atomic mass

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atomic mass,

the mass of a single atomatom
[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
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, usually expressed in atomic mass unitsatomic 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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 (amu). Most of the mass of an atom is concentrated in the protons and neutrons contained in the nucleus. Each proton or neutron weighs about 1 amu, and thus the atomic mass is always very close to the mass numbermass number,
often represented by the symbol A, the total number of nucleons (neutrons and protons) in the nucleus of an atom. All atoms of a chemical element have the same atomic number (number of protons in the nucleus) but may have different mass numbers (from having
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 (total number of protons and neutrons in the nucleus). Atoms of an isotopeisotope
, 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 an 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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 all have the same atomic mass. Atomic masses are usually determined by mass spectrography (see 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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). They have been determined with great relative accuracy, but their absolute value is less certain.
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Atomic mass

The mass of an atom or molecule on a scale where the mass of a carbon-12 (12C) atom is exactly 12.0. The mass of any atom is approximately equal to the total number of its protons and neutrons multiplied by the atomic mass unit, u = 1.660539 × 10-24 gram. (Electrons are much lighter, about 0.0005486 u.) No atom differs from this simple formula by more than 1%, and stable atoms heavier than helium all lie within 0.3%. See Atomic mass unit

This simplicity of nature led to the confirmation of the atomic hypothesis—the idea that all matter is composed of atoms, which are identical and chemically indivisible for each chemical element. In 1802, G. E. Fischer noticed that the weights of acids needed to neutralize various bases could be described systematically by assigning relative weights to each of the acids and bases. A few years later, John Dalton proposed an atomic theory in which elements were made up of atoms that combine in simple ways to form molecules.

In reality, nature is more complicated, and the great regularity of atomic masses more revealing. Two fundamental ideas about atomic structure come out of this regularity: that the atomic nucleus is composed of charged protons and uncharged neutrons, and that these particles have approximately equal mass. The number of protons in an atom is called its atomic number, and equals the number of electrons in the neutral atom. The electrons, in turn, determine the chemical properties of the atom. Adding a neutron or two does not change the chemistry (or the name) of an atom, but does give it an atomic mass which is 1 u larger for each added neutron. Such atoms are called isotopes of the element, and their existence was first revealed by careful study of radioactive elements. Most naturally occurring elements are mixtures of isotopes, although a single isotope frequently predominates. Since the proportion of the various isotopes is usually about the same everywhere on Earth, an average atomic mass of an element can be defined, and is called the atomic weight. Atomic weights are routinely used in chemistry in order to determine how much of one chemical will react with a given weight of another. See Isotope

In contrast to atomic weights, which can be defined only approximately, atomic masses are exact constants of nature. All atoms of a given isotope are truly identical; they cannot be distinguished by any method. This is known to be true because the quantum mechanics treats identical objects in special ways, and makes predictions that depend on this assumption. One such prediction, the exclusion principle, is the reason that the chemical behavior of atoms with different numbers of electrons is so different.

McGraw-Hill Concise Encyclopedia of Physics. © 2002 by The McGraw-Hill Companies, Inc.
The following article is from The Great Soviet Encyclopedia (1979). It might be outdated or ideologically biased.

Atomic Mass


atomic weight; the value of the mass of the atom expressed in atomic mass units. The use of a particular unit for the measurement of the atomic mass is connected with the fact that the masses of atoms are extremely small (10-22 to 10-24g), and to express them in grams is inconvenient. The atomic mass unit is taken as 1/12 of the mass of the carbon isotope12 C. The mass of the carbon unit (CU) is equal to(1.66043 ± 0.00031) x 10-24g. In indicating atomic mass the symbol “CU” is generally omitted.

