ChemistryAtomsFirst2e-WEB-162

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Key Equations
ΔG = −nFEcell
Q = I t = n F
Summary
16.1 Review of Redox Chemistry
Redox reactions are defined by changes in reactant
oxidation numbers, and those most relevant to
electrochemistry involve actual transfer of electrons.
Aqueous phase redox processes often involve water
or its characteristic ions, H+ and OH−, as reactants in
addition to the oxidant and reductant, and equations
representing these reactions can be challenging to
balance. The half-reaction method is a systematic
approach to balancing such equations that involves
separate treatment of the oxidation and reduction
half-reactions.
16.2 Galvanic Cells
Galvanic cells are devices in which a spontaneous
redox reaction occurs indirectly, with the oxidant
and reductant redox couples contained in separate
half-cells. Electrons are transferred from the
reductant (in the anode half-cell) to the oxidant (in
the cathode half-cell) through an external circuit,
and inert solution phase ions are transferred
between half-cells, through a salt bridge, to maintain
charge neutrality. The construction and composition
of a galvanic cell may be succinctly represented
using chemical formulas and others symbols in the
form of a cell schematic (cell notation).
16.3 Electrode and Cell Potentials
The property of potential, E, is the energy associated
with the separation/transfer of charge. In
electrochemistry, the potentials of cells and half-
cells are thermodynamic quantities that reflect the
driving force or the spontaneity of their redox
processes. The cell potential of an electrochemical
cell is the difference in between its cathode and
anode. To permit easy sharing of half-cell potential
data, the standard hydrogen electrode (SHE) is
assigned a potential of exactly 0 V and used to define
a single electrode potential for any given half-cell.
The electrode potential of a half-cell, EX, is the cell
potential of said half-cell acting as a cathode when
connected to a SHE acting as an anode. When the
half-cell is operating under standard state
conditions, its potential is the standard electrode
potential, E°X. Standard electrode potentials reflect
the relative oxidizing strength of the half-reaction’s
reactant, with stronger oxidants exhibiting larger
(more positive) E°X values. Tabulations of standard
electrode potentials may be used to compute
standard cell potentials, E°cell, for many redox
reactions. The arithmetic sign of a cell potential
indicates the spontaneity of the cell reaction, with
positive values for spontaneous reactions and
negative values for nonspontaneous reactions
(spontaneous in the reverse direction).
16.4 Potential, Free Energy, and Equilibrium
Potential is a thermodynamic quantity reflecting the
intrinsic driving force of a redox process, and it is
directly related to the free energy change and
equilibrium constant for the process. For redox
processes taking place in electrochemical cells, the
maximum (electrical) work done by the system is
easily computed from the cell potential and the
reaction stoichiometry and is equal to the free
energy change for the process. The equilibrium
constant for a redox reaction is logarithmically
related to the reaction’s cell potential, with larger
(more positive) potentials indicating reactions with
greater driving force that equilibrate when the
reaction has proceeded far towards completion
(large value of K). Finally, the potential of a redox
process varies with the composition of the reaction
mixture, being related to the reactions standard
potential and the value of its reaction quotient, Q, as
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described by the Nernst equation.
16.5 Batteries and Fuel Cells
Galvanic cells designed specifically to function as
electrical power supplies are called batteries. A
variety of both single-use batteries (primary cells)
and rechargeable batteries (secondary cells) are
commercially available to serve a variety of
applications, with important specifications
including voltage, size, and lifetime. Fuel cells,
sometimes called flow batteries, are devices that
harness the energy of spontaneous redox reactions
normally associated with combustion processes.
Like batteries, fuel cells enable the reaction’s
electron transfer via an external circuit, but they
require continuous input of the redox reactants (fuel
and oxidant) from an external reservoir. Fuel cells
are typically much more efficient in converting the
energy released by the reaction to useful work in
comparison to internal combustion engines.
16.6 Corrosion
Spontaneous oxidation of metals by natural
electrochemical processes is called corrosion,
familiar examples including the rusting of iron and
the tarnishing of silver. Corrosion process involve
the creation of a galvanic cell in which different sites
on the metal object function as anode and cathode,
with the corrosion taking place at the anodic site.
Approaches to preventing corrosion of metals
include use of a protective coating of zinc
(galvanization) and the use of sacrificial anodes
connected to the metal object (cathodic protection).
