CHM 1046
General Chemistry II
Dr. Michael Blaber


Electrochemistry

Balancing Oxidation-Reduction Reactions


In balancing redox reactions, gains and losses of electrons must be balanced

Zn(s) + 2H+(aq) ó Zn2+(aq) + H2(g)

Half-Reactions

Although oxidation and reduction must take place simultaneously (i.e. for something to be reduced, something else must be oxidized) it is often convenient to consider them as separate processes

Sn2+(aq) + 2Fe3+(aq) ® Sn4+(aq) + 2Fe2+(aq)

Oxidation: Sn2+(aq) ® Sn4+(aq) + 2e-

Reduction: 2Fe3+(aq) + 2e- ® 2Fe2+(aq)

Equations that show either oxidation, or reduction, alone are called half-reactions

Balancing Equations By The Method Of Half-Reactions

The use of half-reactions provides a general way to balance redox reactions

The reaction between permanganate ion (MnO4-) and oxalate ion (C2O42-) in acidic solution:

MnO4- + C2O42- ® Mn2+(aq) + CO2(g)

To balance this redox reaction using the method of half-reactions, begin by writing the incomplete oxidation and reduction half-reactions:

MnO4- ® Mn2+(aq)

C2O42- ® CO2(g)

Which compound is being reduced and which one is being oxidized?

Now, balance the atoms undergoing oxidation or reduction by the appropriate addition of coefficients.

MnO4- ® Mn2+(aq)

C2O42- ® 2CO2(g)

If the reaction is done under acidic aqueous solution, H+ and H2O can be added to reactants or products to balance H and O atoms. Likewise, if the reaction is done under basic aqueous solution, OH- and H2O can be added to balance H and O atoms

MnO4- ® Mn2+(aq) + 4H2O(l)

MnO4- + 8H+ ® Mn2+(aq) + 4H2O(l)

MnO4- + 8H+ + 5e- ® Mn2+(aq) + 4H2O(l)

(balanced atoms, balanced charge)

C2O42- ® 2CO2(g)

C2O42- ® 2CO2(g) + 2e-

(balanced atoms, balanced charge)

 

2MnO4- + 16H+ + 10e- ® 2Mn2+(aq) + 8H2O(l)

5C2O42- ® 10CO2(g) + 10e-

5C2O42- ® 10CO2(g) + 10e-

2MnO4- + 16H+ + 10e- ® 2Mn2+(aq) + 8H2O(l)

In other words, the oxalate ion is oxidized by the permananate ion (and the permanganate ion is reduced)

2MnO4- + 16H+ + 10e- ® 2Mn2+(aq) + 8H2O(l)
5C2O42-
® 10CO2(g) + 10e-

2MnO4- + 16H+ + 5C2O42- ® 10CO2(g) + 2Mn2+(aq) + 8H2O(l)


Summary of balancing half-reactions in acidic solutions:

1. Divide reaction into two incomplete half-reactions

2. Balance each half-reaction by doing the following:

a. Balance all elements, except O and H

b. Balance O by adding H2O

c. Balance H by adding H+

d. Balance charge by adding e-as needed

3. If the electrons in one half-reaction do not balance those in the other, then multiply each half-reaction to get a common multiple

4. The overall reaction is the sum of the half-reactions

5. The oxidation half-reaction is the one that produces electrons as products, and the reduction half-reaction is the one that uses electrons as reactants


Balancing Redox Reactions That Occur in Basic Solution

In a basic aqueous solution the half-reactions must be balanced using OH- and H2O (instead of H+ and H2O)

Here's an example of an unbalanced redox reaction that occurs under basic conditions:

CN-(aq) + MnO4-(aq) ® CNO-(aq) + MnO2(s)

CN-(aq) ® CNO-(aq)

MnO4-(aq) ® MnO2(s)

CN-(aq) ® CNO-(aq)

CN-(aq) + H2O(l) ® CNO-(aq)

CN-(aq) + H2O(l) ® CNO-(aq) + 2H+(aq)

CN-(aq) + H2O(l) ® CNO-(aq) + 2H+(aq) + 2e-

Now the other half-reaction:

MnO4-(aq) ® MnO2(s)

MnO4-(aq) ® MnO2(s) + 2H2O(l)

MnO4-(aq) + 4H+(aq) ® MnO2(s) + 2H2O(l)

MnO4-(aq) + 4H+(aq) + 3e- ® MnO2(s) + 2H2O(l)

Now balance electrons by using a common multiple:

3CN-(aq) + 3H2O(l) ® 3CNO-(aq) + 6H+(aq) + 6e-

2MnO4-(aq) + 8H+(aq) + 6e- ® 2MnO2(s) + 4H2O(l)

 

3CN-(aq) + 3H2O(l) ® 3CNO-(aq) + 6H+(aq) + 6e-

(to balance the 6H+ we need to add 6 OH-, to keep things balanced we add to both sides)

3CN-(aq) + 3H2O(l) + 6OH-(aq) ® 3CNO-(aq) + 6H+(aq) + 6OH-(aq) + 6e-

(the six H+ and OH- ions on the right side will neutralize to form 6H2O)

3CN-(aq) + 3H2O(l) + 6OH-(aq) ® 3CNO-(aq) + 6H2O(l) + 6e-

(we have three H2O on the left, and six on the right. We can simplify by just having three on the right)

3CN-(aq) + 6OH-(aq) ® 3CNO-(aq) + 3H2O(l) + 6e-

Now the other half-reaction:

2MnO4-(aq) + 8H+(aq) + 6e- ® 2MnO2(s) + 4H2O(l)

2MnO4-(aq) + 8H+(aq) + 8OH-(aq) + 6e- ® 2MnO2(s) + 4H2O(l) + 8OH-(aq)

2MnO4-(aq) + 8H2O(l) + 6e- ® 2MnO2(s) + 4H2O(l) + 8OH-(aq)

2MnO4-(aq) + 4H2O(l) + 6e- ® 2MnO2(s) + 8OH-(aq)

3CN-(aq) + 6OH-(aq) ® 3CNO-(aq) + 3H2O(l) + 6e-

2MnO4-(aq) + 4H2O(l) + 6e- ® 2MnO2(s) + 8OH-(aq)

with a net redox reaction of:

3CN-(aq) + H2O(l) + 2MnO4-(aq) ® 3CNO-(aq) + 2MnO2(s) + 2OH-(aq)


2000 Dr. Michael Blaber