Chapter 8: Journey Inside the Atom

Acharya Kanada said that repeated division ends in the smallest particles, the parmanus, and the Greek thinkers Leucippus and Democritus called them atomos, meaning indivisible. These were ideas of thought, not of experiment. In 1808 Dalton gave the first scientific atomic theory. Then radioactivity showed that atoms contain smaller particles. In 1897 J. J. Thomson studied cathode rays in a tube and found that they are negatively charged particles, much lighter than atoms, and the same for every material. They were named electrons, with charge −1. Thomson pictured the atom as a sphere of positive charge with electrons embedded in it, the plum pudding model.

In 1911 Geiger and Marsden, working under Rutherford, aimed alpha particles at a thin gold foil. Most passed straight through, some were deflected through large angles, and a very few bounced back. This is called scattering. Rutherford concluded that most of the atom is empty space and that the positive charge and most of the mass are packed in a tiny nucleus, with the electrons revolving outside it like planets round the Sun. An atom is about 10⁻¹⁰ m across and a nucleus about 10⁻¹⁵ m, so the nucleus is about 100000 times smaller. The model could not explain why the atom is stable, since a revolving electron should lose energy and fall into the nucleus.

In 1913 Niels Bohr proposed that electrons move only in fixed circular paths called shells, orbits or energy levels, labelled K, L, M, N or n = 1, 2, 3, 4. In such a stationary state the electron does not lose energy. The K shell is nearest to the nucleus and has the least energy, and the energy rises outward. An electron can jump from one shell to another only by absorbing or releasing a fixed amount of energy equal to the difference between the two levels. Bohr's model has a nucleus with electrons moving in circles, as in Rutherford's, but it explains the stability of the atom. It was later improved by the quantum mechanical model.

An atom contains the electron, the proton and the neutron. Their relative charges are −1, +1 and 0. Protons and neutrons are packed in the nucleus, and the electrons move outside it. A neutral atom has equal numbers of protons and electrons, so its total charge is zero. A proton and a neutron have almost the same mass, while the mass of an electron can be ignored. Chadwick discovered the neutron in 1932, and it explained why a helium atom, with two protons, is about four times as heavy as hydrogen. Neutrons also help to hold the protons together in the nucleus, and heavier atoms generally have more neutrons than protons.

In 1803 Dalton used pictures as symbols, and in 1813 Berzelius suggested letters from the Latin names. Today IUPAC approves the symbols. The first letter is a capital and the second, if any, is small, as in Al and Co. Some symbols come from other languages, such as Fe from ferrum, Na from natrium, K from kalium, Cu, Ag, Au, Pb and Hg, and W from the German wolfram. The atomic number Z is the number of protons and it fixes the identity of the element. The mass number A is the number of protons plus neutrons, which are called nucleons. The number of neutrons is A − Z. The standard notation puts A at the top left and Z at the bottom left of the symbol.

The distribution of electrons among the shells is the electronic configuration. By the Bohr-Bury rules, the maximum number of electrons in the nth shell is 2n², so K holds 2, L holds 8, M holds 18 and N holds 32. The outermost shell can hold at most 8 electrons. Electrons fill the shells step by step, starting from the one nearest the nucleus, K then L then M, and a shell is filled only after the inner shells are full. For example, carbon (6 electrons) is 2, 4, sodium (11) is 2, 8, 1 and argon (18) is 2, 8, 8.

The outermost shell that has electrons is the valence shell, and its electrons are valence electrons. A complete shell of 8 electrons, an octet, or 2 electrons in helium, makes an atom stable and largely unreactive. Other atoms lose, gain or share electrons to complete the octet. The number of electrons gained, lost or shared is the valency, or combining capacity, which is measured in hydrogen or chlorine atoms. An atom with fewer than four valence electrons tends to lose them, and one with more than four tends to gain. Sodium (2, 8, 1) has valency 1, oxygen (2, 6) has valency 2, and carbon (2, 4) shares four electrons, so its valency is 4.

Isotopes are atoms of the same element with the same atomic number but different mass numbers, so they differ in the number of neutrons. Hydrogen has protium (0 neutrons), deuterium (1) and tritium (2), and carbon has carbon-12, carbon-13 and carbon-14. Isotopes have the same chemical properties because they have the same electrons, but different physical properties. Uranium-235 is a reactor fuel, cobalt-60 treats cancer, iodine-131 treats thyroid problems and carbon-14 dates fossils. Isobars are atoms of different elements with the same mass number, such as argon, potassium and calcium, each with mass number 40. The average atomic mass of an element is a weighted average that takes the abundance of each isotope into account.