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Electromotive Force (EMF) Calculator

Ready to calculate
EMF = E°_c − E°_a.
ΔG = −nFE.
23 Reference Half-Cells.
100% Free.
No Data Stored.

How it Works

01Find Reduction Potentials

Look up E° (vs SHE) for each half-cell in a standard table — both values must be reduction potentials

02Identify Anode and Cathode

Cathode = species being reduced (higher E°); anode = species being oxidized (lower E°). Don't flip signs!

03Apply EMF = E°_c − E°_a

Subtract anode reduction potential from cathode reduction potential. The formula handles the sign change

04Read Cell Type & ΔG

EMF > 0 = galvanic (battery); EMF < 0 = electrolytic (forced). Add n electrons → ΔG = −nFE in kJ/mol

What is an Electromotive Force (EMF) Calculator?

Electromotive force (EMF), also called the standard cell potential E°cell, is the central quantity in electrochemistry — it tells you exactly how much voltage a galvanic cell can deliver under standard conditions, and equivalently, how much voltage you need to apply to drive an electrolysis in the reverse direction. The formula is the foundation of every electrochemistry course: EMF = E°cathode − E°anode, where both potentials are entered as standard reduction potentials taken directly from CRC or Bard tables (no sign-flipping needed — the subtraction handles the algebra). Our Electromotive Force Calculator implements this formula with full unit flexibility (mV, V, kV, MV for inputs and output), 6-band strength classification (non-spontaneous through extreme), an optional Gibbs free-energy calculation ΔG = −nFE if you provide the electrons transferred, and a 23-entry reference table of standard reduction potentials spanning the full electrochemistry range.

Just enter the standard reduction potential E° of the anode (the electrode where oxidation occurs) and the cathode (the electrode where reduction occurs), in any voltage unit. The calculator subtracts and reports EMF in your chosen output unit. The sign tells the cell type: positive EMF means a spontaneous galvanic cell — current flows from anode to cathode through the external circuit, electrons go from anode to cathode through the wire, and the cell can power a device; negative EMF means the reaction is non-spontaneous — to make it proceed in this direction you must apply external voltage greater than |EMF|, which is electrolysis. The 6-band classification (very-weak / weak / moderate / strong / extreme for galvanic; non-spontaneous for negative) gives instant context for whether your cell is suitable for a particular application.

Designed for general chemistry students learning electrochemistry, electrochemists characterizing new battery chemistries, materials scientists evaluating corrosion potentials, environmental scientists modeling redox reactions in groundwater, fuel-cell engineers calculating theoretical maximum voltages for hydrogen and methanol cells, and physical chemistry students preparing for the GRE or qualifying exams, the tool runs entirely in your browser — no data is stored or transmitted.

Pro Tip: Pair this with our Nernst Equation Calculator to compute cell potential under non-standard conditions (concentrations ≠ 1 M, T ≠ 25 °C), or our Gibbs Free Energy Calculator for the related ΔG = ΔH − T·ΔS thermochemistry.

How to Use the EMF Calculator?

Look Up Reduction Potentials: Find E° (standard reduction potential, vs SHE) for both half-cells in a CRC handbook or your textbook. Always use REDUCTION potentials — never flip the sign of the anode. The 23-entry reference table on this page covers most common couples.
Identify Anode vs Cathode: The cathode is where REDUCTION occurs (gain of electrons) — usually the species with the higher E°. The anode is where OXIDATION occurs (loss of electrons) — the lower E°. For a Daniell cell: Cu²⁺/Cu (E° = +0.34) is the cathode; Zn²⁺/Zn (E° = −0.76) is the anode.
Enter Both Potentials: Input the E° values with any voltage unit (mV, V, kV, MV). The calculator handles unit conversion to SI volts internally.
Pick Output Voltage Unit: Display the EMF in mV, V, kV, or MV — separate from input units, so you can enter in mV and see V output.
Optional — Enter n (Electrons Transferred): Number of moles of electrons per mole of cell reaction. Daniell cell (Cu²⁺ + Zn → Cu + Zn²⁺): n = 2. Hydrogen fuel cell (H₂ + ½O₂ → H₂O): n = 2. Required for ΔG = −nFE calculation.
Press Calculate: Get EMF (in your chosen unit), cell type (galvanic / electrolytic / equilibrium), 6-band strength classification, optional ΔG in kJ/mol, and full step-by-step breakdown.

