2.1.3 Atomic masses

2.1.3 Atomic masses

1. Measurement of relative masses

   • Carbon-12 isotope is the international standard for measurement of relative mass
   • Atomic masses are measured using the unit called unified atomic mass unit, u
   • 1u is a tiny mass
   • The mass of an atom of carbon-12 is 12u
   • The mass of 1/12th of an atom of Carbon-12 is 1u

2. Relative isotopic mass

   • The mass of an atom of an isotope compared to 1/12th of the mass of an atom of Carbon-12
   • Relative isotopic mass= mass number
   • E.g. Oxygen-16 = 16.0

3. Relative atomic mass, Ar

   • The weighted mean mass of an atom of an element compared with 1/12th of an atom of Carbon-12
   • Weighted mean mass accounts for the contribution made by each isotope to the overall mass of an element
   • Contribution made by an isotope to overall mass depends on:
     • % abundance
     • Relative mass
   • EXAMPLE: A sample of Br contains 53.0% of Br-79 and 47.0% of Br-81. Work out the relative atomic mass
     • ANSWER: (53x79) + (47x81) / 100 = 79.94

4. Relative molecular mass

   • The mass of a molecule is measured as their relative molecular ,ass by comparison with C-12
   • You find Mr by adding together all the Ar of each atoms in the molecules
   • E.g. Cl2 = 35.5 x 2 = 71.0

5. Relative formula mass

   • Compounds with giant structures dont exist as simple molecules
   • It's the same as finding relative molecular mass but is called relative formula mass
   • Relative molecular mass would imply counting all atoms and adding the masses which gives an extremely large number
   • E.g. CaBr2 = 40.1 + (79.9 x 2) = 199.9