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CHEMISTRY • ELECTROCHEMISTRY

Electrochemical Cell

Change electrode potentials, temperature and reaction quotient to see how nonstandard conditions shift voltage, free energy and equilibrium.

● NERNST EQUATIONElectrochemical Cell
Cell voltage E
Standard E°
ΔG
log₁₀ K

Reasoning inspector

live model state
Cathode E°—
Anode E°—
Standard cell E°—
Reaction quotient Q—
Temperature—
Electrons n—
Nernst correction—
Cell potential E—
ΔG—
Equilibrium K—
Favored direction—

Connect voltage and chemical driving force

Standard reduction potentials define E°, while Q and temperature shift the actual cell potential. The same voltage directly determines the reaction Gibbs free energy.

LEARN WITH THE SIMULATION

Understand Electrochemical Cell by experimenting.

START HERE

An electrochemical cell links redox thermodynamics to an electrical potential. The standard cell potential E° comes from subtracting the anode reduction potential from the cathode reduction potential, while the Nernst equation shifts the actual voltage according to reaction quotient Q, temperature and electron count n. This laboratory keeps E°, the Nernst correction, actual E, ΔG and equilibrium constant K visible together so concentration effects can be interpreted thermodynamically.

BUILD THE IDEA

What you should understand first

  • Calculate E°cell from tabulated reduction potentials using cathode minus anode.
  • Use the Nernst equation to predict how Q, T and n alter nonstandard cell voltage.
  • Connect the sign of E with the sign of ΔG for the modeled forward reaction.
  • Relate standard voltage to the equilibrium constant through ln K = nFE°/(RT).

TRY IT YOURSELF

What does changing Q alter, and what does it leave unchanged?

  1. Set log₁₀Q = 0 and record E°, E and ΔG.
  2. Increase Q by several decades while keeping half-cell potentials fixed.
  3. Observe that E° and K stay tied to the standard reaction while E and ΔG shift.
  4. Lower Q below 1 and predict the sign of the Nernst correction before checking.
Read the visual:

Q describes the current composition, so it changes the nonstandard driving force E without changing E°. The standard equilibrium constant remains tied to E° at the selected temperature in this model.

WHY IT WORKS 01

E° is constructed from two reduction potentials

Tables usually list both half-reactions as reductions. In a galvanic cell the cathode actually undergoes reduction while the anode undergoes oxidation, so E°cell = E°cathode − E°anode. Reversing which half-cell is the anode changes the sign of the overall reaction and therefore reverses the sign of the cell potential.

WHY IT WORKS 02

The Nernst term is a logarithmic composition correction

For the reaction as written, E = E° − (RT/nF)ln Q. Increasing Q makes the logarithm larger and lowers E, while decreasing Q below 1 makes ln Q negative and raises E above E°. Because the dependence is logarithmic, decade changes in Q produce additive voltage shifts rather than multiplicative voltage changes.

WHY IT WORKS 03

Voltage, Gibbs free energy and equilibrium are three views of one driving force

Electrical work connects cell voltage to reaction free energy through ΔG = −nFE. Under standard conditions, combining ΔG° = −RT ln K with ΔG° = −nFE° gives ln K = nFE°/(RT). A large positive E° therefore corresponds to a large equilibrium constant for the forward reaction at that temperature.

DO NOT MIX THESE UP

Common misconception

Standard potential is not the voltage under every composition. Reduction potentials are listed using reduction convention, so the anode reduction potential is subtracted when forming E°cell. A positive E for the written reaction corresponds to negative ΔG under the modeled conditions, but real batteries also include kinetics, resistance and transport losses that this equilibrium model does not include.

PRACTICE CHALLENGE

Can you reach a target cell voltage by changing composition and then explain the sign of the Nernst correction?

  • Choose a target E in the premium controls.
  • Adjust log₁₀Q while holding the standard half-cell potentials fixed.
  • Use the inspector to compare E with E° and identify the Nernst correction.
  • Explain how the corresponding sign of ΔG determines the favored modeled direction.
Reveal the reasoning

With half-cell potentials fixed, E° remains constant and changing Q alters only the logarithmic Nernst correction. Q > 1 lowers E relative to E°, while Q < 1 raises it. Once E is known, ΔG = −nFE gives the thermodynamic driving-force sign for the reaction as written.

MODEL BOUNDARIES

What this simulation assumes

Every simulation simplifies reality or a mathematical system. These are the assumptions and limits to keep in mind while interpreting the results.

Model assumptions

  • Electrode inputs are reduction potentials and the overall standard voltage is Ecathode minus Eanode.
  • Activities are represented through a dimensionless reaction quotient Q and ideal thermodynamic behavior is assumed.
  • The Nernst equation uses a uniform cell temperature and a whole-number electron transfer count n.

Model limitations

  • Liquid junction potentials, activity coefficients, electrode kinetics, overpotential and internal resistance are not modeled.
  • Real batteries and electrochemical cells can depart substantially from equilibrium thermodynamics under load.

CONNECT IT

Where this matters

  • Battery chemistry
  • Corrosion and electroplating
  • Electrochemical sensors and concentration cells
  • Thermodynamic analysis of redox reactions

WORDS THAT MATTER

Key vocabulary

Standard cell potential
E°cell = E°cathode − E°anode under standard-state conditions.
Nernst equation
Relation E = E° − (RT/nF)lnQ connecting nonstandard voltage to composition and temperature.
Reaction quotient
Instantaneous activity ratio Q for products and reactants in the written cell reaction.
Gibbs free energy
Thermodynamic driving-force quantity linked to voltage by ΔG = −nFE.
Equilibrium constant
Value K of Q at thermodynamic equilibrium, related to E° through lnK = nFE°/(RT).