Download Hydrolysis Reactions

REACTIONS OF INORGANIC COMPOUNDS
IN AQUEOUS SOLUTION
Lewis Acids & Bases
In 1938, Lewis proposed a definition of acids and bases which is
much broader than the earlier definition by Bronsted and Lowry.
A Lewis acid is an electron pair acceptor.
A Lewis base is an electron pair donor.
Gilbert N. Lewis
The Lewis definition subsumes the Bronsted and Lowry definition,
but in addition it enables reactions which do not involve protons to
be classed as acid-base reactions.
For example, when ammonia reacts with boron trifluoride, the nitrogen atom donates an
electron pair to boron, forming a co-ordinate bond; ammonia is acting as a Lewis base
and boron trifluoride as a Lewis acid.
NH3 + BF3
[H3N:
BF3]
In the chemistry of transition metals, ligands are electron pair donors and are therefore
Lewis bases; metal cations are electron pair acceptors and are therefore Lewis acids.
As a result of the electron pair donation, a co-ordinate bond is formed between the
ligand and the metal.
Metal-aqua Ions
A water molecule has two lone pairs of electrons on the oxygen atom, so therefore
water can behave as an electron pair donor (Lewis base or ligand). When a transition
metal ion is placed in water, it becomes surrounded by six water molecules, each of
which donates one electron pair to the metal ion, forming a hexaaqua complex ion. This
reaction results in a colour change, for example:
CuSO4
CuBr2
CuCl2
(white)
(black)
(yellow)
CoCl2
CoBr2
(blue)
(green)
+ water
+ water
[Cu(H2O)6]2+
(blue)
[Co(H2O)6]2+
(pink)
The anion forms a colourless hydrated ion (SO42-(aq), Br-(aq), Cl-(aq)) and plays no part
in the formation of the colour.
The colour changes when water is added to cobalt chloride paper (blue to pink) or
anhydrous copper(II) sulphate (white to blue) are used as tests for water.
TOPIC 13.20: HYDROLYSIS REACTIONS 1
H
H
2+
..
O
H
:O
O:
H
H
H
Fe
O:
:O
H
H
..
H
[Fe(H2O)6]2+ is green
H
[V(H2O)6]3+ is grey-blue
O
H
H
Hexaaqua ions have octahedral geometry.
When a transition metal salt is crystallised from aqueous solution, the hexaaqua ion
frequently remains intact in the crystal, demonstrating the strength of the co-ordinate
bonding between the metal ion and water molecules. For example:
FeSO4.7H2O is green and contains the ion [Fe(H2O)6]2+
CoCl2.6H2O is pink and contains the ion [Co(H2O)6]2+
Fe(NO3)3.9H2O is pale violet and contains the ion [Fe(H2O)6]3+
Cr2(SO4)3.6H2O is red-violet and contains the ion [Cr(H2O)6]3+
When an anhydrous metal salt dissolves in water, the energy released on hydration of
the metal ion is enough to overcome the lattice enthalpy. In the case of metal(II)-aqua
ions, these energy changes are comparable in size and no significant temperature
change takes place. However, the higher charge density on metal 3+ ions results in a
more exothermic energy change on hydration; the reactions which form metal(III)-aqua
ions are therefore often violent and release a lot of heat. For example,
AlCl3(s) + aq.
[Al(H2O)6]3+ + 3Cl-(aq)
colourless
In the case of metal(IV) compounds, an exothermic hydrolysis reaction takes place and
the metal hydroxide is often precipitated. For example, when TiCl 4 is added to water,
Ti(OH)4 is precipitated.
TOPIC 13.20: HYDROLYSIS REACTIONS 2
Reactions of Metal-aqua Ions
H
H
n+
..
O
H
:O
O:
H
H
H
M
O:
:O
H
H
H
H
..
O
H
H
The octahedral metal-aqua ions contain only two types of bond:
 O-H bonds in the co-ordinated water molecules (
)
 M-O co-ordinate bonds between the metal ion and the water ligands (
)
There are thus only three types of reaction which metal-aqua ions can undergo:
Hydrolysis or acidity reaction
This reaction occurs when the O-H bond in a co-ordinated water molecule is broken,
releasing a proton (H+(aq)).
[Fe(H2O)6]3+ + H2O
[Fe(H2O)5(OH)]2+ + H3O+
Substitution reaction
When the M-O co-ordinate bond is broken, the water ligand is lost and is replaced by
another ligand. This is a ligand substitution reaction.
[Cr(H2O)6]3+ + 6NH3
[Cr(NH3)6]3+ + 6H2O
Redox reaction
In a redox reaction, the O-H and M-O bonds remain intact, but the metal ion loses or
gains an electron.
