Nernst Equation Calculator

Nernst Equation Calculator

Calculate the reduction potential using the Nernst equation.
Enter the standard reduction potential \(E^{\circ}\), temperature \(T\) (K), number of electrons \(n\), and the reaction quotient \(Q\).

* All values should be in SI units.

Step 1: Enter Parameters

Example: 0.00 V

Default: 298.15 K (25°C)

Example: 2

Example: 1 (for standard conditions)

Derived Equation:
$$E = E^{\\circ} – \\frac{RT}{nF} \\ln Q$$
where \(R = 8.314\\,\\text{J/(mol·K)}\) and \(F = 96485\\,\\text{C/mol}\).


Example:
For \(E^{\\circ} = 0.00\\,\\text{V}\), \(T = 298.15\\,\\text{K}\), \(n = 2\), and \(Q = 1\), the reduction potential is calculated.

The Equation:

E = E° – (RT/nF) ln Q

  • E: Electrode potential under non-standard conditions (volts, V)
  • E°: Standard electrode potential (volts, V)
  • R: Universal gas constant (8.314 J/(mol·K))
  • T: Absolute temperature in Kelvin (K)
  • n: Number of electrons transferred
  • F: Faraday’s constant (96,485 C/mol)
  • Q: Reaction quotient from concentrations or pressures of reactants/products

Understanding the Nernst Equation:

The Nernst equation describes how electrode potentials change when reaction conditions differ from the standard state (1 M for solutions, 1 atm for gases).

  • When Q > 1 (more products), E decreases and the reaction is less spontaneous.
  • When Q < 1 (more reactants), E increases and the reaction is more spontaneous.

Simplified Form at 25°C (298 K):

E = E° – (0.0592 V/n) log Q

Practical Applications:

  1. Calculating Electrode Potentials: Predict electrode behavior under varied conditions.
  2. Predicting Reaction Direction: Determine reaction spontaneity and direction.
  3. Determining Equilibrium Constants: Calculate equilibrium constants from electrode potentials.
  4. Analyzing Concentration Effects: Understand how changes in concentration affect cell potentials.

Examples of Use:

  • Batteries: Evaluate voltage during partial discharge.
  • Fuel Cells: Assess efficiency based on fuel and oxidant concentrations.
  • Analytical Chemistry: Electrochemical sensors and potentiometric titrations.

Conclusion:

The Nernst equation connects theoretical electrochemistry with practical applications, highlighting its importance in research, industry, and education.