LAB 3:
DETERMINATION OF AN EMPIRICAL FORMULA
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This lab requires a lab report.
You will work with a partner for this lab.
I Introduction
To find the empirical formula of a compound, we need to find the moles of each
element in a given sample of that compound. For example, we might resolve a 40gram sample into its elements and obtain about 23 grams of sodium, 16 grams of
oxygen, and 1.0 gram of hydrogen. Dividing by the molar mass of each element, we
find the sample contains 1 mole of Na, 1 mole of O, and 1 mole of H.
23 g Na
16 g O
1.0 g H
1.0 mol Na
1.0 mol O
1.0 mol H
Taking the ratio of these three elements, we determine the compound’s empirical
formula to be NaOH: we probably have sodium hydroxide.
1.0 mol Na
1.0 mol O
1.0 mol H
Na1O1H1
Usually, however, the numbers are not so simple, not to mention the question that
keeps nagging you: how do we determine the mass of each element in the first place?
In this experiment, you’ll do just that.
Over the next two days, we will use a three-step process to determine the empirical
formula of a hydrated copper salt, CuxClyzH2O (you will recall that, although the
chloride ion always has a charge of –1, the charge of copper ions can vary between +1
and +2).
 In the first step, we will dehydrate the sample, turning the enchanting bluegreen crystals of the hydrate into dehydrated crystals of an ever-so-lovely
shade of tobacco brown (not that you minors would know what tobacco looks
like). From the mass lost, we will be able to determine the moles of water lost;
this will later be useful in determining z in the mystery formula above.
 In the second step, we will reduce the copper ions (which are dissolved in
solution) into copper atoms (which are solid and are not dissolved in solution).
After dissolving the copper chloride in water, we will add aluminum foil to the
solution, causing the aluminum to precipitate out the copper in a fairly
exothermic single replacement reaction, as shown by the following unbalanced
equation:
Al (s) + Cuy+(aq)  Al3+(aq) + Cu(s)
By measuring the mass of copper metal produced, we can determine the moles
of copper present in the sample. This will help us in determining the x in the
mystery formula.
 Finally, after we have subtracted the masses of water and copper, we can
determine the mass and moles of chloride present in the sample. This will
allow us to calculate y in the mystery formula.
At last, with the moles of water, copper, and chloride in the initial sample, we can
find the simplest whole-number ratios of the three and thus deduce the empirical
formula of our mysterious hydrated copper chloride compound.
Equipment and Reagents
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Evaporating dish
100-mL beaker
Bunsen burner
Ring stand with clay triangle
Stir rods, etc.
Filtering apparatus
Analytical balance
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Hydrated copper chloride, CuxClyzH2O(s)
Aluminum foil, Al(s)
Hydrochloric acid, 6 M HCl(aq)
Ethanol, 95% CH3CH2OH(aq)
! Warnings!
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The 6 Molar hydrochloric acid is extremely corrosive. Use goggles and apron(s).
Ethanol is very flammable. Keep away from open flames and sources of heat.
Procedure
This lab is straightforward and excellent results can be obtained if you are careful and
thorough. When transferring substances, be sure to get every last drip, drop, and scrap!
Part 1: Determining the Moles of Water
Set up your lab notebook. Using the same technique we practiced as a class,
determine the mass and moles of water present in about 1 gram of the hydrated
copper chloride. To do this, dry an evaporating dish, record its mass, record the mass
of your hydrated copper chloride sample, and heat until the aforementioned blue-tobrown color change is complete. Use gentle heat and stir frequently with a glass stir
rod. Be careful that none spatters out. It’s easy to tell when the end product is
reached because it will be a uniform brown color. Re-weigh the dish and nowdehydrated copper chloride. Be sure to record all necessary data, procedures, and
observations.
Part 2: Determining the Moles of Copper
Transfer the dehydrated product to a 100-mL beaker and rinse the residue from the
evaporating dish into the beaker with distilled water. Swirl the beaker to dissolve the
crystals; the solution should return to a pleasant blue-green hue as the copper ions
are rehydrated. Add more distilled water if necessary until all of the crystals are
dissolved. Record your observations and procedure.
Measure out about 0.25 g of aluminum foil and tear it into small pieces (smaller
pieces make the ensuing reaction go faster). Drop these pieces into your solution. Sit
back and enjoy the show after recording your observations. The brown substance
forming is the fine particles of copper precipitating out.
Our goal is to measure the mass of solid copper only, so first we must get rid of the
excess aluminum. Once the reaction is finished (how can you tell when all the copper
ions are gone?), carefully add several mL of 6 M HCl (CAREFUL!). If the reaction
between aluminum and HCl stops and there are still visible aluminum chunks, add HCl
and repeat as necessary until all of the aluminum is gone.
At this point, you have pure, solid copper setting in a solution of aluminum chloride.
Filter to isolate the pure copper metal, being sure to weigh your fluted filter paper in
advance. Once the copper has filtered, “wash” the copper with at least two ~5 mL
portions of distilled water. Then rinse with ~10 mL of 95% ethanol. Carefully set your
filter on a paper towel labeled with your name and allow to dry overnight. STOP HERE
ON DAY 1.
The next day, weigh your copper metal and perform the necessary calculations to
determine the empirical formula of your hydrated copper chloride compound.
Data and Calculations (to be recorded on your lab Data Sheet)
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Procedure used
Mass of hydrated sample and dehydrated sample (and any necessary
measurements needed to calculate those)
Observations of dehydration
Observations of reactions
Mass of dry copper product
Mass and moles of water, copper, and chloride in sample
Empirical formula of compound
? Questions to Answer
1) Oops! In Part 2, suppose you accidentally dissolved the copper chloride in 100
mL of water instead of 10 mL. The voice on the left side of your head says you
screwed it up and need to start over; the voice on the right side of your head
says it doesn’t matter, just keep on going. Which of your voices should you
listen to? Why?
2) Why do we use hydrochloric acid to react with the excess aluminum metal
instead of, for example, nitric acid? (Hint: Think back to Lab 2 and the sterling
silver necklace)