What’s in an atom?


The Ancients Greeks believed that matter was infinitely indivisible. No matter how many times you cut a sheet of copper in half, it would you would still get ever smaller sheets of copper. One Greek, however, took a different view. Democritus suggested that matter was made up fundamentally of discreet particles, beyond which, you could break down an object no further. These particles were indivisible, or in Greek, átomos.

As it turned out, Democritus was half right. Matter is indeed composed of fundamental chemical building blocks, which we know as atoms. The copper sheet only remains a copper sheet until you break it down into individual atoms of copper. Go any further and you no longer have copper. The part he got wrong, was to assume that these atoms were truly indivisible.

Figure 1- Rutherford model of the atom with a central nucleus surrounded by orbiting electrons.

In a long process of discovery starting in the late 19th century, physicists mapped the structure of the atom. It started in 1897 with the discovery of the electron by the British physicist J. J. Thomson. Thomson was working at Cambridge University, studying cathode ray tubes (the devices used in the old tube television sets and monitors before being largely replaced by LCD technology). By playing around with magnets, he determined that the mysterious “cathode rays” were in fact a stream of particles, identical no matter the material of the cathode, smaller than even the lightest atom, and possessing a negative electrical charge. This led scientists to conclude that these particles are contained within atoms themselves; they are “sub-atomic”. When it was then determined that these particles were what also carries an electrical current, Thomson’s original name for them, “corpuscle”, was replaced with the more familiar name, electron.

The next big leap came in 1908 when Hans Geiger and Ernest Marsden, under the direction of Ernest Rutherford, performed the famous gold foil experiment, where they measured the scattering of alpha particles (described on the next page) from a very thin sheet of gold foil. Much to everyone’s astonishment, the magnitude of the scattering was enormous, with alpha particles being deflected in all directions. Only having all the positive charge concentrated in a tiny volume would explain the results of the gold foil experiment.

So Rutherford created a new vision of the structure of the atom in 1911. At the centre was a tiny but concentrated positive charge. The electrons, attracted by their opposite electrical charge (like the attraction between the north and south poles of two magnets), would orbit this central charge, like planets orbiting the Sun. And much like solar system, the vast majority of the atom’s volume would be empty space.

The scientific community would later give a name to the positive charge, concentrated at the heart of the atom. It would be known as the nucleus.


Figure 2- Model of a nucleus, comprised of protons and neutrons, collectively called nucleons.

So the atom was now known to contain a number of electrons, each with a negative charge of universal size and a nucleus of positive charge. In order to fully balance out the negative charge of the electrons, the nucleus must have a charge of discreet packets (i.e. it could be 2 or 3 but not 2.5 since you cannot have half an electron). Therefore, Rutherford went further with his model. He speculated that the nucleus itself might be composed to smaller particles with a fixed positive charge.

In 1918, Rutherford discovered this particle when he performed yet more alpha particle bombardment of nitrogen gas. He observed nuclei of hydrogen being emitted from the nitrogen. Hydrogen is the lightest element and its nucleus has the smallest charge. The charges of all other nuclei are a whole number of the charge of the hydrogen nucleus. This led Rutherford to conclude that the hydrogen nucleus is a fundamental particle of a nucleus and he named it the proton.

This tied into work performed 5 year earlier by Henry Moseley at the University of Manchester. Moseley had used x-ray diffraction to show a direct link between the charge of a nucleus of an atom, and a chemical property of that atom called the atomic number, a unique identifier of the chemical element of an atom. Now, atomic number could be understood as specifying the number of protons in the nucleus. For example, the chemical element oxygen has an atomic number of 8 and so has 8 protons in its nucleus.

But one mystery still remained. No physicist could adequately explain the weight of a nucleus for it bore no connection to the charge. For example, hydrogen has an atomic number of 1, 1 proton in its nucleus and helium has an atomic number of 2, 2 protons in its nucleus. But the vast majority of helium weighs 4 times the vast majority of hydrogen. The led the Austrian theoretician, Wolfgang Pauli, in 1930 to postulate a third sub-atomic particle, which he named the neutron. This hypothetical particle would have no electrical charge, but would merely add mass the nucleus, approximately the same mass each as a proton. The particle was officially discovered two years later at Cambridge by James Chadwick.

This completed the basic picture of the nucleus, a tiny object at the centre of the atom, comprising most of its mass, composed of two types of particles of similar, though not identical masses, protons and neutrons. The effect of different proton numbers was clear. The number of protons directly determined the chemical properties of the atom and so it became the definition of the atomic number.

