1. Chapter 13: Electrochemistry
University Chemistry
Chapter 13: Electrochemistry
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2. Electrochemistry: The branch of chemistry that studies the interconversion of electric energy and chemical energy.
Electrochemical processes are oxidation-reduction reactions in which:
the energy released by a spontaneous reaction is converted to electricity or
electrical energy is used to cause a nonspontaneous reaction to occur
An oxidation-reduction can be considered with
two separate half-reactions
reducing agent
oxidizing agent 2+ 2-
2Mg (s) + O2 (g) MgO (s)
oxidized
reduced
Half-reactions:
2Mg Mg2+ + 4e-
Oxidation half-reaction (lose e-)
O20 + 4e O2-
Reduction half-reaction (gain e-)

3. Metal Displacement Reactions
Zinc strip is placed in a blue CuSO4 solution.

4. Cu wire is placed in a colorless AgNO3 solution.

5. Oxidation Number (Oxidation State)
H2(g) + Cl2(g) HCl(g)
Electrons are not completely transferred from H to Cl. However, if electrons were assumed to be completely transferred, the oxidation number is the charge that the atom would have in a molecule (or an ionic compound. It is a means to keep track of electrons in redox reactions.
Rule:
Free elements (uncombined state) have an oxidation number of zero.
Na, Be, K, Pb, H2, O2, P4 = 0
In monatomic ions, the oxidation number is equal to the charge on the ion.
Li+, Li = +1; Fe3+, Fe = +3; O2-, O = -2
The oxidation number of oxygen is usually –2. In H2O2 and O22- it is –1.

6. HCO3- O = -2 H = +1 3x(-2) + 1 + ? = -1 C = +4
The oxidation number of hydrogen is +1 except when it is bonded to metals in binary compounds. In these cases, its oxidation number is –1.
The oxidation number of the halogens is –1 except in interhalogen compounds and when combined with oxygen.
6. The sum of the oxidation numbers of all the atoms in a molecule or ion is equal to the charge on the molecule or ion.
HCO3-
What are the oxidation numbers of all the atoms in HCO3- ?
O = -2
H = +1
3x(-2) ? = -1
C = +4

8. More common oxidation numbers are in red.

9. Based on the periodic table:
Metallic elements have only positive oxidation numbers, whereas nonmetallic elements may have either positive or negative oxidation numbers.
The highest oxidation number an element in groups 1A - 7A can have is its group number. For example, the halogens are in group 7A, so their highest possible oxidation number is +7.
The transition metals (groups 1B and 3B - 8B) usually have several possible oxidation numbers.

10. Balancing Redox Equations
Balance the equation for the oxidation of Fe2+ to Fe3+ by Cr2O72- in acid solution.
Write the unbalanced equation for the reaction ion ionic form.
Fe2+ + Cr2O Fe3+ + Cr3+
Separate the equation into two half-reactions.
Fe Fe3+ +2 +3
Oxidation:
Cr2O Cr3+ +6 +3
Reduction:
Balance the atoms other than O and H in each half-reaction.
Cr2O Cr3+

11. Add the two half-reactions together and balance the final equation by inspection. The number of electrons on both sides must cancel.
Oxidation:
6Fe Fe3+ + 6e-
Reduction:
6e- + 14H+ + Cr2O Cr3+ + 7H2O
14H+ + Cr2O Fe Fe3+ + 2Cr3+ + 7H2O
Verify that the number of atoms and the charges are balanced.
14x1 – 2 + 6x2 = 24 = 6x3 + 2x3
For reactions in basic solutions, add OH- to both sides of the equation for every H+ that appears in the final equation.

15. Galvanic (Volatic) Cells
spontaneous
redox reaction

16. Zn (s) + Cu2+ (aq) Cu (s) + Zn2+ (aq)
salt bridge.
cell voltage (cell potential): The difference in electrical potential (in Volts, V) between the anode and cathode is called:
Electromotive force or emf ( ); the limit when the cell is operated reversibly
Zn (s) + Cu2+ (aq) Cu (s) + Zn2+ (aq)
[Cu2+] = 1 M and [Zn2+] = 1 M
phase boundary
salt bridge
Cell Diagram
Zn (s) | Zn2+ (1 M) || Cu2+ (1 M) | Cu (s)
anode
cathode

17. Standard Emf of Electrochemical Cells
standard reduction potential ( ), or the electromotive force, is associated with a reduction reaction at an electrode when reactant and product species are in their standard states.
Standard Hydrogen Electrode (SHE)
Reduction Reaction

18. Standard cell emf ( )
Anode (oxidation):
Cathode (reduction):

19. Meaurement of Standard cell emf ( )

20. Meaurement of Standard cell emf ( )

21. is for the reaction as written
The more positive the greater the tendency for the substance to be reduced
The half-cell reactions are reversible
The sign of changes when the reaction is reversed
Changing the stoichiometric coefficients of a half-cell reaction does not change the value of .

22. Like DH, DG, and DS, the sign of
Under standard-state conditions, any species on the left of a given half-cell reaction
will react spontaneously with a species that appears on the right of any half-cell reaction located below it in Table 13.1.
diagonal rule.
Like DH, DG, and DS, the sign of
changes, but its magnitude remains the same when we reverse a reaction.

