Why complexes are coloured
Why complexes are coloured
- In a free ion the five d orbitals are degenerate (same energy).
- Ligands split them into two sets.
- Absorbing light across the gap gives the colour we see.
Practice
In a free transition-metal ion, the five d orbitals are:
They are degenerate until ligands approach and split them into two sets.
Splitting the d orbitals
- ligands split the d orbitals, separated by an energy gap $\Delta E$:
- octahedral: three lower + two higher.
- tetrahedral: two lower + three higher.
Practice
In an octahedral complex, the d orbitals split into:
Octahedral: 3 lower + 2 upper; tetrahedral is the reverse (2 lower + 3 upper).
The colour we see
- A complex absorbs light whose frequency matches $\Delta E$, promoting an electron to a higher d orbital.
- We see the complementary colour of the light absorbed.
- Different ligands give a different $\Delta E$ → different colour. This is why ligand exchange changes the colour.
Practice
The colour we see for a complex is:
The complex absorbs light of energy ΔE; we see the complementary colour of what was absorbed.
Practice
Ligand exchange changes the colour because different ligands:
A different ligand changes the size of the d-orbital splitting (ΔE), so a different colour is seen.
You've got it
Key idea
- a free ion's five d orbitals are degenerate; ligands split them by a gap $\Delta E$
- octahedral = 3 lower + 2 upper; tetrahedral = 2 lower + 3 upper
- the complex absorbs light matching $\Delta E$; we see the complementary colour
- different ligands → different $\Delta E$ → different colour (why ligand exchange changes colour)