AQA A-Level Chemistry: Reactions of ions in aqueous solution (7405)

This study guide covers section 3.2.6, Reactions of ions in aqueous solution, of the AQA AS and A-level Chemistry specification (7404/7405), version 1.2, July 2026, for AS and A-level exams from June 2016 onwards. The whole of 3.2.6 is marked “A-level only”, so it is not part of AS. The specification lists inorganic chemistry (section 3.2) under A-level Paper 1, and Paper 3 can assess any content. Section 3.2.6 includes required practical 11.

Use it with the revision notes and the practice questions. This topic builds on transition metals (ligands, complexes and ligand substitution), acids and bases (Brønsted–Lowry acids) and the amphoteric oxide in Period 3 elements and their oxides. For the rest of the course, see the AQA A-Level Chemistry hub and the printable checklist. For exam technique, see exam preparation, and find your weak spots with the free diagnostics.

What this section covers

Spec section	What you must be able to do	Status
3.2.6 Metal-aqua ions	Know the ions [M(H₂O)₆]²⁺ (M = Fe, Cu) and [M(H₂O)₆]³⁺ (M = Al, Fe)	A-level only
3.2.6 Acidity	Explain, using charge/size ratio, why [M(H₂O)₆]³⁺ is more acidic than [M(H₂O)₆]²⁺	A-level only
3.2.6 Amphoteric hydroxides	Show that some hydroxides (for example, aluminium’s) dissolve in both acids and bases	A-level only
3.2.6 Test-tube reactions	Describe and explain the reactions of Fe²⁺, Cu²⁺, Al³⁺ and Fe³⁺ (aq) with OH⁻, NH₃ and CO₃²⁻	A-level only
Required practical 11	Carry out simple test-tube reactions to identify transition metal ions in aqueous solution	A-level only

The specification limits you to these four ions and these three bases.

The metal-aqua ions

When a salt of one of these metals dissolves in water, six water molecules act as ligands. Each forms a co-ordinate bond from a lone pair on its oxygen atom to the metal ion, giving an octahedral complex.

Ion	Solution colour
[Fe(H₂O)₆]²⁺	pale green
[Cu(H₂O)₆]²⁺	blue
[Fe(H₂O)₆]³⁺	yellow-brown as usually seen
[Al(H₂O)₆]³⁺	colourless

Colour is a first clue only. Confirm the ion with a reagent.

Acidity of metal-aqua ions

A metal-aqua ion can give a proton to a water molecule. It acts as a Brønsted–Lowry acid:

[Fe(H₂O)₆]³⁺ + H₂O ⇌ [Fe(H₂O)₅(OH)]²⁺ + H₃O⁺
[Fe(H₂O)₆]²⁺ + H₂O ⇌ [Fe(H₂O)₅(OH)]⁺ + H₃O⁺

The H⁺ comes from one of the water ligands, so one H₂O ligand becomes an OH⁻ ligand and the charge on the complex falls by one.

Why 3+ ions are more acidic than 2+ ions

Use charge/size ratio in your answer.

1. The M³⁺ ion has a higher charge than the M²⁺ ion and is smaller, so it has a higher charge/size ratio (charge density).
2. It is therefore more polarising: it attracts electron density from the oxygen atoms of the water ligands more strongly.
3. This weakens the O–H bonds in the co-ordinated water.
4. So H⁺ is released more easily, the equilibrium lies further to the right, and the solution has a lower pH.

Worked example 1: comparing two iron solutions

Two solutions have the same concentration: iron(II) sulfate and iron(III) sulfate. Which has the lower pH, and why?

• Fe³⁺ has a 3+ charge and a smaller radius than Fe²⁺. Its charge/size ratio is greater.
• It polarises the co-ordinated water more, so the O–H bonds are weaker.
• [Fe(H₂O)₆]³⁺ releases H⁺ more readily, so iron(III) sulfate has the lower pH.

How bases make precipitates

All three bases in this section (OH⁻, NH₃ and CO₃²⁻) can remove H⁺ ions from the water ligands. Each H⁺ removed lowers the charge on the complex by one. When the complex reaches zero charge it is no longer attracted strongly to water molecules, so it comes out of solution as a precipitate.

• M²⁺ needs two H⁺ removed: [M(H₂O)₄(OH)₂]
• M³⁺ needs three H⁺ removed: [M(H₂O)₃(OH)₃]

Write these neutral hydroxides with square brackets and no charge. They can also be written M(OH)₂ and M(OH)₃.

Reactions with hydroxide ions, OH⁻

Add sodium hydroxide solution drop by drop, then in excess.

