Inorganic and Applied Chemistry

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Inorganic and Applied Chemistry

6.2 Galvanic cells

We saw in the previous section how redox reactions involves the transfer of electrons and how the oxidation

means loss of electrons (increase in level of oxidation) while the reduction means increase in electrons

(decrease in level of oxidation). In order to understand how a redox reaction may generate current we first

look at the following reaction between MnO 4 - and Fe2+ that was balanced in example 6-B.

8 H+(aq) + MnO 4 - (aq) + 5 Fe2+(aq)  Mn2+(aq) + 5 Fe2+(aq) + 4 H 2 O(l)

In this reaction Fe2+ is oxidised while MnO 4 - is reduced. Electrons are thereby being transferred from Fe2+ to

MnO 4 -. We will look further into each of the half-reactions separately. We have the reduction reaction:

8 H+(aq) + MnO 4 - (aq) + 5 e-  Mn2+(aq) + 4 H 2 O(l)

and similarly the oxidation reaction:

5 Fe2+(aq)  5 Fe3+(aq) + 5 e-

Note how the oxidation reaction must take place five times each time the reduction reaction takes place.

When MnO 4 - and Fe2+ are present in the same solution the electrons are transferred directly when the

reactants collide. Under these conditions no energy may be extracted as all the chemical energy is lost as heat.

How may we then extract energy from the reaction above? The answer is to separate the oxidation reaction

from the reduction reaction which requires a wire (e.g. of cobber). The current that runs between the two

solutions may be let through e.g. an electrical bulb or through an electrical motor. Thereby we have extracted

energy from a chemical reaction. This principle requires that a salt bridge is established in addition to the

wire in order to allow transport of ions without completely mixing the solutions significantly. This principle

is sketched in figure 6-1.

Electrochemistry