Lab #11: Composition of a Hydrate
[VIRTUAL MAKE-UP VERSION]
Objective: To determine the formula of a hydrated salt experimentally.
Introduction: A hydrate is an ionic compound (i.e. a salt) that has water molecules incorporated into its crystal
lattice. This “water of hydration” is released as vapor when the hydrate is heated. The remaining solid is known as
the anhydrous salt or anhydrate. The general reaction for heating a hydrate is below. The Δ sign means “heat.”

hydrate 

anhydrate  water
For example, epsom salt (MgSO4) is a hydrated salt that contains seven moles of water per 1 mole of its formula
unit.

MgSO 4  7H 2O 

MgSO 4  7H 2 O
Note: The dot (•) just means that the 7 water molecules are incorporated into the crystal lattice of magnesium
sulfate. It is treated as a “plus” sign for molar mass, not as a multiplication symbol.
In this lab, you will heat a hydrated salt of copper sulfate. By weighing it before and after heating, and doing some
simple mass to mole conversions, you will be able to determine the coefficient for the water of hydration of copper
sulfate.

CuSO 4  xH 2O 

CuSO 4  xH 2 O
blue
white
It is easy to tell when you have driven off the water, since there is a noticeable change in color.
Pre-Lab Questions:
1. Calculate the molar mass of
a) H2O
b) CuSO4
2. Read the entire lab and then construct a Data Table in your lab notebook for the data you will collect.
Safety:

Copper sulfate is a skin and respiratory irritant. It is toxic by ingestion and inhalation. Wear goggles at all
times. Be sure to wash your hands thoroughly at the end of the experiment.

Use caution when working with the Bunsen burner. Tie back hair and secure loose clothing. Do not reach over
the flame, and never leave your lit burner unattended. If you are unable to light the burner within 5 seconds,
shut off the gas, wait 10 seconds, and try again.
Materials:
crucible & cover
ring stand
glass stirring rod
scoopula
watch glass
beaker
crucible tongs
iron ring
wire gauze
bunsen burner and tubing
striker
plastic pipet
clay triangle
hydrated CuSO4
electronic balance
thermal gloves
Procedure:
[GO to
http://group.chem.iastate.edu/Greenbowe/sections/projectfolder/flashfiles/stoichiometry/empirical.html]
1. Measure the mass of the empty crucible and cover and record the mass. You will need this information later.
2. Place the empty crucible on the balance pan and then press the tare/zero button.
Once you see all zeros, carefully add about between 4.80-5.20 grams of the
hydrate and record the exact mass of hydrate.[INITIAL MASS]
3. Set up the ring stand, iron ring, clay triangle, Bunsen burner, and crucible as
shown in the diagram. Position the cover of the crucible slightly tipped, leaving
only a small opening for gases to escape.
4. Heat the crucible and its contents with the Bunsen burner for 10 minutes on
medium to high heat or until all of the blue color has disappeared.
a. Start with a low blue flame and gradually increase the heat applied.
b. During the heating, use a glass stirring rod to break up any large clumps of
hydrate. Try not to pick up any of the solid on the stirring rod.
c. If the edges of the solid appear to be turning brown or grey, remove the heat momentarily and resume
heating with a cooler flame. It is important not to heat too strongly, since above 600oC the anhydrate will
decompose into copper (II) oxide and sulfur dioxide (2CuSO4  2CuO + 2SO3).
5. Using crucible tongs with both hands, remove the crucible from the apparatus. Place it on a square of wire
gauze, so that it is not in contact with the table. Allow the crucible to cool for several minutes.
6. After cooling, immediately find the mass of the crucible + the anhydrate[FINAL MASS] and record it in the
Data Table. The anhydrous compound must be weighed before the compound reabsorbs water from the
atmosphere.
7. When you have finished, scrape the anhydrous salt out of the crucible onto a watch glass. Use a dropper to
place a few drops of water onto the salt to observe the reverse of the dehydrating process. Record your
observations.
8. Clean up and disposal:
a. Put the solid salt into the garbage and wash all equipment with soap and water.
b. Arrange supplies neatly on table and wash your hands.
Calculations:
1. From your experimental data, determine the
a) mass of the anhydrate (CuSO4).
b) mass of water that left the hydrate when you heated it (i.e. the “water of hydration”).
2. According to the formula CuSO4 • xH2O, for every 1 mole of CuSO4 there are x moles of H2O. Using your
masses above, determine the # of moles of anhydrate (CuSO4) and the # of moles of H2O. Then express the
ratio in the form 1:x , where the x is the number of moles of water in the formula CuSO4 • xH2O.
Conclusions:
1. Based on your experiment, what is the formula of copper sulfate hydrate?
2. Obtain the accepted value for the formula, and calculate your percent error for the water of hydration.
% 𝐞𝐫𝐫𝐨𝐫 = (𝐦𝐞𝐚𝐬𝐮𝐫𝐞𝐝 𝐯𝐚𝐥𝐮𝐞 − 𝐚𝐜𝐜𝐞𝐩𝐭𝐞𝐝 𝐯𝐚𝐥𝐮𝐞)/(𝐚𝐜𝐜𝐞𝐩𝐭𝐞𝐝 𝐯𝐚𝐥𝐮𝐞) × 𝟏𝟎𝟎%
3. Discuss three significant sources of error in this experiment. Explain how these errors would affect your
final calculated value of x, the water of hydration.
MPHS Chemistry