Grams per Mole to AMU Converter
Common Conversions
| g/mol | amu |
|---|---|
| 1 | 1 |
| 2 | 2 |
| 12 | 12 |
| 16 | 16 |
| 18.015 | 18.015 |
| 28 | 28 |
| 44 | 44 |
| 58.44 | 58.44 |
| 98.08 | 98.08 |
| 180.16 | 180.16 |
| 342.3 | 342.3 |
Why this conversion matters in chemistry
Molar mass in g/mol describes a mole of particles; atomic mass in u (amu) describes a single particle. The numerical values are identical by construction — water at 18.015 g/mol is also 18.015 u per molecule, and the same identity holds for every other compound. The equivalence falls out of the mole definition itself: Avogadro's number was chosen so that 12 g of carbon-12 atoms is exactly Nₐ atoms × 12 u each. The conversion is the identity, useful at the junction between molar-scale stoichiometry and single-molecule mass-spectrometric measurements.
Formula
Where the factor comes from
The bridge between the two scales is a quantity with a name: the molar mass constant, Mᵤ, defined as the molar mass of carbon-12 divided by twelve. Any molar mass can be written M(X) = Aᵣ(X) × Mᵤ, where Aᵣ is the dimensionless relative atomic or molecular mass — precisely the number an amu value carries. If Mᵤ is exactly 1 g/mol, the two figures are identical and there is nothing to compute. Until 2019 it was, because the mole was defined as the number of atoms in 12 g of ¹²C, which forced the equality by construction. The 2019 revision fixed the Avogadro constant instead, at 6.02214076 × 10²³ mol⁻¹, leaving Mᵤ to be determined experimentally. It now sits within about one part in 10⁹ of 1 g/mol, so the identity survives as a measurement rather than a definition.
Precision and significant figures
That part-in-10⁹ gap between the scales is smaller than anything that will constrain you. The standard atomic weights you started from are what constrain you. They are measured values, and for more than a dozen elements they are published as intervals rather than single numbers, because natural isotopic composition varies with the source of the material — carbon spans roughly 12.0096 to 12.0116 depending on where it came from. Quoting an organic molar mass to four decimals therefore claims more than the atomic weights support, unless the isotopic makeup of that particular sample is known. Two decimals is honest for synthetic work. Digits past that belong to monoisotopic masses, a different quantity built from single nuclide masses rather than abundance-weighted averages.
Worked Examples
Water — the textbook example, where the molar mass and the per-molecule mass write as the same number.
Carbon-12 by definition — the calibration anchor that pins the entire u-scale.
Sodium chloride formula mass — useful in any calculation bridging salt stoichiometry and per-ion mass spectrometry.
Glucose — the standard molar-mass reference for any monosaccharide-related calculation.
Common mistakes
The equality is numerical, not dimensional
g/mol is mass per amount of substance; u is mass. They share a number and nothing else. Saying a water molecule weighs 18.015 g/mol is dimensionally wrong — one molecule is 18.015 u, which is about 2.99 × 10⁻²³ g. The slip is harmless in conversation and expensive inside a calculation where the units are supposed to cancel.
Old literature on a pre-1961 amu scale
Two amu scales coexisted before 1961: physicists referenced ¹⁶O = 16, chemists referenced natural oxygen = 16, and the two differed by roughly 275 parts per million. Unification onto ¹²C settled it. Values copied out of older sources can still carry that offset, which matters only in high-precision work but is entirely invisible when it does.
Formula mass treated as a molecular mass
58.44 for sodium chloride is a formula unit, not a molecule — the crystal contains no discrete NaCl entity to weigh. The number converts to u perfectly well and means the summed mass of one Na and one Cl, which is the right quantity for stoichiometry and the wrong one for describing a species in a mass spectrum.