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 simple
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
See under each
element and even more detail in Electron Configurations in Periodic Table section 2.2
Transition metal electron
arrangements are listed
The number of
orbitals per sub–shell, 1 for s, 3 for p, and 5 for d sub–shell.
PLEASE watch out for the two
‘quirks’ for Cr and Cu atoms and the order of
electron removal when forming positive ions e.g. for the 3d block of
transition metals, you remove the 4s electrons first, before any
of the 3d electrons.
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
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
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.
electronically excited states of titanium(III) and copper(II) are
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.
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, i.e. the minimum pre–conditions for
an electronic transition or 'excitation'.
there a lack of such
possible transitions in Sc(III) and Zn(II) 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. For example, copper(II) complexes often
absorb in the red area of the visible spectrum, so the resulting
colour observed is green–blue.
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 is described in Appendix 9.
* Titanium * Vanadium
* Manganese Iron * Cobalt
* Copper *
* Silver & Platinum
Introduction 3d–block Transition Metals * Appendix
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.
* Appendix 8. Stability Constants and entropy
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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