atomic weight


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Related to atomic weight: molecular weight

atomic weight,

mean (weighted average) of the masses of all the naturally occurring isotopesisotope
, in chemistry and physics, one of two or more atoms having the same atomic number but differing in atomic weight and mass number. The concept of isotope was introduced by F.
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 of a chemical elementelement,
in chemistry, a substance that cannot be decomposed into simpler substances by chemical means. A substance such as a compound can be decomposed into its constituent elements by means of a chemical reaction, but no further simplification can be achieved.
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, as contrasted with atomic massatomic mass,
the mass of a single atom, usually expressed in atomic mass units (amu). Most of the mass of an atom is concentrated in the protons and neutrons contained in the nucleus.
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, which is the mass of any individual isotope. Although the first atomic weights were calculated at the beginning of the 19th cent., it was not until the discovery of isotopes by F. Soddy (c.1913) that the atomic mass of many individual isotopes was determined, leading eventually to the adoption of the atomic mass unitatomic mass unit
or amu,
in chemistry and physics, unit defined as exactly 1-12 the mass of an atom of carbon-12, the isotope of carbon with six protons and six neutrons in its nucleus. One amu is equal to approximately 1.66 × 10−24 grams.
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 as the standard unit of atomic weight. For many elements with two or more stable isotopes, atomic weight is now expressed as a variable interval with a lower and upper bound instead of a single number; in the case of hydrogen, for example, this is typically written as: [1.00784; 1.00811].

Atomic weights were formerly determined directly by chemical means; now a mass spectrographmass spectrograph,
device used to separate electrically charged particles according to their masses; a form of the instrument known as a mass spectrometer is often used to measure the masses of isotopes of elements. J. J. Thomson and F. W. Aston showed (c.
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 is usually employed. The atomic mass and relative abundance of the isotopes of an element can be measured very accurately and with relative ease by this method, whereas chemical determination of the atomic weight of an element requires a careful and precise quantitative analysis of as many of its compounds as possible.

Development of the Concept of Atomic Weight

J. L. Proust formulated (1797) what is now known as the law of definite proportions, which states that the proportions by weight of the elements forming any given compound are definite and invariable. John Dalton proposed (c.1810) an atomic theory in which all atoms of an element have exactly the same weight. He made many measurements of the combining weightscombining weight,
the proportion (by weight) in which a chemical element combines with other elements to form compounds. The determination of combining weights was a very important part of early chemical endeavor.
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 of the elements in various compounds. By postulating that simple compounds always contain one atom of each element present, he assigned relative atomic weights to many elements, assigning a weight of 1 to hydrogen as the basis of his scale. He thought that water had the formula HO, and since he found by experiment that 8 weights of oxygen combine with 1 weight of hydrogen, he assigned an atomic weight of 8 to oxygen. Dalton also formulated the law of multiple proportions, which states that when two elements combine in more than one proportion by weight to form two or more distinct compounds, their weight proportions in those compounds are related to one another in simple ratios. Dalton's work sparked an interest in determining atomic weights, even though some of his results—such as that for oxygen—were soon shown to be incorrect.

While Dalton was working on weight relationships in compounds, J. L. Gay-Lussac was experimenting with the chemical reactions of gases, and he found that, when under the same conditions of temperature and pressure, gases react in simple whole-number ratios by volume. Avogadro proposed (1811) a theory of gases that holds that equal volumes of two gases at the same temperature and pressure contain the same number of particles, and that these basic particles are not always single atoms. This theory was rejected by Dalton and many other chemists.

P. L. Dulong and A. T. Petit discovered (1819) a specific-heat method for determining the approximate atomic weight of elements. Among the first chemists to work out a systematic group of atomic weights (c.1830) was J. J. Berzelius, who was influenced in his choice of formulas for compounds by the method of Dulong and Petit. He attributed the formula H2O to water and determined an atomic weight of 16 for oxygen. J. S. Stas later refined many of Berzelius's weights. Stanislao Cannizzaro applied Avogadro's theories to reconcile atomic weights used by organic and inorganic chemists.

The availability of fairly accurate atomic weights and the search for some relationship between atomic weight and chemical properties led to J. A. R. Newlands's table of "atomic numbers" (1865), in which he noted that if the elements were arranged in order of increasing atomic weight "the eighth element, starting from a given one, is a kind of repetition of the first." He called this the law of octaves. Such investigations led to the statement of the periodic lawperiodic law,
statement of a periodic recurrence of chemical and physical properties of the elements when the elements are arranged in order of increasing atomic number.
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, which was discovered independently (1869) by D. I. Mendeleev in Russia and J. L. Meyer in Germany. T. W. Richards did important work on atomic weights (after 1883) and revised some of Stas's values.

atomic weight

[ə′täm·ik ′wāt]
(chemistry)
The relative mass of an atom based on a scale in which a specific carbon atom (carbon-12) is assigned a mass value of 12. Abbreviated at. wt. Also known as relative atomic mass.
References in periodicals archive ?
Berzelius published his table of atomic weights in 1826 with quite a few that were not an integral multiple of that of hydrogen.
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By the time researchers recognized in 1913 that elements should be arranged by atomic number (the number of protons in their nuclei) rather than by atomic weight, only seven gaps remained in the list of naturally occurring elements.