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Chem 121 Lab
Clark College
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Experiment 5: Percent Composition of a Hydrate
The concept of % composition is often used to determine how many grams of an element might be
produced when a compound is decomposed, or how many grams of an element is necessary to
produce a given quantity of compound (in grams). By using the mole relationship to get
mass/number conversion factors, it is possible to determine the number of moles of water present
per mole of copper sulfate.
A. Definition
Hydrates are compounds having water incorporated within the crystal structure in specific whole
number ratios.
The bonds holding the water are weak bonds and are easily broken when heated.
Δ
BaCl 2 •2 H2 O ⎯⎯→ BaCl 2 + 2 H2 O
hydrate
solid + water
Today you will determine both the percent of water present in a hydrated copper (II) sulfate
(CuSO4 • n H2O) and “n,” the number of moles of water molecules present per mole of CuSO4.
B. Calculating Percent
€ Composition
Percent means parts per 100.
⎛ Part ⎞
% = ⎜
⎟ x 100
⎝ Total ⎠
To calculate the percent composition (by mass) of each element in a compound:
1. Determine the compound’s molar mass (or formula mass).
2. Divide the total mass
€ of each element by the compound’s total mass.
3. Multiply that fractional mass by 100.
Example I: Find the mass percent of oxygen in one mole of H3PO4.
Solution:
⎛ 1.0g ⎞
Mass H : 3 mole H × ⎜
⎟ = 3.0g H
⎝ 1 mol H ⎠
⎛ 31.0g ⎞
Mass P : 1 mole P × ⎜
⎟ = 31.0g P
⎝ 1 mol P ⎠
⎛ 16.0g ⎞
Mass O : 4 mole O × ⎜
⎟ = 64.0g O
⎝ 1 mol O ⎠
⎛ Mass O ⎞
%O = ⎜
⎟ × 100%
⎝ Total mass ⎠
⎛ 64.0g ⎞
= ⎜
⎟ × 100%
⎝ 98.0g ⎠
= 65.3%
Molar mass of compound = 98.0g
€
Percent Composition of a Hydrate
€
Revised S10 SB
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The percent composition values can be used to calculate the number of grams of an element in a
given mass of compound.
Example II: Calculate the number of grams of oxygen that would be present in a 4.00 gram
sample of phosphoric acid, H3PO4.
Solution:
From the previous example, 65.3% of phosphoric acid (H3PO4) is oxygen by mass. This means that
there are 65.3 grams of O in 100. grams of the compound phosphoric acid.
⎛
65.3g O
Therefore: 4.00g H3PO 4 × ⎜
100.0g
H 3PO 4
⎝
⎞
⎟ = 2.61g O
⎠
€
Finding the mass percent water in a hydrate is very similar. The only difference is the replacement
of the element mass by the mass of water (H2O).
Example III: The formula of a hydrate of barium chloride is BaCl2 • 2H2O. Find the mass
percent water in this hydrate.
Solution:
Total Mass of BaCl 2 • 2H 2O =
⎛ 137.3g ⎞
⎛ 35.5g ⎞ ⎡
⎛ 1.0g ⎞
⎛ 16.0g ⎞⎤
1 mol Ba⎜
⎟ + 2 mol Cl⎜
⎟ + 2⎢2 mol H⎜
⎟ + 1 mol O⎜
⎟⎥ = 244.3g BaCl 2 • 2H 2O
⎝ 1 mol Ba ⎠
⎝ 1 mol Cl ⎠ ⎣
⎝ 1 mol H ⎠
⎝ 1 mol O ⎠⎦
⎛ 1.0g ⎞
⎛ 16.0g ⎞
Mass of Water = 2 mol H⎜
⎟ + 1 mol O⎜
⎟ = 18.0g H 2O
⎝ 1 mol H ⎠
⎝ 1 mol O ⎠
⎛
⎞
⎛ 2 × mass H O ⎞
2 × 18.0g H 2O
2
%Water in Hydrate = ⎜
⎟ × 100%= 14.7%
⎟ × 100%= ⎜
⎝ Total mass ⎠
⎝ 244.3g BaCl 2 • 2H 2O ⎠
€
C. Experimental Considerations
In this experiment, you will be heating a hydrate of copper (II) sulfate (CuSO4 • nH2O) to drive off
the water. Masses are taken before heating to determine the mass of the original sample (the
hydrate) and after heating to determine the mass of copper (II) sulfate (CuSO4) remaining. The
difference between these two masses is equal to the mass of the water lost. Heating time and
temperature are critically important for this experiment. If not enough heat is applied, some water
will remain attached to the copper sulfate producing a low calculated mass percent water for the
hydrate. If too much heat is applied, the anhydrous copper (II) sulfate (CuSO4), which has a
grayish white color, can be decomposed to copper (II) sulfide, a black colored compound.
