Standard Electrode Potential
The standard electrode potential (E°) is the potential difference developed at an electrode when it is in contact with a 1 mol L⁻¹ solution of its ions at 25 °C (298 K) and 1 atm pressure, measured relative to the standard hydrogen electrode (SHE), which is assigned a potential of exactly 0.00 V. Positive values of E° indicate a greater tendency for reduction (the species is a stronger oxidising agent), while negative values indicate a tendency for oxidation. Standard electrode potentials are tabulated and used to predict the feasibility of redox reactions and to calculate cell EMFs.
Key Formula
E_cell = E_red(cathode) - E_red(anode)
LaTeX: E°_{cell} = E°_{reduction,\,cathode} - E°_{reduction,\,anode}
| Symbol | Meaning | Unit |
|---|---|---|
| E°_{cell} | Standard cell potential | V |
| E°_{reduction,\,cathode} | Standard reduction potential of cathode half-cell | V |
| E°_{reduction,\,anode} | Standard reduction potential of anode half-cell | V |
Worked Example
Problem
Use standard electrode potentials to determine whether the reaction Fe³⁺(aq) + Cu(s) → Fe²⁺(aq) + Cu²⁺(aq) is spontaneous under standard conditions. Given E°(Fe³⁺/Fe²⁺) = +0.77 V and E°(Cu²⁺/Cu) = +0.34 V.
Solution
Step 1: Identify the half-reactions. Cu is oxidised at the anode; Fe³⁺ is reduced at the cathode. Step 2: E°_cell = E°_cathode − E°_anode = 0.77 − 0.34 = +0.43 V Step 3: Since E°_cell > 0, ΔG° < 0, so the reaction is spontaneous under standard conditions.
Answer
E°_cell = +0.43 V; reaction is spontaneous
Selected Standard Reduction Potentials at 25 °C
| Half-reaction | E° (V) | Tendency |
|---|---|---|
| F₂ + 2e⁻ → 2F⁻ | +2.87 | Strong oxidising agent |
| MnO₄⁻ + 8H⁺ + 5e⁻ → Mn²⁺ + 4H₂O | +1.51 | Strong oxidising agent |
| Cu²⁺ + 2e⁻ → Cu | +0.34 | Moderate oxidising agent |
| 2H⁺ + 2e⁻ → H₂ | 0.00 | Reference (SHE) |
| Zn²⁺ + 2e⁻ → Zn | −0.76 | Strong reducing agent |
| Li⁺ + e⁻ → Li | −3.04 | Strongest common reducing agent |
Interactive Tools
NIST Electrochemical Data
Authoritative database of thermochemical and electrochemical data including electrode potentials
Open ToolKhan Academy – Standard Electrode Potentials
Concept lessons and practice on using reduction potential tables
Open ToolWolfram Alpha – Electrode Potential
Look up and compare standard electrode potentials for any redox couple
Open Tool
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Related Terms
Standard Hydrogen Electrode
The standard hydrogen electrode (SHE) is the primary reference electrode against which all other standard electrode potentials are measured, assigned an electrode potential of exactly 0.00 V by convention. It consists of a platinum electrode immersed in a 1 mol L⁻¹ solution of H⁺ ions (pH = 0) at 25 °C (298 K) with hydrogen gas at 1 atm bubbling over the platinum surface, establishing the equilibrium H⁺(aq) + e⁻ ⇌ ½H₂(g). Because it is experimentally difficult to set up, the SHE is often replaced in practice by secondary reference electrodes such as the saturated calomel electrode (SCE) or the Ag/AgCl electrode, which have known potentials relative to SHE.
Cell EMF
The electromotive force (EMF) of an electrochemical cell is the maximum potential difference between the two electrodes when no current is flowing (open-circuit condition), representing the driving force for electron transfer in the external circuit. It is determined by the difference in the electrode potentials of the cathode and anode and is directly related to the Gibbs free energy change of the reaction by ΔG° = −nFE°. A positive cell EMF indicates a spontaneous reaction (ΔG < 0), while a negative cell EMF indicates a non-spontaneous reaction under the given conditions.
Nernst Equation
The Nernst equation relates the electrode potential of a half-cell (or full cell) to the standard electrode potential and the reaction quotient Q, accounting for the actual concentrations or partial pressures of reactants and products at any temperature. It shows that the cell potential decreases as products accumulate and reactants are consumed, reaching zero at equilibrium when the cell is fully discharged. The equation is critical for predicting cell behaviour under non-standard conditions and forms the basis of pH measurement using electrochemical sensors.
The word "electrode" was coined by Michael Faraday from Greek "elektron" (amber/electricity) and "hodos" (path or way). "Standard" here refers to the IUPAC-defined reference conditions (25 °C, 1 mol L⁻¹, 1 atm).