[Co(H2O)6]2+
TOPIC 13.20: HYDROLYSIS REACTIONS 3
[Co(H2O)6]3+ + e-
Hydrolysis or Acidity Reaction
When a metal-aqua ion is placed in water, the O-H bond in a co-ordinated water
molecule is broken, releasing a proton (H+(aq)), and an equilibrium is established:
[M(H2O)6]2+ + H2O
[M(H2O)5(OH)]+ + H3O+
For a metal(II) ion, the equilibrium lies well to the left hand side; the concentration of
hydrogen ions is low and the solution is, therefore, only very weakly acidic (pH ~ 6). The
acid dissociation constant, Ka typically lies between 10-6 and 10-11.
Ka =
[M(H2O)5(OH)]+ [H3O+]
[M(H2O)6]2+
Metal(III) ions have a greater charge: size ratio and are therefore more polarising. This
makes the O-H bond weaker and easier to break. The equilibrium lies further to the right
hand side and the solution is more acidic (pH ~ 3).
[M(H2O)6]3+ + H2O
[M(H2O)5(OH)]2+ + H3O+
Ka typically lies between 10-2 and 10-5.
It is important to realise that, in both the above cases, the principal species present in
the equilibrium mixture is the hexaaqua ion.
Metal(IV) ions are hydrolysed so extensively that the equilibrium lies almost entirely to
the right hand side; M(OH)4 is formed. [M(H2O)6]4+ does not normally exist in aqueous
solution.
A solution of a metal-aqua ion becomes more acidic:
 the greater the charge on the metal ion
 the smaller the size of the metal ion
Further hydrolysis can occur:
for metal(II) ions
[M(H2O)6]2+ + H2O
[M(H2O)5(OH)]+ + H2O
for metal(III) ions
[M(H2O)6]3+ + H2O
[M(H2O)5(OH)]+ + H3O+
[M(H2O)4(OH)2] + H3O+
[M(H2O)5(OH)]2+ + H3O+
[M(H2O)5(OH)]2+ + H2O
[M(H2O)4(OH)2]+ + H3O+
[M(H2O)4(OH)2]+ + H2O
[M(H2O)3(OH)3] + H3O+
The final product, the metal hydroxide, is not usually obtained unless a base has been
added to remove H+ ions and displace the equilibria further to the right hand side.
TOPIC 13.20: HYDROLYSIS REACTIONS 4
Since the hydroxide has no overall charge, the enthalpy of hydration is insufficient to
overcome the lattice enthalpy. The hydroxide is therefore insoluble and appears as a
(usually gelatinous) precipitate.
Example:
Ammonium iron(III) alum, (NH4)2SO4.Fe2(SO4)3.24H2O, is pale violet in colour because
it contains the ion [Fe(H2O)6]3+. When the salt dissolves in water, it forms a yellow
brown solution. This is because the brown, more intensely coloured ion
[Fe(H2O)5(OH)]2+ is formed by hydrolysis.
[Fe(H2O)6]3+ + H2O
[Fe(H2O)5(OH)]2+ + H3O+
N.B. [Fe(H2O)6]3+ is still the principal iron-containing species.
The position of the equilibrium can be altered, in accordance with Le Chatelier’s
principle, by changing the pH. If a source of H+ ions is added, for example dilute nitric
acid, the equilibrium is displaced to the left hand side and the solution becomes almost
colourless. If H+ ions are removed by adding an alkali such as NaOH, all three
successive hydrolysis equilibria are moved further to the right hand side. The solution
becomes increasingly brown in colour until eventually a brown precipitate of iron(III)
hydroxide [Fe(H2O)3(OH)3] is formed.
Reaction of Metal-aqua Ions with NaOH
When sodium hydroxide is added to a metal-aqua ion in solution, a precipitate
of the metal hydroxide always appears. If the hydroxide is amphoteric, it will
dissolve in excess alkali to form a clear solution.
Removal of H+ ions by the addition of OH- ions causes the successive hydrolysis
equilibria to be displaced to the right hand side.
For metal(II)-aqua ions, the overall equation is:
[M(H2O)6]2+ + 2OH-
[M(H2O)4(OH)2] + 2H2O
[Cu(H2O)6]2+
blue solution
[Cu(H2O)4(OH)2]
blue precipitate
[Co(H2O)6]2+
pink solution
[Co(H2O)4(OH)2]
blue-green precipitate
[Fe(H2O)6]2+
green solution
[Fe(H2O)4(OH)2]
dark green precipitate
TOPIC 13.20: HYDROLYSIS REACTIONS 5
For metal(III)-aqua ions, the overall equation is:
[M(H2O)6]3+ + 3OH-
[M(H2O)3(OH)3] + 3H2O
[Cr(H2O)6]3+
red-violet solution
[Cr(H2O)3(OH)3]
green precipitate
[Fe(H2O)6]3+
pale violet solution
[Fe(H2O)3(OH)3]
brown precipitate
[Al(H2O)6]3+
colourless solution
[Al(H2O)3(OH)3]
white precipitate
Amphoteric Character
The ability of metal hydroxides to dissolve in both strong acids and strong
alkalis, with the formation of a salt, is known as amphoteric character.
