Electrolysis and Its Applications

Lessons

Senior 3 · Chemistry

Electrolysis and Its Applications

5 lessons

Lesson 1 of 5

Electrolysis and Molten Sodium Chloride

25 min

Metals that chemistry alone could not free

In 1807 the English chemist Humphry Davy passed the current from a large battery through melted potash (potassium hydroxide). Small silvery globules appeared at the negative wire, and some burst into flame as they reached the air: the first potassium ever seen. A few days later he made sodium the same way. These metals are so reactive that no ordinary chemical can pull them out of their compounds, but electricity can. Today almost all the world's sodium, aluminium and chlorine is made by electrolysis.

Definition

Electrolysis: the decomposition of an ionic compound, when molten or dissolved in water, by passing a direct electric current through it.

Electrolyte: the molten or dissolved compound that conducts the current; its ions are free to move. Example: molten sodium chloride.

Electrodes: the rods that carry the current into and out of the electrolyte. The anode is the electrode joined to the positive terminal of the supply; the cathode is joined to the negative terminal. Inert electrodes, of graphite or platinum, do not react themselves.

Why the ions must be free to move

Solid sodium chloride does not conduct: its ions are locked in the crystal lattice. When it melts (at 801 °C) or dissolves, the ions can move, and the current is carried by ions drifting through the liquid. In electrolysis:

Cations (positive ions) are attracted to the cathode (negative). There they gain electrons.
Anions (negative ions) are attracted to the anode (positive). There they lose electrons.

In the wires the current is a flow of electrons, from the anode round to the cathode (Figure 1.1); in the electrolyte it is a flow of ions.

Half-equations

What happens at each electrode is shown by a half-equation, which includes the electrons, $\mathrm{e^-}$ . For molten sodium chloride:

$\text{cathode } (-): \quad \mathrm{Na^+(l) + e^- \rightarrow Na(l)}$

$\text{anode } (+): \quad \mathrm{2Cl^-(l) \rightarrow Cl_2(g) + 2e^-}$

The cathode reaction is a reduction (gain of electrons) and the anode reaction an oxidation (loss of electrons). Two electrons leave at the anode for every two that arrive at the cathode, so the overall equation is found by doubling the cathode half-equation and adding:

$\mathrm{2NaCl(l) \rightarrow 2Na(l) + Cl_2(g)}$

Sodium melts at 98 °C, so at the temperature of the cell it forms as a liquid.

Definition

Oxidation: loss of electrons. Reduction: gain of electrons. In electrolysis, reduction always happens at the cathode and oxidation at the anode. Example: $\mathrm{Na^+ + e^- \rightarrow Na}$ is a reduction.

Molten binary compounds

A molten compound of one metal and one non-metal (a binary compound) always gives the metal at the cathode and the non-metal at the anode.

Table 1.1: Products of electrolysing molten binary compounds (inert electrodes)

Molten electrolyte	Ions present	At the cathode (−)	At the anode (+)
sodium chloride, NaCl	$\mathrm{Na^+}$ , $\mathrm{Cl^-}$	sodium	chlorine
lead(II) bromide, $\mathrm{PbBr_2}$	$\mathrm{Pb^{2+}}$ , $\mathrm{Br^-}$	lead	bromine
zinc chloride, $\mathrm{ZnCl_2}$	$\mathrm{Zn^{2+}}$ , $\mathrm{Cl^-}$	zinc	chlorine
aluminium oxide, $\mathrm{Al_2O_3}$	$\mathrm{Al^{3+}}$ , $\mathrm{O^{2-}}$	aluminium	oxygen

Worked example

Problem. Molten zinc chloride is electrolysed with graphite electrodes. Write the half-equation at each electrode, say which is oxidation, and write the overall equation.

Solution. Cathode (−), reduction: $\mathrm{Zn^{2+}(l) + 2e^- \rightarrow Zn(l)}$

Anode (+), oxidation: $\mathrm{2Cl^-(l) \rightarrow Cl_2(g) + 2e^-}$

The electrons already balance (2 and 2), so the overall equation is $\mathrm{ZnCl_2(l) \rightarrow Zn(l) + Cl_2(g)}$ .

