Science with Richard Bleil
Somebody once told Richard Feynman that they pitied him for being unable to appreciate art. Richard Feynman, if you do not know, is one of the preeminent physicists of recent times. This individual explained that the reason for the pity is that, surely, he couldn’t appreciate the skill and art as he only sees the electrons popping around making the colors. He smiled, and explained that, yes, he does appreciate the beauty, the skill, the brush strokes and the time it took to create the art, but because he understands how electrons pop, he has a deeper appreciation for the art.
Electrons popping around is predicted by quantum theory, something that fascinates so many people and yet is truly understood by so few. But even before quantum mechanics, popping electrons was predicted by Niels Bohr. “Electron popping” refers to electrons jumping between available energy levels.
Atoms and molecules have a myriad of energy levels available to them, perhaps hundreds of thousands or more, but they typically only use maybe half a dozen or so. What makes these energy levels so odd in our perspective is that they are discrete. They might have an energy level of thirty, or forty, but nothing in between. We say those energy levels, such as thirty-five, are forbidden. The electrons are free to use any of these energy available energy levels, but when we discuss electronic configuration, we typically are referring to the absolute lowest possible energy, called the “ground state”.
To get to these higher energy levels, electrons required the input of energy, such as light. Certain wavelengths of electromagnetic energy correspond exactly to the energy required to transition from one of these energy levels, say the ground state, to a higher one, say the fifth energy level. When this happens, the light is absorbed, resulting in the electron “popping” from the ground state to the fifth. When this wavelength of light happens to occur in the visible light spectrum, that color, say, for example, indigo, is trapped. All other colors are reflected back. We call these molecules “pigments”, and the color we see is basically a combination of all visible light wavelengths that are not absorbed. So by absorbing indigo, we would see a color more in the reds and yellows, the colors reflected back.
Now, if light is absorbed as the electron jumps from the ground to a higher energy level, then it makes sense that light is emitted as the electron “relaxes” from a higher energy level back to a lower energy level. This is called “fluorescence”, but the light fluoresced does not have to be the same color as the light absorbed. As it turns out, as the electron relaxes, it does not have to return immediately to the ground state. So if electrons popping from the ground state to the fifth state corresponds to the indigo wavelength, then the electron might relax from, say, the fifth energy state to the third. Because this energy transition is smaller, the wavelength of light released will correspond to a lower energy color such as yellow. Relaxing from the third to the ground state, then, would likely correspond to a wavelength in the infrared region, making it invisible to the naked eye. But because of the transition from the fifth to the third state, the compound might absorb indigo light, but release yellow light.
Fluorescence has been known for a very long time. You see it in those shirts that seem to glow in the back of stores with blacklights on. But electrons can never emit light of higher energy than it absorbs. It’s worth keeping in mind the name ROY G. BIV as it gives, from low to high energy, the colors of the visible light spectrum (red, orange, yellow, green, blue, indigo and violet). This is why the higher energy blue light tends to make things fluoresce at the lower energy colors green or yellow.
These energy transitions are so rapid that in our human time scale they seem instantaneous. Knowing that light absorbs in the visible light spectrum, it’s really not a surprise that it can also emit in these frequencies. In fact, measuring the frequencies of light absorbed (spectroscopy) or emitted (fluorescence) is a key technique in identifying chemicals. What did come as a shock, however, were the earliest compounds where the relaxation of the light was delayed. Certain stabilizing factors in molecules that can do this actually holds the electrons in the higher energy states for seconds, or even minutes. This is incredible.
When you play with glow-in-the-dark toys, that’s what’s happening. Electrons get “trapped” in higher energy levels for a noticeable amount of time. The longer these items are exposed to light, the more electrons get into these higher energy levels, and the more intense the light will be when it’s finally in the dark. When these trapped electrons relax to lower energy levels (which occurs in accordance with a Gaussian distribution which is why it glows intensely for a short time and seems to drag out dimly much longer) they release light, and your toy glows.