The standard hydrogen half-cell

This is the apparatus set-up that is used to compare the electrode potentials of metal and other electrode systems. It comprises hydrogen gas at atmospheric pressure brought into contact with a platinum black electrode in 1 molar (mol dm^-3) acid solution. The hydrogen ions from the acid solution set up an equilibrium with the hydrogen gas. This set of conditions is known as standard conditions.

H₂ ⇋ 2H⁺ + 2e

If the other half-cell connected to the standard hydrogen electrode releases electrons, it can be said to be relatively negative with respect to the standard hydrogen half-cell (which is assigned a value of 0 volts)

For example, metals that are more reactive than hydrogen gas preferentially release electrons (forcing the standard hydrogen half-cell to accept the electrons) and have standard electrode potentials with negative values.

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The standard electrode potential

This is the 'potential' or tendency of a redox system to lose, or gain, electrons when compared to the standard hydrogen electrode - assigned a value of 0 volts. In any reduction-oxidation half-equation the electrons are gained by the species on the left hand side:

Cu²⁺ + 2e ⇋ Cu

This is an equilibrium, and so if a more powerful reducing agent enters into electrical contact with the above system it can force the copper ions to accept electrons and push the equilibrium to the right hand side. Conversely, if a weaker reducing agent is brought into contact with the above equilibrium then the copper can force it to accept electrons allowing its own equilibrium to move to the left hand side.

The electrode potential measures the tendency of electrons to flow away from or towards a redox equilibrium. They are always measured with respect to the standard hydrogen electrode (which is assigned a value of zero volts).

Equilibrium redox systems with the reduced side (usually a metal) more reactive than hydrogen have a negative electrode potential, i.e. they can lose electrons more easily than hydrogen. Equilibrium redox systems with the reduced side less reactive than hydrogen have a positive electrode potential, i.e. they can lose electrons less easily than hydrogen.

Note: Details of the standard hydrogen electrode are not required for first examinations 2025.

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Using the Standard Hydrogen half-cell

In the actual experimental setup, the two half-cells are connected together under standard conditions, via an external circuit and a salt bridge to make the whole cell.

In the above apparatus set-up, the zinc|zinc sulfate(aq) half-cell is connected to the SHE via an external circuit which includes a high resistance voltmeter (high resistance to prevent passage of current). The salt bridge completes the circuit - it allows ions to flow from one side to another to equalise the movement of charge.

Once the apparatus is setup, the reading on the high resistance voltmeter records the standard electrode potential of the Zn|Zn²⁺_(aq) system

In this particular case the voltmeter reads -0.76 V, indicating that zinc is more reactive than hydrogen; that there is a force pushing electrons around the external circuit to the hydrogen half-cell from the zinc half-cell.

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Cell representation

The whole cell can be represented by showing the half cells in order of phase (solid, | solution, |salt bridge | solution | solid)

The above cell diagram can be represented as:

Zn_(s) |Zn²⁺_(aq) ||H⁺_(aq) | H_2(g),Pt

By convention the half cell that provides the electrons, i.e. the best reducing agent, is written on the left hand side. In this representation the Zn|Zn²⁺(aq) is the most negative potential and is behaving as the anode, the electrode where oxidation takes place:

Zn → Zn²⁺ + 2e

The right hand half cell, H⁺_(aq) | H_2(g),Pt, is the electrode where reduction is forced to take place, i.e. the cathode, where electrons are available.

2H⁺ + 2e → H₂

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Electrochemical series

By comparing many redox systems with the SHE and other reference electrodes, a series can be drawn up showing the reductions in order of their standard electrode potential (usually from negative at the top to positive at the bottom)

Using the electrochemical series, the species at the top on the right hand side are reducing agents, and the species at the bottom on the left hand side are oxidising agents.

redox half-equation	Eº/ V
Mg2+(aq) + 2e- → Mg(s)	-2.38
Zn2+(aq) + 2e- → Zn(s)	-0.76
Fe2+(aq) + 2e- → Fe(s)	-0.41
Pb2+(aq) + 2e- → Pb(s)	-0.13
2H+(aq) + 2e- → H2(g)	0.00
Sn4+(aq) + 2e- → Sn2+(aq)	+0.15
Cu2+(aq) + 2e- → Cu(s)	+0.34
Ag+(aq) + e- → Ag(s)	+0.80
Cl2(g) + 2e- → 2Cl-(aq)	+1.36

The value of the electrode potential is a relative value for the equilibrium compared with the standard hydrogen electrode. Its sign never changes regardless of which way round the equilibrium is written. By convention, the equilibria are written as reductions (left to right) but they could just as easily be written as oxidations, without changing the sign of the electrode potential.

Full standard electrode potentials table

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