Atoms being extremely small, cannot be seen or weighed directly. But indirect methods of physics have enabled us to know the absolute mass of nearly all kinds of atoms. The mass of a hydrogen atom is found to be 1.66 × 10^-24 g while that of a carbon atom is 1.9926 × 10^-23 g. As these masses are too small, it is not convenient to use kilograms or grams as unit. It has, therefore, been considered appropriate to use the mass of some standard atom as a unit and then relate masses of other atoms to it.

(a) What do you understand by the statement "R.A.M. of silver is 108" ?

(b) Which element is considered as the standard for atomic masses ?

(c) What is the difference between R.A.M. and R.M.M.? Give an example to explain.

(d) What is the significance of amu or μ ?

(e) How is 1 amu related to grams ?

(f) Is the atomic mass expressed in amu the actual mass of an atom of that element ?

(g) Why is amu preferred instead of grams or kilogram for calculating the atomic mass of an atom ?

(a) “R.A.M. of silver is 108” means that one atom of silver is 108 times as heavy as 1⁄12^th the mass of one carbon-12 atom.

(b) The element considered as the standard for atomic masses is carbon-12.

(c) R.A.M. means Relative Atomic Mass. It tells us how many times one atom of an element is heavier than $\frac{1}{12}$^th the mass of one carbon-12 atom.

Example:
R.A.M. of oxygen = 16.
This means one atom of oxygen is 16 times heavier than $\frac{1}{12}$^th the mass of one carbon-12 atom.

R.M.M. means Relative Molecular Mass. It tells us how many times one molecule of a substance is heavier than $\frac{1}{12}$^th the mass of one carbon-12 atom.

Example:
R.M.M. of water, H₂O:

R.M.M. of H₂O = 2 × R.A.M. of H + 1 × R.A.M. of O
= 2 × 1 + 16 = 18

So, the R.M.M. of water is 18. This means one molecule of water is 18 times heavier than $\frac{1}{12}$^th the mass of one carbon-12 atom.

Thus, R.A.M. is used for atoms of elements, while R.M.M. is used for molecules of elements or compounds.

(d) The significance of amu is that it provides a convenient unit for expressing the extremely small masses of atoms and molecules. It simplifies calculations by avoiding extremely small values in grams.

(e) 1 amu = 1.66 × 10^-24 g

(f) No, the atomic mass expressed in amu is not the actual mass of an atom in grams. It is a relative mass expressed by comparing the mass of the atom with $\frac{1}{12}$^th the mass of a carbon-12 atom.

(g) The amu is preferred because the actual masses of atoms in grams or kilograms are extremely small and inconvenient to use. The amu gives a simple, practical and manageable scale for comparing atomic masses.