We all know colors, they are everywhere, they define concepts. Trees? Green. Sky? Blue. Fire? Orange. We can also play and have fun with colors by painting. Colors give us the tools to be creative, to explore the artistic world and our own imagination.
The color wheel (see Figure 1) is a simple circle of colors invented by Isaac Newton in 1666. Since then, artists and scientists have created multiple versions of this model to understand color and color combinations.
Colors can be classified as primary, secondary and tertiary. Primary colors are red, blue and yellow. Traditional color theory states that they are the only three colors that cannot be created by combining others. Secondary colors are results of mixing the primary colors in different ratios, for instance green, orange and purple. Tertiary colors are obtained by mixing primary with secondary colors, such as red-orange, blue-green, yellow-purple, etc. Black and white are used to create shades, to either enlighten or darken colors.
Maybe someone has wondered how do we produce different colors? Which components should we mix to synthesize different paints? In other words, what determines a color? The answer is pigments.
Pigments define colors. So… what are pigments?
Well, pigments are usually inorganic compounds. Many of the pigments are colored due to their capacity to absorb visible light. When a compound absorbs a certain wavelength of visible light, what we see is the color that results from mixing the rest of wavelengths in the spectrum. If the compound absorbs the green wavelength, it emits the violet wavelength and therefore we see violet color!
Many pigments are based on transition metals, which are characterized by having d-orbitals. When a metal is not bond to anything else, these d-orbitals are degenerate (they all have the same energy). However, when a transition metal binds to ligands, the situation changes. Due to the different symmetries, the d-orbitals split apart and become non-degenerate (they have different energy levels), see Figure 2.
D-orbitals are often partly occupied, meaning that electrons can move from one d-orbital to another d-orbital. In a naked metal, all d-orbitals are identical (i.e. degenerate) and movements of electrons do not require any energy input (see left-hand side of Figure 2). This is where the other relevant part of the pigments enters: ligands. When a ligand (or ligands) bind to a metal, the d-orbitals break their symmetry and change their energies. If you check the right-hand side of Figure 2, the degeneracy got broken to adopt an octahedral geometry (this is just only one possible geometry, there are plenty more, hence the richness of colors!). Now, the metal will need an energy input to promote one electron from one orbital to a different one.
From where do metals get the energy to promote electrons? From electromagnetic radiation! That is, light! However, the light spectrum is complex and contains different kinds of radiation. Actually, most substances are only able to absorb frequencies of radiation which are outside the visible light spectrum, for example they might be able to absorb infrared or ultraviolet radiation. This means that these compounds reflect all other types of radiation, including the full spectrum of visible light, and our eyes see a mixture of all the colors: red, green, blue, violet, yellow… = white light. This explains why, for instance, most organic compounds are white.
However, transition metals are special. The energy difference between the non-degenerate d-orbitals usually correspond to the energy of radiation of certain wavelengths in the visible light spectrum. This means that when we look at the metal complex, we don’t see the entire visible light spectrum, but only a part of it, one wavelength. For example, if the electrons in an octahedral (Oh) metal complex are able to absorb blue light and get promoted from the low-energy orbitals to the high-energy orbitals, the compound will reflect all colors except of blue. Upon relaxation (going back to the low energy orbital), the electron will emit the complementary wavelength which eventually determines the color of the compound. By using the color wheel (see Figure 3), we can predict the complementary color for each frequency. Following with the example, if a compound absorbs blue light, it will emit orange light and therefore look orange to us.
Not all transition metal complexes are colorful. For example, copper sulfate is a bright blue compound, however zinc sulfate, which is also a transition metal, is white. The reason behind this is because zinc’s d-orbitals are completely filled up with electrons, meaning that it is not possible for any electron to make a d –> d transition.
In figure 4 there are examples of inorganic compounds that give the respective colors. Notice that all of them (except black), contain transition metals.
To all artists out there: next time you paint something, think of the d-orbitals! 😉
Written by Alba Monferrer, PhD student in Prof. Hendrik Dietz’s lab at TUM.