Empirical and Molecular Formulae

Chemists use two types of formulae to describe compounds. They help us understand what atoms are present and in what ratios.

Key ideas

• Empirical formula: the simplest whole-number ratio of atoms in a compound. Example: hydrogen peroxide is H₂O₂, but its empirical formula is HO.
• Molecular formula: the actual number of each type of atom in one molecule. Example: glucose is C₆H₁₂O₆; its empirical formula is CH₂O.

Think of a recipe: the empirical formula is the simplest version of the recipe, while the molecular formula is the full-size recipe used.

Finding an empirical formula from data

Use masses or percentages of each element. If you have percentages, assume 100 g so the numbers become grams.

Convert each element to moles:

$$\text{moles} = \frac{\text{mass (g)}}{A_r}$$

• Divide all mole values by the smallest to get a ratio.
• If you get simple fractions (like 1.5 or 1.33), multiply all by 2 or 3 to make whole numbers.

Finding a molecular formula

You need the empirical formula and the relative molecular mass, $M_r$, of the compound.

Calculate the empirical formula mass (add the $A_r$ values). Then find the multiplier $n$:

$$n = \frac{M_r(\text{molecule})}{M_r(\text{empirical})}$$

Multiply all subscripts in the empirical formula by $n$.

Common misconceptions

• Using percentages directly as ratios without converting to moles.
• Rounding too early; keep at least two decimal places until the final ratio.
• Forgetting to multiply ratios to remove simple fractions (e.g., 1:1.5:2 must become 2:3:4).
• Confusing empirical (simplest ratio) with molecular (actual numbers).