Hydrate Lab
Purpose:
• To experimentally confirm the properties of a hydrate.
(reversible dehydration and color change, solubility of residue)
• To identify substances as hydrates based on their known properties.
• To determine the percent water in the given hydrate
• To determine the empirical formula of the given hydrate
Background Information:
Hydrates are solid ionic compounds that contain water chemically bound in the crystal.
They are crystalline compounds that have a specific number of water molecules trapped
within the crystal lattice.
Salts are an example of such compounds. A salt is an ionic compound that forms as a
product of a reaction between an acid and a base. An example of a salt is sodium
chloride, NaCl, which forms from the reaction of sodium hydroxide, a strong base, with
hydrochloric acid, a strong acid.
NaOH + HCl  H2O + NaCl
base + acid  water + salt
Water molecules can become incorporated within the crystal lattice of the salt. These
water molecules associated with the salt are known as waters of hydration. The waters of
hydration can be driven off as water vapor with heating, leaving the salt. The remaining
salt is referred to as anhydrous (without water).
The salt and water are combined in definite molar proportions in a hydrate. The hydrate
formula indicates the mole ratio of water to salt. A dot is used between the formula of
the salt and the formula of water. An example of a hydrate formula is CuSO4·5 H2O ,
copper (II) sulfate pentahydrate. The formula indicates that five moles of water are
associated with every one mole of CuSO4. Another example of a hydrate formula is
AlCl3·6 H2O, aluminum chloride hexahydrate.
Some anhydrous compounds have a strong tendency to absorb water vapor from the air,
thus becoming hydrated compounds. These anhydrous compounds find use as moisture
reducing agents. You may have noticed that containers of such compounds are often
found in bottles containing pills that would decompose if moisture were present. Such
compounds are said to be hygroscopic.
Some of these compounds absorb water to such an extent that they actually dissolve in
the water that they take up. When this is the case, the compounds are said to be
deliquescent. Sodium hydroxide, NaOH, is an example of this type of compound.
On the other hand, some hydrated compounds tend to spontaneously lose their waters of
hydration when they are placed in a dry environment. These compounds are said to be
efflorescent. Sodium sulfate decahydrate, Na2SO4•10H2O is an example of this type of
compound.
The fact that some compounds can be hygroscopic, even undergo deliquescence, and
some can undergo efflorescence, is one good reason to always keep jars of reagents
tightly closed!
The water molecules in a hydrate crystal can be removed by heating. In this dehydration
reaction the crystal structure of the solid will change slightly and the color of the solid
salt may also change. Hydrated cobalt chloride is sometimes used in hygrometers because
it loses water to the atmosphere upon standing. The amount of water lost depends on the
relative humidity of the air. In warm moist air, cobalt chloride is fully hydrated and is
red; at intermediate humidity, cobalt chloride is semi hydrated and is violet; in dry cold
air, cobalt chloride is anhydrous and is blue: Some anhydrous (without water) ionic
compounds will absorb water from the air so strongly that they can be used to dry liquids
or gases. These are referred to as desiccants.
Some compounds evolve water when heated but are not true hydrates. The water is
produced by decomposition of the compound rather than by loss of water of hydration.
Carbohydrates behave this way. Decomposition is irreversible dehydration.
It is possible to experimentally determine the number of waters of hydration present in a
hydrate. The mass of a sample of a hydrate is carefully measured. The hydrate sample
is heated to remove all the waters of hydration and the mass of the remaining anhydrous
salt is measured. The difference in mass between the hydrate and the anhydrous salt
represents the mass of the waters of hydration. The masses of the anhydrous salt and the
water are converted to moles and the mole ratio is determined.
It is very important that all of the waters of hydration be removed. To ascertain that they
are, a technique known as heating to a constant mass is used. The sample is repeatedly
heated, cooled and massed until the sample no longer loses mass when heated.
Several precautions must be observed in order to obtain the required accurate
results.
• The mass measurements must be as accurate as possible.
• Use the same balance throughout the experiment.
• Be certain the balance doors are closed when taring and taking a mass reading.
