Evaluating the Composition of Compounds
Part I: %Composition
Purpose: To determine the % composition of copper(II) sulfate hydrate. (What % is
water and what % is copper(II) sulfate?)
Procedure:
1.
Obtain a clean dry test tube and beaker. Weigh the test tube in the beaker.
Record the mass.
2.
Using your scoopula, transfer about 2 grams of the copper(II) sulfate
hydrate to the test tube.
3.
Weigh the test tube with the copper(II) sulfate hydrate in the beaker.
Record the mass.
4.
Attach the test tube to a test tube clamp on the ring stand. Heat the
copper(II) sulfate until the entire sample has changed color (all the water
has been driven off.)
5.
Allow the test tube to cool. Weigh the test tube with the copper(II) sulfate
(minus the water) in the beaker.
Data:
Mass of beaker and test tube
____________________
Mass of beaker, test tube, and hydrate
____________________
Mass of beaker, test tube, and copper(II) sulfate
____________________
Calculations:
Part I Questions:
1.
In 50.0 grams of the hydrate, how many grams is copper(II) sulfate?
2.
A reaction required 12.5 grams of copper(II) sulfate. How many grams of the
hydrate would you have to use?
3.
A reaction requires pure anhydrous calcium sulfate. Only an unidentified
hydrate of the compound is available.
We heat 67.5 grams of the hydrate until all the water has been driven off. The
resulting mass of pure calcium sulfate is 53.4 grams. What is the %
composition of the hydrate?
If we need 95.5 grams of calcium sulfate, how much hydrate should we use?
4.
What is the percent composition of water (H2O)? (What % is hydrogen and
what % is oxygen)
5.
How could you determine how many water molecules there are per formula
unit of zinc sulfate with your data from the lab?
Part II: Empirical Formula
In Part II, you will determine the empirical formula manganese chloride. You are using a
transition metal, manganese, which does not always have the same charge in an ionic
compound. Manganese may combine with chlorine to form MnCl, MnCl2, MnCl3, or
MnCl4, depending on reaction conditions.
Purpose: What is the empirical formula for a compound containing manganese and
chlorine?
Prelab Questions:
When 5.49 g of manganese are reacted with an excess of fluorine, 14.99 g of
manganese fluoride are produced.
a.
What is the mass of fluorine in this compound?
b.
Calculate the number of moles of manganese reacted.
c.
Calculate the number of moles of fluorine reacted.
d.
What is empirical formula for the compound (ratio of moles).
Procedure:
Day 1:
1. Measure the mass of a clean, dry beaker labeled with your name. Record the
mass on the report sheet.
2. Add between 2 and 3 grams of manganese metal to the beaker, measure the
mass, and record the mass on the report sheet.
3.
Add 150 mL of 1 M HCl slowly. Record your observations on the Report
Sheet.
4. Place the beaker on the hot plate in the fume hood.
Day 2:
1. Record observations of the beaker contents on the Report Sheet.
2. Measure the mass of the beaker and the manganese chloride. Record the mass
on the report sheet.
3. Set up a ring stand, ring, wire gauze, and burner.
4. Heat the beaker and contents for three minutes and allow it to cool.
5. Measure the mass of the beaker and contents. Record the mass on the report
sheet.
6. Repeat step 4 until a constant mass is obtained.
Data:
Day 1
Mass of beaker
__________
Mass of beaker and manganese
__________
Day 2
Mass of beaker and manganese chloride
__________
Mass of beaker and manganese chloride after heating
__________
__________
Calculations:
a.
Calculate the mass of manganese reacted.
b.
Calculate the mass of manganese chloride produced.
c.
Calculate the mass of chlorine used.
d.
Use the masses of manganese and chlorine to determine the empirical
formula (mole ratio) for the compound produced.
What is the empirical formula?
_______________________
Questions:
1.
If the empirical formula is not a whole number ratio within experimental
uncertainty, analyze your experiment to determine possible causes.
2.
Compare your data to the class data.
2. What is the name of the compound that you formed?
Part III: Empirical Formula of a Hydrate
Purpose: What is the empirical formula for a hydrate of barium chloride?
Prelab Calculations Lab Period 1:
1.
What is the molar mass of Cu(NO3)2 * 3H2O?
Determine the % water in Cu(NO3)2 * 3H2O.
2.
Cobalt sulfate hydrate has an unknown chemical formula. If 23.61g of cobalt
sulfate hydrate loses 10.91g when heated, what is the % by mass water in the
compound?
3.
A student accidentally spilled some of a sample from a crucible while transferring
it from the ring stand to the balance for the final weighing. What effect will this
spill have on the students calculated value for % by mass? (too high or too low).
Explain.
Procedure Day 1:
1.
Place a clean dry crucible (no cover) in a clay triangle mounted on an iron
ring. Heat with a Bunsen burner flame for two or three minutes. Remove
the flame and allow the crucible to cool.
2.
When the crucible is cool enough to touch, transfer it to the balance and
record its mass to the nearest 0.01g.
