Common Cell Presets
More positive = reduction occurs here
Enter cathode standard potential (V).
Standard reduction potential of anode half-cell
Enter anode standard potential (V).
Moles of electrons per mole of reaction
Enter n (1–8).
Cell EMF
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Sources & Methodology
Standard electrode potentials from NIST Electrochemical Series. Constants: F=96485.33212 C/mol, R=8.31446 J/mol/K (CODATA 2018).
NIST Electrochemical Series & Standard Potentials
Standard reduction potentials vs SHE at 25°C. nist.gov
IUPAC Recommendations — Electrochemical Conventions
Standard sign conventions for electrode potentials and cell notation. iupac.org
Atkins’ Physical Chemistry, 12th Ed. — Electrochemistry (Ch. 19)
Standard reference for Nernst equation, Gibbs-EMF relationship, and equilibrium electrochemistry.
E°_cell = E°_cathode − E°_anode
Nernst: E = E° − (RT/nF) ln(Q) = E° − (0.05916/n) log(Q) at 25°C
ΔG = −nFE | ΔG° = −nFE°
log(K) = nE° / 0.05916 (at 25°C)
Nernst: E = E° − (RT/nF) ln(Q) = E° − (0.05916/n) log(Q) at 25°C
ΔG = −nFE | ΔG° = −nFE°
log(K) = nE° / 0.05916 (at 25°C)
E_cell = E_cathode - E_anode | delta-G = -nFE
Daniell cell: E° = +0.34 − (−0.76) = +1.10 V
ΔG° = −2 × 96485 × 1.10 = −212.3 kJ/mol (spontaneous)
ΔG° = −2 × 96485 × 1.10 = −212.3 kJ/mol (spontaneous)
Electromotive Force and the Nernst Equation
Electromotive force (EMF) of a galvanic cell is the maximum potential difference between the electrodes when no current flows. It equals the standard cell potential E°_cell = E°_cathode − E°_anode, where both potentials are standard reduction potentials. The cathode (where reduction occurs) always has the more positive reduction potential in a spontaneous galvanic cell.
For non-standard conditions (concentrations other than 1 M, temperatures other than 25°C), the Nernst equation corrects the standard EMF: E = E° − (0.05916/n) × log(Q) at 25°C, where Q is the reaction quotient.
Selected Standard Electrode Potentials (vs SHE, 25°C)
| Half-Reaction | E° (V) | Notes |
|---|---|---|
| F₂ + 2e⁻ → 2F⁻ | +2.87 | Strongest oxidizing agent |
| MnO₄⁻ + 8H⁺ + 5e⁻ → Mn²⁺ | +1.51 | Permanganate reduction |
| Cl₂ + 2e⁻ → 2Cl⁻ | +1.36 | Chlorine gas |
| O₂ + 4H⁺ + 4e⁻ → 2H₂O | +1.23 | Oxygen in acid |
| Cu²⁺ + 2e⁻ → Cu | +0.34 | Copper deposition |
| 2H⁺ + 2e⁻ → H₂ | 0.000 | SHE reference |
| Zn²⁺ + 2e⁻ → Zn | −0.76 | Zinc oxidation in Daniell cell |
| Al³⁺ + 3e⁻ → Al | −1.66 | Aluminum metal |
| Na⁺ + e⁻ → Na | −2.71 | Alkali metal |
| Li⁺ + e⁻ → Li | −3.04 | Strongest reducing agent |
💡 Spontaneity Rules:
E_cell > 0 → Spontaneous (ΔG < 0, K > 1)
E_cell = 0 → Equilibrium (ΔG = 0, Q = K)
E_cell < 0 → Non-spontaneous (ΔG > 0, K < 1)
All three criteria are equivalent expressions of spontaneity.
E_cell > 0 → Spontaneous (ΔG < 0, K > 1)
E_cell = 0 → Equilibrium (ΔG = 0, Q = K)
E_cell < 0 → Non-spontaneous (ΔG > 0, K < 1)
All three criteria are equivalent expressions of spontaneity.
Frequently Asked Questions
EMF = maximum potential difference between galvanic cell electrodes at zero current. E°_cell = E°_cathode − E°_anode (using reduction potentials). Positive EMF = spontaneous reaction.
E_cell = E_cathode − E_anode (both as reduction potentials). Cathode has more positive E°. Daniell cell: E° = +0.34 − (−0.76) = +1.10 V.
E = E° − (RT/nF) ln(Q). At 25°C: E = E° − (0.05916/n) log(Q). Q = reaction quotient. As Q increases (products accumulate), E decreases until E=0 at equilibrium.
ΔG = −nFE. For standard conditions: ΔG° = −nFE°. Positive E → negative ΔG = spontaneous. This links electrochemistry directly to thermodynamics.
At equilibrium, E=0 and Q=K. From Nernst: E° = (0.05916/n) log(K) at 25°C. So log(K) = nE° / 0.05916. Daniell cell: log(K) = 2(1.10)/0.05916 = 37.2; K = 10³⁷·² ≈ 10³⁷.
Reduction potential under standard conditions (1 M, 1 atm, 25°C) vs standard hydrogen electrode (SHE = 0.000 V). More positive = stronger oxidizing agent. More negative = stronger reducing agent.
The SHE (Pt electrode, 1 M H⁺, 1 atm H₂) with E° = 0.000 V exactly by convention. All other standard potentials are measured relative to SHE.
A cell that converts spontaneous chemical energy into electrical energy. Anode (oxidation) and cathode (reduction) connected by salt bridge. Electrons flow from anode to cathode through external circuit.
When E_cell > 0 (positive EMF). Equivalent to ΔG < 0 and K > 1. All three are interlinked via ΔG = −nFE = −RT ln(K).
Zn | Zn²⁺ || Cu²⁺ | Cu. E° = +0.34 − (−0.76) = +1.10 V. ΔG° = −2(96485)(1.10) = −212 kJ/mol. K ≈ 10³⁷. First practical galvanic cell (Daniell, 1836).
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