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5. Simple empirical formula and formula mass from reacting masses (easy start, no moles!)
This page describes and explains, with fully worked out examples, how to work out the empirical formula of a compound. The empirical formula of a compound is defined and explained.
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5. Empirical formula and formula mass from reacting masses (easy start, no moles!)
The EMPIRICAL FORMULA of a compound can be worked out by knowing the exact masses of the elements that combine to form a given mass of a compound.
The empirical formula of a compound is the simplest whole number ratio of atoms present in a compound. (see section 3. for some simpler examples). Here the word 'empirical' means from experimental data.
Do not confuse with molecular formula which depicts the actual total numbers of each atom in a molecule. The molecular formula and empirical formula can be different or the same. They are the same if the molecular formula cannot be simplified on a whole number basis.
molecular formula = empirical formula for sodium sulfate Na2SO4, propane C3H8
where different: butane mol. form. C4H10, emp. form. C2H5, or glucose mol. form. C6H12O6, emp. form CH2O
The following examples illustrate the ideas using numbers more easily appreciated than in real experiments.
In real laboratory experiments only a fraction of a gram or a few grams of elements would be used, and a more 'tricky' mole calculation method is required than shown here (dealt with later for higher students in section 8).
However the examples below show in principal how formulae are worked out from experiments.
Any calculation method must take into account the different relative atomic masses of the elements in order to get to the actual ratio of the atoms in the formula. For example, just because 10g of X combines with 20g of Y, it does not mean that the formula of the compound is XY2 !
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