The uv and copper(II) complex ions absorption spectrum of copper complex ions and compounds

(a) Octahedral complexes with water and ammonia ligands

If you add conc. ammonia to pale blue copper(II) sulfate solution, you get a stepwise ligand exchange reaction to give a whole series of copper(II) ion ammine complexes as each water molecule is replaced by an ammonia molecules.

This is accomplished by gradually adding excess ammonia solution after the initial  hydrated copper(II) hydroxide complex precipitate dissolves (see The chemistry of copper for more details and equations)

[Cu(H₂O)₆]²⁺_(aq)  +  4NH₃(aq)  ==>  [Cu(NH₃)₄(H₂O)₂]²⁺_(aq)  +  4H₂O(l)

[Cu(NH₃)₄(H₂O)₂]²⁺_(aq)  +  2NH₃(aq)  ==> [Cu(NH₃)₆]²⁺_(aq)  +  2H₂O(l)

_{The complexes can be octahedral or tetrahedral/square planar.}

 

_{Below are the very simplified 'idealised' diagrams of the spectra of three copper(II) ion complexes, so the} λ_max values are very approximate, but the trend is correct i.e. from a very pale blue to a much deeper intense violet-blue.

The original [Cu(H₂O)₆]²⁺_(aq) hexaaquacopper(II) ion, octahedral, gives a pale blue-cyan aqueous solution, that absorbs strongly in the red with a λ_max of 780 nm

 The [Cu(NH₃)₄(H₂O)₂]²⁺_(aq) diaquatetraamminecopper(II) ion, is octahedral, and dark blue, absorbing in yellow-orange-red region, with a λ_max of 600-650 nm (recently found internet quote).

Structure shown on the right.

This ion exhibits E/Z isomerism (cis/trans geometrical isomerism), in this case, the more stable isomer is the E isomer (trans) shown.

It worth noting (for 'concept' reasons), that in E/Z isomers of transition metal complex ions, the central metal ion experiences a slightly different ligand field effect, hence a slightly different colour is observed.

There will be differences in the absorption spectra of E/Z isomers of complexions.

For more on these complexes see The chemistry of copper

The final substitution product is [Cu(NH₃)₆]²⁺_(aq), the hexaamminecopper(II) ion, octahedral, an even deeper violet-indigo-blue. Unfortunately I could find a λ_max for this ion?  I've since found a quote of λ_max of ~610 nm for this complex ion.

However, seems obvious, from colour observations, the greater the amount of NH₃ replacing the H₂O ligand, the deeper the blue colour of the resulting copper(II) complex ion, so I would expect the λ_max to be less than 650 nm?, with the violet and deep blue light is transmitted and less of the yellow-orange transmitted?.

The ammonia causes a large splitting of the 3d orbitals than the water ligand, increasing the energies of absorptions at shorter wavelengths, giving a more violet-blue colour.

The above electronic energy diagram shows the octahedral field splitting effect of the 6 water/ammonia ligands on the 3d orbitals of the central copper ion (Cu²⁺) which has a 3d⁹ configuration - meaning there is partially filled 3d orbital in the complex ion (compare this with the 3d¹⁰ configuration of a Cu⁺ ion).

The ∆E_elec excitation energy, promoting an electron from a lower 3d orbital to a higher 3d orbital is within the copper(II) complex ions light photon range, so the ion is coloured.

At this pre-university level, the diagram applies to the square planar complexes of Cu²⁺ with H₂O or NH₃ i.e. for a 4 ligand square complexes like [Cu(NH₃)₄]²⁺ or [Cu(H₂O)₄]²⁺, the electronic diagram is the same for an octahedral complex described above.

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(b) A tetrahedral/square planar copper(II) ion complex with a chloride ion ligand

In the presence of excess chloride ions, the octahedral hexaaquacopper(II) ion forms the tetrahedral (or square planar?) ion, which theoretically is colourless, but isn't in reality!

[Cu(H₂O)₆]²⁺_(aq) + 4Cl^–_(aq) [CuCl₄]^2–_(aq) + 6H₂O_(l)

The above electronic energy diagram shows the tetrahedral field splitting effect of the 4 chloride ion ligands on the 3d orbitals of the central copper ion (Cu²⁺).

In salts, depending on the size of the cation, you observe an orange colour for the distorted tetrahedral [CuCl₄]^2- ion or pale yellow for the same ion in a square planar configuration.

This shows a change in crystal structure for the same ion can cause a difference in the ∆E_elec excitation energy, hence a change in colour of the copper(II) complex ion.

For a 4 ligand square complexes like [Cu(NH₃)₄]²⁺ or [Cu(H₂O)₄]²⁺, the electronic diagram is the same for an octahedral complex described above, noting the 'middle' of an octahedral complex is square planar.

For more on these complexes see The chemistry of copper

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(c) Complexes of the copper(I) ion

Not all, but many of copper(I) complex ions are colourless - you get colour with copper(I) charge transfer complexes.

The reason why you can get a colourless copper(I) complex ions/compounds is explained in the diagrams below.

So, I'm presenting the 'electronic' argument as to why you might expect copper(I) compounds to be colourless e.g. CuCl, copper(I) chloride solid is white, but readily oxidises in air or water to give a cyan colour of a copper(II) compound.

The above electronic energy diagrams show the field splitting effect of the 4 or 6 ligands on the 3d orbitals of the central copper ion (Cu⁺) which has a 3d¹⁰ configuration (not the difference from the 3d⁹ configuration for a Cu2+ ion).

This means there is NO partially filled 3d orbital in the complex ion.

Therefore, there is no partially filled 3d orbital to which an electron can be excited by a copper(II) complex ions light photon.

For more on copper(I) chemistry see The chemistry of copper

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