configurations and the theory and variation of complex ion colour
Transition metal ions can be
identified by their colour.
The colour arises when some of the
wavelengths of visible light are absorbed and the remaining wavelengths
of light are transmitted or reflected, so you experience the net effect.
In transition metal species d
electrons move from the ground state to an excited state when light is
The energy difference between the
ground state and the excited state of the d electrons is given by the
Changes in oxidation state,
co-ordination number and ligand alter ∆E and this leads to a change in
colour e.g. of the transition metal complex ion.
The absorption of visible light is
used in spectroscopy e.g. using a
colorimeter to determine the concentration of coloured ions in
solution and also the
chemical formula of
a complex ion.
For the 3d–block know the
complete order of filling of the sub–shells from Z=21 to 30 and be
able to write out the full or abbreviated electron
Transition metals can be
identified by the colour of their complexes which of course is a
very characteristic feature of their chemistry (e.g. the
hydroxide precipitates which
are, of course, all neutral complexes).
can varies with
change in (i) oxidation state, (ii) ligand and (iii) co–ordination
number or shape (which in turn depends on the ligand and oxidation
state) and obviously changing the transition metal itself, will give
another range of differently coloured compounds.
All of these
factors are linked to the electronic state of the central metal ion, so,
if the electronic levels are changed by change in oxidation state or
ligand, the difference between quantum
levels changes, therefore the wavelength of the light photons absorbed
changes, i.e. the observed colour changes e.g.
e.g. (i) The
same ligand (H2O), shape and co–ordination number but
different oxidation state.
e.g. (ii) The
same oxidation state, shape and co–ordination number but
e.g. (iii) The
same oxidation state but with a different ligand, shape and
pale blue hexaaquacopper(II) ion and [CuCl4]2–,
yellow tetrachlorocuprate(II) ion.
state +2, but different ligands (water and chloride ion), different
shape (octahedral and tetrahedral) and different co–ordination
number (6 and 4).
THEORY for transition element complexes:
The argument is
presented from the point of view of an octahedral complex, but
similar arguments apply for a tetrahedral or square planar
What is presented below is a
simplified version of crystal field theory for the 3d sub-orbitals.
The full explanation involves ligand
field molecular orbital theory and both theories would be dealt with at
university level courses.
There are five 3d
sub–shell orbitals whose 3D spatial representations are shown
below. Theoretically it is considered that the ligands in an
octahedral complex approach the central metal ion along the x, y
and z axis, which would minimise the repulsion between the
orbitals of bonding electrons in the six M–ligand dative covalent bonds (note that
4s and 4p orbitals are involved in complex ion bonding).
electronic ground states
of scandium(III), titanium(III), copper(II) and zinc(II) are illustrated below.
Left: Electronic diagrams for
octahedral complexes of
scandium, titanium(III), copper(II) and zinc and colour of hexaaqua
electronically excited states of titanium(III) and copper(II) -
the colours are shown with the electronic diagrams for their hexaaqua metal cations
The colours/colourless of the four complex ions shown
above are illustrated below.
It is the ligands in transition metal
complexes that cause a splitting of the d orbitals in the d sub-shell (see
The observed colour resulting
from the ∆Eelec changes due to the 3d (or any
d) orbital splitting depend on:
(i) the d electron configuration of
the central metal ion and its oxidation state - the electronic state of
the d orbital sub-shell,
(ii) the nature of the
ligand (L) and strength of its bond with the central metal ion
different ligands have different effects on the relative energies of the
d orbitals of a particular ion - the Mn+-L binding can be
weak or strong.
(iii) and the number and spatial
arrangement of the ligands - this affects the splitting of the d
The colour arises from
electronic transitions from the ground state to excited states, the
energy needed can be calculated using
Equation, ΔE = hv
, E = energy of a single photon (J), h
= Planck's Constant (6.63 x 10–34 JHz–1),
v = frequency (Hz).
Therefore if the
photo energy/frequency is equal to
then energy is absorbed and an electron can be promoted from the
lower 3d level to the higher 3d level.
is in the visible light frequency range the complex will be
In the case
of coloured transition metal complexes, the colour arises from
visible light energy absorption on promoting electrons from the
lower 3d levels to the higher 3d levels.
The colour you observe is
derived from the wavelengths of visible light which are NOT
However, this can only
occur if there is at least one electron in the 'lower' 3d orbitals and
at least one half–filled 'higher' 3d quantum level orbital, i.e. the minimum pre–conditions for
an electronic transition or 'excitation'.
These are known as d-d transitions
and vary a great deal in terms of colour intensity - often not that
intense e.g. hexaaqua complex of Mn2+(aq) is a
very pale pink but the ammonia complexes of Cu2+(aq)
are much more deeply coloured.
However, in some cases the electron
is promoted from a metal d orbital to a ligand orbital (or vice versa),
the transition is called a charge transfer and a generally more intense
colour results e.g. the deep purple manganate(VII) ion, MnO4-(aq)
there is a lack of possible d-d transitions in the Sc(III) Sc3+ and Zn(II)
Zn2+ ions, their compounds are usually
colourless i.e. no light absorbed in the visible region of the spectrum.
In the true transition metals
from Ti to Cu,
it is possible for the electromagnetic radiation energy to produce this excitation from the lower to the higher
3d sub–levels and it is usually in the visible region.
ranges of visible
radiation are absorbed and the perceived colour arises from the frequencies not
absorbed i.e. the transmitted visible light.
The electronic structure
and colour of some typical 'simple' aqueous ions is shown below. They
are all hexa–aqua ions of an octahedral shape except ...
form a stable simple Cu+(aq) ion, but copper(I)
compounds tend to be colourless when pure e.g CuCl, copper(I) chloride,
but copper(II) forms the blue square planar [Cu(H2O)4]2+
The colour you see in a
transition metal compound is the visible light that isn't absorbed by
the 3d electronic transitions.
in transition metal reactions can arise from change of ligand,
change in co–ordination number or change in metal oxidation state
(sometimes several of these simultaneously.
The colours are
quite useful for simple transition metal ion identification tests e.g.
precipitates with sodium hydroxide and ammonia (see pictures) and the
thiocyanate test for iron(III)
visible absorption spectra
Dyes and pigments
Ultraviolet and visible
spectroscopy can be used to determine the concentration of metal
ions in solution, usually after the addition of a suitable ligand to
intensify the colour using the more elaborate technique of
spectrophotometry or the simpler technique of
– appendix 9.
Colorimetric analysis of coloured solutions for quantitative analysis
using a colorimeter described in Appendix 9.
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