The colour of the [Cr(H₂O)₆]³⁺_(aq) hexaaquachromium(III) ion is a sort of violet-blue-grey, but tends to be greenish in the presence of chloride ions due to the formation of aqua-chloro complexes e.g. [Cr(H₂O)₄Cl₂]⁺_(aq) which is dark green.

[Cr(H₂O)₅Cl]²⁺_(aq) is green and note how to work out the overall charge on these three chromium(III) complex ions.

Remember a change of ligand usually results in a change of 3d orbital electronic levels, and the ligand field splitting, hence a change in absorption spectrum and resulting colour.

A diagram showing the 3d orbital ligand/crystal field splitting effect in the hexaaquachromium(III) complex ion and the excitation by a visible light photon of an electron from a lower to a higher 3d orbital level.

There must be at least two possible excitations to produce two absorbance peaks in the visible region.

For more see the chemistry of chromium

 

(b) The colour of the hexaamminechromium(III) ion

If you change the ligand, you usually change the ∆E_elec , in this case excitation of an electron from the lower 3d orbital to a higher 3d orbital in the hexaamminechromium(III) complex ion.

The change of ligand is from water to ammonia - both lone pair donors to form dative covalent bonds with the central metal ion, the result is the dark green coloured hexaamminechromium(III) ion, formed a simple ligand exchange reaction.

[Cr(H₂O)₆]³⁺_(aq) + 6NH_3(aq) ===> [Cr(NH₃)₆]³⁺_(aq) + 6H₂O_(l)

(c) The colour of the hexaaquachromium(II) ion

The blue hexaaquachromium(II) ion, [Cr(H₂O)₆]²⁺_(aq)

The change in oxidation state changes the electronic environment of the theoretical central Cr²⁺ ion in the complex compared to the Cr³⁺ ion.

This in turn changes the ∆E_elec energy needed to excite an electron from the lower 3d orbital to a higher one from the field splitting effect of the six ligands surrounding the central Cr²⁺ ion.

The ligand field electronic splitting diagram for the aqueous Cr²⁺ ion.

For more see the chemistry of chromium

 

(d) The uv-visible light absorption spectra of the dichromate(VI) ion and chromate(VI) ion

The yellow chromate(VI) ion and orange dichromate(VI) ion form an equilibrium, one species dominates depending on the pH of the solution.

2CrO₄^2–_(aq) +  2H⁺_(aq) Cr₂O₇^2–_(aq)  +   H₂O_(l) 

There is no change in oxidation state, but the electronic environment of the central chromium 'ion' is changed so they have different uv-visible absorption spectra and the different field splitting ∆E_elec values result in the two different colours.

The chromate(VI) ion strongly absorbs in the violet region and gives a yellow colour.

The dichromate(VI) ion absorbs in the blue and produces an orange colour.

For more see the chemistry of chromium

 

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