1.1
V
1
0
V
1.1
V
-212.27
kJ/mol
100
1.1
V
1
0
V
1.1
V
-212.27
kJ/mol
100
The Galvanic Cell Calculator models a spontaneous electrochemical cell (also called a voltaic cell) that converts chemical energy into electrical energy. By selecting an anode and cathode metal, specifying ion concentrations, and setting the temperature, you can compute both the standard cell EMF and the actual cell potential under non-standard conditions using the Nernst equation. This calculator also determines the Gibbs free energy change and verifies whether the cell is truly spontaneous. Galvanic cells are the basis for all batteries, from the classic Daniell cell to modern lithium-ion technology. Understanding how electrode choice and concentration affect voltage is crucial for battery design, electrochemical engineering, and corrosion science.
A galvanic cell consists of two half-cells connected by a salt bridge. The anode (more negative E°) undergoes oxidation, releasing electrons that flow through the external circuit to the cathode (more positive E°), where reduction occurs.
$$E°_{cell} = E°_{cathode} - E°_{anode}$$
Under non-standard conditions, the Nernst equation gives the actual potential:
$$E = E°_{cell} - \frac{RT}{nF} \ln Q$$
For a simple metal-ion cell, the reaction quotient is approximated as:
$$Q = \frac{[\text{anode ions}]}{[\text{cathode ions}]}$$
The Gibbs free energy under actual conditions is:
$$\Delta G = -nFE$$
A positive E means the reaction is spontaneous and the cell can do electrical work. The salt bridge maintains electrical neutrality by allowing counter-ion migration between the half-cells.
If the actual cell potential E > 0, the cell operates spontaneously as a galvanic cell. Electrons flow from the anode to the cathode through the wire. A larger E means more driving force and greater maximum work. When anode concentration is high relative to cathode concentration (Q > 1), the cell potential decreases. When cathode concentration is high (Q < 1), the cell potential increases beyond E°. If you select two metals where the "anode" has a higher E° than the "cathode," the calculator will show a negative EMF, indicating the cell is non-spontaneous as configured — swap the electrodes.
Inputs
Results
E°cell = 0.34 − (−0.76) = 1.10 V. At standard concentrations (Q = 1), the actual potential equals E°. Zinc dissolves at the anode and copper deposits at the cathode. ΔG = −212.27 kJ/mol confirms spontaneity.
Inputs
Results
Q = 1.0/0.001 = 1000. E = 1.10 − (0.01285) × ln(1000) = 1.10 − 0.0887 = 1.0113 V. Lower cathode concentration reduces the driving force, giving a lower cell voltage.
A galvanic (voltaic) cell is an electrochemical device that converts chemical energy into electrical energy through spontaneous redox reactions. It consists of two half-cells with different electrode potentials connected by an external circuit and a salt bridge.
The metal with the more negative standard reduction potential is the anode (oxidation occurs). The metal with the more positive potential is the cathode (reduction occurs). This arrangement gives a positive E°cell and a spontaneous reaction.
The salt bridge maintains electrical neutrality by allowing ions to migrate between the half-cells. Without it, charge buildup would quickly stop the cell from operating. Common salt bridges contain KCl or KNO₃ in agar gel.
According to Le Chatelier's principle and the Nernst equation, changing concentrations shifts the equilibrium. Higher reactant concentration (lower Q) increases voltage; higher product concentration (higher Q) decreases it.
Cell notation describes the cell compactly: Anode | Anode ion || Cathode ion | Cathode. For example, Zn(s)|Zn²⁺(aq)||Cu²⁺(aq)|Cu(s). The single line represents a phase boundary and the double line represents the salt bridge.
Yes, any two metals with different reduction potentials can form a galvanic cell. The greater the difference in E° values, the higher the cell voltage. Even metals close in the series (like Ni and Sn) produce a small but measurable voltage.
As the cell operates, anode metal dissolves (increasing anode ion concentration) and cathode ions deposit (decreasing cathode ion concentration). Q increases, E decreases. Eventually Q = K and E = 0 — the cell is dead.
Maximum electrical work equals −ΔG = nFE. This is achieved only at infinitesimally small current (reversible operation). Real cells deliver less work due to internal resistance and overpotential losses.
Temperature affects both E° and the Nernst correction term (RT/nF). Generally, higher temperatures increase the magnitude of the concentration correction. The direction depends on whether ΔS° is positive or negative for the cell reaction.
The Daniell cell is the classic galvanic cell with a zinc anode in ZnSO₄ solution and a copper cathode in CuSO₄ solution. Invented by John Frederic Daniell in 1836, it produces about 1.1 V and was one of the first practical batteries.
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