What makes color so colorful?
July 12, 2004 4:08 PM Subscribe
Physics of Color: I understand that color arises from the uneven absorption of the different wavelengths of electromagnetic radiation, and is decoded in the eye by the rods and cones &c., but what is it about the physical structure of materials that makes them have distinct absorption profiles in the first place? In other words, why is grass green?
Response by poster: Great answer, thanks!
One more question, if the "allowed" absorption frequencies are quantized, why are absorption curves smooth? Shouldn't they be discontinous?
posted by signal at 4:37 PM on July 12, 2004
One more question, if the "allowed" absorption frequencies are quantized, why are absorption curves smooth? Shouldn't they be discontinous?
posted by signal at 4:37 PM on July 12, 2004
P.S.: Electron excitation is not the only way for molecules to absorb electromagnetic energy, but it is the primary way energy is absorbed from the visible portion spectrum. For example, a transition from one vibrational mode of a molecule to another--that is, the atoms themselves vibrating in different ways--also absorbs electromagnetic energy, but that's generally in the infrared portion of the spectrum, so it doesn't affect the color that we perceive.
posted by DevilsAdvocate at 4:40 PM on July 12, 2004
posted by DevilsAdvocate at 4:40 PM on July 12, 2004
One more question, if the "allowed" absorption frequencies are quantized, why are absorption curves smooth? Shouldn't they be discontinous?
They are discontinuous, under idealized conditions--e.g., light passing through a gas of individual atoms. Under everyday conditions, the issue is that molecules are surrounded by lots of other molecules, and that alters the allowable energy levels. This chlorophyll molecule, surrounded by such-and-such other molecules in certain ways, absorbs slightly different frequencies of light than that chlorophyll molecule, with a different set of molecules around it. When you're talking millions of millions of molecules, each with a slightly different environment, it adds up to a smooth curve. I think.
posted by DevilsAdvocate at 4:48 PM on July 12, 2004
They are discontinuous, under idealized conditions--e.g., light passing through a gas of individual atoms. Under everyday conditions, the issue is that molecules are surrounded by lots of other molecules, and that alters the allowable energy levels. This chlorophyll molecule, surrounded by such-and-such other molecules in certain ways, absorbs slightly different frequencies of light than that chlorophyll molecule, with a different set of molecules around it. When you're talking millions of millions of molecules, each with a slightly different environment, it adds up to a smooth curve. I think.
posted by DevilsAdvocate at 4:48 PM on July 12, 2004
They are discontinuous, under idealized conditions--e.g., light passing through a gas of individual atoms.
Nitpick, which actually only reinforces your point: Even a sample of gas will exhibit a "continuous spectrum", since the motion of the molecules themselves will cause some molecules to absorb photons of a higher frequency, some at a lower frequency. (This phenomenon is called Doppler broadening.) Moral of the story: there's always something going on to muck up the works. The art of being a physicist is balancing your desire for precision with the desire to actually be able to predict stuff...
posted by Johnny Assay at 5:55 PM on July 12, 2004
Nitpick, which actually only reinforces your point: Even a sample of gas will exhibit a "continuous spectrum", since the motion of the molecules themselves will cause some molecules to absorb photons of a higher frequency, some at a lower frequency. (This phenomenon is called Doppler broadening.) Moral of the story: there's always something going on to muck up the works. The art of being a physicist is balancing your desire for precision with the desire to actually be able to predict stuff...
posted by Johnny Assay at 5:55 PM on July 12, 2004
another process for broad absorption is that a photon can kick an electron from within an atom to flying off into free space. since the energy of free-flying atoms are effectively not quantised the "Y" can vary smoothly, so the photon energy isn't restricted to a fixed value. however, i don't think this occurs much at optical wavelengths.
posted by andrew cooke at 6:23 AM on July 13, 2004
posted by andrew cooke at 6:23 AM on July 13, 2004
Best answer: I take umbarage with the question. Physicists don't now a damn thing about colour---this one is pure quantum chemistry. Here, let me show you my union card.
Devil's Advocate is correct in a limited way, that pehnomoenon called pressure-broadening, but what's happening in the visible spectrum is more complex than that.
We need a bit of background first. There are essentially three things a molecule can do to absorb energy (and still stay a neutral molecule). In order of lowest to highest energies they are: spinning (usually called rotation---microwave energy), bond-streching (vibration---in the IR) and electron motion (or excitation---in the visible). There's about an order of magnitue or two difference between each one of these types of energy absorption.
Particularly important for this question are the vibrational and electronic states (or energy levels). Electronic states are generally widely spaced for smallish molecules--the ones that are most brightly coloured. However at each electronic state, there are a (very large) number of vibrational states the molecule can be in.
