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Atomic mass

Atomic mass

Atomic mass is the mass of an atom. The unit of measure of the SI is one kilogram; In fact, a unit outside the system, an atomic mass unit, is generally used.

One of the fundamental properties of an atom is its mass. The absolute mass of an atom is an extremely small amount. Therefore, a hydrogen atom has a mass of approximately 1.67 × 10 -2 4 g. Therefore, in chemistry (mainly for practical purposes) it is much more convenient to use a relative (conditional) value, which is called relative atomic mass or simply atomic mass and that shows how many times the atomic mass of a given element is greater than Atomic mass of another element taken as a unit of mass.

As a unit of measurement of atomic and molecular masses, 1 ⁄ 12 of the mass of the neutral atom of the most common carbon isotope 12 C was adopted. This unit of non-systemic mass measurement is called the atomic unit of mass (AU) or Dalton (Yes).

The difference between the atomic mass of the isotope and its mass number is called excess mass (usually expressed in M ​​eV). It can be both positive and negative; The reason for its appearance is the non-linear dependence of the binding energy of the nuclei on the number of protons and neutrons, as well as the difference in the masses of protons and neutrons.

The dependence of the atomic mass of the isotope on the mass number is as follows: the excess mass is positive for hydrogen-1, with an increase in the mass number decreases and becomes negative until it reaches a minimum for iron- 56, then begins to grow and increases to positive values ​​for heavy nuclides. This corresponds to the fact that fission of nuclei heavier than iron releases energy, while fission of light nuclei requires energy. On the contrary, the fusion of lighter nuclei than iron releases energy, while the fusion of heavier elements than iron requires additional energy.

The atomic mass of a chemical element (also "average atomic mass", "standard atomic mass") is the weighted average atomic mass of all stable and unstable isotopes of this chemical element that exist in nature, taking into account its natural prevalence (percentage) in the earth's crust and atmosphere. It is this atomic mass that is presented in the periodic table of DI Mendeleev; It is used in stoichiometric calculations. The atomic mass of an element with a violated isotope ratio (for example, enriched in some isotopes) differs from the standard. For monoisotopic elements. (such as iodine, gold, etc.) the atomic mass of an element coincides with the atomic mass of its only isotope present in a natural mixture. For chemical elements that are absent in nature (synthetic chemical elements), such as technetium, curium, etc., the atomic mass of an element conventionally indicates the mass number of the most stable isotopes known from this element; these values ​​in the periodic table are traditionally indicated in square brackets.

Relative atomic mass

Relative atomic mass (obsolete name - atomic weight): the value of the mass of an atom, expressed in atomic units of mass. It is defined as the ratio of the mass of the element of the atom to 1/12 the mass of the neutral atoms of the carbon isotope 12 C. From the definition it follows that the relative atomic mass is a dimensionless quantity.

Molecular Mass (Molar)

The molecular mass of a chemical compound is the sum of the atomic masses of the elements that compose it, multiplied by the stoichiometric coefficients of the elements by the chemical formula of the compound. Strictly speaking, the mass of a molecule is less than the mass of its constituent atoms in an amount equal to the binding energy of the molecule (see above). However, this mass defect is 9-10 orders of magnitude smaller than the mass of the molecule, and can be neglected.

The definition of a mole (and Avogadro's number) is chosen so that the mass of a mole of a substance (molar mass), expressed in grams (per mole), is numerically equal to the atomic mass (or molecular) of this substance For example, the atomic mass of iron is 55,847 a. E. m. Therefore, one mole of iron (that is, the number of iron atoms equal to Avogadro's number, .06.022⋅10 23 ) has a mass of 55.847 g.

Direct comparisons and measurements of the masses of atoms and molecules are made using mass spectrometry methods.

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Last review: December 2, 2019