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## Atomic and Molecular Masses You know about the terms atoms and molecules. Thus, it is appropriate here to understand what we mean by atomic and molecular masses. ### Atomic Mass Every element has a characteristic atomic mass. Atomic mass is the mass of an atom. It is actually very very small....

## Atomic and Molecular Masses You know about the terms atoms and molecules. Thus, it is appropriate here to understand what we mean by atomic and molecular masses. ### Atomic Mass Every element has a characteristic atomic mass. Atomic mass is the mass of an atom. It is actually very very small. For example, the mass of one hydrogen atom is 1.6736 x 10-24g. This is very small quantity and not easy to measure. In the present system, the mass of an atom is determined relative to the mass of a carbon-12 atom as the standard, and this has been agreed upon in 1961 by IUPAC. In this system, an atom of carbon-12 is assigned a mass of exactly 12.00000 atomic mass unit (amu) and all other atoms of other elements are given a relative atomic mass, to that of carbon-12. The atomic masses are expressed in amu. **One amu** is defined as a mass exactly equal to one twelth of the mass of one carbon-12 atom. Later on, the exact value of atomic mass unit in grams was experimentally established. $1 \ amu = \frac{1}{12} * Mass \ of \ one \ C-12 = \frac{1}{12} * 1.992648 * 10^{-23} g = 1.66056 * 10^{-24} g$ Recently, amu has been replaced by the unified mass unit called dalton (symbol 'u' or 'Da'). 'u' means unified mass. **Problem 1.1:** Mass of an atom of oxygen in gram is 26.56896 x 10-24g. What is the atomic mass of oxygen in u? **Solution:** Mass of an atom of oxygen in gram is 26.56896 x 10-24g, and $\frac{26.56896 * 10^{-24} g}{1.66056 * 10^{-24} g/u} = 16.0 u$ Similarly, the mass of an atom of hydrogen is 1.0080 u. ### Average Atomic Mass Many naturally occurring elements exist as a mixture of more than one isotope. Isotopes have different atomic masses. The atomic mass of such an element is the weighted average of atomic masses of its isotopes (taking into account the atomic masses of isotopes and their relative abundance i.e., percent occurrence). This is called average atomic mass of an element. For example, carbon has the following three isotopes with relative abundances and atomic masses, as shown against each of them: | Isotope | Atomic Mass (u) | Relative Abundance (%) | |---|---|---| | 12C | 12.00000 | 98.892 | | 13C | 13.00335 | 1.108 | | 14C | 14.00317 | 2 x 10-10 | From the above data, the average atomic mass of carbon is: (12 u) (98.892/100) + (13.00335 u) (1.108/100) + (14.00317) (2 x 10-10/100) = 12.011 u Similarly, average atomic masses for other elements can be calculated. **Remember** In the periodic table of elements, the atomic masses mentioned for different elements are actually their average atomic masses. For practical purposes, the average atomic mass is rounded off to the nearest whole number when it differs from it by a very small fraction. | Element | Isotopes | Average Atomic Mass | Rounded off Atomic Mass | |---|---|---|---| | Carbon | 12C, 13C, 14C | 12.011 u | 12.0 u | | Nitrogen | 14N, 15N | 14.007 u | 14.0 u | | Oxygen | 16O, 17O, 18O | 15.999 u | 16.0 u | | Chlorine | 35Cl, 37Cl | 35.453 u | 35.5 u | | Bromine | 79Br, 81Br | 79.904 u | 79.9 u |