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Appendix
1.
Hydrated
salts, metal–aqua
complex ions and their relative
acidity, salt hydrolysis
(not necessarily just
transition metal ions)
-
All metal ions in
solution are 'associated' with water. The water molecules can also be
weakly bonded or more strongly as a ligand to form a complex
ion, and these can also present in solid 'hydrated' salts on
crystallisation e.g.
-
FeSO4.7H2O(s),
CoCl2.6H2O(s), CuSO4.5H2O(s)
etc.
-
Iron(II)
sulfate heptahydrate, cobalt(II) chloride hexahydrate and
copper(II) sulfate pentahydrate.
-
The above
crystals contain 7, 6 and 5 molecules of water of crystallisation
respectively.
-
A hexa–aqua ion
is present in the first two, [M(H2O)6]2+
(M = Fe, Co)
-
What is the
difference between water of crystallization and the co–ordinated
water molecules bonded to the central metal ion via the dative
covalent bonds?
-
There may or not be a
difference!
-
The water of
crystallisation is the total number of water molecules incorporated
into the crystal structure irrespective of the nature of the
chemical bonding involved OR any intermolecular associations.
-
The case of copper(II)
sulfate pentahydrate is considered below, where 4/5 water molecules
are ligands and the 5th water molecule is held in place by hydrogen
bonding.
-
However in the case of
magnesium chloride (magnesium isn't a transition metal), the crystal
lattice consists of hexaaqua magnesium ions and chloride ions.
-
Cl–
<=> [Mg(H2O)]2+ <=> Cl– are all
bonded together by electrostatic attraction in the crystal lattice.
-
No hydrogen bonding is
involved, and the number of molecules of water of crystallisation
is equal to the ligand coordination number of the central metal ion.
-
Just in passing, a
complex ion is a complex ion, it doesn't matter if the central metal
ion isn't a 3d block transition metal ion etc. Its still the same
sort of structure with the same sort of bonding and shape.
-
In the case of
copper(II) sulfate, 4 water molecules are covalently bonded to form
effectively a square
planar complex ion,
-
[Cu(H2O)4]2+
and the 5th water molecule H2O is hydrogen bonded to this ion and
hydrogen bonded to a
neighbouring sulfate ion
-
SO42–
thus helping to hold the crystal lattice
together, though the main force of attraction is the electrostatic
attraction between copper complex ion and the sulfate ion.
-
Therefore, the water
of crystallisation number doesn't equal the co–ordination number of
the central metal ion.
-
 ...
H–O–H ...
 are
the three components of the crystal structure of copper(II) sulfate
pentahydrate, and all three are linked by hydrogen bonds. The full
structure is a bit complicated to draw but the 5th and 6th
octahedral positions of the Cu2+ ion are occupied by
oxygen atoms of the sulfate ion and the 5th water molecule is held
in position by hydrogen bonding.
-
However,
this blue crystal lattice is readily broken down on heating, a
classical demonstration of a reversible reaction, since the
white anhydrous solid turns blue on adding water (a simple
test for water.
-
CuSO4.5H2O(s)
CuSO4(s) + 5H2O(g/l)
-
So three words–phrases
to know ...
-
Water of
crystallisation – the molecules of water incorporated into the
crystal structure either by acting as a ligand to a metal ion (e.g.
four in hydrated CuSO4) or just hydrogen bonded into the
lattice (one in hydrated CuSO4), so a total of five molecules of water
of crystallisation as discussed above.
-
Hydrated salt –
lattice contains molecules of associated water e.g. water of
crystallisation in the case of salts.
-
Anhydrous salt –
devoid of water molecules in the crystal lattice e.g. dehydrated
salts which can crystallise with water of crystallisation.
-
It should be pointed out
that the term anhydrous is used quite generally to mean a substance
has had all water–moisture removed from it.
-
See also
Calculations of water
of crystallization – % composition & simple experimental
determination.
-
Lewis acid–base theory
reminders:
-
A base is an electron pair donor and an
acid is an electron pair
acceptor.