• Allow samples to cool before taking a mass reading. Balances will not give
correct measurements of warm substances.
• Do not touch vessels, beaker and crucible, with fingers. Oils from the skin can be
left on the vessels and introduce error into the data.
• All of the waters of hydration must be driven off as water vapor.
• Heat to a constant mass. If the mass of a sample has decreased by more than
0.003 grams between successive heatings, the heating process is not complete.
Reheat until the sample loses less than 0.003 grams during a single heating.
• Minimize the risk of waters of hydration reincorporating into the salts during
cooling by covering the compounds during cooling.
• Heating must be controlled to avoid decomposition of the anhydrous salt.
• Heat the crucible dried hydrates as directed in the procedure.
Pre-laboratory questions:
1) State two reasons why you should always use crucible tongs to handle the crucible in
this experiment.
2) How will you determine that your crucible has returned to room temperature after
heating?
3) Why is it important to allow your crucible to cool to room temperature before
weighing it?
4) What do we mean by the phrase “water(s) of hydration”?
5) Describe the difference between a hydrate and its anhydrous form.
6) A student followed the procedure of this experiment, using her unknown hydrate, and
obtained the following data.
Mass of crucible = 10.439 g
Mass of crucible and hydrate = 11.844 g
Mass of crucible and dehydrated sample = 11.213 g
Calculate the mass of hydrate heated =
Calculate the mass of water in the hydrate =
Calculate the percent of water in the hydrate =
Below are three "mini labs" examining characteristics of hydrates. Write each of them up
as a unit keeping their purpose, procedure, observations/data/calculations, and
conclusions together (in your lab notebook).
Mini Lab 1:
Determine whether or not each of these compounds is a true hydrate.
1. Place about 1.0 g of nickel (II) chloride, sucrose, chromium (III) chloride, and
barium chloride in separate dry test tubes. Heat each sample gently. If droplets of
water condense on the cool upper walls of the test tube, this is evidence that the
compound might be a hydrate. Note the texture and color of the residue. Let the
tubes cool.
2. Try to dissolve the residues in a few drops of water, warming gently, if necessary.
A true hydrate will dissolve in water, producing a solution with a color very
similar to that of the original hydrate.
However, if the compound is a carbohydrate, it will give off water on heating and
then char.
Observations:
Conclusion:
Which ones are hydrates? Explain your decision based on your data. Classify every
compound as a hydrate, anhydrate or carbohydrate.
Write the balanced equations for the dehydration process of the true hydrates. Write the
equation for the decomposition of sucrose as well.
Mini Lab 2:
Is the dehydration of cobalt (II) chloride reversible?
Gently heat about 1.0 g of hydrated cobalt (II) chloride hexahydrate in an evaporating
dish until the color change appears to be complete. Dissolve the residue in a few drops of
water. Heat the resulting solution to boiling (CAUTION spattering could occur!) and
carefully boil it to dryness.
Note any color changes. Put the evaporating dish on the lab bench and let it cool.
Observations :
Conclusion:
Explain, based on your observations, why you think the dehydration of cobalt (II)
chloride is either reversible or irreversible. Write appropriate equation showing the
process.
Mini-Lab 3:
What is the % H2O (by mass) in a given hydrate?
What is the empirical formula of the given hydrate?
1. Prepare the Crucible
• Place a crucible and lid on a clay triangle on a ring stand. The lid should be placed
beside the crucible.
• Heat the crucible and lid strongly with a Bunsen burner for 5 minutes.
• Remove the heat. Using crucible tongs, place the lid and the crucible on wire gauze to
cool.
• Allow the crucible to cool to room temperature. (Hold your hand about 1 cm above the
crucible to test.) Handle the hot crucible and cover with tongs. Transport the cooled
crucible and cover in a large, clean, dry beaker.
• Measure and record the combined mass of the cooled crucible and lid. Also record the
mass of the crucible lid separately. This will be useful should you accidentally break the
lid.
