What is an Electrochemical Cell?
Imagine a tiny battery that can turn a chemical reaction into electric energy, or the other way round. That little powerhouse is called an electrochemical cell. It’s the heart of everything from your phone’s battery to the corrosion you see on a rusty nail.
💡 In Simple Words: An electrochemical cell is a device where a chemical reaction either makes electricity (like a battery) or uses electricity to drive a reaction (like electroplating). Think of it as a two‑room house: one room gives away electrons, the other room takes them in, and a wire connects the rooms.
Two Main Types
- Galvanic (voltaic) cell: chemical energy → electrical energy. It powers devices.
- Electrolytic cell: electrical energy → chemical change. It’s used for things like plating jewelry.
How a Galvanic Cell Works
Picture a playground slide. The top is the anode, where oxidation (loss of electrons) happens. The bottom is the cathode, where reduction (gain of electrons) occurs. Electrons roll down the slide through a wire, lighting a bulb along the way.
The two half‑cells are linked by a salt bridge – think of it as a tiny hallway that lets ions move to keep the charge balanced, just like a hallway lets people move between rooms without crowding the slide.
Key Vocabulary (First Time)
- Anode: the electrode where oxidation occurs; electrons leave this side.
- Cathode: the electrode where reduction occurs; electrons arrive here.
- Oxidation: a process where a substance loses electrons.
- Reduction: a process where a substance gains electrons.
- Salt bridge: a tube filled with electrolyte that lets ions travel to maintain charge neutrality.
- Electrode potential: the tendency of a half‑cell to gain or lose electrons, measured in volts.
Worked Example: Calculating Cell EMF
Suppose we have a zinc‑copper galvanic cell. The half‑reactions are:
- Zn(s) → Zn²⁺(aq) + 2e⁻ (E° = –0.76 V)
- Cu²⁺(aq) + 2e⁻ → Cu(s) (E° = +0.34 V)
The cell potential (E°cell) = E°cathode – E°anode = 0.34 – (‑0.76) = 1.10 V. That means the cell can push 1.10 joules of energy per coulomb of charge that flows.
Electrolytic Cell in a Nutshell
Flip the playground slide upside down. Now you need to push electrons up the slide using a battery. That’s an electrolytic cell. It’s how we make aluminium from bauxite or coat a spoon with silver.
Comparison Table
| Feature | Galvanic Cell | Electrolytic Cell |
|---|---|---|
| Spontaneity | Spontaneous (energy released) | Non‑spontaneous (energy required) |
| Electrode where oxidation occurs | Anode (negative) | Anode (positive) |
| Electrode where reduction occurs | Cathode (positive) | Cathode (negative) |
| Typical use | Powering devices, batteries | Electroplating, electro‑refining |
| Cell notation | Zn|Zn²⁺||Cu²⁺|Cu | Cu|Cu²⁺||Zn²⁺|Zn |
Why It Matters for ISC Exams
Questions often ask you to identify the anode/cathode, write half‑reactions, or calculate the overall cell potential. Knowing the flow of electrons and the role of the salt bridge helps you sketch neat diagrams that earn marks.
Quick Summary
- Electrochemical cells turn chemical energy ↔ electrical energy.
- Galvanic cells generate electricity; electrolytic cells consume electricity.
- Electrons move from anode (oxidation) to cathode (reduction) through an external wire.
- Salt bridge keeps charge balanced by moving ions.
- Cell potential = E°cathode – E°anode.
📝 Likely Exam Questions
- Write the half‑reactions for a Zn–Cu galvanic cell and calculate its standard cell potential.
Answer: Anode: Zn → Zn²⁺ + 2e⁻ (‑0.76 V). Cathode: Cu²⁺ + 2e⁻ → Cu (0.34 V). E°cell = 0.34 – (‑0.76) = 1.10 V. - Explain why the anode is negative in a galvanic cell but positive in an electrolytic cell.
Answer: In a galvanic cell the anode loses electrons spontaneously, so it has a surplus of negative charge, making it negative. In an electrolytic cell an external battery forces electrons into the anode, giving it a positive charge. - Describe the purpose of a salt bridge and give one example of a material used.
Answer: The salt bridge allows ions to move between half‑cells, maintaining electrical neutrality. A common material is a KNO₃‑filled agar gel. - Predict what happens to the cell potential if the concentration of Cu²⁺ is increased.
Answer: Raising Cu²⁺ concentration makes the reduction half‑reaction more favorable, increasing the overall cell potential according to the Nernst equation. - Distinguish between a primary and a secondary battery with examples.
Answer: A primary battery (e.g., alkaline) cannot be recharged; a secondary battery (e.g., lead‑acid) can be recharged by applying external voltage.