The Hydrogen Electrode (HL) (DP IB Chemistry): Revision Note

The hydrogen electrode

• The absolute value of a half-cell potential cannot be measured, only differences in potential between pairs of half-cells

• A standard reference electrode is required to compare these potentials

• The standard hydrogen electrode is used as a reference electrode and consists of:

  • Hydrogen gas at 100 kPa

  • In equilibrium with H⁺ ions at 1.00 mol dm^-3 

  • An inert platinum electrode in contact with the hydrogen gas and H⁺ ions

• The half-equation for the standard hydrogen electrode is:

2H⁺ (aq) + 2e^– ⇌ H₂ (g)

• It is given an arbitrary value of E^θ = 0.00 volts

• When the standard hydrogen electrode is connected to another half-cell, the standard electrode potential of that half-cell can be read from a high-resistance voltmeter

Standard hydrogen electrode diagram

• In practice, the standard hydrogen electrode is rarely used for a number of reasons:

  • The electrode reaction is slow

  • The electrodes are not easily portable

  • It is difficult to maintain a constant pressure

• Once one standard electrode potential is known relative to the standard hydrogen electrode, it can be used as a reference for measuring others

  • So, any half-cell with a known potential can then serve as a secondary reference

Measurements using the hydrogen electrode

• If a hydrogen electrode is used to measure the electrode potentials of zinc and copper half reactions, the conventional cell diagrams would be:

Pt 丨H₂(g), 2H⁺(aq)  ∥  Zn²⁺ (aq), Zn (s)          E^θ = -0.76 V

Pt 丨H₂(g), 2H⁺(aq)  ∥  Cu²⁺ (aq), Cu (s)           E^θ = +0.34 V

• The hydrogen electrode is always written on the left

• The polarity of the other half-cell is measured relative to hydrogen

• The half-reaction is always shown as a reduction

  • This is why these values are called standard reduction potentials

• Standard reduction potential tables list half-cells from most negative to most positive

  • The IB Chemistry data booklet (Section 19) contains a table of standard reduction potentials at 298.15 K

Table of standard electrode potentials

Interpreting electrode potential values

• The electrode potential of a half-cell tells us how easily the ions and atoms involved in the half-equation are reduced or oxidised

• For example, in the half-equation Cu²⁺ + 2e^- ⇌ Cu, the species being reduced is Cu²⁺ and the species being oxidised is Cu

• The electrode potential reflects how readily electrons are gained or lost in this equilibrium

• A more negative E^θ value:

  • The species loses electrons more readily

  • The equilibrium lies to the left

  • Favours oxidation

  • Indicates a stronger reducing agent

• A more positive E^θ value:

  • The species gains electrons more readily

  • The equilibrium lies to the right

  • Favours reduction

  • Indicates a stronger oxidising agent