Metal Aqua Ions: Acidity, Ligand Exchange and Precipitation
A metal ion in water is already a complex, and a highly charged one pulls hard enough on its water to make the solution acidic. Adding alkali, ammonia or an excess of either swaps ligands and drops or redissolves a hydroxide.
What a learner can do afterwards
- Explains why a 3 plus aqua ion gives a more acidic solution than a 2 plus one
- Writes the equation for a hydroxide precipitating when sodium hydroxide is added
- Predicts which precipitates redissolve in excess alkali and which redissolve in excess ammonia
- Identifies an unknown transition metal ion from the colours in a sequence of tests
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A metal ion in water is already a complex, and a 3+ ion packs more charge into its pull than a 2+ one. It drags harder on the water around it, weakens the O-H bonds, and frees more H+, so its solution is more acidic.
Add sodium hydroxide to copper(II) and blue copper hydroxide falls out: Cu2+ in water plus 2OH- gives Cu(OH)2 solid. Pour in concentrated HCl instead and the pale blue water complex swaps to the yellow-green chloro complex.
Some hydroxides redissolve in excess. Copper, nickel and zinc hydroxides clear in excess ammonia, with copper turning deep blue, while aluminium, chromium, zinc and lead hydroxides clear in excess alkali. Iron(II) green stays put in both.
A green precipitate that refuses to redissolve in excess alkali points to iron(II), since chromium(III) is green too but redissolves. Back it up with ammonia, which also gives green with no redissolve for iron(II).
Charge sets acidity, alkali drops hydroxides, excess sorts them, and green that stays put means iron(II).
More to read
14.4 Hydrolysis of Salt Solutions
A metal ion in water never sits bare. It arrives wrapped in six water molecules as an aqua ion like [Fe(H2O)6]3+, which is itself a complex. A highly charged small ion drags electron density off those waters, weakening their O-H bonds until protons drop off into solution. That is why 3+ ions such as Fe3+ and Al3+ make distinctly acidic solutions. Charge density rules the trend. A 3+ ion pulls far harder than a 2+ ion of similar size, so it releases more H+ and the pH lands lower. This same polarising power explains why some salt solutions are acidic: the cation is quietly acting as a weak acid through its water shell.
19.1 Occurrence, Preparation, and Properties of Transition Metals and Their Compounds
Drip sodium hydroxide into Cu2+ solution and hydroxide ligands muscle out the water: Cu2+(aq) + 2OH- becomes Cu(OH)2(s), a blue precipitate. Ammonia does the same at first, since it supplies OH- from its reaction with water. Precipitation is just ligand exchange where the new ligand happens to build an insoluble lattice. Excess reagent sorts the hydroxides out. With excess NaOH, only the amphoteric ones redissolve (Al3+, Cr3+, Zn2+, Pb2+). With excess ammonia, the ones that form stable ammine complexes redissolve instead (Cu2+ deep blue, Ni2+, Zn2+). Memorise the two lists separately and tests become lookups.
19.1: Properties of Transition Metals and Their Compounds
Unknown ions give themselves away by colour. Fe2+ drops a green precipitate that refuses to redissolve in excess alkali and darkens on top as air oxidises it. Cu2+ drops blue and redissolves in excess ammonia to an inky deep blue. Cr3+ drops green but redissolves in excess alkali to a dark green solution. Run the sequence the same way every time: NaOH dropwise, then excess, then fresh sample with ammonia dropwise and excess. A green precipitate that stays solid in excess alkali shouts Fe2+, confirmed if ammonia repeats the result. Two reagents, four observations, one name.
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8 questions wait behind this lesson, each with its answer explained. Every answer feeds the sky: stars light as they are learned, and dim when it is time to come back.