What is EMF?
EMF is an abbreviation of the electromotive force. By definition, EMF is a force or an electric pressure that generates a current in a circuit. The EMF units are volts [V].
To understand this definition better, consider a typical galvanic or voltaic cell consisting of two electrodes. The electromotive force is the potential difference between them. Additionally, the electromotive force of a galvanic cell specifies the ability of an electrochemical cell to cause the flow of electrons through the electric circuit.
EMF Equation
The electromotive force of the electrochemical cell can be calculated using the equation:
where Ecathode is the standard reduction potential of the cathode (in volts) and Eanode is the standard reduction potential of the anode (in volts). Remember that in a galvanic cell, the potential of the cathode is higher than the potential of the anode.
How to Calculate EMF — Daniell Cell Example
Let's consider a Daniell cell — a simple electrochemical cell invented in 1836 by John Frederic Daniell. In this cell, copper is immersed in a solution of copper(II) sulfate, and zinc is immersed in a solution of zinc sulfate.
Cell schema: (−) Zn | Zn²⁺ ‖ Cu²⁺ | Cu (+)
- Check the electrode potentials: Zn = −0.76 V, Cu = +0.34 V
- Because Zn has the lower potential, zinc is the anode and copper is the cathode
- Calculate the EMF: EMF = +0.34 V − (−0.76 V) = 1.10 V
Electropotential of Anode and Cathode
If you want to find the potential of a metal electrode, you can check the galvanic series or calculate it using the Nernst equation (cell potential equation). The Nernst equation allows you to calculate the reduction potential of a half-cell or full-cell reaction.
In our calculator, you can select from 20 common electrode materials or enter a custom potential value. All standard potentials listed are measured at 25°C (77°F) and 1 atm pressure relative to the Standard Hydrogen Electrode (SHE).
Standard Electrode Potentials Table
| Electrode | E° (V) | E° (mV) |
|---|---|---|
| Li / Li⁺ | −3.04 | −3040 |
| Ca / Ca²⁺ | −2.86 | −2860 |
| Mg / Mg²⁺ | −2.36 | −2360 |
| Al / Al³⁺ | −1.69 | −1690 |
| Mn / Mn²⁺ | −1.18 | −1180 |
| Zn / Zn²⁺ | −0.76 | −760 |
| Cr / Cr³⁺ | −0.74 | −740 |
| Fe / Fe²⁺ | −0.44 | −440 |
| Cd / Cd²⁺ | −0.40 | −400 |
| Co / Co²⁺ | −0.28 | −280 |
| Ni / Ni²⁺ | −0.26 | −260 |
| Sn / Sn²⁺ | −0.14 | −140 |
| Pb / Pb²⁺ | −0.14 | −140 |
| Fe / Fe³⁺ | −0.04 | −40 |
| H₂ / 2H⁺ | 0.00 | 0 |
| Bi / Bi³⁺ | +0.32 | +320 |
| Cu / Cu²⁺ | +0.34 | +340 |
| Ag / Ag⁺ | +0.80 | +800 |
| Hg / Hg²⁺ | +0.85 | +850 |
| Au / Au³⁺ | +1.52 | +1520 |
Electromotive Force Sources
The most common sources of EMF include:
- Electrochemical cells — galvanic (voltaic) cells that convert chemical energy to electrical energy
- Batteries — series of electrochemical cells (e.g., alkaline, lithium-ion)
- Fuel cells — convert chemical energy of a fuel into electricity via electrochemical reaction
- Thermoelectric generators — convert heat difference into electrical energy (Seebeck effect)
- Electromagnetic induction — generators that produce EMF via changing magnetic flux (Faraday's law)
- Solar cells — photovoltaic conversion of light energy to electrical energy
Electrochemical Cell Types
Galvanic Cell (Voltaic Cell)
A galvanic cell (also called a voltaic cell) is an electrochemical cell that derives electrical energy from spontaneous redox reactions. The EMF of a galvanic cell is positive (EMF > 0).
- The anode is the negative electrode where oxidation occurs
- The cathode is the positive electrode where reduction occurs
- Electrons flow from anode to cathode through the external circuit
Electrolytic Cell
An electrolytic cell uses electrical energy to drive a non-spontaneous chemical reaction. The EMF of the reaction would be negative (EMF < 0), meaning external energy must be supplied. Examples include electroplating, electrolysis of water, and charging rechargeable batteries.
Redox Reaction
In an electrochemical cell, a redox reaction (reduction-oxidation) takes place. The overall cell reaction consists of two half-reactions:
- Oxidation half-reaction (at the anode): the metal loses electrons — M → Mn+ + ne⁻
- Reduction half-reaction (at the cathode): metal ions gain electrons — Mn+ + ne⁻ → M
For the Daniell cell example:
- Anode (oxidation): Zn → Zn²⁺ + 2e⁻
- Cathode (reduction): Cu²⁺ + 2e⁻ → Cu
- Overall: Zn + Cu²⁺ → Zn²⁺ + Cu
Unit Systems
Our EMF calculator supports two display unit options:
- Volts (V) — the SI metric unit of electromotive force. Used in standard electrochemistry and physics.
- Millivolts (mV) — 1/1000 of a volt. Commonly used in practical electrochemistry, biochemistry, and precision measurements (1 V = 1000 mV).
Both units are internationally recognized; the choice depends on the scale of measurement. Small biological cell potentials are typically expressed in mV, while standard electrode potentials use V.