Structure of the Atom

By the end of the nineteenth century, scientists found that atoms are not the simplest, indivisible particles of matter, as had earlier been believed. Experiments with electric discharge through gases at low pressure revealed that atoms are made up of even smaller charged particles.

J.J. Thomson discovered the electron in 1897 while studying the rays produced in a discharge tube. He showed that these rays were streams of negatively charged particles, far lighter than the atom itself. Every atom, regardless of the element, contains electrons.

E. Goldstein discovered new radiation in a discharge tube, called canal rays, which led to the discovery of positively charged particles named protons. A proton carries a positive charge equal in magnitude to the negative charge of an electron.

Later, J. Chadwick discovered a third sub-atomic particle present in the nucleus that carries no charge at all, called the neutron. For example, an ordinary hydrogen atom has no neutron, while a carbon atom has six neutrons along with six protons in its nucleus.

J.J. Thomson proposed that an atom consists of a sphere of positive charge with negatively charged electrons embedded in it, much like seeds in a watermelon or plums in a pudding. Since the positive and negative charges were equal in magnitude, the atom as a whole was electrically neutral. This is often called the plum pudding model.

Rutherford tested this idea with his famous alpha-particle scattering experiment. A beam of fast-moving, positively charged alpha particles was directed at a thin sheet of gold foil surrounded by a fluorescent screen. Three observations were made: most alpha particles passed straight through the foil undeflected; a small fraction were deflected through small angles; and a very small number (about 1 in 12000) bounced back almost the way they had come, deflected by nearly 180°.

From these results, Rutherford concluded that most of the space inside an atom is empty (since most particles passed through undisturbed), that almost all the mass and the entire positive charge is concentrated in a very small volume at the centre called the nucleus (since only a rare, head-on encounter could reverse a fast alpha particle), and that electrons revolve around this nucleus in circular paths, just as planets revolve around the sun. This is called the nuclear model of the atom.

However, Rutherford's model had a serious flaw: an electron moving in a circular orbit is continuously accelerating, and an accelerating charged particle should radiate energy according to the laws of electromagnetism. If the revolving electron kept losing energy this way, it would gradually spiral inward and fall into the nucleus, making the atom highly unstable. Since atoms are known to be stable, this model could not explain the stability of the atom.

To overcome the instability problem of Rutherford's model, Niels Bohr proposed a new model of the atom. According to Bohr's postulates, electrons revolve around the nucleus only in certain fixed circular paths called orbits or shells, each having a definite energy. As long as an electron stays in a particular orbit, it does not lose or radiate energy; such orbits are therefore called stationary states or energy levels.

These shells are numbered 1, 2, 3, 4 counting outward from the nucleus, and are also labelled with the letters K, L, M, N respectively. The energy of an electron increases as its distance from the nucleus increases, so the K shell (closest to the nucleus) has the lowest energy and shells further out have progressively higher energy.

For example, in a sodium atom, electrons are not scattered randomly but occupy the K, L and M shells in a definite, stepwise pattern, giving the atom a stable, well-defined structure. Bohr's model successfully explained why atoms do not collapse, something Rutherford's model had failed to do.

Bohr and Bury gave a set of rules for how electrons are distributed among the shells of an atom. The maximum number of electrons that can be accommodated in a shell with number n is given by 2n². So the K shell (n = 1) can hold at most 2 electrons, the L shell (n = 2) at most 8, the M shell (n = 3) at most 18, and the N shell (n = 4) at most 32.

Two further rules apply: the outermost shell of an atom cannot hold more than 8 electrons, and electrons are not filled into a new shell until the inner shells nearer the nucleus are completely filled. For example, sodium has 11 electrons, distributed as 2 in the K shell, 8 in the L shell and the remaining 1 in the M shell, written as 2, 8, 1.

The electrons present in the outermost shell of an atom are called valence electrons, and they decide the valency of the element, that is, its combining capacity with other atoms. If the outermost shell is completely filled (a stable, "full" arrangement of 8 electrons, or 2 for the K shell alone, as in noble gases), the valency is zero. If the number of valence electrons is 1, 2, 3 or 4, the valency is usually equal to that number. If the number of valence electrons is 5, 6 or 7, the valency is usually 8 minus that number, since the atom needs that many more electrons to complete its outer shell of 8.

For example, sodium (2, 8, 1) has 1 valence electron, so its valency is 1. Oxygen (2, 6) has 6 valence electrons, so its valency is 8 − 6 = 2.

The atomic number (Z) of an atom is the number of protons present in its nucleus. Since an atom is electrically neutral, the atomic number is also equal to the number of electrons in the atom. Every element has a unique atomic number, which is why it is used to identify the element. For instance, every atom with atomic number 6 is a carbon atom, and every atom with atomic number 8 is an oxygen atom.

The mass number (A) of an atom is the total number of protons and neutrons present in its nucleus, since these particles (called nucleons) contribute almost all the mass of the atom. It is written as: mass number = number of protons + number of neutrons.

Sometimes atoms of the same element have different numbers of neutrons, giving them different mass numbers while their atomic number (and hence their chemical identity) stays the same. Such atoms are called isotopes. For example, hydrogen has three isotopes: protium (no neutron), deuterium (1 neutron) and tritium (2 neutrons), all with atomic number 1 but mass numbers 1, 2 and 3 respectively. Isotopes of an element have the same chemical properties (since they have the same number of valence electrons) but different physical properties (since their masses differ).

In contrast, atoms of different elements can sometimes have the same mass number but different atomic numbers; these are called isobars. For example, an atom of calcium (atomic number 20) and an atom of argon (atomic number 18) both have mass number 40.