Empirical and molecular formula problems look intimidating because they combine percent composition, moles, and ratios into one multi step problem, but each individual step is something students already know how to do.
Finding the empirical formula
Assume 100 grams if given percentages
If a compound is 40% carbon and 60% oxygen, assuming a 100 g sample means those percentages convert directly into 40 g and 60 g.
Convert each mass to moles
Divide each element's mass by its molar mass to get moles of each element in the sample.
Divide every mole value by the smallest one
This gives the simplest whole number ratio between elements, which is exactly what an empirical formula represents.
Round to whole numbers
A result very close to a whole number can be rounded. A result like 1.33 is actually 4/3, so multiplying every ratio by 3 clears it instead of rounding incorrectly.
Going from empirical to molecular formula
The molecular formula is always a whole number multiple of the empirical formula. Divide the compound's actual molar mass (usually given separately in the problem) by the empirical formula's mass to find that whole number multiplier, then multiply every subscript in the empirical formula by it.
Where students actually lose points here
Dividing by the wrong mole value
Every mole value in the problem needs to be divided by the smallest one, not an average or an arbitrary one. Getting this wrong skews the entire ratio.
Rounding a repeating decimal too aggressively
A ratio like 1.5 or 1.33 is a signal to multiply the whole ratio by 2 or 3, not to round straight to a whole number, which would give the wrong formula entirely.
Stopping at the empirical formula when the molecular formula was asked for
If the problem gives an actual molar mass, it's asking for the molecular formula, and skipping that last multiplication step leaves the answer incomplete.
Where percent composition actually goes
This is one of the clearest examples in the whole course of several smaller skills, percent composition, moles, and ratios, combining into one multi step problem.
Common questions
What's the difference between an empirical and a molecular formula?
An empirical formula is the simplest whole number ratio of atoms in a compound. A molecular formula is the actual number of atoms, and it's always a whole number multiple of the empirical formula.
What if the mole ratio doesn't round to a whole number?
A value like 1.5 or 1.33 usually means multiplying every ratio in the problem by 2 or 3 respectively, rather than rounding directly, since those decimals represent simple fractions.
Do you need percent composition, or can you use masses directly?
Either works. If you're given actual masses of each element in a real sample, skip the assume 100 grams step and convert those masses to moles directly.