Part 1: Aqua-ion acidity and hydroxide reactions

Explain why metal aqua ions are acidic, write deprotonation equations and distinguish precipitation from amphoteric dissolution.

A dissolved metal ion has water ligands

For this topic, represent the specified ions as [Cu(H₂O)₆]²⁺, [Fe(H₂O)₆]²⁺, [Fe(H₂O)₆]³⁺ and [Al(H₂O)₆]³⁺. Each water donates a lone pair from oxygen, giving six coordinate bonds in an octahedral arrangement; copper(II) is distorted octahedral. Short formulae such as Fe²⁺(aq) are useful shorthand, but a question about water-ligand acidity requires the complete complex.

Typical dilute laboratory copper(II) solutions are pale blue, iron(II) pale green and aluminium colourless. Iron(III) solutions commonly look yellow/brown because hydrolysed species and other ligands can be present; an isolated hexaaquairon(III) complex is pale violet. Use the actual conditions and observations rather than assuming every iron(III) solution has one fixed colour.

Aluminium is included in this section even though it is not a transition metal. The chemistry here is about charge density, coordinated water and acid–base reactions, not a requirement that every central ion has d electrons.

The metal ion makes an O–H bond easier to deprotonate

A small, highly charged metal ion attracts electron density from coordinated water. This polarises and weakens the water ligand’s O–H bond, making transfer of H⁺ to a neighbouring water molecule easier. The metal centre itself is not releasing a proton. A water ligand loses H⁺ to become an OH⁻ ligand; the metal oxidation state does not change.

For comparable ions, the 3+ charge gives a greater charge-to-size ratio and more acidic aqua ions than the 2+ charge. Explain the link all the way to O–H polarisation and proton loss. Size matters as well as charge; “three plus is more acidic” on its own is an incomplete causal explanation.

Hydroxide removes protons in stages

OH⁻ accepts a proton from a coordinated water molecule to form water. For a 3+ aqua ion, each deprotonation reduces the complex charge by one. After three deprotonations the neutral hydrated hydroxide precipitates. In this A-level representation, six O donor sites remain around the metal as water ligands become hydroxide ligands.

Small additions of hydroxide produce four useful precipitates

Write the neutral solid without an overall ionic charge. The hydrated formulae below make the water-ligand chemistry explicit; the simpler forms Cu(OH)₂, Fe(OH)₂, Fe(OH)₃ and Al(OH)₃ are often used as shorthand. Do not mix the two representations within an equation without rebalancing the waters.

Aluminium hydroxide reacts with acid and excess base

Amphoteric aluminium hydroxide dissolves in excess OH⁻ to give [Al(OH)₄]⁻, a colourless solution. Its white precipitate also dissolves in acid as hydroxide ligands are protonated back to water ligands. Aluminium stays +3 throughout.

In the shorthand representation, Al(OH)₃ + OH⁻ → [Al(OH)₄]⁻. In the hydrated representation used below, three water molecules are additionally released. These describe the same stoichiometric aluminium conversion using different conventions. The precipitation-stage coordination number is six; the tetrahydroxoaluminate product has four donor bonds.

Worked example: precipitation is not the final excess-base product

Constructed stoichiometric model: a sample contains 0.500 mmol Al³⁺ aqua ions, with no additional acid consuming hydroxide. Forming the hydroxide precipitate requires 3 × 0.500 = 1.50 mmol OH⁻. Converting all that precipitate to [Al(OH)₄]⁻ requires a further 0.500 mmol OH⁻. The total is therefore 2.00 mmol.

With 0.200 mol dm⁻³ NaOH, the idealised precipitation requirement is 7.50 cm³ and the additional dissolution requirement is 2.50 cm³, giving 10.0 cm³ overall. Actual observed boundaries depend on equilibria, concentrations and other acid present; these numbers are stoichiometric quantities, not a universal practical titration curve.

Quick checks

Original Finesse questions. Reveal the indicative worked solutions after attempting each question; these are not official AQA mark allocations.

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