Molar mass

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Molar mass is the mass of one mole of a chemical element or chemical compound.^[1] It is commonly used in stoichiometric calculations of bulk substances in chemistry. Its primary purpose is as a conversion factor between the number of grams of a pure substance, which can be measured directly, and the number of moles of that substance, which has greater chemical significance. This allows for using the appropriate number of molecules of a substance regardless of the mass. For example, if an equal number of molecules of two substances are needed for a reaction but the molar mass of one substance is twice that of the other, twice as many grams will be needed of that substance to give the same number of molecules.

The molar mass of a chemical substance may be computed from the standard atomic weights listed for the elements on a standard periodic table. A mole of a substance is defined to be approximately 6.023x10²³ (see Avogadro's number) of particles of the substance. Thus the molar mass is the mass of 6.023x10²³ particles of the substance. Molar mass is different from Molecular mass which is the mass of one molecule.

In chemistry, the unit of molar mass is g/mol due to chemical utility. In physics, molar mass is usually defined in kilograms per mole (kg/mol) because the base SI unit of mass is the kilogram.

In linear polymers not every polymer chain consists of the same amount of repeating units. A given polymer sample is said to be made up of a mixture of macromolecules with a certain molar mass distribution.^[2]

[edit] Example

Let us see approximately how many grams are in 2.3 moles of table sugar with a chemical formula of C₁₂H₂₂O₁₁. The standard atomic weights of carbon, hydrogen, and oxygen are approximately 12.011 , 1.008 , and 15.999 g/mol respectively. Thus the molar mass of sucrose is its sum: (12.011 * 12) + (1.008 * 22) + (15.999 * 11) = 342.297 g/mol. The mass of 2.3 moles of sugar is then 2.3 mol * 342.297 g/mol = 787.2831 g.

(Note: For simplicity this example ignores the proper use of significant figures)

[edit] Molar mass versus molecular mass

Molar mass is sometimes confused with the related but distinct molecular mass. This is largely due to that when the molar mass and molecular mass are expressed in g/mol and u respectively they will almost always have similar but not identical numerical values. The molar mass is generally computed from isotopically weighted averages, whereas the molecular mass is the mass of a single molecule consisting of well-defined isotopes. The isotopically weighted averages used to compute molar masses are those found in most versions of the periodic table and are numbers recommended by IUPAC. They represent the most likely weights of substances found in the laboratory. The averaging takes into account the natural abundance of, usually heavier, isotopes as well as the variation in their natural abundance in different places on earth. Additionally the confidence, or number of significant figures after the decimal, is different. The significant figures in the standard atomic weights and thus the computed molar masses are often limited by the natural variations in the isotopic distributions and not necessarily by our ability to measure accurately. The confidence in the isotopic masses and resulting molecular masses are only limited by the accuracy of measurement of the invariable isotopic masses.

It is common, even amongst professional chemists, to use the terms interchangeably since for most common applications the difference is insignificant. This can, however, on occasion lead to substantive confusion. Due to this common practice some areas of chemistry have developed their own more specific terms such as monoisotopic mass and average mass. Due to these subtle differences and the inherent nature of the molar mass it is always more correct, accurate and consistent to use molar mass in any bulk stoichiometric calculations.