2. Prepare the Hydrate
• Obtain about 1 g of your crucible dried assigned hydrate and transfer the sample to the
crucible. After you obtain your sample, quickly replace the cap to the reagent bottle and
tighten securely.
• Measure and record the mass of the crucible, hydrate sample, and crucible lid.
• Using crucible tongs, place the compound containing crucible back on the clay
triangle. Partially cover the opening of the crucible with the lid so that water vapors can
escape.
3. Heat the Hydrate
• Heat the crucible and its contents gently by brushing the bottom of the crucible with
the burner flame for a minimum of 5 min. Continue to heat gently until all evidence of
water is absent. Be careful, if the compound is initially heated too quickly, a significant
amount of water may vaporize suddenly and cause the compound to splatter from the
crucible.
• Increase the flame temperature and heat with a medium flame for 5 min. Further
increase the flame and heat the sample for an additional 10 min. Do not allow the
crucible to turn red.
(Overheating may lead to decomposition of your sample! A decomposed sample will
appear dark gray or black.)
4. Cool and Mass the Hydrate
• Using crucible tongs, remove the crucible and lid from the clay triangle to wire gauze
on the lab bench. Cover the crucible completely with the lid.
• Allow the crucible to cool to room temperature. (Hold your hand about 1 cm above the
crucible to test.)
• Measure and record the mass of the crucible lid and contents.
5. Heat to a Constant Mass
• Reheat crucible with the cover ajar as before, for 10 minutes using a hot flame.
• Allow the crucible to cool to room temperature and measure and record the mass.
• Repeat the heating and cooling procedure until the mass difference between
consecutive heatings is less than 0.003 g.
Observations: Record observations while you perform the lab
Determination of the Waters of Hydration
Data Table
1) Mass of crucible and lid ___________________ g
2) Mass of lid
________________
3) Mass of crucible, lid and hydrate ___________________ g
Heat #1
Heat #2
4) Mass of crucible, lid, and dried sample ________ g _______ g
Heat #3
_______ g
5) Mass of dried sample
________ g
_______ g
_______ g
6) Mass lost by sample (water)
________ g
_______ g
_______ g
Calculations:
Formula of hydrate = ___________________ (Given by teacher)
7) Mass of the hydrate =
______ g
8) Mass of the anhydrous residue = _____ g
9) Mass of the water lost (mass of hydrate - mass of residue) = _____ g
10) Mass percent water in the hydrate =
(Mass of water lost/mass of hydrate) x 100
11) Moles of anhydrous residue = _____ mols
12) Moles of waters of hydration = ______ mols
13) Determine the ratio of moles of waters of hydration to moles of anhydrous salt. Note:
Round the mole ratio to a whole number. Ratios ending in 0.3 or greater should be rounded
UP.
Moles of water/ Moles of anhydrous residue =
(Look at the example given in the next page)
There are 6 moles of water for each mole of NiSO4. The formula for the hydrate is
NiSO4.6H2O.
14) Calculate the theoretical (actual) percent water in the given hydrate (using
the formula that your teacher gives) =
15) Calculate the percent error in the your % water determination =
% Error = [Experimental % water – Theoretical % water ] X 100
Theoretical % water
Conclusion:
The experimentally determined % water in the given hydrate is ______
The experimentally determined empirical formula of the given hydrate is _____
The % error in % water determination is ______
Write the balanced equation to show the process of dehydration.
Post- laboratory questions:
1) If heating cycles are stopped before the sample reaches a constant mass will the
calculated waters of hydration be too high, too low, or not affected? Explain.
2) During heating of the hydrate, loss due to splattering can cause a serious error. If a
significant amount of sample is lost, will the calculated waters of hydration be too high or
too low? Explain.
3) Why is it important to heat the hydrate gently for five minutes and then strongly for 10
minutes?
4) Why did we place the lid over the crucible containing the anhydrous salt as it cooled?
5) If the bunsen burner left some black soot on the bottom of your crucible, how would
this change your answer?
6) What chemistry theme(s) explain the concepts used in the lab?
7) Suggest improvements and extensions for the lab you performed.