-
Ligands like
water, can donate a pair of non–bonding electrons (lone pair) into a
vacant orbital of a central metal ion and so dative covalent
(co–ordinate) bonds
hold a complex together.
-
The central metal
ion with vacant bonding orbitals can act as a Lewis acid.
-
Ligands act as Lewis
bases by electron pair donation to form the metal–ligand bond.
-
Bronsted–Lowry
acid–base theory reminders (essentially a sub–set of Lewis Theory)
-
Many hexa–aqa complex ions can
undergo acid–base reactions with water to produce solutions of pH less
than 7.
-
Usually group
2, 3 and transition metal ions.
-
The positive central
metal ion polarises a water molecule, releasing a proton, H+.
-
In the
deprotonation reaction the proton
transfers to water and the overall charge on the complex falls by 1
unit since the H2O – H+ = OH–,
i.e. one of the ligands is now a hydroxide ion.
-
In these
reactions the hydrated ions act as
Bronsted Lowry acids and
water
acts as a Bronsted–Lowry base.
-
These reactions are
examples of what is termed 'salt hydrolysis' because
the metal ion (of usually a salt) reacts with water to give, in
this case, two products.
-
These are acid–base
reactions NOT redox reactions, even if they involve
transition metal ions – there is NO change in oxidation state of
the metal!
-
 e.g.
for hexaaquametal(II) ions ...
-
[M(H2O)6]2+(aq)
+ H2O(l)
[M(H2O)5(OH)]+(aq)
+ H3O+(aq)
-
e.g. when M = Mn, Fe,
Co, Ni, Cu, Mg etc. gives a very weak acid solutions with
pH's just less than 7.
-
Ti(II), V(II)
and Cr(II) M2+ ions are redox unstable in the presence of air, but
theoretically their salts give very weakly acid solutions.
-
They are usually prepared by zinc–acid reduction from higher
oxidation states.
-
–
-
e.g. for
hexaaquametal(III) ions ...
-
 [M(H2O)6]3+(aq)
+ H2O(l)
[M(H2O)5(OH)]2+(aq)
+ H3O+(aq)
-
e.g. when M = Ti, V,
Cr, Fe, Al etc. give very weak acids solutions (but generally
stronger than for M2+) of pH's in the 3–5 region.
-
In the
presence of alkali, OH–, removing H3O+
ions, the equilibrium moves more to the right and more
protons are lost from the complex in stages until the hydroxide
precipitate is formed e.g. for iron(III), chromium(III) or
aluminium.
-
Some of the M3+
hydroxides are amphoteric and dissolve in excess alkali
(1.) or acid (2.) e.g. to eventually form for chromium(III) or
aluminium, 1. the soluble hexa–hydroxo complex ion or 2. the
original hexa–aqua ion.
-
[M(H2O)3(OH)3]0(aq)
+ 3OH–(aq) 
[M(OH)6]3–(aq)+
3H2O(l)
-
[M(H2O)3(OH)3]0(aq)
+ 3H3O+(aq)
[M(H2O)6]3+(aq)
+ 3H2O(l)
-
As a general rule
the greater the
polarising power of the central metal ion, the lower the pH of the resulting aqueous solution, i.e.
the acid–base equilibrium is shifted more to the right causing an
increase in acidity of the solution.
-
This effect
and process facilitated by the central metal ion on one
water ligand molecule can be envisaged as
-
[M–O–H2]n+
==> [M–O–H]n+–1 + H+ (proton transferred to a water
molecule)
-
One of the O–H
bond pairs is 'attracted' onto the oxygen atom by the electric field
effect of the central metal ion of charge n+, allowing proton
transfer to the base water.
-
Polarising
power is a function of ionic charge/ionic radius ratio.
therefore ...
-
the greater the metal
ion charge (n+), the more acidic the hexaaqua ion, hence a lower
pH solution,
-
and the smaller the
radius of the metal ion, the more acidic the hexaaqua ion, hence a
lower pH solution,
-
so ...
