review Questions
18.21 Explain the difference between AG and LGo.
18.22 Explain why Equation (18.10) is of great importance
is 8.6 kJ. Calculate the "equilibrium constant" for the
18.32 Calculate LGo for the process
C(diamond) -+ C(graphite)
in chemistry.
18.23 Calculate Kp for the following reaction at 25oC:
H2(g) + I2(g) ? 2HI(,!,) fsGo = 2.60 kJ
N8.24 For the autoionization of water at 25oC,
H20(I) ? H"(aq) + OH (aq:)
KW is 1.0 X 10 14. What is LGo for the process?
Is the reaction spontaneous at 25oC? If so, why is it
that diamonds do not become graphite on standing?
Review Questions
18.33 What is a coupled reaction? What is its importance
in biological reactions?
18.34 What is the role of ATP in biological reactions?
18.25 Consider the following reaction at 25oC:
Fe(OH);i(s) ? Fe2'(aq) + 20H (aq)
b9o for the reaction. Ksp for Fe(OH)2 is
1.6 X 10 "'.
N8.26 Calculate LGo and Kp for the following equilibrium
reaction at 25oC.
18.35 Referring to the metabolic process involving glucose
on p. 747, calculate the maximum number of moles
of ATP that can be synthesized from ADP from the
breakdown of one mole of glucose.
18.36 In the metabolism of glucose, the first step is the con-
19.17 Which species in each pair is a better oxidizing agent
under standard-state conditions ? (a) Br2 or Au3+, (b)
Hz or Ag+, (c) Cd2+ or Cr3+, (d) 02 in acidic media or 02 in basic media
19,10 Which species in each pair is a better reducing agent
under standard-state conditions ? (a) Na or Li, (b) Hz
or I2, (c) Fe2+ or Ag, (d) Br or Co2+
fleview Questions
19.19 Write the equations relating AGo and K to the standard emf of a ceM. Define all the terms.
19.20 Compare the ease of measuring the equilibrium constant electrochemically with that by chemical means
[see Equation (18.10)].
19.21 What is the equilibrium constant for the following
reaction at 25oC?
Review Questions
19.27 Write the Nernst equation and explain all the terms.
19.28 Write the Nernst equation for the following
processes at some temperature 7:a
(a) Mg(s) + Sn"+(a4) ? Mg2+(aq) + Sn(s)
(b) 2Cr(s) + 3Pb"+(aq) ? 2Cr3+(aq) + 3Pb(s)
19.29 What is the potential of a cell made
le up of 2Zn/Zn"+
and Cu/Cu2' half-cells at 25oC if [Zn2+]
(zn2+),= 0.25 M
and [Cu2'] = 0.15 A/?
19,30 Calculate Eo, E, and AG for the following cell reactions.
(a) Mg(s) + Sn"+(aq) -+ Mg"+(aq) + Sn(s)
[Mg2"] = 0.045 M, [Sn2+] = 0.035 M
(b)-3Zn(s) + 2Cr3+(aq) ? 3Zn"+(aq) + 2Cr(s)
(cr")4 o.oio rvr, (zn2j = 0.0085 w
19.31 Calculate the standard potential of the cell consist-
ing of the Zn/Zn2+ halfHcell and the SHE. What will
Mg(s) + Zn2+(aq) ? Mg2+(aq) + Zn(s)
19,22 The equilibrium constant for the reaction
Sr(s) + Mg2+(aq) ? Sr"+(aq) + Mg(s)
is 2.69 X [email protected] at 25oC. Calculate Eo for a cell
made up of Sr/Sr2+ and Mg/Mg2+ half-cells.
19.23 Use the standard reduction potentials to find the
equilibrium constant for each of the following re-
th; emf of the cell be if tzn2+) = 0.45 M, PH2 =
2.0 atm, and [H+] = 1.8 M?
19,32 What is the ernf of a cell consisting of a Pb/Pb2+
half-cell and a Pt/H2/H+ half-cell if [Pb2+] =
0.10 M, [H+] = 0.050 M, and PH2 = 1.0 atm?
19.33 Referring to the arrangement in Figure 19.1, calcu-
late the [Cu"+]/[Zn"+] ratio at which the following
reaction is spontaneous at 25oC:
Cu(s) + Znz+(aq) ? Cu2+(aq) + Zn(s)
actions at 25oC:
(a) Br2(}) + 2I (aq) -? 2Br (aq) + I2(s)
(b) 2Ce"+(aq) + 2Cl (aq) ?
- Cl2(!) + 2Ce3+(aq)
19,34 Calculate the ernf of the following concentration
Mg(s)l Mg?'+(0.24 M) 11 Mg2+(0.53 M)i Mg(s)
(c) 5Fe2+(aq) + Mn04 (aq) + 8H"(aq) ?
Mn2+(aq) + 4H20 + 5Fe3+(aq)
19,24 Calculate LGo and Kc for the following reactions at
(a) Mg(s) + Pb"+(aq) ? Mg2+(aq) + Pb(S)
(b) Br2(I) + 2I (aq) ? 2Br (aq) + I2(s)
(c) 0:?(g) + 4H+(aq) + 4Fe"+(aq) ?
(d) 2Al(s) + 3I2(s) ? 2Al3+(aq) + bI (aq)
19.25 Under standard-state conditions, what spontaneous
reaction will occur in aqueous solution among the
ions Ce4+, Ce3+, Fe3+, and Fe"+? Calculate AGo
Review Questions
19.35 Explain the differences between a primary electrochemical cell-one that is not rechargeable-and a
storage cell (for example, the lead storage battery),
which is rechargeable.
19.36 Discuss the advantages and disadvantages of fuel
cells over conventional power plants in producing
and Kc for the reaction.
19,26 Given that Eo = 0.52 V for the reduction Cu+(aq)
+ e -? Cu(s), calculate Eo, AGo, and K for the
following reaction at 25oC:
2Cu+(aq) ? Cu"+(aq) + Cu(s).
19.37 The hydrogen-oxygen fuel cell is described in
Section 19.6. (a) What volume of H2(g), stored at
25oC at a pressure of 155 atm, would be needed to