Formulas — common mistakes

In calculations, do not round your intermediate values too early. Keep at least 3-4 significant figures for mole calculations. Only round at the final step when determining the whole-number ratio. If a ratio is very close to a whole number (e.g., 2.01 or 1.99), you can round it. If it's like 1.5 or 1.33, you must multiply to get a whole number.

You cannot round these to the nearest whole number. You must find the smallest integer to multiply the entire ratio by to make all numbers whole. For 1.5 (which is 3/2), multiply by 2. For 1.25 (which is 5/4), multiply by 4. For 1.33 (which is 4/3), multiply by 3.

Yes, frequently. This happens when the simplest ratio of atoms is also the actual ratio in the molecule. For example, water (H₂O) and methane (CH₄) have molecular formulas that cannot be simplified further, so their empirical and molecular formulas are identical.

Ionic compounds form giant lattices with a repeating pattern of ions, not individual molecules. The formula we write, called a formula unit, must represent the simplest whole-number ratio of ions in that continuous structure. This is, by definition, the empirical formula.

You use the charges of the ions. The overall charge of the formula unit must be zero. For example, to combine Al³⁺ and O²⁻, you need two Al³⁺ ions (total charge +6) and three O²⁻ ions (total charge -6) to balance the charges. This gives the formula Al₂O₃.