Discoveries: The Bohr Model

After Rutherford's discovery of the nucleus, the next major advance in atomic theory came from Danish physicist Niels Bohr. To discuss Bohr's model of the atom properly, we will need to start with some background on atomic spectra.

Emission Spectra

When a material is heated to high enough temperature or when enough electric current flows through it, it will begin to emit light. If this light is spread out, the result is called an emission spectrum – it is the spectrum of the light emitted from the material. (Note: "spectra" is the plural form of "spectrum" - it's a Latin thing).

Two kinds of emission spectra are possible. One, called a continuous spectrum, has a continuous range of colors. This is the kind of spectrum the filament of a light bulb emits. The range of colors emitted depends on the temperature. At low temperatures, light at the red/yellow end of the spectrum is emitted. Candles burn at rather low temperature, hence the light from the flame appears yellowish. As the temperature is raised, more and more of the spectrum is emitted and the light becomes whiter. At very high temperatures, more light at the blue end of the spectrum is emitted, and the color takes on a bluish cast.

The other kind of spectrum is called a discrete spectrum, or line spectrum. A line spectrum contains on specific, ‘discrete’ colors of light that appear as lines. As you can see, a line spectrum looks very different from a continuous spectrum.

At the time Niels Bohr was working, scientists had become aware of many important facts about line spectra. They found that, for almost any element, you could get it to emit light by heating it up or subjecting it to high voltage. More interestingly, if you look at the specific arrangement of lines, each element had its own unique pattern. That is, different elements emitted different colors, had different numbers of lines in their spectra, etc.

Physicists had recently learned that the color of light is tied to its energy; so if elements were emitting different colors, it was because they were emitting different amounts of energy. But they didn't have a good explanation for why. Furthermore, they had trouble explaining why the emission spectra of gases were discrete at all - their knowledge of physics suggested that they should be continuous, like those of solids. There was good agreement that the electrons in atoms were responsible for the behavior ... but no good explanations for how it arose.

The Bohr Model

In 1913 Niels Bohr proposed an explanation. He proposed (as most physicists did at the time) that electrons orbited the nucleus of an atom the way planets orbit the sun in a solar system. However, Bohr went a step further: Bohr proposed that for electrons, only certain, specific orbital distances were possible.

The essence of Bohr’s idea was that an electron, like light, had a wavelength, and that the only orbits it could be in were those that had room for a whole number of wavelengths. For example, if the wavelength of an electron were 2 nm, it could only fit in orbits that were 2 nm, 4 nm, 6 nm, etc., in circumference, but not in an orbit with a circumference of, say, 3 nm.

Bohr further suggested that each “allowed” orbit had a different energy – the further from the nucleus, the higher the energy. When an atom absorbed energy, an electron moved from a lower energy orbit to a higher energy orbit. Since only certain orbits were allowed, only certain amounts of energy could be absorbed – amounts corresponding to the differences in energy between the allowed orbits.

When electrons dropped from higher to lower energy orbits, energy was emitted as light. Since only certain orbits were allowed, only certain amounts of energy could be emitted – amounts corresponding to the differences in energy between the allowed orbits.

Bohr’s model correctly predicted the spectrum of hydrogen. It also explained why different atoms showed different spectral lines: their different atomic numbers would alter the energies of the allowed orbits. The Bohr model failed to correctly predict the spectra of atoms with more than one electron. That would await the development of wave mechanics, which is the topic of our next section.