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Why are transition metal complex ions coloured?

Firstly, transition metal ions are not coloured on their own - it is only when they become complex ions that they become coloured. 

In a transition metal ion, all the orbitals are at the same energy level (there's no difference in energy level between them whatsoever). However, when it forms a complex ion (which is when ligands surround the ion & the no. of coordinate bonds is greater than the oxidation state), the d orbitals split into 2 sets as the electrons from the ligands interfere with it.

Accordingly, there is now a set of 3 orbitals and a set of 2 orbitals. These sets are at different energy levels - the 3-orbital set is at a lower energy level than the 2-orbital set. Hence, a specific amount of energy is needed for an electron to move from the lower orbitals to the higher orbitals. Remember from AS that [Delta]E = hf and so, the energy required corresponds to a specific frequency of light.

Hence, when white light is shone onto the complex ion, certain frequencies of light are absorbed (which promotes an electron to a ‘higher’ orbital).

Any light which is not absorbed is the colour of the complex ion as the non-absorbed light is reflected back into our eyes.

For example, if the colour of the complex ion is blue - this means that red light is absorbed and that blue light is reflected. It is helpful to look at/think of the visible spectrum of light for this (the colours of the rainbow basically!).In essence, the colour that we see consists of the frequencies of light which are not absorbed. 

Finally, different ligands interfere with the ion’s orbitals differently and thus, cause different splitting of the d-orbitals. Therefore, different frequencies of light are absorbed and this leads to different colours.

For example, Cu2+ in water is pale blue (the water molecules are the ligands) but when ammonia is added, the complex ion turns a royal blue colour. 

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