In simple terms
A friendly intro before the formal notes — no formulas yet.
Forcing Reactions with Electricity
Electrolysis uses a direct current power supply to force a non-spontaneous redox reaction to occur. This process is fundamental to splitting compounds and producing elements like chlorine and aluminium.
Imagine you have a ball at the bottom of a hill. It won't roll up on its own – that's a non-spontaneous process. Electrolysis is like using a conveyor belt (the power supply) to provide the energy needed to carry the ball (the reactants) to the top of the hill (to form the products).
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Electrolysis uses a direct current (DC) power supply to drive a non-spontaneous redox reaction at electrodes within an electrolyte.
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Positive cations move to the negative cathode for reduction (gain of electrons), while negative anions move to the positive anode for oxidation (loss of electrons).
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In molten electrolytes, only the compound's ions react. In aqueous solutions, water can also be oxidised (to O₂) or reduced (to H₂), so H⁺ and OH⁻ ions compete with the solute ions at the electrodes.
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Faraday's laws state that the mass of product is proportional to the total electric charge passed. We use Q = It and n = Q/F to calculate the moles of product formed.
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Key formulas
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$Charge (C) = Current (A) × Time (s) \ Q = I \times t$
$Amount of electrons (mol) = \frac{\text{Charge (Q)}}{\text{Faraday constant (F)}} \ n(e^{-}) = \frac{Q}{F} \quad \text{where F} \approx 96500 \mathrm{C mol}^{-1}$
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Full topic notes
Formal explanation with the rigour you need for the exam.
The Fundamentals of an Electrolytic Cell
An electrolytic cell consists of three main components. First, the electrolyte, which is an ionic compound in either a molten state or dissolved in a suitable solvent (usually water), providing mobile ions. Second, two electrodes, which are electrical conductors providing a surface for the reaction to occur. Finally, an external DC power supply (like a battery) which acts as an 'electron pump', driving electrons to one electrode and removing them from the other.
Cathode: The negative electrode. It is connected to the negative terminal of the power supply. Positive ions (cations) migrate here and are reduced (gain electrons).
Anode: The positive electrode. It is connected to the positive terminal of the power supply. Negative ions (anions) migrate here and are oxidised (lose electrons).
Electron Flow: Electrons flow from the anode, through the external circuit to the power supply, and from the power supply to the cathode. The ions carry the charge within the electrolyte itself.
Electrolysis of Molten Salts
The electrolysis of a molten binary ionic compound is the simplest case. Only two types of ions are present: the metal cation and the non-metal anion. There is no competition, so the products are predictable: the metal is formed at the cathode, and the non-metal is formed at the anode. For this to work, a high temperature is required to melt the ionic compound and overcome the strong electrostatic forces of the crystal lattice.
Electrolysis of Aqueous Solutions
When an ionic salt is dissolved in water, the situation becomes more complex. Water itself can be weakly ionised () and can also be oxidised or reduced. This means there is a competition at each electrode to determine which species is discharged.
- At the cathode: The metal cation competes with H⁺ ions (from water).
- At the anode: The non-metal anion competes with OH⁻ ions (from water).
Cation Discharge Rule: The cation with the more positive (less negative) standard electrode potential () is preferentially reduced. In practice: Cations of reactive metals (e.g., Na⁺, K⁺, Mg²⁺) are not discharged; H⁺ is reduced instead, producing H₂(g). Cations of less reactive metals (e.g., Cu²⁺, Ag⁺) are discharged in preference to H⁺.
Anion Discharge Rule: For concentrated solutions, halide ions (Cl⁻, Br⁻, I⁻) are discharged in preference to OH⁻. For dilute solutions or solutions containing complex ions (SO₄²⁻, NO₃⁻), OH⁻ is discharged, producing O₂(g) and H₂O.
Half-equation for water reduction (at cathode): 2H₂O(l) + 2e⁻ → H₂(g) + 2OH⁻(aq)
Half-equation for water oxidation (at anode): 2H₂O(l) → O₂(g) + 4H⁺(aq) + 4e⁻ OR 4OH⁻(aq) → O₂(g) + 2H₂O(l) + 4e⁻
Always include state symbols in your half-equations. They are crucial and often worth a mark. Pay special attention to (l) for molten substances, (aq) for aqueous species, and (g) or (s) for products. Forgetting them is a common way to lose easy marks.
Quantitative Electrolysis: Faraday's Laws
Faraday's laws allow us to calculate the amount of substance produced or consumed during electrolysis. The key principle is that the amount of chemical change is directly proportional to the quantity of electricity that passes through the electrolyte. This quantity of electricity, or charge, depends on the current and the time for which it flows.
The link between charge and the amount of substance in moles is the Faraday constant, F. One Faraday is the charge of one mole of electrons.
Worked examples
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Predict the products and write the half-equations for the electrolysis of molten lead(II) bromide (PbBr₂) using inert graphite electrodes.
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Identify ions present: In molten PbBr₂, the ions are Pb²⁺(l) and Br⁻(l).
Calculate the mass of copper deposited at the cathode during the electrolysis of aqueous copper(II) sulfate, when a current of 1.50 A is passed for 40.0 minutes. (Ar of Cu = 63.5)
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Convert time to seconds:
How it all connects
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Glossary
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Revision flashcards
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What is electrolysis?
The decomposition of a chemical compound (the electrolyte) by passing a direct electric current through it. It forces a non-spontaneous redox reaction to occur.
Key takeaways
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Cathode: The negative electrode. It is connected to the negative terminal of the power supply. Positive ions (cations) migrate here and are reduced (gain electrons).
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Anode: The positive electrode. It is connected to the positive terminal of the power supply. Negative ions (anions) migrate here and are oxidised (lose electrons).
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Electron Flow: Electrons flow from the anode, through the external circuit to the power supply, and from the power supply to the cathode. The ions carry the charge within the electrolyte itself.
Practice — then mark it
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Test Your Knowledge on Electrolysis
Test Your Knowledge on Electrolysis
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