Color of chemicals

The colour of chemicals is a physical property of chemicals that in most cases comes from the excitation of electrons due to an absorption of energy performed by the chemical. What is seen by the eye is not the colour absorbed, but the complementary colour from the removal of the absorbed wavelengths.

The study of chemical structure by means of energy absorption and release is generally referred to as spectroscopy.

Theory

The UV-vis spectrum for a compound that appears orange in Dimethylformamide

All atoms and molecules are capable of absorbing and releasing energy in the form of photons, accompanied by a change of quantum state. The amount of energy absorbed or released is the difference between the energies of the two quantum states. There are various types of quantum state, including, for example, the rotational and vibrational states of a molecule. However the release of energy visible to the human eye, commonly referred to as visible light, spans the wavelengths approximately 380 nm to 760 nm, depending on the individual, and photons in this range usually accompany a change in atomic or molecular orbital quantum state. The perception of light is governed by three types of colour receptors in the eye, which are sensitive to different ranges of wavelength within this band.

The relationship between energy and wavelength is determined by the equation:

$E = hf = \frac{hc}{\lambda} \,\!$

where E is the energy of the quantum (photon), f is the frequency of the light wave, h is Planck's constant, $\lambda$ is the wavelength and c is the speed of light.

The relationships between the energies of the various quantum states are treated by atomic orbital, molecular orbital, and Ligand Field Theory. If photons of a particular wavelength are absorbed by matter, then when we observe light reflected from or transmitted through that matter, what we see is the complementary colour, made up of the other visible wavelengths remaining. For example, beta-carotene has maximum absorption at 454 nm (blue light), consequently what visible light remains appears orange.

Colours by wavelength

Below is a rough table of wavelengths, colours and complementary colours. This utilizes the scientific CMY and RGB colour wheels rather than the traditional RYB colour wheel.^[1]

Wavelength (nm)	Colour		Complementary Colour
400-424		Violet		Green-yellow
424-491		Blue		Yellow
491-570		Green		Violet
570-585		Yellow		Blue
585-647		Orange		Cyan-Blue
647-700		Red		Cyan

This can only be used as a very rough guide, for instance if a narrow range of wavelengths within the band 647-700 is absorbed, then the blue and green receptors will be fully stimulated, making cyan, and the red receptor will be partially stimulated, diluting the cyan to a greyish hue.

By category

The vast majority of simple inorganic (e.g. sodium chloride) and organic compounds (e.g. ethanol) are colorless. Transition metal compounds are often colored because of transitions of electrons between d-orbitals of different energy. (see Transition metal#Coloured compounds). Organic compounds tend to be colored when there is extensive conjugation, causing the energy gap between the HOMO and LUMO to decrease, bringing the absorption band from the UV to the visible region. Similarly, color is due to the energy absorbed by the compound, when an electron transitions from the HOMO to the LUMO. Lycopene is a classic example of a compound with extensive conjugation (11 conjugated double bonds), giving rise to an intense red color (lycopene is responsible for the color of tomatoes). Charge-transfer complexes tend to have very intense colors for different reasons.

Examples

Ions in aqueous solution

Name	Formula	Colour
Alkali metals	M^+	Colourless
Alkaline earth metals	M^2+	Colourless
Scandium(III)	Sc^3+	Colourless
Titanium(III)	Ti^3+	Violet
Titanium(IV)	Ti^4+	Colourless
Titanyl	TiO^2+	Colourless
Vanadium(II)	V^2+	Lavender
Vanadium(III)	V^3+	Dark grey/green
Vanadyl	VO^2+	Blue
Hypovanadate	V4O9^2-	Brown
Pervanadyl	VO2^+	Yellow
Metavanadate	VO3^−	Colourless
Orthovanadate	VO4^3−	Colourless
Chromium(III)	Cr^3+	Blue-green-grey
Chromate	CrO4^2−	Yellow
Dichromate	Cr2O7^2−	Orange
Manganese(II)	Mn^2+	Very light Pink
Manganate(VII) (Permanganate)	MnO4^−	Deep violet
Manganate(VI)	MnO4^2−	Dark green
Manganate(V)	MnO4^3−	Deep blue
Iron(II)	Fe^2+	Very pale green
Iron(III)	Fe^3+	Very pale violet/brown/yellow
Iron(III) tetrachloro complex	FeCl4^-	Yellow/brown
Cobalt(II)	Co^2+	Pink
Cobalt(III)-ammine complex	Co(NH3)6^3+	Yellow/orange
Nickel(II)	Ni^2+	Light green
Nickel(II)-ammine complex	Ni(NH3)6^2+	Lavender/blue
Copper(I)-ammine complex	Cu(NH3)2^+	Colourless
Copper(II)	Cu^2+	Blue
Copper(II)-ammine complex	Cu(NH3)4^2+	Royal Blue
Copper(II) tetrachloro complex	CuCl4^2−	Yellow/green
Zinc(II)	Zn^2+	Colourless
Silver(I)	Ag^+	Colourless
Silver(III) in conc. HNO3	Ag^3+	Dark brown

^[2] It is important to note, however, that elemental colours will vary depending on what they are complexed with, often as well as their chemical state. An example with vanadium(III); VCl₃ has a distinctive reddish hue, whilst V₂O₃ appears black.

