Experiment 14 - Hydrate Lab

Experiment 14 - Hydrate Lab

Name ______Date ______


AP Chemistry



When an ionic solid is crystallized from water solution, the crystal which forms often contains chemically-bound water molecules. The number of moles of water per mole of ionic substance is usually an integer. Compounds of this sort are called hydrates. One of the commonly encountered hydrates is magnesium sulfate heptahydrate, otherwise known as Epsom salt:


The water in a hydrate is bound loosely, and so is relatively easily removed by heating. Most hydrates lose their water of hydration at temperatures slightly above 100 oC. Sometimes the water is liberated in stages, with one or more lower hydrates being observed during the heating process. Thus, MgSO4 may also be prepared with 1 mole of H2O per mole of ionic solid. If all the hydrated water is removed, as it will be if the solid is heated sufficiently, the ionic solid is said to be anhydrous (without water).

Given the mass of a sample of the hydrate and the mass of anhydrous salt of known formula obtained on heating, it is easy to find the formula of the hydrate. One simply needs to determine the number of moles of water per mole of anhydrous compound in the hydrate. Once you have your mole values, use your knowledge of finding empirical formulas to determine the formula of the hydrated salt (in our case hydrated copper(II) sulfate).

Finding the percent of water in the hydrate will be done in a similar fashion. From the difference in the mass of the hydrated and the anhydrous magnesium sulfate, one can find the mass of water liberated. Dividing this number by the original mass will give the percent water in the hydrate.

Thus, in this experiment it is our goal to determine the percent of water in an unknown hydrate as well as the formula of the hydrate. This will be done through a knowledge of finding empirical formulas and percent composition.


Along with you introduction and procedure, construct a data table for the lab. Keep in mind you are to complete two trials with the unknown magnesium sulfate hydrate.


  1. Wash and dry a porcelain crucible and lid. Heat both the crucible and cover for about 2-3 minutes to ensure that any water that is adsorbed to the walls of the crucible or cover is driven off. Allow the crucible and cover to cool on a piece of wire gauze.
  1. When the crucible and cover are cool, obtain the mass (at this point you should clean and cool a second crucible for your duplicate measurement. It is completely acceptable to run two trials at the same time. Rarely as scientists are we satisfied with one set of data.)
  1. Put about 2-3 g of the hydrated copper(II) sulfate in the crucible, replace the cover and mass (carry the crucible and cover in a beaker as described in class.)
  1. Place the crucible on the clay triangle on a ring stand. With the cover on the crucible slightly ajar, heat the crucible and its contents, gently at first and then strongly, for approximately 8-10 minutes. The hydrate will probably have a different appearance when the water has been driven off.
  1. Allow the crucible and cover to cool, and obtain a mass.
  1. Place the crucible and cover back onto the clay triangle and heat strongly for another 3 minutes. Allow the system to cool and obtain a second mass. Repeat this step until a constant mass is obtained.
  1. Complete two trials on the unknown copper(II) sulfate hydrate.
  1. Dispose of the solid product in the trash can, and wash the crucible and cover for the next group to use.


Show the following calculations in your lab report for each trial:

  • moles of water
  • moles of anhydrous salt
  • ratio of water/salt (actual value, and then the closest integer)
  • formula of hydrate


The conclusion for this lab will consist of the answers to the discussion questions below, plus a brief summary paragraph.

  1. What is the formula for your hydrate? Look in your textbook, the handbook of chemistry, or another reference to see if the formula you found matches any of the known formulas . How many hydrates of your salt are known? Is your hydrate among them? Give a reference to the source in which you found it.
  1. Determine the experimental % of water in the hydrate.
  1. Examine each of the following sources of error. Determine and explain the direction of the error in the determination of the ratio, water/anhydrous salt. The direction of the error is considered positive when the ratio exceeds the theoretical value and negative when the ratio is less.
  2. Insufficient heating of the hydrated salt.
  3. Absorption of moisture by the dried salt before weighing.
  4. Weighing a warm crucible (which gives a low mass) before introducing the sample.
  5. Weighing a warm crucible after heating the sample.
  6. Overheating to cause decomposition, such as decomposing CuSO4(s) to CuO(s) and SO3(g).