Inside the Atom


Atoms are so small that we cannot see them directly, even with the aid of a powerful microscope. The use of the word atom to describe the smallest particles of matter is thought to have originated with the Greek philosopher Democritus. It is derived from the Greek word atomos, which means that which cannot be split.

Subatomic particles

Today, we know that the atom can in fact be split into even smaller components, which we collectively call subatomic particles. All atoms - with the exception of the hydrogen atom * - consist of three types of subatomic particle - protons, neutrons, and electrons. Actually, several other kinds of subatomic particle are known to exist, but we don't need to worry about them at the moment.

The important characteristics of protons, neutrons and electrons are summarised in the table below. You can see from the table that most of an atom's mass is provided by protons and neutrons. These particles have almost the same mass as one another. The electron, by comparison, has a mass of roughly  1/1840  the mass of either a proton or a neutron.

Note that we have shown mass in both grams and atomic mass units (also known as relative atomic mass). One atomic mass unit (u) is defined as one-twelfth of the mass of an unbound neutral atom of carbon-12. We often use atomic mass units to specify the mass of atoms and their constituent particles because it gives us a meaningful number to work with, as opposed to simply expressing their weight in grams.

Location: NucleusNucleusOrbital
Charge (electron volts): +1e0-1e
Mass (grams): 1.6727 × 10 -24 g1.6750 × 10 -24 g9.110 × 10 -28 g
Mass (atomic mass units): 1.0073161.0087010.000549

* Although there are naturally occurring isotopes (see below) of hydrogen that have either one or two neutrons in their nucleus, the vast majority of hydrogen atoms do not possess a neutron. Most hydrogen atoms (99.98% in fact) consist of one proton and one electron.

Basic structure

The protons and neutrons in an atom (collectively known as nucleons) are all found in the core of the atom, in a structure we call the nucleus. The electrons, which are very much smaller than the protons and neutrons, are to be found some distance away from the nucleus, and are constantly in motion. The illustration below shows a helium atom. The nucleus of a helium atom consists of two protons and two neutrons. The fuzzy area around the nucleus represents the electron cloud, which in the case of the helium atom contains two electrons.

A graphical representation of a helium atom

A graphical representation of a helium atom

We sometimes talk about electrons "orbiting" the nucleus in the same way that the planets in the solar system orbit the Sun, but this is an over simplification. The path of an electron is pretty unpredictable, and the electrons move incredibly fast - close to the speed of light, in fact. In truth, therefore, it is virtually impossible to predict exactly where an electron will be at any given time, which is why an atom’s electrons are often represented diagrammatically as a "cloud".

The regions around the nucleus of an atom in which an electron is most likely to be found are called orbitals. The number of orbitals, together with their shape, size and orientation, will depend on the kind of atom we are dealing with. The overall shape of the atom (the nucleus and its electron orbitals) will also depend on the type of atom. It is generally believed to be symmetrical in nature, although the application of external forces such as electrical fields may cause deformations to occur.

The negatively charged electrons are bound to the nucleus because of the electromagnetic attraction between them and the positively charged protons. The closer the electron is to the nucleus, the greater that attraction will be, so you may well ask the question: why does an electron not simply fall into the nucleus? In essence, the loss of potential energy suffered by the electron as it approaches the nucleus is balanced by a gain in kinetic energy (momentum and velocity); the total amount of energy within the atom remains constant.

The electrons furthest away from the nucleus experience only a relatively weak attraction to the nucleus. Under the right conditions (the application of an external force, for example), they may be able to break free of their parent atom, leaving the atom with an overall positive electrical charge. It is also the possible for an atom to gain one or more additional electrons, in which case the atom will acquire an overall negative charge. Atoms that become positively or negatively charged in this way are called ions.

How big is an atom?

Defining the size of an atom is not a straightforward task, because most of the atom is empty space - more than ninety-nine percent of its mass is concentrated in the nucleus. As a rough guide, if we compare the overall size of an atom to that of a football stadium, the nucleus would be about the size of a pea. The electrons would be no bigger than specks of dust, and would be whizzing around all over the stadium at incredible speeds.

The size of an atom will depend on what type of atom we are dealing with, but as a very rough approximation we could put the diameter of a typical atom at about one ten-billionth of a metre (10 -10 m). To try and put this into context, a sugar lump contains roughly as many atoms as there are stars in the universe. For comparison, the diameter of a proton or neutron (both approximately equal in size) has been estimated to be in the order of 2.5 × 10 -15 m.

Elements and their isotopes

The number of protons in a (non-ionised) atom is always matched by the number of electrons orbiting the atom, and will determine to which element the atom belongs. We refer to this number as the atomic number. Hydrogen has an atomic number of one (1), because hydrogen atoms have one proton and one electron. Carbon has an atomic number of six (6), because carbon atoms have six protons and six electrons.

In some of the lighter elements (e.g. helium, carbon, nitrogen and oxygen), the number of neutrons in an atom is the same as the number of protons - at least for the most commonly occurring isotopes of those elements. In heavier elements, we find significantly greater numbers of neutrons than protons. The only stable isotope of gold, for example, has seventy-nine (79) protons and one hundred and eighteen (118) neutrons.

If we add the number of protons and neutrons in the nucleus together, we get the element's mass number. The mass number is important, because it allows us to distinguish between different isotopes of the same element. The most abundant isotope of carbon, for example, has six protons and six neutrons for a total of twelve nucleons. You will often see this isotope referred to as either carbon-12 or 12 C. Another isotope of carbon - the one that scientists use for carbon dating - has eight neutrons in its nucleus, and is designated as carbon-14 (or 14 C).

You may be wondering why there are more neutrons than protons in these heavier elements. Without going into too much detail, it has to do with the forces at work inside the atom. Positively charged protons generate electrostatic forces that cause them to repel one another, but there is also an attractive nuclear force at work (sometimes called the residual strong force) between all of the nucleons that overcomes these electrostatic forces and keeps them together.

As long as there are only a small number of protons, the attractive nuclear forces are stronger than the repulsive electrostatic forces. The range of the nuclear force is limited, however. As the number of protons grows, the repulsive electrostatic forces threaten to overcome the attractive nuclear force. Because neutrons carry no charge, they do not repel other nucleons, but they do contribute to the nuclear force. Additional neutrons therefore compensate for the growing imbalance between the electrostatic and nuclear forces.

When the number of protons grows beyond a certain point, even the addition of more neutrons will not be enough to counter the electrostatic forces trying to push the protons apart. The magic number seems to be eighty-two (82) - there are no stable isotopes for elements that have more than eighty-two protons in their nucleus. Lead (Pb), with an atomic number of eighty-two, is the heaviest element that can exist as a stable isotope.

Radioactive decay

Every element beyond lead in the periodic table is radioactive. This means that they are not stable and will eventually decay, becoming different elements in the process. Carbon-14 eventaully decays to become nitrogen-14 through a process of beta minus decay, in which one of the neutrons is transformed into a proton and an electron (and an infinitesimally small particle called an antineutrino, which we will ignore for the purposes of this discussion). The electron is ejected from the nucleus at high speed.

The rate at which unstable isotopes decay varies from one isotope to the next. Because radioiactive decay is a negative exponential function (the rate of decay slows down over time), we tend to talk about the half-life of a radioisotope - the time required for half of the material to decay. In naturally occurring radioisotopes found on Earth, the half-life can be anything from a few minutes to many billions of years!