The first quantitative information about atomic masses came from the work of Dalton, Gay Lussac, Lavoisier, Avogadro, and Berzelius. By observing the proportions in which elements combine to form various compounds, nineteenth-century chemists calculated relative atomic masses.
An atom is an extremely small particle, therefore, we cannot determine the mass of a single atom. However, it is possible to determine the mass of one atom of an element relative to another experimentally. This can be done by assigning a value to the mass of one atom of a given element so that it can be used as a standard.
By international agreement in 1961, the light isotope of carbon C-12 has been chosen as a standard. This isotope of carbon (C-12) has been assigned a mass of exactly 12 atomic mass units. This value has been determined accurately using a mass spectrometer.
The mass of atoms of all other elements is compared to the mass of C-12. Thus, "the mass of an atom of an element relative to the mass of an atom of C-12 is called its relative atomic mass".
One atomic mass unit (amu) is defined as a mass exactly equal to one-twelfth the mass of one C-12 atom.
Mass of one C-12 atom = 12 amu
1 amu = Mass of one C-12 atom / 12
A hydrogen atom is 8.40% as massive as the standard C-12 atom. Therefore, relative atomic mass of hydrogen.
= (8.40 / 100) × 12 amu
= 1.008 amu
Similarly, relative atomic masses of O, Na, Al are 15.9994 amu, 22.9898 amu, 26.9815 amu respectively. Table 3.1 shows the relative atomic masses of some elements.
| Element | Relative Atomic Mass |
|---|---|
| H | 1.008 amu |
| N | 14.0067 amu |
| O | 15.9994 amu |
| Na | 22.9898 amu |
| Al | 26.9815 amu |
| S | 32.06 amu |
| Cl | 35.453 amu |
| Fe | 55.847 amu |