Taking the Quantum Leap

The common view of the atom is that of a compact nucleus with electrons swirling around it like little planets. This view is similar to the model proposed by Ernest Rutherford in the early 1900s when his famous experiment demonstrated that most of the mass of an atom is densely concentrated at its center. This is similar to the way most of the mass of the solar system is concentrated in the Sun, so Rutherford wondered if an atom had a similar structure.

But this “mini solar system” idea leads to a serious problem. Electrons have a negative charge, and moving charges create electromagnetic waves. If an electron orbited the nucleus like a little planet, it would radiate electromagnetic energy, and that loss of energy would cause it to spiral inward until it collided with the nucleus. If electrons orbited, in a fraction of a second the electron’s orbit collapse in a flash of light.

The first reasonable solution to this problem was proposed by Niels Bohr, who proposed that the orbits of an electron were quantized. This would mean that rather than continuously radiating energy, electrons would only emit energy when they jumped from a larger orbit to a smaller orbit. Likewise electrons could absorb energy to jump from a smaller to a larger orbit. In other words, the orbits of electrons were quantized, and electrons must make a quantum leap to move from one orbit to another.

We now know that electrons are not particles in the way we traditionally view them, but rather quanta with both wave and particle properties. Instead of swirling around the nucleus, they exist in regions around the nucleus called orbitals. Although Bohr’s model is simplistic, his idea of electrons existing in quantized energy levels still holds. This means that an individual atom or molecule cannot radiate light at any possible wavelength, but only at particular wavelengths depending on the type of element or molecule. They can also only absorb light at the same particular wavelengths.

This turns out to be an extremely useful tool for astronomy. Objects such as stars are hot enough and dense enough that the light they give off is almost a continuous spectrum. (Things like hot metal and incandescent lights also give off continuous spectra.) But the upper layers of a star are in a gaseous or plasma state, so their energy levels are quantized. This means when light from the interior of a star passes through the cooler layer of the star’s atmosphere, certain wavelengths of light are absorbed. This creates what is known as a dark line or absorption spectra, as you can see in the figure above.

The specific wavelengths absorbed depends on the type of atoms or molecules, so we can look at the absorption lines of a star and see what type of elements exist in its atmosphere. From that we can determine things like its age and size.

We can do the same thing with nebula and other diffuse clouds of gas in the universe. If they are heated by some source of energy, then they emit light at specific wavelengths. This creates what is known as a bright line or emission spectra. Interstellar clouds can also absorb light passing through them, so absorption spectra are also possible.

Since the pattern of line spectra is unique to particular elements, we can also see when the spectra are shifted in wavelength toward the red or blue. This redshift or blueshift effect is due to the motion of the star or cloud relative to us. So we can use line spectra to determine the motion of a star or nebula. Line spectra are also affected by temperature, ionization, magnetic fields, pressure, and lots of other things, which lets us study those properties purely by the light we observe.

There’s a great deal we have learned from the universe, once we learned to take the quantum leap.