Ion	Few drops of NaOH(aq)	Excess NaOH(aq)
Fe²⁺	green precipitate, [Fe(H₂O)₄(OH)₂]; darkens to brown on standing in air	no further change
Cu²⁺	pale blue precipitate, [Cu(H₂O)₄(OH)₂]	no further change
Fe³⁺	brown precipitate, [Fe(H₂O)₃(OH)₃]	no further change
Al³⁺	white precipitate, [Al(H₂O)₃(OH)₃]	precipitate dissolves to give a colourless solution

The iron(II) precipitate darkens because air oxidises iron(II) to iron(III) at the surface.

Equations:

[Fe(H₂O)₆]²⁺ + 2OH⁻ → [Fe(H₂O)₄(OH)₂] + 2H₂O
[Cu(H₂O)₆]²⁺ + 2OH⁻ → [Cu(H₂O)₄(OH)₂] + 2H₂O
[Fe(H₂O)₆]³⁺ + 3OH⁻ → [Fe(H₂O)₃(OH)₃] + 3H₂O
[Al(H₂O)₆]³⁺ + 3OH⁻ → [Al(H₂O)₃(OH)₃] + 3H₂O

Amphoteric aluminium hydroxide

An amphoteric substance reacts with both acids and bases. Aluminium hydroxide is the example the specification gives.

With excess OH⁻, a fourth H⁺ is removed, giving a soluble negative ion:

[Al(H₂O)₃(OH)₃] + OH⁻ → [Al(H₂O)₂(OH)₄]⁻ + H₂O

The product is also written [Al(OH)₄]⁻.

With acid, the hydroxide ligands gain H⁺ and the aqua ion re-forms:

[Al(H₂O)₃(OH)₃] + 3H⁺ → [Al(H₂O)₆]³⁺

Worked example 2: hydroxide needed for aluminium

A solution contains 0.0150 mol of [Al(H₂O)₆]³⁺. How many moles of OH⁻ are needed (a) to precipitate all of it and (b) to redissolve the precipitate completely?

(a) Al³⁺ : OH⁻ = 1 : 3     n(OH⁻) = 3 × 0.0150 = 0.0450 mol
(b) [Al(H₂O)₃(OH)₃] : OH⁻ = 1 : 1   extra n(OH⁻) = 0.0150 mol
    total from the start = 0.0450 + 0.0150 = 0.0600 mol

Reactions with ammonia, NH₃

Ammonia does two jobs here.

• As a base, in small amounts, it removes H⁺ from water ligands. You see the same hydroxide precipitates as with NaOH.
• As a ligand, in excess, it replaces water ligands on copper, so the copper precipitate dissolves.

Ion	Few drops of NH₃(aq)	Excess NH₃(aq)
Fe²⁺	green precipitate	no further change
Cu²⁺	pale blue precipitate	precipitate dissolves to give a deep blue solution, [Cu(NH₃)₄(H₂O)₂]²⁺
Fe³⁺	brown precipitate	no further change
Al³⁺	white precipitate	no further change

As a base, ammonia becomes the ammonium ion:

[Fe(H₂O)₆]³⁺ + 3NH₃ → [Fe(H₂O)₃(OH)₃] + 3NH₄⁺
[Cu(H₂O)₆]²⁺ + 2NH₃ → [Cu(H₂O)₄(OH)₂] + 2NH₄⁺

In excess, ammonia replaces the hydroxide and some water ligands on copper:

[Cu(H₂O)₄(OH)₂] + 4NH₃ → [Cu(NH₃)₄(H₂O)₂]²⁺ + 2H₂O + 2OH⁻

The transition metals study guide shows the overall substitution from [Cu(H₂O)₆]²⁺ and explains why the co-ordination number stays at 6.

Aluminium hydroxide does not dissolve in excess ammonia. Ammonia is a weak base, so it does not remove the fourth H⁺ the way excess OH⁻ does.

Reactions with carbonate ions, CO₃²⁻

Add sodium carbonate solution. Here the 2+ and 3+ ions behave differently, and this is the part of the section you must explain, not just describe.

Ion	Na₂CO₃(aq)	Product
Fe²⁺	green precipitate	FeCO₃
Cu²⁺	blue-green precipitate	CuCO₃
Fe³⁺	brown precipitate and effervescence	[Fe(H₂O)₃(OH)₃] and CO₂
Al³⁺	white precipitate and effervescence	[Al(H₂O)₃(OH)₃] and CO₂

2+ ions are only weakly acidic. They are not acidic enough to protonate carbonate ions to carbon dioxide, so the metal carbonate precipitates:

[Fe(H₂O)₆]²⁺ + CO₃²⁻ → FeCO₃ + 6H₂O
[Cu(H₂O)₆]²⁺ + CO₃²⁻ → CuCO₃ + 6H₂O

3+ ions are more acidic, because of their higher charge/size ratio. They give up H⁺ to the carbonate ions, which form carbon dioxide and water. The metal ion is left as the neutral hydroxide. No iron(III) or aluminium carbonate forms.

