Galvanic cell vs electrolytic cell — complete comparison with cell diagrams

Question

Compare galvanic cells and electrolytic cells in terms of energy conversion, electrode polarity, spontaneity, and cell potential. Why is the anode negative in a galvanic cell but positive in an electrolytic cell?

(CBSE 12 boards ask this comparison almost every year; NEET 2023 tested electrode sign convention)

Solution — Step by Step

Galvanic cell: Chemical energy → Electrical energy. The reaction is spontaneous ( $\Delta G < 0$ , $E_{cell} > 0$ ). It generates current on its own.

Electrolytic cell: Electrical energy → Chemical energy. The reaction is non-spontaneous ( $\Delta G > 0$ , $E_{cell} < 0$ ). It needs an external power source to drive the reaction.

In both cells: oxidation happens at the anode, reduction at the cathode. This never changes.

But the polarity flips:

Galvanic cell: Anode is negative (electrons flow OUT from here), Cathode is positive
Electrolytic cell: Anode is positive (connected to + terminal of battery), Cathode is negative

In a galvanic cell, the anode spontaneously releases electrons (oxidation), making it the source of electrons — so it is negative.

In an electrolytic cell, the external battery forces electrons out of the anode. The positive terminal of the battery pulls electrons away from one electrode (making it positive = anode) and pushes them towards the other (making it negative = cathode).

A galvanic cell uses a salt bridge (or porous partition) to maintain electrical neutrality. The standard cell notation:

\text{Zn}(s) | \text{Zn}^{2+}(aq) || \text{Cu}^{2+}(aq) | \text{Cu}(s)

Anode on the left, cathode on the right, double line for the salt bridge. An electrolytic cell has no salt bridge — both electrodes are in the same solution.

flowchart LR
    subgraph Galvanic["Galvanic Cell"]
        GA["Anode (−)<br/>Oxidation<br/>Zn → Zn²⁺"] -->|"e⁻ flow through wire"| GC["Cathode (+)<br/>Reduction<br/>Cu²⁺ → Cu"]
        GA -.->|"Salt bridge"| GC
    end
    subgraph Electrolytic["Electrolytic Cell"]
        EA["Anode (+)<br/>Oxidation"] -->|"e⁻ forced by battery"| EC["Cathode (−)<br/>Reduction"]
        BAT["External Battery"] --> EA
        BAT --> EC
    end

Why This Works

The core principle is the same in both: oxidation at anode, reduction at cathode. The difference is whether the process is spontaneous or forced. A galvanic cell is like a ball rolling downhill (energy released), while an electrolytic cell is like pushing a ball uphill (energy consumed).

The sign convention confusion disappears if you remember: follow the electrons. In a galvanic cell, electrons naturally flow from the more reactive metal (anode, negative) to the less reactive metal (cathode, positive). In electrolysis, the battery reverses this natural tendency.

Alternative Method

Feature	Galvanic Cell	Electrolytic Cell
Energy conversion	Chemical → Electrical	Electrical → Chemical
Spontaneity	Spontaneous	Non-spontaneous
$E_{cell}$	Positive	Negative (before external voltage)
Anode sign	Negative (−)	Positive (+)
Cathode sign	Positive (+)	Negative (−)
Salt bridge	Required	Not needed
Example	Daniel cell (Zn-Cu)	Electrolysis of brine

Mnemonic for galvanic cell: “AN OX” and “RED CAT” — ANode is where OXidation occurs, REDuction at CAThode. This is true for both types of cells. Only the sign flips.

Common Mistake

Students memorise “anode is positive, cathode is negative” without specifying which cell type. This is only true for electrolytic cells. In galvanic cells, the signs are reversed. The safest approach: always remember that oxidation = anode, reduction = cathode (this never changes), and then determine the sign based on whether the process is spontaneous (galvanic) or forced (electrolytic).

Question

Solution — Step by Step

Why This Works

Alternative Method

Common Mistake

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