Electrode Potential Calculators

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Electrode potential is the voltage developed at the interface between an electrode and its electrolyte solution, resulting from the tendency of a metal to lose electrons (oxidation) or a solution to donate electrons to an electrode (reduction). Standard electrode potential (E°) is measured at 25°C, 1 M concentration, and 1 atm, relative to the standard hydrogen electrode (SHE), which is assigned E° = 0.00 V. Standard cell potential: E°_cell = E°_cathode − E°_anode. The Nernst equation corrects for non-standard concentrations: E = E° − (RT/nF) × ln(Q). Biological redox couples drive the electron transport chain.

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Standard Reduction Potential Table (Selected)

  • Li⁺/Li: −3.04 V (strong reducing agent)
  • Zn²⁺/Zn: −0.76 V
  • Fe²⁺/Fe: −0.44 V
  • H⁺/H₂: 0.00 V (SHE reference)
  • NAD⁺/NADH: −0.32 V (biological)
  • Cu²⁺/Cu: +0.34 V
  • O₂/H₂O: +1.23 V (strongest common oxidizing agent)
  • F₂/F⁻: +2.87 V

Cell Potential

E°_cell = E°_cathode − E°_anode

Spontaneous when E°_cell > 0. Zn-Cu cell: E° = +0.34 − (−0.76) = +1.10 V.

Nernst Equation

E = E° − (RT/nF)ln(Q). At 25°C: E = E° − (0.0592/n)log(Q). Equilibrium: E = 0 → log(Keq) = nE°/0.0592. ΔG = −nFE.

Biological Electron Transport

ETC drives electrons from NADH (E°'= −0.32 V) to O₂/H₂O (+0.82 V). ΔE°' = 1.14 V → ΔG°' = −nFΔE°' = −220 kJ per NADH.

Glossary

Standard Electrode Potential (E°)
Voltage of a half-cell at 25°C, 1 M, 1 atm relative to SHE (= 0 V); more positive = stronger oxidizing agent; E°_cell = E°_cathode − E°_anode; spontaneous when E°_cell > 0.
Nernst Equation
E = E° − (RT/nF)ln(Q) = E° − (0.0592/n)log(Q) at 25°C; corrects for non-standard concentrations; E = 0 at equilibrium → log(Keq) = nE°/0.0592.
Biological Redox Potential
Standard reduction potentials at pH 7 (E°'); NADH/NAD⁺ = −0.32 V; O₂/H₂O = +0.82 V; ΔE°' = 1.14 V per NADH; drives ETC and ATP synthesis; ΔG°' = −nFΔE°' = −220 kJ/mol NADH.

Frequently Asked Questions

Standard electrode potential (E°) = the voltage of a half-cell at standard conditions (25°C, 1 M concentration, 1 atm) relative to the standard hydrogen electrode (SHE = 0.00 V). Measurement: set up a half-cell (e.g., Cu²⁺ + 2e⁻ → Cu electrode); connect via a salt bridge to the SHE; measure the voltage. Cu²⁺/Cu: E° = +0.34 V (Cu²⁺ tends to be reduced — copper is a better oxidizing agent than H⁺). Zn²⁺/Zn: E° = −0.76 V (Zn tends to be oxidized — zinc is a better reducing agent than H₂). The more positive E°, the stronger the oxidizing agent (tendency to accept electrons). The more negative E°, the stronger the reducing agent (tendency to donate electrons).

E°_cell = E°_cathode − E°_anode. Both values are taken from the standard reduction potential table. The half-reaction with higher E° is the cathode (reduction). Lower E° is the anode (oxidation). Example: Zn-Cu cell. E°(Cu²⁺/Cu) = +0.34 V; E°(Zn²⁺/Zn) = −0.76 V. Cu²⁺ has higher E° → cathode: Cu²⁺ + 2e⁻ → Cu. Zn → anode: Zn → Zn²⁺ + 2e⁻. E°_cell = +0.34 − (−0.76) = +1.10 V. E°_cell > 0 → spontaneous reaction → ΔG° = −nFE°_cell = −2 × 96,485 × 1.10 = −212,267 J/mol = −212.3 kJ/mol.

The Nernst equation: E = E° − (RT/nF) × ln(Q). At 25°C: E = E° − (0.0592/n) × log(Q). Q = reaction quotient = [products]^coeff / [reactants]^coeff. When Q < 1 (more reactants than products): ln(Q) < 0 → E > E° (reaction is more spontaneous than standard). When Q > 1 (more products): E < E°. At equilibrium: E = 0; Q = Keq: 0 = E° − (0.0592/n)log(Keq) → log(Keq) = nE°/0.0592. pH meter application: for a pH-sensitive electrode, E depends on [H⁺] concentration: E = constant + (0.0592/1) × log[H⁺] = constant − 0.0592 × pH → electrode potential changes 59.2 mV per pH unit (the Nernst slope).

In cellular respiration, electrons flow 'downhill' through increasingly positive standard reduction potentials: NADH (E°' = −0.32 V) → Ubiquinone (E°' = +0.045 V) → Cytochrome c (E°' = +0.23 V) → O₂ (E°' = +0.82 V). Each step releases free energy: ΔG° = −nFΔE°'. NAD⁺/NADH → O₂: ΔE°' = 0.82 − (−0.32) = +1.14 V; n = 2 electrons. ΔG° = −2 × 96,485 × 1.14 = −220 kJ per mole of NADH. This energy is captured as the proton motive force (ΔμH⁺) across the inner mitochondrial membrane → ATP synthesis by ATP synthase → ~2.5 ATP per NADH. The ETC is essentially a biological electrochemical cell, with NADH as the fuel and O₂ as the terminal electron acceptor.