Chemistry - A Molecular Science

powderedPbO cathode

2 (+)

spongyPb anode

(-)

HSO^24

+

• 
  

metallicconductors

0V

2V

4V

6V

Figure11.7 Lead-storage battery The potential of the cathode of

each cell is 2 V higher than

that of the anode. Metallic conductors keep the anode of each cell at the same potential as the

cathode of the previous cell.

For example, the anode of the

first cell is at 0 V, so its

cathode is at 2 V. The anode of

the second cell is held at 2 V

also, so the cathode potential is 4 V. The battery in the figure contains three cells and each cell increases the cell potential by 2 V, so the battery is a 6-V battery.

LEAD-STORAGE BATTERY Also known as the lead-acid battery, the lead-storage battery is the source of power for starting automobiles. A 12-V battery consists of six cells, each with a potential difference of 2V, connected in series. As shown in Figur

e 11.7, each cell consists of two lead grids,

one packed with spongy Pb (anode) and the other with powdered PbO

(cathode). The 2

cells are immersed in ~4.5 M sulfuric acid (H

SO 2

). The half-reactions for the lead storage 4

battery are

Anode reaction:

Pb(s) + SO

2- 4

(aq)

→

PbSO

(s) + 2e 4

1-^

Cathode reaction:

PbO

(s) + 4H 2

1+(aq) + SO

2-(aq) + 2e 4

1-^

→

PbSO

(s) + 2H 4

O(l) 2

Thus, when you start your car, you do so by

utilizing the free energy of the reaction

PbO

(s) + Pb(s) + 4H 2

1+(aq) + 2SO

2- 4

(aq)

→

2PbSO

(s) + 2 H 4

O(l) 2

oE

= 2.05 V

Because the products are deposited on the elec

trodes, this reaction can be reversed by

applying an external voltage across the battery. Thus, while your car is running, some of the energy of the combustion of gasoline is us

ed to turn the generator (alternator), which

generates the required voltage to reverse th

e battery reaction, thereby recharging the

battery. Recharging the lead-storage ba

ttery is discussed in Section 11.8.

11.7

CORROSION Many metals are good reducing agents, and t

hose with reduction potentials more negative

than -0.41 V (the non-standard reduction pote

ntial of pure water) react with water and

oxygen. Such reactions can be very benefici

al, but some can also be extremely costly.

Corrosion

is the

unwanted

oxidation of a metal. Approximately 20 to 25% of the steel

produced in the United States is for the repl

acement of corroded steel! Although most

metals corrode, we focus our discussion on th

e corrosion of iron, as it is the one most

commonly encountered. We begin with the following observations:

1. A piece of corroded iron contains regions where rust (Fe

O 2

) has accumulated and 3

regions where holes have formed in the iron

. Due to the presence of holes in the

corroded iron, it is often said to be ‘pitt

ed’. The two regions are often separated.

2. Iron will not rust in a dry climate; water is required. 3. Iron will not rust in water in the absence of O
   . 2
   4. Rusting is enhanced in

the presence of acid.

Chapter 11 Electron Transfer and Electrochemistry