How is EMF calculated?

EMF is the most direct experimental measure of redox spontaneity — a single number that tells you whether two half-cells will combine to drive current spontaneously, or instead require external voltage to be forced. Here's the complete framework:

Walther Nernst formalized the electrochemistry framework in the 1890s, work that earned him the 1920 Nobel Prize. The standard cell potential E°_cell ties cell EMF to the underlying half-cell reduction potentials via a simple sign convention.

The Formula

For any electrochemical cell at standard conditions (25 °C, 1 M, 1 atm):

EMF = E°cell = E°cathode − E°anode

where both E° values are STANDARD REDUCTION POTENTIALS (taken from tables as written, with reduction direction). The subtraction handles the sign conversion — the anode physically undergoes oxidation, but you don't manually flip its E° sign in this formula.

Why "Cathode − Anode" Instead of "Anode + Cathode"

Because both E° values are written in the REDUCTION direction. To get the cell EMF, you need:

E°cell = E°reduction(cathode) + E°oxidation(anode) = E°reduction(cathode) + (−E°reduction(anode)) = E°cathode − E°anode

The minus sign comes from converting the anode's reduction potential into its oxidation potential. The compact form (cathode − anode) is what every electrochemistry textbook uses.

Sign Convention and Cell Types

  • EMF > 0: Galvanic (voltaic) cell — spontaneous reaction, generates current. The cell can do work on external circuit. All batteries operate this way during discharge.
  • EMF = 0: Cell is at equilibrium — no net reaction. A "discharged" battery has reached this state.
  • EMF < 0: Electrolytic — the reaction is non-spontaneous in this direction. To make it proceed, you must apply external voltage > |EMF|. Electrolysis (water → H₂ + O₂, electroplating, aluminum production) all run with negative EMF as written.

Connection to Gibbs Free Energy

EMF is connected to thermodynamic spontaneity by:

ΔG = −nFE

where n is the moles of electrons transferred per mole of cell reaction, F = 96,485 C/mol is Faraday's constant, and E is the cell EMF in volts. Result is in joules per mole. Each volt of EMF corresponds to ΔG = −96.5 kJ/mol per mole of electrons. For n = 2 (typical for many cells like Daniell): ΔG = −193 kJ/mol per volt.

Connection to the Equilibrium Constant

At equilibrium, EMF = 0 and ΔG = 0. Combining ΔG° = −nFE° with ΔG° = −RT·ln(K) gives:

E°cell = (RT/nF) · ln(K) = (0.0257/n) · ln(K) at 25 °C

Or equivalently, K = exp(nFE°/RT). For a Daniell cell with E° = 1.10 V and n = 2: K = exp(2 × 96485 × 1.10 / (8.314 × 298.15)) ≈ 10³⁷ — staggeringly large, confirming that the reaction goes essentially to completion.

The Nernst Equation (Non-Standard Conditions)

EMF as defined here applies at standard conditions. For real cells with non-1-M concentrations, use the Nernst equation:

Ecell = E°cell − (RT/nF) · ln(Q) = E°cell − (0.0592/n) · log10(Q) at 25 °C

where Q is the reaction quotient (products/reactants). Doubling [Cu²⁺] in a Daniell cell shifts EMF by (0.0592/2) × log(2) ≈ 9 mV — small but measurable.

