Colors of chemicals

Most research into chemical structure is some form of spectroscopy. Spectroscopy is the careful measurement of the absorbance by unknown materials by a range of electromagnetic radiation. As humans, our sensitivty to electromagnetic radiation is limited to visible light. The chemist uses instruments to extend the detectable range, but in a broad sense, the question is 'what color is this?'.

Fundamental to physics and chemistry is the rather non-intuitive idea of the electromagnetic spectrum. The wavelengths correspond to energies (i.e. the energy of one photon). The visible light, of which we are so accustomed, is about in the middle of the energy range, at ca. 400 to 800 nm.

There are also energies associated with chemical structures; in fact, discreet energies may be associated with every facet of chemistry.

Chemical behavior is described by the model of quantum mechanics. This model says that molecular structures go from one form to another with no intermediate. There is no way to intuitively understand this. Students and teachers can almost always relate new learning to past experience, but not so in this case.

These discontinuous events, however, in chemical structure have corresponding specific energy changes. And these energy changes have a corresponding electromagnetic frequency.

So, what causes color? How about some examples.

The working part of hemoglobin is a heme group. The structure is such that the electrons circle around in quite large clouds. If the heme is in the form of oxyhemoglobin the clouds are a little smaller. The bluer wavelengths are absorbed and the longer red wavelengths predominate. The blood is red. If the hemoglobin is not oxygenated, the clouds are a little larger. The redder wavelengths are absorbed and the blood appears bluish.

Acid/base indicators are familiar to most students. Phenolphthalein is a large organic compound. Its structure changes depending on the acidity of the environment. In an acid environment it has no electron clouds "small enough" to absorb visible light. In a basic environment, the structure changes just a little and the electon clouds shrink just a little and the phenolphthalein turns pink.

Permanganate, MnO₄^-, is intensely purple/violet. Even a very diluted solution will hardly be translucent. The reason is that Permanganate-ions absorb green light, and what is left of the spectrum appears violet. Permanganate is imagined to be a so-called charge transfer complex, which means that the electrons from the oxygen's full 2p-orbitals are temporarily "lifted" into the manganese's empty 4s- and 3d-orbitals by photons, which in that process are absorbed.

Visible light corresponds to the energy absorbed in many large organic structures. Evolution has converged on these as the most useful wavelengths for our eyes.

It is all very interesting. What causes the colors of compounds? The structure causes the color. Changes in color are a reflection of changes in structure.

Some examples

Ions in aqueous solution

name formula colour

• alkali metals M⁺ none
• alkaline earth metals M²⁺ none
• scandium 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
• chrome(II) Cr²⁺ blue
• chrome(III) Cr³⁺ violet, mostly green when coordinated
• chromate CrO₄ ^2- yellow
• dichromate Cr₂O₇^2- orange
• manganese(II) Mn²⁺ VERY light pink
• permanganate MnO₄^- deep purple/violet
• manganate (VI) MnO₄ ^2- very dark green
• manganate (V) MnO₄ ^3- blue
• iron (II) Fe²⁺ very light green
• iron (III) Fe³⁺ yellow/brown
• cobalt Co²⁺ light red, pink
• nickel Ni²⁺ fairly light green
• nickel-ammonia complex [Ni(NH₃)₆]²⁺ lavender/blue
• copper Cu ²⁺ blue
• copper-ammonia complex [Cu(NH₃)₄]²⁺ deep blue
• Zinc Zn²⁺ none
• Silver Ag⁺ none

Salts

• copper vitriol (sulfate) CuSO₄ none
• copper vitriol hydrate CuSO₄·5H₂O blue
• cobalt chloride CoCl₂ deep blue
• cobalt chloride hydrate CoCl₂·6H₂O light red

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