Topic 16 · page 8 of 19

Why the conditions matter

Every electrode potential comes from an equilibrium, so concentration, pressure and temperature all change it. E equals the standard value only under standard conditions.

10 min read16.9

Read first: Core Practical 12: electrochemical cells

An electrode potential is an equilibrium 16.9

At every electrode a reversible reaction reaches equilibrium between the metal (or other reduced form) and its ions. For copper:

CuX2+(aq)+2 eX−⇌Cu(s) \ce{Cu^2+(aq) + 2e- <=> Cu(s)}

The more the position of equilibrium lies to the right, the more readily the electrode takes electrons, and the more positive its potential. Anything that shifts the equilibrium therefore changes the potential. Use Le Chatelier's principle to decide which way.

Change to the Cu²⁺/Cu half-cell Equilibrium shifts Potential, E
increase [Cu²⁺] right, using up Cu²⁺ more positive
decrease [Cu²⁺] left, making Cu²⁺ less positive

A change in temperature, or in the pressure of a gas in a half-cell, also shifts the equilibrium, so E changes then as well.

This is why every half-cell in a table is measured under the same, standard conditions. It is the only way the values can be compared fairly.

What happens to a cell 16.9

Cell emf is the difference between two electrode potentials, so a change that moves one electrode changes the emf.

Worked exampleIncreasing [Zn²⁺] in the zinc-copper cell

  1. The zinc equilibrium is ZnX2++2 eX−⇌Zn\ce{Zn^2+ + 2e- <=> Zn}.
  2. More Zn²⁺ shifts it to the right, towards Zn.
  3. So EE(Zn) becomes less negative: it moves closer to the copper value.
  4. The gap between the electrodes shrinks, so EcellE_\text{cell} decreases.

Worked exampleLowering [Cu²⁺] in the zinc-copper cell

  1. The copper equilibrium is CuX2++2 eX−⇌Cu\ce{Cu^2+ + 2e- <=> Cu}.
  2. Less Cu²⁺ shifts it to the left.
  3. So EE(Cu) becomes less positive.
  4. The gap between the electrodes shrinks, so EcellE_\text{cell} decreases.

The same idea explains why real readings in Core Practical 12 are close to, but not exactly, the standard values.

Check yourself

  1. Why does the potential of an electrode depend on the concentration of its ions?
  2. What happens to the potential of the Ag⁺/Ag half-cell if the Ag⁺ concentration is increased? Explain.
  3. In the zinc-copper cell, the concentration of Cu²⁺ is increased. What happens to EcellE_\text{cell}? Explain.
  4. A student measures a copper half-cell against a hydrogen electrode in the lab and gets +0.31 V, not +0.34 V. Suggest one reason.
  5. Why must standard electrode potentials all be measured under the same conditions?

Answers cover these until you have tried the questions

  1. The electrode reaction is a reversible equilibrium that involves the ions. Changing their concentration shifts the equilibrium (Le Chatelier), which changes how readily the electrode takes electrons.
  2. AgX++eX−⇌Ag\ce{Ag+ + e- <=> Ag} shifts to the right, so the electrode takes electrons more readily and its potential becomes more positive.
  3. CuX2++2 eX−⇌Cu\ce{Cu^2+ + 2e- <=> Cu} shifts right, so EE(Cu) becomes more positive. The gap to zinc widens, so EcellE_\text{cell} increases.
  4. The conditions were not standard: for example [Cu²⁺] was below 1.00 mol dm⁻³, or the temperature was not 298 K. A lower [Cu²⁺] shifts the equilibrium left and makes E less positive.
  5. E changes with concentration, pressure and temperature. Using the same standard conditions means the values can be compared fairly.