Electrolysis

What is electrolysis?

Electrolysis is the breaking down (decomposition) of an ionic compound using electricity. It only works when the compound is molten (melted) or dissolved in water, because the ions must be free to move and carry charge.

In a solid ionic compound the ions are locked in a fixed lattice, so the substance cannot conduct electricity. Melting it or dissolving it sets the ions free.

Key terms Electrolyte — the molten or aqueous ionic compound that conducts electricity and is broken down.

Electrodes — the two solid conductors (usually graphite or a metal) that carry the current into and out of the electrolyte.

Anode — the positive (+) electrode.

Cathode — the negative (−) electrode.

Cation — a positive ion (e.g. $Na^+$, $H^+$, $Cu^{2+}$).

Anion — a negative ion (e.g. $Cl^-$, $OH^-$, $Br^-$).

How ions move

Opposite charges attract, so:

Cations (positive ions) move to the cathode (negative electrode).
Anions (negative ions) move to the anode (positive electrode).

At the electrodes the ions gain or lose electrons and turn into neutral atoms or molecules. This is where the compound is actually broken down.

At the cathode, cations gain electrons — this is reduction.
At the anode, anions lose electrons — this is oxidation.

Exam tip Remember OIL RIG (Oxidation Is Loss, Reduction Is Gain) of electrons, and PANIC — Positive Anode, Negative Is Cathode. Cations go to the cathode (both start with the same idea: a cation is cathode-bound).

Electrolysis of molten compounds

When a molten ionic compound is electrolysed, the rule is simple: there are only two types of ion present, so the metal forms at the cathode and the non-metal forms at the anode.

Example: molten lead(II) bromide, $PbBr_2$

The melt contains $Pb^{2+}$ ions and $Br^-$ ions.

At the cathode (−), lead ions gain electrons:

$$Pb^{2+} + 2e^- \rightarrow Pb$$

A bead of molten lead forms at the bottom of the cathode.

At the anode (+), bromide ions lose electrons:

$$2Br^- \rightarrow Br_2 + 2e^-$$

Red-brown bromine vapour is given off.

Worked example Write the cathode half equation for electrolysing molten aluminium oxide.

Aluminium ions are $Al^{3+}$. At the cathode they gain 3 electrons each:

$$Al^{3+} + 3e^- \rightarrow Al$$

Electrolysis of aqueous solutions

Water complicates things. As well as the dissolved compound's ions, water itself provides a few $H^+$ and $OH^-$ ions. So at each electrode there is a choice of ion to discharge. Use these rules:

At the cathode (positive ions):

1. If the metal is more reactive than hydrogen (e.g. $Na^+$, $K^+$, $Ca^{2+}$, $Mg^{2+}$, $Al^{3+}$), then hydrogen gas is produced instead.

At the anode (negative ions):

1. If a halide ($Cl^-$, $Br^-$, $I^-$) is present in reasonable concentration, the halogen is produced.
2. Otherwise (e.g. sulfate, nitrate solutions), oxygen is produced from $OH^-$.

Watch out Concentration matters at the anode. Concentrated sodium chloride gives chlorine; very dilute sodium chloride tends to give oxygen. In exams, assume reasonably concentrated halide solutions give the halogen.

Worked examples of aqueous electrolysis

Dilute sulfuric acid (effectively electrolysing water)

Ions present: $H^+$, $OH^-$ (and $SO_4^{2-}$).

Cathode (−): hydrogen is the only positive ion that can discharge.

$$2H^+ + 2e^- \rightarrow H_2$$

Anode (+): sulfate is not discharged, so oxygen forms from hydroxide.

$$4OH^- \rightarrow O_2 + 2H_2O + 4e^-$$

You get twice the volume of hydrogen as oxygen — matching water, $H_2O$.

Copper(II) sulfate solution with inert (graphite/platinum) electrodes

Ions present: $Cu^{2+}$, $H^+$, $OH^-$, $SO_4^{2-}$.

Cathode (−): copper is less reactive than hydrogen, so copper is deposited (a pink/brown coating).

$$Cu^{2+} + 2e^- \rightarrow Cu$$

Anode (+): sulfate stays put, so oxygen is given off.

$$4OH^- \rightarrow O_2 + 2H_2O + 4e^-$$

The blue colour of the solution fades as $Cu^{2+}$ ions are removed.

Copper(II) sulfate solution with copper electrodes

If the electrodes are made of copper, the anode now dissolves instead of giving off oxygen:

Anode (+): copper dissolves into solution.

$$Cu \rightarrow Cu^{2+} + 2e^-$$

Cathode (−): copper is deposited.

$$Cu^{2+} + 2e^- \rightarrow Cu$$

So copper is transferred from the anode to the cathode. This is the principle behind purifying copper and electroplating.

Concentrated sodium chloride solution (brine)

Ions present: $Na^+$, $H^+$, $Cl^-$, $OH^-$. This is industrially very important.

Cathode (−): sodium is more reactive than hydrogen, so hydrogen forms.

$$2H^+ + 2e^- \rightarrow H_2$$

Anode (+): chloride is discharged in preference to oxygen.

$$2Cl^- \rightarrow Cl_2 + 2e^-$$

The $Na^+$ and $OH^-$ ions are left behind, so the solution becomes sodium hydroxide ($NaOH$).

Uses of electrolysis

Purifying copper. Impure copper is the anode and a thin sheet of pure copper is the cathode, in copper(II) sulfate solution. Copper dissolves from the impure anode and pure copper is deposited on the cathode. Impurities (such as gold and silver) drop to the bottom as anode sludge.

Electroplating. A thin layer of one metal is coated onto another to improve appearance or resist corrosion (e.g. silver-plated cutlery, chromium-plated taps). The object to be plated is the cathode, the plating metal is the anode, and the electrolyte contains ions of the plating metal.

Real world Electroplating is everywhere: rings plated with gold, bumpers plated with chromium, and "tin" cans which are steel electroplated with tin to stop rust.

Extracting aluminium (briefly). Aluminium is too reactive to extract by heating with carbon, so it is obtained by electrolysing molten aluminium oxide (dissolved in molten cryolite to lower the melting point). Aluminium forms at the carbon cathode; oxygen forms at the carbon anode and slowly burns the anodes away.

Exam tip For any electrolysis question, always: (1) list the ions present, (2) decide what's discharged at each electrode using the rules, (3) write balanced half equations with electrons. Marks are nearly always awarded for the half equations.