ABSORPTION AND COLOR

We have seen that atoms at low-energy states can absorb energy and that this energy can be light. If an atom at a low-energy state absorbs a photon with a particular amount of energy, that atom will reach a high energy state with a final energy equal to its original energy plus the energy of the photon absorbed. This is why substances appear in certain colors. For example, a red liquid appears red when you look through it because red photons are not absorbed by the liquid as are other colors of light (such as blue, green, and yellow). The same is true of a blue liquid, in which blue photons are not absorbed but greens, yellows, and reds are. So why does the blue liquid absorb the lower-energy red photon yet allow the higher energy blue photon through? The answer lies in the nature of absorption itself, in which the energy of an absorbed photon must match the energy of a transition: If energy levels in a particular atom exist 2.1 eV apart, a 2.1-eV photon is required to excite the atom from the lower level to the upper level. A photon with 1.7 eV of energy falls short of the energy needed to make the jump, whereas a 2.3-eV photon has too much energy. Another way of looking at absorption is as the opposite of emission. This is expected in nature: If atoms can emit light to lose energy, they must be able to absorb it to gain energy. Because of the quantized nature of energy levels, it is possible to find atomic and molecular species that have energy levels allowing the absorption of photons of essentially any energy. In the case of an atom such as hydrogen, energy levels are specific and sharply defined; however, molecules have broad energy bands that allow absorption (or emission) over a wide spectrum of wavelengths. This is why liquids absorb a range of wavelengths (such as all red and orange light) instead of a specific wavelength such as a low pressure gas would.