Atomic mass: Encyclopedia - Atomic mass

Atomic mass

The atomic mass of a chemical element (also known as the relative atomic mass or average atomic mass or atomic weight) is the average atomic mass of all the chemical element's isotopes as found in a particular environment, weighted by isotopic abundance. Periodic tables usually list these with reference to the local environment of Earth's crust and atmosphere. For artificial elements the nucleon count of the most stable isotope is listed in parentheses as the atomic mass.

The atomic mass of an isotope is the relative mass of the isotope, scaled with carbon-12 as exactly 12. No other isotopes have whole number masses due to the different mass of neutrons and protons, as well as loss/gain of mass to binding energy. However, since mass defect due to binding energy is minimal compared to the mass of a nucleon, rounding the atomic mass of an isotope tells you the total nucleon count. Neutron count can then be derived by subtracting the atomic number.

The pattern in the amounts the atomic masses deviate from their mass numbers is as follows: the deviation starts positive at hydrogen-1, becomes negative until a minimum is reached at iron-56, then increases to positive values in the heavy isotopes, with increasing atomic number. This corresponds to the following: nuclear fission in an element heavier than iron produces energy, and fission in any element lighter than iron requires energy; the opposite is true of nuclear fusion reactions - fusion in elements lighter than iron produces energy, and fusion in elements heavier than iron requires energy.

A similar definition applies to molecules; it is then called molecular mass. One can compute the molecular mass of a compound by adding the atomic masses of its constituent atoms multiplied by the ratios of elements given in the chemical formula. A similar formula mass can be calculated for those compounds which do not form molecules.

Direct comparison and measurement of the masses of atoms and molecules is achieved with mass spectrometry.

The standard scientific unit for dealing with atoms in macroscopic quantities is the mole (mol), which is defined arbitrarily as the amount of a substance with as many atoms or other units as there are in 12 grams of the carbon isotope C-12. The number of atoms in a mole is called Avogadro's number, the value of which is approximately 6 x 10²³. One mole of a substance always contains almost exactly the atomic or molecular mass of that substance, expressed in grams. For example, the atomic mass of iron is 55.847, and therefore one mole of iron has a mass of 55.847 grams.

Atomic mass - History

Before the 1960s, this was expressed so that the oxygen-16 isotope received the atomic weight 16, however, the proportions of oxygen-17 and oxygen-18 present in natural oxygen, which were also used to calculate atomic mass led to two different tables of atomic mass.

Formerly chemists and physicists used two different atomic mass scales. The chemists used a scale such that the natural mixture of oxygen isotopes had an atomic mass 16, while the physicists assigned the same number 16 to the atomic mass of the most common oxygen isotope (containing eight protons and eight neutrons). The unified scale based on carbon-12, ¹²C, met the physicists' need to base the scale on a pure isotope, while being numerically close to the old chemists' scale.

The term atomic weight is being phased out slowly and being replaced by relative atomic mass, in most current usage. The term standard atomic weight refers to the mean relative atomic mass of an element.

See also

• atomic mass unit
• isotope
• molecular mass
• Jean Stas

Adapted from the Wikipedia article "Atomic mass", under the G.N U Free Docmentation License. Please also see http://en.wikipedia.org/wiki