Metal-Aqua Ions & their Reactions with Bases (AQA A Level Chemistry): Video

Hi, I'm Eleanor with 3 years of experience teaching A-Level Chemistry, and this video is about metal-aqua ions and their reactions with bases.

Metal-aqua ions form when transition metal salts dissolve in water, and how acidic those ions are — set by the metal's charge density — is exactly what decides how they react with bases like hydroxide, ammonia and carbonate.

Metal-aqua ions form when a metal salt dissolves in water, with water molecules coordinating to the metal ion. These aqua ions can act as acids because the metal ion withdraws electron density from the oxygen-hydrogen bonds in the coordinated water molecules, making it easier for H plus ions to be released. Their acidity can vary and this affects how they react with bases such as hydroxide, ammonia and carbonate, often producing precipitates or, with excess ammonia, soluble complex ions.

So this video will cover: how metal-aqua ions form in the first place, what happens when they meet hydroxide and ammonia, and what happens with carbonate — including the amphoteric behaviour of aluminium hydroxide.

When a transition metal salt dissolves in water, the metal ion doesn't just float free — it forms a hexaaqua complex, with six water molecules acting as ligands, bonded to the metal by dative covalent bonds. Water here is behaving as a Lewis base by donating its lone pairs to the metal. The plus 3 ions, like iron(III) and aluminium, are noticeably more acidic in water than the plus 2 ions. That's because their higher charge density pulls more strongly on the surrounding water molecules, polarising them and making it easier for a proton to break away.

Add hydroxide or ammonia to any of these metal-aqua ions and you get a hydroxide precipitate. The base is simply removing protons from the water ligands. But the details differ by metal. Iron(II) gives a dark green precipitate that slowly turns orange-brown at the surface. Copper(II) gives a pale blue precipitate that, in excess ammonia, redissolves into a deep blue solution as ammonia substitutes for some of the water ligands. Aluminium and iron(III) both give simple precipitates — white for aluminium and red-brown for iron(III). Aluminium’s precipitate will redissolve to a colourless solution in excess hydroxide but iron(III)’s remains insoluble even in excess hydroxide or ammonia.

With carbonate ions, the plus2 metals behave simply — iron(II) and copper(II) just form their carbonate precipitate. Green for iron(II) and blue-green for copper(II). The plus 3 ions are different, because they're more acidic. The carbonate reacts with the acidic metal-aqua ion, releasing bubbles of carbon dioxide as the hydroxide precipitates out instead. Aluminium hydroxide has one more trick — it's amphoteric, meaning it shows both acidic and basic properties. Add dilute hydrochloric acid and it redissolves as the aluminium ion; add hot concentrated sodium hydroxide and it dissolves the other way, into the tetrahydroxoaluminate ion.

Two things worth remembering: plus 3 ions are more acidic, so they don't form a simple carbonate precipitate the way plus 2 ions do — expect fizzing instead. And this kind of proton-splitting from a water ligand is called a hydrolysis reaction.

So metals dissolve in water as hexaaqua ions, and the plus 3 ions are more acidic than the plus 2 ones. That acidity is exactly what shapes how they react with bases — simple precipitates form for plus 2 ions, fizzing carbonate reactions for plus 3 ions. And aluminium hydroxide's amphoteric nature follows from that same underlying chemistry.