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Cell EMF Calculator

Chemistry

Calculate standard cell potential E°cell = E°cathode − E°anode, Gibbs free energy ΔG°, and equilibrium constant K for any electrochemical cell.

0.34 V
V
-0.44 V
V
2

E°cell

0.78
ΔG°
-150.517
log₁₀(K)
26.371
Spontaneity
Spontaneous (ΔG° < 0)

This calculator computes your E°cell, ΔG°, log₁₀(K), Spontaneity from the values you enter.

Inputs
Cathode Reduction Potential (E°)Anode Reduction Potential (E°)Electrons Transferred (n)
Outputs
E°cellΔG°log₁₀(K)Spontaneity

What is a Cell EMF?

The Cell EMF Calculator computes the standard cell potential E°cell = E°cathode − E°anode for any galvanic or electrolytic cell, along with the Gibbs free energy ΔG° and equilibrium constant K. Enter the standard reduction potentials (from an electrochemical series table) for the two half-reactions and the number of electrons transferred.

Standard electrode potentials are the fundamental language of electrochemistry. Every galvanic cell — from a simple zinc-copper Daniell cell to a lithium-ion battery — can be characterised by E°cell, which encapsulates the entire thermodynamic driving force of the cell reaction. A positive E°cell means the cell does work spontaneously (like a battery); a negative E°cell means external work must be done to drive the reaction (electrolysis). The Nernst Equation Calculator extends this to non-standard conditions.

The three outputs — E°cell, ΔG°, and log₁₀(K) — are all equivalent representations of the same thermodynamic information. ΔG° = −nFE°cell tells you the maximum work the cell can do; log₁₀(K) = nE°cell/0.05916 tells you where the equilibrium lies. A cell voltage of just 0.1 V per electron transferred corresponds to K ≈ 10^1.7 ≈ 50 — products already heavily favoured. This calculator makes these conversions automatic.

How to use this Cell EMF calculator

  1. Look up the Cathode Reduction Potential (E°) from an electrochemical series table for the half-reaction that undergoes reduction. The species with the higher E° always acts as cathode.
  2. Look up the Anode Reduction Potential (E°) for the half-reaction that undergoes oxidation. Use the reduction potential value (not the oxidation potential) — the formula subtracts it automatically.
  3. Enter n, the number of moles of electrons transferred in the balanced overall equation. (Count electrons in the balanced half-reactions.)
  4. Read E°cell — if positive, the reaction is spontaneous.
  5. Use ΔG° and log₁₀(K) to quantify the thermodynamic driving force and equilibrium position.

Formula & Methodology

Standard cell potential:

E°cell = E°cathode − E°anode    (both as reduction potentials)

Gibbs free energy:

ΔG° = −nFE°cell F = 96,485 C/mol  (Faraday constant)

Equilibrium constant:

log₁₀(K) = nFE° / (R × T × ln10) = nE°cell / 0.05916   (at 25°C)

Worked example — Daniell cell (Cu²⁺/Cu || Zn²⁺/Zn):

Cathode: Cu²⁺ + 2e⁻ → Cu, E° = +0.34 V
Anode: Zn²⁺ + 2e⁻ → Zn, E° = −0.76 V (Zn is oxidised, but we use E°reduction = −0.76 V)
n = 2

E°cell = 0.34 − (−0.76) = 1.10 V ΔG° = −2 × 96485 × 1.10 / 1000 = −212.3 kJ/mol log₁₀(K) = 2 × 96485 × 1.10 / (8.314 × 298.15 × 2.3026) = 37.21 K = 10^37.21 ≈ 1.6 × 10³⁷

The Daniell cell delivers 1.10 V, releases 212 kJ per mole of reaction, and has an equilibrium constant so large that the reaction is effectively irreversible. This explains why zinc spontaneously corrodes in contact with copper sulfate solution.

