Structure of the Atom

1. Charged particles in matter

For a long time people thought the atom was the smallest, indivisible particle of matter (the Greek word atomos means “uncuttable”). But experiments with electricity showed that atoms contain even smaller sub-atomic particles. When two objects are rubbed together they become charged, which means charge can be moved around — so atoms must contain charged parts. Three sub-atomic particles were discovered: the electron (negative charge), the proton (positive charge) and the neutron (no charge). An atom as a whole is electrically neutral because it has an equal number of protons and electrons.

2. Discovery of the electron and proton

J. J. Thomson discovered the electron while studying the flow of electricity through gases in a discharge tube (cathode rays). He showed these rays were streams of negatively charged particles present in all atoms. E. Goldstein later discovered new positively charged rays called canal rays, which led to the discovery of the proton. The electron carries a charge of −1 and the proton a charge of +1; the proton is about 2000 times heavier than the electron. Because the two charges are equal and opposite, a neutral atom has equal numbers of each.

3. Thomson’s model of the atom

Based on his discovery, J. J. Thomson (1898) proposed that an atom is a sphere of positive charge with electrons embedded inside it — rather like the seeds inside a watermelon, or the plums in a pudding (the “plum-pudding model”). The positive and negative charges were equal in amount, so the atom was neutral. This model explained neutrality but could not explain later experimental results, so it was soon replaced.

4. Rutherford’s alpha-particle experiment

Ernest Rutherford fired fast-moving, positively charged alpha (α) particles at a very thin sheet of gold foil. He expected them to pass almost straight through. The surprising results were: (i) most alpha particles passed straight through undeflected; (ii) some were deflected by small angles; and (iii) a very few (1 in 12000) bounced straight back. From this he concluded that most of the atom is empty space (most particles pass through), there is a tiny, dense, positively charged centre called the nucleus (the rare bounce-backs), and the nucleus is extremely small compared with the whole atom.

5. Rutherford’s nuclear model

Rutherford proposed that an atom has a tiny positively charged nucleus at the centre where nearly all the mass is concentrated, and electrons revolve around the nucleus in circular paths. The size of the nucleus is extremely small compared with the size of the atom. The biggest drawback of this model is that, according to physics, a charged particle moving in a circle should continuously give out energy, spiral inward, and crash into the nucleus — which would make the atom unstable. But atoms are stable, so this model needed correction.

6. Bohr’s model of the atom

Niels Bohr solved the stability problem. He proposed that electrons revolve only in certain fixed special orbits (also called energy levels or shells) and that while moving in these shells an electron does not lose energy. These shells are numbered from the nucleus outwards and named K, L, M, N… (corresponding to n = 1, 2, 3, 4…). An electron can move from one shell to another only by absorbing or releasing a fixed packet of energy.

7. How electrons are arranged (Bohr–Bury rules)

The arrangement of electrons in the shells of an atom is called its electronic configuration. The rules are: (i) the maximum number of electrons a shell can hold is 2n², where n is the shell number — so K holds 2, L holds 8, M holds 18, N holds 32; (ii) the outermost shell cannot hold more than 8 electrons; and (iii) shells are filled step by step — a new outer shell is started only after the inner shells are filled to a stable level. For example, magnesium (12 electrons) is arranged as 2, 8, 2.

8. Discovery of the neutron and the nucleus

The masses of atoms were greater than the mass of the protons alone could explain. In 1932 James Chadwick discovered the neutron, a particle with no charge and a mass nearly equal to that of a proton. Neutrons are present in the nucleus of all atoms (except ordinary hydrogen, which has just one proton and no neutron). Therefore the mass of an atom is the sum of the masses of its protons and neutrons, which together are called nucleons.

9. Valency

The electrons in the outermost shell are called valence electrons. The combining capacity of an atom is its valency. Atoms react in order to complete their outermost shell (8 electrons, or 2 for the first shell — the stable octet/duplet). If the outermost shell has 1, 2 or 3 electrons, the valency equals the number of valence electrons (electrons it tends to lose). If it has 5, 6 or 7 electrons, the valency equals 8 minus the number of valence electrons (electrons it tends to gain). Atoms with a complete outermost shell (like the noble gases) have a valency of zero.

10. Atomic number, mass number, isotopes and isobars

The atomic number (Z) is the number of protons in the nucleus — it decides which element an atom is. The mass number (A) is the total number of protons and neutrons. Isotopes are atoms of the same element (same atomic number) but different mass numbers (different numbers of neutrons), for example hydrogen has three isotopes — protium, deuterium and tritium; chlorine exists as Cl-35 and Cl-37. Isobars are atoms of different elements with the same mass number but different atomic numbers, such as calcium (Z = 20) and argon (Z = 18), both with mass number 40. Isotopes have many uses: an isotope of uranium is used as fuel in nuclear reactors, an isotope of cobalt is used to treat cancer, and an isotope of iodine is used to treat goitre.