DRAFT DOCUMENT
LAB 19
DETERMINING THE SIMPLEST FORMULA OF A COMPOUND
Chemical formulas provide a great deal of information that serves as the basis for many
computations in chemistry. The simplest formula gives the smallest whole number ratio
of atoms present in a compound. This simplest formula is known as the empirical
formula. A molecular formula represents the actual number of atoms of different
elements in whole number ratios. For example, the empirical formula of tin oxide, SnO2,
tells us that for every atom of tin, Sn, there are two atoms of oxygen, O. In this activity,
you will determine the mass of given elements and use the mole ratio method to establish
the simplest formula of a compound.
OBJECTIVES
When you have completed this activity, you will be able to:
1. Prepare a compound from its elements.
2. Determine the mass of the elements present in the compound.
3. Compute the simplest formula of a compound.
4. Construct a balanced equation for the chemical reaction.
in order to determine quantitative information about a compound.
MATERIALS
ring stand
iron ring
clay triangle
crucible and cover
ceramic pad
crucible tongs
goggles
glass stirring rod
burner
balance, sensitive to 0.01g
deionized water
magnesium ribbon, Mg
wash bottle
PROCEDURE
CAUTION: Wear goggles and apron for all steps.
1. Set up a ring stand, ring, and clay triangle. Place a ceramic pad on the lab table to
be used for placing objects that are cooling.
2. Place a crucible and cover on the clay triangle on the ring stand. Tilt the crucible
cover to leave a small opening.
3. Heat the crucible and cover strongly for 3 minutes to remove any impurities.
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CAUTION: Carefully use the crucible tongs to move and handle the
crucible and cover throughout the lab.
4. Remove the crucible and cover from the ring stand with the crucible tongs,
placing them on the ceramic pad to cool for three minutes.
5. Measure the combined mass of the crucible and cover. Record the mass on the
Data Table. Repeat heating, cooling, and mass measurement until constant mass
is reached. Record measurements. Do not remove crucible and cover from the
balance after the last measurement.
6. Obtain 25 cm of magnesium; cut it into small pieces, and place them in the
crucible. Determine and record the combined mass.
CAUTION: Do not look directly at the magnesium as it burns.
7. Heat the crucible containing Mg with the cover ajar. Continue heating until the
magnesium ignites with the white flame. Move the burner from under the
crucible. Lift the cover with tongs and inspect the contents. If magnesium metal
remains, continue heating until the change is complete.
8. Heat strongly for 5 additional minutes with the lid off. Place the crucible and
cover on the ceramic pad until cool.
9. Use a glass stirring rod to carefully grind the product in the crucible into small
pieces. Rinse the remaining particles from the glass stirring rod into the crucible
with deionized water from a wash bottle.
CAUTION: Do not breathe the vapors.
10. Add just enough deionized water to the crucible to make the product damp. There
should be no excess water in the bottom of the crucible. When water is added,
magnesium nitride will react to form magnesium oxide, ammonia gas, and heat.
Hold the crucible 15 cm away from your face and waft the fumes toward your
nose to check for the odor of ammonia.
11. Replace the crucible in the clay triangle with the lid ajar. Heat gently until the
water begins to boil. Continue to heat until all of the water is boiled off and steam
no longer escapes.
12. Place the crucible on the ceramic pad to cool. Repeat heating, cooling, and mass
measurement until constant mass is reached.
13. Clean the crucible as directed by your teacher.
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DATA TABLE
Mass of crucible and cover (1 heating)
Mass of crucible and cover (2nd heating)
Constant mass of crucible and cover
Mass of crucible, cover, and magnesium
Mass of magnesium
Mass of crucible, cover, and magnesium oxide (1st heating)
Mass of crucible, cover, and magnesium oxide (2nd heating)
Constant mass of crucible, cover, and magnesium oxide
Mass of magnesium oxide
st
ANALYSIS AND CONCLUSIONS
1. Compute the mass of oxygen that reacted with the magnesium.
2. Use the mass of oxygen, from question #1, to calculate the moles of oxygen.
3. Use the mass of magnesium, from the data table, to calculate the moles of
magnesium.
4. The mole ratio of the magnesium to oxygen used in the experiment can be
determined using the following formula. Round the ratio to the nearest whole
number integers such as 1:1, 2:1, 1:2, 3:2, etc.
moles of magnesium : moles of oxygen
5. Using the mole ratio calculated above to write the simplest formula for
magnesium oxide.
6. Using the information from the data table, compute the experimental mass
percentage of Mg in magnesium oxide.
7. Obtain the accepted formula from your teacher. Calculate the percent
composition of magnesium from the accepted value.
8. Calculate the percent error of magnesium using your experimental percent
composition and the accepted value.
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9. Write the balanced equation for the reaction of magnesium with oxygen gas to
form solid magnesium oxide. Use an (s) or (g) as subscripts in the formulas to
denote if the substance is a solid or a gas.
10. Write the equation for the reaction between magnesium and nitrogen gas to form
magnesium nitride, Mg3N2. This reaction occurs due to the exposure of the
magnesium to the nitrogen of the air when Mg is burned. (NOTE: Air is a
mixture of gases of which approximately 78% is nitrogen and 20% is oxygen.)
11. In step 10 of the lab procedures, water was added to magnesium nitride. The
mixture was heated, and an oxide was formed. The detected odor was due to
ammonia gas, NH3, produced when the reaction was heated to boiling. Write a
balanced equation for the reaction which occurred between magnesium nitride
and water forming ammonia as one of the products.
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PRE-LAB: DETERMINING THE SIMPLEST FORMULA OF A COMPOUND
1. Define empirical formula.
2. Which of the following is not an empirical formula?
a. Na2SO4
c. N2H4
b. C6H5Cl
d. Sn3(PO4)4
3. How do you determine the empirical formula of a compound?
4. Determine the empirical formula of the compound with the percent composition
of: 29.1% Na, 40.5% S, and 30.4% O.
5. What is the empirical formula of a compound that is 3.05% C, 0.26% H, and
96.69% I?
a. C2HI7
c. C3H2I11
b. CH2I5
d. CHI3
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