An empirical formula gives the simplest whole-number ratio of atoms of each element in a compound. A molecular formula gives the actual number of atoms of each element in one molecule and is a whole-number multiple of the empirical formula.
This distinction is part of IB Chemistry Structure 1.4 for both SL and HL.
| Feature | Empirical formula | Molecular formula |
|---|---|---|
| Information shown | Simplest ratio of atoms | Actual number of each type of atom |
| Example for glucose |
| Required data | Elemental composition | Empirical formula and molar mass |
| Applies to | Molecular and ionic compounds | Discrete molecules only |
From composition data, convert each element's mass or percentage to moles, divide by the smallest amount, then scale to whole-number ratios. For a compound containing 40.0% carbon, 6.7% hydrogen, and 53.3% oxygen, assume a 100 g sample. The mole amounts are mol C, mol H, and mol O. Dividing by 3.33 gives approximately , so the empirical formula is . If a ratio contains 1.5, multiply every ratio by 2; do not round it directly to 2.
The molecular formula is an integer multiple of the empirical formula:
The multiplier is calculated using:
For example, suppose a compound has the empirical formula and a molar mass of .
A common misconception is that the empirical formula always represents a real molecule. It does not: is only the simplest ratio for glucose, not its actual molecular formula. Ionic compounds such as form lattices rather than discrete molecules, so their formulas represent empirical ratios.
For a determine question, show the mole ratio, simplify it to obtain the empirical formula, calculate , and multiply every subscript. Do not multiply only one element, and check that is a whole number.