16.7 Electrolysis
Nonspontaneous redox processes may be forced to
occur in electrochemical cells by the application of
an appropriate potential using an external power
source—a process known as electrolysis. Electrolysis
is the basis for certain ore refining processes, the
industrial production of many chemical
commodities, and the electroplating of metal
coatings on various products. Measurement of the
current flow during electrolysis permits
stoichiometric calculations.
Exercises
16.1 Review of Redox Chemistry
1. Identify each half-reaction below as either oxidation or reduction.
(a)
(b)
(c)
(d)
2. Identify each half-reaction below as either oxidation or reduction.
(a)
(b)
(c)
(d)
3. Assuming each pair of half-reactions below takes place in an acidic solution, write a balanced equation for
the overall reaction.
(a)
(b)
(c)
(d)
4. Balance the equations below assuming they occur in an acidic solution.
(a)
(b)
(c)
5. Identify the oxidant and reductant of each reaction of the previous exercise.
16 • Exercises 793
6. Balance the equations below assuming they occur in a basic solution.
(a)
(b)
(c)
(d)
7. Identify the oxidant and reductant of each reaction of the previous exercise.
8. Why don’t hydroxide ions appear in equations for half-reactions occurring in acidic solution?
9. Why don’t hydrogen ions appear in equations for half-reactions occurring in basic solution?
10. Why must the charge balance in oxidation-reduction reactions?
16.2 Galvanic Cells
11. Write cell schematics for the following cell reactions, using platinum as an inert electrode as needed.
(a)
(b)
(c)
(d)
12. Assuming the schematics below represent galvanic cells as written, identify the half-cell reactions
occurring in each.
(a)
(b)
13. Write a balanced equation for the cell reaction of each cell in the previous exercise.
14. Balance each reaction below, and write a cell schematic representing the reaction as it would occur in a
galvanic cell.
(a)
(b)
(c)
(d)
15. Identify the oxidant and reductant in each reaction of the previous exercise.
16. From the information provided, use cell notation to describe the following systems:
(a) In one half-cell, a solution of Pt(NO3)2 forms Pt metal, while in the other half-cell, Cu metal goes into a
Cu(NO3)2 solution with all solute concentrations 1 M.
(b) The cathode consists of a gold electrode in a 0.55 M Au(NO3)3 solution and the anode is a magnesium
electrode in 0.75 M Mg(NO3)2 solution.
(c) One half-cell consists of a silver electrode in a 1 M AgNO3 solution, and in the other half-cell, a copper
electrode in 1 M Cu(NO3)2 is oxidized.
17. Why is a salt bridge necessary in galvanic cells like the one in Figure 16.3?
18. An active (metal) electrode was found to gain mass as the oxidation-reduction reaction was allowed to
proceed. Was the electrode an anode or a cathode? Explain.
19. An active (metal) electrode was found to lose mass as the oxidation-reductionreaction was allowed to
proceed. Was the electrode an anode or a cathode? Explain.
20. The masses of three electrodes (A, B, and C), each from three different galvanic cells, were measured
before and after the cells were allowed to pass current for a while. The mass of electrode A increased, that
of electrode B was unchanged, and that of electrode C decreased. Identify each electrode as active or inert,
and note (if possible) whether it functioned as anode or cathode.
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16.3 Electrode and Cell Potentials
21. Calculate the standard cell potential for each reaction below, and note whether the reaction is
spontaneous under standard state conditions.
(a)
(b)
(c)
(d)
22. Calculate the standard cell potential for each reaction below, and note whether the reaction is
spontaneous under standard state conditions.
(a)
(b)
(c)
(d)
23. Write the balanced cell reaction for the cell schematic below, calculate the standard cell potential, and
note whether the reaction is spontaneous under standard state conditions.
24. Determine the cell reaction and standard cell potential at 25 °C for a cell made from a cathode half-cell
consisting of a silver electrode in 1 M silver nitrate solution and an anode half-cell consisting of a zinc
electrode in 1 M zinc nitrate. Is the reaction spontaneous at standard conditions?
25. Determine the cell reaction and standard cell potential at 25 °C for a cell made from an anode half-cell
containing a cadmium electrode in 1 M cadmium nitrate and a cathode half-cell consisting of an
aluminum electrode in 1 M aluminum nitrate solution. Is the reaction spontaneous at standard
conditions?
26. Write the balanced cell reaction for the cell schematic below, calculate the standard cell potential, and
note whether the reaction is spontaneous under standard state conditions.