The concept of atomic mass was introduced by J. Dalton in 1803; he was the first to define atomic mass. A vast amount of work to establish atomic mass was carried out in the first half of the 19th century by J. Berzelius and later by J. S. Stas and T. W. Richards. In 1869, D. I. Mendeleev discovered the law of periodic dependence of the properties of the elements on the atomic mass and on this basis revised the atomic masses of many elements that were well known at that time (Be, U, La, and others), and in addition predicted the atomic masses of the still undiscovered elements Ga, Ge, and Sc. After the discovery by F. Soddy in 1914 of the phenomenon of isotopy, the concept of atomic mass came to be connected both with elements which consist of a mixture of isotopes and with individual isotopes. For elements which are represented in nature by one isotope (for example, F and Al), the atomic mass of the element coincides with the atomic mass of that isotope. If an element is a mixture of isotopes, then its atomic mass is calculated as the average of the values of the atomic masses of the individual isotopes, taking into account the relative percentage of each of them. Thus, natural chlorine consists of the isotopes 35CI (75.53 percent) and 37CI (24.47 percent), the masses of whose atoms are equal to 34.964 and 36.961 respectively. The atomic mass of the element Cl is

The fluctuation of the natural isotopic composition in the majority of elements is negligibly small (less than 0.003 percent); therefore each element has a practically constant atomic mass, which is one of the most important characteristics of the element. The closeness to integers of the atomic masses of elements that are represented in nature by one isotope is explained by the fact that almost all the mass of the atom is contained in its nucleus, and the masses of the protons and neutrons which compose the nucleus are close to I. At the same time, the atomic mass values of the isotopes (other than 12C, whose mass is taken to be 12.00000) are never exactly equal to a whole number. This is explained first by the fact that the relative masses of the neutron and proton are somewhat greater than 1 (1.0086654 and 1.00727663 respectively), second by the mass defect, and third by the small contribution to the total atomic mass of the masses of the electrons.

According to a proposal by J. Dalton (1803), the first unit of atomic mass was the mass of the hydrogen atom (the hydrogen scale). In 1818, Berzelius published a table of atomic masses relative to the atomic mass of oxygen, which was taken to be 100. The Berzelius system of atomic masses predominated until the 1860’s when chemists again adopted the hydrogen scale. In 1906, however, they converted to the oxygen scale, in which the atomic mass unit was taken to be 1/16 of the atomic mass of oxygen. After the discovery of the oxygen isotopes (16O, 17O, 18O) the atomic mass came to be indicated according to two scales—the chemical, based on 1/16 of the average mass of the natural oxygen atom; and the physical, with a mass unit equal to 1/16 of the mass of the l6O atom. The use of two scales had a number of shortcomings; as a consequence, in 1961 there was a conversion to a single (carbon) scale.

Various methods are used for the calculation of atomic mass. Some are based on the experimental determination of the molecular mass of some compound of a given element. In this case the atomic mass is some fraction of the molecular mass of this element divided by the number of the element’s atoms in the molecule. The exact values for atomic mass can be found by determining by chemical analysis the chemical equivalent of the element (the atomic mass is equal to the product of the equivalent times the valence). The atomic mass can be determined with the greatest accuracy (up to 0.001 percent and higher) by the method of mass spectroscopy; the mass spectrum of an element yields information concerning the quantitative isotopic composition and masses of atoms of individual isotopes, on the basis of which it is easy to compute the atomic mass (see example above with 35Cl and 37Cl). The atomic masses of newly synthesized elements are evaluated on the basis of observation of the atomic reaction accompanying their formation.


Mendeleev, D. I. Osnovy khimii, 13th ed., vols 1–2. Moscow-Leningrad, 1947.
Nekrasov, B. V. Osnovy obshchei khimii, vol. 1. Moscow, 1965.
Pauling, L. Obshchaia khimiia. Moscow, 1964. (Translated from English.)
Remy, H. Kurs neorganicheskoi khimii, vol. 1. Moscow, 1963. (Translated from German.)
Gina, M. Istoriia khimii. Moscow, 1966. (Translated from Italian.)


The Great Soviet Encyclopedia, 3rd Edition (1970-1979). © 2010 The Gale Group, Inc. All rights reserved.

atomic mass

[ə′täm·ik ′mas]
The mass of a neutral atom usually expressed in atomic mass units.
McGraw-Hill Dictionary of Scientific & Technical Terms, 6E, Copyright © 2003 by The McGraw-Hill Companies, Inc.
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