3.2 Advanced Inorganic Chemistry · Year 13
3.2.6 Reactions of Ions in Aqueous Solution
Use acid-base and precipitation reactions of aqueous metal ions to predict observations and equations.
What you need to know
Open a line for a quick recap. If it feels obvious, move straight to the linked practice.
3.2.6 Know the course metal-aqua ions: [Fe(H₂O)6]2+, [Cu(H₂O)6]2+, [Al(H₂O)6]3+ and [Fe(H₂O)6]3+. Quick revision
Represent the course ions as hexaaqua complexes: [Fe(H₂O)₆]²⁺, [Cu(H₂O)₆]²⁺, [Al(H₂O)₆]³⁺ and [Fe(H₂O)₆]³⁺. The water molecules are ligands coordinated to the metal ion.
Writing the full aqua ion matters in this topic because the coordinated O–H bonds take part in acid–base reactions; a bare “Fe³⁺(aq)” formula hides the chemistry you are explaining.
3.2.6 Recognise that 3+ hexaaqua ions are more acidic than comparable 2+ ions. Quick revision
Compare similar hexaaqua ions: the 3+ ion is more acidic than the 2+ ion. Its higher charge density pulls electron density more strongly from the coordinated water molecules, making their O–H bonds easier to polarise and break as H⁺ is lost.
The O–H bonds become more polarised, so loss of H⁺ is easier. This is acidity of the coordinated water, not the metal ion somehow releasing a proton of its own.
3.2.6 Recognise amphoteric hydroxides, especially aluminium hydroxide, as able to dissolve in both acid and excess base. Quick revision
For aluminium hydroxide, use its behaviour to remember the word amphoteric: Al(OH)₃ reacts with acid and also dissolves in excess strong base. You need both sides of that behaviour.
If equations are requested, show the hydroxide reacting with both acid and base. Merely stating “amphoteric” is not enough. Keep amphoterism separate from the initial precipitation of Al(OH)₃.
3.2.6 Explain the greater acidity of 3+ metal-aqua ions using their larger charge-to-size ratio and stronger polarisation of O-H bonds. Quick revision
For the greater acidity of a 3+ aqua ion compared with a similar 2+ ion, follow the chain all the way through: higher charge density pulls more strongly on the O atoms of coordinated H₂O, polarising the O–H bonds and making H⁺ loss easier.
Do not stop at ‘the 3+ ion attracts more strongly’. The mark-earning chemistry is what that stronger attraction does to coordinated water: it polarises O–H bonds and makes proton loss easier.
3.2.6 Describe and explain the test-tube reactions of Fe2+, Cu2+, Al3+ and Fe3+ aqueous ions with OH⁻, NH₃ and CO3^2-, including observations and equations where required. Quick revision
For Fe²⁺, Cu²⁺, Al³⁺ and Fe³⁺, learn the observations and equations for adding OH⁻, NH₃ and CO₃²⁻ to the hydrated ions. OH⁻ commonly removes protons from coordinated water to give hydroxide precipitates; NH₃ can act as a base and, in excess for suitable ions, as a ligand.
I’d learn each ion as an observation + equation + excess-reagent behaviour package. With the 3+ aqua ions, keep their acidity in mind because carbonate can give CO₂ as well as a hydroxide precipitate.