Empirical Formula
Introduction
For a generic chemical formula AxBy, x gives the number of
atoms of element A that combine with y atoms of element
B. The empirical formula for a compound is the smallest
whole number ratio of x to y. If the combining masses of
elements in a compound are known, the empirical formula
is found by converting the mass of each element into its
corresponding number of moles, and then dividing through
by the number of moles which has the smallest value.
equipment
From the storeroom, a kit containing:
• one large test tube.
• one straight, one bent glass tube.
• one two-holed rubber stopper.
• two quick-connect plastic couplings.
• one black rubber tube.
From your locker:
• one 10 cm, one 30 cm amber rubber tube
• 100 ml, 150 ml beaker
From the large drawer:
• one Bunsen burner.
• one flame striker.
• one adjustable clamp (not vinyl coated).
• one ring stand.
Many analytical techniques have been developed to
determine the mass distribution of elements in a compound.
In this experiment, the simple process of reduction is
used.
One definition of reduction is the removal of oxygen from
a compound. Oxygen can be removed from a number of
metal oxides using carbon or hydrogen with heat. For iron
oxide, Fe2O3, for example, the reactions are:
experiment
Make sure the test tube is dry. Use a piece of folded paper
towel to dry the tube. Do not use compressed air.
2 Fe2O3 +3 C → 4 Fe + 3 CO2
Fe2O3 +3 H2 → 2 Fe + 3 H2O
The draft shield should be on the balance whenever you
record a weight.
The setup for the balance is:
Some metal oxides are reduced by hot CH4, methane, the
prime component in natural gas. The reaction of iron oxide
with methane is:
4 Fe2O3 + 3 CH4 → 8 Fe + 3 CO2 + 6 H2O
Copper oxide will be reduced with methane in this
experiment. The mass of the copper oxide will first be
measured, and then, after the oxygen is removed, the mass
of the remaining copper will be measured. The difference
in the mass before and after the oxygen is removed will
give the mass of oxygen in the compound. The number of
moles of both copper and oxygen will be calculated from
the masses, and the simplest whole number ratio of moles
will give the empirical formula.
Two sources of error to watch for: 1) All of the water
produced in the reduction must be removed before the
final weighing, and 2) The copper must be allowed to
cool after the reduction before coming in contact with air,
otherwise it might re-oxidize. A physical problem to avoid:
If the test tube is not kept close to level during the heating,
condensed water can flow down the tube and cause it to
break if it hits hot glass.
100 ml
beaker
ON/OFF
MENU
TARE
You must "trick" the balance to get weighings to the
nearest 0.001 g. Here's how:
1
Put your 150 ml beaker on the balance, and zero the
weight. Remove the 150 ml beaker. Now put the 100 ml
beaker with the test tube in it on the balance, and record
the weight. Add the copper oxide until the weight has
increased by about one gram. Use your micro spatula to
add the copper oxide to the test tube. Record the weight
to the nearest 0.001 g. You should have between 0.8 and
1.2 grams of copper oxide in the test tube. After you have
heated the copper oxide, repeat the procedure to weigh
the copper. Put the 150 ml beaker on the balance, zero the
balance, then remove the beaker. Put the 100 ml beaker
with the test tube on the balance, and record the weight.
See the discussion in the balance handout to see the
reason for this procedure.
Pay particular attention to the angle of the test tube and
the placement of the clamp holding the test tube. Do not
use vinyl coated clamps.
Refer to the Rubber and Glass handout for instructions on
making connections. That discussion will also help you
understand why two different kinds of rubber hoses are
used.
Use significant digits properly when you record data and
when you do the calculations.
Connect the test tube into the apparatus. Call the instructor
to examine the apparatus and initial the data table before
you turn on the gas.
If all the balances are in use, begin assembling the
apparatus until a balance is free.
Follow the instructions in the Bunsen Burner handout to
get a hot flame. Ask the instructor for help if needed.
Tube is s
li
When the gas is first turned on, it will take a few seconds
before it gets through the test tube to the Bunsen burner,
so initial efforts to light the flame might fail. Move the
flame over the area of the test tube containing the copper
oxide. Move slowly and consistently until the black oxide
has turned into the red metal. If this has not happened in
15 minutes, call the instructor over for advice. Heat the
upper portion of the tube on occasion to allow the water
condensing to vaporize.
ghtly tilt
ed
10 cm amber tubing
30 cm amber tubing
After the copper oxide has been reduced, set the burner
down away from the tube; turn the air vent down (see the
Bunsen Burner Use discussion ) and reduce the gas flow
to get a very small, gentle flame.
black tubing from kit
To
Gas
Spigot.
Red
Bunsen
Burner
Tubing
When the tube containing the copper has cooled to room
temperature, turn the gas off and remove the rubber
stopper, being careful that no condensed water rolls down
the tube towards the copper. If there is any water on the
upper portion of the tube, wipe it off with a paper towel,
then weigh the tube with the copper, using the method
described above, and record the weight in the data table.
base facing
towards you
Pay careful attention to the diagram. Your setup should
match exactly. Points will be taken off of your grade for
this experiment if your setup has substantial deviations
from the diagram. Carefully read the instructions in the
next column before you start.
Return all items to their place of origin and clean up your
work space.
You may put the copper in a baggie and keep it as a
trophy.
2
Instructor's initials, setup OK _______
DATA
1.
mass of empty test tube
g
2.
mass of test tube with copper oxide
g
3.
mass of test tube with copper
g
CALCULATIONS
4.
mass of copper oxide (data 2 – data 1)
g
5.
mass of copper (data 3 – data 1)
g
6.
mass of oxygen (data 2 – data 3)
g
7.
moles of copper (show work)
8.
moles of oxygen atoms (show work)
9.
whole number ratio, mol Cu to mol O
10.
Formula of copper oxide (experimental)
mol
mol
3
Name_________________________________________ Grade___________ Date ___________
questions
1. The substance used was copper (II) oxide. Use the name of the substance to write its formula:
2. Look at the calculated number of moles of copper and of oxygen from your data. Subtract the smaller number from
the larger number. (Use significant digits properly.) Divide this difference by the smaller number. Multiply by 100.
Now write the sentence: “The number of moles of _______ (the larger number of moles substance, either copper
or oxygen) was ­___ % larger than a 1 to 1 ratio of moles requires.” (Note that you could alternately say that the
smaller number substance was some percent smaller than a 1 to 1 ratio requires.)
3. For each of the following possible experimental errors, state what the effect would be on data entries 1, 2, and/or
3. That is, would the errors make the data entries larger or smaller? Also state whether the error would cause the
calculated moles of copper and the calculated moles of oxygen to be too high or too low.
• Incomplete reduction of the oxide:
• Re-oxidation of the copper after the reduction because of opening the test tube while still hot:
• Incomplete removal of the water after the reduction:
• Water in the test tube before the first weighing:
• Can you think of any other problems that might occur?
Did any of these problems occur during your experiment?
4. Write out the balanced equations for the reduction of copper oxide with C, H2, and CH4.
(use the reactions shown on page 1 for iron(III) oxide as a template.)
4