Percent Composition of a Hydrate
Revised S10 SB
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D. Experimental Procedure
After you have washed the crucible and lid, use only tongs (not your hands) to handle them
1. Place a clean, empty crucible with lid in a clay triangle on a ring stand (as shown in diagram).
Tilt the lid so that it is slightly ajar, then heat strongly (bottom of crucible should turn red) for
about 3 minutes. Turn off the Bunsen burner, and use tongs to close the lid so that water from
the air does not get inside the dry crucible. Allow the crucible and lid to cool (this should take 5
minutes).
2. Using tongs, transfer the crucible and lid (still closed) to a wire gauze and carry them to the
balance. It’s okay if you need to remove the lid momentarily to transfer the crucible & lid separately to
the wire gauze. Mass the crucible and lid (together) carefully to the nearest 0.01g and record the
mass in the Data Table provided.
3. Add approximately 2.5- 3.5g of the copper (II) sulfate hydrate to the crucible and mass the
crucible with the hydrate and the lid again to the nearest 0.01g. Record the mass in the Data
Table.
4. With the lid slightly ajar, heat the crucible gently (crucible should NOT glow red; use the top of
the outer flame, not the inner flame) for about 12 minutes. Turn off the Bunsen burner and use
tongs to close the lid. Allow the crucible with sample and closed lid to cool for 5 minutes, then
mass to the nearest 0.01g, recording the mass in the Data Table. During the time you are
cooling the sample and crucible you can be working on the Exercise questions at the end of this
experiment.
5. Reheat the sample for 2-3 minutes (with lid slightly ajar), cool (with lid closed) and mass again.
This should be repeated until the successive masses are constant within 0.03g. When the masses
are constant, record the lowest mass as the final mass.
Waste Disposal: Place compound remaining in your crucible in the jar labeled Copper Sulfate
Collection. Note the color of the solid closest to the crucible bottom: ____________
Percent Composition of a Hydrate
Revised S10 SB
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DATA TABLE.
1.
Mass of empty, dry crucible and lid
(after heating & cooling)
2.
Mass of lid, crucible and hydrate
(before heating the sample)
3.
Mass of hydrate (original sample)
4.
Mass of crucible, lid, and dehydrated sample:
after first heating & cooling
after second heating & cooling
after third heating & cooling (if needed)
final mass
5.
Mass of the water driven off
6.
Mass percent water in the hydrate
(show calculations below)
E. Analysis
1. How many moles of copper (II) sulfate (CuSO4) and how many moles of water did you have in
your original sample? Hint: Convert each mass using molar mass (g/mol).
Percent Composition of a Hydrate
Revised S10 SB
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2. Examine the formula for the hydrate: CuSO4 • n H2O. Notice that “n” is “the molar ratio of
water to copper sulfate.” Find the numerical value for n in this sample (use your numbers
from part a above). Hint: what does “molar ratio of water to copper sulfate” mean? Write it as a
fraction – then solve it, since n = “the molar ratio of water to copper sulfate.”
3. The actual mass percent of water in the hydrated copper (II) sulfate compound should have
been 36.1%. Compare this value to the experimental percentage you obtained. Hint: how do you
quantitatively compare an experimental value to an “actual” value?
4. In the experiment involving hydrated copper sulfate, overheating causes a high calculated
percent value for water. Why is the high reading obtained? Hint: Overheating causes copper (II)
sulfate, CuSO4 to turn into copper (II) sulfide, CuS. What is lost from the CuSO4 in this process?
Where does it go?
5. What experimental evidence would you have to indicate you inadvertently overheated the
hydrated copper sulfate compound? Hint: Re-read the “Experimental Considerations” section.
Percent Composition of a Hydrate
Revised S10 SB
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F. Exercise Questions
6. Determine the mass percent of each element present in CaCO3.
7. Determine the number of grams of Na present in 4.0g of NaOH.
8. Calculate the mass percent water in MgSO4 • 2 H2O.
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Revised S10 SB
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