know that transition metal ions can be identified by their colour,
limited to the complexes in this unit
know that colour changes arise from changes in oxidation state, co-ordination
number and ligand
know that colour arises from electronic transitions from the ground
state to excited states: ΔE = hf
appreciate that this absorption of visible light is used in spectrometry
to determine the concentration of coloured ions
Colour in materials or compounds is caused when the light reaching the eyes
has some of the wavelengths removed by absorption. When we see an object as
red it is only this colour that is reflected from the object while the other
colours or wavelengths are absorbed.
A coloured solution is caused by the white light passing through it and losing
some of its wavelengths by absorption. If the solution appears blue it means
that the complementary colours are absorbed by the solution.
If colour is caused by the absorption of certain wavelengths from white light,
the question remains - how are these wavelengths absorbed?
The colour in the transition metals (d-block) is usually due to the 'splitting'
of the 'd' shell orbitals into slightly different energy levels. As a result,
certain wavelengths of energy can be absorbed by the d-block elements (with
electrons jumping between these slightly different energy levels), resulting
in the complementary colour being visible.
The theory that explains the production of colour in this way is called crystal
When ligands bond to transition metals they do so by donating electrons into
empty hybridised orbitals. The inner 3d orbitals are not involved in the bonding.
They are, however, affected by the repulsive force of the donated electron pairs
and the transition metal orientates itself to minimise the repulsions between
the 3d orbital electrons and the ligands electron pairs.
When the ligands are arranged octahedrally around the transition metal ion
the lowest energy orientation for the ion involves keeping the dxy, dyz and
dxz orbitals in the gaps between the ligands. However, this puts the other two
orbitals, the dx2-y2 and the dz2 orbital close to the ligands, raising their
energy with respect to the other three orbitals. The 3d orbitals are now no
longer degenerate. This is called crystal field splitting.
If the crystal field (the electrostatic repulsion) is strong, then the 'd'
orbitals are split further with more energy between the two different sets.
This depends on the ligands. Ligands can be arranged according to the strength
of their electrostatic effect into an spectrochemical series.
Anything that can affect the electrostatic field around the transition metal
ion can affect the wavelengths of light absorbed and hence the colour transmitter
by a solution, or reflected by a solid.
Colour in transition metal complexes is affected by three factors:
1 the transition metal
2 the oxidation state of the transition metal
3 the type of ligand
The transition metal
The transition metals have certain colours, or colour ranges that are typical
of that metal. Copper salts, for example, are usually blue or green, iron has
salts that are pale green, yellow or orange.
The oxidation state
The oxidation state is important. Copper (II) salts are coloured, whereas copper
(I) salts are white solids.
The reason for this lies in the electronic configurations of the two oxidation
Copper (II) [Ar] 4s0 3d9
Copper (I) 4s0 3d10
The copper (I) ions cannot absorb energy for d-d transitions as there are no
empty, or partially empty, orbitals available to accept a promoted electron.
Table of some common transition metal complex ions and their
Oxidation state of metal
The nature of the ligand
The ligand can affect the colour in two ways. Ligands all have different crystal
field strenghts and will split the 'd' orbitals by differing amounts. However,
apart from this, the actual shape of the complex is defined by the type of ligand
and the oxidation state of the transition metal. If the shape of the complex
changes, then this also causes a change in the type of crystal field splitting.
Tetrahedral fields split the 3d orbitals in a different manner to an octahedral
Example: Hexaaqua copper (II) ions are light blue in solution.
The complex is octahedral, and there are six water molecules coordinated
to the copper ion. If concentrated hydrochloric acid is added to the solution,
the water ligands get replaced by chloride ligands and a new complex,
tetrachlorocuprate (II) is formed.
This new complex has a tetrahedral geometry and a different crystal field
splitting pattern. Its colour is now a deep green as the new energy difference
between the non-degenerate 'd' orbitals absorbs light of a different wavelength.
If concentrated ammonia solution is now added to the tetrachlorocuprate
(II) solution, the colour fades and a light blue precipitate is formed,
which then dissolves to form a deep blue solution. This final solution
contains the tetrammine copper (II) complex, which is square planar.