Electrochemical Cell Voltage Calculator

Electrochemical Cells Guide

What do I need to know about Cell Potential Calculation?

E°cell = E°cathode − E°anode. The cathode is the electrode where reduction occurs (gains electrons); the anode is where oxidation occurs (loses electrons). By convention, standard electrode potentials are all written as reduction potentials. The more positive the reduction potential, the stronger the oxidising agent. Example: zinc-copper cell. Cathode (Cu²⁺/Cu): +0.34 V. Anode (Zn²⁺/Zn): −0.76 V. E°cell = 0.34 − (−0.76) = +1.10 V. A positive E°cell means the reaction is spontaneous in the forward direction as written, releasing energy that can do electrical work — a negative E°cell would mean the reverse reaction is spontaneous instead, and the cell as drawn would not function as a spontaneous battery.

What should I know about Gibbs Free Energy and Spontaneity?

ΔG° = −nFE°cell. Where n = moles of electrons transferred, F = Faraday's constant (96,485 C/mol), E°cell = standard cell potential. If E°cell > 0: ΔG° < 0 → spontaneous. If E°cell < 0: ΔG° > 0 → non-spontaneous (would require external energy — electrolysis). If E°cell = 0: ΔG° = 0 → equilibrium. For the zinc-copper cell (n=2): ΔG° = −2 × 96485 × 1.10 = −212 kJ/mol. Relationship to equilibrium: ΔG° = −RT ln K, so ln K = nFE°/RT = nE°/0.02569 at 298 K.

What do I need to know about The Electrochemical Series?

The electrochemical series lists half-cells in order of standard electrode potential, from most negative (strongest reducing agents: Li, K, Na, Mg, Al) to most positive (strongest oxidising agents: F₂, MnO₄⁻, O₂, Cl₂). Rules: a more positive half-cell always oxidises a more negative one. Displacement reactions: zinc displaces copper from copper sulfate solution because Zn has a more negative E° (−0.76 V) than Cu (+0.34 V). Electroplating: the metal to be deposited is the cathode in an electrolytic cell — for example, electroplating an object with silver connects it as the cathode in a cell containing a silver salt solution, so silver ions are reduced and deposited onto its surface as a coherent metal layer.

What's the difference between Real Cells and Standard Conditions?

Standard electrode potentials apply at 298 K, 1 mol/L concentrations, and 1 atm for gases. Real cells deviate due to concentration effects (Nernst equation): E = E° − (RT/nF) × ln Q. At 298 K: E = E° − (0.0257/n) × ln Q. This is important for: batteries as they discharge (reactant concentrations decrease), pH cells (glass electrode measures H⁺ concentration), and biological systems where concentrations are far from 1 mol/L. The Daniell cell (Zn/Cu) in everyday conditions operates at approximately 1.1 V under standard conditions, but the actual voltage falls as the cell discharges and reactant concentrations drop, following the Nernst equation rather than staying fixed at the standard value — this is why a battery's voltage sags gradually over its working life rather than switching off abruptly at zero.