Practical: Measuring EMF of an Electrochemical Cell Revision Notes | A Level AQA Chemistry

1. Why Measuring EMF Matters

Every standard electrode potential value in your data booklet was measured experimentally — and this practical is how it is done. Getting the technique right is not just about passing the required practical; it underpins your ability to predict whether reactions are feasible and to interpret electrochemical data with confidence.

The core challenge is this: the moment current flows through a cell, the electrode potentials shift away from their standard values. So the entire practical is designed around one goal — measuring voltage while keeping current as close to zero as possible.

2. Key Definitions and the Standard Conditions

Standard conditions are non-negotiable for comparing electrode potentials fairly. They are: temperature 298 K (25 °C), all ion concentrations at 1 mol dm$^{-3}$, and any gases at 1 atm pressure. Deviate from these and you are no longer measuring $E^\circ$ — you are measuring a non-standard potential.

The platinum electrode in the SHE is inert — it does not take part in the reaction. Its role is purely to conduct electrons and provide a surface for the $\text{H}_2 / \text{H}^+$ equilibrium to be established. This is why platinum is also used in any half-cell involving gases or non-conducting species.

3. Apparatus and Cell Construction

Picture two beakers side by side. Each contains a metal electrode dipping into a solution of its own ions — for example, zinc in zinc sulphate solution and copper in copper sulphate solution. These are your two half-cells. On their own, they do nothing measurable; you need two connections to complete the circuit.

• A high-resistance voltmeter connected between the two metal electrodes — this measures the potential difference while allowing negligible current to flow.
• A salt bridge connecting the two solutions — typically filter paper soaked in saturated potassium nitrate (KNO$_3$) solution, or a U-tube filled with the same. It completes the ionic circuit without allowing the two solutions to mix.
• Clean metal electrodes (or platinum for gas/ion-only half-cells) with no oxide layer that could introduce a contact resistance.

The salt bridge deserves particular attention. Without it, charge would build up in each beaker — the zinc beaker would become positively charged as Zn²⁺ ions enter the solution, and the copper beaker would become negatively charged as Cu²⁺ ions are removed by deposition. The salt bridge allows K⁺ and NO₃⁻ ions to migrate and neutralise this charge imbalance, keeping the cell running.

Potassium nitrate is chosen because both $\text{K}^+$ and $\text{NO}_3^-$ ions are unlikely to react with the solutions in either half-cell, and they have similar mobilities, which minimises any junction potential introduced at the salt bridge interface.

4. Why a High-Resistance Voltmeter is Essential

Think of electrode potential like the pressure in a sealed tank. The moment you open a valve and let fluid flow, the pressure drops. The same happens here: if current flows through the cell, the electrode reactions proceed, concentrations change, and the potential difference you measure is no longer the equilibrium value — it is a lower, non-standard reading.

A high-resistance voltmeter (typically greater than 10 MΩ) draws negligible current — so negligible that the electrode reactions are effectively frozen. The reading you get is the true EMF, not a loaded voltage. A standard low-resistance ammeter connected instead would allow significant current to flow and give a meaningless, lower reading.

5. Calculating the Theoretical EMF

Before you can evaluate your experimental result, you need to know what value to expect. The theoretical cell EMF is calculated from standard electrode potentials using a single equation. The half-cell with the more positive $E^\circ$ acts as the cathode (reduction occurs there); the half-cell with the less positive $E^\circ$ acts as the anode (oxidation occurs there).

Worked example: Calculating the EMF of a Zn/Cu Cell

Calculate the standard EMF of a cell constructed from a zinc half-cell and a copper half-cell. Standard electrode potentials: $\text{Zn}^{2+}(\text{aq}) + 2e^- \rightleftharpoons \text{Zn(s)}$, $E^\circ = -0.76$ V; $\text{Cu}^{2+}(\text{aq}) + 2e^- \rightleftharpoons \text{Cu(s)}$, $E^\circ = +0.34$ V.