In the presence of a strong alkali further equilibria may be established:
[M(H2O)3(OH)3] + OH-
[M(OH)4]- + 3H2O
[M(H2O)3(OH)3] + 3OH-
[M(OH)6]3- + 3H2O
When sodium hydroxide solution is added to a solution of an aluminium salt, a white
precipitate of aluminium hydroxide is formed. This re-dissolves in excess to give a
colourless solution containing the tetrahedral aluminate ion [Al(OH)4][Al(H2O)3(OH)3] + OH-
[Al(OH)4]- + 3H2O
Most transition metal hydroxides are amphoteric and behave in this way.
When sodium hydroxide solution is added to a solution of a chromium(III) salt, a green
precipitate of chromium(III) hydroxide is formed. This re-dissolves in excess to give a
dark green solution. The behaviour of chromium(III) is unusual. It never forms
tetrahedral complexes, so the ion [Cr(OH)4]- does not exist. Instead, the octahedral ion
[Cr(OH)6]3- is formed.
[Cr(H2O)3(OH)3] + 3OH-
[Cr(OH)6]3- + 3H2O
In high oxidation states, transition metals often exist in the form of oxo anions, for
example vanadate(V) VO43-, chromate(VI) CrO42-, dichromate(VI) Cr2O72-,manganate(VI)
MnO42-, manganate(VII) MnO4-.
TOPIC 13.20: HYDROLYSIS REACTIONS 6
O
V
O-
O
OO-
Cr
O
O
Mn
O
O
Cr
OO-
Mn
O-
O
Cr
O
O
O
O
O
O
O
-
O
-
OO-
Chromate(VI) and dichromate(VI) exist in a pH dependent equilibrium:
2CrO42- + 2H+
yellow
Cr2O72- + H2O
orange
The position of the equilibrium may be altered by changing the concentration of one of
the reactants or products. The direction in which the equilibrium shifts is in accordance
with Le Chatelier’s Principle.
If an acid is added to chromate(VI) ions (the concentration of H + is increased), the
equilibrium is displaced to the right hand side, and the mixture turns orange due to the
formation of dichromate(VI) ions. If an alkali is then added, the equilibrium is displaced
towards chromate(VI), and the solution turns back to yellow.
Cr2O72- + 2OH-
2CrO42- + H2O
Reaction of Metal(II)-aqua Ions with Na2CO3
When sodium carbonate solution is added to a metal(II)-aqua ion in solution, a
precipitate of the metal carbonate appears.
Carbonate ions react with hydrogen ions in a two-stage process, forming first
hydrogencarbonate ions and then carbon dioxide and water.
H+ + CO32HCO3- + H+
HCO3H2O + CO2
Therefore, provided that it is strong enough, an acid will liberate carbon dioxide when
added to a carbonate. Metal(II)-aqua ions are not sufficiently acidic to displace carbon
dioxide from carbonates, so instead the insoluble metal(II) carbonate is precipitated.
M2+(aq) + CO32-(aq)
TOPIC 13.20: HYDROLYSIS REACTIONS 7
MCO3(s)
Reaction of Metal(III)-aqua Ions with Na2CO3
When sodium carbonate solution is added to a metal(III)-aqua ion in solution, a
precipitate of the metal hydroxide appears and carbon dioxide is evolved.
Metal(III)-aqua ions are more extensively hydrolysed in aqueous solution than metal(II)aqua ions and are therefore more acidic. They are sufficiently acidic to displace carbon
dioxide from carbonates.
[M(H2O)6]3+ + H2O
[M(H2O)5(OH)]2+ + H3O+
[M(H2O)5(OH)]2+ + H2O
[M(H2O)4(OH)2]+ + H3O+
[M(H2O)4(OH)2]+ + H2O
[M(H2O)3(OH)3] + H3O+
As the hydrogen ions produced in the above equilibria react with carbonate ions:
H+ + CO32HCO3- + H+


HCO3H2O + CO2
effervescence is seen as carbon dioxide is evolved
the hydrolysis equilibria are displaced to the right hand side, forming a precipitate
of the metal(III) hydroxide
This means that metal(III) carbonates, such as Al2(CO3)3, Cr2(CO3)3 and Fe2(CO3)3,
cannot be prepared in aqueous solution and do not exist.
If the above equations are added together in the correct ratio, the overall equation can
be represented as:
2[M(H2O)6]3+ + 3CO32-
2[M(H2O)3(OH)3] + 3CO2 + 3H2O
Reaction of Metal-aqua Ions with NH3(aq)
In an aqueous solution of ammonia, the following equilibrium is set established:
NH3 + H2O
NH4+ + OH-
Therefore when ammonia is added to an aqueous solution of a metal-aqua ion, the H+
ions in the hydrolysis equilibria can be removed by either OH - or NH3 acting as a
Bronsted-Lowry base, displacing the equilibria to the right.
When ammonia solution is added to a metal-aqua ion in solution, a precipitate
of the metal hydroxide always appears.
With an excess of ammonia, a subsequent ligand substitution reaction can occur.
TOPIC 13.20: HYDROLYSIS REACTIONS 8