Ions in Aqueous Solution (Edexcel International A Level Chemistry): Revision Note
Substitutions with Hydroxide & Ammonia
When transition metal ions in an aqueous solution react with aqueous sodium hydroxide and aqueous ammonia they form precipitates
However, some of these precipitates will dissolve in an excess of sodium hydroxide or ammonia to form complex ions in solution
The Reactions of Aqueous Transition Metal Ions with Aqueous Sodium Hydroxide
Transition Metal Ion | Metal-aqua ion | With NaOH (aq) | With excess NaOH (aq) |
|---|---|---|---|
Cr3+ | Green solution [Cr(H2O)6]3+ (aq) | Green precipitate Cr(OH)3(H2O)3 (s) | Green solution [Cr(OH)6]3– (aq) |
Fe2+ | Green solution [Fe(H2O)6]2+ (aq) | Green precipitate Fe(OH)2(H2O)4 (s) | No further change |
Fe3+ | Yellow-brown solution [Fe(H2O)6]3+ (aq) | Brown precipitate Fe(OH)3(H2O)3 (s) | No further change |
Co2+ | Pink solution [Co(H2O)6]2+ (aq) | Blue precipitate Co(OH)2(H2O)4 (s) | No further change |
Mn2+ | Pale pink solution [Mn(H2O)6]2+ (aq) | Pale brown precipitate Mn(OH)2(H2O)4 (s) | No further change |
Ni2+ | Green solution [Ni(H2O)6]2+ (aq) | Green precipitate Ni(OH)2(H2O)4 (s) | No further change |
Cu2+ | Blue solution [Cu(H2O)6]2+ (aq) | Blue precipitate Cu(OH)2(H2O)4 (s) | No further change |
Zn2+ | Colourless solution [Zn(H2O)6]2+ (aq) | White precipitate Zn(OH)2(H2O)4 (s) | Colourless solution [Zn(OH)4(H2O)2]2– (aq) |
Examples of ionic equations for the reactions in the table above
[Fe(H2O)6]2+ (aq) + 2OH- (aq) → [Fe(H2O)4(OH)2] (s) + 2H2O (l)
[Cu(H2O)6]2+ (aq) + 2OH- (aq) → [Cu(H2O)4(OH)2] (s) + 2H2O (l)
[Fe(H2O)6]3+ (aq) + 3OH- (aq) → [Fe(H2O)3(OH)3] (s) + 3H2O (l)
The Reactions of Aqueous Transition Metal Ions with Ammonia
Transition Metal Ion | Metal-aqua ion | With NH3 (aq) | With excess NH3 (aq) |
|---|---|---|---|
Cr3+ | Green solution [Cr(H2O)6]3+ (aq) | Green precipitate Cr(OH)3(H2O)3 (s) | Purple solution [Cr(NH3)6]3+ (aq) |
Fe2+ | Green solution [Fe(H2O)6]2+ (aq) | Green precipitate Fe(OH)2(H2O)4 (s) | No further change |
Fe3+ | Yellow-brown solution [Fe(H2O)6]3+ (aq) | Brown precipitate Fe(OH)3(H2O)3 (s) | No further change |
Co2+ | Pink solution [Co(H2O)6]2+ (aq) | Blue precipitate Co(OH)2(H2O)4 (s) | Yellow solution [Co(NH3)6]2+ (aq) |
Mn2+ | Pale pink solution [Mn(H2O)6]2+ (aq) | Pale brown precipitate Mn(OH)2(H2O)4 (s) | No further change |
Ni2+ | Green solution [Ni(H2O)6]2+ (aq) | Green precipitate Ni(OH)2(H2O)4 (s) | Dark blue solution [Ni(NH3)6]2+ (aq) |
Cu2+ | Blue solution [Cu(H2O)6]2+ (aq) | Blue precipitate Cu(OH)2(H2O)4 (s) | Dark blue solution [Cu(NH3)4(H2O)2]2+ (aq) |
Zn2+ | Colourless solution [Zn(H2O)6]2+ (aq) | White precipitate Zn(OH)2(H2O)4 (s) | Colourless solution [Zn(NH3)4]2+ (aq) |
Solutions of metal aqua ions react as acids with aqueous ammonia
This means that the ammonia solution will initially act as a base to remove one H+ ion per ammonia molecule used
Examples of ionic equations for this type of reaction from the table above:
[Fe(H2O)6]2+ (aq) + 2NH3 (aq) → Fe(H2O)4(OH)2 (s) + 2NH4+ (aq)
