Empirical and Ionic Compound Formulae

Chemical formulae are like recipe ratios. They tell us how many of each atom or ion are joined together. In this topic you will learn how to find the empirical formula of a compound and how to write the correct formula of an ionic compound from ion charges.

Empirical formula

Empirical formula = the simplest whole-number ratio of atoms in a compound. It does not have to match the actual number in a molecule (that is the molecular formula).

To find it from composition data:

• Step 1: Assume 100 g (so % becomes grams).
• Step 2: Convert each mass to moles using $A_r$ (relative atomic mass).

$$\text{moles} = \frac{\text{mass}}{A_r}$$

• Step 3: Divide all mole values by the smallest to get a ratio.
• Step 4: If any ratio is not a whole number, multiply all by the same factor to make whole numbers.

Ionic compound formulae

Ionic compounds are made of positive ions (cations) and negative ions (anions). Their formulas show the simplest whole-number ratio of ions in the giant lattice. Total positive charge must equal total negative charge (overall charge = 0).

How to write an ionic formula:

• Write the ion symbols with charges (from the Periodic Table or data): e.g. Mg$^{2+}$, Cl$^{-}$, SO₄$^{2-}$.
• Balance charges so they cancel out.
• Use brackets for polyatomic ions when more than one is needed (e.g. NO₃).

Examples: Mg$^{2+}$ and Cl$^{-}$ → MgCl₂; Al$^{3+}$ and O$^{2-}$ → Al₂O₃; Na$^{+}$ and SO₄$^{2-}$ → Na₂SO₄.

Common misconceptions

• Empirical vs molecular: Glucose has molecular formula C₆H₁₂O₆ but empirical formula CH₂O.
• Ionic formulas are ratios, not molecules. NaCl means 1:1 ions in a lattice.
• Always use brackets for multiple polyatomic ions: Ca(NO₃)₂, not CaNO₃₂.
• Do not leave decimal ratios in empirical formulas; multiply to make whole numbers.