Electrolysis of Molten Salts
Electrolysis is a non-spontaneous redox process driven by an external power supply. In electrolysis of molten salts, the ionic compound is melted so ions can move freely and no water is present, so only the salt ions are discharged. The main quantity computed is the charge passed (and the resulting moles of electrons) to predict products and amounts at each electrode.
Core definitions and essential formulas
\[
Q = I \cdot t
\]
\[
n_{e^-} = \frac{Q}{F}
\]
\[
n_{\text{product}} = \frac{Q}{zF}
\]
\(Q\) is total charge (C), \(I\) is current (A), and \(t\) is time (s). \(F \approx 96485~\mathrm{C\,mol^{-1}}\) is the Faraday constant. \(z\) is the number of electrons needed per mole of product in the electrode half-reaction. In electrolysis, the cathode is negative (reduction of cations) and the anode is positive (oxidation of anions), so electrons are consumed at the cathode and produced at the anode.
How to interpret results
Larger \(I\) or longer \(t\) increases \(Q\), which increases \(n_{e^-}\) and therefore increases the moles (and mass) of products formed. Amounts are reported in moles, and mass follows from molar mass; gases can be related to volume if conditions are specified. Product identity depends on which ions are discharged: in a single molten salt, the only cation is reduced at the cathode and the only anion is oxidized at the anode; in a molten mixture, discharge is governed by relative electrode potentials and oxidation ease (for halides, \(\mathrm{I^-}\) is typically oxidized more readily than \(\mathrm{Br^-}\), which is more readily than \(\mathrm{Cl^-}\)).
- Wrong time units: minutes or hours must be converted to seconds for \(Q=I\cdot t\).
- Wrong \(z\): \(z\) comes from the balanced half-reaction, not the ionic charge alone.
- Including water chemistry: molten salts have no competing \(\mathrm{H_2O}\) or \(\mathrm{H^+}/\mathrm{OH^-}\) discharge.
- Electrode mix-up: cations reduce at the cathode (−) and anions oxidize at the anode (+).
Micro example: Molten \(\mathrm{MgCl_2}\), \(I=2.00~\mathrm{A}\), \(t=30.0~\mathrm{min}=1800~\mathrm{s}\).
\(Q=3600~\mathrm{C}\), so \(n_{e^-}=\dfrac{3600}{96485}=0.0373~\mathrm{mol}\).
For \(\mathrm{Mg^{2+}+2e^- \rightarrow Mg}\), \(z=2\) and \(n(\mathrm{Mg})=\dfrac{Q}{2F}=0.0186~\mathrm{mol}\).
Use this tool to predict electrode products and calculate amounts for molten ionic compounds and molten-salt mixtures using Faraday’s law. Avoid using it for aqueous electrolysis, where water and pH-dependent species can compete at the electrodes; a common next step is applying electrode potentials and the Nernst equation to compare discharge preferences under nonstandard conditions.
