RP 8: Measuring the EMF of an electrochemical cell

Hazards

Method

1. Build the two half-cellsClean strips of zinc and copper with emery paper (then a propanone rinse) so the surface is fresh metal. Stand each metal in a beaker of a 1.00 mol dm⁻³ solution of its own ions: Zn in ZnSO₄(aq), Cu in CuSO₄(aq).
2. Bridge them with a salt bridgeSoak a strip of filter paper in saturated potassium nitrate solution and drape it between the beakers. It completes the circuit by letting ions migrate — the electrons take the wire. KNO₃ is the safe choice because every common nitrate is soluble; KCl would precipitate AgCl in a silver half-cell.
3. Measure with a high-resistance voltmeterConnect the electrodes to a high-resistance voltmeter: it draws essentially no current, so you read the true EMF at zero current. Record the value and note which electrode is positive.
4. Compare with the data bookE°cell = E°(positive electrode) − E°(negative electrode). For Zn/Cu the standard value is +1.10 V at 298 K, 1.00 mol dm⁻³, 100 kPa — quote the conditions whenever you quote the number.
5. Change a concentrationDilute the Cu²⁺ solution tenfold and re-measure: the Cu²⁺ + 2e⁻ ⇌ Cu equilibrium shifts left, the copper electrode potential falls, and the cell EMF drops — le Chatelier applied to electrochemistry.

Mistakes examiners see every year

What to expect

Zn/Cu gives ≈ +1.10 V with copper positive; classroom values fall a little short through non-standard concentrations and junction potentials. Tenfold dilution of the Cu²⁺ drops the EMF by roughly 30 mV (indicative) — the direction, not the size, is what the spec asks you to predict.