Science of plant color (and color in general)

Why do pumpkins look orange? BETA-CAROTENE! Orange-yah glad you asked? Hopefully readers will delight if I discuss how these pumpkins steal sunlight! (but only specific wavelengths of it!)

blog form: http://bit.ly/carotenoidschromaphores & http://bit.ly/lightleafcolor

Light is “just” little packets of energy called photons traveling in waves. You can think of them kinda like baseballs wiggling through the air. The balls are all thrown at the same speed (the speed of light) so, when a group of balls is thrown by the pitcher (light source), all the balls will reach the catchers at the same time. But some of the balls have more energy, so they take a wavier route, oscillating up and down more as they travel giving you closer together peaks (higher frequency (f), shorter wavelength (λ)).If you think of slalom racing, it’s like you have a bunch of skiers that all get to the finish line at the same time, but the skiers with higher energy make more S’s during that time.

White light has a combination of “all the colors” - so it’s like the sun is a pitcher throwing lots of baseballs at us, but molecules can “catch” some of the balls and “hide them” so that things look colored. (ROYGBIV is white, but OYGBIV or ROYBIV, etc. isn’t).

Molecules can only catch those photon “baseballs” if they have the right “catcher’s mitt.” These “mitts” - the photon-absorbing molecules (or parts of bigger molecules) - are called chromophores, and different ones absorb different photons (catch different balls). Why?

They can only “catch a ball” if the photon’s energy is just right for exciting an electron in a molecule. An electron is a type of negatively-charged subatomic particle that molecules use to interact & bond with one another. They can live in different “orbitals” and the highest energy electrons live in the furthest orbitals from the nucleus (the central hub of the atom where the positively-charged protons and the neutral neutrons live). More here: http://bit.ly/33RznDA

It’s kinda like the “electron housing” landlords charge more in rent (in the form of energy) to live further from the central nucleus - out in the outer atomic suburbs. If they get some extra energy income, electrons can afford to move out from a “ground state” to a higher orbital - a so-called “excited state.” But they can only do this if they pay “exact” change, so they can only absorb photons that have an energy amount equal to the difference between the higher orbital and the current one.

So who determines the rent? The atoms making up the molecules themselves (with some influence from the local environment). Different molecules have different housing arrangements and different differences in rent between the different housing levels, so “moving up” costs different amounts and, as a result, different molecules absorb photons with different amounts of energy.

And if we go back to that energy-wavelength-frequency-color relationship, this means that different molecules are absorbing different colors of light. And since they’re stealing different slices of the rainbow, the leftovers they leave us with look different.

For example, if the orbitals are really far apart, it takes a lot of energy for an electron to move to those outer suburbs, so they’ll absorb high-energy (and thus high-frequency, short wavelength) light like blue light. And since they’re stealing that blue light, the light they leave us with (either going through it (transmission) or bouncing off (reflection)) looks yellow-orange-y. (To see what color something will look if it absorbs a color, look across from the absorbed color in a color wheel).

If a chromophore’s orbitals are closer together, the difference in rent is smaller, so they’ll absorb lower-energy photons (lower-frequency, longer-wavelength) - for example, they might absorb red red light & look green.

I put “exact” energy quotes because there’s a little wiggle room because electrons can have different “vibrational levels” within orbitals and stuff. So if you look at an absorption spectrum for a molecule (which shows you what wavelengths the molecule “steals”) - instead of sharp peaks you see more of bell curves - peaks at the maximum absorption wavelength and some absorption on either side, petering out the further you get from that “optimal wavelength.” And you’ll likely also see multiple peaks because, for example, molecules can have multiple chromophores in the same molecule,

Most of the molecules I study are invisible to us because their orbitals are arranged such that the photons they absorb are outside of the visible range (either too low in energy or too high in energy for us to see) - there’s only a segment of the electromagnetic radiation (EMR) spectrum that we can actually see. We call this the visible spectrum - it spans from wavelengths of about 380-740 nanometers (nm)(a nanometer is 1 billionth of a meter so even the longest of these are pretty dang short).

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