A-level Applied Science/Colour Chemistry/Colour

The Chemistry of Colour

Coloured chemicals absorb electromagnetic waves in the visible part of the spectrum. The absorbed energy

causes changes in the energy of the molecules’ electrons. The electrons change from a ‘ground

state’ to an ‘excited state’.

Most transitions are not caused by visible light. Many absorb ultra-violet

radiation. Chemicals which absorb UV radiation are colourless (unless they fluoresce). The

energy changes when molecules of a coloured compound and of a colourless compound are illustrated below:

UV (right) and visible absorbance.

Remember that the apparent colour is caused by absorbing photons of a complementary colour. A

blue compound is blue because it absorbs yellow light.

The hexagon has the three primary colours of light (RGB), and their complementary colours (CYM).

Chromophores

Chemical structures which have excited states corresponding to visible light are called

chromophores. There are two main types:

1. Transition Metal Complexes.

Transition metals form complex ions – the metal binds to small molecules

or anions called ligands. The ligands allow the electrons of the metal ion to enter an excited

state if the electrons absorb a photon of visible light.

e.g. tetrachlorocuprate (II) and hexaaquacopper (II) ions:

The partially-occupied d-orbitals of transition metal compounds are important in giving colour to

transition metal complexes. See diagram (which could represent V⁺², Cr⁺³,

Mn⁺⁴, etc.):

①. In an uncomplexed ion, all the d-orbitals have the same energy.

②. When ligands surround the ion, the negative charges of the ligands make the d-orbitals

less stable (higher energy).

③. Critically, the ligands will come closer to some d-orbitals than to others. Typically,

two or three of the orbitals will be destabilised more than the remainder.

An electron in one of the lower d-orbitals can acquire the energy to be excited into a higher

d-orbital:

This mechanism allows transition metal complexes to absorb photons of visible light.

2. Conjugated/Delocalised Electron Systems.

When single and double bonds alternate, the electrons in the double bonds can enter an excited

state if they absorb a photon of visible light. e.g. β-carotene (above) has ten conjugated C=C bonds:

The diagram above shows the excitation energies of conjugated aldehydes. n is the number of C=C

double bonds which are conjugated. The simplest (n=1) is CH₃-CH=CH-CH=O.

Note how the excitation energy is lower with higher numbers of conjugated bonds.

n	Wavelength (nm)	Energy (kJ mol−1)
1	220	544
2	270	443
3	312	384
4	343	349
5	370	324
6	393	305
7	415	289

Chromophores of dye molecules often contain unsaturated groups such as >C=O and -N=N-, which are

part of a conjugated bonding system, usually involving aromatic rings. Chrysoidine, a

basic dye, is shown below:

chrysoidine

Note how the –N=N- group is just the centre of a conjugated system which extends across all

twelve carbon atoms and includes seven double bonds. All

azo dyes contain the -N=N- arrangement.

Auxochromes: Attached to the chromophore are two -NH₂ groups which interact with

the chromophore to modify the orange colour. A group of atoms attached to a chromophore which

modifies the ability of that chromophore to absorb light is called an auxochrome. They can modify

or enhance the colour of the dye. Examples: -OH, - NH₂, aldehydes.

Added functional groups can also:

• alter the solubility of the dye in water or other solvents.
• bind the dye molecules to cloth, paper or other substrates.

Reactive Red 6 has both organic (conjugated system) and inorganic (complex metal ion) chromophores.

References

Notes on colour chemistry by elecuter.

See also

Notes on colour chemistry at Bristol University.