Electrical conductivity is the movement of charged particles.
In this case, charged particles means either delocalised electrons or ions.
These particles need to be free to move in a substance for that substance to be conductive.
Covalent compounds do not conduct electricity because there are no charged particles that are free to move.
Ionic compounds only conduct electricity only when molten or in solution.
When solid the ions are not free to move.
When molten or in solution the ions are free to move.
A negative ion is called an anion. Examples are the bromide ion (Br⁻) and the oxide ion (O²⁻).
A positive ion is called a cation. Examples are the sodium ion (Na⁺) and the aluminium ion (Al³⁺).
A trick to remember this is to write the ‘t’ as a ‘+’ in the word cation: ca+ion
Electrolysis: The breaking down of a substance caused by passing an electric current through an ionic compound which is molten or in solution. New substances are formed.
ELECTROLYSIS OF MOLTEN IONIC COMPOUNDS
Example: The electrolysis of molten lead bromide (PbBr2)
Solid lead bromide is heated and becomes molten. Explanation: ions become free to move.
Overall Reaction
word equation: lead bromide –> lead + bromine
chemical equation: PbBr2(l) –> Pb(l) + Br2(g)
Remember
OILRIG : Oxidation Is the Loss of electrons and Reduction Is the Gain of electrons
PANCAKE : Positive Anode, Negative Cathode
ELECTROLYSIS OF IONIC SOLUTIONS
Rules for working out elements formed from electrolysis of solutions
Follow these rules to decide which ions in solution will react at the electrodes:
At the cathode
Metal ions and hydrogen ions are positively charged. Whether you get the metal or hydrogen during electrolysis depends on the position of the metal in the reactivity series:
At the anode
At the anode, the product of electrolysis is always oxygen gas (O2) unless the solution contains a high concentration of Cl–, Br- or I– ions, in which case a halogen is produced, e.g. chlorine gas (Cl2), bromine gas (Br2), and iodine gas (I2).
The electrolysis of sodium chloride solution (NaCl(aq))
Solid sodium chloride is dissolved in water. Explanation: The sodium ions and chloride ions become free to move.
H2O(l) ⇋ H+(aq) + OH–(aq)
Electron half equation: 2Cl–(aq) –> Cl2 (g) + 2e–
Electron half equation: 2H+(aq) + 2e– –> H2 (g)
The electrolysis of copper sulfate solution (CuSO4(aq))
Electron half-equation: Cu2+(aq) + 2e– –> Cu (s)
Electron half-equation: 4OH–(aq) –> O2 (g) + 2H2O(l) + 4e–
The electrolysis of sulfuric acid (H2SO4(aq))
Electron half-equation: 2H+(aq) + 2e– –> H2(g)
Electron half-equation: 4OH–(aq) –> O2 (g) + 2H2O(l) + 4e–
2H2O(l) –> 2H2(g) + O2(g)
There are twice the amount (in moles) of H2 compared to O2
Oxidation: the loss of electrons or the gain of oxygen
Reduction: the gain of electrons or the loss of oxygen
Example: The electrolysis of lead (II) bromide, PbBr2
At the cathode (negative electrode): Pb2+ (l) + 2e– → Pb (l) (reduction)
At the anode (positive electrode): 2Br– (l) → Br2 (g) + 2e– (oxidation)
Example: The electrolysis of aluminium oxide, Al2O3
At the cathode: Al3+ + 3e– → Al (reduction)
At the anode: 2O2- → O2 + 4e– (oxidation)
Example: The electrolysis of sodium chloride solution (NaCl (aq))
At the cathode: 2H+ (aq) + 2e– → H2 (g) (reduction)
At the anode: 2Cl– (aq) → Cl2 (g) + 2e– (oxidation)
Example: The electrolysis of copper sulfate solution (CuSO4 (aq))
At the cathode: Cu2+ (aq) + 2e– → Cu (s) (reduction)
At the anode: 4OH– (aq) → O2 (g) + 2H2O (l) + 4e– (oxidation)
The diagram shows an electrolytic cell.
The electrolyte is an aqueous solution. For example it might be concentrated sodium chloride, NaCl (aq).
The test tubes over the electrodes must not completely cover them to make sure the ions are free to move throughout the solution.
In the case of NaCl (aq) bubbles of gas will be seen forming at the electrodes. These float up and collect in the test tubes when each gas can be tested to assess its identity.