A-level Applied Science/Colour Chemistry/Colour

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The Chemistry of Colour[edit | edit source]

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[edit | edit source]

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:

Tetrachlorocuprate (ii).png

Hexaaquacopper (ii).png

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.):

Ligand field splitting.png

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


Excitation in TM complex.png

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:

Excited aldehydes.png

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:


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) chromopohores.

References[edit | edit source]

Notes on colour chemistry by elecuter.

  1. Streitwieser, A & Heathcock, CH (1985) Introduction to organic chemistry (3rd ed) p 628, Macmillan, New York

See also[edit | edit source]

Notes on colour chemistry at Bristol University.