Reduction Potential Calculator

Calculator Inputs

Enter standard reduction potentials as tabulated (V). For the anode, still enter its reduction potential; the calculator handles the subtraction.

Cathode E° (reduction) *

Typical oxidant half-reaction reduction potential (V).

Anode E° (reduction) *

Enter as reduction potential; anode oxidation is implied.

Electrons transferred, n *

Number of electrons balanced in the overall reaction.

Temperature (°C) *

Used in the Nernst term and K computation.

Use Nernst correction

Compute Ecell at non‑standard conditions

If unchecked, Ecell = E°cell and Q is ignored.

Q input mode

Use activities (≈ concentrations) for dilute solutions.

Reaction quotient Q (dimensionless)

Only used when Q mode is direct and Nernst is enabled.

Q Builder (optional)

Q = Π(a_products^ν) / Π(a_reactants^ν)

Enter up to three products and reactants. Leave unused rows blank. Use stoichiometric powers (ν) as positive coefficients.

Products

Example: for [Cu2+] with coefficient 1, a=0.01, ν=1.

Reactants

Example: for [Zn2+] coefficient 1, a=1, ν=1.

Quick guidance

Enable Nernst to use Q in Ecell.
Choose “Build Q” to use these fields.
Use activities when available; else use concentrations.
For gases, use partial pressures in atm as activity.
Keep Q dimensionless (use reference states).

Reset

Example Data Table

Sample inputs and outputs for quick verification. Values are illustrative and may not match your exact system.

Case	E°cath (V)	E°an (V)	n	T (°C)	Q	E°cell (V)	Ecell (V)	ΔG (kJ/mol)
Standard-like	1.229	0.000	2	25	1	1.229	1.229	-237.2
Non‑standard	0.340	-0.763	2	25	0.010	1.103	1.163	-224.3
Near equilibrium	0.100	0.090	1	25	10	0.010	-0.049	4.7

Formulas Used

Standard cell potential
E°cell = E°cathode − E°anode
Both terms are entered as reduction potentials from tables.

Nernst equation (non‑standard)
Ecell = E°cell − (RT/nF) ln(Q)
R = 8.314 J/mol·K, F = 96485 C/mol, T in Kelvin.

Gibbs free energy change
ΔG = −n F Ecell
Negative ΔG indicates thermodynamic favorability as written.

Equilibrium constant
K = exp(n F E°cell / RT)
Computed at the entered temperature using E°cell.

How to Use This Calculator

Find the two half-reactions and their tabulated E° values (reduction form).
Set the stronger oxidant as the cathode (higher E°) and the other as anode.
Balance electrons to determine n for the overall reaction.
For non‑standard conditions, enable Nernst and enter Q directly or build it.
Press compute to see E°cell, Ecell, ΔG, and K.
Download CSV/PDF to document design reviews and scenario comparisons.

Engineering caution: A positive Ecell indicates thermodynamic driving force, not reaction rate. Kinetics, mass transfer, passivation, and overpotential can dominate real performance.

FAQs

1) What is “reduction potential” in practice?

It is the tendency of a species to gain electrons. Higher reduction potential means a stronger oxidizing agent under the stated conditions.

2) Why do I enter the anode value as a reduction potential?

Tables list standard potentials as reductions. The cell voltage uses subtraction, so using two reduction values avoids sign mistakes.

3) When should I enable the Nernst correction?

Enable it when concentrations, activities, or partial pressures differ from standard states. It adjusts Ecell to match your operating conditions.

4) What is Q and how accurate is the “Q builder”?

Q is the reaction quotient from activities of products and reactants. The builder assumes ideal behavior; for concentrated electrolytes, use measured activities or activity coefficients.

5) How do I interpret ΔG from the output?

ΔG is the electrical work per mole of reaction. Negative ΔG suggests the reaction can proceed as written, but real systems may require overpotential.

6) What does the equilibrium constant K tell me?

K estimates how far a reaction favors products at equilibrium. Very large K implies product-favored equilibrium, though kinetics may still be slow.

7) Why can a positive Ecell still fail in experiments?

Thermodynamics does not guarantee speed. Surface films, sluggish charge transfer, limited mass transport, and competing side reactions can stop observable conversion.