Half-equations and HT skills for electrolysis
Chemical changes • Electrolysis
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Key concepts
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Definition and purpose of a half-equation
A half-equation describes either the oxidation or the reduction that occurs at a single electrode. The equation shows the chemical species that lose or gain electrons and therefore isolates the electron transfer for that electrode. A complete cell reaction is the combination of the two half-equations, with electrons cancelling between them. A half-equation uses electrons explicitly to balance charge. Electrons appear on the product side for oxidation and on the reactant side for reduction. The presence and number of electrons reveal the change in oxidation state for the element involved.
Oxidation and reduction at the electrodes
Oxidation takes place at the anode because atoms or ions lose electrons there. Loss of electrons increases the positive charge on the species, so electrons appear on the product side of the half-equation. For example, chloride ions lose electrons at the anode to form chlorine atoms, which pair to give Cl2. Reduction takes place at the cathode because atoms or ions gain electrons there. Gain of electrons lowers the oxidation state and places electrons on the reactant side of the half-equation. For example, silver ions gain electrons to form silver atoms that deposit on the electrode.
Systematic steps to write and balance half-equations
Step 1: Identify the species undergoing change and its oxidation states before and after the reaction. Step 2: Write the basic skeletal ionic change without electrons. Step 3: Add electrons to the more positive side to balance charge. Step 4: Balance other atoms by adding H2O, H+ or OH− depending on the medium, and re-check charge. Step 5: Multiply half-equations if necessary so electrons cancel when combining. Charge balancing always determines the number of electrons. Atom balancing follows after electrons are placed. In aqueous solutions, addition of H+ and H2O balances hydrogen and oxygen; in alkaline solutions, OH− and H2O perform that role instead.
Aqueous versus molten electrolytes and limiting factors
Aqueous electrolytes supply H+ and OH− from water, which can undergo reduction or oxidation and therefore compete with dissolved ions. Presence of hydrogen or oxygen evolution changes the electrode half-equations from those that would occur in the molten state. For example, molten lead bromide gives Pb2+ reduction and Br− oxidation, whereas the aqueous solution of a salt can reduce water at the cathode if metal ions are less easily reduced. Electrode material also limits reactions. Inert electrodes (graphite, platinum) do not react and therefore allow only electrolyte species to be oxidised or reduced. Reactive electrodes can participate in electrode reactions and alter the half-equations by providing or consuming species.
HT skills: completing and balancing supplied half-equations
Completion tasks present partially written ionic changes that require insertion of electrons, H+, OH− or H2O to make the equation chemically and electrically balanced. Balancing tasks require checking atom counts and net charge and adjusting coefficients or adding appropriate species. Multiplying half-equations by integers ensures electron numbers match before combination. Consequence-based thinking aids completion: if charge is not equal on both sides, electrons provide the necessary change; if hydrogen or oxygen atoms are missing, H+ or H2O (or OH− in alkaline solutions) restore atom balance. Final verification requires that both mass and charge balance.
Key notes
Important points to keep in mind