3.9
Experimental Determination of Chemical Formula
The elemental makeup of a compound defines its chemical identity, and chemical formulas are the most concise way of representing this elemental makeup…
Chemical compounds are usually described using an empirical- or molecular formula. These formulas provide information about the amount of the different atoms of elements involved. But how are these formulas established?
Experimental analysis, like decomposition of compounds, is used to estimate the relative masses of constituent elements in the compound. These relative masses are then used to calculate the number of moles of each element to determine the formula of a chemical compound.
For example, a sample of a compound is experimentally determined to contain 43.64 grams of phosphorus and 56.36 grams of oxygen. Using molar masses as conversion factors, the relative masses from the experimental data are converted to 1.41 moles for phosphorus and 3.52 moles for oxygen. These mole values, when assigned as provisional subscripts to the elements, yield a pseudo-formula of the compound.
Dividing the mole values by the smallest mole value provides the mole ratios of approximately 2.5 moles oxygen to 1 mole phosphorus, which directly relates to the actual proportion of elements in the compound.
If one of the quotients is still a decimal, then all the numbers are multiplied by the smallest counting number that gives the smallest whole-number ratio of subscripts, generating the empirical formula of P2O5.
The molecular formula of compounds can be determined from their empirical formula and either the molar mass or molecular weight.
For example, the chemical compound with the empirical formula of P2O5 is experimentally measured to have a molar mass of 283.89 g/mol. Its molecular formula is a whole-number multiple of its empirical formula, while its molar mass is a whole-number multiple of its empirical formula mass.
The ratio of molar mass and empirical formula mass yields the number of formula units. Multiplying empirical formula with the number of formula units gives the molecular formula. Hence, from the molecular formula P4O10, the compound is identified as tetraphosphorus decaoxide, or more commonly known by its empirical formula name as phosphorus pentoxide.
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Q1: How do you convert experimental mass data into an empirical 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.
Q2: What is the difference between empirical and molecular formulas?
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.
Q3: How do you determine a molecular formula from an empirical formula?
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.
Q4: Why must chemical formulas contain whole-number subscripts?
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.
Q5: Can you derive an empirical formula from percent composition data?
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.
Q6: What experimental techniques are used to determine molecular mass?
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.
Q7: Why is molar mass essential for determining chemical formulas experimentally?
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.