3.
Using your scoopula, transfer 2.3 to 2.5 grams of the hydrate to the
crucible. Record the combined mass of the crucible and hydrate to the
nearest 0.01g.
4.
Place the crucible with the sample on the triangle and heat them very
gently until any popping or sizzling has stopped. Gradually increase the
flame until the bottom of the crucible is at a dull red. Maintain this
temperature for five minutes.
5.
Use a pair of tongs to transfer the hot crucible carefully to a clean, safe
spot to cool.
6.
Wait about 5 minutes until the crucible is cool enough to touch. Transfer
if to the balance and weigh it.
7.
To make sure that all the water is driven off, heat the crucible to dull
redness again. Carefully cool it and weigh again. If your results do not
agree within 0.03g, consult your instructor.
8.
Carefully transfer your anhydrous sample from the crucible to a clean
250mL beaker. Using a stream of tap water from your wash bottle, rinse
the crucible thoroughly allowing the rinse water to flow into the beaker.
Shake the crucible to transfer all of the rinse water to the beaker. Add
enough water to bring the volume of the solution in the beaker to about
140mL.
9.
Add 2 mL of 6M hydrochloric acid to the solution, cover the beaker with a
clean watch glass and label the beaker with your initials.
10.
Store the beaker in the tray indicated by your instructor until the next lab
period.
Data:
Mass of empty crucuble
_______________
Combined mass of crucible and hydrate
_______________
Mass of crucible and dry salt (1st mass)
_______________
Mass of crucible and dry salt (2nd mass)
_______________
Calculations:
1.
Determine the mass of the original hydrate.
2.
Determine the mass of water removed from the hydrate.
3.
Determine the percent by mass of water in your hydrate sample. (you will use
this in day 2 calculations)
4.
List all the class values for % water. Determine the average for the class.
Prelab Calculations Lab Period 2:
1.
What is the % by mass of Ba in BaSO4?
2.
How many grams of barium will be found in 1.72g or barium sulfate, BaSO4?
3.
A student analyzes the compound copper sulfate hydrate and obtains the
following data:
mass of original hydrate
2.37g
mass of water driven off
0.86g
mass of copper present
0.61g
Calculate the % water in the compound.
Calculate the % copper in the compound.
Calculate the % sulfate in the compound.
What is the mole ratio of Cu : SO4 : H2O?
What is the correct way to write the chemical formula of copper sulfate hydrate?
Procedure Day 2:
1.
Heat your solution containing the barium salt (from part 1) until the
solution just begins to boil.
2.
Rapidly add 10mL of 3M sulfuric acid.
3.
Stir the solution and leave the stirring rod in the beaker (protruding from
the spout). Cover the beaker with the watch glass, stir frequently, use a
small flame, and keep the temperature of the solution just below boiling
for about 10 minutes. Severe bumping and loss of precipitate will occur if
the flame is too hot.
4.
After 10 minutes, remove the heat and allow the precipitate to settle. This
enlarges the BaSO4 that has been produced and makes filtering easier.
5.
Obtain a piece of filter paper. Weigh the paper to the nearest 0.01g. Fold
the paper in half twice. Place it in a filter funnel so that there is one
thickness of paper on one side of the funnel and three on the other side.
Wet the filter paper with a small portion of water to keep it in place.
6.
Using a wash bottle, rinse the watch glass covering the beaker with a small
amount of water, allowing the rinse water to fall into a 250 or 400mL
beaker. Filter the rest of the beaker contents, never letting the funnel get
more than two-thirds full. Discard the clear filtrate occasionally, since
you only intend to collect the precipitate. Using a small stream of water
from your wash bottle, carefully remove any precipitate adhering to the
walls or bottom of the beaker. If the filtrate is cloudy, place a clean
beaker below the funnel, and re-filter the filtrate.
7.
Discard any clear filtrate from the beaker. Replace the empty beaker
under the funnel and rinse the precipitate with two 10mL portions of tap
water. Then rinse with one or two 10mL portions of acetone. (Caution:
Do not bring acetone to your lab desk until all of the students have
finished using their burner. Acetone is flammable and should never be
used by an open flame.)
8.
Place the filter paper with precipitate on a watch glass. Label your sample
and store it carefully on the tray indicated by your teacher until it is dry
the next day.
9.
Mass the dry barium sulfate precipitate on filter paper the next day.
Data:
Mass of filter paper
_______________
Mass of filter paper and dry BaSO4 _______________
Calculations:
1.
Determine the mass of barium sulfate precipitate.
2.
Calculate the mass of barium that is contained within the barium sulfate
precipitate. (Use the % by mass of Ba in BaSO4)
3.
Use your answer to question 2 and the mass of the original hydrate from day 1 to
determine the % barium in your original sample. (Remember all the barium in the
hydrate was converted to barium sulfate and trapped in that new form.)
4.
Calculate the % chlorine in the hydrate by subtraction.
5.
Determine the mole ratio Ba ; Cl ; H2O.
6.
Report the empirical formula of the compound.