Colour of complexes

Chemistry A2 · Chemistry of transition elements. The idea in plain words, the rules to remember, the mistake to avoid and a worked example, written and checked by Mega Lecture.

Test this topic free All Chemistry A2 topics

The idea

In a free transition metal ion the five d orbitals are degenerate: they all have the same energy. When ligands bond to the ion, they repel the d electrons by different amounts, so the d orbitals split into two sets with different energies (non-degenerate). The energy gap between the sets is ΔE.

An electron in a lower d orbital can absorb light and move up to a higher d orbital. The light absorbed has exactly the energy ΔE, so only some frequencies of visible light are removed. We see the light that is left, which is the complementary colour.

A different ligand gives a different ΔE, so a different frequency is absorbed and the colour changes.

Rules to remember

Degenerate orbitals have the same energy; non-degenerate orbitals have different energies.
Octahedral complex: two d orbitals higher (3d_z², 3d_x²−y²), three lower. Tetrahedral complex: three higher, two lower.
Energy gap = Planck constant × frequency absorbed (ΔE = hf): a larger ΔE means light of higher frequency (shorter wavelength) is absorbed.
Colour seen = complementary colour of the light absorbed, e.g. orange-red absorbed, blue seen.
Copper(II): [Cu(H₂O)₆]²⁺ blue; Cu(OH)₂(H₂O)₄ pale blue precipitate; [Cu(NH₃)₄(H₂O)₂]²⁺ deep blue; [CuCl₄]²⁻ yellow (the solution often looks yellow-green).
Cobalt(II): [Co(H₂O)₆]²⁺ pink; Co(OH)₂(H₂O)₄ blue precipitate; [Co(NH₃)₆]²⁺ yellow-brown; [CoCl₄]²⁻ blue.
No colour if the d orbitals are empty (3d⁰) or full (3d¹⁰): no d electron can be promoted.

Common mistake

Saying the colour seen is the colour absorbed, or that electrons give out light as they fall back. The colour comes from light being absorbed; what you see is the light that is not absorbed.

Worked example

Complexes of copper(I), such as [Cu(NH₃)₂]⁺, are colourless, but complexes of copper(II) are coloured. Why?

Cu⁺ is [Ar] 3d¹⁰; Cu²⁺ is [Ar] 3d⁹.
In Cu⁺ all five d orbitals are full, so there is no space in the higher set for an electron to move into.
In Cu²⁺ one higher d orbital is half-filled, so an electron can be promoted by absorbing visible light of energy ΔE.

Answer: Cu⁺ (3d¹⁰) cannot absorb visible light by promoting a d electron, so it is colourless; Cu²⁺ (3d⁹) can, so it is coloured.

Ask about this

Type what you are stuck on. The helper finds the explanation, worked examples and the notes to read, free and instant.

Read the full notes

Watch: 2 | Transition Metals 2 | Coloured Compounds & Complexes

Keep going

← General characteristic chemical properties of the transition elements, titanium to copper Stereoisomerism in transition element complexes →Take a Chemistry A2 test