Salts

Predicting the colour of a compound can be extremely complicated. Some examples include: Cobalt chloride is pink or blue depending on the state of hydration (blue dry, pink with water) so it is used as a moisture indicator in silica gel. Zinc oxide is white, but at higher temperatures becomes yellow, returning to white as it cools.

Name	Formula	Color	Picture
Copper(II) sulfate	CuSO4	White
Copper(II) sulfate pentahydrate	CuSO4 · 5H2O	Blue
Copper(II) benzoate	C14H10CuO4	Blue
Cobalt(II) chloride	CoCl2	Deep blue
Cobalt(II) chloride hexahydrate	CoCl2 · 6H2O	Deep magenta
Manganese(II) chloride tetrahydrate	MnCl2 · 4H2O	Pink
Copper(II) chloride dihydrate	CuCl2 · 2H2O	Blue-green
Nickel(II) chloride hexahydrate	NiCl2 · 6H2O	Green
Lead(II) iodide	PbI2	Yellow

Ions in Flame

Flame Tests on cations for Alkali, Alkali Earth Metals, and Hydrogen (see atomic spectroscopy) (see also Flame test)

Metals

Name	Formula	Color
Potassium	K	Lilac/Purple
Sodium	Na	Yellow/orange
Lithium	Li	Red
Cesium	Cs	Blue
Calcium	Ca	Brick Red
Strontium	Sr	Red[3]
Barium	Ba	Green/Yellow

Gases

Name	Formula	Color
Hydrogen	H2	colorless
Oxygen	O2	colorless
Fluorine	F2	very pale yellow/brown
Chlorine	Cl2	greenish yellow
Bromine	Br2	red/brown
Iodine	I2	dark purple
Chlorine dioxide	ClO2	intense yellow
Dichlorine monoxide	Cl2O	brown/yellow
Nitrogen dioxide	NO2	dark brown
Dinitrogen tetroxide	N2O4	colorless
Trifluoro nitroso methane	CF3NO	deep blue
Diazomethane	CH2N2	yellow

Bead tests

A variety of colours, often similar to the colours found in a flame test, are produced in a bead test, which is a qualitative test for determining metals. A platinum loop is moistened and dipped in a fine powder of the substance in question and borax. The loop with the adhered powders is then heated in a flame until it fuses and the colour of the resulting bead observed.

Metal[4]	Oxidizing flame	Reducing flame
Aluminium	colourless (hot and cold), opaque	colourless, opaque
Antimony	colourless, yellow or brown (hot)	gray and opaque
Barium	colourless
Bismuth	colourless, yellow or brownish (hot)	gray and opaque
Cadmium	colourless	gray and opaque
Calcium	colourless
Cerium	red (hot)	colourless (hot and cold)
Chromium	Dark yellow (hot), green (cold)	green (hot and cold)
Cobalt	blue (hot and cold)	blue (hot and cold)
Copper	green (hot), blue (cold)	red, opaque (cold), colourless (hot)
Gold	golden (hot), silver (cold)	red (hot and cold)
Iron	yellow or brownish red (hot and cold)	green (hot and cold)
Lead	colourless, yellow or brownish (hot)	gray and opaque
Magnesium	colourless
Manganese	violet (hot and cold)	colourless (hot and cold)
Molybdenum	colourless	yellow or brown (hot)
Nickel	brown, red (cold)	gray and opaque (cold)
Silicon	colourless (hot and cold), opaque	colourless, opaque
Silver	colourless	gray and opaque
Strontium	colourless
Tin	colourless (hot and cold), opaque	colourless, opaque
Titanium	colourless	yellow (hot), violet (cold)
Tungsten	colourless	brown
Uranium	Yellow or brownish (hot)	green
Vanadium	colourless	green

References

1. http://www.sapdesignguild.org/resources/glossary_color/index1.html
2. http://www.sanjuan.edu/webpages/dkrenecki/files/ColorsofSubstances2.pdf
3. Flame Tests at chemguide.co.uk
4. CRC Handbook of Chemistry and Physics. CRC Press. 1985. ISBN 0-8493-0466-0.