TITLE: Moles of Iron and Copper
PURPOSE:
To determine the number of moles of copper produced in the reaction of iron and copper (II) chloride.
To determine the number of moles of iron used up in the reaction of iron and copper (II) chloride.
To determine the ratio of moles of iron used to moles of copper produced.
BACKGROUND INFORMATION:
The mole is a convenient unit for analyzing
chemical reactions. The mole is equal to 6.02 x 10 23
particles, or Avogadro’s number of particles. More
importantly, however, the mass of a mole of any
components or element is the mass in grams that
corresponds to the chemical formula. Simply stated, the
atomic mass of copper is 63.5 amu, which means that the
PROCEDURE:
1. Put on your lab apron, safety goggles and plastic
gloves.
2. Find the mass of a clean, empty, dry 250-mL
beaker. Record the mass.
3. Add approximately 8 g of copper (II) chloride
crystals to the beaker. Find the mass and record it
on the Report Sheet.
4. Add 50-mL of distilled water to the beaker. Swirl
the beaker around to dissolve all of the copper (II)
chloride crystals. Record your observations of
the liquid.
5. Obtain two clean, dry nails. If the nails are not
clean, use a piece of sandpaper to make the
surface of the nail shiny. Find the mass of the
nails together and record it on the Report Sheet.
6. Place the nails into the copper (II) chloride
solution. Leave them undisturbed for
approximately 20 minutes. During this time, you
should see the formation of copper in the beaker.
At the same time, some of the iron will be used
up.
7. Use the tongs to carefully pick up the nails, one at
a time. Keeping the nails over the large beaker
and try to get all of the copper off of the nail and
into the beaker (NOT DOWN THE DRAIN)!!
Use distilled water in a wash bottle to rinse the
copper from the nails BEFORE removing them
mass of one mole of copper is equal to 63.5 g. Likewise,
the molecular mass of water is equal to 18.0 amu, and the
mass of one mole of water molecules is equal to 18.0 g.
The mole is the common language in chemical
reactions. In this experiment, you will observe the reaction
of iron nails with a solution of copper (II) chloride and
determine the number of moles of each in the reaction.
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completely from the beaker. If necessary, use a
stirring rod to scrape any excess copper from the
nails. Set the nails aside to dry on a paper towel.
After the nails are completely dry, find the mass
of the nails and record it on the Report Sheet.
Decant means to pour off only the liquid from a
container that is holding both solid and liquid.
Decant the liquid into the small beaker and then if
it does not have any solids in it you can discard
the decanted liquid.
After decanting, rinse the solid with about 25-mL
of distilled water. Decant again. Repeat the step
three or four times.
Next, wash the solid with about 15-ml of HCl.
Decant and discard the liquid.
After the final washing with water, place the
copper into a drying over to dry. (The copper
should be in the original beaker that you have the
empty mass of)
Allow the copper to become completely dry, then
find the mass of the beaker plus the copper and
record it on the Report Sheet. Record your
observations of the solids. Discard the solids into
a trash can (NOT DOWN THE DRAIN).
Clean up all of your materials. Before you leave
the lab, wash your hands thoroughly with soap
and water.
DATA/OBSERVATIONS:
BEFORE THE REACTION
Mass of the empty, dry beaker
Mass of the beaker and copper (II)
chloride
Mass of the two iron nails (together)
AFTER THE REACTION
Mass of two iron nails (together)
Mass of beaker and copper after
drying
g
G
G
G
G
OBSERVATIONS
Beaker with the copper (II) chloride
solution (before the reaction)
Solution after the reaction is
completed
Solids after drying
CALCULATIONS:
Perform the following calculations. SHOW ALL WORK.
1. Calculate the mass of iron used in the reaction. (HINT: the nails are the source of iron).
2. Calculate the moles of iron used.
3. Calculate the mass of copper solid produced.
4. Calculate the moles of copper solid produced.
5. Based on the number of moles in #2 and #4, calculate your experimental ratio of moles of iron used to moles of
copper produced.
6. Write a balanced equation for this reaction assuming the products are copper and iron (III) chloride.
7. Using this equation and its coefficients and the mass of iron, calculate the theoretical yield of copper.
8. Calculate the percent yield for copper.
9. Show by calculation which reactant was limiting and which was in excess at the beginning of th reaction. Clearly
identify each.
CONCLUSION:
Don’t forget to do this.
QUESTIONS:
1. How does the ratio you calculated in #5 above compare to the actual ratio in the balanced equation?
2. How many grams of iron would need to be used if 44.45 grams of copper were to be produced?
3. Give some possible sources of error that may have caused your ratio to differ from the expected ration in the balanced
equation and your percent yield to be other than 100%