Cell Potential
E°cell = E°cathode − E°anode
Gibbs Free Energy
ΔG° = −nFE°cell
Nernst Equation
E = E° − (RT/nF) ln Q
Disclaimer
Results are based on standard conditions (25°C, 1 atm, 1 M concentrations) unless otherwise specified. Actual experimental values may vary.
| Half-Reaction | E° (V) |
|---|---|
| F₂ + 2e⁻ → 2F⁻ | +2.87 |
| Au³⁺ + 3e⁻ → Au | +1.50 |
| Cl₂ + 2e⁻ → 2Cl⁻ | +1.36 |
| O₂ + 4H⁺ + 4e⁻ → 2H₂O | +1.23 |
| Br₂ + 2e⁻ → 2Br⁻ | +1.07 |
| Ag⁺ + e⁻ → Ag | +0.80 |
| Fe³⁺ + e⁻ → Fe²⁺ | +0.77 |
| I₂ + 2e⁻ → 2I⁻ | +0.54 |
| Cu²⁺ + 2e⁻ → Cu | +0.34 |
| 2H⁺ + 2e⁻ → H₂ | +0.00 |
| Pb²⁺ + 2e⁻ → Pb | -0.13 |
| Sn²⁺ + 2e⁻ → Sn | -0.14 |
| Ni²⁺ + 2e⁻ → Ni | -0.26 |
| Fe²⁺ + 2e⁻ → Fe | -0.44 |
| Zn²⁺ + 2e⁻ → Zn | -0.76 |
| Al³⁺ + 3e⁻ → Al | -1.66 |
| Mg²⁺ + 2e⁻ → Mg | -2.37 |
| Na⁺ + e⁻ → Na | -2.71 |
| Ca²⁺ + 2e⁻ → Ca | -2.87 |
| K⁺ + e⁻ → K | -2.93 |
| Li⁺ + e⁻ → Li | -3.04 |
Electrochemistry is the branch of chemistry that studies the relationship between electrical energy and chemical reactions. It deals with redox (reduction-oxidation) reactions where electrons are transferred between species. These reactions form the basis of batteries, fuel cells, electrolysis, corrosion, and many industrial processes.
In an electrochemical cell, oxidation (loss of electrons) occurs at the anode, while reduction (gain of electrons) occurs at the cathode. The cell potential measures the driving force of the reaction, with positive values indicating a spontaneous reaction and negative values indicating a non-spontaneous reaction that requires external energy.
The standard cell potential (E°cell) is calculated by subtracting the anode's reduction potential from the cathode's reduction potential. A positive E°cell indicates that the reaction will proceed spontaneously under standard conditions. The more positive the cell potential, the greater the tendency for the reaction to occur.
Galvanic Cell (E° > 0)
Spontaneous reaction that converts chemical energy to electrical energy. Used in batteries.
Electrolytic Cell (E° < 0)
Non-spontaneous reaction that requires electrical energy input. Used in electroplating.
The Gibbs free energy change (ΔG°) provides a thermodynamic measure of whether a reaction is spontaneous. It is directly related to cell potential through the equation ΔG° = −nFE°, where n is the number of electrons transferred and F is the Faraday constant (96,485 C/mol).
ΔG° < 0 (Negative)
Reaction is spontaneous and releases energy. The more negative ΔG°, the more favorable the reaction.
ΔG° = 0
System is at equilibrium. No net reaction occurs in either direction.
ΔG° > 0 (Positive)
Reaction is non-spontaneous and requires energy input to proceed.
The Nernst equation allows calculation of cell potential under non-standard conditions. It accounts for the effect of temperature and concentration on cell potential. The equation is E = E° − (RT/nF) ln Q, where Q is the reaction quotient representing the ratio of product to reactant concentrations.
At 25°C, this simplifies to E = E° − (0.0592/n) log Q. As the reaction proceeds and Q increases, the cell potential decreases until equilibrium is reached (E = 0). This is why batteries lose voltage as they discharge—the concentrations of reactants and products change, affecting the cell potential.