Don’t Fall into the trap of thinking science is pretty boring – leave it to the bumbling biochemist to make leaf color-changing pretty but not boring! There’s cool chemistry behind the colors you (finally) see! (once all that chlorophyll’s gone…)

Sunlight contains a lot of energy – and the leaves of plants are designed to capture and use the energy from sunlight (along with carbon dioxide from the air) to make sugar (of the glucose variety) in a process called photosynthesis. But in order to make food from it, plants need to be able to capture it – so they produce chemicals called chlorophyll which can absorb sunlight – but they only absorb specific “slices of the rainbow” – which don’t include green – so green is seen! But chlorophylls aren’t the only color-absorbers (chromophores) they make. And, as the days get shorter, plants stop making chlorophyll, allowing underlying pigment colors to shine through. To understand how they do this it helps to think about what sunlight – or any light for that matter – actually is.

“Light” (aka ElectroMagnetic Radiation (EMR)) can be thought of as little packets of energy called photons traveling in waves. I like to think of these photons as baseballs being thrown wiggly-ly. And what happens to light when it hits something else depends on what “catchers” are there. Different molecules have different “catcher’s mitts” – parts of molecules called CHROMOPHORES, that, because of their electronic housing situations which we’ll look at more later, can only catch (absorb) photons with certain energies. 

That energy is directly related to the color we perceive – different colors of light have photons with different amounts of energy. Higher-energy photons give you colors on the backside of ROYGBIV (e.g. blues & purples) and lower-energy photons correspond to the frontside (reds & oranges). If the photons have less energy than the “reddest” we can see, you get into “infrared light” and if the photons have more energy than the “violetest” you can see you’ve gone into ultraviolet (UV) territory.

Light can have different combinations of photons. White light contains all visible photon energies so it’s like a pitcher throwing all the balls – when molecules absorb certain colors of light we can’t see that color anymore – instead we see the sum of the remaining colors – so, depending on which color(s) get absorbed, objects will look different colors –  e.g. steal red, get green

I’ve been speaking in terms of energy, which is helpful for some things, like when looking at the electronic level, but you’ll more commonly see light discussed in terms of wavelength (λ) and frequency (f). Thankfully, you can easily convert between them because they’re all interrelated. 

All light travels at the same speed (the speed of light, c) – but imagine how bored Usain Bolt would be if he had to keep pace with a slow person just casually walking… So kinda like Bolt might run in a jig-zag to burn off some of that extra energy while keeping the same linear pace, higher energy photons will bounce up & down more (have a higher frequency) so their peaks will be closer together (shorter wavelength). Another way to think of it is as 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.

higher frequency light requires higher energy photons and corresponds to shorter wavelengths⠀

lower frequency light -> lower energy photons -> longer wavelengths⠀

Formally, we can write mathematize this with the equation E=hf=hc/λ, where h is Planck’s constant, 6.626*10-34 m² kg / s.

So, if we go back to our baseball analogy, the balls are all thrown at the same speed (the speed of light) so, when a group of balls is thrown by the pitcher, 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. (so the purple balls oscillate more than the red ones).

When a photon encounters something you can have a few different things happen, depending on the “something’s” catcher-fielding: transmission (photon goes through) (imagine a field with catchers really far apart & with the wrong mitts), reflection (photon bounces off)(imagine lots of catchers close together, but with the wrong mitts), or absorption (photon “caught”)(catchers with right mitts)

Most of the time, when a photon gets “caught” the energy of the photon is just gradually given off through electron wiggling, heat dissipation, etc. – but sometimes certain molecules can actually give a photon back (usually with a lower energy because the process isn’t totally efficient) and we call that FLUORESCENCE.

So, what do I mean by “right” or “wrong” “mitts”? Not all the catchers are equipped to capture all the balls. Different catchers have “special mitts” (chromophores) that only allow them to catch specific color balls. 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:

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”5 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.

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 “only” 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 the 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 (wavelengths of about 380-740 nanometers (nm)). 

note: a nm is a billionth of a meter or, somewhat easier to picture (though still pretty impossible…) a millionth of a millimeter. Sometimes, instead of nm, we use “angstrom” (Å) which is a unit of distance equal to 0.1 nm. Sorry for the segue, just don’t want you to get confused if you see this! Anyway, back to the story… 

As I was saying, lots of chemicals have too high of rent differences so the light they absorb is too high energy to be seen. But, if you look at chemicals that we *can* see you’ll often see a bunch of alternating single and double bonds (at least they’re drawn like this). Atoms link up to form molecules by sharing electrons with their next-door neighbors – pairs of electrons “moving in together” – one pair shared is a single bond & 2 pairs makes a double bond. 

These covalent bonds normally just involve next door neighbors. But in the case of resonance – aka electron delocalization – aka conjugation – which can happen when you have alternating single and double bonds – atoms can share electrons with others in the neighborhood. Basically the atoms do some “redistricting” so that they can all evenly share some of their electrons. And this often lowers the rent differences, so these molecules often absorb visible light. 

