Into the Atom

A Little History Of Science: Into the Atom

The chemists liked the atom. It was what entered into chemical reactions. It had definite positions within compounds. It had properties that were roughly defined by its place on the periodic table.

Each atom had its tendency to be either negative or positive in its relationships with other atoms, and to have the joining patterns called valence. Chemists also appreciated the difference between a single atom and the grouping of atoms into molecules (collections of atoms bound together). They realised that whereas most were happy to exist as single atoms, some atoms – hydrogen and oxygen, for instance – naturally existed in the molecular form (H2 or O2).

Atoms’ relative weights, with hydrogen always as 1, were also measured with increasing accuracy. Yet none of this gave chemists much clue about the finer structures of atoms. They found they could manipulate atoms in their laboratories, but could not say much about what these units of matter actually were.

For much of the nineteenth century, physicists were rather more interested in other things: how energy was conserved, how electricity and magnetism could be measured, the nature of heat, and why gases behaved as they did. The physicists’ theory of gases – called the kinetic theory – also involved thinking about atoms and molecules. But physicists, like chemists, agreed that although atomic theory was extremely useful in explaining what they saw and measured, the real nature of atoms was difficult to understand.

The first sign that atoms were not simply the smallest unit of matter came with the momentous discovery of one of its components, the electron. Experiments had already shown that atoms could possess electric charges, because electric currents in a solution attracted some atoms to the positive and others to the negative pole.

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Physicists were not so sure that an atom’s electrical properties played any part in chemical reactions. But they measured their electrical charges and found they came in definite units. These units had been named ‘electrons’ in 1894, just after J.J. Thomson (1856–1940) in Cambridge began to use the cathode tube in his experimental work.

The cathode tube is quite simple. It is amazing, really, that some- thing so simple could begin to tell us about the fundamental structure of the atom and the universe. This tube has had most of the air sucked out to create a partial vacuum, and electrodes have been inserted at each end. When an electric current is sent though the tube, all sorts of interesting things happen, including the production of rays (radiations). Radiations are streams of energy or particles, and those made in the cathode tube consisted mostly of fast-moving, charged particles. Thomson and his colleagues at the Cavendish Laboratories began to measure the electrical charge and the weight of some of these radiations. They tried to decide how these two measurements were related to each other. In 1897 Thomson proposed that these rays were streams of charged sub- atomic particles: bits of the atoms. He estimated that they weighed only a tiny fraction of the lightest atom, hydrogen. It took physicists several years to agree that Thomson had indeed found the electron – and that it was the unit of charge that he and others had been measuring for some time.

So, atoms have electrons. What else do they have? The answer came gradually, from the results of more experiments with the cathode tube. The vacuums within the tube became better, and stronger electric currents could be passed through. Among those who exploited these technical advances was Thomson’s one-time student, collaborator and eventual successor at the Cavendish Laboratories in Cambridge, the New Zealander Ernest Rutherford (1873–1937). In the late 1890s Rutherford and Thomson identified two different kinds of rays given off by uranium, an element that had acquired great importance for physicists.

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One of the uranium rays could be bent in a magnetic field; the other could not. Not knowing what they were, Rutherford called them simply ‘alpha’ and ‘beta’ rays – just ‘A’ and ‘B’ in Greek. The names stuck. Rutherford continued to experiment with both strange rays for decades. It turned out that not just uranium, but a whole group of elements, gave off (emitted) these rays. These elements created great excitement in the early years of the twentieth century, and they remain very important today. They are the ‘radioactive’ elements, and uranium, radium and thorium occur most commonly. When scientists began to investigate their special properties, they learned crucial things about atomic structure.

The alpha ‘ray’ was fundamental. (It is also called the alpha ‘particle’ – the distinctions sometimes blur in the very small and very fast world of atomic physics.) Rutherford and his colleagues aimed them at very thin sheets of metals, measuring what happened.

Normally, the particles passed through the metal sheets. But occasionally, they bounced straight back. Imagine Rutherford’s surprise when he considered what had happened. It was as if he had fired a heavy cannon-ball into a sheet of paper and discovered that it had bounced right back at him. What the experiment meant was that the alpha particle had encountered a very dense part of the atoms that were making up the metal sheet. This dense area was the nucleus of the atom. His experiments showed that atoms consist mostly of empty space, and that was why most of the alpha particles passed straight through. It was only when they hit the highly concentrated mass in this central nucleus that they bounced back.

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More work showed that the nucleus is positively charged. Physicists began to suspect that the nucleus’s positive charge is balanced by the electron’s negative charges, and that the electrons circle around in the largely empty space surrounding the nucleus.

Rutherford is now considered the founder of nuclear physics. In 1908 he won the Nobel Prize in Chemistry for his discoveries. These prizes were named after their Swedish founder. They became the highest accolade in science after their introduction in 1901, and winning one remains the goal of many ambitious scientists. Rutherford was good at finding outstanding students and colleagues, and several of them won Nobel Prizes too.

