Relative Atomic Mass and Relative Molecular Mass/Relative Formula Mass

By international agreement the mass of an atom of carbon-12 is given as 12 unified atomic mass units (u). Therefore, 1 u is ¹/₁₂th the mass of an atom of carbon-12.

In most chemistry work, you can:

• neglect the tiny contribution to atomic mass from electrons

• take both the mass of a proton and a neutron as being equal to 1 u.

In a pure isotope all the atoms have the same atomic structure and the same mass.

Relative isotopic mass is then simply the mass number of the isotope, e.g. the relative isotopic mass of ³²S is 32.

However, most elements contain a mixture of isotopes, each with a different mass number. To find the relative atomic mass of an element we must know:

• The relative isotopic mass of each isotope.

• The proportion of each isotope in the element.

   

For example:

Copper has 2 isotopes, ⁶³Cu and ⁶⁵Cu. The proportions of each isotope are 69.17% and 30.83%.

These can be rounded to 70% and 30%.

The relative atomic mass of Copper is therefore (⁷⁰/₁₀₀ x 63) + (³⁰/₁₀₀ x 65) = 63.6

There is no unit as it is a relative value.

Molecular Mass (M_r) is the sum of all the relative atomic masses for all the atoms in a given formula.

For covalent compounds it is called the Relative Molecular Mass.

For ionic compounds it is called the Relative Formula Mass.

Methane (CH₄) for example, contains 1 Carbon atom and 4 Hydrogen atoms.

Therefore, the molecular mass is calculated as follows:

C = 12, H = 1

CH₄ = C + (H x 4) = 12 + (1 x 4) = 16