atom

The arrangement of electrons in a sodium atom and a sulphur atom. The number of electrons in a neutral atom gives that atom its atomic number: sodium has an atomic number of 11 and sulphur has an atomic number of 16.

The structure of a sodium atom. The nucleus is composed of 11 protons and 12 neutrons. Eleven electrons orbit the nucleus in 3 orbits: 2 in the inner orbit, 8 in the middle, and 1 in the outer.

Smallest unit of matter that can take part in a chemical reaction, and which cannot be broken down chemically into anything simpler. An atom is made up of protons and neutrons in a central nucleus (except for hydrogen, which has a single proton as its nucleus) surrounded by electrons (see atomic structure). The atoms of the various elements differ in atomic number, relative atomic mass, and chemical behaviour.

Atoms are much too small to be seen by even the most powerful optical microscope (the largest, caesium, has a diameter of 0.0000005 mm/0.00000002 in), and they are in constant motion. However, modern electron microscopes, such as the scanning tunnelling microscope (STM) and the atomic force microscope (AFM), can produce images of individual atoms and molecules.

Two Greek theories Among the ancient Greeks there were two theories as to the nature of matter, or substance. Some, such as Anaxagoras and Aristotle, held that matter was infinite and continuous, and that therefore any substance could theoretically be divided and subdivided to an infinite extent. Others, such as Democritus and Epicurus, taught that matter was grained, that is, consisted of minute particles which could not be divided. Both theories were based on naturally slender experimental evidence.

The conservation of matter Towards the end of the 18th century, the development of experimental chemistry demanded greater quantitative exactness, and experimental evidence, primarily from studies in combustion, led to the principle of the conservation of matter. The value of this principle has been enormous, particularly in the direction of detecting new elements.

Dalton's theory John Dalton, in the 19th century, believed that gases consisted of particles or ‘corpuscles’. Particles of a compound must therefore be divisible into atomic particles of the atoms combined. Dalton enunciated the law of constant proportions, which states that when two elements unite to form a compound they do so in a constant ratio that is characteristic of that compound. For instance, when oxygen and hydrogen combine to form water, the weights combining always take the same ratio.

Determining atomic weights Shortly after Dalton's atomic theory had been enunciated, Joseph Gay-Lussac investigated the volumetric conditions of gases in combination, with the result that he discovered and published the law that when gases combine, they do so in volumes which bear a simple ratio to one another and to that of their product (if gaseous). In 1811 Amadeo Avogadro published his hypothesis on the molecular constitution of gases, which asserts that under the same conditions of temperature and pressure equal volumes of all gases contain the same number of molecules whether those molecules consist of single atoms or many atoms in combination. Both hypotheses were well supported by experimental evidence, and were used to determine the relative atomic masses of the elements. Much of the progress in chemistry has been based on quantitative analysis using atomic weights.

Rutherford and Moseley It became apparent that atoms have structure and are not indivisible, around 1900. From his experiments with alpha-particles, Ernest Rutherford and others (1911–13) showed that practically the whole mass of any atom is concentrated in an extremely small central nucleus bearing a positive electrical charge. With Henry Moseley in 1913 he showed that the nucleus contains a number of positive charges dependent on the element, and called the atomic number of the element. Around the nucleus move an equal number of electrons at a relatively great distance. The lightest nucleus, the hydrogen nucleus, contains a single positive charge, and is called a proton.

Bohr In 1913, Niels Bohr proposed that the electrons move in orbits around the nucleus-like planets round the Sun, and suggested how atoms might emit or absorb light. These ideas were developed and applied with great success by Bohr and others using quantum theory, to the full elucidation of atomic structure, and the explanation of the properties of matter in bulk, and of the substructure of the nucleus itself.

Chadwick In 1932, James Chadwick discovered that the bombardment of beryllium by alpha-particles produced neutral particles which he called neutrons. From the atomic weights of atoms, and the known weights of the proton and the electron it became clear that (1) protons and neutrons have essentially equal masses, and (2) that atomic nuclei contain approximately equal numbers of protons and neutrons, the protons carrying the nuclear charge.