Electrolysis of Molten Compounds
Section: 4. Electrochemistry | Syllabus: Cambridge AS Level Physics 9702
What are Molten Compounds? Molten compounds are ionic compounds that have been heated until they melt. The heat breaks down the ionic lattice, allowing the ions to move freely. When an ionic compound is molten: The solid lattice structure breaks down Ions are free to move The compound can conduct electricity Electrolysis becomes possible Why Electrolyse Molten Compounds?
Molten electrolysis is simpler than aqueous electrolysis because: Only the compound's ions are present - no water molecules to complicate things Predictable products - you get the elements from the compound No competing reactions - unlike aqueous solutions where water can be electrolysed Main use: Extracting reactive metals that can't be obtained by reduction with carbon.
General Pattern for Molten Binary Compounds A binary ionic compound contains just two elements (e.g., NaCl, MgO, PbBr₂) At the cathode (negative electrode): Metal ions gain electrons → Metal forms Example: Na⁺ + e⁻ → Na At the anode (positive electrode): Non-metal ions lose electrons → Non-metal forms Example: 2Cl⁻ → Cl₂ + 2e⁻ Example 1: Molten Lead Bromide (PbBr₂) Setup: Electrolyte: Molten lead bromide (must be hot ~370°C) Electrodes: Inert graphite electrodes Ions present: Pb²⁺ and Br⁻ At the cathode: Pb²⁺ ions attracted to negative electrode Each Pb²⁺ gains 2 electrons Pb²⁺ + 2e⁻ → Pb Product: Grey molten lead metal At the anode: Br⁻ ions attracted to positive electrode Each Br⁻ loses 1 electron 2Br⁻ → Br₂ + 2e⁻ Product: Brown bromine gas (toxic fumes) Overall equation: PbBr₂ → Pb + Br₂ Example 2: Molten Sodium Chloride (NaCl) Setup: Electrolyte: Molten sodium chloride (must be hot ~800°C) Electrodes: Inert graphite or steel electrodes Ions present: Na⁺ and Cl⁻ At the cathode: Na⁺ ions attracted to negative electrode Each Na⁺ gains 1 electron Na⁺ + e⁻ → Na Product: Molten sodium metal (very reactive, silvery) At the anode: Cl⁻ ions attracted to positive electrode Each Cl⁻ loses 1 electron 2Cl⁻ → Cl₂ + 2e⁻ Product: Chlorine gas (greenish-yellow, toxic) Overall equation: 2NaCl → 2Na + Cl₂ Example 3: Molten Aluminium Oxide (Al₂O₃) Setup: Electrolyte: Molten aluminium oxide mixed with cryolite (lowers melting point from 2000°C to ~900°C) Electrodes: Graphite (carbon) Ions present: Al³⁺ and O²⁻ At the cathode: Al³⁺ ions attracted to negative electrode Each Al³⁺ gains 3 electrons Al³⁺ + 3e⁻ → Al Product: Molten aluminium metal (sinks to bottom) At the anode: O²⁻ ions attracted to positive electrode Each O²⁻ loses 2 electrons 2O²⁻ → O₂ + 4e⁻ Product: Oxygen gas Problem: Hot oxygen reacts with carbon anode: C + O₂ → CO₂ Anodes must be replaced regularly!
Overall equation: 2Al₂O₃ → 4Al + 3O₂ Example 4: Molten Magnesium Chloride (MgCl₂) At the cathode: Mg²⁺ + 2e⁻ → Mg Product: Magnesium metal At the anode: 2Cl⁻ → Cl₂ + 2e⁻ Product: Chlorine gas Overall equation: MgCl₂ → Mg + Cl₂ Predicting Products Simple rule for molten binary compounds: Cathode product: The metal from the compound Anode product: The non-metal from the compound Molten compound Cathode product Anode product PbBr₂ Pb (lead) Br₂ (bromine) NaCl Na (sodium) Cl₂ (chlorine) MgO Mg (magnesium) O₂ (oxygen) CaI₂ Ca (calcium) I₂ (iodine) Observations During Electrolysis Compound Cathode observation Anode observation PbBr₂ Grey molten lead forms Brown bromine gas/vapour NaCl Silvery molten sodium Greenish-yellow chlorine gas Al₂O₃ Silvery molten aluminium sinks Oxygen gas (anode burns away) MgCl₂ Silvery molten magnesium Greenish-yellow chlorine gas Industrial Extraction of Aluminium The extraction of aluminium is the most important industrial use of molten electrolysis: The Process: Aluminium ore (bauxite) is purified to get Al₂O₃ Al₂O₃ mixed with cryolite (Na₃AlF₆) to lower melting point Electrolysis at ~900°C instead of 2000°C (saves energy) Molten aluminium collects at the bottom (tapped off) Oxygen forms at carbon anodes (which burn away as CO₂) Why cryolite?
Reduces energy costs (lower temperature needed) Dissolves Al₂O₃ to make it conduct better Makes the process economically viable Why Can't We Use Carbon Reduction? For very reactive metals (like sodium, magnesium, aluminium), carbon reduction doesn't work: These metals are more reactive than carbon Carbon can't displace them from their oxides Electrolysis is the only option Reactivity series reminder: Potassium > Sodium > Calcium > Magnesium > Aluminium > Carbon > Zinc > Iron > Tin > Lead Metals above carbon must be extracted by electrolysis!
Balancing Electrons In electrolysis, the number of electrons lost must equal the number gained: Example: Magnesium chloride Cathode: Mg²⁺ + 2e⁻ → Mg (gains 2 electrons) Anode: 2Cl⁻ → Cl₂ + 2e⁻ (loses 2 electrons) Electrons balanced: 2 gained = 2 lost ✓ Example: Aluminium oxide Cathode: Al³⁺ + 3e⁻ → Al (gains 3 electrons per Al) Need 4 Al atoms: 4Al³⁺ + 12e⁻ → 4Al (gains 12 electrons total) Anode: 2O²⁻ → O₂ + 4e⁻ (loses 4 electrons per O₂) Need 3 O₂ molecules: 6O²⁻ → 3O₂ + 12e⁻ (loses 12 electrons total) Electrons balanced:…
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