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atomic structure

Internal structure of an atom.

The nucleus

The core of the atom is the nucleus, a dense body only one ten-thousandth the diameter of the atom itself. The simplest nucleus, that of hydrogen, comprises a single stable positively charged particle, the proton. Nuclei of other elements contain more protons and additional particles, called neutrons, of about the same mass as the proton but with no electrical charge. Each element has its own characteristic nucleus with a unique number of protons, the atomic number. The number of neutrons may vary. Where atoms of a single element have different numbers of neutrons, they are called isotopes. Although some isotopes tend to be unstable and exhibit radioactivity, all those of a single element have identical chemical properties.

Electrons

The nucleus is surrounded by a number of moving electrons, each of which has a negative charge equal to the positive charge on a proton, but which has a mass of only 1/1,836 times as much. In a neutral atom, the nucleus is surrounded by the same number of electrons as it contains protons. According to quantum theory, the position of an electron is uncertain; it may be found at any point. However, it is more likely to be found in some places than others. The region of space in which an electron is most likely to be found is called an atomic orbital. The chemical properties of an element are determined by the ease with which its atoms can gain or lose electrons.

Attraction and repulsion

According to the theory of fundamental forces, atoms are held together by the electrical forces of attraction between each negative electron and the positive protons within the nucleus. The latter repel one another with enormous forces; a nucleus holds together only because an even stronger force, called the strong nuclear force, attracts the protons and neutrons to one another. The strong force acts over a very short range - the protons and neutrons must be in virtual contact with one another. If, therefore, a fragment of a complex nucleus, containing some protons, becomes only slightly loosened from the main group of neutrons and protons, the natural repulsion between the protons will cause this fragment to fly apart from the rest of the nucleus at high speed. It is by such fragmentation of atomic nuclei (nuclear fission) that nuclear energy is released.

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 nucleuslike 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.

Subatomic particles Research into high-energy particle physics has established the existence of subatomic particles other than the proton, neutron, and electron. More than 300 kinds of particle are now known, and these are classified into several classes according to their mass, electric charge, spin, magnetic moment, and interaction. The elementary particles, which include the electron, are indivisible and may be regarded as the fundamental units of matter; the hadrons, such as the proton and neutron, are composite particles made up of either two or three elementary particles called quarks.

atom Bohr, Niels Henrik David Deisenhofer, Johann electron essence exclusion principle

Kleppner, Daniel matter Pauli exclusion principle Sidgwick, Nevil Vincent X-ray diffractometer