-
both
increasing charge, or decreasing the central cation radius intensify the electric
field polarising effect on a water
ligand which facilitates proton donation from the complex ion to a
free water molecule.
-
The acidity of the
hexaaqua ions
M3+(aq) due to the polarising influence of the
central highly charged M3+ ion accounts for the lack of
stability/existence of e.g. aluminium carbonate, iron(III)
carbonate or chromium(III) carbonate, whereas MgCO3 , ZnCO3 and FeCO3
etc. with the less polarising M2+ ion exist. It
also accounts for why you see bubbles of carbon dioxide when
carbonates/hydrogencarbonates are mixed with aluminium chloride, iron(III) chloride or chromium(III) chloride solutions e.g.
-
2[Fe(H2O)6]3+(aq)
+ CO32–(aq)
2[Fe(H2O)5(OH)]2+(aq)
+ H2O(l) + CO2(g)
-
or [M(H2O)6]3+(aq)
+ HCO3–(aq)
[M(H2O)5(OH)]2+(aq)
+ H2O(l) + CO2(g)
-
There several
legitimate permutations based on these equations.
-
Amphoteric nature of
ions such as Al3+ and Fe3+
-
In the above chemistry
the acidic nature of the hexa–aqua ions was emphasised, BUT as
soon as one proton has been lost the resulting complex ion can then act
as a base.
-
e.g. in solutions of the
weakly acidic Fe3+ or Al3+ species with
excess strong acid the hexa–aqua ion would predominate
-
[M(H2O)4(OH)2]+(aq)
+ 2H+(aq)
 [M(H2O)6]3+(aq)
-
and hydroxide precipitates
will readily dissolve on adding an acid which is stronger than the
hexa–aqua ion
-
e.g. Fe(OH)3(s)
+ 3H3O+(aq) ==> [Fe(H2O)6]3+(aq)
-
or Al(OH)3(s)
+ 3H3O+(aq) ==> [Al(H2O)6]3+(aq)
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The sequence of
chromium(III) hydroxide precipitate formation and its subsequent
dissolving in excess strong alkali. Each step is essentially one of
proton removal from each complex (from 3+ to 3–). |
| 1 |
2 |
3 |
4 |
From 1 to 7 happen
as you add more alkali, increasing pH and the OH–
concentration, removing protons from the complex. |
| 5 |
6 |
7 |
* |
From 7 back to1
represents what happens when you add acid, decreasing pH, increasing
H+/H3O+ concentration and
protonating the complex. |
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The sequence of
aluminium hydroxide precipitate formation and its subsequent
dissolving in excess strong alkali. Each step is essentially one of
proton removal from each complex (from 3+ to 3–). |
Scandium
* Titanium * Vanadium
* Chromium
* Manganese
Iron * Cobalt
* Nickel
* Copper *
Zinc
*
Silver & Platinum
A level Revision notes for GCE Advanced
Subsidiary Level AS Advanced Level A2 IB
Revise AQA GCE Chemistry OCR GCE Chemistry Edexcel GCE Chemistry Salters
Chemistry CIE Chemistry, WJEC GCE AS A2 Chemistry, CCEA/CEA GCE AS A2 Chemistry revising courses for pre–university students
(equal to US grade 11 and grade 12 and AP Honours/honors level courses)
Introduction 3d–block Transition Metals * Appendix
1.
Hydrated salts, acidity of
hexa–aqua ions * Appendix 2. Complexes
& ligands * Appendix 3. Complexes and isomerism * Appendix 4.
Electron configuration & colour theory * Appendix 5. Redox
equations, feasibility, Eø * Appendix 6.
Catalysis * Appendix 7.
Redox
equations
* Appendix 8. Stability Constants and entropy
changes *
Appendix 9. Colorimetric analysis
and complex ion formula * Appendix 10 3d block
– extended data
* Appendix 11 Some 3d–block compounds, complexes, oxidation states
& electrode potentials * Appendix 12
Hydroxide complex precipitate 'pictures',
formulae and equations
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