# General Chemistry/Formulas and Numbers

 ← Naming Substances · General Chemistry · Stoichiometry → Book Cover · Introduction ·  v • d • e

## Calculating Formula Masses

In molecules but not ionic compounds, the formula mass is also known as the molecular mass.

The calculation of a compound's formula mass (the mass of its molecule or formula unit) is straightforward. Simply add the individual mass of each atom in the compound (found on the periodic table). For example, the formula mass of glucose (C6H12O6) is 180 amu.

Molar masses are just as easy to calculate. The molar mass is equal to the formula mass, except that the unit is grams per mole instead of amu.

## Calculating Percentage Composition

Percentage composition is the relative mass of one substance in a compound compared to the whole. For example, in methane (CH4), the percentage mass of hydrogen is 25% because hydrogen makes up a total of 4 amu out of 16 amu overall.

### Using Percentage Composition

Percentage composition can be used to find the empirical formula of a compound, which shows the ratios of elements in the compound. However, this is not the same as the molecular formula. For example, many sugars have the empirical formula CH2O, which could correspond to a molecular formula of CH2O, C2H4O2, C6H12O6, etc.

To find the empirical formula from percentage composition, follow these procedures for each element.
1. Convert from percentage to grams (for simplicity, assume a 100 g sample).
2. Divide by the element's molar mass to find moles.
3. Simplify to lowest whole-number ratio.

For example, a compound is composed of 75% carbon and 25% hydrogen by mass. Find the empirical formula.

• 75g C / (12 g/mol C) = 6.25 mol C
• 25g H / (1 g/mol H) = 25 mol H
• 6.25 mol C / 6.25 = 1 mol C
• 25 mol H / 6.25 = 4 mol H

Thus the empirical formula is CH4.

### Calculating Molecular Formula

If you find the empirical formula of a compound and its molar/molecular mass, then you can find its exact molecular formula. Remember that the molecular formula is always a whole-number multiple of the empirical formula. For example, a compound with the empirical formula HO has a molecular mass of 34.0 amu. Since HO would only be 17.0 amu, which is half of 34.0, the molecular formula must be H2O2.