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Half Cell Reactions

Skills to develop

Half Cell Reactions

A half cell is one of the two electrodes in a galvanic cell or simple battery. For example, in the Zn-Cu battery, the two half cells make an oxidizing-reducing couple.

Placing a piece of reactant in an electrolyte solution makes a half cell. Unless it is connected to another half cell via an electric conductor and salt bridge, no reaction will take place in a half cell.

On the cathode, reduction takes place.

Oxidant + n e- ® Reductant
Example: Cu2+ + 2 e ® Cu
Cu2+ is the oxidizing agent and Cu the reducing agent.
On the anode, oxidation takes place. Reductant ® n e- + Oxidant
Example: Zn ® Zn2+ + 2 e-.
Zn is the reducing agent, and Zn2+ the oxidizing agent.

A battery requires at least two electrodes, the anode at which oxidation occurs, and the cathode at which reduction occurs. Reduction and oxidation are always required in any battery setup.

A battery operation requires an anode, a cathode, a load, and a salt bridge (if the salt bridge is not there already). These are the key elements of a battery.

Examples

Some example problems are given below to illustrate the kind of problems you are expected to solve.

Example 1

Example 2

Example 3

Example 4

The Hydrogen Half Cell

A half cell consists of an electrode and the species to be oxidized or reduced. If the material conducts electricity, it may be used as an electrode. The hydrogen electrode consists of a Pt electrode, H2 gas and H+. This half cell, is represented by:

Pt(s) | H2(g) | H+(aq)

where the vertical bars represent the phase boundaries. Conventionally, the cell potential for the hydrogen electrode is defined to be exactly zero if it has the condition as given below:

Pt | H2 (g, 1 atm) | H+(aq), 1 M

The notations for half cells are not rigid, but a simplified way to represent a rather complicated setup.

Standard Reduction Potential

The tendency for a reduction reaction is measured by its reduction potential.

Oxidant + n e- ® Reductant . . . Eo
For example: Cu2+ + 2 e ® Cu . . . Eo = 0.339 V
The reduction potential is a quantity measured by comparison. As mentioned earlier, the reduction potential of the standard hydrogen electrode (SHE) is arbitrary defined to be zero as a reference point for comparison. When a half cell Cu2+ || Cu for the reaction Cu2+ + 2 e ® Cu is coupled with the Standard Hydrogen Electrode (SHE), the copper electrode is a cathode, where reduction takes place. The potential accross the cell Pt | H2 (g, 1 atm) | H+(aq), 1 M || Cu2+ | Cu has been measured to be 0.339 V. This indicates that Cu2+ ions is easier to reduce than the hydrogen ions, and we usually represent it by Cu2+ + 2 e ® Cu . . . Eo = 0.339 V A positive cell potential indicates a spontaneous reaction.

When the cell Zn | Zn2+ is coupled with the SHE,

Zn | Zn2+> (aq) 1 M || H+(aq), 1 M | H2 (g, 1 atm) | Pt The potential has been measured to be 0.76 V. However, in this cell, Zn is oxidized, and its electrode is the anode. Therefore, the reduction potentail has a negative value for the reduction reaction Zn2+ + 2 e ® Zn . . . Eo = - 0.76 V This means that Zn2+ ions are less readily to accept electrons than hydrogen ions.

Ideally, for every redox couple, there is a reduction potential. Reduction potentials of standard cells have been measure against the SHE or other standards, their potentials are measured. This values are usually tabulated in handbooks. A short Standard Reduction Potentials table is available from the HandbookMenu, but you may also click the live link to see one.

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