An electrochemical cell converts chemical energy directly into electrical energy by separating the two half-reactions of a displacement reaction so that electrons must travel through an external circuit. The greater the difference in reactivity between the two metals involved, the larger the voltage the cell can produce.
What is an electrochemical cell?
When a more reactive metal is dipped directly into a solution containing ions of a less reactive metal, a displacement reaction occurs — electrons are transferred directly between the metals and the energy is released as heat. An electrochemical cell exploits this same chemistry, but separates the two metals so that electrons must flow through an external wire rather than jumping directly.
A simple electrochemical cell consists of:
- Two different metal electrodes (rods or strips) dipping into electrolyte solutions.
- A salt bridge (or a porous plug) connecting the two solutions, allowing ions to flow between them to maintain electrical neutrality.
- An external wire connecting the two electrodes, along which electrons flow.
- A voltmeter (or other device) in the external circuit to detect or use the voltage.
What happens at each electrode?
In an electrochemical cell, the two half-reactions of the overall displacement are separated:
- Negative electrode (anode) — the more reactive metal. Metal atoms lose electrons and enter solution as positive ions. This is oxidation: M(s) → M²⁺(aq) + 2e⁻
- Positive electrode (cathode) — the less reactive metal. Metal ions from solution gain electrons and are deposited as atoms. This is reduction: M²⁺(aq) + 2e⁻ → M(s)
Electrons flow from the more reactive metal (anode) to the less reactive metal (cathode) through the external circuit — in the opposite direction to conventional current (conventional current flows from + to −).
Example — Zinc–copper cell (Daniel cell):
Anode (zinc): Zn(s) → Zn²⁺(aq) + 2e⁻ (oxidation) Cathode (copper): Cu²⁺(aq) + 2e⁻ → Cu(s) (reduction) Overall: Zn(s) + Cu²⁺(aq) → Zn²⁺(aq) + Cu(s)
How does reactivity difference affect cell voltage?
The voltage produced by an electrochemical cell is related to how far apart the two metals are in the reactivity series. The greater the difference in reactivity, the greater the tendency for electron transfer, and the higher the voltage measured.
| Metal pair (more reactive + less reactive) | Approximate cell voltage |
|---|---|
| Zinc + copper | ~1.1 V |
| Magnesium + copper | ~2.7 V |
| Iron + copper | ~0.78 V |
| Zinc + iron | ~0.32 V |
| Zinc + zinc (same metal) | 0 V |
Using the same metal for both electrodes gives zero voltage — there is no difference in reactivity, so no net electron transfer occurs.
Using different electrolyte solutions also affects the voltage, but the dominant factor at GCSE level is the identity of the two metals.
What is the role of the salt bridge?
As the cell operates, positive metal ions (e.g. Zn²⁺) enter the solution at the anode, making it progressively more positive. At the cathode, Cu²⁺ ions are removed from solution, making it progressively more negative. Without a pathway for ion flow, these charge imbalances would halt electron flow almost immediately.
The salt bridge — typically a tube filled with a saturated inert salt solution (e.g. potassium nitrate) — allows negative ions to migrate towards the anode compartment and positive ions to migrate towards the cathode compartment, neutralising the charge build-up. This allows electron flow in the external circuit to continue.
How does an electrochemical cell differ from electrolysis?
Students sometimes confuse electrochemical cells and electrolysis. The key distinction:
| Feature | Electrochemical cell | Electrolysis |
|---|---|---|
| Energy conversion | Chemical energy → electrical energy | Electrical energy → chemical energy |
| Energy source | The cell itself (spontaneous redox reaction) | External power supply (non-spontaneous) |
| Everyday example | Battery, fuel cell | Electroplating, extracting aluminium |
| Reaction driven? | Spontaneous — occurs on its own | Non-spontaneous — requires energy input |
Both involve oxidation at an anode and reduction at a cathode — but in opposite energy directions.
What are fuel cells and how do they differ from simple cells?
A fuel cell is a type of electrochemical cell that generates electricity by the controlled oxidation of a fuel (typically hydrogen) without combustion. Unlike a battery, it does not run down — it continues to generate electricity as long as fuel and oxygen are supplied.
Hydrogen fuel cell:
- Anode: H₂ → 2H⁺ + 2e⁻ (oxidation)
- Cathode: O₂ + 4H⁺ + 4e⁻ → 2H₂O (reduction)
- Overall: 2H₂ + O₂ → 2H₂O
The only product is water. Fuel cells have high efficiency (around 60–80 % compared with ~30–35 % for a petrol engine) and produce no direct CO₂ emissions at the point of use. The main challenges for widespread use are the cost of producing and storing hydrogen safely, and the expense of the platinum catalyst at the electrodes.
Frequently asked questions
What is a simple electrochemical cell in GCSE chemistry?
A simple electrochemical cell consists of two different metal electrodes connected by an external wire and a salt bridge. The more reactive metal acts as the negative electrode (anode) and loses electrons by oxidation; the less reactive metal acts as the positive electrode (cathode) and gains electrons by reduction. Electrons flow through the wire from anode to cathode, producing an electric current. The voltage produced depends on how far apart the two metals are in the reactivity series.
Why does a greater difference in reactivity produce a higher voltage?
A metal higher in the reactivity series has a stronger tendency to lose electrons and form ions. The voltage of an electrochemical cell reflects the difference in this tendency between the two electrode metals. If both metals have similar reactivities, there is little driving force for electron transfer and the voltage is low. If one metal is much more reactive than the other (e.g. magnesium vs copper), the driving force is large and the voltage is high.
What is the difference between oxidation and reduction at a cell electrode?
At the negative electrode (anode), the more reactive metal undergoes oxidation — metal atoms lose electrons and enter solution as positive ions. At the positive electrode (cathode), metal ions from solution undergo reduction — they gain electrons and are deposited as solid metal. A useful memory aid is OIL RIG: Oxidation Is Loss (of electrons), Reduction Is Gain. The electrons released by oxidation at the anode travel through the external circuit and are used for reduction at the cathode.
Why does using the same metal for both electrodes give zero voltage?
If both electrodes are the same metal in contact with the same electrolyte, there is no difference in the tendency to lose or gain electrons. The two half-cells are identical, so no net electron transfer occurs and the voltmeter reads zero. A voltage is only generated when there is a difference in reactivity — i.e. a difference in the tendency of the two electrode metals to be oxidised. This is why different metals are required for a functioning cell.
For Socratic GCSE chemistry with Professor Curie — reasoning from electron transfer and particle-level half-reactions before wiring any circuit — visit aitutors.me.