# How to Find Molecular Mass

Molecular mass is the mass of one molecule. Usually, you report molecular mass (m) in units of daltons (Da) or atomic mass units (amu or u). Here is how to find molecular mass, example calculations, and a look at the difference between molecular mass and molar mass.

• Write the molecular formula for the molecule or compound.
• Look up the atomic masses of each element in the formula.
• Multiply each atomic mass by its subscript in the formula.
• Add up the atomic masses of each element.

### How to Find Molecular Mass

Calculating molecular mass requires that you know the number and type of atoms that make up the molecule. Once you know the atomic makeup of the molecule, you use the periodic table to find the mass of each atom and add them together. The number of atoms of each element is its subscript in the molecular formula. If there is no subscript following an element symbol, it means there is one atom of that element in the formula.

For example, water has the molecular formula H2O. This means one water molecule contains two hydrogen atoms and one oxygen atom.

The periodic table shows the atomic mass of hydrogen is 1.0008 amu and the atomic mass of oxygen is 15.999 amu. The molecular mass of water will be 2 masses of hydrogen plus the mass of oxygen.

molecular mass of H2O = (2 x 1.008 amu) + (15.999 amu)
molecular mass of H2O = 2.016 amu + 15.999 amu
molecular mass of H2O = 18.015 amu

The molecular mass of water is 18.015 amu.
Remember that one amu is equivalent to 1 gram/mole. This is how you can find the molecular weight of a molecule. This means one mole of water weighs 18.015 grams.

### Molecular Mass Example Calculation

Example:
What is the molecular mass of sulfuric acid (H2SO4)?

Solution:
Each sulfuric acid molecule contains 2 hydrogen atoms, 1 sulfur atom, and 4 oxygen atoms. So the molecular mass is:

Molecular mass of H2SO4 = (2 x mass of hydrogen) + (1 x mass of sulfur) + (4 x mass of oxygen)

From the periodic table, we see

mass of H = 1.008 amu
mass of S = 32.066 amu
mass of O = 15.999 amu

Plug these values into the formula.

Molecular mass of H2SO4 = (2 x 1.008 amu) + (1 x 32.066 amu) + (4 x 15.999 amu)
Molecular mass of H2SO4 = 2.016 amu + 32.066 amu + 63.996 amu
Molecular mass of H2SO4 = 98.078 amu

The molecular mass of sulfuric acid is equal to 98.078 amu or 98.078 grams/mole.

### Difference Between Molecular Mass and Molar Mass

The molar mass of a compound is the sample mass divided by the number of moles. So, while you report molecular mass in daltons or amu, molar mass is in kilograms per mole (kg/mol) or grams per mole (g/mol). Technically, the usual molecular mass calculation is really a molar mass calculation. This is because the calculation uses the average atomic masses from the periodic table, which are weighted mass averages based on isotope abundance of natural elements.

Most of the time, molecular mass and molar mass are interchangeable. But, there are two exceptions.

1. You might not use the relative atomic mass values from the periodic table. For example, consider the molecular mass of D2O, where D stands for the hydrogen isotope deuterium. Here, the atomic mass of the isotope is 2.014 (not 1.008, like you use for “hydrogen” in general). The molecular mass is (2 x 2.014) + 15.999 = 20.027 amu. The same situation applies any time you use samples with known isotopic ratios.
2. If you really want the mass of a single molecule, you might call the value “molecular mass.” Here, you take the molar mass and divide it by Avogadro’s number. For example, the mass of a single water molecule is 18.015 g/mol ÷ 6.022 x 1023 molecules/mol. So, the mass of one individual water molecule is approximately 2.991×10−23 g.

### References

• International Bureau of Weights and Measures (2006). The International System of Units (SI) (8th ed.). ISBN 92-822-2213-6.
• International Union of Pure and Applied Chemistry (1993). Quantities, Units and Symbols in Physical Chemistry (2nd ed.). Oxford: Blackwell Science. ISBN 0-632-03583-8.
• Possolo, Antonio; van der Veen, Adriaan M. H.; Meija, Juris; Hibbert, D. Brynn (2018). “Interpreting and propagating the uncertainty of the standard atomic weights (IUPAC Technical Report)”. Pure and Applied Chemistry. 90 (2): 395–424. doi:10.1515/pac-2016-0402
• Wieser, M. E. (2006). “Atomic Weights of the Elements 2005.” Pure and Applied Chemistry. 78 (11): 2051–66. doi:10.1351/pac200678112051