Worked example 3: balancing the 3+ carbonate equation

Write the equation for [Al(H₂O)₆]³⁺ with carbonate ions.

Step 1: each Al³⁺ loses 3H⁺ to reach [Al(H₂O)₃(OH)₃]
Step 2: each CO₃²⁻ takes 2H⁺ to give CO₂ + H₂O
Step 3: 2 Al³⁺ release 6H⁺, which 3 CO₃²⁻ take up
2[Al(H₂O)₆]³⁺ + 3CO₃²⁻ → 2[Al(H₂O)₃(OH)₃] + 3CO₂ + 3H₂O

Check charge: left 2(+3) + 3(−2) = 0; right 0. The iron(III) equation has the same pattern.

Worked example 4: mass of carbonate precipitate

25.0 cm³ of 0.160 mol dm⁻³ copper(II) sulfate is added to excess sodium carbonate solution. Calculate the mass of CuCO₃ formed. (Ar: Cu = 63.5, C = 12.0, O = 16.0)

n(Cu²⁺) = 0.160 × 25.0/1000 = 4.00 × 10⁻³ mol
Cu²⁺ : CuCO₃ = 1 : 1
Mr(CuCO₃) = 63.5 + 12.0 + 3(16.0) = 123.5
mass = 4.00 × 10⁻³ × 123.5 = 0.494 g

Required practical 11

The specification’s wording is: carry out simple test-tube reactions to identify transition metal ions in aqueous solution.

A sound method:

1. Put about 1 cm³ of the unknown solution in a test tube. Note its colour.
2. Add sodium hydroxide solution one drop at a time, shaking after each drop. Record the colour of any precipitate.
3. Keep adding until the reagent is in excess, and record whether the precipitate dissolves.
4. Repeat steps 1 to 3 with fresh samples and aqueous ammonia, then sodium carbonate solution.
5. Leave any green iron precipitate to stand, and watch for darkening at the surface.

Safety: wear eye protection. Sodium hydroxide is corrosive, and ammonia solution is an irritant with a sharp smell, so use it in a well-ventilated room.

Fresh samples matter. If you add NH₃ to a tube that already holds NaOH, you cannot say which reagent caused what.

Worked example 5: identifying an unknown

Solution P is colourless. It gives a white precipitate with a few drops of NaOH(aq) that dissolves in excess. With Na₂CO₃(aq) it gives a white precipitate and bubbles.

• Colourless and a white precipitate: Al³⁺ (of the four ions in this section).
• Dissolves in excess NaOH: amphoteric hydroxide, so Al³⁺ again.
• Effervescence with carbonate: a 3+ ion.
• P contains [Al(H₂O)₆]³⁺.

A white precipitate with NaOH can also come from Group 2 ions such as Mg²⁺, which you meet in Group 2. Only the aluminium precipitate dissolves in excess NaOH.

A classification key for the four aqua ions

Each step uses one feature only.

1. (a) The solution is colourless → Al³⁺. (b) The solution is coloured → go to 2.
2. (a) The solution is blue → Cu²⁺. (b) The solution is not blue → go to 3.
3. (a) NaOH(aq) gives a green precipitate → Fe²⁺. (b) NaOH(aq) gives a brown precipitate → Fe³⁺.

Common errors

• Writing [Fe(H₂O)₆]³⁺ + 3OH⁻ → Fe(OH)₃ with no water. Either balance with 3H₂O and the full complex, or use the simple formulas consistently on both sides.
• Giving a charge to the neutral hydroxide, such as [Cu(H₂O)₄(OH)₂]²⁺.
• Saying iron(III) carbonate or aluminium carbonate precipitates. They do not form here.
• Forgetting the effervescence with 3+ ions and carbonate. “Bubbles” or “fizzing” is a separate observation.
• Saying the copper precipitate dissolves in excess NaOH. It dissolves in excess ammonia.
• Saying aluminium hydroxide dissolves in excess ammonia. It dissolves only in excess NaOH (or in acid).
• Explaining acidity with “Fe³⁺ is bigger”. It is smaller, with a greater charge.
• Describing a solution as “clear” when you mean “colourless”. A blue solution is clear too.

Official syllabus

AQA AS and A-level Chemistry specification (7404/7405), version 1.2, July 2026, for AS and A-level exams June 2016 onwards (AQA). Section 3.2.6 Reactions of ions in aqueous solution (A-level only), including required practical 11.