When EMF Differs from Cell Voltage

EMF is the open-circuit voltage (no current flowing). When current is drawn, the actual terminal voltage drops due to: (1) internal resistance — V = EMF − I·R_int; (2) activation overpotential at electrodes; (3) concentration overpotential from depletion at electrode surfaces. A real lithium-ion cell with EMF 4.2 V might deliver only 3.7 V under typical load.

Real-World Example

EMF Calculator – Worked Examples

Example 1 — The Classic Daniell Cell. Zinc + copper(II) → copper + zinc(II). Half-cells:
  • Cathode (reduction): Cu²⁺ + 2e⁻ → Cu(s), E° = +0.34 V.
  • Anode (oxidation, but reduction E° entered): Zn²⁺ + 2e⁻ → Zn(s), E° = −0.76 V.
  • EMF = E°cathode − E°anode = 0.34 − (−0.76) = +1.10 V.
  • Positive → galvanic cell, spontaneous. The classroom Daniell cell delivers ~1.10 V open-circuit.
  • ΔG with n = 2: ΔG = −2 × 96485 × 1.10 = −212,267 J/mol = −212.3 kJ/mol. Strongly spontaneous.
  • K (equilibrium): K = exp(2 × 96485 × 1.10 / (8.314 × 298.15)) ≈ 1.4 × 10³⁷ — reaction goes essentially to completion.

Example 2 — Hydrogen Fuel Cell. H₂ + ½O₂ → H₂O.

  • Cathode: O₂ + 4H⁺ + 4e⁻ → 2H₂O, E° = +1.23 V.
  • Anode: 2H⁺ + 2e⁻ → H₂, E° = 0 V (SHE reference).
  • EMF = 1.23 − 0 = +1.23 V. Theoretical maximum voltage of a hydrogen-oxygen fuel cell.
  • Real PEM fuel cells deliver only ~0.7 V under typical load due to overpotentials. The 0.5 V "lost" appears as waste heat.
  • ΔG = −2 × 96485 × 1.23 = −237,353 J/mol = −237.4 kJ/mol per mol H₂ — closely matches the standard ΔG° of formation of liquid water.

Example 3 — Lithium-Ion Cell. A simplified analysis: Li → Li⁺ + e⁻ (anode); CoO₂ + Li⁺ + e⁻ → LiCoO₂ (cathode).

  • Cathode (reduction): Co⁴⁺ + e⁻ → Co³⁺ in LiCoO₂ matrix, effective E° ≈ +0.86 V vs SHE (corresponding to ~3.9 V vs Li/Li⁺).
  • Anode (reduction E°): Li⁺ + e⁻ → Li(s), E° = −3.04 V vs SHE.
  • EMF = 0.86 − (−3.04) = +3.90 V. Matches the typical lithium-ion cell voltage (3.7-4.2 V depending on state of charge).
  • Strong cell — band classification: "Strong Cell". This high voltage is why lithium-ion batteries dominate consumer electronics: more energy per cell than alkaline (1.5 V) or lead-acid (2.05 V/cell).

Example 4 — Electrolysis of Water (Negative EMF). 2H₂O → 2H₂ + O₂. To run as a cell:

  • Cathode (the species we'd reduce): 2H⁺ + 2e⁻ → H₂, E° = 0 V.
  • Anode (the species we'd oxidize): O₂ + 4H⁺ + 4e⁻ → 2H₂O, E° = +1.23 V (entered as reduction).
  • EMF = 0 − 1.23 = −1.23 V. Negative → non-spontaneous as written. To electrolyze water, apply at least 1.23 V (in practice ~1.7-2 V due to overpotentials at typical electrodes).
  • This is exactly why hydrogen production via water electrolysis requires ~50 kWh/kg H₂ of electricity — the reverse of the fuel-cell reaction.

Example 5 — Aluminum Anodic Protection (Galvanic Anode). Sacrificial Mg anode protecting steel pipeline.

