Electrochemical cell voltages are not fixed numbers: changing ion concentrations or gas pressures changes the driving force for electron flow.

Labeled galvanic cell diagram highlighting anode vs. cathode, the direction of electron flow through the wire, and ion migration through the salt bridge to maintain charge balance. This visual helps connect concentration changes in either half-cell to changes in the reaction quotient $Q$, which in turn shifts $E_{\text{cell}}$. Source

This section explains how and why concentration changes shift measured cell potential.

Core idea: concentration affects the “push” of the redox reaction

For a given cell reaction, $E_{\text{cell}}$ depends on how strongly the system “wants” to move from its current composition toward products. That driving force changes as concentrations of active species change in either half-cell.

What counts as an “active species”

Only species whose effective concentration can change appear in the concentration dependence:

• Aqueous ions and solutes (use molar concentration as an AP-level approximation)

• Gases (use partial pressure as an AP-level approximation)

• Exclude pure solids and pure liquids (their activity is effectively constant)

This matters because adding more solid metal electrode does not directly change $E_{\text{cell}}$, but changing the concentration of its dissolved ions does.

Photograph of a Daniell cell (Zn/Cu) connected to a voltmeter, illustrating how cell potential is measured in practice. Because actual ion concentrations may differ from standard conditions, the observed voltage can differ from $E^\circ_{\text{cell}}$, motivating the need for the Nernst equation. Source

Direction of the change: how concentration shifts $E_{\text{cell}}$

Concentration changes influence the tendency for oxidation and reduction by changing the “availability” of reactants/products in each half-reaction.

General qualitative rules (galvanic conditions)

For the overall cell reaction written as:

• Reactants → Products

Then, at constant temperature:

• Increasing reactant concentration (or pressure, for gases) tends to increase $E_{\text{cell}}$

• Increasing product concentration tends to decrease $E_{\text{cell}}$

• Diluting reactants tends to decrease $E_{\text{cell}}$

• Removing products (lowering their concentration) tends to increase $E_{\text{cell}}$

These trends reflect the syllabus idea that $E_{\text{cell}}$ depends on concentrations and changes as the system composition shifts.

“Farther from equilibrium” means a larger magnitude potential

When reactant and product concentrations are very unequal (for example, reactants high and products low), the cell has a stronger driving force, so the magnitude of the potential is larger. As the reaction proceeds, reactants are consumed and products accumulate, so the driving force weakens and the measured potential typically drops in magnitude.

The Nernst relationship (linking voltage to composition)

A compact way to express concentration dependence is the Nernst equation.

In this form, concentration effects enter through $\ln Q$: changing concentrations changes $Q$, which shifts $E_{\text{cell}}$ relative to $E^\circ_{\text{cell}}$.

Interpreting the sign of the concentration term (without calculation)

• If $Q$ increases (more products relative to reactants), $\ln Q$ increases, so the subtraction term grows and $E_{\text{cell}}$ becomes smaller.

• If $Q$ decreases (fewer products relative to reactants), $\ln Q$ decreases, making the subtraction term smaller (or effectively adding), so $E_{\text{cell}}$ becomes larger.

Common concentration-related pitfalls

• Forgetting to omit solids and liquids from $Q$

• Using concentrations for species that are not meaningfully “active” (for AP, assume dilute solution behaviour unless told otherwise)

• Ignoring that changing either half-cell concentration can change the overall $E_{\text{cell}}$