The effect of different neutron number was merely to change the mass of the nucleus. In 1913, J. J. Thomson discovered that indeed some atoms of the same element could have different masses. These variations were later understood to be due to nuclei having the same number of protons but different number of neutrons. That same year, Margaret Todd also recognised these variations and saw in this, atoms of different mass belonging to the same place on the Periodic Table of Elements. She suggested a term for these mass variants based on the Greek for “at the same place” and named them isotopes.


Atoms of the same element but with different masses are said to be isotopes of that element. They are specified by their mass number, the sum total of protons and neutrons (together known as nucleons) in the nucleus. By placing this number after the name of the element, any nucleus can be given a unique name. For example, carbon-12, oxygen-16, polonium-210, uranium-238. The first has 12 nucleons in its nucleus. Carbon has an atomic number of 6 so 6 of those nucleons are protons. The other 6 are therefore neutrons. Chemical symbols from the Periodic Table can also be used with mass number preceding it in superscript e.g. 12C, 16O, 210Po, 238U.

Hydrogen is a special case and due to its importance in nuclear physics, its known isotopes are given individual names. The vast majority of hydrogen in the universe is hydrogen-1, known, as you would expect, as hydrogen. A fraction of a percent has any extra neutron and this hydrogen-2 is given the name deuterium. Although technically, it should be denoted 2H like any other isotope, it is colloquially given the symbol, D. A similar situation exists for the third isotope, hydrogen-3, named tritium, which unofficially is given the symbol, T.

The difference in chemical behaviour of isotopes is usually negligible. The vast majority of times, the isotopes of an element are chemically indistinguishable from each other. Uranium-235 and uranium-238 will undergo chemical reactions in exactly the same way. However, the difference of just one neutron can make all the difference in nuclear reactions. For example, oxygen-16 is the stable isotope we all know and love, but oxygen-15, with one less neutron, is very radioactive.


Although the concept had been given weight before, the big moment came in 1905 when Albert Einstein wrote his papers on the Theory of Special Relativity. From his thought experiments, he had concluded that there was equivalence between mass and energy. The two were the same, but in a slightly different form. A particle gaining energy will automatically gain mass and a particle losing energy will automatically lose mass according to that famous equation, E=mc².

The German physicist Max Planck suggested that this implied something about atoms, which get bound together in molecules. It was known that energy must be put into a molecule to break it apart into to its component atoms and that similarly, energy is given off when atoms are bound together. This difference in energy was called the binding energy. Planck said that if mass-energy equivalence was true, then it would mean the bound molecule, say methane, would have less mass than its constituent atoms indiviudally, in this case one carbon atom and four hydrogen atoms individually.

Figure 3- Binding energy per nucleon. (Image courtesy of NASA GSFC)

The problem for Planck was that the binding energies of chemical reactions were miniscule. The change in mass would be too small to be measured. The energy released from the combustion of one gram of methane with four grams of oxygen was known to be 55,000 joules, enough to keep a 500MW power station going for a 20,000th of a second. This change in mass from this release of energy would be 0.0000000006g. So a total of 5g of methane and oxygen forms carbon dioxide and water, which would collectively have a mass of 4.9999999994g. This difference remains immeasurable.

However, with the discovery of the neutron, something significant was discovered as scientists measured its mass. The mass of both the proton and the neutron alone was measurably greater than their masses within the nucleus. Helium-4 was known to have an atomic weight of 4.0026. But with a proton having an atomic weight of 1.0073 and the neutron having an atomic weight of 1.0087, helium-4 would be expected to have a total weight of 4.032. In forming 1g of helium-4, the nucleons had lost 0.0073g of mass. Given out as energy, this translates to 660 billion joules, enough to keep a 500MW power station going for over 7 minutes!

One thing was clear. The energy of the nucleus was unlike anything ever before conceived.

Over the years that followed, a combination of theory and experiment and an understanding of the myriad of forces and geometric principles at work within the nucleus, allowed physicists to draw the graph of binding energy. The binding energy released by nucleons upon coming together to form a nucleus increases dramatically starting from deuterium up to iron-58. From that point onwards, the energy tails off gradually. This means that energy can be released from either the fusion of elements lighter than iron or the fission of elements heavier than iron. It is fission, which is the basis for what has become known as the nuclear power station.