Ever wondered how a tiny battery can light up a torch or charge your phone?
💡 In Simple Words: An electrochemical cell is a little box where a chemical reaction creates a flow of electrons. One side loses electrons (oxidation) and the other side gains them (reduction). The moving electrons become the electric current you can use.
What is an Electrochemical Cell?
An electrochemical cell is any device that changes chemical energy into electrical energy, or the other way round. Think of it as a rulebook that tells electrons where to go, just like a game guide tells players which moves are allowed.
Types of Electrochemical Cells
- Galvanic (or voltaic) cell: makes electricity on its own because the reaction happens spontaneously.
- Electrolytic cell: needs an outside power source to push a non‑spontaneous reaction.
How a Galvanic Cell Works
In a galvanic cell, the oxidation half‑reaction occurs at the anode (the negative side) and the reduction half‑reaction at the cathode (the positive side). Electrons travel through an external wire from anode to cathode, while ions move inside the cell to keep charge balanced.
Imagine water flowing through a pipe: the water leaves a high‑point (anode), travels through the pipe (wire), and enters a low‑point (cathode). The pipe’s pressure difference is like the cell’s voltage.
Worked Example: The Daniell Cell
Consider a classic Daniell cell: Zn|Zn²⁺||Cu²⁺|Cu. Zinc metal oxidizes, copper ions reduce.
Half‑reactions
- Oxidation: Zn → Zn²⁺ + 2e⁻ (E° = +0.76 V for the reverse, so we use –0.76 V)
- Reduction: Cu²⁺ + 2e⁻ → Cu (E° = +0.34 V)
Cell potential (E°cell) = E°cathode – E°anode = 0.34 V – (‑0.76 V) = 1.10 V. That 1.10 V is what you would measure with a voltmeter.
Electrolytic Cell Basics
In an electrolytic cell, we hook the cell up to a battery. The external power forces electrons to move opposite to their natural direction, so a reaction that wouldn’t happen on its own is forced to occur. A common example is the electrolysis of water, where electricity splits H₂O into H₂ and O₂.
Galvanic vs. Electrolytic Cells – Quick Comparison
| Feature | Galvanic Cell | Electrolytic Cell |
|---|---|---|
| Spontaneity | Reaction occurs spontaneously (ΔG | Reaction is non‑spontaneous (ΔG > 0) |
| Electron flow | From anode to cathode on its own | Forced by external power source |
| Energy source | Chemical energy → electrical | Electrical energy → chemical |
| Typical example | Daniell cell, dry cell | Electrolysis of water, metal plating |
Important Formulas and Exam Tips
- Cell potential (Ecell) = E°cathode – E°anode.
- If the reaction is not standard, use the Nernst equation: E = E° – (0.0592 V/n) log Q, where n = number of electrons transferred and Q = reaction quotient.
- Remember: oxidation always happens at the anode, reduction at the cathode – even in an electrolytic cell.
- For a salt bridge, choose ions that don’t react with the half‑cells (e.g., KNO₃).
📝 Likely Exam Questions
- Explain why the anode of a galvanic cell is negative.
Answer: At the anode oxidation occurs, releasing electrons into the external circuit. The buildup of electrons makes the anode negative relative to the cathode. - Calculate the standard cell potential for a cell with the half‑reactions: Fe³⁺ + e⁻ → Fe²⁺ (E° = +0.77 V) and Cu²⁺ + 2e⁻ → Cu (E° = +0.34 V).
Answer: The more positive reduction potential is Fe³⁺/Fe²⁺, so it acts as cathode. E°cell = 0.77 V – 0.34 V = 0.43 V. - What is the role of the salt bridge in a galvanic cell?
Answer: It allows ions to move between the two half‑cells, maintaining electrical neutrality and completing the circuit. - Describe one practical application of an electrolytic cell.
Answer: Electroplating uses an electrolytic cell to deposit a thin metal layer (e.g., chrome on car parts) by forcing metal ions onto a conductive surface. - Write the overall reaction for the Daniell cell and state its cell potential.
Answer: Zn(s) + Cu²⁺(aq) → Zn²⁺(aq) + Cu(s); E°cell = 1.10 V.