History of the Atom Part 2:
The Modern View
Joseph John Thomson: 1856 – 1940
The atom, now soundly established, needed further investigation. Thomson developed what would come to be known as the “Plum Pudding model”. In this model the electrons, or negative charge, were spread throughout the atom like plums in a plum pudding. He proposed various structures for their arrangements, but none of these was able to convincingly explain all the observed properties of different elements.
Thomson used cathode ray tubes (vacuum tubes with two electrodes) to prove that negative charges were part of the structure of every element tested, and could not be separated from them. He also later used similar tubes to prove the existence of two types of Neon, and hence the concepts of the neutron and the isotope were born.
The discovery of negatively charged particles was a huge leap forward. It also proved the existence of positive charge, since some elements had been proved to be electrically neutral. However, it was up to one of Thomson’s students to take the theory to the next level of detail, the New Zealander Ernest Rutherford.
Ernest Rutherford: 1871 – 1937
Rutherford also had a string of outstanding discoveries to his name including discovering radioactive half lives, which he used to accurately age the Earth. He also involved in electrical technology and was the first person in history to transmute one element into another – he turned nitrogen into oxygen. He also heavily investigated the existence of neutrons.
His biggest achievement in terms of atomic theory was his now famous Gold Foil experiment. In this he fired a beam of highly positively charged particles called alpha particles (later to be known as helium nuclei) at a sheet of gold foil a few particles thick.
Rutherford was expecting results in line with Thomson’s model. Given that Thomson’s model said that atoms were a spread out positive charge with some negative particles buried in them, Rutherford expected to see only tiny amounts of deflection of the stream of positive particles as a result of the dilute positive charge.
In short, he expected to see this:
Certainly, this is what the Thomson model predicted. The thinly spread positive charge would have a small but consistent effect on the super high positive charge of the alpha particles. Therefore most of the beams should have their course changed, but only slightly.
But that was NOT what was observed. What he actually observed was this:
What did that difference mean? Clearly something with a strong positive charge was repelling some of the alpha particles, but most went through the foil unchanged. There was only one explanation.
The only thing that could cause the positive particles to be deflected this much was a dense, very positive charge. This led Rutherford to the conclusion that the positive charge was all condensed into one place. Also, since the majority of the beams passed through without any deflection at all, the rest of the atom had to be made up of empty space. The electrons, or negative charges, must be occupying an equally small space since they would also cause some minor deflections.
By now elements were known to absorb and give off very specific amounts of energy. These are known as the atomic absorption spectrum and emission spectrum respectively (see next section). Rutherford did not meet with any real success in explaining these with his model. That would be tackled by the next big name in the field, Neils Bohr.
Niels Bohr 1885 – 1962: The Modern Atom
Niels Bohr further developed the emerging modern atomic model by proposing that electrons occupied fixed orbits around the nucleus, called electron shells.
These shells were at set distances from the nucleus and these were the same for all elements. The shells become larger the further they away they are from the nucleus, similar to the layers of skin that make up an onion; the outermost layer is considerably larger than the ones closest to the core.
He reasoned that larger shells could hold more electrons and proposed that each shell could hold 2 n squared electrons where n is the shell number. So the first shell could hold 2x(1 squared)=2 electrons, the second shell 2x(2 squared) electrons and so on. This gives the following sequence of numbers of electrons in shells: 2, 8, 18, 32, 50 etc.
He also showed that electrons can move from a lower orbit (close to the nucleus) to a higher orbit (further from the nucleus) by absorbing energy when heated, accounting for already known absorption spectra. These electrons would then lose that energy when allowed to cool, giving the emission spectra for different elements.
This was a great leap forward, and is still the basic theory taught in middle high school years. His theory worked well for Hydrogen and reasonably well for Helium, but did not fit the emission spectra of larger elements.
This theory would be refined in the final step of this story by a scientist called Erwin Schrodinger, who developed the currently used model.