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Standard electrode potentials provide the quantitative framework that turns the qualitative redox chemistry of lessons 0 and 1 into predictive thermodynamics. Every redox couple — a half-equation written as a reduction — is assigned a single number, the standard electrode potential E°, measured in volts and referenced to the standard hydrogen electrode (SHE), which by international convention is assigned E° = 0.00 V. The SHE consists of a platinised platinum electrode immersed in 1.00 mol dm⁻³ H⁺(aq) with hydrogen gas at 100 kPa bubbling over it at 298 K. This lesson develops the operational definition of E°, the four standard conditions that fix its value (298 K, 100 kPa, 1.00 mol dm⁻³ ions, Pt for non-metallic couples), the electrochemical series of representative values, and the construction of a salt-bridge cell that allows E° to be measured in the laboratory.
Spec mapping (AQA 7405): This lesson anchors §3.1.11 (electrode potentials and electrochemical cells), focusing specifically on the definition of E°, the standard hydrogen electrode, the standard conditions, and the electrochemical series. It builds directly on lessons 0 and 1 of this course (oxidation states, half-equations, oxidising and reducing agents), and is required prior reading for lesson 3 (electrochemical cell construction and notation), lesson 4 (cell EMF and the prediction of feasibility), and any later treatment of fuel cells and storage cells. The thermodynamic bridge to §3.1.8 — ΔG° = −nFE° — is introduced at A* level here and developed quantitatively in lesson 4. Refer to the official AQA specification document for the exact wording of each section.
Assessment objectives: AO1 recall items include defining E°, describing the construction and operation of the SHE, stating the standard conditions, and quoting the sign convention (a more positive E° identifies the stronger oxidising agent on the left of the reduction half-equation). AO2 calculations dominate the algebra: E°cell = E°(cathode) − E°(anode), with both half-equations expressed as reductions, and the identification of cathode and anode from the two given E° values. AO3 reasoning appears whenever a question asks for the direction of spontaneous reaction, the feasibility of a displacement, or a synoptic link to ΔG°. The high-mark items typically combine a numerical E°cell calculation with a written explanation of what the sign tells you about the chemistry.
The SHE is the universal reference half-cell against which every other electrode potential is measured. Its construction is fixed by convention:
The half-equation occurring at the platinum surface is:
2H⁺(aq) + 2e⁻ ⇌ H₂(g) E° = 0.00 V (by definition)
The platinum itself takes no part in the redox chemistry — it provides a chemically inert electronic conductor and a catalytic surface for the equilibrium between protons and hydrogen molecules. The value E° = 0.00 V is assigned by international agreement; it is not measured. Every other electrode potential in the electrochemical series is reported relative to the SHE.
Key Point: The SHE is awkward to operate (hydrogen gas at controlled pressure, platinum black surfaces that are easily poisoned, the safety hazard of flowing H₂) and is rarely used outside primary metrology laboratories. In routine practice it is replaced by secondary reference electrodes such as the silver–silver chloride (Ag/AgCl) electrode or the calomel electrode (see Going Further). All such secondary electrodes are themselves calibrated against the SHE.
Standard electrode potentials are only defined under a specific set of conditions. Outside these conditions the measured potential changes and is denoted simply E (no superscript). The four standard conditions are:
| Parameter | Standard value |
|---|---|
| Temperature | 298 K (25 °C) |
| Pressure of any gaseous species | 100 kPa |
| Concentration of any aqueous ion | 1.00 mol dm⁻³ |
| Electrode for non-metallic couples | Inert platinum |
Two further constructional features are required of every measurement:
Common Misconception: "Standard" does not mean "STP". The standard temperature for electrochemistry is 298 K, not 273 K, and the standard pressure is 100 kPa (the modern thermodynamic standard), not 101.325 kPa (atmospheric). Confusing these conventions with the molar-gas-volume conventions of lesson 4 of the atomic-structure course is a common source of error.
By international convention, every electrode potential is reported for the half-equation written as a reduction — that is, with the electrons on the left-hand side:
oxidised form + n e⁻ ⇌ reduced form E°
The numerical sign of E° then carries unambiguous chemical information:
If the half-equation is reversed (written as an oxidation, with electrons on the right), the sign of E° is reversed. This is a notational change, not a physical change — the chemistry is unaltered.
