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4.3 Electrolysis

4.3.1 The process of electrolysis

- When an ionic compound is melted or aqueous, the ions are free to move about within the liquid or solution.
- These liquids and solutions are able to conduct electricity and are called electrolytes.
        Some examples of electrolytes are molten lead bromide (PbBr₂) and aqueous sodium chloride (NaCl).
- Electrolysis is the process of using electricity to break down an ionic compound into its elements.
- Passing an electric current through electrolytes causes the ions to move to the electrodes.
- The positive ions (cations) move to the negative electrode (cathode) and the negative ions (anions) move to the positive electrode (anode).
- Ions are discharged at the electrodes producing elements (substances formed of just one type of atom).

Writing half equations

- Half equations show the reactions at each electrode during electrolysis.
- At the cathode, reduction occurs (gain of electrons).
- At the anode, oxidation occurs (loss of electrons).
- To write a half-equation: identify the ion at the electrode, decide whether it is reduced (gains e^-) or oxidised (loses e^-), add electrons to balance charge and then balance atoms.
- Electrons appear on the left for reduction (cathode) and on the right for oxidation (anode).
- Example – molten lead(II) bromide:
    Pb²⁺ + 2 e^- → Pb (cathode, reduction)
    2 Br^- → Br₂ + 2 e^- (anode, oxidation)
- Example – aqueous sodium chloride (NaCl):
    Cathode (water reduced): 2 H₂O + 2 e^- → H₂ + 2 OH^-
    Anode (chloride oxidised): 2 Cl^- → Cl₂ + 2 e^-
- When combining half-equations, multiply so the electrons cancel, then add the half-equations to give the overall reaction.

4.3.2 Electrolysis of molten ionic compounds

- When a simple ionic compound (e.g. lead bromide) is electrolysed in the molten state using inert electrodes, the metal (lead) is produced at the cathode and the non-metal (bromine) is produced at the anode.
- The metal ion is reduced at the cathode and the non-metal ion is oxidised at the anode.

4.3.3 Using electrolysis to extract metals

- Electrolysis is used if the metal is too reactive to be extracted by reduction with carbon or if the metal reacts with carbon.
- Large amounts of energy are used in the extraction process to melt the compounds and to produce the electrical current.
- For example, aluminium is extracted from its ore (bauxite) by electrolysis of molten aluminium oxide mixed with molten cryolite, using carbon as the anode material.
- The cryolite lowers the melting point of aluminium oxide, reducing energy costs.
- At the cathode, aluminium ions are reduced to aluminium atoms.
- The carbon anode must be replaced regularly as it reacts with the oxygen produced to form carbon dioxide (oxidises).

4.3.4 Electrolysis of aqueous solutions

- The ions discharged when an aqueous solution is electrolysed using inert electrodes depend on the relative reactivity of the elements involved.
- At the negative electrode (cathode), hydrogen is produced if the metal is more reactive than hydrogen; otherwise the metal is produced.
- At the positive electrode (anode), oxygen is produced unless the solution contains halide ions, in which case the halogen is produced.
- In aqueous solution water molecules break down to give hydrogen ions (H⁺) and hydroxide ions (OH^-), which may be discharged at the electrodes.
- To predict products for a solution containing a single ionic compound: list the ions present (metal or non-metal/halide plus H⁺ and OH^-), then use reactivity to decide whether hydrogen or the metal forms at the cathode and whether oxygen or a halogen forms at the anode.

Examples

Required practical 3: investigate what happens when aqueous solutions are electrolysed using inert electrodes.

- Aim: Investigate what happens when different aqueous solutions are electrolysed using inert electrodes and test a hypothesis about which species are discharged at each electrode.
- Example hypothesis: "For aqueous solutions with halide ions, the halogen will be produced at the anode; otherwise oxygen will be produced. At the cathode, hydrogen is produced if the metal is more reactive than hydrogen, otherwise the metal is deposited."
- Variables:
  - Independent: type of electrolyte (e.g. NaCl, CuSO₄, KBr).
  - Dependent: identity / volume of gas produced at each electrode, mass change of electrodes, gas rate (cm³ s^-1).
  - Controls: concentration, electrode material & area (inert electrodes: carbon/platinum), cell geometry, temperature, applied potential/current, time.
- Apparatus (suggested): beaker, inert electrodes (graphite or Pt), DC power supply, ammeter, gas syringes or inverted graduated tubes for gas collection, connecting wires, stopwatch, balance, thermometer, measuring cylinder, safety equipment (goggles, gloves, fume hood for halogens).
- Method (condensed):
  1. Prepare a 0.1–1.0 M solution of the chosen electrolyte and pour into the cell.
  2. Insert two inert electrodes connected to the DC supply; record electrode area and initial masses if measuring mass change.
  3. Connect ammeter and set a constant current (record current). Start timer and collect gases in gas syringes or inverted tubes over water above each electrode for a fixed time (e.g. 5–10 min).
  4. Record volumes of gas from each electrode and any visible deposits/discolouration. Repeat for other electrolytes.
  5. Optionally test collected gases (pop test for H₂, damp litmus/indicator for O₂, detect halogen smell/bleach test for Cl₂/Br₂ with care).
- Expected results / interpretation:
  - Cathode: If metal is more reactive than H, H₂ forms; if less reactive, metal ions are reduced to metal deposit (e.g. Cu²⁺ → Cu).
  - Anode: If solution contains halide ions (Cl^-, Br^-, I^-) the halogen is produced; otherwise O₂ is formed from OH^-/H₂O oxidation.
  - Example: Aqueous NaCl → H₂ (cathode) + Cl₂ (anode). CuSO₄ → Cu (cathode) + O₂ (anode).
- Data analysis / calculations:
  - Convert collected gas volumes to moles to compare amounts and calculate rates (mol s^-1).
  - Chlorine and bromine are toxic/irritant - work in a fume hood, avoid inhalation, wear PPE. Handle electricity safely, avoid short circuits, and dispose of solutions correctly.
- Reliability & improvements:
  - Repeat trials, use gas syringes for accurate volumes, measure and control current precisely, keep electrode area constant, run for longer times for better signal, use analytical tests to identify gases rather than just smell/colour.

4.3.5 Representation of reactions at electrodes as half equations

- Reactions at electrodes can be represented by half equations, for example:
    At the cathode: Cu²⁺ + 2 e^- → Cu
    At the anode: 2 OH^- → O₂ + 2 H₂O + 4 e^-