Grams to Atomic Mass Units Converter
Common Conversions
| g | amu |
|---|---|
| 1e-27 | 0.000602 |
| 1e-26 | 0.00602 |
| 1e-25 | 0.0602 |
| 1.661e-24 | 1 |
| 1e-23 | 6.022 |
| 1.993e-23 | 12 |
| 2.656e-23 | 15.999 |
| 1e-22 | 60.22 |
| 1e-20 | 6022 |
| 1e-10 | 60220000000000 |
| 1 | 6.022e+23 |
Why this conversion matters in chemistry
Grams and amu describe mass at two radically different scales — benchtop and atomic — but they're tied together by a clean definitional relationship. Avogadro's number of amu is exactly 1 gram. So 1.000 g of carbon-12 is 6.022 × 10²³ u in total mass, which works out to exactly 5.018 × 10²² atoms each weighing 12.000 u. The conversion is multiplying by Avogadro's number (6.022 × 10²³), which is the same as dividing by the amu-to-grams factor (1.66054 × 10⁻²⁴ g/u). The real use case is computational: bulk-experimental measurements come in grams, quantum-mechanical calculations live in per-particle u, and the conversion is what lets the two meet.
Formula
Where the factor comes from
Watch what happens to exactness in this one. The factor is one gram divided by the atomic mass constant, and that constant — a twelfth of the mass of a free ¹²C atom — has to be measured in grams, because the kilogram is defined through the Planck constant and retains no tie to any atom. Its value is near 1.660539 × 10⁻²⁴ g, known to a few parts in 10¹⁰. Invert that and one gram comes to about 6.022141 × 10²³ u. The resemblance to the Avogadro constant is not accidental: before 2019 the two were the same number by definition. They no longer are. Nₐ is now exact at 6.02214076 × 10²³ mol⁻¹, while this factor is experimental, and the two agree to roughly a part in 10⁹ rather than perfectly.
Precision and significant figures
The factor is good to better than a part in 10⁹, several orders of magnitude beyond any balance, so it never limits an answer. The mass on the pan does. A four-place analytical balance reading 1.0000 g offers five figures at best, and only once buoyancy, static charge and drift have been dealt with; a top-loading balance offers three or four. Carrying 6.02214076 × 10²³ through a calculation that opened with 0.25 g is theatre — 6.022 × 10²³ says everything the input supports. Round the factor to the figures the mass had, and remember the result is still a mass, just written on the atomic scale, so it inherits the weighing’s uncertainty unchanged.
Worked Examples
The defining mass of one atomic mass unit — 1/12 of a carbon-12 atom, pinned by convention.
Avogadro's number of amu. The elegant equivalence that lets macroscopic mass and atomic mass share numerical values.
The mass of a single oxygen atom. Small, exact, and the anchor for most elemental mass calculations.
The mass of a single neon atom, using the standard atomic weight of 20.180 u.
Common mistakes
A mass in u is not a count
One gram is about 6.022 × 10²³ u of mass, but one gram of carbon-12 holds only 5.018 × 10²² atoms — a twelfth as many, because each atom carries 12 u. The two figures look alike enough to be swapped in a hurry, and the resulting error is the size of the atomic mass.
Standard atomic weight applied to one atom
15.999 u is an abundance-weighted average across the oxygen isotopes present in terrestrial material. No individual atom has that mass: ¹⁶O is 15.9949 u, with ¹⁷O and ¹⁸O above it. For a bulk sample the average is the right number to convert; for a single-particle statement or a high-resolution spectrum it is not.
Summing nucleon masses to get atomic mass
Six protons, six neutrons and six electrons add to roughly 12.099 u, yet a ¹²C atom is exactly 12 u by definition. The missing 0.099 u is nuclear binding energy showing up as mass. Atomic masses have to come from measurement rather than being assembled from constituents, and the gap is far too large to be rounding.