Worked example. The copper half-cell:
Cu²⁺(aq) + 2e⁻ ⇌ Cu(s) E° = +0.34 V
reversed reads
Cu(s) ⇌ Cu²⁺(aq) + 2e⁻ E° = −0.34 V
Both expressions describe the same equilibrium. The convention to quote E° for the reduction direction is what makes electrochemical series tables internally consistent.
The electrochemical series ranks redox couples in order of their standard electrode potentials. Representative AQA-relevant values, all quoted relative to the SHE at 298 K, 100 kPa and 1.00 mol dm⁻³, are listed below. The values are arranged from the most positive (strongest oxidising agents on the left) to the most negative (strongest reducing agents on the right).
| Half-equation (reduction) | E° / V |
|---|---|
| F₂(g) + 2e⁻ ⇌ 2F⁻(aq) | +2.87 |
| MnO₄⁻(aq) + 8H⁺(aq) + 5e⁻ ⇌ Mn²⁺(aq) + 4H₂O(l) | +1.51 |
| Cl₂(g) + 2e⁻ ⇌ 2Cl⁻(aq) | +1.36 |
| O₂(g) + 4H⁺(aq) + 4e⁻ ⇌ 2H₂O(l) | +1.23 |
| Br₂(l) + 2e⁻ ⇌ 2Br⁻(aq) | +1.07 |
| Fe³⁺(aq) + e⁻ ⇌ Fe²⁺(aq) | +0.77 |
| I₂(s) + 2e⁻ ⇌ 2I⁻(aq) | +0.54 |
| Cu²⁺(aq) + 2e⁻ ⇌ Cu(s) | +0.34 |
| 2H⁺(aq) + 2e⁻ ⇌ H₂(g) | 0.00 (SHE reference) |
| Pb²⁺(aq) + 2e⁻ ⇌ Pb(s) | −0.13 |
| Sn²⁺(aq) + 2e⁻ ⇌ Sn(s) | −0.14 |
| Fe²⁺(aq) + 2e⁻ ⇌ Fe(s) | −0.44 |
| Zn²⁺(aq) + 2e⁻ ⇌ Zn(s) | −0.76 |
| Mg²⁺(aq) + 2e⁻ ⇌ Mg(s) | −2.37 |
| Li⁺(aq) + e⁻ ⇌ Li(s) | −3.04 |
Reading the table from top to bottom: F₂ is the strongest oxidising agent in the list (most positive E°); Li(s) is the strongest reducing agent (most negative E°). The two extremes bracket the range of redox behaviour that aqueous chemistry can sustain. Outside this window (E° significantly more positive than +1.23 V or more negative than 0.00 V, both relative to water), the solvent itself is reduced or oxidised — see the depth misconceptions section.
The sequence of metal couples (Mg, Zn, Fe, Pb, Cu …) reproduces the reactivity series that GCSE students learnt qualitatively. The thermodynamic basis for that series is now explicit: a metal with a more negative E° than another can reduce the latter's ions in displacement reactions.
When two half-cells are connected by a salt bridge and a wire, electrons flow through the external circuit and ions migrate through the bridge. The cell's electromotive force (EMF), denoted E°cell when measured under standard conditions, is the potential difference between the two electrodes at zero current. The formula is:
E°cell = E°(cathode) − E°(anode)
with both E° values written as reduction potentials taken directly from the electrochemical series. The cathode is the electrode where reduction occurs — the half-cell with the more positive E°. The anode is the electrode where oxidation occurs — the half-cell with the more negative E°. Identifying which half-cell is which is therefore the first algorithmic step in any cell-EMF calculation.
Two half-cells are combined: a zinc electrode in 1.00 mol dm⁻³ Zn²⁺(aq) and a copper electrode in 1.00 mol dm⁻³ Cu²⁺(aq), linked by a KNO₃ salt bridge.
From the table:
The copper couple has the more positive E°, so reduction occurs at the copper electrode (the cathode) and oxidation at the zinc electrode (the anode).
