Molecular mass

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The molecular mass (m) of a substance is the mass of one molecule of that substance, in unified atomic mass unit(s) u[1] (equal to 1/12 the mass of one atom of the isotope carbon-12[2]). This is numerically equivalent to the relative molecular mass (Mr) of a molecule, frequently referred to by the term molecular weight, which is the ratio of the mass of that molecule to 1/12 of the mass of carbon-12 and is a dimensionless number. Thus, it is incorrect to express relative molecular mass (molecular weight) in daltons (Da). Unfortunately, the terms molecular weight and molecular mass have been confused on numerous websites, which often state that molecular weight was used in the past as another term for molecular mass.

Molecular mass differs from more common measurements of the mass of chemicals, such as molar mass, by taking into account the isotopic composition of a molecule rather than the average isotopic distribution of many molecules. As a result, molecular mass is a more precise number than molar mass; however it is more accurate to use molar mass on bulk samples. This means that molar mass is appropriate most of the time except when dealing with single molecules.

The concept of molecular mass is important for all molecules, especially for complex molecules like polymers and biopolymers such as proteins and carbohydrates. The determination of their molecular mass is often difficult, and is usually inferred from gel permeation chromatography and mass spectrometry.



There are varying interpretations of this definition. Many chemists use molecular mass as a synonym of molar mass,[3] differing only in units (see average molecular mass below). A stricter interpretation does not equate the two, as the mass of a single molecule is not the same as the average of an ensemble. Because a mole of molecules may contain a variety of molecular masses due to natural isotopes, the average mass is usually not identical to the mass of any single molecule. The actual numerical difference can be very small when considering small molecules and the molecular mass of the most common isotopomer in which case the error only matters to physicists and a small subset of highly specialized chemists; however it is always more correct, accurate and consistent to use 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. Also the molar mass comes into comparison with the number of particles contained within the structure by multiplying the number of moles by Avogadro's constant: NA=6.022 141 79(30)·1023 mol−1.

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