Determining an Empirical Formula Name ____________________________________

Many ionic compounds, when crystallized from water solution, take up definite proportions of water as an integral part of their crystal structures. This water of crystallization may be driven off by the application of heat.

Since the law of definite composition holds for crystalline hydrates, the number of moles of water of crystallization driven off per mole of the anhydrous compound is some simple, whole number. If the formula of the anhydrous compound is known, you can then determine the formula of the hydrate.


After completing this experiment, you will be able to determine the empirical formula of a hydrate when the formula of the anhydrous compound is known.


balance, centigram crucible and cover tongs ring stand

desiccator sparker iron ring clay triangle

spatula burner and tubing


barium chloride, crystals


Take the necessary safety precautions before beginning this experiment. Wear safety goggles, and an apron. Read all safety cautions and discuss them with your teacher. It is important to use good safety techniques while conducting experiments.

Recording Your Observations

Write the procedure and make Data and Calculations tables in your notebook. After completing each procedure, record your observations in your Data Table.


1. Throughout the experiment handle the crucible and cover with clean crucible tongs only.

2. Place the crucible and cover on the triangle as shown in Figure 1. Position the cover slightly tipped, leaving only a small opening for any gases to escape. Preheat the crucible and cover to redness.



Before lighting the burner, remember to confine loose clothing and long hair. Remember to handle the crucible and cover only with tongs. The crucible and cover are very hot after each heating!

Using the tongs, transfer the crucible and cover to a desiccator. Allow them to cool 5 minutes in the desiccator. Never place a hot crucible on a balance. When cool, determine the mass of the crucible and cover to the nearest 0.01 g. Record this mass in your Data Table.

3. Using a spatula, add approximately 5 g of fine barium chloride hydrate crystals to the crucible. Determine the mass of the covered crucible and crystals to the nearest 0.0001 g. Record this mass in your Data Table.

4. Place the crucible with the barium chloride hydrate on the triangle and again position the cover so there is only a small opening. Too large an opening may allow the hydrate to spatter out of the crucible. Heat the crucible very gently on a low flame to avoid spattering any of the hydrate. Increase the temperature gradually for 2 or 3 minutes. Then heat strongly (red-hot) for at least 5 minutes. Allow the crucible, cover, and contents to cool for 5 minutes in the desiccator and then determine their mass. Enter all masses, properly labeled, in your Data Table.

5. Heat the covered crucible and contents again to redness for 5 minutes. Allow the crucible, cover, and contents to cool in the desiccator and then determine their mass. If the last two mass determinations differ by no more than 0.01 g, you may assume that the water has all been driven off. Otherwise repeat the process, heating to constant mass. Record this mass in your Data Table. The dehydrated compound left in the crucible should be returned to your instructor, since it can be used in the preparation of solutions.

6. At the end of this experiment, clean all apparatus. Check to see that the gas valve is completely shut off before leaving the laboratory and remember to wash your hands.

7. Repeat this experiment with a different hydrate.

8. Calculate the mass of the anhydrous barium chloride, moles of anhydrous barium chloride, mass of water driven off, moles of water driven off, percent water, mole ratio of water to anhydrous salt and empirical formula of barium chloride hydrate.

9. Repeat the calculations for the second salt.

10. Record the results of your calculations on the spreadsheet on the computer. Tomorrow, we will analyze the results.