THE HISTORY OF ATOMS

All matter is made up of atoms. This is something we now take as a given and one of the
things you learn right back at the beginning of high school or secondary school chemistry
classes. Despite this, our ideas about what an atom is are surprisingly recent: as little as one
hundred years ago, scientists were still debating what exactly an atom looked like. This graphic
takes a look at the key models proposed for the atom, and how they changed over time.
Though our graphic starts in the 1800s, the idea of atoms was around long before. In fact,
we have to go all the way back to Ancient Greece to find its genesis. The word ‘atom’ actually
comes from Ancient Greek and roughly translates as ‘indivisible’. The Ancient Greek theory has
been credited to several different scholars but is most often attributed to Democritus (460–370
BC) and his mentor Leucippus. Though their ideas about atoms were rudimentary compared to
our concepts today, they outlined the idea that everything is made of atoms, invisible and
indivisible spheres of matter of infinite type and number.
These scholars imagined atoms as varying in shape depending on the type of atom. They
envisaged iron atoms as having hooks which locked them together, explaining why iron was a
solid at room temperature. Water atoms were smooth and slippery, explaining why water was a
liquid at room temperature and could be poured. Though we now know that this is not the case,
their ideas laid the foundations for future atomic models.
It was a long wait, however, before these foundations were built upon. It wasn’t until 1803
that the English chemist John Dalton started to develop a more scientific definition of the atom.
He drew on the ideas of the Ancient Greeks in describing atoms as small, hard spheres that are
indivisible, and that atoms of a given element are identical to each other. The latter point is one
that pretty much still holds true, with the notable exception being isotopes of different elements,
which differ in their number of neutrons. However, since the neutron wouldn’t be discovered until
1932, we can probably forgive Dalton for this oversight. He came up with theories about how
atoms combine to make compounds and also came up with the first set of chemical symbols for
the known elements.
Dalton’s outlining of atomic theory was a start, but it still didn’t really tell us much about
the nature of atoms themselves. What followed was another, shorter lull where our knowledge of
atoms didn’t progress all that much. There were some attempts to define what atoms might look
like, such as Lord Kelvin’s suggestion that they might have a vortex-like structure, but it wasn’t
until just after the turn of the 20th Century that progress on elucidating atomic structure really
started to pick up
The first breakthrough came in the late 1800s when English physicist Joseph John (JJ)
Thomson discovered that the atom wasn’t as indivisible as previously claimed. He carried out
experiments using cathode rays produced in a discharge tube and found that the rays were
attracted by positively charged metal plates but repelled by negatively charged ones. From this,
he deduced the rays must be negatively charged.
By measuring the charge on the particles in the rays, he was able to deduce that they
were two thousand times lighter than hydrogen, and by changing the metal the cathode was
made from he could tell that these particles were present in many types of atoms. He had
discovered the electron (though he referred to it as a ‘corpuscle’), and shown that atoms were
not indivisible, but had smaller constituent parts. This discovery would win him a Nobel Prize in
1906.
In 1904, he put forward his model of the atom based on his findings. Dubbed ‘The Plum
Pudding Model’ (though not by Thomson himself), it envisaged the atom as a sphere of positive
charge, with electrons dotted throughout like plums in a pudding. Scientists had started to peer
into the atom’s innards, but Thomson’s model would not hang around for long – and it was one of
his students who provided the evidence to consign it to history.
Ernest Rutherford was a physicist from New Zealand who studied at Cambridge University
under Thomson. It was his later work at the University of Manchester which would provide further
insights into the insides of an atom. This work came after he had already received a Nobel Prize
in 1908 for his investigations into the chemistry of radioactive substances.
Rutherford devised an experiment to probe atomic structure which involved firing
positively charged alpha particles at a thin sheet of gold foil. The alpha particles were so small
they could pass through the gold foil, and according to Thomson’s model which showed the
positive charge diffused over the entire atom, they should do so with little or no deflection. By
carrying out this experiment, he hoped to be able to confirm Thomson’s model, but he ended up
doing exactly the opposite.
