Revision notes: Redox Equilibrium

Redox Equilibrium is one of the largest Form 5 chapters, and it rewards two skills above all: assigning oxidation numbers accurately, and writing correct half-equations. Everything else, displacement, rusting, electrolysis, cells and extraction, is built on those two ideas. These notes work through the chapter one content standard at a time so you can check your definitions against what the marking scheme expects.

9.1 Oxidation and reduction

There are three complementary ways to define oxidation and reduction, and the exam can test any of them.

• In terms of oxygen: oxidation is the gain of oxygen; reduction is the loss of oxygen.
• In terms of hydrogen: oxidation is the loss of hydrogen; reduction is the gain of hydrogen.
• In terms of electrons: oxidation is the loss of electrons; reduction is the gain of electrons.

Remember OIL RIG. Oxidation Is Loss, Reduction Is Gain, of electrons. This is the definition you use most in Form 5, so make it automatic.

A redox reaction is one in which oxidation and reduction happen together, because the electrons lost by one species are gained by another. The oxidising agent is the substance that causes oxidation and is itself reduced; the reducing agent causes reduction and is itself oxidised. Naming the agent correctly, and saying what happens to it, is a frequent source of easy marks.

9.2 Oxidation number and its changes

The oxidation number is a signed value assigned to an atom that tracks electron control. Learn these rules in order and apply them from the top:

• The oxidation number of an uncombined element is 0 (e.g. Cu, O₂, Cl₂).
• For a simple monatomic ion, it equals the ion charge (Na⁺ is +1, S²⁻ is −2).
• Group 1 metals are +1, Group 2 metals are +2, aluminium is +3.
• Hydrogen is usually +1 (but −1 in metal hydrides such as NaH).
• Oxygen is usually −2 (but −1 in peroxides such as H₂O₂).
• The sum of oxidation numbers in a neutral compound is 0; in a polyatomic ion it equals the ion charge.

Use these to find an unknown value: in SO₄²⁻, oxygen contributes 4 × (−2) = −8, so sulfur must be +6 for the total to be −2. When an atom’s oxidation number increases, it has been oxidised; when it decreases, it has been reduced. Oxidation numbers also drive Stock nomenclature, iron(II) means Fe with oxidation number +2, iron(III) means +3.

Exam habit. To identify a redox reaction, assign oxidation numbers to every atom on both sides and look for one that rises and one that falls. No change means no redox.

9.3 Redox in displacement reactions

A more reactive element displaces a less reactive one, and this is redox because electrons transfer.

• Metal displacement: a more reactive metal displaces a less reactive metal from its salt solution, e.g. zinc displaces copper from copper(II) sulfate. Zinc is oxidised (loses electrons); copper(II) ions are reduced (gain electrons).
• Halogen displacement: a more reactive halogen displaces a less reactive halide from solution, e.g. chlorine displaces bromine from potassium bromide.
• Electron transfer at a distance: if the oxidising and reducing halves are separated but joined by a wire and a salt bridge, electrons flow through the wire and can be detected with a galvanometer, proving that displacement is electron transfer.

Write each change as a half-equation: Zn → Zn²⁺ + 2e⁻ (oxidation) and Cu²⁺ + 2e⁻ → Cu (reduction). Electrons must balance when the half-equations are combined.

9.4 Rusting and its prevention

Rusting is the oxidation of iron to hydrated iron(III) oxide, and it needs both water and oxygen, remove either and rust does not form. Iron rusts faster when in contact with a less reactive metal (e.g. tin), and slower when in contact with a more reactive metal.

Prevention methods and their reasons:

• Barrier methods (paint, oil, grease, plastic, galvanising, tin-plating): keep out water and oxygen.
• Sacrificial protection: attach a more reactive metal such as zinc or magnesium. It corrodes in place of the iron because it loses electrons more readily.
• Galvanising works twice over: the zinc coat is a barrier and, if scratched, still protects the iron sacrificially.
• Alloying (stainless steel) resists corrosion by changing the metal’s surface chemistry.

9.5 Electrolysis of molten and aqueous compounds

Electrolysis decomposes an electrolyte using direct current. Cations move to the cathode (negative electrode) and are reduced; anions move to the anode (positive electrode) and are oxidised.

• Molten compounds: only the ions of the compound are present, so products are predicted directly. Molten lead(II) bromide gives lead at the cathode and bromine at the anode.
• Aqueous solutions: water provides H⁺ and OH⁻ as well, so selective discharge applies. Three factors decide which ion is discharged: position in the electrochemical series (the ion of the less reactive element is discharged more readily), concentration (a very concentrated halide ion may be discharged in preference to OH⁻), and type of electrode (inert carbon versus a reactive electrode such as copper, which dissolves).

Product prediction routine. List all ions present, take each electrode separately, apply the three factors, then write the half-equation. This routine earns full marks in electrolysis questions.

9.6 Voltaic and electrochemical cells

A voltaic (galvanic) cell converts chemical energy to electrical energy. Two metals of different reactivity dip into an electrolyte; the more reactive metal is the negative terminal (it is oxidised and releases electrons), and the less reactive metal is the positive terminal. Electrons flow through the external wire from the more reactive to the less reactive metal. The greater the difference in reactivity, the higher the voltage.

An electrolytic cell does the opposite, it uses electrical energy to drive a chemical change. Comparing the two (energy change, electrode signs, whether a reaction is spontaneous) is a standard exam task.

9.7 The reactivity series and metal extraction

The reactivity series ranks metals by how readily they lose electrons. Its position decides how a metal is extracted from its ore:

• Very reactive metals (potassium, sodium, calcium, magnesium, aluminium) are extracted by electrolysis of the molten compound, because they hold their oxide too strongly for carbon to remove.
• Moderately reactive metals (zinc, iron, tin, lead) are extracted by reduction with carbon.
• Unreactive metals (silver, gold) occur native or need only gentle treatment.

The thermite reaction, in which aluminium reduces iron(III) oxide, is a displacement that shows aluminium above iron in the series.

How to use these notes

Revise one standard at a time and test yourself aloud. Practise assigning oxidation numbers until it is automatic, then rehearse half-equations for displacement and electrolysis. A one-to-one teacher can check that your electrode signs and half-equations are consistent, the two places students most often slip in this chapter of SPM Chemistry. Then attempt the practice questions and mark your working against these definitions.