Chapter 11 Electron Transfer and Electrochemistry

Discharging (Galvanic) Anode

Pb(s) + SO

2- 4

U

PbSO

(s) + 2e 4

1-^

Cathode PbO

(s) + 4H 2

1+ + SO

2- 4

+ 2e

1-^ U

PbSO

(s) + 2H 4

O 2

Electron flow:

anode

→

starter

→

cathode

Charging (Electrolytic) Cathode PbSO

(s) + 2e 4

1-^ U

Pb(s) + SO

2- 4

Anode PbSO

(s) + 2H 4

O 2

U

PbO

(s) + 4H 2

1+ + SO

2- 4

+ 2e

1-^

Electron flow:

anode

→

alternator

→

cathode

battery discussed in Section 11.6 and represented in the margin. When the starter is used, the battery is a galvanic cell because elect

rons spontaneously flow from the Pb anode

through the starter, where their energy is

used to start the car, and then to the PbO

(^2)
cathode. As this occurs, lead, lead oxide, and
sulfuric acid are converted to lead sulfate
and water. The ‘cranking power’ of
the battery is the rate at which the reaction occurs. A
battery rated with 550 amps of ‘cra
nking power’ generates ~1.7 g of PbSO
per second 4
while the starter is in use. Clearly, the battery
would not be able to perform for very long
at that rate before all of the chemical reac
tants would be consumed, and the battery would
be ‘dead’. To increase the life of the battery
, part of the free energy derived from the
combustion of the gasoline when the car is opera
ting is used to run an electrical generator
(alternator), which supplies the electrical n
eeds of the car and recharges the battery. To
recharge the battery, a voltage greater than the 12 V the battery delivers as a galvanic cell is applied across the electrodes in the reve
rse direction. The large voltage forces the
battery to operate as an
electrolytic cell
as electrons are forced through the battery in the
reverse direction, oxidizing PbSO
to produce PbO 4
at the anode and reducing PbSO 2
to 4
Pb at the cathode. Thus, some of the free en
ergy from the combustion of the gasoline is
used to pump the electrons back uphill fro
m lower energy orbitals into higher energy
orbitals.
11.9
CHAPTER SUMMARY AND OBJECTIVES Reduction is the gain of electrons, and oxida
tion is the loss of electrons. Spontaneous
redox reactions involve a transfer of electrons from a donor (reducing agent) to an acceptor (oxidizing agent) that is lower in free energy. The free energy released by the reaction can be harnessed if the oxidation and
reduction half-reactions are separated into
compartments. Oxidation always occurs in
the anode compartment and reduction always
occurs in the cathode compartment. By measu
ring the cell potentials of the various half-
reactions relative to the standard hydrogen
electrode (SHE), we obtain the standard
reduction potentials of the half-reactions. A large and negative standard reduction potential means that the reduced form is a very good reducing agent, while a large and positive standard reduction potential implies that the oxidized form is a good oxidizing agent.
The cell potential is the potential differen
ce (voltage) between the cathode and the
anode:
(^) Ecell

Ecathode

(^) Eanode

. If

(^) Ecell

0, the reaction is spontaneous, which means that
energy can be extracted from it (
G < 0). This type of cell is a galvanic cell. Galvanic Δ