How to Calculate Molecular Mass From the Periodic Table

To calculate molecular mass, look up the atomic weight of each element in the formula, multiply it by that element’s subscript, then add the results. If an element has no subscript, count it once.

For example, water is H₂O. Using atomic weights of about 1.008 for hydrogen and 15.999 for oxygen:

(2 × 1.008) + (1 × 15.999) = 18.015 g/mol

The same arithmetic applies to molecular mass, formula mass, and molar mass. The terms differ by context: molecular mass refers to a molecule, formula mass is used for ionic compounds, and molar mass is the mass of one mole. The numerical results are commonly reported in g/mol.

Use atomic weight, not atomic number

A periodic table lists an atomic number and an atomic weight for each element. The atomic number is the whole number that identifies the element by its number of protons. It is not the value used in a mass calculation.

Use the atomic weight, usually shown as a decimal. These values account for the isotopes found in natural samples, so they are often not whole numbers. The International Union of Pure and Applied Chemistry (IUPAC) Commission on Isotopic Abundances and Atomic Weights publishes standard atomic weights CIAAW.

For ordinary calculations, use the value on your periodic table. If you need the mass of a specific isotope, consult an isotope reference such as the NIST Chemistry WebBook. An isotope-specific mass is different from the standard atomic weight.

Count every atom in the formula

Subscripts show how many atoms of an element are present. In H₂SO₄, for example, there are two hydrogen atoms, one sulfur atom, and four oxygen atoms. Multiply each element’s atomic weight by its count before adding.

Parentheses change how you read the count. In Ca(OH)₂, the subscript 2 applies to both elements inside the parentheses. The formula contains one calcium atom, two oxygen atoms, and two hydrogen atoms.

A hydrate uses a dot to separate the main compound from water molecules. In CuSO₄·5H₂O, the 5 applies to the whole H₂O unit. Count five water molecules, then add their mass to the CuSO₄ portion.

For any formula, use this sequence:

1. List each element and its atomic weight.
2. Count its atoms, including multipliers from parentheses or hydrate notation.
3. Multiply each atomic weight by its atom count.
4. Add the products.

Worked example: glucose

Glucose has the formula C₆H₁₂O₆. Using approximate atomic weights of 12.011 for carbon, 1.008 for hydrogen, and 15.999 for oxygen:

• Carbon: 6 × 12.011 = 72.066
• Hydrogen: 12 × 1.008 = 12.096
• Oxygen: 6 × 15.999 = 95.994

Add the contributions:

72.066 + 12.096 + 95.994 = 180.156 g/mol

The method is the same for small and large formulas. A written tally helps prevent missed subscripts and arithmetic errors.

Units and rounding

Atomic weights are relative values. When you use them to calculate the mass of one mole, the numerical result is expressed in g/mol. The molecular mass in u and the molar mass in g/mol have the same numerical value for a given formula, though they describe different quantities.

Keep the digits provided by your periodic table during the calculation, then round the final result to suit the precision of your inputs or the instructions for your problem. Do not add digits your source does not provide. For example, a table that lists carbon as 12.01 and oxygen as 16.00 gives CO₂ as:

12.01 + (2 × 16.00) = 44.01 g/mol

A table with more precise values may give a slightly different result. Use one set of atomic weights consistently.

If you are unsure which atomic weight to use, check a reliable periodic table, such as the IUPAC Periodic Table of Elements. For coursework, the values supplied in the problem or on the assigned periodic table usually take precedence.

Key Takeaways

• Multiply each element’s atomic weight by its count in the formula, then add the results.
• Use atomic weight, not atomic number.
• An omitted subscript means one atom; a subscript outside parentheses applies to every atom inside.
• In a hydrate, multiply the entire water unit by the number after the dot.
• Carry through the precision provided by your periodic table and round the final answer accordingly.