Atomic mass
The atomic mass is the mass of an atom. It can be like any mass specified in the SI unit kilogram ( kg). However, for calculations it is often convenient to move (Da) to use atomic mass unit u (formerly with amu, atomic mass unit called, ) or Dalton. This is the twelfth of the mass of an atom of carbon isotope 12C. In SI units is 1 u = 1 Da = 1.660 538 921 (73 ) × 10-27 kg.
The numerical value of in u given atomic mass but without the unit of measure is often referred to as relative atomic mass ( engl. atomicweight ) and regarded formally as a separate, non-dimensional parameter, namely, as the mass ratio of each atom u to an imaginary atom of mass 1 or 1 Da.
From the atomic masses, the resulting predictable molecular weights and the basis of the derived molar mass, the mass and volume ratios of the substances involved in a chemical reaction can be calculated.
Historical
The first table of relative atomic masses was published in 1805 by John Dalton. He received based on the mass ratios in chemical reactions, in which he, as a "mass unit " chose the lightest atom, the hydrogen atom (see atomic mass unit ).
Later, the calculation of the relative masses of atoms and molecules of gaseous elements or compounds on the basis of Avogadro's law carried, that is obtained by weighing a known volume of gas, then with the help of Faraday's law. For Avogadro, the smallest conceivable parts were still referred to as molecules. Berzelius introduced the term atom for the smallest conceivable part of a substance. Arbitrarily, he put the atomic weight of oxygen equal to 100 Subsequent researchers chose the lightest, hydrogen, as standard, but put the hydrogen molecule is equal to 1 for carbon they then received the equivalent weight of 6, 8 for oxygen
The real pioneer of correct atomic weights of elements was Jean Baptiste Dumas. He determined for 30 elements very accurately the atomic weights and found that 22 elements had atomic weights that are multiples of the atomic weight of hydrogen.
Stanislao Cannizzaro was only introduced in 1858, today a distinction between atom and molecule. He assumed that a hydrogen molecule consists of two atoms of hydrogen. For the single hydrogen atom he sat arbitrary atomic weight 1, a hydrogen molecule thus has a molecular mass of 2
In 1865, oxygen, whose atoms have approximately the average 16 times the mass of the hydrogen atom, proposed by Jean Servais Stas as a reference element and assigned him the mass 16.00. As the physicist later this value to the oxygen isotope 16O, the chemist but zuordneten the oxygen in its natural isotopic composition, so that two slightly different mass scales were used until about 1960.
Since the decision of the IUPAP, 1961 approved the proposal of their atomic mass Commission from 1960, the carbon isotope 12C is used as the reference point with the mass of 12.00. The relative atomic mass indicates how many times larger is the mass of each atom as 1/12 the mass of the 12C atom. Since the atom has 12 nucleons, 6 protons and 6 neutrons corresponds to the atomic mass of any nuclide almost exactly the number of nucleons contained in the nucleus, the mass number; the small deviation is caused by the mass difference between the proton and neutron, and the atomic mass defect.
Measurement
Accurate atomic masses are now determined with mass spectrometers. Here, the atomic masses of the individual isotopes yield very precise. To determine the relative atomic masses of the elements in their natural isotopic composition ( average atomic mass, also called atomic weight ) then has to be determined yet the isotopic ratio. For purposes of this chemistry average atomic mass of natural isotope mixture is specified in the earth's crust; in special cases must be considered the origin of the isotopic mixture.
Other examples of the relative atomic masses of some chemical elements:
- Silicon (Si): 28.0855
- Gold ( Au): 196.966569
- Iron (Fe): 55.845