Niels Bohr (1885–1962) from Denmark was one of these. He took Rutherford’s idea that the atom’s mass is almost all squeezed into its small nucleus and applied an exciting new tool called ‘quantum’ physics to develop something called the ‘Bohr atom’ in 1913. This was a model visualising what was going on inside the atom, using the best information scientists had at the time. It imagined that an atom had a structure something like our solar system, with the sun/nucleus in the middle and the planets/electrons spinning around it in their orbits. In Bohr’s model, the weight of the positively charged nucleus gave the atom its atomic weight, and therefore its place in the periodic table.

The nucleus was made up of positively-charged protons. The heavier the atom, the more protons were in the nucleus. The number of protons and electrons had to match so that the atom as a whole is electrically neutral. The electrons swirled around the nucleus in different orbits and this was where the ‘quantum’ bit came in. One of the brilliant parts of the whole package of ideas that scientists called ‘quantum physics’ was the idea that things in nature come in definite, individual packets (‘quanta’). The things can be mass, energy or whatever you are interested in. In the Bohr model, the electrons orbit in different, individual quantum states. The electrons nearest the nucleus are more strongly attracted to it. Those furthest away are the less strongly bound and it is these electrons that are available to participate in chemical reactions or to generate electricity or magnetism.

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If this all seems rather difficult – well, it is. Bohr knew that. But he also knew that his Bohr model allowed physicists and chemists to talk in the same language. It was grounded on experiments by physicists, but also went far in explaining what the chemists observed in their own laboratories. In particular, it helped explain why elements in the periodic table behaved as they did, with their differing joining patterns, or valence. Those that joined singly did so because they had only one ‘free’ electron. Others had different patterns because of the number of ‘free’ electrons they had. His model of the atom has become one of the modern icons of science, even if we know now that the atom is much more complicated even than Bohr thought.

All sorts of new questions arose. First, how is it that the positively-charged protons can co-exist in the tight space that is the atomic nucleus? With electric charge, like repels like, and opposites attract (think of two magnets). So why don’t the protons push each other apart, and why don’t the electrons get sucked in? Second, the lightest known atom was hydrogen, so let’s assume that hydrogen, with its atomic weight of 1, consists of a single proton and an almost weightless electron. That means it’s reasonable to assume that the proton has an atomic weight of 1. So why don’t the atomic weights of the atoms in the periodic table simply go up in a nice steady flow: 1, 2, 3, 4, 5, and so on?

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An answer to the first puzzle had to wait until quantum mechanics was further developed. The second puzzle, about skips in the sequence of atomic weights, was solved much sooner, by another of Rutherford’s Cambridge colleagues, James Chadwick (1891–1974). In 1932, Chadwick announced the results of his bombarding experiments. Ever since Rutherford, this method had been a vital tool for physicists at work on the structure of the atom.

Chadwick had been sending streams of alpha particles at his favourite metal, beryllium. He found that beryllium sometimes emitted a particle with an atomic weight of one, and no charge. He used Rutherford’s name for the particle – the neutron – but it soon became clear that it was not simply a proton and electron combined, as Rutherford had thought, but a fundamental particle of nature.

The neutron was a kind of missing link for physicists, explaining puzzling atomic weights and positions in the periodic table. Mendeleev’s representation of the earth’s elements was continuing to prove its worth in charting the basic materials of our planet.

Chadwick’s neutron also led to the discovery of isotopes. Sometimes atoms of the same element have different atomic weights – if they have a different number of neutrons, which are these neutral particles in the atom’s nucleus. Isotopes are thus atoms of the same element with different atomic weights. Even hydrogen can occasionally have an atomic weight of 2 instead of 1, when it has a neutron along with its single proton. Chadwick won the Nobel Prize for his discovery of neutrons and what they could do, only three years after he discovered them.

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The neutron was a powerful tool for bombarding the nuclei of other atoms. Lacking either a positive or negative charge, it isn’t naturally repulsed by the heavily positive atomic nucleus, with its tightly packed protons. Chadwick recognised this, and saw that if you were going to smash atoms, you needed a machine that could accelerate them to high speeds and energies: a cyclotron or a synchrotron. These use very strong magnetic fields to propel atoms and their particles almost as fast as light. To do this kind of research, Chadwick left Cambridge for the University of Liverpool, because there he was given funds to build a cyclotron. There he saw that smashing high-speed neutrons into heavy atoms, such as uranium, could generate terrific energies. If such energies were harnessed, they could start a chain reaction leading to a momentous outcome: atomic ‘fission’, the splitting of the atom. The atomic bombs that were made and used to end the Second World War were the result of this work, and Chadwick was the leader of the British side of this project.

Many thought that Chadwick’s discovery of the neutron solved the problems of the structures of atoms (the building-blocks of the universe). But they were wrong. There were many surprises still to be discovered. Yet even the basic understanding of the electron, proton and neutron had involved physicists with several waves and particles, such as the alpha, beta and gamma rays. They had had to understand other mysterious phenomena, such as X-rays, and the discovery that nature trades in those small packets, called quanta. Nuclear physics and quantum physics: these were the areas of physics at the cutting edge of knowledge for much of the twentieth century.