For my first real post on QUANTUM: Einstein, Bohr, And The Great Debate About The Nature Of Reality by Manjit Kuman, I’ve decided I can best use the time and space to try to clarify for myself the big-picture development of the idea of the “quantum” atom.
A little over a hundred years ago, about the time scientists were beginning to agree that, yeah, maybe atoms really do exist, experimenters discovered that they could cause atoms to release tiny electrically charged particles. They called these particles electrons, and described their charge as “negative.” But the question arose: why don’t these negative charges repel each other and rip atoms apart?
J.J. Thomson suggested that electrons are dispersed in, and held together by, a cloud of positive charge, like negative plums in a positive pudding. This theory not only explained how the atom stayed together, but why it could, itself, have no net charge.
Thomson’s model was almost immediately contradicted by reality… Rutherford discovered “alpha rays” that were actually particles of positive charge, to be called protons. In his model of the atom, he located at its center the protons, with the negatively charged electrons revolving around them.
This is where the real trouble began.
According to an old scientific theory no one wanted to jettison, a revolving charged particle like the electron should steadily lose energy as it revolves— and thus collapse into the center of the atom. That would be the end of the atom, and the end of the atom would mean the end of the matter. Needless to say, this threw some doubt on Rutherford’s “improved” atomic model.
Bohr postulated that electrons did not collapse inward because they were NOT steadily radiating energy. Bohr said that energy came in packets that could only be so small, and this smallest packet of energy, he called a “quantum.” If an electron did not have a quantum’s worth of energy to release, it could not radiate energy, and thus, could not begin its spiral downward through the lower energy levels.
Also, according to the latest thinking, there were only certain, very specifically allowed orbits within which an electron could revolve. An electron could not occupy a space inbetween these orbits even for a milli-second. Thus, to get from one orbit to another, an electron would have to “jump” all at once. The electron would be in one orbit and then—wham!—it would “magically” appear in another. This is the infamous “quantum leap.”
Why must electrons leap between orbits at all? When it was discovered that an atom bombarded with energy will often radiate energy back out, physicists had to come up with a theory to explain this. The consensus opinion that emerged was that the energy coming into the atom was being absorbed by electrons. This extra energy allowed the electrons to obtain the quanta of energy they needed to move between orbits. If, however, an electron moved up an orbit and that new orbit proved unstable, the electron would drop back down to a lower energy orbit, releasing its now-excess energy as a quantum packet. These quanta would then be picked up by the instruments of observers as energy the atom was radiating back out.
Another thing scientists worried about in their developing model of the atom was: what keeps the positively charged protons at the center of the atom from repelling each other? To keep the protons together, physicists invented the “strong” force—an attractive force between protons that is even stronger than the mysterious force we call “charge” that would otherwise repel protons from each other and break up the atom.
Also, when it was discovered that electrons had practically no mass and protons only accounted for half the weight of the atom, neutrons were added to the mix of subatomic particles. In this version of the atom, the protons and neutrons are in the center of the atom, making up what is called the nucleus, and the electrons revolve around the nucleus.
As quantum theory developed, physicists were able to satisfy themselves that they could describe very exactly the number and distances of the electron orbits circling an atom’s nucleus. Unfortunately, their nice neat atomic model was discovered not to contain enough electrons. So, to add more electrons without having to change the size of the orbits they had settled on, they decided that different electrons could revolve at roughly the same distance from the nucleus as long as their orbits were of different elliptical shapes.
When this still failed to provide enough electrons, theorists contended that the ellipses of the same orbit-distance could not only be of different shapes but could also be revolving at different angles— in other words, not all the orbits had to be on the same plane; some orbits might be vertical, some horizontal, and some inbetween. This idea of electrons being freed from a single plane of orbit is why we often speak now NOT of electron “orbits” but of electron “shells”—groups of orbits.
After all this, it was discovered there were still not enough spaces for electrons in the shells; the theoretical atoms had only half the number that they needed to match the experimental evidence. So, to double the number of electrons at one stroke, physicists came up with “spin.”
When they first came up with spin, theorists really did imagine the electrons as spinning on their little axes. It was maintained that some electrons could spin clockwise and some could spin counterclockwise. By allowing two electrons of opposite “spins” to occupy the exact same orbits, the number of electrons allowed in the shells was thus doubled. Rather quickly, physicists decided they didn’t like the idea of electrons actually spinning, so nowadays, “spin” just refers to some quality of the electron that makes two electrons capable of occupying the same orbit—though no one knows exactly what that certain quality is. However, the idea of visualizing electrons as truly spinning around their axes still come up sometimes—especially in theories of magnetism.