To some extent, the properties of electrons can be compared with the properties of a closely related phenomenon, radiant energy or light. Radiant energy, also known as electromagnetic radiation, travels through a vacuum in waves at a constant speed of 3.0 X 10⁸ m/sec.

In many ways the waves of electromagnetic radiation are like waves in water. You have probably seen closely spaced, choppy waves on a small lake. You may also have seen, on larger bodies of water like the ocean, waves that are farther apart. The difference between these two types of waves is in their wavelength (lambda), which is the distance from crest to crest. Three waves of different wavelength are shown in Figure 5.1.

FIGURE 5.1 Different waves have different wavelengths (distance from crest to crest).

A wave can also be characterized by its frequency - that is, the number of wave crests that pass a given point in a unit time. The frequency (nu) of a wave is related to its wavelength by the equation

C =

where c is the constant speed of light, 3.0 X 10⁸ m/sec. From this equation you can see that as wavelength increases, the frequency of the wave decreases. The energy associated with a wave is directly proportional to its frequency. Hence, the higher the frequency, the shorter the wavelength and the higher the energy of the wave.

Figure 5.2 shows the wide range of electromagnetic radiation from AM radio waves with a wavelength of 10⁴ m to gamma waves with a wavelength of 10 ^-12 m. This range is called the electromagnetic spectrum. The common names of the other kinds of electromagnetic radiation and their wavelengths are also given. Notice that visible light, electromagnetic radiation with wavelengths between 4 X 10 ^-7 and 7 X 10 ^-7 m, comprises only one small part of the electromagnetic spectrum. In Figure 5.2 you can see that red light has a longer wavelength than blue light. Red light, then, has a lower frequency and is associated with less energy than blue light. Infrared light, microwaves, television waves, and radio waves are invisible forms of electromagnetic radiation; their wavelengths are greater than that of visible light, and thus their energies are lower than that of visible light.

FIGURE 5.2 The electromagnetic spectrum. The visible range has been expanded to show the individual colors.

Ultraviolet light, X rays, and gamma rays - all of which have wavelengths shorter and energies higher than those of visible light (see Figure 5.2) - are also invisible forms of radiant energy.

When an object is heated, it radiates energy, often in the form of visible light. Our sun is probably the most familiar example of a heated body giving off light. The "white" light from the sun is a collection of light of all wavelengths and is called continuous light. When light passes through a prism, it is separated into its various wavelengths. You have probably seen how sunlight, when passed through a prism, separates into all the colors and wavelengths of the rainbow. However, if a gaseous sample of a single element is heated, the light emitted is not continuous and is only of a few wavelengths. When this light is passed through a prism, instead of a rainbow we see a series of brightly colored lines, each line corresponding to a particular wavelength of the emitted light. The pattern of wavelengths (or lines) of light thus produced is unique for each element and is called its emission spectrum (plural, spectra). Because the pattern is characteristic, it can be used to show the presence of that element in even the tiniest amounts. Figure 5.3 shows the emission spectra of hydrogen, neon, and sodium in the visible range. The light from sodium-vapor street lamps is that of the yellow-orange lines of the sodium spectrum. Neon signs use the red lines of the neon spectrum to produce their color. The spectra of these elements, as well as those of all the other elements, show other lines in the invisible parts of the electromagnetic spectrum.

FIGURE 5.3 Emission spectra of hydrogen, neon, and sodium in the visible range.

The spectra shown in Figure 5.3 are called emission spectra because they show the light (energy) given off (emitted) by an unusually energetic atom. Atoms can also absorb energy. If continuous light is passed through the vapor of an element, some wavelengths of light are absorbed. Analysis of the emerging light shows a rainbow interspersed with some black lines. This spectrum, called an absorption spectrum, shows that some energy (measured by the wavelength at which the black lines appear) has been absorbed by the vaporized element. If the vaporized element is hydrogen, the black lines will appear at exactly the same places in the absorption spectrum as did the bright lines in the emission spectrum of hydrogen shown in Figure 5.3. Similarly for sodium, neon, and other elements, the bright lines in their emission spectra are of the same wavelengths as the black lines in their absorption spectra.

We conclude from these observations that, first, atoms can lose or gain energy in amounts similar to the energy of light and, second, the energy of an electron can change only by certain increments and not by random amounts.