Electrolysis Experiments

Molten electrolysis involves binary ionic compounds that have been heated until liquid, ensuring the only ions present are those from the compound itself. This makes the prediction of products straightforward as there is no competition between ions.

At the cathode, the metal cation is discharged to form the pure metal. For example, in molten lead bromide, lead ions ($Pb^{2+}$) gain electrons to form grey lead metal which deposits at the bottom of the electrode.

At the anode, the non-metal anion is discharged to form a non-metallic element, often a gas. In the lead bromide example, bromide ions ($Br^-$) lose electrons to form brown bromine gas ($Br_2$).

Aqueous solutions are more complex because they contain ions from the solute as well as $H^+$ and $OH^-$ ions from the partial dissociation of water. The specific ion discharged at each electrode depends on the Reactivity Series.

At the cathode, hydrogen gas is produced if the metal in the solution is more reactive than hydrogen (e.g., sodium, magnesium, or aluminum). If the metal is less reactive than hydrogen (e.g., copper, silver, or gold), the metal itself is produced.

At the anode, if a halide ion ($Cl^-$, $Br^-$, or $I^-$) is present, the corresponding halogen gas is formed. If no halide is present, the hydroxide ions (or water molecules) are discharged to produce oxygen gas.

Oxidation is defined as the loss of electrons, and it always occurs at the positive anode. For a chloride ion, the half-equation is $2Cl^- ightarrow Cl_2 + 2e^-$, indicating that two ions lose one electron each to form a neutral molecule.

Reduction is defined as the gain of electrons, and it always occurs at the negative cathode. For a copper ion, the half-equation is $Cu^{2+} + 2e^- ightarrow Cu$, where the ion gains two electrons to become a neutral metal atom.

To ensure a balanced description of the system, the number of electrons lost at the anode must equal the number of electrons gained at the cathode. This conservation of charge is a fundamental requirement for all electrochemical half-equations.

Strategy: When predicting products, always list all four ions present in an aqueous solution ($H^+$, $OH^-$, and the solute ions) before applying the reactivity rules. This prevents the common mistake of ignoring the contributions from water.

Gas Tests: Identifying products in the lab is crucial. Hydrogen gas is confirmed by a 'squeaky pop' with a lighted splint, oxygen by relighting a glowing splint, and chlorine by its ability to bleach damp blue litmus paper white.

Visual Checks: Metals produced at the cathode often have distinct appearances, such as the pink-brown deposit of copper. Bubbling or effervescence indicates the production of a gas at an electrode.