[Cu(H2O)6]2+ (aq) + 2NH3 (aq) → Cu(H2O)4(OH)2 (s) +2NH4+ (aq)
[Fe(H2O)6]3+ (aq) + 3NH3 (aq) → Fe(H2O)3(OH)3 (s) +3NH4+ (aq)
Some metal aqua ions will react further and undergo ligand substitution with excess ammonia
Two specific examples to be aware of are Co2+ and Cu2+ with excess ammonia:
Starting with the hexa aqua ions:
[Cu(H2O)6]2+ (aq) + 4NH3 (aq) → [Cu(H2O)2(NH3)4]2+ (aq) + 4H2O (l)
[Co(H2O)6]2+ (aq) + 6NH3 (aq) → [Co(NH3)6]2+ (aq) + 6H2O (l)
Starting with the precipitates:
[Cu(H2O)4(OH)2] (s) + 4NH3 (aq) → [Cu(H2O)2(NH3)4]2+ (aq) + 2H2O (l) + 2OH- (aq)
[Co(H2O)4(OH)2] (s) + 6NH3 (aq) → [Co(NH3)6]2+ (aq) + 4H2O (l) + 2OH- (aq)
Examiner Tips and Tricks
It is easiest to remember the formulas of the precipitates by remembering that the number of OH- ions substituted is the same as the value of the charge on the initial ion
Ionic Equations
Reaction with limited OH- and limited NH3
The bases OH- and ammonia when in limited amounts form the same hydroxide precipitates.
They form in deprotonation acid base reactions
For example, consider the reaction that occurs when aqeuous sodium hydroxide is added to copper(II) sulfate solution
[Cu(H2O)6]2+ (aq) + 2OH- (aq) → [Cu(H2O)4(OH)2] (s) + 2H2O (l)
This seems like a ligand substitution reaction - two hydroxide ions replacing two water molecules
However this is actually a deprotonation reaction - two hydroxide ions removing hydrogen ions from two of the water ligands converting them into water molecules
The two ligands that have lost hydrogen ions are now hydroxide ligand
Reaction with excess OH-
From above, we have seen how hydrated transition metal ions can be deprotonated by adding a base such as aqueous sodium hydroxide to form a metal hydroxide precipitate
For example
[Cr(H2O)6]3+ (aq) + 3OH- (aq) → [Cr(H2O)3(OH)3] (s) +3H2O (l)
When an excess of sodium hydroxide is added further deprotonation takes place
[Cr(H2O)3(OH)3] (s) + 3OH- (aq) → [Cr(OH)6]3- (aq) + 3H2O (l)
In this reaction, chromium(III) hydroxide acts as an acid, as it is reacting with a base
Chromium(III) hydroxide can also act as a base because it can react with acids as follows
[Cr(H2O)3(OH)3] (s) + 3H+ (aq) → [Cr(H2O)6]3+ (aq)
A metal hydroxide that can act as both an acid and a base is called an amphoteric hydroxide
This is an example of amphoteric behaviour
Reaction with excess NH3
With excess NH3 ligand substitution reactions occur with Cu, Co and Cr and their precipitates dissolve
The ligands NH3 and H2O are similar in size and are uncharged
Ligand exchange occurs without a change of co-ordination number for Co and Cr
For example when excess aqueous ammonia is added to a copper(II) hydroxide precipitate it dissolves forming a deep blue solution
[Cu(H2O)4(OH)2] (s) + 4NH3 (aq) → [Cu(NH3)4(H2O)2]2+ (aq) + 2H2O (l) + 2OH- (aq)
This is a ligand substitution - four ammonia molecules replace two water molecules and two hydroxide ions
In these reactions NH3 is acting as a Lewis base donating an electron pair
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