So, when a photon of the right colour comes along to kick an electron from one state to another, it can come from a (very large) number of vibrational states in the ground state to a (very large) number of vibrational states in the excited state. There are restrictions (selection rules) for which state-to-state jump can be made, but that's not important for us right now. If you'll allow me an equation, we can write the energy of the transition like this:
Ground State (Eelectron + Evib) -> Excited State (E'electron + E'vib)
There are only one of Eelectron and E'electron, but there are lots (hundreds, if not thousands) of Evib and E'vib's. Remember Eelectron >> Evib. This means that the transition occurs from large multidue of ground states to a larger multitue of excited states. Each one of these transitions has a slightly different energy and thus is triggered by a slightly different colour photon. Thus you get broad visible absorption spectra. This mechanism is called vibronic coupling.
Bonus answer: what happens to the energy that is abosorbed? Most of the time molecules can dump it in a series of short jumps down, through overlapping vibrational levels, rather than in a single jump. This we know as heat. Of course, it can do other things too. Chlorophyl uses abosrbed light to power a chemical reaction that results in simple sugars. If there's a state available for a long jump down, somethimes this is favoured---this is fluorescence. Sometimes the energy has to bleed down through a series of low-probability transitions before it finds a ground state it can jump down two---this is phosphorescence.
posted by bonehead at 7:23 AM on July 13, 2004 [1 favorite]
Devil's Advocate is correct in a limited way, that pehnomoenon called pressure-broadening, but what's happening in the visible spectrum is more complex than that.
We need a bit of background first. There are essentially three things a molecule can do to absorb energy (and still stay a neutral molecule). In order of lowest to highest energies they are: spinning (usually called rotation---microwave energy), bond-streching (vibration---in the IR) and electron motion (or excitation---in the visible). There's about an order of magnitue or two difference between each one of these types of energy absorption.
Particularly important for this question are the vibrational and electronic states (or energy levels). Electronic states are generally widely spaced for smallish molecules--the ones that are most brightly coloured. However at each electronic state, there are a (very large) number of vibrational states the molecule can be in.
So, when a photon of the right colour comes along to kick an electron from one state to another, it can come from a (very large) number of vibrational states in the ground state to a (very large) number of vibrational states in the excited state. There are restrictions (selection rules) for which state-to-state jump can be made, but that's not important for us right now. If you'll allow me an equation, we can write the energy of the transition like this:
Ground State (Eelectron + Evib) -> Excited State (E'electron + E'vib)
There are only one of Eelectron and E'electron, but there are lots (hundreds, if not thousands) of Evib and E'vib's. Remember Eelectron >> Evib. This means that the transition occurs from large multidue of ground states to a larger multitue of excited states. Each one of these transitions has a slightly different energy and thus is triggered by a slightly different colour photon. Thus you get broad visible absorption spectra. This mechanism is called vibronic coupling.
Bonus answer: what happens to the energy that is abosorbed? Most of the time molecules can dump it in a series of short jumps down, through overlapping vibrational levels, rather than in a single jump. This we know as heat. Of course, it can do other things too. Chlorophyl uses abosrbed light to power a chemical reaction that results in simple sugars. If there's a state available for a long jump down, somethimes this is favoured---this is fluorescence. Sometimes the energy has to bleed down through a series of low-probability transitions before it finds a ground state it can jump down two---this is phosphorescence.
posted by bonehead at 7:23 AM on July 13, 2004 [1 favorite]
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So, in order for an electron to be excited--to move from having X energy to having Y energy--it has to interact with a photon having (Y-X) energy. More or less energy than that, and the electron would have a total energy that isn't allowed, so it can only absorb energy from photons having certain specific energies.
The energy of a photon is directly proportional to its frequency, and inversely proportional to its wavelength. Light that we perceive as "red" carries a certain amount of energy per photon; what we perceive as "green" carries a different amount per photon.
Chlorophyll has allowable levels of electron energy such that the difference between one level (occupied when all the electrons are in their lowest energy states) and another (unoccupied when all electrons are in their lowest energy states) corresponds to the energy of photons which we perceive as red. Red photons kick an electron in chlorophyll from one allowed energy state to another, so chlorophyll absorbs red light. It doesn't absorb green light, because to do so would give an electron an amount of energy that isn't allowed for chlorophyll. When red light is removed from a white light source, what remains is perceived by us as green, so grass is green.
I know that's not a complete answer, because I haven't gotten into why the allowable electron levels (and the difference between levels) are different in different molecules, but I can't really explain that without going into pretty heavy quantum mechanics--and amittedly, it's been so long since I did that anyway I'm not sure I still could even if I wanted to.
posted by DevilsAdvocate at 4:30 PM on July 12, 2004 [1 favorite]