Electrolytic Cells and Electrolysis

What is an electrolytic cell?

A galvanic cell lets a spontaneous redox reaction generate electricity on its own.

An electrolytic cell works the other way round: an external DC source is

connected across two electrodes, and the electrical energy supplied from

outside is what forces a reaction to occur — a reaction that would not

happen by itself. This is the basis of electrolysis, and it underlies a

large part of industrial and laboratory electrochemistry.

The simplest picture is two copper strips dipped into an aqueous solution of

copper sulphate, with a DC voltage applied across them.

• At the electrode connected to the negative terminal (the cathode), $\text{Cu}^{2+}$ ions from solution pick up electrons and deposit as metal:

• At the other electrode (the anode), metallic copper goes into solution as ions, releasing electrons:

A note on NCERT's printed form: the NCERT textbook prints this anode

reaction (its equation 2.29) with the copper ion labelled $\text{Cu}^{2+}(s)$.

That state symbol is a misprint — the copper dissolves off the anode as aqueous ions going into solution, so the correct label is $(aq)$, as written

above.

So copper is oxidised (dissolved) at the anode and reduced (deposited) at the

cathode. This simple pairing is exploited industrially: making an impure lump of copper the anode, it dissolves as current flows, while an equal

mass of pure copper deposits on the cathode. This is how blister copper is

refined to the very high purity needed for electrical wiring.

The same idea — using electrical energy to force a reduction that has no

convenient chemical reducing agent — is how several reactive metals are

produced on an industrial scale:

• Sodium and magnesium are obtained by electrolysing their fused (molten) chlorides.
• Aluminium is obtained by electrolysing molten aluminium oxide dissolved in cryolite.

These metals sit so high in reactivity that no ordinary chemical reductant can

pull their cations down to the metal — only supplied electrical energy can.

Faraday's laws of electrolysis

Michael Faraday was the first to put the quantitative side of electrolysis on

a firm footing. Working through the 1830s, he summarised his results as two

laws.

First Law — the amount of chemical change (reaction) occurring at an

electrode during electrolysis is directly proportional to the quantity of

electricity (charge) passed through the electrolyte, whether it is a solution

or a molten salt.

Second Law — when the same quantity of charge is passed through

different electrolytes, the masses of the different substances liberated at

the electrodes are proportional to their chemical equivalent weights, i.e.

Relating charge to moles of electrons

With a steady current, the charge passed is simply

where $Q$ is the charge in coulombs, $I$ is the current in amperes, and $t$ is

the time in seconds. (Before constant-current sources existed, Faraday

measured $Q$ indirectly, by connecting a "coulometer" — a standard

electrolytic cell — in series and weighing the silver or copper it

deposited.)

The charge required for an electrode reaction depends on its stoichiometry.

For example, reducing one mole of silver ions,

needs exactly one mole of electrons. Since the charge on a single electron is

$1.6021 \times 10^{-19}\ \text{C}$, the charge on one mole of electrons is

This fixed quantity of charge — the charge carried by one mole of electrons —

is called one Faraday, symbol $F$.

$1\ F = 96487\ \text{C mol}^{-1} \;\;(\text{used as} \approx 96500\ \text{C mol}^{-1}\ \text{for approximate work})$ …