Molecular mass

Definition

There are varying interpretations of this definition. Many chemists use molecular mass as a synonym of molar mass in any bulk stoichiometric calculations. The size of this error becomes much larger when considering larger molecules or less abundant isotopomers. The molecular mass of a molecule which happens to contain heavier isotopes than the average molecule in the sample can differ from the molar mass by several mass units.

Average molecular mass

The average molecular mass (sometimes abbreviated as average mass) is another variation on the use of the term molecular mass. The average molecular mass is the abundance weighted mean (average) of the molecular masses in a sample. This is often closer to what is meant when "molecular mass" and "molar mass" are used synonymously and may have derived from shortening of this term. The average molecular mass and the molar mass of a particular substance in a particular sample are in fact numerically identical and may be interconverted by standard atomic weights, whereas the average molecular mass, in fields that need the term, is often a measured figure specific to a sample. Therefore, they often vary since one is theoretical and the other is experimental. Specific samples may vary significantly from the expected isotopic composition due to real deviations from earth average isotopic abundances.

Computing the Molecular Mass

The molecular mass can be calculated as the sum of the individual isotopic masses (as found in a table of isotopes) of all the atoms in any one molecule. This is possible because molecules are created by chemical reactions which, unlike nuclear reactions, have very small binding energies compared to the rest mass of the atoms ( < 10^-9) and therefore create a negligible mass defect. Note that the use of average atomic masses derived from the standard atomic weights found on a standard periodic table will result in an average molecular mass, whereas the use of isotopic masses will result in a molecular mass consistent with the strict interpretation of the definition, i.e. that of a single molecule. Note that any given molecule may contain any given combination of isotopes, so there may be multiple molecular masses for each chemical compound.

Measuring the Molecular Mass

The molecular mass can also be measured directly using standard atomic weights found on a typical periodic table, since there is likely to be a statistical distribution of atoms representing the isotopes throughout the molecule. This however may differ from the true average molecular mass of the sample due to natural (or artificial) variations in the isotopic distributions.

Example: Average Molecular Mass versus Molecular Mass versus Molar Mass

The Avogadro's constant approximately 6.022*10²³. The most common units of molar mass are g/mol because in those units the numerical value equals the average molecular mass in units of u.

Conversion Factor of average molecular mass to molar mass:

molar mass = average molecular mass * (6.022*10^-23g/u)*(6.022*10²³/mol)

or

molar mass in g/mol= average molecular mass in u

(Note that these relations are true for theoretical and experimental values, but not between experimental and theoretical values. Molar mass is most often theoretical and average molecular mass is most often experimental)

The average atomic mass of natural oxygen is 15.9994 u; therefore, the molecular mass of natural water with formula H₂O is (2 × 1.00794 u) + 15.9994 u = 18.01528 u. Therefore, one mass spectrometry and particle physics (where the mixture of isotopes does not act as an average).

There are also situations where the isotopic distributions are not typical such as with standard atomic weights, will not be the same as the actual molar mass or average molecular mass of the sample. In this case the mass of deuterium is 2.0136 u and the average molecular mass of this water (assuming 100% deuterium enrichment) is (2 × 2.0136 u) + 15.9994 u = 20.0266 u. This is a very large difference of ~11% error from the expected average molecular mass based on the standard atomic weights. Furthermore the most abundant molecular mass is actually slightly less than the average molecular mass since oxygen-16 is still the most common. (2 × 2.0136 u) + 15.9949 u = 20.0221 u. Although this is an extreme artificial example, natural variation in isotopic distributions do occur and are measurable.

See also

• Avogadro's number
• Atomic mass unit
• Molar mass
• Monoisotopic mass
• gram-molecular weight