The Empirical Formula of a Compound

Maxine Wagg

April 27th, 2016

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The Empirical Formula of a Compound

Introduction

        In chemistry, chemical formulas communicate which elements are found in a compound, as well as their proportions. This stands true for the two branches in which chemical formulas can be represented: molecular and empirical formulas. The formula for glucose, C6H12O6, indicates that a single molecule of glucose contains six atoms of carbon, twelve atoms of hydrogen, and six atoms of oxygen. Thus, this formula shows the actual number of atoms of each element in one molecule of the compound. This formula is known as the molecular formula. Molecular formulas often can be as complex as glucose, which contains many atoms, where as simple molecules such as water (H2O) contain a little amount of atoms. To calculate the molecular formula, you must first determine the molar mass of the empirical formula. You do this by adding the molar mass of each element. You then determine the molecular formulas “multiplier”, or the number of times the empirical formula is multiplied to change it to the molecular formula. The multiplier is found by dividing the mass of the molecular formula by the mass of the empirical formula. Lastly, you use this number by multiplying it with the subscripts of each element, which gives you the molecular formula.

However popular mass is used within formulae in chemistry, the relationships of mass reveals little order or sense. The ratio of the masses of the elements in a compound, while constant, does not tell anything about the formula of a compound. For instance, while water always contains the same amount of hydrogen (11.11% by mass) and oxygen (88.89% by mass), these figures tell us nothing about how the formula H2O is obtained, but it does tell us the percentage compositions of each element. Percent composition is the relationship of mass of a specific element is in respect to the remainder of its compound. It is calculated by dividing the mass of a specific element by the mass of the remainder of the compound, then converting the resulting decimal result into a percentage by multiplying it by 100(%). It is shown here: , where x is the mass of a specific element. [pic 1]

Mass must be considered when working with empirical and molecular formulas, as it is used in the standard of chemical quantity, which is not mass but the number of moles (mol) of particles (ions, atoms or molecules). Therefore, when solving for empirical or molecular formulas when given quantities of compounds in a unit of mass such as grams, you must first convert it to moles. The mass of one mole of any substance is numerically equal to its formula weight. Thus, one mole of carbon atoms has a mass equal to 12.011 grams. One mole of glucose (C6H12O6) has a mass equal to 180.18 grams.