So back to our fall leave colors… To capture usable energy from sunlight, plants make a bunch of a chemical called chlorophyll. During the summer there’s long days with lots of sunlight, so it’s all about glucose-making – they want to capture as much of the sunlight as they can, use it to make glucose (a sugar), and store that glucose as starch in the “hardier” parts of plants like the trunk & roots. It’s kinda like the tree equivalent of squirrels hoarding acorns to prepare for winter

To take full advantage of the summer sun, plants produce lots of chlorophyll “catcher’s mitts” and spread them out over a large surface area so that mitt and light have a chance to meet! They grow big flat leaves with compartments called chloroplasts that contain lots and lots of a chlorophyll which can absorb specific photons that the plant can “cash in” for sugar through the process of photosynthesis. 

There are actually multiple types of chlorophyll, with chlorophyll a & B being the main photosynthetic pigments. They absorb blue & red wavelengths, but not green – so we see green. Chlorophyll a is the most abundant plant pigment & it absorbs at maxima of 430nm (blue) & 662nm (red). Chlorophyll b has a slightly shifted absorption spectrum (453 & 642 maxima) which helps expand the range of light that can be captured & used to make sugar. 

Chlorophyll’s photon-catching is done by a “tetrapyrrolic ring” that holds a metal that helps it with electron transferring. It also has another, not-visible-light-abasorbing part – a long hydrophobic (water-avoided) tail that lets it embed in membranes, but the tetrapyrrolic (4 pyrorole-like) ring is where the electron action’s at! 

This ring forms a “porphyrin” structure that is like a big electron orgy – 18 atoms participating in one of those conjugated systems where electrons are shared among them. This porphyrin is similar to the heme of of hemoglobin – the oxygen-carrying molecule in our blood. But even though these “mitts” may look similar at the chemical level, hemoglobin makes our blood look red while chlorophyll makes plants look green – clearly little differences in chemical makeup can cause big differences in color! If you compare heme & chlorophyll you can see that heme holds iron, while chlorophyll has magnesium, and there are also some other differences that affect electronic “rent costs” and thus the absorbed colors. If you want to learn more about heme & hemoglobin, check out but for now, let’s get back to the plants…

As the days start getting shorter and there’s less available sunlight, it gets to the point where, if the plants kept up stocking chlorophyll at the current levels, they’d spend more energy preparing to catch sunlight than they’d get back in usable energy captured. Unfavorable energy economics… So the plants stop making chlorophyll. What’s left gets degraded (and parts recycled), and the red and blue light stop getting stolen.

But instead of looking white, the leaves usually look orange, yellow, or red. This is because, even though the leaves had “only looked green” before, when there was sooo much chlorophyll, there were actually other chromophores present with “different mitts”- they were just hidden.

A while back, we looked at one of these – beta-carotene, which makes pumpkins look orange because it absorbs violet & blue-green light, with its highest absorption at 450nm. is just one of numerous “carotenoids” – plants have lots of them because, in addition to helping out chlorophyll with collecting usable energy, by absorbing high-energy light, they keep that light from damaging the plant. They also help protect the plant by sopping up extra “free radicals” (unpaired electrons). And their pretty colors attract seed-spreaders. Carotenoids also start degrading in the fall, but the chlorophyll greatly “out-degrade” them – so the green’s gone long before these are.

In addition to beta-carotenoids, other common carotenoids are of the “xanthophyll” subclass – with a major one being lutein which makes leaves (and egg yolks) look yellow. But it’s not just lutein doing the yellowing – another class of hidden yellow pigments is the flavonoids (flavonols & flavones), which usually have more phenyl rings & more oxygen than the other pigments.

Those yellow & orange chromophores were already there all along (just hidden by the chlorophyll) but red & purple colors are “made new” – they come from chemicals called anthocyanins – another type of flavonoid made just before leaves fall off. Just how red or purple leaves look depends on a few things including the acidity of the tree sap & soil, as well as how much anthocyanins are made – plants actually have to use some of their precious glucose to make them, so the more glucose they’ve stocked up, the more anthocyanins they can make & the prettier the leaves – so it might be worth it to spend some glucose in order to attract seed-spreaders. They might also play a similar high-energy-light absorbing-to-keep-from-plant-harming role as the carotenoids. 

Another reason to spend that glucose on color? The leaves physically fall off because the plant starts building a barrier between the leaf and the rest of the tree & this traps whatever glucose is already in the leaves there. So might as well use it for something, right?!

Since the red’s not alone – you have all those other pigments there too – you can get different shades & different color patterns throughout leaves.  Wondering about the brown? The brown colors are due to tannins which are waste products of metabolism, and predator-ward-offers, that get trapped in the leaves when that barrier forms. 

Not all trees lose their leaves – the ones that don’t either have really good conditions year-round so they don’t need to (this is the case with tropical plants) – or they have really bad conditions year-round so they’ve adapted to it (the case with “evergreens” including the conifers pine, spruce, & fir) Not only do those trees have to deal with low sunlight, they’re also facing colder temps & less water – storing enough glucose for winter’s not gonna cut it. So they’ve adapted with things like long waxy needles – the wax helps minimize water loss – and many make “antifreeze” molecules that prevent water from freezing inside of the cells. 

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