3.5.3 Redox Equilibria - Electrode potentials


Students should:
  • know the IUPAC convention for writing half-equations for electrode reactions
  • know and be able to use the conventional representation of cells
  • understand how cells are used to measure electrode potentials by reference to the standard hydrogen electrode
  • know the importance of the conditions when measuring the electrode potential, E (Nernst equation not required)
  • know that standard electrode potential, E , refers to conditions of 298 K, 100 kPa and 1.00 mol dm-3 solution of ions

Half-equation convention

These are equations representing either the reduction or the oxidation that takes place in a redox reaction. However, as these haf equations are reversible, the convention is to always write them as reductions. In other words a species gaining electrons:

Zn2+(aq) + 2e Zn(s)


Cell representation

Electrochemical cells are represented using a convention of vertical lines to indicate phase or state barriers and commas to show similar phases in contact. The salt bridge is shown as a double vertical line.

Hence the following voltaic cell:

The Zinc - Copper electrochemical cell. The Voltaic cell

Can be represented by the following convention: Zn|Zn2+(aq)||Cu2+(aq)|Cu

The anode (where oxidation takes place) is shown at the left hand side and the cathode at the right hand side.


The standard hydrogen electrode

The position of equilibrium of an electrochemical half cell cannot be measured directly. We know that one half cel has a tendency to push electrons towards another and we can measure the potential difference between two half cells, but there is no way of knowing the actual absolute position of the equilibrium set up between a species and its reduced (or oxidised) counterpart.

The actual position of the equilibrium:

Zn2+(aq) + 2e Zn(s)

cannot be measured.

So, what we do is measure the potential difference between each half cell and a standard half cell called the standard hydrogen electrode:

the standard hydrogen electrode

2H+(aq) + 2e H2(g)

The standard hydrogen electrode is defined as the reference with a value of exactly zero volts electrical potential, 0.0 V. Any other half cell can then be attached to the standard hydrogen electrode and its electrode potential recorded:

measurement of electrode potential of a zinc half cell using the standard hydrogen electrode

The conditions are important, because the electrode potential is affected by both concentration of the solutions and the temperature. In the case of the zinc half cell it has an electrode potential of -0.76 V. This means that compared to the standard hydrogen electrode, it forms the relatvely negative electrode (it provides electrons) by the process:

Zn(s) Zn2+(aq) + 2e

While the standard hydrogen elctrode is relatively more positive and undergoes the reduction reaction

2H+(aq) + 2e H2(g)

And the overall reaction for the zinc-standard hydrogen electrode cell: Zn|Zn2+(aq)||H+(aq)|Pt|H2(g):

Zn(s) + 2H+(aq) Zn2+(aq) + H2(g)