25. Dental Filling Discomfort
Dental amalgam actually consists of three solid phases having stoichiometries approximately corresponding to Ag2Hg3, Ag3Sn, and Sn8Hg.
Hg2 /Ag2Hg V 2+
Sn /Ag3Sn V 2+
Sn /Sn8Hg V 2+
Discomfort occurs when an electrochemical cell has been created in the mouth,
Corrosion of a dental filling brought about by contact with a gold inlay.

26. Emf and Gibbs Free Energy
nFEcell = -Grxn ?
Ecell (T, conc.) ?
E0cell = ?
Ecell vs Q ?; E0cell vs K ?
pH measurement from Ecell?
Batteries? Fuel cells? Electrolytic cells?

27. Emf and Gibbs Free Energy
In a galvanic cell, chemical energy is converted to electric energy.

28. Relationship of to DG, DGo, and K

29. Spontaneity of Redox Reactions

32. Temperature Dependence of Emf
Assume that both DHo and DSo are independent of temperature.

35. The Nernst Equation and Concentration Dependence of the Emf
At 25ºC ( K),
At equilibrium, no net transfer of electrons, so = 0, Q = K.

40. Measurement of pH: the glass electrode
The emf of the cell made up of the glass electrode and the reference electrode is measured with a voltmeter that is calibrated in pH units.
Very thin glass membrane that is permeable to H+ ions.
A potential difference develops between the two sides of the membrane.

41. Concentration Cells
A galvanic cell from two half-cells composed of the same material but differing in ion concentrations is called a concentration cell.

42. Determining Activities from Emf Measurements-
mean activity coefficient

43. Batteries
A battery is a galvanic cell, or a series of combined galvanic cells, that can be used as a source of direct electric current at a constant voltage.
use electrochemical reactions to produce a ready
supply of electric current
similar in principle to that of the galvanic cells
has the advantage of being completely self-contained
requires no auxiliary components such as salt bridges.

44. The Dry Cell Leclanché cell
Anode:
Zn (s) Zn2+ (aq) + 2e-
Cathode:
2NH4 (aq) + 2MnO2 (s) + 2e Mn2O3 (s) + 2NH3 (aq) + H2O (l) +
Zn (s) + 2NH4 (aq) + 2MnO2 (s) Zn2+ (aq) + 2NH3 (aq) + H2O (l) + Mn2O3 (s)
Output is 1.5 V, but voltage will decrease somewhat over time.

45. The Mercury Battery
Anode:
Zn(Hg) + 2OH- (aq) ZnO (s) + H2O (l) + 2e-
Cathode:
HgO (s) + H2O (l) + 2e Hg (l) + 2OH- (aq)
Zn(Hg) + HgO (s) ZnO (s) + Hg (l)
More constant voltage (1.35 V) than the Leclanché cell.

46. Lead Storage Battery Anode: Pb (s) + SO2- (aq) PbSO4 (s) + 2e-
Discharge can be checked by measuring the density of the electrolyte
with a hydrometer.
An increase in the viscosity of the electrolyte as the temperature
decreases causes cold weather battery problems.
Anode:
Pb (s) + SO2- (aq) PbSO4 (s) + 2e- 4
Cathode:
PbO2 (s) + 4H+ (aq) + SO2- (aq) + 2e PbSO4 (s) + 2H2O (l) 4
Pb (s) + PbO2 (s) + 4H+ (aq) + 2SO2- (aq) PbSO4 (s) + 2H2O (l) 4
Each cell produces 2 V; a total of 12 V from the six cells
The lead storage battery can be recharged by reversing the normal electrochemical reaction by applying an external voltage at the cathode and the anode.

47. Solid State Lithium Battery
Advantages:
Lithium has the most negative standard reduction potential value
Lithium is the lightest metal – smallest mass per mole of electrons.
Can be recharged literally hundreds of times without deterioration.

48. Fuel Cell
A fuel cell is an electrochemical cell that requires a continuous supply of reactants to keep functioning
In its simplest form:
Anode:
2H2 (g) + 4OH- (aq) H2O (l) + 4e-
Cathode:
O2 (g) + 2H2O (l) + 4e OH- (aq)
2H2 (g) + O2 (g) H2O (l)

49. A hydrogen-oxygen fuel cell used in the space program.
Unlike batteries, fuel cells do not store chemical energy.
Reactants must be constantly
be supplied.
Products must be constantly removed from a fuel cell.
The fuel cell does not operate like a heat engine and therefore is not subject to the same kind of thermodynamic limitations in energy conversion.
Fuel cells may be as much as 70 percent efficient, about
twice as efficient as an internal combustion engine.

50. Electrolysis is the process in which electrical energy is used to cause a nonspontaneous chemical reaction to occur.
Simplified Version
Downs Cell
Minimum of 4 V must be supplied
by the battery to carry out the reaction.
= -4 V

51. Electrolysis of Water

52. Electrolysis of an Aqueous Sodium Chloride Solution
Possible oxidation reactions:
Cl2 actually forms due to the high overvoltage of O2.
Overvoltage is the difference between the electrode potential and the actual voltage required to cause electrolysis.
Possible reduction reactions:
Due to low [H+]

55. Steps in calculating amounts of substances reduced or oxidized in electrolysis.
Molten CaCl2in an electrolytic cell.
Current of A is passed through the cell for 1.50 hr.