Be sure to read and understand the material in chapters 2 (page 68), section 2.2, and the mole concept in chapter 3.
Hydrates are crystalline solids that contain a fixed number of water molecules as an integral part of their crystalline structure. The number of water molecules bound per metal ion is often characteristic of that particular metal ion. One of the more common hydrates is copper(II) sulfate pentahydrate, which contains 5 moles of water per 1 mol of copper(II) sulfate, written as CuSO4·5H2O. It is used as a catalytic precursor, fungicide, and as a source of copper in chemical manufacturing processes. Epsom salt is magnesium sulfate heptahydrate, MgSO4·7H2O. Epsom salt is used to reduce inflammation when applied externally.
Many hydrates can be transformed to the anhydrous compound when heated strongly. For example, copper sulfate pentahydrate can be converted into anhydrous copper sulfate. This change can be followed visually. The blue crystalline copper sulfate pentahydrate is converted when heated to a white, powdery, anhydrous salt, according to:
CuSO4·5 H2O ® CuSO4 + 5 H2O
hydrated salt anhydrous salt + water vapor
It is also possible to reverse the above process, as shown in the equation below:
CuSO4 + 5 H2O ® CuSO4·5 H2O
If water is added to the white anhydrous copper sulfate, a blue color is obtained indicating that the blue pentahydrate is regenerated. The property of reversibility can be used to distinguish true hydrates from other compounds that produce water when heated.
Since many hydrates contain water in a stoichiometric quantity, it is possible to determine the molar ratio of water to salt. This is exactly what you will be doing in Part II of today’s experiment. We will begin by heating an accurately weighed sample of the hydrate to drive out the water. The compound formed is now anhydrous. By determining the mass of the anhydrous sample and subtracting this mass from that of the hydrate, we can determine the amount of water in the original substance.
In part I of this experiment, we will look at properties of several crystalline hydrates.
Some anhydrous salts are capable of becoming hydrated on exposure to the moisture in their surroundings. These salts are called hygroscopic and can be used as chemical drying agents or desiccants. Some salts are some excellent desiccants and are able to absorb so much moisture from their surroundings that they can eventually dissolve themselves. These salts are called deliquescent.
Nickel(II) chloride hexahydrate
Cobalt(II) chloride hexahydrate
Copper(II) sulfate pentahydrate
Potassium aluminum sulfate dodecahydrate
Porcelain crucible and lid
Unknown hydrate sample for percentage determination
Part I: Properties of Hydrates
1.Place about 0.1 g of the following compounds in each one test tube:
CuSO4·5 H2O, CoCl2·6 H2O, NiCl2·6 H2O, and KAl(SO4)2·12 H2O.
2. Heat each test tube gently over a Bunsen burner flame and record your observations in your notebook.
3. After the sample has cooled, add a few drops of deionized water. What happens and what can be concluded?
Part II: Formula of a Hydrate
You and your partner will perform two trials of dehydration of a copper (II) sulfate hydrate.
During the course of the experiment, handle the crucible and lid only with crucible tongs as shown here or as demonstrated by your instructor.
Clean two crucibles with soap and water. Rinse the crucibles with distilled water and dry them with a paper towel. Check your crucible for cracks.
Heating your crucible first without the hydrate.
Prepare two set-ups as shown below using a clay triangle on a ring stand. Place each crucible on a clay triangle and heat the crucibles until red hot or for five minutes. Once the heating is complete, place the crucible on a clean wire gauze and let it cool to room temperature. Determine the mass of the crucible and lid to the nearest 0.001g.
Obtain an unknown hydrate from your Instructor. Record the number of the unknown in your notebook. Place about 1 to 2 g (to the nearest 0.001g) of the unknown hydrate into the crucible. Weigh the crucible, lid, and unknown to the precision indicated above.
Place the crucible and its lid onto the clay triangle. Arrange the lid so it is slightly ajar and begin heating very gently for about five minutes. If heated too strongly, some of the sample may decompose. Beware of spattering during the heating process. If spattering occurs put the lid on the crucible and remove the flame. Continue heating for another 10 minutes. Cover the crucible with the lid, cool to room temperature, determine the mass of the anhydrous salt in the crucible, and report it in your notebook.
Repeat the heating-cooling cycle a second time, this time heating the crucible for another 5 minutes. Determine the mass of the anhydrous salt. Continue the heating-cooling cycle until two successive weighing agree within ±0.040 grams.
Calculate the mass percent water in the hydrate based upon the mass of the hydrate and that of the anhydrous salt. Then determine the molar composition of the unknown compound.
Dispose of any used and unused chemicals in the proper waste container located in the hood. Solids are placed in the container labeled “Solid Waste” and all other solutions in the beaker.