Experiment #3 / Unit 3
Determining the Empirical Formula of a Compound (MgxOy)
Introduction:
In a compound, the atoms are present in whole number ratios as given by
the subscripts. The empirical (simplest) formula of the compound provides that
ratio. For example, the empirical formula for potassium chlorate, KClO 3, tells us
that in that compound, for every 1 K atom there is 1 Cl atom and 3 O atoms. The
K:Cl atom ratio is 1:1, and the K:O atom ratio is 1:3. An important job of the
chemist is to find the simplest formulas of any new compounds he or she
discovers. Scientists are constantly finding new genes, proteins, hormones, etc.
which play important roles in human health and whose complex array of atoms
determine their function.
Empirical formulas are determined by finding the mass of each element
present in the compound. From these masses one determines the number of
moles (remember, this tells us "how many") of each element. The mole ratio
must equal the atom ratio in the compound so the mole ratio provides us with
the subscripts for the empirical formula.
To find the mass of each element in a compound, one must carry out a
chemical reaction. Sometimes it is possible to form the compound directly from
the elements. This is called a synthesis reaction. In this experiment, we will form,
or synthesize, magnesium oxide by burning magnesium in the oxygen present in
the air.
magnesium + oxygen  magnesium oxide (synthesis reaction)
By measuring the mass of the magnesium before the reaction occurs, and the
magnesium oxide produced by the reaction, one can calculate the mass of
oxygen that reacted with the magnesium. Using the masses of magnesium and
oxygen in the magnesium oxide we will be able to find the mole ratio between the
elements. The mole ratio will give us the empirical formula of the compound.
Procedure:
Day 1
1. Obtain a piece of magnesium ribbon approximately 40-45 cm in length.
Magnesium will corrode over time and its surface will become coated with
compounds. Remove these compounds by rubbing the surface with steel
wool until it is shiny (please do not shine it on the benchtop because it will
remove the gray surface).
2. Coil the ribbon around a pencil, but make sure that it will be able to lay flat.
Mass the coil and put it in a labeled cup for overnight storage.
Day 2
1. Obtain a clean, dry crucible and lid and record their mass. Place the
magnesium coil in the crucible so that it lies flat on the bottom.
*Caution: do not look directly at the burning magnesium. The intense light
may hurt your eyes.
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2. With the lid off, heat the crucible on a clay triangle set up with a ring stand.
The flame must be a hot one (not necessarily a tall one) and the magnesium
must sit at the top of the pale blue inner cone. When the magnesium ribbon
ignites, you can control the reaction in several ways:
 slow down the reaction by covering the crucible to prevent oxygen from
fueling the reaction.
 slow down the reaction by partially covering the crucible to control the level
of oxygen.
 slow down the reaction by taking away the flame to remove the source of
heat.
 speed up the reaction by blowing gently into the reaction to move some
fresh air (and oxygen) to the reaction.
Note: if the magnesium burns too quickly, some of the compound will be
vaporized (you’ll see white “smoke”) and lost from the crucible. If you have
trouble getting the magnesium to ignite, try keeping the lid on in order to trap
more heat. After heating for a few minutes, remove the lid to give the oxygen
access to the hot magnesium.
3. Your reaction is complete when you no longer observe the white/orange glow
or sparks coming from the magnesium. Heat with the lid on for another minute
to ensure a complete reaction. Now remove the lid to let in more oxygen. If the
reaction is complete, turn off the burner and allow the crucible to cool.
*Caution: the crucible gets very hot- handle with tongs until it has had several
minutes to cool down.
4. When the crucible is cool, take a stirring rod and very gently grind the
contents of the crucible into small particles. Rinse the particles remaining on
the stirring rod and any particles on the sides of the crucible into the bottom
with a minimum amount of water from a dropper. All of the solid must come in
contact with water. Heat gently until all the water has evaporated and the
residue is completely dry. Avoid spattering.
How do I know when the compound is dry?
 Touch it with the glass stirring rod. Is it sticky or powdery?
 The sure way: heat it until you get a constant mass.
5. Allow the crucible to cool and then mass the crucible, its lid, and the
magnesium oxide. Now re-heat and re-mass the crucible.
6. Clean and dry the crucible.
Data:
 mass of dry crucible and lid
 mass of magnesium ribbon
 mass of crucible, lid, and compound(s) after first massing
 additional massings after reheatings
 observations
 error log
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Calculations and Questions:
1. a. Calculate the mass of the magnesium oxide produced.
b. Calculate the mass of oxygen that reacted with the magnesium sample.
c. How many moles of magnesium atoms are there in the compound?
How many moles of oxygen atoms are there in the compound?
d. What is the mole ratio between the two atoms (show in decimal form to
the correct number of sig figs)?
e. Determine how many of each type of atom were present. What is the
ratio between the number of atoms in the compound?
f. Using the ratio calculated, what is the empirical or simplest formula
for magnesium oxide (show any rounding and/or multiplying required)?
2. a. Using the masses of magnesium and magnesium oxide obtained
experimentally, calculate the %Mg by mass in magnesium oxide.
b. Calculate your percent error from the value obtained in 2a if the
accepted value is 60.38% Mg.
3. Write electron configurations for magnesium and oxygen before and after the
reaction. Explain what happened between Mg and O during the reaction.
4. Write equations for the formation of the magnesium and oxygen ions.
5. Write the equation for the reaction of magnesium with oxygen gas to form
magnesium oxide. Since oxygen gas is a diatomic molecule, we use its
molecular formula, O2, in chemical equations. Use (s) or (g) after the
formulas to identify the substance as a solid or a gas.
6. Air is a mixture of gases which contains primarily nitrogen and oxygen gases.
When magnesium burns in air, a small amount of solid magnesium
nitride, Mg3N2, is formed along with the oxide. Noting that nitrogen gas is
diatomic and has the formula N2, write the balanced chemical equation for the
reaction between magnesium and nitrogen gas to form solid magnesium
nitride. Include the physical states of the substances.
7. If water is added to a mixture of magnesium oxide (Mg xOy) and magnesium
nitride (Mg3N2) and heated, the nitride is converted to the oxide . This was
done to make sure we got only magnesium oxide as a product. Ammonia
is produced, leaving the pure magnesium oxide. Write the equation for the
reaction between magnesium nitride and water. This equation does not need
to be balanced. Include the physical states of the substances.
8. Explain the effect on your result for %Mg (high, low, or unchanged) if the
following errors had been made (2 pts each….must explain for credit):
a. a small amount of Mg remained unreacted.
b. some of the product was spilled after heating but before taking the mass.
c. all of the water added was not boiled off.
d. the product contained some Mg3N2. (tricky, be careful)
Lab Report #3.3:
 title page
 abstract
 procedure
 data
 calculations and questions
 results and discussion (minimum of four errors analyzed)
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