16.4 Potential, Free Energy, and Equilibrium
27. For each pair of standard cell potential and electron stoichiometry values below, calculate a
corresponding standard free energy change (kJ).
(a) 0.000 V, n = 2
(b) +0.434 V, n = 2
(c) −2.439 V, n = 1
28. For each pair of standard free energy change and electron stoichiometry values below, calculate a
corresponding standard cell potential.
(a) 12 kJ/mol, n = 3
(b) −45 kJ/mol, n = 1
29. Determine the standard cell potential and the cell potential under the stated conditions for the
electrochemical reactions described here. State whether each is spontaneous or nonspontaneous under
each set of conditions at 298.15 K.
(a)
(b) The cell made from an anode half-cell consisting of an aluminum electrode in 0.015 M aluminum
nitrate solution and a cathode half-cell consisting of a nickel electrode in 0.25 M nickel(II) nitrate solution.
(c) The cell comprised of a half-cell in which aqueous bromine (1.0 M) is being oxidized to bromide ion
(0.11 M) and a half-cell in which Al3+ (0.023 M) is being reduced to aluminum metal.
30. Determine ΔG and ΔG° for each of the reactions in the previous problem.
31. Use the data in Appendix L to calculate equilibrium constants for the following reactions. Assume 298.15
K if no temperature is given.
(a)
(b)
(c)
(d)
16 • Exercises 795
16.5 Batteries and Fuel Cells
32. Consider a battery made from one half-cell that consists of a copper electrode in 1 M CuSO4 solution and
another half-cell that consists of a lead electrode in 1 M Pb(NO3)2 solution.
(a) What is the standard cell potential for the battery?
(b) What are the reactions at the anode, cathode, and the overall reaction?
(c) Most devices designed to use dry-cell batteries can operate between 1.0 and 1.5 V. Could this cell be
used to make a battery that could replace a dry-cell battery? Why or why not.
(d) Suppose sulfuric acid is added to the half-cell with the lead electrode and some PbSO4(s) forms. Would
the cell potential increase, decrease, or remain the same?
33. Consider a battery with the overall reaction:
(a) What is the reaction at the anode and cathode?
(b) A battery is “dead” when its cell potential is zero. What is the value of Q when this battery is dead?
(c) If a particular dead battery was found to have [Cu2+] = 0.11 M, what was the concentration of silver ion?
34. Why do batteries go dead, but fuel cells do not?
35. Use the Nernst equation to explain the drop in voltage observed for some batteries as they discharge.
36. Using the information thus far in this chapter, explain why battery-powered electronics perform poorly in
low temperatures.
16.6 Corrosion
37. Which member of each pair of metals is more likely to corrode (oxidize)?
(a) Mg or Ca
(b) Au or Hg
(c) Fe or Zn
(d) Ag or Pt
38. Consider the following metals: Ag, Au, Mg, Ni, and Zn. Which of these metals could be used as a sacrificial
anode in the cathodic protection of an underground steel storage tank? Steel is an alloy composed mostly
of iron, so use −0.447 V as the standard reduction potential for steel.
39. Aluminum is more easily oxidized than iron and yet when
both are exposed to the environment, untreated aluminum has very good corrosion resistance while the
corrosion resistance of untreated iron is poor. What might explain this observation?
40. If a sample of iron and a sample of zinc come into contact, the zinc corrodes but the iron does not. If a
sample of iron comes into contact with a sample of copper, the iron corrodes but the copper does not.
Explain this phenomenon.
41. Suppose you have three different metals, A, B, and C. When metals A and B come into contact, B corrodes
and A does not corrode. When metals A and C come into contact, A corrodes and C does not corrode.
Based on this information, which metal corrodes and which metal does not corrode when B and C come
into contact?
42. Why would a sacrificial anode made of lithium metal be a bad choice
16.7 Electrolysis
43. If a 2.5 A current flows through a circuit for 35 minutes, how many coulombs of charge moved through the
circuit?
44. For the scenario in the previous question, how many electrons moved through the circuit?
45. Write the half-reactions and cell reaction occurring during electrolysis of each molten salt below.
(a) CaCl2
(b) LiH
(c) AlCl3
(d) CrBr3
46. What mass of each product is produced in each of the electrolytic cells of the previous problem if a total
charge of 3.33 105 C passes through each cell?
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	Chapter 16 Electrochemistry
	Key Equations
	Summary
	Exercises