Worked example

Problem. Aluminium oxide contains $\mathrm{Al^{3+}}$ and $\mathrm{O^{2-}}$ ions. Write the half-equations for its electrolysis and balance the electrons to get the overall equation.

Solution. Cathode: $\mathrm{Al^{3+} + 3e^- \rightarrow Al}$ . Anode: $\mathrm{2O^{2-} \rightarrow O_2 + 4e^-}$ .

The lowest common multiple of 3 and 4 is 12, so multiply the cathode equation by 4 and the anode equation by 3:

$\mathrm{4Al^{3+} + 12e^- \rightarrow 4Al} \qquad \mathrm{6O^{2-} \rightarrow 3O_2 + 12e^-}$

Adding: $\mathrm{2Al_2O_3(l) \rightarrow 4Al(l) + 3O_2(g)}$ .

Warning

Common mistakes and safety. (1) In electrolysis the anode is the positive electrode: it is the one wired to the positive terminal. (2) Electrons flow in the wires, never through the electrolyte; in the liquid the charge is carried by ions. (3) In a half-equation, electrons appear on the left for reduction and on the right for oxidation. (4) Molten salts cause severe burns and chlorine and bromine are toxic, so molten electrolysis is a teacher demonstration in a fume cupboard, if it is done in school at all.

Tip

PANIC and OIL RIG. Positive is Anode, Negative Is Cathode. Oxidation Is Loss, Reduction Is Gain (of electrons). Cations go to the cathode.

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Practice

Define electrolysis, electrolyte and inert electrode.
Why does solid sodium chloride not conduct electricity, while molten sodium chloride does?
To which electrode do (a) cations, (b) anions move? Is each one oxidised or reduced there?
Write the half-equations, with state symbols, for the electrolysis of molten lead(II) bromide, and name the products.
Molten potassium iodide, KI, is electrolysed. Predict the product at each electrode and write the overall equation.
Copy and balance: $\mathrm{Mg^{2+} + \_e^- \rightarrow Mg}$ and $\mathrm{\_Cl^- \rightarrow Cl_2 + \_e^-}$ . What is formed at each electrode when molten magnesium chloride is electrolysed?
In Figure 1.1, which way do electrons flow in the wire joined to the anode? Explain.
A student says: "The negative ions carry electrons from the cathode to the anode through the liquid." Evaluate this statement.

Show answersHide answers

Electrolysis: decomposing a molten or dissolved ionic compound with a direct current. Electrolyte: the molten or dissolved compound that conducts and is decomposed. Inert electrode: one, such as graphite or platinum, that carries the current but does not react.
In the solid the ions are held in fixed positions in the lattice and cannot move; in the melt they are free to move and carry the charge.
(a) To the cathode, where they gain electrons: reduced. (b) To the anode, where they lose electrons: oxidised.
Cathode: $\mathrm{Pb^{2+}(l) + 2e^- \rightarrow Pb(l)}$ , lead. Anode: $\mathrm{2Br^-(l) \rightarrow Br_2(g) + 2e^-}$ , bromine (a brown vapour).
Potassium at the cathode, iodine at the anode: $\mathrm{2KI(l) \rightarrow 2K(l) + I_2(g)}$ .
$\mathrm{Mg^{2+} + 2e^- \rightarrow Mg}$ ; $\mathrm{2Cl^- \rightarrow Cl_2 + 2e^-}$ . Magnesium at the cathode, chlorine at the anode.
Away from the anode, towards the positive terminal of the supply. Chloride ions give up electrons at the anode, and the supply pulls them round the circuit and pushes them out of its negative terminal to the cathode.
Not correct. No electrons travel through the liquid. The ions themselves move: anions carry negative charge towards the anode, cations carry positive charge towards the cathode, and electrons are only exchanged at the electrode surfaces.

Sources and credits

Written by Ganzaa. The diagram, table, worked examples and Practice items were drawn and written by Ganzaa.

1Metals that chemistry alone could not free

2Why the ions must be free to move

3Half-equations

4Molten binary compounds

Metals that chemistry alone could not free

Why the ions must be free to move

Half-equations

Molten binary compounds