Standard/ Class/ Grade XI ChemistryChapter 1 Basic ConceptsGurudatta K Wagh, gkwagh@gmail.com

Isotopes, Atomic mass, Molecular mass

Isotopes

Atoms of the same elements having same atomic number containing same number of protons and electrons but different number of neutrons

Have different mass numbers

Hydrogen – Three isotopes, H-1 (protium 1H, 1 proton and no neutron, 99.98 %), H-2 (1 proton and 1 neutron) (deuterium 2H), H-3 (tritium 3H) (1 proton and 2 neutrons), observed atomic mass 1.008 u

Neon-20 (19.9924 u), Ne-22 (21.9914 u), Ne-21 (20.9940 u). Average atomic mass of Ne = 20.1707 u

Observed atomic mass is the average atomic mass taking into consideration the natural abundance (per cent occurrence) of the isotopes

Atomic mass

Masses of atoms of elements is determined relative to mass of a standard

Carbon-12 is chosen as standard

Atomic mass unit (amu) or unified mass (u) of carbon is 12

Masses of other atoms are determined relative to the mass of an atom of carbon-12

Carbon-12 (98.89 %), C-13, C-14 are the three isotopes of carbon

Atomic mass unit (amu) or unified mass (u) of carbon is 12

Masses of other atoms are determined relative to the mass of an atom of Carbon-12

Molecular/ molar mass

Relative mass compared to mass of standard atom of Carbon-12

Ratio of mass of one molecule of a substance to 1/12th of mass of one atom of Carbon-12

Algebraic sum of atomic masses of constituent atoms

Unitless quantity

Molar mass expressed in gram is gram molar mass

1 gram molar mass = 1 gram molecule = 1 mole = 1 gram atom

It is calculated as the sum of the atomic mass of each constituent atom multiplied by the number of atoms of that element in the molecular formula

Molecular mass or molecular weight is the mass of a molecule. The adjective 'relative' is omitted as it is universally assumed that atomic and molecular masses are relative to the mass of 12C