By SeanPublished On: August 6, 2018Last Updated: November 23, 2022
Someone commented on one of our Instagram posts that the only reason they had studied chemistry was because of the ‘amazing colors’. While this is solid motivation for many of us, an organic chemist might find it harder to relate. Most small organic molecules are either a white powder or are colorless crystals. For a science that is known to be filled with color, we take a look at why some of these organic compounds refuse to conform.
Organic compounds are usually white or colorless because the molecules do not absorb and emit photons with wavelengths that correspond to visible light. This means that sunlight that reaches them reflects all visible light (appearing white) or passes through organic crystals (appearing colorless). Organic molecules typically require higher energy photons in the ultraviolet range to excite them.
Where Color Comes From
Photons with Energy
To understand the absence of color in organic compounds, we must first define color. From a human’s perspective, we perceive color when photons reach our cone cells. These are types of receptors in our eyes that, in turn, produce signals for our brains to identify different colors.
Photons, in this case, can be defined as tiny ‘packets’ of energy, with a corresponding wavelength that is inversely proportional to its energy (i.e. lower energy = larger wavelength).
Light from the sun is made up of photons with many different wavelengths across the electromagnetic spectrum. Some of these correspond to the colors we see, while others aren’t even visible to the human eye, such as infrared, ultraviolet and even X-rays.
The Colors of Sunlight
The photons with wavelengths corresponding to visible colors (between 380 and 780 nm) combine to produce the sunlight we all know and love. The sun emits yellow photons at the highest intensity, which is why sunlight isn’t completely white.
Our atmosphere helps by scattering shorter wavelengths of visible light (violet and blue); longer wavelengths like yellow pass through unhindered. (This also explains the color of our blue sky!) How light is created—well converted, because energy—in the sun involves blackbody radiation when something is heated to a extreme temperatures, but that’s a topic for another day.
Color Through Wavelengths
Okay, so the sun produces white light. But we see a whole spectrum of different colors in the world around us (except blue, nature’s most hated color)!
The myriad of colors we observe in our everyday lives comes from electrons! When photons strike a molecule, its electrons may decide to move up an energy level or two, depending on the photon’s energy.
If the photon has sufficient energy, it can ‘kick’ a molecule’s electron up its energy level. The energy required to do so is specific for every molecule, influenced by its orbitals.
Once an electron is kicked up an energy level, it wants to move back down to its more relaxed ground state by losing energy in the form of a photon. Through a series of steps, the molecule emits a photon of a single wavelength.
We see this as color if the photon has a wavelength that falls within the visible region. The colors and their corresponding wavelength are in the chart below. Remember that lower energy photons have a larger wavelength (it’s an inversely proportional relationship).
Why Organic Compounds Are White/Colorless
Boring Organic Compounds
White powders look white because when light hits these molecules, they are scattered at all angles without being absorbed. In their crystalline form, this light simply passes through —again, without being absorbed—rendering them transparent. For a compound to possess color, some form of light absorption must occur.
It turns out that for smaller molecules, visible light just doesn’t carry enough energy to ‘kick’ its electrons up their energy levels. We can blame this on molecular orbitals and how smaller molecules have a larger energy level gap that needs to be overcome for an electron to be kicked up.
Organic compounds are usually small molecules, usually requiring higher energy light in the ultraviolet range, which is why our eyes don’t pick up on their transitions.
Double Bonds = Color
Hang on, so if I decrease this energy gap and therefore the energy required to excite electrons, maybe visible light could excite my compound? Exactly!
A way to reduce the energy is to introduce double bonds; C=C double bond electron orbitals are higher in energy than their C-C single bond counterparts. Even better would be to conjugate (alternate in the molecule) these double bonds to absorb energy at an even higher wavelength.
Size Matters Too!
To give your white organic compounds color, you can also choose to make it very big. How does this work? From a molecular orbital perspective, solving Schrodinger’s equation for an ‘electron in a box’ (a great pastime, by the way) gives:
Notice that increasing the size L of the ‘box’ (the molecule in this case), leads to a smaller value for energy E. With small organic molecules, E is high because L is small. To put it simply: mashing many orbitals together by increasing the size of a molecule will lower the overall energy, allowing the highest energy electron to transition more easily!
Indeed, many dyes that we use today are structured both conjugated and large, meaning that visible light is enough to make them show off their pretty colors. And yup, they’re actually classified as ‘organic’ dyes.
A Transition to Inorganic Chemistry
Organic chemistry mostly deals with smaller elements such as carbon and nitrogen; inorganic chemistry is all about transition metals that are much larger and contain more electrons. This means transition metal complexes often transition with a whole range of beautiful colors!
So, while organic compounds are white and colorless, on the other side of the spectrum (ha!), inorganic chemists get to spend their days running out of colors to label their compounds with. Chartreuse, anyone?
Lapedes, D. N. (1976). Dictionary of Scientific and Technical Terms McGraw-Hill.
Sklar, A. L. (1937). Theory of color of organic compounds. The Journal of Chemical Physics, 5(9), 669-681.
About the Author
Sean is a consultant for clients in the pharmaceutical industry and is an associate lecturer at La Trobe University, where unfortunate undergrads are subject to his ramblings on chemistry and pharmacology.