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Inorganic reactions
24.1 The acid-base chemistry of aqueous transition metal Replacing water ligands, can happen with neutral,
ions: negative or chelates.
3+ ions are more acidic than 2+ ions, this is because With neutral ligands-
3+ ions are smaller and highly charged so it’s more [M(H2O)6]2+ + 6NH3 <—> [M(NH3)6]2+ + 6H2O
polarising and attracts the lone pair on the water With ammonia you need excess as it initially forms a
molecules, this causes a H+ ion to be released and so the precipitate as it acts like a base but it can re dissolve.
complex is acidic. 2+ ions are less polarising and so less Eg. Cobalt(II) and Copper(II)
O-H bonds break so less H+ ions are releases, making it
less acidic. Replacement with Cl, this will change the charge and the
coordination number as CL- s larger and negatively
Hydrolysis reactions, when aqua ions release H+ ions. charged.
[Fe(H2O)6]3+ <——-> [Fe(H2O)5(OH)]2+ + H+ [M(H2O)6]2+ + 4Cl- <—> [MCl4]2- + 6H2O
Lewis acid: an electron pair acceptor With chelates, are usually more stable due to the increase
Lewis base: an electron pair donator (atoms who dative in entropy.
bond) [M(H2O)6]2+ + 3en —> [M(en)3]2+ + 6H2O
[M(H2O)6]2+ + EDTA4- —> [MEDTA]2- + 6H2O
When adding base (OH- ions) it will remove protons
from the aqueous complex.
[M(H2O6]3+ + OH- —> [M(H2O)5(OH)]2+ + H2O
24.3 A summary of acid-base and substitution reactions
[M(H2O)5(OH)]2+ +OH- --> [M(H2O)4(OH)2]+
with metal ions:
+H2O
[M(H2O)4(OH)2]+ + OH- —> [M(H2O)3(OH)3]
Base [Fe(H2O)2+ [Cu(H2O)6]2+
+ H2O
As the [M(H2O)3(OH)3] had no charge it is insoluble OH- [Fe(H2O)4(OH)2]2+ [Cu(H2O)4(OH)2]2+
and so forms a precipitate. The same happens with 2+ = green = pale blue
ions until the charge is removed producing
Excess [Fe(H2O)4(OH)2]2+ [Cu(H2O)4(OH)2]2+
[M(H2O)4(OH)2] precipitates.
= green = pale blue
Carbonates of 3+ transition metals don’t exist as the NH3 [Fe(H2O)4(OH)2]2+ [Cu(H2O)4(OH)2]2+
carbonate ion removes protons to form aqua complexes = green = pale blue
whereas carbonates of 2+ transition metals do exist. Excess [Fe(H2O)4(OH)2]2+ [Cu(NH3)4(H2O)2]2+
2[Fe(H2O]63+ + 3CO32- —> 2[Fe(H2O)3(OH)3] +
=green =deep blue
3CO2 + 3H2O
CO32- FeCO3 CuCO3
To distinguish between ions, =green =blue-green
1. Add dilute alkali
2. A hydroxide of each transition metal is precipitated Base [Fe(H2O)6]3+ [Al(H2O)6]3+
and so colour differences are more obvious
OH- [Fe(H2O)3(OH)3 [Al(H2O)3(OH)3]
3. The [Fe(H2O)3(OH)3](s) is a brown colour, whilst
[Fe(H2O)4(OH)2](s) is green =brown =white
Excess [Fe(H2O)3(OH)3 [Al(OH)4]-
Aluminium hydroxide is amphoteric and so will act as =brown =colourless
an acid and a base. [Fe(H2O)3(OH)3 [Al(H2O)3(OH)3]
NH3
Al(H2O)3(OH)3 + 3HCl —> Al(H2O)63+ + 3Cl-
Al(H2O)3(OH)3 + OH- —> [Al(OH)4]- + 3H2O =brown =white
Excess [Fe(H2O)3(OH)3 [Al(H2O)3(OH)3]
=brown =white
24.2 Ligand Substitution reactions: CO32- [Fe(H2O)3(OH)3 [Al(H2O)3(OH)3]
=brown =white
Ligand substitution occurs because another ligand
will form a stronger dative bond with the metal ion,
causing the weaker ones to break. Or because the other
ligands are in higher concentration and so the
equilibrium is displaced.
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