So Ernest Rutherford used alpha particles, the same particles he showed were really just helium atoms in disguise, to discover the nature of all atoms. Rutherford pictured the atom as a sort of miniature solar system, with the compact and massive nucleus in the middle and electrons whizzing about in orbits like tiny planets. But scientists soon realized that Rutherford’s model was unstable.
Rutherford’s student Niels Bohr found a solution, of sorts. Bohr decided that electrons weren’t like planets, which can orbit at any distance from the Sun. Instead, electrons could exist in only certain energy states, called orbitals, and moved from orbital to orbital in discreet jumps (quantum leaps), during which they would either absorb or radiate specific amounts of energy. Bohr’s model worked wonderfully well for hydrogen. But as soon as Bohr tried to apply the model to helium, the wheels fell off. With its two interacting electrons, helium proved a far tougher atom to crack.
But Bohr was on the right track. Today a theory called quantum electrodynamics (QED) explains with great precision the orbital of helium and the behavior of its pair of electrons. QED shows us that chemistry is just number. And helium’s key number, 2, makes it unique as the most standoffish of all the elements. Today we know why, and helium was the clue that showed the way.
Wolfgang Pauli wondered why chemistry worked at all. Why, Pauli wondered, didn’t all electrons fall into the lowest energy state of Bohr’s atom? After all, if you excite the electron in a hydrogen atom, it eventually falls back to the ground state, releasing light energy along the way. Why didn’t all electrons behave this way?
Pauli guessed that it must be due to the other electrons in the way. And yet helium still didn’t make sense, because in helium it seemed that two electrons, not just one, existed and could fall back into that ground state. If two, why not more?
To explain this strange behavior of the helium atom, Pauli conjured an effect called the “Pauli exclusion principle” (he probably didn’t call it that right away). Essentially, Pauli saw that two identical electrons couldn’t occupy the same state – that’s what kept all atoms that aren’t helium from behaving like helium. But, Pauli said, the two electrons in helium’s ground state aren’t really in the same state. There’s one crucial difference between them. That difference was spin.
Pauli said that as long as the spins of two electrons are opposite, there is an attractive force between them (yes, even though they’re both negatively charged). If, however, the spins are the same, there is a repulsive force between them. This force (attractive or repulsive) is called a coupling force, and it happens with all particles that, like electrons, have a kind of spin called half-integral spin. None of that is particularly important. What is important is this. The Pauli exclusion principle explains that there’s not room for just one electron in the ground state. There’s room for two.
Once helium has its two electrons, it is a happy camper. The electrons form an extremely symmetric cloud around the nucleus, blocking out (“masking”) virtually all of the positive charge deep within. There’s no magnetic field produced, either, and since both kinds of electron spin are present, the coupling forces are all repulsive. This makes helium extremely non-reactive, and is why, unlike most other elements, it is always found in its simplest form – a single atom, disconnected from its own kind and from all other elements.
The Pauli exclusion principle became a crucial tool in helping scientist explain how all the other atoms in the periodic table are constructed. And, as we’ll see later, its absence for another kind of matter leads to some of the strangest behaviors we’ve encountered yet. Once again, helium lit the way.