Why do batteries power our phones and flashlights?
Ever wondered what actually makes a battery light up a LED? The secret lies in an electrochemical cell – a tiny factory that turns chemical energy into electric energy.
💡 In Simple Words: An electrochemical cell is a set‑up where a chemical reaction pushes electrons through a wire, giving you electricity. Think of it like a water pump: the reaction is the pressure that moves water (electrons) from one place to another.
Key Parts of a Cell
Every cell has four main players:
- Anode – the side where oxidation happens. Oxidation means a substance loses electrons, just like a person handing over money.
- Cathode – the side where reduction happens. Reduction is the opposite: a substance gains electrons, like receiving that money.
- Electrolyte – a liquid or gel that lets ions (charged atoms) move around. It’s the pipe that lets the water (ions) flow.
- Salt bridge or porous membrane – a path that lets ions travel between the two halves, keeping the charge balanced.
How It Works – Step by Step
Imagine a sandwich. The bread slices are the electrodes (anode and cathode) and the filling is the electrolyte. When the sandwich is built, a tiny reaction at the anode releases electrons. Those electrons jump onto the external wire, travel to the cathode, and complete the reaction there. Meanwhile, ions drift through the salt bridge so the whole system stays neutral.
Galvanic (Voltaic) Cell vs. Electrolytic Cell
Both are electrochemical cells, but they behave opposite ways. Use the table to keep them straight.
| Feature | Galvanic (Voltaic) Cell | Electrolytic Cell |
|---|---|---|
| Spontaneity | Reaction occurs on its own (spontaneous) | Reaction needs outside electricity (non‑spontaneous) |
| Energy Flow | Chemical → Electrical | Electrical → Chemical |
| Electrode Polarity | Anode = negative, Cathode = positive | Anode = positive, Cathode = negative |
| Common Use | Batteries, fuel cells | Electroplating, metal extraction |
Worked Example: Calculating Cell Potential
Let’s find the EMF (electromotive force) of a Daniell cell: Zn|Zn²⁺||Cu²⁺|Cu.
Step 1: Write the half‑reactions.
- Zn → Zn²⁺ + 2e⁻ (oxidation)
- Cu²⁺ + 2e⁻ → Cu (reduction)
Step 2: Look up standard reduction potentials (E°).
Cu²⁺/Cu = +0.34 V, Zn²⁺/Zn = –0.76 V.
Step 3: Use the formula E°cell = E°cathode – E°anode.
E°cell = (+0.34 V) – (–0.76 V) = +1.10 V.
Positive value tells us the reaction is spontaneous – exactly what a galvanic cell needs.
Quick Summary
- Electrochemical cell = chemical reaction that moves electrons through a wire.
- Anode = oxidation site, usually negative in a galvanic cell.
- Cathode = reduction site, usually positive in a galvanic cell.
- Salt bridge keeps charge balanced by letting ions travel.
- Galvanic cells give electricity; electrolytic cells need electricity.
- Cell potential (E°) predicts whether a reaction will run on its own.
📝 Likely Exam Questions
- Define an electrochemical cell and name its four essential components.
Answer: A device where a redox reaction produces electric current. Components: anode, cathode, electrolyte, and salt bridge (or porous membrane). - Write the overall reaction and calculate the standard cell potential for the cell: Fe|Fe²⁺||Cu²⁺|Cu.
Answer: Oxidation: Fe → Fe²⁺ + 2e⁻ (E° = –0.44 V). Reduction: Cu²⁺ + 2e⁻ → Cu (E° = +0.34 V). E°cell = 0.34 – (–0.44) = +0.78 V. - State two differences between a galvanic cell and an electrolytic cell.
Answer: (i) Galvanic cells operate spontaneously, electrolytic cells require external voltage. (ii) In galvanic cells the anode is negative; in electrolytic cells the anode is positive. - Explain why a salt bridge is necessary in a galvanic cell.
Answer: It allows ions to move between half‑cells, maintaining electrical neutrality; otherwise charge buildup would stop the flow of electrons.