  • Cathode (steel surface): Fe²⁺ + 2e⁻ → Fe, E° = −0.44 V (the rusting reaction we want to prevent).
  • Anode (Mg block): Mg²⁺ + 2e⁻ → Mg, E° = −2.37 V.
  • EMF = −0.44 − (−2.37) = +1.93 V. Strong cell forms with Mg as anode, sacrificially corroding to protect the steel.
  • Mg is consumed over months/years; periodic replacement extends pipeline life by decades. Same chemistry protects boat hulls and water-heater tanks.

Who Should Use the EMF Calculator?

1
General Chemistry Students: Solve textbook EMF problems, predict cell directions, compute equilibrium constants for redox reactions via E° = (RT/nF)·ln(K).
2
Battery Chemists: Estimate theoretical maximum voltage for new battery chemistries; compare lead-acid (2.05 V), nickel-cadmium (1.35 V), lithium-ion (3.7 V), and lithium-air (3.0 V) cells.
3
Fuel-Cell Engineers: Calculate maximum theoretical efficiency from EMF / ΔH ratio; identify overpotential losses by comparing measured voltage to EMF.
4
Corrosion Engineers: Design sacrificial anode systems (Mg, Zn) to protect steel infrastructure; predict galvanic corrosion rates from EMF differences between metal pairs.
5
Environmental Chemists: Model redox transformations in groundwater (Fe²⁺/Fe³⁺, Mn²⁺/MnO₂, NO₃⁻/N₂) using the same EMF framework with pE values (= E/0.0592).
6
Electroplating / Electrowinning: Determine minimum applied voltage for industrial electrolysis (Al production, Cu refining, chlor-alkali process); compute energy requirements per kg product.

Technical Reference

Founding Work. Walther Nernst formalized electrochemistry in the 1890s, building on earlier work by Volta (the first battery, 1800), Daniell (the Daniell cell, 1836), and Faraday (the laws of electrolysis, 1834). Nernst's 1889 paper "Die elektromotorische Wirksamkeit der Ionen" derived the equation now bearing his name and earned him the 1920 Nobel Prize in Chemistry. The standard hydrogen electrode (SHE) was adopted as the universal reference (E° = 0 by definition) at the 1900 IUPAC convention.

The Standard Hydrogen Electrode (SHE). By international convention, all standard reduction potentials are tabulated relative to: 2H⁺(aq, 1 M) + 2e⁻ → H₂(g, 1 atm), E° = 0.000 V exactly. This is a definition, not a measurement. All other E° values are differences relative to this reference. Practical reference electrodes (silver-silver chloride, calomel) are used in real measurements, with known offsets to SHE.

Key Constants:

  • Faraday's constant F: 96,485.33212 C/mol — the charge of one mole of electrons. Named after Michael Faraday.
  • Universal gas constant R: 8.314 J/(mol·K).
  • RT/F at 25 °C: 0.02569 V (the "Nernst voltage" — the natural unit of voltage in electrochemistry).
  • (RT/F) × ln(10): 0.0592 V at 25 °C (the slope of the Nernst equation when using log₁₀).

Standard Reduction Potentials (E° at 25 °C, 1 M, vs SHE):

  • Strongest oxidizers (E° > +1.5 V): F₂ (+2.87), O₃ (+2.07), MnO₄⁻/Mn²⁺ (+1.51 acidic), Cr₂O₇²⁻/Cr³⁺ (+1.33 acidic), Cl₂ (+1.36)
  • Moderate oxidizers (E° = 0 to +1.5 V): O₂/H₂O (+1.23), Br₂ (+1.07), Ag⁺/Ag (+0.80), Fe³⁺/Fe²⁺ (+0.77), I₂ (+0.54), Cu²⁺/Cu (+0.34), 2H⁺/H₂ (0.00 reference)
  • Moderate reducers (E° = 0 to −1 V): Pb²⁺/Pb (−0.13), Sn²⁺/Sn (−0.14), Ni²⁺/Ni (−0.25), Fe²⁺/Fe (−0.44), Zn²⁺/Zn (−0.76)
  • Strongest reducers (E° < −1 V): Al³⁺/Al (−1.66), Mg²⁺/Mg (−2.37), Na⁺/Na (−2.71), Li⁺/Li (−3.04 — most negative)

