Atomic mass

The atomic mass, the mass of a single atom, can be expressed like any mass in the SI unit kilogram (kg). As a rule, however, the mass an atom is expressed in atomic mass units

It is .

The unit , formerly denoted by (atomic mass unit), is one twelfth of the mass of an atom of the carbon isotope ^12C. It is thus about 8‰ smaller than the mass of a hydrogen atom ^1H. The choice of just this size is motivated, among other things, by the fact that then the numerical values for all known nuclides are close to an integer.

In biochemistry, and in the USA also in organic chemistry, the atomic mass unit is also called Dalton (unit symbol: Da), named after the English naturalist John Dalton.

In chemistry, on the recommendation of the IUPAC, the numerical value referred to as the relative atomic weight on its own, without a unit, and is formally understood as a separate, dimensionless quantity, namely as the mass ratio of the respective atom to an imaginary atom of mass . In contrast to this relative atomic mass, the mass expressed in kg, g or u is called absolute atomic mass.

The atomic masses of the nuclides are approximately integer multiples of the mass of the hydrogen atom. The deviations to the nearest integer are explained by the different masses of the proton and neutron and the mass defect. In lists such as Atomic Mass Adjustment 2012 and in interactive nuclide maps, the mass excess is often given instead of the atomic mass, sometimes both mass excess and atomic mass.

The mass ratios of the substances involved in a chemical reaction can be calculated from the atomic masses, the molecular masses that can be calculated from them and on the basis of the molar mass derived from them.

The average atomic mass of a mixed element is calculated as the weighted arithmetic mean of the atomic masses of the isotopes with the natural abundances of the isotopes as weights. In chemistry, this average atomic mass is called the atomic weight of the element.

Historical

The first table with relative atomic masses was published by John Dalton in 1805. He obtained them on the basis of the mass ratios in chemical reactions, choosing the lightest atom, the hydrogen atom, as the "mass unit" (see Atomic mass unit) - this, however, in ignorance of the property of hydrogen as a diatomic molecule.

Further relative atomic and molecular masses were calculated for gaseous elements and compounds on the basis of Avogadro's law, i.e. by weighing and comparing known gas volumes, later also with the help of Faraday's laws. Avogadro still referred to the smallest conceivable parts as molecules. Berzelius then introduced the term atom for the smallest conceivable part of a substance. He arbitrarily set the atomic weight of oxygen equal to 100. Later researchers chose the lightest substance, hydrogen, as the standard, but set the hydrogen molecule equal to 1. For carbon they then obtained the "equivalent weight" 6, for oxygen 8.

The actual pioneer for correct atomic weights of elements was Jean Baptiste Dumas. He determined the atomic weights for 30 elements very precisely and found that 22 elements had atomic weights that were multiples of the atomic weight of hydrogen.

It was not until 1858 that Stanislao Cannizzaro introduced today's distinction between atoms and molecules. He assumed that a molecule of hydrogen consisted of two atoms of hydrogen. For the individual hydrogen atom he arbitrarily set the atomic weight 1, a hydrogen molecule consequently has a molecular mass of 2. In 1865, oxygen, whose atoms have on average approximately 16 times the mass of the hydrogen atom, was proposed as a reference element by Jean Servais Stas and assigned a mass of 16.00. The mass of the hydrogen atom is 16.00.

In 1929, W. F. Giauque and H. L. Johnston discovered that oxygen has three isotopes. This prompted IUPAP to introduce a mass scale based on m(^16O), while IUPAC continued to use _Ar(O) = 16, i.e. oxygen in its natural isotopic composition.

In 1957, A. O. Nier and A. Ölander independently proposed that _Ar(^12C) and m(^12C) = 12 u should replace the old atomic mass units. IUPAP and IUPAC then agreed on this in 1959-1961. Until that time, therefore, physicists and chemists had two slightly different mass scales. In 1960, F. Everling, L. A. König, Josef Mattauch and Aaldert Wapstra published masses of nuclides.

Until today, the carbon isotope ^12C with the mass of 12 u serves as a reference base. The atomic mass indicates how many times greater the mass of the respective atom is than 1/12 of the mass of the ^{12C atom}. As mentioned above, the atomic masses of the nuclides are approximate, but not exact, integer multiples of the mass of the hydrogen atom.

The following table shows some average (see below) relative atomic masses, i.e. atomic weights, depending on the four different reference masses:

Element	related to
	natH = 1	natO = 16	16O = 16	12C = 12
natH	01,000	01,008	01,008	01,008
natCl	35,175	35,457	35,464	35,453
natO	15,872	16,000	16,004	15,999
natN	13,896	14,008	14,011	14,007
natC	11,916	12,011	12,015	12,011

Table with atomic weights in Johann Samuel Traugott Gehler's Physical Dictionary 1840