A single half-cell has no measurable voltage on its own, so chemists compare each one with the standard hydrogen electrode. The number that comes out, E⊖, ranks every redox couple by how strongly it oxidises or reduces.
The standard electrode potential, E⊖, of a half-cell is the potential difference between it and the standard hydrogen electrode (SHE), measured under standard conditions: 298 K, 1.00 mol dm−3 for every ion in solution and 100 kPa (Cambridge also accepts 101 kPa) for any gas. The SHE is H2 gas bubbling over platinum (coated with platinum black) in 1.00 mol dm−3 H+, and is defined as 0.00 V. A metal dips into its own ions; an ion/ion couple such as Fe3+/Fe2+, or a non-metal such as Cl2/Cl−, needs an inert platinum electrode. Half-equations are written as reductions, so the more positive E⊖, the stronger the oxidising agent on the left; the more negative, the stronger the reducing agent on the right.
Set up the standard hydrogen electrode against a half-cell in the simulator, then build reduction half-equations: balance atoms, then add electrons to balance charge. Each balanced half-equation is a half-cell whose E⊖ you can look up.
| Half-cell type | Example | Electrode | Solution |
|---|---|---|---|
| hydrogen (reference) | 2H+ + 2e− ⇌ H2 | Pt, H2 at 100 kPa | 1.00 mol dm−3 H+ (e.g. HCl) |
| metal / metal ion | Cu2+ + 2e− ⇌ Cu | the metal itself | 1.00 mol dm−3 Cu2+ |
| ion / ion | Fe3+ + e− ⇌ Fe2+ | platinum | 1.00 mol dm−3 of both Fe3+ and Fe2+ |
| non-metal / ion | Cl2 + 2e− ⇌ 2Cl− | platinum | Cl2 at 100 kPa, 1.00 mol dm−3 Cl− |
Four quick checks. Each correct answer earns XP and records this skill on your profile.
What is the standard electrode potential of the standard hydrogen electrode?
Which set of conditions is standard for measuring E⊖?
Which half-cell must use a platinum electrode?
Using E⊖ values, which is the strongest oxidising agent: Cl2 (+1.36 V), I2 (+0.54 V), Cu2+ (+0.34 V) or Zn2+ (−0.76 V)?
Pick oxidising and reducing agents in the simulator and read the table from top to bottom: oxidising agents on the left get stronger; reducing agents on the right get stronger the other way.
| Half-equation | E⊖ / V | What it tells you |
|---|---|---|
| Zn2+ + 2e− ⇌ Zn | −0.76 | Zn is a good reducing agent |
| 2H+ + 2e− ⇌ H2 | 0.00 | reference |
| Cu2+ + 2e− ⇌ Cu | +0.34 | Cu does not reduce H+ |
| I2 + 2e− ⇌ 2I− | +0.54 | weak oxidising agent |
| Fe3+ + e− ⇌ Fe2+ | +0.77 | Fe3+ oxidises I− |
| Ag+ + e− ⇌ Ag | +0.80 | Ag+ oxidises Cu |
| Cr2O72− + 14H+ + 6e− ⇌ 2Cr3+ + 7H2O | +1.33 | strong oxidising agent |
| Cl2 + 2e− ⇌ 2Cl− | +1.36 | stronger still |
| MnO4− + 8H+ + 5e− ⇌ Mn2+ + 4H2O | +1.52 | strongest here |
E⊖ values are always quoted for the reduction. Do not double E⊖ when a half-equation is multiplied up: it is a property of the couple, not of the amount. An ion/ion or gas/ion half-cell needs a platinum electrode and, for ion/ion, equal 1.00 mol dm−3 concentrations of both ions. A more positive E⊖ means the species on the left is a stronger oxidising agent, not that the species on the right is.
Unlocks once the four checks above are done. Worth more XP, written to A Level Paper 4 standard.
Using the Data Booklet values, which species is the strongest reducing agent: Zn, Cu, Fe2+ or I−?
When the Cu2+/Cu half-cell is connected to the SHE, the voltmeter reads 0.34 V. Which statement is correct?
Why is the platinum in the standard hydrogen electrode coated with platinum black?
Which half-cell would be the correct standard set-up for the Cr2O72−/Cr3+ couple?
This skill is now recorded on your profile. Keep the chain going.