Chemical Formulas

Overview: This section provides a review of chemical formulas and the information that is available from the different, but equivalent representations (formulas) of a molecule/compound.


New terms:

First what is a chemical formula? They have been used numerous times already in this tutorial but so far, they have not been defined.

Chemical Formula: A representation of the chemical composition of a substance. This can be either a molecular or empirical formula.

Basically a chemical formula gives scientists a variety of information about a particular compound.
Let's look at iso-octane. Iso-octane is the component of gasoline that burns the smoothest. Iso-octane can be written in numerous ways.

Percent Composition by Mass (Weight)

Experimentally, the mass percentage of a compound is obtained by means of combustion analysis or other types of elemental analysis. We can use mass percentages to determine empirical formulas, but not molecular formulas.


Please note that the weight fraction is equal to the mass fraction so that the percent weight also equals the percent mass. This is because weight and mass are proportional to each other.

Example 1.

What is the percent elemental composition by weight (mass) of each iso-octane molecule, C8H18?

What do we need to do first? We need to compute the mass of each of the elements, C and H, which can be done with the equation above and the information from the molecular formula. We know that in each molecule of iso-octane, there are 8 atoms of carbon and 18 atoms of hydrogen. This tells us that in each mole of iso-octane, there are 8 moles of C and 18 moles of H. By now, we can easily go from moles to mass by using molar masses. Unfortunately, we do not know how many moles of iso-octane there are. So what can we do? Remember what it is that we are looking for in the problem. We want the percent mass. The percent mass remains the same for iso-octane whether there is 1 molecule or 5000 moles of iso-octane. So to make things easy, we'll assume that there is one mole of iso-octane. First we will find the mass of each element, then find the percent composition by dividing the molar mass of each element by the total mass of the molecule?

g H

This means that iso-octane is 84.12% carbon. The percent hydrogen must be 100% - 84.12% since the total elemental percentages must add up to 100. Therefore hydrogen accounts for 15.88% of the molecular mass. Notice that we could have found the % H first and then subtracted to solve for the percent carbon. Try solving for the percent hydrogen first and then the percent carbon to verify that you get the same answer.
Once you have the percent elemental compositions, you can derive the empirical formula.

Example 2:

What is the empirical formula of a compound containing 84.12% carbon and 15.88% hydrogen by mass?

To find the empirical formula, we need to find the number of moles of each element. We are able to go from mass to moles by the molecular mass but we do not have an initial mass. We can assume a mass of the sample because, once again, the percent mass is the same no matter what the mass of the sample is. To make things easy, we will assume 100 g of compound. Therefore, there are 84.12 g (this is 84.12% of 100 g) of carbon and 15.88 g (15.88% of 100 g) of hydrogen.


Now determine the smallest mole ratio. Do this by taking the number of moles of the element over the number of moles of the element that has the least number of moles in the compound. In this case, carbon has the least number of moles.


Now we need to find the smallest integer ratio. Carbon is already in integer form but what factor multiplied by 2.25 gives an integer. The factor will be 4. This means we need to multiply each of these mole ratios by 4 to get the empirical formula.


C: 1 x 4 = 4

H: 2.25 x 4 = 9

So the empirical formula is C4H9.

From the information given in this example, can we determine the molecular formula?

Example 3:

Given that the molecular mass for a compound with an empirical formula of C4H9 is approximately 114 g/mole, what is the molecular formula of the compound?

We know the molecular formula is a multiple of the empirical formula: (C4H9)x

The molecular mass will be the sum of the individual molecular masses


Therefore the molecular formula is (C4H9)2 = C8H18, which is iso-octane.

Combustion Train

When a hydrocarbon, a compound that contains only carbon and hydrogen, is burned in oxygen, it yields CO2 and H2O. This is done using a combustion train. In the combustion train, the sample is weighed and then heated in the presence of oxygen. The products, carbon dioxide and water, are allowed to pass into another chamber where the water is absorbed. A dessicant such as magnesium perchlorate absorbs the water. The carbon dioxide continues into the final chamber where it is absorbed by NaOH and CaCl2. Once the original hydrocarbon is completely burned, the absorbers are weighed. The difference in masses is due to the CO2 and H2O (the H and C come only from the hydrocarbon). From this information, one can determine the empirical formula.

Example 4:

A certain hydrocarbon, when burned completely in oxygen, yields 3.38 g of CO2 and 1.384 g of H2O and no other products. What is the empirical formula of the compound?

Need help with a strategy?


Advanced Applications: A trip back in time. Advanced Applications circa 1890.


Now you should have a better understanding of chemical formulas and the different ways chemists represent a compound. Also, you should be able to determine percent elemental compositions and know how to calculate empirical formulas from the percent elemental composition.

Practice Quiz: Chemical Formulas

Note: You will need a pencil, scratch paper, calculator, periodic table and equation sheet to work the practice quiz. Quizzes are timed (approximately 4 minutes per question).

We suggest that you print out the periodic table and the constants/equations page before starting the quizzes.

Open Periodic Table in a separate browser window.

Open Equations/Constants Page in a separate browser window.

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© Washington University in St. Louis, 2005.