Empirical versus Molecular Formulas

Glucose (C_6H_12O₆) and formaldehyde (CH_2O) share the same ratio of carbon to hydrogen to oxygen: 1 to 2 to 1. Reduce glucose's subscripts by dividing by 6 and you land right back on CH_2O. The empirical formula only records that ratio, stripped down as far as it will go. It throws away size information on purpose, which is exactly why it cannot tell compounds apart on its own.

Molar mass restores the size information. Calculate the mass of one empirical unit, then divide the compound's actual molar mass by that empirical mass. The result is a whole number n, the multiplier, and you multiply every subscript in the empirical formula by n to get the molecular formula. If the division does not come out close to a whole number, something in the data is off, because real molecules can't stack together a fractional number of empirical units.

This sits right before percent composition and combustion analysis, which is where empirical formulas actually get built from lab data: mass percentages become mole ratios, which become the empirical formula, and then a measured molar mass pins down the molecular formula. A particulate drawing of a molecule shows exactly the molecular formula's atom count, not the reduced ratio. If a drawing shows six carbons, six oxygens, and twelve hydrogens clustered into one particle, that particle is glucose, not formaldehyde, even though both would reduce to the same empirical formula.