Revise Formulas for Cambridge 9701 (2.3)

It's crucial to distinguish between two types of formulas. The molecular formula gives the actual number of atoms of each element in one molecule of the compound. It is the 'true' formula. The empirical formula, on the other hand, provides the simplest whole-number ratio of atoms of each element in the compound. Sometimes they are the same, but often they are different.

$\text{Molecular Formula} = (\text{Empirical Formula})_n \\ \text{where } n = \frac{\text{Molar Mass}}{\text{Empirical Formula Mass}}$

We can determine the empirical formula of a new substance from experimental data, typically percentage composition by mass or the masses of elements in a sample. The process involves converting mass to moles and then finding the simplest whole-number ratio of moles.

If you know the empirical formula and the relative molecular mass ($M_r$) of a compound, you can determine its molecular formula. The $M_r$ is usually found experimentally using techniques like mass spectrometry. By comparing the $M_r$ with the mass of the empirical formula, you can find the integer 'n' which scales the empirical formula up to the molecular formula.

For covalent substances that exist as simple molecules (like H₂O or C₄H₈), the molecular formula is a very useful description. However, ionic compounds like sodium chloride (NaCl) or magnesium oxide (MgO) do not form discrete molecules. Instead, they form a vast, three-dimensional structure called a giant ionic lattice. For these substances, it is impossible to define a 'molecule'. Therefore, we use a formula unit, which represents the simplest whole-number ratio of ions in the lattice. The formula unit for an ionic compound is, by definition, its empirical formula.