What is a complex ion?

A complex ion has a metal ion at its centre with a number of other molecules or ions surrounding it. These can be considered to be attached to the central ion by co-ordinate (dative covalent) bonds. (In some cases, the bonding is actually more complicated than that.)

The molecules or ions surrounding the central metal ion are called ligands.

Simple ligands include water, ammonia and chloride ions.







What all these have got in common is active lone pairs of electrons in the outer energy level. These are used to form co-ordinate bonds with the metal ion.

Some examples of complex ions formed by transition metals
[Fe(H2O)6]2+
[Co(NH3)6]2+
[Cr(OH)6]3-
[CuCl4]2-

Other metals also form complex ions - it isn't something that only transition metals do. Transition metals do, however, form a very wide range of complex ions.

Why do we see some compounds as being coloured?

White light

You will know, of course, that if you pass white light through a prism it splits into all the colours of the rainbow. Visible light is simply a small part of an electromagnetic spectrum most of which we can't see - gamma rays, X-rays, infra-red, radio waves and so on.

Each of these has a particular wavelength, ranging from 10-16 metres for gamma rays to several hundred metres for radio waves. Visible light has wavelengths from about 400 to 750 nm. (1 nanometre = 10-9 metres.)

The diagram shows an approximation to the spectrum of visible light.





Why is copper(II) sulphate solution blue?

If white light (ordinary sunlight, for example) passes through copper(II) sulphate solution, some wavelengths in the light are absorbed by the solution. Copper(II) ions in solution absorb light in the red region of the spectrum.

The light which passes through the solution and out the other side will have all the colours in it except for the red. We see this mixture of wavelengths as pale blue (cyan).

The diagram gives an impression of what happens if you pass white light through copper(II) sulphate solution.






Working out what colour you will see isn't easy if you try to do it by imagining "mixing up" the remaining colours. You wouldn't have thought that all the other colours apart from some red would look cyan, for example.


Sometimes what you actually see is quite unexpected. Mixing different wavelengths of light doesn't give you the same result as mixing paints or other pigments.
You can, however, sometimes get some estimate of the colour you would see using the idea of complementary colours.


Complementary colours


If you arrange some colours in a circle, you get a "colour wheel". The diagram shows one possible version of this. An internet search will throw up many different versions!







Colours directly opposite each other on the colour wheel are said to be complementary colours. Blue and yellow are complementary colours; red and cyan are complementary; and so are green and magenta.


Mixing together two complementary colours of light will give you white light.


Beware: That is NOT the same as mixing together paint colours. If you mix yellow and blue paint you don't get white paint. (:


What this all means is that if a particular colour is absorbed from white light, what your eye detects by mixing up all the other wavelengths of light is its complementary colour. Copper (II) sulphate solution is pale blue (cyan) because it absorbs light in the red region of the spectrum. Cyan is the complementary colour of red.


Some sample colours
The diagrams show the approximate colours of some typical metal ions. The charge on these ions is typically 2+ or 3+.


Non-transition metal ions





These ions are all colourless. (Sorry, we can't do genuinely colourless!)

Transition metal ions




The corresponding transition metal ions are coloured. So . . . what causes transition metal ions to absorb wavelengths from visible light (causing colour) whereas non-transition metal ions don't? And why does the colour vary so much from ion to ion?



The origin of colour in complex ions containing transition metals

Complex ions containing transition metals are usually coloured, whereas the similar ions from non-transition metals aren't. That suggests that the partly filled d orbitals must be involved in generating the colour in some way. Remember that transition metals are defined as having partly filled d orbitals.

Octahedral complexes


For simplicity we are going to look at the octahedral complexes which have six simple ligands arranged around the central metal ion. The argument isn't really any different if you have multidentate ligands - it's just slightly more difficult to imagine!


When the ligands bond with the transition metal ion, there is repulsion between the electrons in the ligands and the electrons in the d orbitals of the metal ion. That raises the energy of the d orbitals.


However, because of the way the d orbitals are arranged in space, it doesn't raise all their energies by the same amount. Instead, it splits them into two groups.


The diagram shows the arrangement of the d electrons in a Cu2+ ion before and after six water molecules bond with it.








Whenever 6 ligands are arranged around a transition metal ion, the d orbitals are always split into 2 groups in this way - 2 with a higher energy than the other 3.


The size of the energy gap between them (shown by the blue arrows on the diagram) varies with the nature of the transition metal ion, its oxidation state (whether it is 3+ or 2+, for example), and the nature of the ligands.

When white light is passed through a solution of this ion, some of the energy in the light is used to promote an electron from the lower set of orbitals into a space in the upper set.

Each wavelength of light has a particular energy associated with it. Red light has the lowest energy in the visible region. Violet light has the greatest energy.
Suppose that the energy gap in the d orbitals of the complex ion corresponded to the energy of yellow light.




Each wavelength of light has a particular energy associated with it. Red light has the lowest energy in the visible region. Violet light has the greatest energy.



Suppose that the energy gap in the d orbitals of the complex ion corresponded to the energy of yellow light.








The yellow light would be absorbed because its energy would be used in promoting the electron. That leaves the other colours.

Your eye would see the light passing through as a dark blue, because blue is the complementary colour of yellow.





What about non-transition metal complex ions?

Non-transition metals don't have partly filled d orbitals. Visible light is only absorbed if some energy from the light is used to promote an electron over exactly the right energy gap. Non-transition metals don't have any electron transitions which can absorb wavelengths from visible light.


For example, although scandium is a member of the d block, its ion (Sc3+) hasn't got any d electrons left to move around. This is no different from an ion based on Mg2+ or Al3+. Scandium(III) complexes are colourless because no visible light is absorbed.

In the zinc case, the 3d level is completely full - there aren't any gaps to promote an electron in to. Zinc complexes are also colourless.


Tetrahedral complexes

Simple tetrahedral complexes have four ligands arranged around the central metal ion. Again the ligands have an effect on the energy of the d electrons in the metal ion. This time, of course, the ligands are arranged differently in space relative to the shapes of the d orbitals.

The net effect is that when the d orbitals split into two groups, three of them have a greater energy, and the other two a lesser energy (the opposite of the arrangement in an octahedral complex).

Apart from this difference of detail, the explanation for the origin of colour in terms of the absorption of particular wavelengths of light is exactly the same as for octahedral complexes.


Note: Coding for the ligand


The table shows some common ligands and the code for them in the name of a complex ion. The old names sometimes differ by a letter or so, but never enough for it to be confusing.








Take care with the code for ammonia as a ligand - it has 2 "m"s in its name. If you miss one of these out so that you are left with "amine" or "amino", you are referring to the NH2 group in an organic compound.


This is probably the only point of confusion with these names.

The factors affecting the colour of a transition metal complex ion

In each case we are going to choose a particular metal ion for the centre of the complex, and change other factors. Colour changes in a fairly haphazard way from metal to metal across a transition series.





The nature of the ligand

Different ligands have different effects on the energies of the d orbitals of the central ion. Some ligands have strong electrical fields which cause a large energy gap when the d orbitals split into two groups. Others have much weaker fields producing much smaller gaps.



Remember that the size of the gap determines what wavelength of light is going to get absorbed.

The list shows some common ligands. Those at the top produce the smallest splitting; those at the bottom the largest splitting.





The greater the splitting, the more energy is needed to promote an electron from the lower group of orbitals to the higher ones. In terms of the colour of the light absorbed, greater energy corresponds to shorter wavelengths.

That means that as the splitting increases, the light absorbed will tend to shift away from the red end of the spectrum towards orange, yellow and so on.