Why are some aqueous metal-aqua ions more acidic than others, and how do their reactions with bases let us identify them?
The acidity of metal-aqua ions in terms of the charge density of the metal ion and the polarisation of coordinated water. The reactions of metal-aqua ions with bases such as sodium hydroxide and ammonia, and with carbonate ions. The amphoteric character of the aluminium hydroxide complex. The use of these reactions to identify metal ions in solution by the colours and behaviour of the precipitates formed.
A focused answer to the AQA A-Level Chemistry 3.2.6 specification points on reactions of ions in aqueous solution. Covers the acidity of metal-aqua ions and the link to charge density, the reactions of 2+ and 3+ aqua ions with sodium hydroxide, ammonia and carbonate, the amphoteric behaviour of aluminium hydroxide, and how the precipitate colours identify metal ions.
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What this dot point is asking
AQA wants you to explain the acidity of metal-aqua ions in terms of charge density, describe the reactions of 2+ and 3+ aqua ions with sodium hydroxide, ammonia and sodium carbonate, explain the amphoteric behaviour of aluminium hydroxide, and use the precipitate colours and behaviour to identify metal ions in solution.
Why aqua ions are acidic
Reactions with sodium hydroxide and ammonia (a small amount)
Adding a base removes and precipitates the neutral hydroxide. The colours identify the ion:
- : blue solution gives a blue precipitate, .
- : gives a green precipitate, , that darkens on standing.
- : gives a brown (rust) precipitate, .
- : gives a white precipitate, .
For a 3+ ion three protons are removed, for example:
Reactions in excess base
Reactions with carbonate ions
The behaviour with sodium carbonate distinguishes 2+ from acidic 3+ ions:
- 2+ ions are only weakly acidic, so they form an insoluble carbonate precipitate, for example (blue-green) and (green).
- 3+ ions are acidic enough to react with carbonate as an acid, releasing carbon dioxide (effervescence) and precipitating the hydroxide, for example (brown) or (white).
Identifying the ions: summary
Try this
Q1. Write an equation for the first acid dissociation of in water. [2 marks]
- Cue. .
Q2. State the colour of the precipitate formed when is added to a solution of . [1 mark]
- Cue. Brown (rust) .
Q3. Explain how you would distinguish from using sodium hydroxide. [2 marks]
- Cue. Both give a precipitate (white , brown ); only the white redissolves in excess .
Exam-style practice questions
Practice questions written in the style of AQA exam questions on this dot point, with worked answer explainers. The year tag is the paper they imitate, not the source.
AQA Paper 1 (style)3 marksExplain why a solution of iron(III) ions is more acidic than a solution of iron(II) ions.Show worked answer →
A 3-mark answer links charge density to the polarisation of water.
has a higher charge and a smaller radius than , so it has a greater charge density. This polarises (weakens) the O-H bonds in the coordinated water molecules more strongly, making it easier to release an ion. The solution therefore has a higher concentration of and is more acidic.
Markers reward higher charge density of , greater polarisation of coordinated water, and easier loss of .
AQA Paper 1 (style)3 marksDescribe what you would observe when sodium hydroxide solution is added dropwise, and then in excess, to a solution containing aluminium ions.Show worked answer →
A 3-mark answer needs the precipitate and its redissolving.
Adding a little gives a white precipitate of . Adding excess redissolves the precipitate to give a colourless solution of the aluminate ion, , because is amphoteric.
Markers reward white precipitate, dissolves in excess alkali, and the amphoteric explanation.
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Sources & how we know this
- AQA A-level Chemistry (7405) specification — AQA (2015)