During the experiment, most of the alpha particles did pass through the foil with little or
no deflection. However, a very small number of the particles were deflected from their original
paths at very large angles. This was completely unexpected; as Rutherford himself observed, “It
was almost as incredible as if you fired a 15-inch shell at a piece of tissue paper and it came
back and hit you”. The only possible explanation was that the positive charge was not spread
throughout the atom, but concentrated in a small, dense centre: the nucleus. Most of the rest of
the atom was simply empty space.
Rutherford’s discovery of the nucleus meant the atomic model needed a rethink. He
proposed a model where the electrons orbit the positively charged nucleus. While this was an
improvement on Thomson’s model, it didn’t explain what kept the electrons orbiting instead of
simply spiralling into the nucleus.
Enter Niels Bohr. Bohr was a Danish physicist who set about trying to solve the problems
with Rutherford’s model. He realised that classical physics could not properly explain what was
going on at the atomic level; instead, he invoked quantum theory to try and explain the
arrangement of electrons. His model postulated the existence of energy levels or shells of
electrons. Electrons could only be found in these specific energy levels; in other words, their
energy was quantised, and couldn’t take just any value. Electrons could move between these
energy levels (referred to by Bohr as ‘stationary states’), but had to do so by either absorbing or
emitting energy.
Bohr’s suggestion of stable energy levels addressed the problem of electrons spiralling
into the nucleus to an extent, but not entirely. The exact reasons are a little more complex than
we’re going to discuss here because we’re getting into the complex world of quantum
mechanics; and as Bohr himself said, “If quantum mechanics hasn’t profoundly shocked you, you
haven’t understood it yet”. In other words, it gets kind of weird.
Bohr’s model didn’t solve all the atomic model problems. It worked well for hydrogen
atoms, but couldn’t explain observations of heavier elements. It also violates the Heisenberg
Uncertainty Principle, one of the cornerstones of quantum mechanics, which states we can’t
know both the exact position and momentum of an electron. Still, this principle wasn’t postulated
until several years after Bohr proposed his model. Despite all this, Bohr’s is probably still the
model of the atom you’re most familiar with since it’s often the one first introduced during high
school or secondary school chemistry courses. It still has its uses too; it’s quite handy for
explaining chemical bonding and the reactivity of some groups of elements at a simple level
At any rate, the model still required refining. At this point, many scientists were
investigating and trying to develop the quantum model of the atom. Chief amongst these was
Austrian physicist Erwin Schrödinger, who you’ve probably heard of before (he’s the guy with the
cat and the box). In 1926 Schrödinger proposed that, rather than the electrons moving in fixed
orbits or shells, the electrons behave as waves. This seems a little weird, but you probably
already recall that light can behave as both a wave and a particle (what’s known as a wave-
particle duality), and it turns out electrons can too.
Schrödinger solved a series of mathematical equations to come up with a model for the
distributions of electrons in an atom. His model shows the nucleus surrounded by clouds of
electron density. These clouds are clouds of probability; though we don’t know exactly where the
electrons are, we know they’re likely to be found in given regions of space. These regions of
space are referred to as electron orbitals. It’s perhaps understandable why high school chemistry
lessons don’t lead in straight with this model, though it’s the accepted model today, because it
takes a little more time to get your head around!
Schrödinger’s wasn’t quite the last word on the atom. In 1932, the English physicist James
Chadwick (a student of Ernest Rutherford) discovered the existence of the neutron, completing
our picture of the subatomic particles that make up an atom. The story doesn’t end there either;
physicists have since discovered that the protons and neutrons that make up the nucleus are
themselves divisible into particles called quarks – but that’s beyond the scope of this post! At
any rate, the atom gives us a great example of how scientific models can change over time and
shows how new evidence can lead to new models.