Common Battery Cell EMFs:

  • Alkaline AA cell (Zn/MnO₂): 1.5 V
  • Carbon-zinc dry cell: 1.5 V
  • Lead-acid cell (per cell, 6 cells = 12V battery): 2.05 V
  • Nickel-cadmium (NiCd): 1.20 V
  • Nickel-metal hydride (NiMH): 1.20 V
  • Lithium-ion (LiCoO₂): 3.7 V (3.0-4.2 V range)
  • Lithium-iron-phosphate (LiFePO₄): 3.2 V
  • Lithium primary (Li-MnO₂): 3.0 V
  • Hydrogen fuel cell (H₂/O₂): 1.23 V (theoretical), ~0.7 V actual
  • Methanol fuel cell: 1.21 V (theoretical), ~0.5 V actual
  • Lithium-air (Li/O₂): 3.0 V theoretical — promises ~10× lithium-ion energy density

Why Real Cell Voltage Differs from EMF. Three loss mechanisms reduce the actual terminal voltage below EMF when current flows: (1) internal resistance (IR drop, ohmic loss); (2) activation overpotential at electrodes (kinetic barrier to electron transfer); (3) concentration overpotential at electrode surfaces (mass-transport limitation). For a Li-ion cell with EMF 4.2 V at full charge, terminal voltage at typical 1C discharge is ~3.7 V — about 12% lost to overpotentials and IR.

Concentration Cells. A special case: same redox couple but different concentrations on the two sides. EMF comes from the concentration difference alone, no chemistry change. Example: Cu | Cu²⁺(1 M) || Cu²⁺(0.001 M) | Cu has EMF = (0.0592/2) × log(1/0.001) = 0.089 V. Concentration cells are weak (millivolt range) but useful for sensors (ion-selective electrodes, pH meters).

Key Takeaways

Electromotive force EMF = E°cathode − E°anode is the master equation of electrochemistry: enter both half-cell standard reduction potentials directly from tables (no sign-flipping for the anode — the formula handles it), and the result tells you whether the cell will spontaneously generate current (positive EMF, galvanic) or require external voltage to drive (negative EMF, electrolytic). The Daniell cell gives +1.10 V; the hydrogen fuel cell gives +1.23 V; modern lithium-ion gives ~3.9 V. ΔG = −nFE connects EMF to thermodynamic spontaneity — every volt of EMF corresponds to ~96.5 kJ/mol per mole of electrons. Use the ToolsACE EMF Calculator with full unit flexibility (mV, V, kV, MV), 6-band strength classification, optional ΔG calculation, and a 23-entry reference table covering all common half-cells from F₂ (+2.87 V) to Li⁺/Li (−3.04 V). Bookmark it for chemistry coursework, battery chemistry, corrosion engineering, fuel-cell analysis, and any redox-spontaneity question.

Frequently Asked Questions

What is the Electromotive Force (EMF) Calculator?
It computes the standard cell potential EMF = E°cathode − E°anode for any electrochemical cell, given the standard reduction potentials of the two half-cells. Inputs: anode and cathode reduction potentials in mV, V, kV, or MV; optional moles of electrons transferred (n) for ΔG = −nFE calculation. Output: EMF in any voltage unit; cell type (galvanic / electrolytic / equilibrium); 6-band strength classification (very-weak through extreme); optional ΔG in kJ/mol; full step-by-step breakdown; 23-entry reference table of standard reduction potentials.

Designed for general chemistry students learning electrochemistry, battery chemists comparing cell chemistries, fuel-cell engineers, corrosion engineers designing sacrificial anode systems, environmental chemists modeling redox in natural waters, and electroplating / electrowinning engineers. Runs entirely in your browser — no data stored.