1. Cu$^{2+}$/Cu: $E^\circ = +0.34$ V (more positive) → cathode Zn$^{2+}$/Zn: $E^\circ = -0.76$ V (less positive) → anode — The half-cell with the more positive $E^\circ$ is more readily reduced, so it acts as the cathode. Reduction occurs at the cathode; oxidation at the anode.
2. Cathode (reduction): $\text{Cu}^{2+}(\text{aq}) + 2e^- \rightarrow \text{Cu(s)}$ Anode (oxidation): $\text{Zn(s)} \rightarrow \text{Zn}^{2+}(\text{aq}) + 2e^-$ — At the cathode, copper ions gain electrons and are deposited as copper metal. At the anode, zinc metal loses electrons and dissolves as zinc ions.
3. $E^\circ_{\text{cell}} = E^\circ_{\text{cathode}} - E^\circ_{\text{anode}}$ $E^\circ_{\text{cell}} = (+0.34) - (-0.76)$ — Subtracting a negative number is the same as adding its magnitude. Take care with the signs here — this is where most errors occur.
4. $E^\circ_{\text{cell}} = +0.34 + 0.76 = +1.10$ V — The positive value confirms the reaction is thermodynamically feasible under standard conditions. The cell produces 1.10 V.
5. Copper is below zinc in the electrochemical series (more positive $E^\circ$), so copper ions should oxidise zinc metal. The positive EMF is consistent with this. — Always verify the sign makes chemical sense. If you get a negative EMF, you have likely swapped cathode and anode.

Answer: $E^\circ_{\text{cell}} = +1.10$ V. The copper electrode is the positive terminal (cathode) and the zinc electrode is the negative terminal (anode).

6. The Standard Hydrogen Electrode in Practice

When you measure the $E^\circ$ of any half-cell, you are always measuring it against the SHE. In the Zn/Cu practical above, you are measuring the difference between two non-hydrogen half-cells — but the SHE underpins the entire scale. Every $E^\circ$ value in your data booklet was originally determined by connecting that half-cell to an SHE and reading the voltmeter.

In school practicals, the SHE is rarely set up directly because it requires pure hydrogen gas, a platinised platinum electrode, and precise pressure control. Instead, you measure the EMF between two non-hydrogen half-cells and compare your result with the difference in their tabulated $E^\circ$ values. The principle is identical — you are still using the SHE scale, just indirectly.

If you were to connect a zinc half-cell directly to an SHE, the voltmeter would read +0.76 V with the SHE as the positive terminal, confirming $E^\circ(\text{Zn}^{2+}/\text{Zn}) = -0.76$ V. The sign tells you that zinc is oxidised relative to hydrogen — zinc is the anode in that pairing.

7. Sources of Error and Evaluation

Your measured EMF will almost certainly differ slightly from the theoretical $E^\circ_{\text{cell}}$ value. This is not a failure — it is an opportunity to demonstrate analytical thinking, which is exactly what AQA mark schemes reward in evaluation questions.

When evaluating reliability, note that repeating the measurement three times and calculating a mean reduces the effect of random errors such as slight variations in salt bridge contact. Systematic errors — such as consistently using a slightly incorrect concentration — cannot be reduced by repetition and must be addressed by improving the method.

8. Exam Technique and Mark-Winning Guidance

AQA questions on this practical tend to cluster around three areas: describing the method, explaining why specific features are necessary (especially the high-resistance voltmeter and salt bridge), and evaluating discrepancies between measured and theoretical EMF. Each of these demands a different approach.

For describe questions, give a sequential account of what you do — apparatus, setup, measurement. Do not explain why at this stage. For explain questions, always include a because-clause that links to an underlying principle: for example, 'a high-resistance voltmeter is used because it prevents current from flowing, which would shift the electrode potentials away from their equilibrium values and reduce the measured EMF.' For evaluate questions, identify the discrepancy, give at least two specific reasons for it, and reach a conclusion about the reliability of the method.

Mastering this practical means understanding not just the steps, but the reason behind each one. Every design choice — the high-resistance voltmeter, the salt bridge composition, the electrode cleaning — connects directly to the principle that EMF must be measured at zero current under standard conditions. That understanding is what separates a descriptive answer from a full-mark analytical response.