Colors of chemicals

From Wikipedia, the free encyclopedia

This article or section is missing citations or needs footnotes.
Using inline citations helps guard against copyright violations and factual inaccuracies. (January 2008)

All chemical compounds have colors, and in some cases these are distinctive or useful for identification. The study of chemical structure by means of energy adsorption and release is generally referred to as spectroscopy.

1 Theory
2 Colors by wavelength
3 Examples

[edit] Theory

All chemical compounds, including atoms, are capable of absorbing and releasing energy. The amount(s) of energy (quanta) absorbed and released is determined by the quantum structure of the chemical. The release of energy visible to the human eye spans the wavelengths 380 nm to 760 nm and is commonly referred to as color. The relationship between energy and wavelength is determined by the equation:

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

where E is the energy of the quanta (photon), h is Planck's constant, $λ$ is the wavelength and c is the speed of light. The relationship between chemical structure and energy can be understood using atomic orbital, molecular orbital, or Ligand Field Theory. In organic compounds the color can be determined by the difference between the Highest Occupied Molecular Orbital and the Lowest Unoccupied Molecular Orbital. The energy absorbed and/or released does not directly correlate to what humans perceive as color. The absorption of a particular wavelength of light effectively subtracts it from the full visible spectrum, and what we see is the complementary color, made up of the other visible wavelengths. Beta-carotene has maximum absorption at 454 nm (blue light) consequently what visible light remains appears orange.

[edit] Colors by wavelength

Below is a rough table of wavelengths, colors and complementary colors.

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

[edit] Examples

[edit] Ions in aqueous solution

Name	Formula	Color
Alkali metals	M⁺	None
Alkaline earth metals	M²⁺	None
Scandium (III)	Sc³⁺	None
Titanium (III)	Ti³⁺	Violet
Titanyl	TiO²⁺	None
Vanadium (II)	V²⁺	Lavender
Vanadium (III)	V³⁺	Dark grey/green
Vanadyl	VO²⁺	Blue
Pervanadyl	VO₂⁺	Yellow
Metavanadate	VO₃^-	None
Orthovanadate	VO₄^3-	None
Chromate	CrO₄ ^2-	Colorless or Yellow(sometimes)
Dichromate	Cr₂O₇^2-	Orange
Manganese (II)	Mn²⁺	Light pink
Manganate (VII) (Permanganate)	MnO₄^-	Deep violet
Manganate (VI)	MnO₄^2-	Dark green
Manganate (V)	MnO₄^3-	Deep blue
Iron (II)	Fe²⁺	Light blue
Iron (III)	Fe³⁺	Yellow/brown
Cobalt (II)	Co²⁺	Light red
Nickel (II)	Ni²⁺	Light green
Nickel-ammonium complex	Ni(NH₃)₆²⁺	Lavender/blue
Copper (II)	Cu ²⁺	Blue
Copper-ammonium complex	Cu(NH₃)₄²⁺	Royal Blue
Zinc (II)	Zn²⁺	None
Silver	Ag⁺	None

It is important to note, however, that elemental colors 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 redish hue, whilst V₂O₃ appears black.

[edit] Salts

Predicting the color 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's 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
Copper (II) sulfate	CuSO₄	Blue
Copper (II) sulfate pentahydrate	CuSO₄ · 5H₂O	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