Pro Tip: Use our Nernst Equation Calculator for non-standard conditions (concentrations ≠ 1 M, T ≠ 25 °C).

What's the EMF formula?
EMF = E°cell = E°cathode − E°anode, where both E° values are STANDARD REDUCTION POTENTIALS (taken from tables exactly as written, with reduction direction). The subtraction handles the sign conversion automatically — you don't manually flip the anode's E° sign. Positive EMF = spontaneous galvanic cell; negative EMF = non-spontaneous (electrolysis required).
Why don't I flip the sign of the anode potential?
Because the formula E°_cathode − E°_anode already does it for you. The cathode undergoes reduction (use its E° as written); the anode physically undergoes oxidation, which is the reverse of reduction, so its contribution should be −E°_anode (with E°_anode being the reduction potential). Combining: E°_cell = E°_cathode + (−E°_anode) = E°_cathode − E°_anode. The minus sign in front of the anode term accomplishes the sign flip without you having to do it manually. Use both values straight from the reduction-potential table.
What does positive vs negative EMF mean?
EMF > 0 (positive): Galvanic (voltaic) cell — spontaneous reaction, generates current. The cell can do work on an external circuit. All batteries operate this way during discharge. EMF = 0: Cell at equilibrium, no net reaction (a fully discharged battery). EMF < 0 (negative): Electrolytic — non-spontaneous direction. To make the reaction proceed forward, apply external voltage greater than |EMF|. Reverse the half-cells (swap anode/cathode roles) and EMF becomes positive — that's the natural galvanic direction.
How is EMF related to Gibbs free energy?
ΔG = −nFE, where n is the moles of electrons transferred per mole of cell reaction, F = 96,485 C/mol is Faraday's constant, and E is the cell EMF in volts. ΔG comes out in joules per mole. Each volt of EMF corresponds to ΔG = −96.5 kJ/mol per mole of electrons; for n = 2 (Daniell cell), ΔG = −193 kJ/mol per volt. ΔG = −nFE is exactly equivalent to ΔG = −R·T·ln(K) since both equal the standard free energy change at equilibrium.
How does EMF connect to the equilibrium constant K?
Setting ΔG = −nFE° equal to ΔG° = −R·T·ln(K) gives E°cell = (RT/nF) · ln(K) = (0.0257/n) · ln(K) at 25 °C, or K = exp(nFE°/RT). For a Daniell cell with E° = 1.10 V and n = 2: K = exp(2 × 96485 × 1.10 / (8.314 × 298.15)) = exp(85.6) ≈ 1.4 × 10³⁷. This staggeringly large K confirms the reaction goes essentially to completion — the cell can be fully discharged and the reverse reaction is negligible.
Why is the actual battery voltage lower than the EMF?
EMF is the OPEN-CIRCUIT voltage (no current flowing). When current is drawn under load, three loss mechanisms reduce the terminal voltage: (1) internal resistance (IR drop) — the battery's own electrolyte and electrode resistance dissipates power as heat; (2) activation overpotential — kinetic barrier to electron transfer at the electrode-electrolyte interface; (3) concentration overpotential — local depletion of reactants near the electrode surface from mass-transport limits. A Li-ion cell with EMF 4.2 V might deliver only 3.7 V at typical 1C discharge — a 12% loss.
What's the difference between a galvanic cell and an electrolytic cell?
Galvanic (voltaic) cell: Uses spontaneous redox to generate current (EMF > 0). The chemistry drives the electrons; the cell is the energy source. All batteries during discharge are galvanic. Electrolytic cell: Uses external current to drive non-spontaneous redox (EMF < 0 as written). External power supply is the energy source; chemistry is being forced. Examples: water electrolysis to make H₂; Hall-Héroult process making Al from Al₂O₃; chlor-alkali electrolysis making Cl₂ + NaOH. The same cell can switch roles: a battery discharges as galvanic, then charges as electrolytic.
Why is fluorine the strongest oxidizer?
Because F₂ + 2e⁻ → 2F⁻ has the most positive reduction potential (E° = +2.87 V). Fluorine is the most electronegative element (electronegativity 4.0), so it pulls electrons more strongly than any other species. Combined with its low bond dissociation energy (158 kJ/mol — F-F bond is unusually weak due to lone-pair repulsion) and high electron affinity, F₂ is essentially impossible to use as an oxidizer in aqueous solution because it tears water apart — F₂ + H₂O → HF + HOF. F₂ chemistry happens in non-aqueous solvents only.
Why is lithium the strongest reducer?
Because Li⁺ + e⁻ → Li(s) has the most negative reduction potential (E° = −3.04 V). The reverse reaction (Li → Li⁺ + e⁻, oxidation) is therefore the most spontaneous — lithium gives up its electron more readily than any other metal. Three reasons: low ionization energy (520 kJ/mol — alkali metals lose their s-electron easily), high hydration enthalpy of Li⁺ in water (520 kJ/mol — small ion, strong hydration), and small atomic size (no inner-shell shielding). This is why lithium-based batteries achieve such high voltages — pairing Li anode with high-E° cathode (like CoO₂ at +0.86 V vs SHE) gives EMFs near 4 V.
How is EMF used in corrosion protection?
Sacrificial anodes work by exploiting EMF differences between metals. To protect steel pipelines, ship hulls, or water-heater tanks from oxidation (Fe²⁺/Fe E° = −0.44 V), attach a metal block with a more negative E° — typically magnesium (E° = −2.37 V) or zinc (E° = −0.76 V). The setup forms a galvanic cell with EMF = (−0.44) − (−2.37) = +1.93 V (for Mg-Fe), which preferentially oxidizes the magnesium instead of the iron. The Mg block is consumed over months/years and replaced periodically — far cheaper than replacing the protected structure. Same principle protects underground pipelines (impressed-current cathodic protection adds external power for higher current).

