Relative Atomic Mass
The mass of an atom of an element as compared to 1/12th mass of an atom of carbon taken as 12 is called relative atomic mass (atomic mass unit).
The mass of an atom is very small. It cannot be weighed by any known physical balance. That is why we use the relative atomic mass unit scale.
Examples:
H = 1.008 a.m.u, Cl = 35.5 a.m.u. These are the average of atomic masses of isotopes proportional to their relative abundances.
Unit:
The unit used to express atomic mass is called atomic mass unit (a.m.u).
- 1 a.m.u = 1.661 × 10–27 kg
- 1 a.m.u = 1.661 × 10–24 g
Relative Isotopic Mass
The mass of an isotope of an element as compared to 1/12th mass of an atom of carbon-12 is called relative isotopic mass.
Examples:
- Relative atomic mass of Chlorine-35 = 35 a.m.u.
- Relative atomic mass of Chlorine-37 = 37 a.m.u.
Relative Molecular Mass
The mass of a molecule as compared to 1/12th mass of an atom of carbon-12 is called relative molecular mass.
Examples:
- Relative molecular mass of water (H2O) = 18 a.m.u
- Relative molecular mass of carbon dioxide (CO2) = 44 a.m.u
Relative Formula Mass
The mass of a formula unit as compared to 1/12th mass of an atom of carbon-12 is called relative formula mass.
Examples:
- Relative formula mass of sodium chloride (NaCl) = 58.5 a.m.u
- Relative formula mass of sodium sulphate (Na2SO4) = 142 a.m.u
Note: Carbon-12 is taken as standard on the C-12 scale because:
- It is a highly stable isotope.
- Its mass is exactly a whole number (12.000).
- It can be handled easily because it is a solid.
Average Atomic Masses
The average atomic mass depends on the number of possible isotopes and their natural abundance.
Formula:
Average atomic mass = Σ (isotopic mass × fractional abundance)
Example (Neon isotopes):
- Neon-20: Abundance = 90.92%, Mass = 20 a.m.u
- Neon-21: Abundance = 0.26%, Mass = 21 a.m.u
- Neon-22: Abundance = 8.82%, Mass = 22 a.m.u
Relative atomic mass of Neon = (20 × 0.9092) + (21 × 0.0026) + (22 × 0.0882) = 20.18 a.m.u