Frequently Asked Questions

The standard cell potential E°cell is the voltage (electromotive force) generated by an electrochemical cell when all species are at standard conditions: 1 M concentration for solutes, 1 atm for gases, and 25°C temperature. It measures the thermodynamic driving force for the redox reaction. A positive E°cell means the reaction is spontaneous under standard conditions; a negative E°cell means it is non-spontaneous.
E°cell = E°cathode − E°anode, where both E°cathode and E°anode are the standard reduction potentials (tabulated in the electrochemical series) for the half-reactions as written in the reduction direction. The cathode is where reduction occurs (higher E°); the anode is where oxidation occurs (lower E°). The cell potential is always the more-positive electrode minus the less-positive electrode.
ΔG° = −nFE°cell (where n = moles of electrons transferred, F = 96,485 C/mol). K = e^(nFE°/RT) = 10^(nE°/0.05916) at 25°C. A positive E°cell gives negative ΔG° (spontaneous, K > 1). A cell potential of 0.0591/n volts per decade of K means a 1-electron cell at 1 V has K ≈ 10^17 — the reaction strongly favours products.
EMF (electromotive force) is the general term for the cell voltage, which depends on conditions. E°cell is specifically the EMF at standard conditions (1 M, 1 atm, 25°C). The actual cell voltage under non-standard conditions is calculated from E°cell using the Nernst equation: E = E°cell − (RT/nF)ln(Q). At standard conditions Q = 1, so E = E°cell. The [Nernst Equation Calculator](/nernst-equation-calculator/) computes the actual EMF at any conditions.
Enter the Standard Reduction Potential (E°) of the cathode half-reaction and the Standard Reduction Potential of the anode half-reaction — both in volts, as reduction potentials from the standard electrode potential table. Enter n, the number of electrons transferred in the balanced overall reaction. The calculator returns E°cell, ΔG°, log₁₀(K), and the spontaneity classification.
Selected E° values vs. SHE (Standard Hydrogen Electrode): F₂/F⁻: +2.87 V. MnO₄⁻/Mn²⁺: +1.51 V. Cl₂/Cl⁻: +1.36 V. O₂/H₂O: +1.23 V. Cu²⁺/Cu: +0.34 V. H⁺/H₂: 0.00 V (reference). Pb²⁺/Pb: −0.13 V. Fe²⁺/Fe: −0.44 V. Zn²⁺/Zn: −0.76 V. Al³⁺/Al: −1.66 V. Li⁺/Li: −3.04 V. More positive = stronger oxidising agent; more negative = stronger reducing agent.
The Daniell cell (zinc-copper galvanic cell) is the classic example of an electrochemical cell. Anode: Zn → Zn²⁺ + 2e⁻ (E° = −0.76 V reduction). Cathode: Cu²⁺ + 2e⁻ → Cu (E° = +0.34 V reduction). E°cell = 0.34 − (−0.76) = 1.10 V. n = 2. ΔG° = −2 × 96485 × 1.10 / 1000 = −212.3 kJ/mol. K = 10^(2 × 1.10/0.05916) = 10^37.2 ≈ 10^37. The reaction strongly favours products.
The Standard Hydrogen Electrode (SHE) is defined as having E° = 0.00 V exactly: 2H⁺(aq, 1M) + 2e⁻ → H₂(g, 1 atm). All standard reduction potentials are measured relative to the SHE. It cannot be set up absolutely; instead, a secondary reference electrode like the saturated calomel electrode (SCE, E = +0.241 V vs SHE) or the silver/silver chloride electrode (E = +0.197 V vs SHE) is used in practical laboratories. CBSE Class 12 (Chapter 3) and JEE cover SHE, galvanic cells, and the electrochemical series.
Feasibility rule: if E°cell > 0, the forward reaction is spontaneous under standard conditions. If E°cell < 0, the reaction requires electrical energy input (electrolysis). For example: can Fe³⁺ oxidise Cu? E°(Fe³⁺/Fe²⁺) = +0.77 V (cathode), E°(Cu²⁺/Cu) = +0.34 V (anode for reverse: Cu → Cu²⁺). E°cell = 0.77 − 0.34 = +0.43 V > 0 → yes, Fe³⁺ can oxidise Cu. This is the standard electrode potential table reading method taught in NCERT Class 12 electrochemistry.
Yes. Lead-acid batteries (E°cell ≈ 2.05 V per cell) power most vehicles including two-wheelers and cars in India. The lithium-ion batteries in electric vehicles (Tata Nexon EV, Ola S1) use LiCoO₂/graphite chemistry with E°cell ≈ 3.7 V per cell. Hydrogen fuel cells (E°cell = 1.23 V, O₂/H₂O vs H₂) are being developed for India's green hydrogen mission. DRDO and ISRO use specialised electrochemical cells for defence and space applications.