E°cell = E°(cathode) − E°(anode) = (+0.34) − (−0.76) = +1.10 V
The positive E°cell tells us the cell reaction proceeds spontaneously in the forward direction as written:
Zn(s) + Cu²⁺(aq) → Zn²⁺(aq) + Cu(s) E°cell = +1.10 V
The two relevant half-equations are:
E°cell = (+1.23) − (0.00) = +1.23 V
Overall reaction: 2H₂(g) + O₂(g) → 2H₂O(l) with E°cell = +1.23 V. The size of this EMF is the thermodynamic ceiling on the voltage a hydrogen fuel cell can deliver; real cells produce less because of overpotentials and ohmic losses (kinetic effects, not thermodynamic).
E°cell = (+0.77) − (+0.34) = +0.43 V
Overall (after balancing electrons by multiplying the iron half-equation by 2):
2Fe³⁺(aq) + Cu(s) → 2Fe²⁺(aq) + Cu²⁺(aq) E°cell = +0.43 V
The Fe³⁺ ion is therefore a strong enough oxidising agent to oxidise copper metal to Cu²⁺ — a prediction that turns out to be experimentally correct.
Exam Tip: When balancing electrons across the two half-equations, you may need to multiply one half-equation by an integer to match electron counts. Do not multiply E° itself. Standard electrode potentials are intensive quantities — they do not scale with the amount of reaction, only with the chemical identity of the couple. This is the single most common error in electrochemistry calculations.
The sign of E°cell is the diagnostic for spontaneity under standard conditions:
The bridge to thermodynamics is the equation ΔG° = −nFE°cell, where n is the number of moles of electrons transferred per mole of reaction and F is the Faraday constant (96 485 C mol⁻¹). This formula is developed quantitatively in lesson 4; here it suffices to note that a positive E°cell guarantees a negative ΔG° and therefore a thermodynamically feasible reaction.
Crucial caveat: thermodynamics is not kinetics. A positive E°cell tells you only that the reaction is thermodynamically feasible. It does not tell you whether the reaction proceeds at an observable rate. Many reactions with large positive E°cell values are kinetically inert — the classic example is the room-temperature decomposition of water into H₂ and O₂, which is wildly favourable thermodynamically (ΔG° large and negative) yet does not occur spontaneously because the activation energy is prohibitive. Catalysts, elevated temperature, or electrical input are required to make kinetically slow but thermodynamically favourable reactions go.
The practical construction of a half-cell depends on the type of redox couple involved.
These are the simplest. A strip of the pure metal is partially immersed in a 1.00 mol dm⁻³ aqueous solution of one of its own salts. The metal itself serves as both electrode and active redox species. Examples: a copper rod in 1.00 mol dm⁻³ CuSO₄(aq); a zinc rod in 1.00 mol dm⁻³ ZnSO₄(aq).
When both the oxidised and reduced forms are aqueous ions, neither species provides a solid electrode. The half-cell is constructed with an inert platinum electrode immersed in a solution containing both species, each at 1.00 mol dm⁻³.
The platinum is electronically conductive, chemically inert, and provides a surface at which the electron-transfer equilibrium can occur. It does not participate stoichiometrically.
For half-cells involving a gas (H₂, Cl₂, O₂), the gas is bubbled at 100 kPa over an inert platinum electrode immersed in a solution containing the corresponding ion at 1.00 mol dm⁻³. The SHE itself is the prototype of this construction.
The salt bridge completes the electrical circuit between the two half-cells without allowing the bulk solutions to mix. Its functions are:
Practical forms include a U-tube of saturated KNO₃ in 5% agar gel, a strip of filter paper soaked in KNO₃ solution, or a porous frit/disc separating two compartments. The choice of bridge material does not affect E°cell provided the bridge ions are inert and the bridge resistance is small compared with the voltmeter's input impedance.
E° values apply only under the four standard conditions listed earlier. Outside those conditions:
For these reasons, electrode potentials are best regarded as a thermodynamic database for predicting direction and feasibility, not absolute rates or non-standard equilibria. Lesson 4 develops the feasibility tests in detail; the Nernst correction for non-standard conditions is touched on under Going Further and developed quantitatively in first-year undergraduate physical chemistry.
For a Required-Practical-adjacent investigation of E°cell:
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