Chemistry - A Molecular Science

11.5

WRITING REDOX REACTIONS The number of electrons gained in the reduction half-reaction must equal the number lost in the oxidation half-reaction, so both the at

oms and the number of electrons gained and

lost must be balanced. The number of electrons transferred must be a multiple of the number shown for the half-reaction in Table 11.1. For example, consider the chemical equation for dissolving metallic copper in

nitric acid. The two half-reactions are

Cu

2+ + 2e

1-^

U

Cu

oE

= +0.34 V

NO

1- + 4H 3

1+ + 3e

1-

U

NO + 2H

O* 2

oE

= +0.96 V

The reactants, copper and nitric acid, are high

lighted in yellow. All of the reactions in

Table 11.1 are written as reductions, so one half-reaction must always be reversed to make it the oxidation. In this case, the Cu

2+/Cu couple must be reversed because the half-

reaction shows the metallic c

opper as a product, but it is a reactant in the reaction with

nitric acid. In addition, both equations mu

st be multiplied by an integer to make the

number of electrons gained in the reduction of

nitric acid equal to the number of electrons

lost in the oxidation of copper. Each nitrat

e ion requires three electrons, but each copper

atom gives up only two. The lowest common multiple (LCM)

† of three and two is six, so

the nitric acid half-reaction must be multiplied

by two to obtain a six-electron gain, and the

copper half-reaction must be reversed and multip

lied by three to obtain a six-electron loss.

Summing the two equations yields the overall equation.

† The lowest common multiple (LCM) of two numbers is the

smallest multiple that is exactly divisible by both numbers.

3Cu

U

3Cu

2+ + 6e

1- anode (oxidation) half-reaction

2NO

1- + 8H 3

1+ + 6e

1-

U

2NO + 4H

O 2

cathode (reduction) half-reaction

3Cu + 2NO

1- 3

+ 8H

1+ + 6e

1-U^

3Cu

2+ + 6e

1- + 2NO + 4H

O 2

The 6e

1- on each side cancel, yielding the net balanced equation: 3Cu + 2NO

1- 3

+ 8H

1+^

U

3Cu

2+ + 2NO + 4H

O 2

Electrons flow spontaneously from lower to

higher potential, so redox reactions are

extensive when the reducing agen

t is at a lower potential than the oxidizing agent. That is,

electron transfer is extensive when any

half-reaction is coupled with the

reverse

of a half-

reaction that is above it in Table 11.1

. This reactivity is summarized in Figure 11.4. The

standard potential of the reaction would be th

e following if it were carried out in an

electrochemical cell:

oE

rxn

=

oE

reduced

• E

ooxidized

Eq.

11.6

WeakerOxidizingAgent

WeakerOxidizingAgent

StrongerOxidizingAgent

StrongerOxidizingAgent

WeakerReducingAgent

WeakerReducingAgent

StrongerReducingAgent

StrongerReducingAgent

1-e

1-e

Standard Reduction Potential

Free Energy

• 
  
  
  
  • 
    
  
  
  
  

Extensive

Not Extensive

(a)

X (b)

Figure 11.4 Relative position of reactants and products on a standard reduction potential chart (a) The reaction of the stronger

oxidizing and reducing agents to

produce the weaker oxidizing and reducing agents is extensive. (b) The reaction of the weaker oxidizing and reducing agents to produce stronger oxidizing and reducing agent

s is not extensive. Electrons are

transferred extensively downhill in Table 11.1 because that is the direction of more positive potential and lower free energy.

oE

reduced

is the standard reduction potential of the couple that is reduced, and

oE

oxidized

is that

of the couple that is oxidized during the reaction.

Chapter 11 Electron Transfer and Electrochemistry