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Comparing and contrasting the two cell types

Why this page is the most important one

The NZQA clarification names comparing and contrasting the oxidation–reduction processes occurring in electrolytic and electrochemical cells as an Excellence requirement.
Everything in this topic exists to make this comparison possible. If you can do this page properly, you can reach Excellence.

The comparison table

	Electrochemical (galvanic)	Electrolytic
Reaction	spontaneous	non-spontaneous
$E°_{cell}$	positive	negative
Energy	produces electrical energy	consumes electrical energy
Power supply	none — the cell is the source	required, and must exceed $\lvert E°_{cell} \rvert$
Anode	oxidation, negative	oxidation, positive
Cathode	reduction, positive	reduction, negative
Electron flow (external)	anode → cathode	anode → cathode
Compartments	usually two, joined by a salt bridge	usually one
Salt bridge	needed	not needed
Electrolyte	one per half cell, matching each electrode	one, molten or aqueous
Examples	batteries, fuel cells, corrosion	electrolysis, electroplating, aluminium smelting

What is the SAME in both

Getting these right is what makes the contrast meaningful — a comparison that only lists differences misses half the point.

Oxidation happens at the anode. Reduction happens at the cathode. In both. Always.
Electrons flow from anode to cathode in the external circuit. In both.
Cations move towards the cathode; anions move towards the anode, in the solution. In both.
The half equations look identical — only the direction they are driven differs.
Charge must balance: electrons lost equal electrons gained. In both.

What is DIFFERENT

The sign of $E°_{cell}$ , and therefore whether energy is produced or consumed. This is the root difference; everything else follows from it.
The sign of the electrodes. In a galvanic cell the anode is negative because it is the source of electrons; in an electrolytic cell the anode is positive because it is connected to the positive terminal, which pulls electrons out.
Whether the reaction runs on its own. A galvanic cell runs until equilibrium is reached; an electrolytic cell stops the moment the power is switched off.
The role of the external circuit. In a galvanic cell the circuit is where the useful work is done. In an electrolytic cell it is where the energy is supplied.

The relationship between them

They are the same reaction run in opposite directions.
The clearest illustration is a rechargeable battery:
- Discharging — it is an electrochemical cell. $E°_{cell}$ is positive, the reaction is spontaneous, and electrical energy is delivered.
- Charging — it is an electrolytic cell. The external charger forces the reaction backwards against a negative $E°_{cell}$ , consuming energy.
- The same electrode is the anode in one mode and the cathode in the other, which is why battery terminals are labelled + and − rather than anode and cathode.
This single example contains the whole comparison and is worth building an Excellence answer around.

How to write the comparison

Do not just list differences. State a point of similarity, then the corresponding difference, then why they differ.
Anchor every point in chemistry. "The anode is positive in one and negative in the other" is a fact; "the anode is negative in a galvanic cell because it is the source of the electrons that flow into the circuit, and positive in an electrolytic cell because it is connected to the terminal that removes electrons" is a comparison.
Support with equations, a calculation and an observation — the criteria name all three, and a comparison carrying all three is a strong Excellence answer.

Worked ExampleA full compare-and-contrast answer

Compare and contrast the oxidation–reduction processes occurring in an electrochemical cell and an electrolytic cell, using the zinc–copper cell and the electrolysis of molten sodium chloride as examples. Support your answer with half equations, calculations and observations.

The two systems

Electrochemical: the zinc–copper cell

Anode: $\text{Zn}(s) \rightarrow \text{Zn}^{2+}(aq) + 2e^-$ , $E° = -0.76$ V
Cathode: $\text{Cu}^{2+}(aq) + 2e^- \rightarrow \text{Cu}(s)$ , $E° = +0.34$ V
$E°_{cell} = (+0.34) - (-0.76) = \mathbf{+1.10 \text{ V}}$

Electrolytic: molten sodium chloride

Anode: $2\text{Cl}^-(l) \rightarrow \text{Cl}_2(g) + 2e^-$ , $E° = +1.36$ V
Cathode: $\text{Na}^+(l) + e^- \rightarrow \text{Na}(l)$ , $E° = -2.71$ V
$E°_{cell} = (-2.71) - (+1.36) = \mathbf{-4.07 \text{ V}}$

What is the same

The location of each process is identical. In both cells, oxidation occurs at the anode and reduction occurs at the cathode. This is definitional, not coincidental — the terms are defined by the chemistry.

The direction of electron flow is identical. Electrons leave the species being oxidised at the anode, travel through the external circuit, and arrive at the species being reduced at the cathode. In the zinc–copper cell they flow from zinc to copper; in the electrolysis they flow from the chloride-oxidising anode to the sodium-reducing cathode.

The ion movement is identical. Cations migrate towards the cathode (Cu²⁺ towards the copper electrode; Na⁺ towards the cathode in the melt), and anions towards the anode.

The electron bookkeeping is identical. In both, electrons lost equal electrons gained, which is why half equations must be multiplied to match before combining.

What is different, and why

The sign of $E°_{cell}$ , which is the root difference. +1.10 V against −4.07 V. Everything else follows from this one fact.

Whether energy is produced or consumed. The positive $E°_{cell}$ of the zinc–copper cell means the reaction is spontaneous and releases energy, which can be drawn off as electrical work. The negative $E°_{cell}$ of the sodium chloride electrolysis means the reaction cannot occur on its own; a power supply delivering more than 4.07 V is required to force it, and it stops instantly if that supply is disconnected.

The sign of the electrodes, and the reason for it. In the galvanic cell the anode is negative, because it is the source of the electrons entering the circuit — electrons pile up there as zinc dissolves. In the electrolytic cell the anode is positive, because it is connected to the positive terminal, which pulls electrons out of the chloride ions. In both cases oxidation happens there; only what drives the electron removal differs.

The role of the external circuit. In the galvanic cell the circuit is where the useful work is extracted — the electrons do work on their way from zinc to copper. In the electrolytic cell the circuit is where the energy is put in.

The physical construction. The galvanic cell needs two separate compartments and a salt bridge, because letting Zn and Cu²⁺ touch would allow the electrons to transfer directly and release the energy as heat instead. The electrolytic cell needs only one compartment, since the reaction will not occur spontaneously even with the reactants in contact.

Observations, contrasted

Zinc–copper cell: the zinc electrode dissolves and loses mass; pink-brown copper deposits on the cathode; the blue colour fades as Cu²⁺ is consumed; the voltmeter reads about 1.10 V and falls gradually.

Sodium chloride electrolysis: silvery molten sodium collects at the cathode; pale yellow-green chlorine gas with a choking smell bubbles off at the anode; nothing happens at all when the power is disconnected.

That last contrast is the most revealing observation of the pair: one cell runs until its chemistry is exhausted, and the other stops the instant the energy supply is removed.

The unifying point

These are not two kinds of chemistry but one kind of chemistry driven in two directions. A rechargeable battery makes this explicit: discharging, it is an electrochemical cell with a positive $E°_{cell}$ delivering energy; charging, it is an electrolytic cell with a negative $E°_{cell}$ consuming it. The same physical electrode is the anode in one mode and the cathode in the other — which is precisely why battery terminals are marked + and − rather than anode and cathode.

Answer: both cells site oxidation at the anode and reduction at the cathode with electrons flowing anode to cathode; they differ in the sign of E°cell, and therefore in whether energy is produced or consumed, which in turn determines the electrode polarity and whether the reaction proceeds without an external supply.