The Formula of a Hydrate - Performance Lab 4352 (35 points)
Complete the procedure and the calculations. If your % error is above 10% you will be
required to write a complete lab report. If your % error is between 1-10% you will submit
your pre lab and results. If your % error is between 0-1% you will receive 35/35!
Introduction
Many salts which have been crystallized from a water solution appear to be perfectly dry, yet when heated
yields large quantities of water. The crystals change form, even color sometimes, as water is driven off.
This suggests that water was present as part of the crystal structure. Such compounds are called hydrates.
The number of moles of water present per mole of anhydrous salt is usually some simple number. One
example of a hydrate is cobalt (II) chloride hexahydrate. Its pink crystals look and feel dry, yet each mole
of the hydrate contains six moles of water. Its formula is CoCl2 • 6H2O. The dot between CoCl2 and
6H2O does not mean multiplication.
It means that water molecules are rather loosely attached to the other atoms. The molecular mass of
CoCl2 • 6H2O is: 58.9 + 2(35.45) + 6(18.0) = 237.8 grams/moles.
The hydrate is heated in a crucible. For accurate results, the crucible is handled with tongs because
fingerprints can add weighable mass to the crucible. Porcelain crucibles absorb a bit of weighable moisture
from the air. For this reason, the porcelain crucible and lid is also pre-fired (pre-heating to high
temperature) to constant mass before the pre-weighing. This determines the mass of the completely dry
crucible and lid. At least two firing, cooling, and weighing resulting in exactly the same mass are needed to
confirm constant (completely dry) mass of the crucible and lid and similarly again for the crucible, lid, and
sample residue inside. This is called massing to constant weight.
A desiccator is a container which contains a desiccant to absorb moisture from the air inside, so the air
inside will be completely dry. The desiccant is usually an anhydrous salt, which combines with the water
from the air to form a stable hydrate. Refer to the figure on the attached page.
Purposes
1. Determine the molecular formula of an unknown hydrate.
Pre-lab
1. Read the entire lab
2. Summarize the procedure in your lab notebook
3. Prepare an appropriate data table
4. Answer pre-lab questions.
Pre-lab Questions: Show your work and units!
Maintain significant figures for all calculations!
1. Experimentally, how would you obtain the mass of the anhydrous salt? Explain.
2. Experimentally, how would you obtain the mass of water? Explain.
3. What is “massing to constant weight”?
4. Why was the crucible placed in the desiccator?
5. Why is the hydrate heated at least twice?
6. In an experiment, 2.3754 g of copper (II) sulfate pentahydrate is heated to drive off all the water of
crystallization.
a) Write a balanced equation.
b) Determine the mass of anhydrous salt that remains. Show your calculations.
7. Cobalt (II) chloride is commonly obtained from chemical supply houses as a hydrate with the formula
CoCl2 ∙6H2O. What is the percent water by weight in this hydrate?
8. A hydrate of sodium phosphate, Na3PO4, contains 49.7% water by weight.
a) How many grams of water and how many grams of anhydrous Na3PO4 are in 1000. grams of this sample?
b) In this same 1000.-gram sample, how many moles of water and how many moles of anhydrous
Na3PO4, are present?
c) What is the formula of the hydrate?
Materials:
Crucible and cover
Pipe stem triangle
Bunsen burner
Desiccator with an anhydrous salt
Balance
Iron ring
Tongs
Chemicals:
Unknown hydrate
Procedure
1. Wear safety glasses and apron.
2. Place a clean, dry crucible with cover in a triangle mounted on an iron ring. Leave the cover askew so
that any water in the crucible can easily escape. (See setup figure below)
3. Heat the empty crucible with a high Bunsen burner flame for two or three minutes to clean and dry the
crucible.
4. Do not touch the clean crucible, nor should you weight it while it is still hot. Use tongs to transfer the
crucible to the counter to cool. After about five minutes, use tongs to transport the crucible and its cover to
a balance and find the mass.
5. Put enough of the hydrate crystals in the crucible to fill it one-fourth to one-third full.
(Less than 5 g). Weigh the crucible with cover and content.
Record the masses and write your observations of the hydrate in your data table.
6. Place the crucible, with its cover only slightly askew, on the triangle. See figure on attached page. Begin
heating gently so as to avoid splattering. Gradually increase the flame until the crucible bottom is at most
a dull red.
Maintain this high temperature for five minutes. Use tongs to lift the cover and observe the color of the salt. If
no more blue hydrate is present, use tongs to transfer the hot crucible
and cover to a desiccator. Be careful! The crucible and its cover are VERY HOT!
7. Wait about five minutes, until the crucible is cool enough to touch. Before touching, put your hand close
to the crucible and feel the heat radiating from it. If it is still hot don’t touch it. When cool transfer the
crucible and its cover to the balance and weigh them. Record your data and observations in your data table.
8. To make sure all of the water is driven off heat the crucible and its cover to dull redness again. Carefully
transfer them to a desiccator. When cool, weigh again. If your results do not agree within 0.03 grams from
the previous evaporation, you will need to repeat the heating and cooling process until it does.
9. Add a few drops of water to the crucible and record you observations.
Cleanup: Dispose of the salt in the trash. Wash the crucible well with soap and water.
Return all equipment to the proper bins. Wash your hands.
Calculations: Show your work, units and Significant figures!
Give a brief description of each part.
1. Calculate the number of moles of the anhydrous salt you obtained.
2. Calculate the number of moles of water driven off from your hydrate.
3. Calculate the whole number mole ratio of water to anhydrous salt.
4. Write the empirical formula for your hydrate, _____•XH2O. Check your result with your teacher!
5. Based on your data of the mass of hydrate and the mass of water evaporated, calculate the % composition of
water in the sample of hydrate you used. Calculate the % composition of water in the hydrate based on its
theoretical empirical formula.
6. Calculate the % error of the water mass you measured based on the calculations in question #5 by using
the formula below.
% error 
Theoretical  Experimental
x 100
Theoretical
Analysis Questions
1. Why is the formula of the water in a hydrated compound separated by a dot (e.g. BaCl2•2H2O) instead of
being included in one formula (BaCl2H4O2)? How did this experiment prove this?
2. Suggest reasons why the procedure used in the experiment might not be suitable for all hydrates?
3. Many different drying agents could have been used in your desiccator. Give at least three characteristics
this desiccant should have.
4. Is the process of removing water from a hydrate endothermic or exothermic? What about the reverse
reaction (of adding water to the anhydrous salt)? Explain.
5. Why did the anhydrous copper sulfate change color when water was added?
6. List three possible sources of error in this experiment, and state whether each would have caused a high
or low value for x in the reported formula.
Conclusions
Write a conclusion paragraph.