Author Spotlight

The ToolsACE Team - ToolsACE.io Team

The ToolsACE Team

Our chemistry tools team implements the foundational electrochemistry equation E°_cell = E°_cathode − E°_anode — the formula every chemistry and engineering student uses to predict whether a galvanic cell will spontaneously generate current. The convention adopted is the IUPAC standard: BOTH electrode potentials are entered as standard REDUCTION potentials (taken directly from CRC or Bard's electrochemistry tables — no sign-flipping needed), and the formula does the algebra. Output includes the EMF in volts (or mV/kV/MV via unit selector), the cell type (galvanic / electrolytic / equilibrium), 6-band strength classification (very-weak through extreme), an optional Gibbs free-energy calculation ΔG = −nFE if you provide the moles of electrons transferred per cell reaction, full step-by-step calculation breakdown, and a reference table of 23 standard reduction potentials spanning the full range from F₂ (E° = +2.87 V, strongest oxidizer) through Li⁺/Li (E° = −3.04 V, most negative). Faraday's constant is taken as F = 96,485.33 C/mol.

ElectrochemistryGalvanic Cells & BatteriesSoftware Engineering Team

Disclaimer

EMF = E°_cathode − E°_anode applies to STANDARD conditions: 1 M concentrations, 1 atm partial pressures, 25 °C, all activities = 1. For non-standard conditions, use the Nernst equation. EMF is the OPEN-CIRCUIT voltage; under current load the actual voltage drops by 5-30% due to internal resistance, activation overpotential, and concentration overpotential. Reduction-potential values are CRC Handbook standards; some sources cite slightly different values (typically within ±0.05 V) due to differences in electrolyte composition or temperature reference. Both inputs MUST be reduction potentials — flipping the anode sign would double-count the conversion and give an incorrect result.