Electrochemistry · equilibrium potentials
From chemical driving force
to cell potential.
Build a cell from two balanced reduction half-reactions. Compare standard and actual potentials, connect voltage to Gibbs energy and K, and explore a concentration cell.
Cell potential and driving force
Activities by compartment
Teal: Nernst relation at fixed T and E°. Amber: supplied composition. Graphite: Q=K, E=0. The sweep varies the aggregate reaction quotient; it is not a solved electroneutral composition path.
Inspect curve data
Electron balance and reaction basis
Both input half-reactions consume electrons. The app checks every declared element and charge including electrons, then uses the least common multiple of the two electron counts. It reverses the left half-reaction and adds the right. E°cell=E°right−E°left: electrode potentials are never multiplied by the stoichiometric scale factors.
Scaling the entire reaction scales n, ΔrG° and ln K together, while leaving E unchanged. Identical species in different compartments retain their labels and activities. Do not cancel ions across a concentration cell.
Atom counts are explicit user declarations, not inferred from the species label. These checks establish bookkeeping, not the existence or stability of the supplied phases.
Activities, temperature and applicability
Both standard reduction potentials must use the same reference electrode, temperature and compatible activity standard states. Solute inputs use a declared molarity standard; molality-based coefficients must first be converted consistently. Individual-ion activities are convention dependent. Counterions and solvent may be omitted from the reaction table, but the physical bulk solutions must be electroneutral.
Pure solid and liquid activities are one only when those phases are present. Gas activity is fugacity divided by its standard pressure, supplied directly. No activity-coefficient model, electrolyte speciation, liquid-junction correction, phase-stability calculation or electrode kinetics is solved here. Liquid junction potential is assumed zero. No current or loaded-cell voltage is predicted.
At equilibrium for the complete cell reaction Q=K and E=0. Each reversible electrode may individually be at local interfacial equilibrium even while the cell has a nonzero open-circuit potential. Concentration cells can have E°=0 and K=1 with nonzero E.
Sources and reproducibility
IUPAC: standard electromotive force · IUPAC: Nernst equation · DeVoe: Nernst equation and junction assumption
Calculation source · Reproduce PDF and 600 dpi PNG. The JSON contains all input tables, standard-state notes, results and curve points. Data stay in this browser until downloaded; edits invalidate previous results.