Experimental Determination of Chemical Formula

Convert experimental mass measurements to moles using molar masses as conversion factors. Divide each mole value by the smallest mole value to obtain whole-number ratios. If decimals remain, multiply all subscripts by the smallest counting number yielding whole numbers. These ratios become the subscripts in the empirical formula, representing the relative proportions of elements in the compound.

An empirical formula shows the simplest whole-number ratio of atoms in a compound, while a molecular formula shows the actual number of atoms per molecule. The molecular formula is a whole-number multiple of the empirical formula. Both formulas may be identical, or the molecular formula may contain subscripts that are 2, 3, or more times larger than the empirical formula subscripts.

Calculate the empirical formula mass by summing atomic masses of all atoms in the empirical formula. Divide the compound's known molar mass by the empirical formula mass to find the number of empirical formula units per molecule. Multiply each subscript in the empirical formula by this whole number to obtain the molecular formula.

Chemical formulas represent the relative numbers of atoms in a compound, and atoms cannot be divided into fractions. Whole-number subscripts reflect the actual count of each element's atoms in one molecule or formula unit. Fractional subscripts are converted to whole numbers by dividing by the smallest value or multiplying by an appropriate factor to maintain the correct atom ratio.

Yes. Assume a 100-gram sample and convert percent values directly to grams. Use molar masses to convert grams to moles for each element. Divide each mole value by the smallest to obtain whole-number ratios, which become the empirical formula subscripts. This method is convenient because percentages scale to 100, simplifying mass calculations.

Molecular mass can be determined using mass spectrometry, which directly measures the mass of individual molecules. Other techniques include freezing-point depression, boiling-point elevation, and osmotic pressure measurements. These experimental methods provide the molar mass or molecular mass needed to distinguish between empirical and molecular formulas.

Molar mass bridges the gap between experimental mass measurements and the actual number of atoms in a compound. It converts grams of an element into moles, allowing calculation of atom ratios. Additionally, comparing molar mass to empirical formula mass reveals whether the empirical formula is also the molecular formula or if it must be multiplied by a whole number.