,2.6. Reactions of Inorganic Compounds in Aqueous Solution
Lewis acids and bases
In the formation of complex ions the ligand is the Lewis
Definitions: Lewis acid: electron pair acceptor base because it is donating a pair of electrons in the
Lewis base: electron pair donator dative covalent bond and the metal ion is the Lewis acid.
Metal-aqua ions
Metal aqua ions are formed in aqueous solution.
In solution Fe(III) appears yellow/brown
[M(H2O)6]2+, limited to M = Fe (green) and Cu (blue); due to hydrolysis reactions. The violet
[M(H2O)6]3+, limited to M = Al (colourless), and Fe (violet) colour is only really seen in solid hydrated
salts that contain these complexes.
Acidity or hydrolysis reactions
The following equilibria happen in aqueous solutions of metal ions.
The equilibria lead to generation of acidic
[M(H2O)6]2+ + H2O [M(H2O)5(OH)]+ + H3O+ solutions with M3+ ions, and very weakly
acidic solutions with M2+ ions. The 3+ ions
[M(H2O)6]3+ + H2O [M(H2O)5(OH)]2+ + H3O+ are noticeably more acidic.
The acidity of [M(H2O)6]3+ is greater than that of [M(H2O)6]2+ because the 3+ metal ions have higher charge
density (charge/size ratio) and have greater polarising power. The greater the polarising power, the more
strongly it attracts the water molecule. This weakens the O-H bond so it breaks more easily releasing H+ ions.
Reaction with limited OH- and limited NH3
The bases OH- and ammonia when in limited
M(OH)2(H2O)4 (s) : Cu blue ppt, Fe (II) green ppt
amounts form the same hydroxide
M(OH)3(H2O)3 (s) : Fe(III) brown ppt, Al white ppt
precipitates. They form in deprotonation
acid base reactions.
[Cu(H2O)6]2+ (aq) + 2OH- (aq) Cu(H2O)4(OH)2 (s) + 2H2O (l) [Fe(H2O)6]2+ (aq) + 2NH3 (aq) Fe(H2O)4(OH)2 (s) + 2NH4+ (aq)
[Al(H2O)6]3+(aq) + 3OH- (aq) Al(H2O)3(OH)3 (s) + 3H2O (l) [Fe(H2O)6]3+ (aq) + 3NH3 (aq) Fe(H2O)3(OH)3 (s) + 3NH4+ (aq)
This process can happen step wise removing one proton at a time. Be able to write equations for this too.
e.g. [Al(H2O)6]3+ (aq) + OH- (aq) [Al(H2O)5(OH)]2+ (aq) + H2O (l)
Here the NH3 and OH- ions are acting as
[Al(H2O)5(OH)]2+ (aq) + OH- (aq) [Al(H2O)4(OH)2]+ (aq) + H2O (l) Bronsted-Lowry bases accepting a proton
Reaction with excess OH-
With excess NaOH the aluminium hydroxide dissolves. This hydroxide is classed as amphoteric
Al becomes [Al(OH)4]- (aq) colourless solution. because it reacts and dissolves in both acids and
bases.
Al(H2O)3(OH)3 (s) + OH- (aq ) [Al(OH)4]- (aq) + 3H2O (l) Al(H2O)3(OH)3 (s) + 3H+ (aq ) [Al(H2O)6]3+ (aq)
Reaction with excess NH3
With excess NH3 a ligand substitution reaction occurs with Cu and its precipitate dissolves to form a deep blue
solution.
This substitution is incomplete with Cu. Cu becomes [Cu(NH3)4(H2O)2]2+ deep blue solution
Cu(OH)2(H2O)4(s) + 4NH3 (aq) [Cu(NH3)4(H2O)2]2+ (aq) + 2H2O (l) + 2OH- (aq) In this reactions NH3 is
acting as a Lewis base
donating an electron pair.
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, Reactions with carbonate solution
The 2+ ions react differently to the 3+ ions with carbonate solutions.
The 2+ ions with carbonate solution results in MCO3 ppt being formed (Cu blue/green, Fe(II) green )
Cu2+ (aq) + CO32- (aq) CuCO3 (s) [Cu(H2O)6]2+ + CO32- CuCO3 + 6H2O These are
Fe2+ (aq) + CO32- (aq) FeCO3 (s) [Fe(H2O)6]2+ + CO32- FeCO3 + 6H2O precipitation
reactions
The 3+ ions with carbonate solution form a M(OH)3 ppt and CO2 gas is evolved.
Al forms white ppt of Al(OH)3 (H2O)3 + CO2
MCO3 is formed with 2+ ions but
Fe(III) forms brown ppt of Fe(OH)3 (H2O)3 + CO2
M2(CO3)3 is not formed with 3+ ions. The
difference is explained by the greater
2[Fe(H2O)6]3+(aq) + 3CO32-(aq) 2Fe(OH)3(H2O)3(s) +3CO2 + 3H2O(l) polarising power of the 3+ ion due to its
higher charge density.
2[Al(H2O)6]3+(aq) + 3CO32-(aq) 2Al(OH)3(H2O)3(s) +3CO2 + 3H2O(l)
These are classed as acidity reactions.
