Why transition metals exhibit so many colours in their
compounds is based on the quantum structure of the transition metal atom. All
transition metals have their valence electrons in the d-subshell, and in the isolated, atomic state the five d-orbitals have the same energy level,
i.e. they are degenerate. However, in a compound/complex ion the ligands
attached to the transition metal ion cause these d-orbitals to split in terms of energy, forming one group of d-orbitals with a lower energy and
another group of d-orbitals
with a higher energy. This happens due to the repulsion between the metal electrons and the electrons from the attached chemical.
Electrons in the lower energy d-orbitals can ‘jump’ to the higher energy d-orbitals via absorption of energy in the form of photons. To make
this jump the electron must generally absorb photons having energies of 2.8 to 5.0
x 10^-19 joules, which corresponds to a wavelength range of 400-700 nanometres,
the same wavelength range of the visible light spectrum. Thus, as visible light
passes through a transition metal compound, photons of a specific wavelength (and
thus colour) are absorbed by the compound, and the remaining photons pass
through the compound to give a coloured light.
For example, a solution
of copper (II) ions is blue as red light is absorbed by the d-electrons in the copper (II) ion (or
more accurately, the copper (II) hexahydrate ion). A solution of iron (II) ions
is green as in this case, the aqueous iron (II) ion absorbs reddish-blue light.
In this activity, CHS Science and Maths will experiment with
chromium compounds, creating as many types of coloured compounds as we can
through various physical and chemical reactions.
Follow the procedure here:
Follow the procedure here:
Starting with the commonly-used chromium (III) chloride (CrCl3) as a base chemical, the following reactions are conducted:
Dehydration - the dark green hydrated chromium (III) chloride hexahydrate is dehydrated via heat to become the violet anhydrous salt:
Amination - dark green aqueous chromium (III) chloride is aminated via concentrated ammonia solution to form the pink hexaammine chromium (III) complex ion, with chromium (III) hydroxide as an intermediate:
Oxidation (alkaline) - first the chromium (III) solution is made basic with the addition of sodium hydroxide solution. Excess sodium hydroxide solution is added to redissolve the insoluble chromium (III) hydroxide as shown below:
Following this hydrogen peroxide is added and the solution warmed. This oxidises the dark green chromate (III) ion to the yellow chromate (VI) ion:
Addition of acid - Adding acid to the alkaline chromate (VI) solution increases the concentration of H+ ions, turning the yellow chromate (VI) ion into the orange dichromate (VI) ion.
As shown by the double-headed arrow, the chromate (VI) and dichromate (VI) ions exist in dynamic equilibrium, with their concentration determined by the acidity (pH) of the solution. By Le Chatelier's principle, a low pH (high H+ conc.) encourages the formation of the orange dichromate (VI) ion whereas a high pH (low H+ conc.) encourages the formation of the yellow chromate (VI) ion.
Addition of peroxide/Ligand exchange reaction
The addition of acidified hydrogen peroxide to a solution of either yellow chromate (VI) or orange dichromate (VI) ions yields the blue chromium (VI) oxide peroxide. It is unstable when dissolved in water, so to obtain a sample of the peroxide the reaction has to be carried out in an organic solvent. In this case, a chromate (VI) compound dissolved(solvated) in ethanol is used.
Addition of chlorine/Ligand exchange reaction
Dehydrated, strongly acidic orange dichromate (VI) ions react with concentrated hydrochloric acid to form the powerfully oxidising red compound of chromyl chloride. Due to the corrosive and toxic nature of the reactants and products involved this reaction is best carried out in a fume hood. Chromyl chloride reacts spontaneously with water, thereby requiring that it be stablised in a dry medium.
