Moles of Iron and Copper

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 1023 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 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.

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 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.

8.  After the nails are completely dry, find the mass of the nails and record it on the Report Sheet.

9.  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.

10.  After decanting, rinse the solid with about 25-mL of distilled water. Decant again. Repeat the step three or four times.

11.  Next, wash the solid with about 15-ml of HCl. Decant and discard the liquid.

12.  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)

13.  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).

14.  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 / g
Mass of the beaker and copper (II) chloride / G
Mass of the two iron nails (together) / G
AFTER THE REACTION
Mass of two iron nails (together) / G
Mass of beaker and copper after drying / 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%