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Lewis acids and bases
In the formation of complex ions the ligand is the Lewis
Definitions: Lewis acid: electron pair acceptor base because it is donating a pair of electrons in the
Lewis base: electron pair donator dative covalent bond and the metal ion is the Lewis acid.
Metal-aqua ions
Metal aqua ions are formed in aqueous solution.
In solution Fe(III) appears yellow/brown
[M(H2O)6]2+, limited to M = Fe (green) and Cu (blue); due to hydrolysis reactions. The violet
[M(H2O)6]3+, limited to M = Al (colourless), and Fe (violet) colour is only really seen in solid hydrated
salts that contain these complexes.
Acidity or hydrolysis reactions
The following equilibria happen in aqueous solutions of metal ions.
The equilibria lead to generation of acidic
[M(H2O)6]2+ + H2O [M(H2O)5(OH)]+ + H3O+ solutions with M3+ ions, and very weakly
acidic solutions with M2+ ions. The 3+ ions
[M(H2O)6]3+ + H2O [M(H2O)5(OH)]2+ + H3O+ are noticeably more acidic.
The acidity of [M(H2O)6]3+ is greater than that of [M(H2O)6]2+ because the 3+ metal ions have higher charge
density (charge/size ratio) and have greater polarising power. The greater the polarising power, the more
strongly it attracts the water molecule. This weakens the O-H bond so it breaks more easily releasing H+ ions.
Reaction with limited OH- and limited NH3
The bases OH- and ammonia when in limited
M(OH)2(H2O)4 (s) : Cu blue ppt, Fe (II) green ppt
amounts form the same hydroxide
M(OH)3(H2O)3 (s) : Fe(III) brown ppt, Al white ppt
precipitates. They form in deprotonation
acid base reactions.
[Cu(H2O)6]2+ (aq) + 2OH- (aq) Cu(H2O)4(OH)2 (s) + 2H2O (l) [Fe(H2O)6]2+ (aq) + 2NH3 (aq) Fe(H2O)4(OH)2 (s) + 2NH4+ (aq)
[Al(H2O)6]3+(aq) + 3OH- (aq) Al(H2O)3(OH)3 (s) + 3H2O (l) [Fe(H2O)6]3+ (aq) + 3NH3 (aq) Fe(H2O)3(OH)3 (s) + 3NH4+ (aq)
This process can happen step wise removing one proton at a time. Be able to write equations for this too.
e.g. [Al(H2O)6]3+ (aq) + OH- (aq) [Al(H2O)5(OH)]2+ (aq) + H2O (l)
Here the NH3 and OH- ions are acting as
[Al(H2O)5(OH)]2+ (aq) + OH- (aq) [Al(H2O)4(OH)2]+ (aq) + H2O (l) Bronsted-Lowry bases accepting a proton
Reaction with excess OH-
With excess NaOH the aluminium hydroxide dissolves. This hydroxide is classed as amphoteric
Al becomes [Al(OH)4]- (aq) colourless solution. because it reacts and dissolves in both acids and
bases.
Al(H2O)3(OH)3 (s) + OH- (aq ) [Al(OH)4]- (aq) + 3H2O (l) Al(H2O)3(OH)3 (s) + 3H+ (aq ) [Al(H2O)6]3+ (aq)
Reaction with excess NH3
With excess NH3 a ligand substitution reaction occurs with Cu and its precipitate dissolves to form a deep blue
solution.
This substitution is incomplete with Cu. Cu becomes [Cu(NH3)4(H2O)2]2+ deep blue solution
Cu(OH)2(H2O)4(s) + 4NH3 (aq) [Cu(NH3)4(H2O)2]2+ (aq) + 2H2O (l) + 2OH- (aq) In this reactions NH3 is
acting as a Lewis base
donating an electron pair.
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, Reactions with carbonate solution
The 2+ ions react differently to the 3+ ions with carbonate solutions.
The 2+ ions with carbonate solution results in MCO3 ppt being formed (Cu blue/green, Fe(II) green )
Cu2+ (aq) + CO32- (aq) CuCO3 (s) [Cu(H2O)6]2+ + CO32- CuCO3 + 6H2O These are
Fe2+ (aq) + CO32- (aq) FeCO3 (s) [Fe(H2O)6]2+ + CO32- FeCO3 + 6H2O precipitation
reactions
The 3+ ions with carbonate solution form a M(OH)3 ppt and CO2 gas is evolved.
Al forms white ppt of Al(OH)3 (H2O)3 + CO2
MCO3 is formed with 2+ ions but
Fe(III) forms brown ppt of Fe(OH)3 (H2O)3 + CO2
M2(CO3)3 is not formed with 3+ ions. The
difference is explained by the greater
2[Fe(H2O)6]3+(aq) + 3CO32-(aq) 2Fe(OH)3(H2O)3(s) +3CO2 + 3H2O(l) polarising power of the 3+ ion due to its
higher charge density.
2[Al(H2O)6]3+(aq) + 3CO32-(aq) 2Al(OH)3(H2O)3(s) +3CO2 + 3H2O(l)
These are classed as acidity reactions.
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