Why does the cell potential drop to zero at equilibrium?

At equilibrium, the reaction quotient Q equals the equilibrium constant K. Substituting Q = K into the Nernst equation gives E = E° − (RT/nF) ln K. Since the standard relationship is E° = (RT/nF) ln K, these two terms cancel, resulting in E = 0. A zero cell potential means there is no net driving force for the reaction in either direction; the system is at balance.

What's the practical use of the Nernst equation?

The Nernst equation is fundamental to the operation of pH meters, ion-selective electrodes, and all potentiometric sensors. These devices measure voltage (E) to determine the concentration of an ion in solution. By holding E°, n, T, and other concentrations constant, the measured potential varies logarithmically with the concentration of the target ion, as described by the Nernst equation.

Does the Nernst equation apply to both galvanic and electrolytic cells?

Yes, it applies to both. For a galvanic (voltaic) cell, E is positive, indicating a spontaneous reaction that produces electrical work. For an electrolytic cell, an applied external voltage forces a non-spontaneous reaction; the Nernst equation calculates the minimum voltage that must be applied to drive the reaction, which corresponds to a negative calculated E for the forward cell reaction.

Why do we use natural log (ln) in the equation, and can we use log base 10?

The thermodynamic derivation naturally yields the natural logarithm. However, a common practical form uses base-10 log: E = E° − (2.303RT/nF) log Q. The constant 2.303 is the conversion factor (ln 10). Both forms are correct; the simulator uses the natural log form as it is the most fundamental, but the behavior is identical in principle.

Nernst Equation

E = E° − (RT/nF) ln Q: sliders for E°, n, T, and reaction quotient.

Adjust sliders to see how standard potential, electron count, temperature, and Q move the reversible cell or half-cell potential.

+818 mV

About this model

The Nernst equation provides a quantitative relationship between the potential of an electrochemical cell under non-standard conditions and its standard potential. It is expressed as E = E° − (RT/nF) ln Q, where E is the cell potential, E° is the standard cell potential, R is the universal gas constant, T is the absolute temperature in Kelvin, n is the number of moles of electrons transferred in the redox reaction, F is Faraday's constant, and Q is the reaction quotient. This simulator allows you to manipulate these key variables—E°, n, T, and Q—to observe their direct impact on the calculated cell potential. The underlying principle is that the driving force for an electrochemical reaction (the cell potential) decreases as the reaction proceeds and the concentrations of products increase relative to reactants, moving the system toward equilibrium where E = 0. The model simplifies real electrochemical cells by assuming ideal behavior, ignoring junction potentials, and treating activity coefficients as unity (so concentrations can be used directly in Q). It also holds R and F as constants. By interacting with the sliders, students learn how the spontaneity of a redox reaction (indicated by a positive E) depends not just on the standard potential but critically on the reaction conditions. They can visualize how increasing product concentration (increasing Q) diminishes the cell's voltage, how temperature influences the sensitivity of E to Q, and why the number of electrons transferred is a crucial scaling factor.

Who it's for: Undergraduate chemistry students studying electrochemistry and thermodynamics, particularly in courses covering galvanic cells, equilibrium, and the Nernst equation.

Key terms

Nernst Equation
Cell Potential
Standard Potential
Reaction Quotient
Redox Reaction
Faraday's Constant
Electrochemical Cell
Equilibrium Constant

How it works

The Nernst equation shifts the equilibrium electrode potential when concentrations depart from standard states. It links electrochemistry to reaction quotients and temperature.

Key equations

E = E° − (RT / nF) ln Q

Frequently asked questions

Why does the cell potential drop to zero at equilibrium?: At equilibrium, the reaction quotient Q equals the equilibrium constant K. Substituting Q = K into the Nernst equation gives E = E° − (RT/nF) ln K. Since the standard relationship is E° = (RT/nF) ln K, these two terms cancel, resulting in E = 0. A zero cell potential means there is no net driving force for the reaction in either direction; the system is at balance.
What's the practical use of the Nernst equation?: The Nernst equation is fundamental to the operation of pH meters, ion-selective electrodes, and all potentiometric sensors. These devices measure voltage (E) to determine the concentration of an ion in solution. By holding E°, n, T, and other concentrations constant, the measured potential varies logarithmically with the concentration of the target ion, as described by the Nernst equation.
Does the Nernst equation apply to both galvanic and electrolytic cells?: Yes, it applies to both. For a galvanic (voltaic) cell, E is positive, indicating a spontaneous reaction that produces electrical work. For an electrolytic cell, an applied external voltage forces a non-spontaneous reaction; the Nernst equation calculates the minimum voltage that must be applied to drive the reaction, which corresponds to a negative calculated E for the forward cell reaction.
Why do we use natural log (ln) in the equation, and can we use log base 10?: The thermodynamic derivation naturally yields the natural logarithm. However, a common practical form uses base-10 log: E = E° − (2.303RT/nF) log Q. The constant 2.303 is the conversion factor (ln 10). Both forms are correct; the simulator uses the natural log form as it is the most fundamental, but the behavior is identical in principle.

Similar topics nearby (including related fields) — or browse all 65 in Chemistry.

View category →

NewSchool

Galvanic (Voltaic) Cell

Launch Simulator

Two half-cells, salt bridge, voltmeter; E°cell and Nernst E from ion concentrations (pedagogical E°).

NewSchool

Nernst Potentials & Membrane

Launch Simulator

Nernst E for K⁺ and Na⁺ from inside/outside concentrations; optional Goldman–Hodgkin–Katz resting estimate.

NewSchool

Le Châtelier Principle (Gas)

Launch Simulator

N₂ + 3H₂ ⇌ 2NH₃: T, V, add/remove species; Q vs K and relaxing mole bars.

NewSchool

Reaction Rate

Launch Simulator

Arrhenius k(T), order 0/1/2, [A](t) curves and particle tank.

NewSchool

Radioactive Decay & Chain

Launch Simulator

N(t) from half-life T₁/₂; optional parent → daughter → stable Bateman ODEs (RK4, normalized).

NewSchool

Henry's Law (Gas Solubility)

Launch Simulator

p = k_H c with k_H(T) from ΔH; c vs T at fixed p (ideal-dilute sketch).

Nernst Equation

About this model

How it works

Key equations

Frequently asked questions

Related simulators

Galvanic (Voltaic) Cell

Nernst Potentials & Membrane

Le Châtelier Principle (Gas)

Reaction Rate

Radioactive Decay & Chain

Henry's Law (Gas Solubility)

Nernst Equation

About this model

How it works

Key equations

Frequently asked questions