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Definition: atom from Dictionary of Energy

Physics. the basic component of all matter; the smallest particle of an element that can still retain all the chemical properties of that element. An atom has a dense central core (the nucleus) consisting of positively charged particles (protons) and uncharged particles (neutrons). Negatively charged particles (electrons) are scattered in a relatively large space around this nucleus and move about it in orbital patterns at extremely high speeds. An atom contains the same number of protons as electrons and thus is electrically neutral and stable under most conditions. See also atomic theory.


Summary Article: atom from The Columbia Encyclopedia

[Gr.,=uncuttable (indivisible)], basic unit of matter; more properly, the smallest unit of a chemical element having the properties of that element.

Structure of the Atom

The atom consists of a central, positively charged core, the nucleus, and negatively charged particles called electrons that are found in orbits around the nucleus.

The Nucleus

Almost the entire mass of the atom is concentrated in the nucleus, which occupies only a tiny fraction of the atom's volume. The nucleus of an atom consists of neutrons and protons, the neutron being an uncharged particle and the proton a positively charged one. Their masses are almost equal. Atoms containing the same number of protons but different numbers of neutrons represent different forms, or isotopes, of the same element.

The Electrons

Surrounding the nucleus of an atom are its electrons; for a neutral atom, the number of electrons is equal to the atomic number. The outermost electrons of an atom determine its chemical and electrical properties. An atom may combine chemically with another atom in various ways, either by giving up or receiving electrons, thus setting up an electrical attraction between the atoms (see ion), or by sharing one or more pairs of electrons (see chemical bond). Because metals have few outermost electrons and tend to give them up easily, they are good conductors of electricity or heat (see conduction).

The electrons are often described as revolving about the nucleus as the planets revolve about the sun. This picture, however, is misleading. The quantum theory has shown that all particles in motion also have certain wave properties. For a particle the size of an electron, these properties are of considerable importance. As a result the electrons in an atom cannot be pictured as localized in space, but rather should be viewed as smeared out over the entire orbit so that they form a cloud of charge. The electron clouds around the nucleus represent regions in which the electrons are most likely to be found. The shapes of these clouds can be very complex, in marked contrast to the simple elliptical orbits of planets. Surprisingly, the sizes of all atoms are comparable, in spite of the large differences in the number of electrons they contain.

Atomic Weight and Number

The atomic number of an atom is simply the number of protons in its nucleus. The atomic weight of an atom is given in most cases by the mass number of the atom, equal to the total number of protons and neutrons combined. An atom may be conveniently symbolized by its chemical symbol with the atomic number and mass number written as subscript and superscript, respectively. For example, the symbol for uranium is U (atomic number 92); the isotopes of uranium with atomic weights 235 and 238 are indicated by and .

Development of Atomic Theory
Early Atomic Theory

The atomic theory, which holds that matter is composed of tiny, indivisible particles in constant motion, was proposed in the 5th cent. B.C. by the Greek philosophers Leucippus and Democritus and was adopted by the Roman Lucretius. However, Aristotle did not accept the theory, and it was ignored for many centuries. Interest in the atomic theory was revived during the 18th cent. following work on the nature and behavior of gases (see gas laws).

From Dalton to the Periodic Table

Modern atomic theory begins with the work of John Dalton, published in 1808. He held that all the atoms of an element are of exactly the same size and weight (see atomic weight) and are in these two respects unlike the atoms of any other element. He stated that atoms of the elements unite chemically in simple numerical ratios to form compounds. The best evidence for his theory was the experimentally verified law of simple multiple proportions, which gives a relation between the weights of two elements that combine to form different compounds.

Evidence for Dalton's theory also came from Michael Faraday's law of electrolysis. A major development was the periodic table, devised simultaneously by Dmitri Mendeleev and J. L. Meyer, which arranged atoms of different elements in order of increasing atomic weight so that elements with similar chemical properties fell into groups. By the end of the 19th cent. it was generally accepted that matter is composed of atoms that combine to form molecules.

Discovery of the Atom's Structure

In 1911, Ernest Rutherford developed the first coherent explanation of the structure of an atom. Using alpha particles emitted by radioactive atoms, he showed that the atom consists of a central, positively charged core, the nucleus, and negatively charged particles called electrons that orbit the nucleus. There was one serious obstacle to acceptance of the nuclear atom, however. According to classical theory, as the electrons orbit about the nucleus, they are continuously being accelerated (see acceleration), and all accelerated charges radiate electromagnetic energy. Thus, they should lose their energy and spiral into the nucleus.

This difficulty was solved by Niels Bohr (1913), who applied the quantum theory developed by Max Planck and Albert Einstein to the problem of atomic structure. Bohr proposed that electrons could circle a nucleus without radiating energy only in orbits for which their orbital angular momentum was an integral multiple of Planck's constant h divided by 2π. The discrete spectral lines (see spectrum) emitted by each element were produced by electrons dropping from allowed orbits of higher energy to those of lower energy, the frequency of the photon of light emitted being proportional to the energy difference between the orbits.

Around the same time, experiments on x-ray spectra (see X ray) by H. G. J. Moseley showed that each nucleus was characterized by an atomic number, equal to the number of unit positive charges associated with it. By rearranging the periodic table according to atomic number rather than atomic weight, a more systematic arrangement was obtained. The development of quantum mechanics during the 1920s resulted in a satisfactory explanation for all phenomena related to the role of electrons in atoms and all aspects of their associated spectra. With the discovery of the neutron in 1932 the modern picture of the atom was complete.

Contemporary Studies of the Atom

With many of the problems of individual atomic structure and behavior now solved, attention has turned to both smaller and larger scales. On a smaller scale the atomic nucleus is being studied in order to determine the details of its structure and to develop sources of energy from nuclear fission and fusion (see nuclear energy), for the atom is not at all indivisible, as the ancient philosophers thought, but can undergo a number of possible changes. On a larger scale new discoveries about the behavior of large groups of atoms have been made (see solid-state physics). The question of the basic nature of matter has been carried beyond the atom and now centers on the nature of and relations between the hundreds of elementary particles that have been discovered in addition to the proton, neutron, and electron. Some of these particles have been used to make new types of exotic "atoms" such as positronium (see antiparticle) and muonium (see muon).

Bibliography
  • See Gamow, G., The Atom and Its Nucleus (1961).
  • H. A. Boorse and L. Motz, ed., The World of the Atom (2 vol., 1966).
  • Bransden, B. H.;Joachain, C. J., Physics of Atoms and Molecules (1986).
The Columbia Encyclopedia, © Columbia University Press 2017

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