Investigation 5 Lab Moles
Synthesis of Two Copper Oxides
Required measurements and calculations: Clearly record all required observations and measurements with
units. Show all calculations and clearly and label your answers. Ask the teacher to initial your original data.
1. Copper metal, when heated, combines with oxygen (O2) in the air to form copper (I) oxide (Cu2O), after
heating to a higher temperature, copper (II) oxide (CuO) forms.
a) Write the balanced equation for the reaction of copper metal with oxygen to form copper (I) oxide
(Cu2O).
b) Write the balanced equation for the reaction of copper (I) oxide (Cu2O) with oxygen to form copper
(II) oxide (CuO).
2. How many grams are equivalent to 0.020 moles of copper?
3. Weigh out 0.020 moles of copper and put it in an evaporating dish.
4. PUT ON GOGGLES BEFORE CONTINUING.
5. Use a Bunsen burner to heat the copper. Heat gently and record observations. Heat more strongly and
observe. Turn off the burner and allow the evaporating dish to cool. You may remove your goggles now.
6. After the evaporating dish is cool to the touch (10-15 minutes). Put the copper oxide in a solid waste
container and rinse and dry the evaporating dish.
7. Use the balanced equations to calculate:
a) How many moles of copper (I) oxide (Cu2O) would be produced from the reaction of 0.02 moles of
copper?
b) How many grams of copper (I) oxide (Cu2O) would that be?
Investigation 5 Lab Moles
Gummy Bear Decomposition of KClO3
Required measurements and calculations: Show all calculations and clearly label your answers.
Ask the teacher to initial your original data.
1. Write the balanced equation for the VERY exothermic decomposition of potassium
chlorate (KClO3) forming potassium chloride (KCl) and oxygen gas (O2).
2. Calculate the molar mass of KClO3.
3. We will put 5.0 grams of KClO3 into a large test tube. Calculate how many moles of
KClO3 this is.
4. Use the balanced equation to determine how many moles of KCl will be produced
from the decomposition of 5.0 grams of KClO3.
5. Calculate how many grams of KCl will be produced from the decomposition of 5.0
grams of KClO3.
6. The Gummy Bear is added to the tube and sacrificed to fuel the reaction. The tube is
heated to provide the activation energy necessary to get the reaction started. Observe
carefully and record complete observations.
7. Draw and clearly label an Energy Diagram for this reaction.
Investigation 5 Lab Moles
Burning Magnesium
Perform the following procedure carefully and then answer each question completely in your data book: Show all calculations
and clearly label your answers. Ask the teacher to initial your original data.
Procedure: Goggles must be worn for this experiment.
 Cut Magnesium ribbon to a length of 7.0 cm. Weigh and record in a data table.
 Weigh and record the mass of an empty dry crucible in your data table.
 Set up a ring stand, with a ring and a clay triangle to support your crucible. Ensure proper height for use of the Bunsen Burner
underneath.
 Roll your magnesium ribbon so that it fits in your crucible and place it in the crucible on the ring stand apparatus.
 Heat your sample until all the magnesium burns completely. Avoid looking directly at the light.
 After letting the crucible cool for at least ten minutes weigh the crucible with product and record in the data table.
 Carefully wash and dry your crucible and return it to the teacher.
 Calculate the mass of only your product, record, and have your teacher initial your data table.
1. List three observations for this reaction.
2. Write the balanced equation for the reaction that takes place when magnesium metal is burned in the
presence of oxygen gas (O2) to form magnesium oxide.
3. What kind of reaction is this? What kind of compound is the product in this reaction?
4. How many moles is your sample of magnesium?
5. How many moles of oxygen are needed to burn your sample? How many grams of oxygen is this?
6. Draw and label an appropriate Energy Diagram for this reaction.
Extension questions:
1. Based on your mass of magnesium, how many grams of magnesium oxide should you be theoretically able to make?
2. Calculate your % Error for this lab. % Error = |Theoretical mass - Experimental mass| X 100
Theoretical mass
3. Calculate your % Yield for this lab.
% Yeild = Experimental mass X 100
Theoretical mass
Investigation 5 Lab Moles
Combining Iron and Sulfur
Perform the following procedure carefully and then answer each question completely in your data book: Show all calculations
and clearly label your answers. Ask the teacher to initial your original data.
Procedure: This experiment must be done in the fume hood and goggles must be worn!
 Weigh out 0.10 grams of both iron powder and sulfur. Place both reactants into a plastic cup. Stir.
 Under the plastic cup, wave a bar magnet and note the results.
 Transfer the mixture into a small test tube. Slowly heat the mixture over a Bunsen Burner while holding the test tube with a test
tube holder. * (sulfur’s fumes are toxic, so turn on the fume hood.)
 After letting the test tube cool for at least ten minutes, carefully break the test tube and release the product.
 Wave the bar magnet over the product. Note results.
1. List four observations for this reaction. (include smell)
2. Write the balanced equation for the reaction that takes place when iron powder is burned in the
presence of solid sulfur to form iron (II) sulfide. (hint: the (II) means a 2+ charge)
3. What kind of reaction is this? What kind of compound is the product in this reaction?
4. How many moles is your sample of iron?
5. How many moles of iron (II) sulfide can be formed from this amount of iron? How many grams of iron
(II) sulfide is this?
6. Draw and label an appropriate Energy Diagram for this reaction.
Extension questions:
1. Based on your mass of sulfur, how many grams of iron (II) sulfide should you be able to make?
2. Compare this mass to the mass in question 5 above. Is it more or less? Which amount do you think is
correct and why?
Investigation 5 Lab Moles
Making Oxygen Gas
Perform the following procedure carefully and then answer each question completely in your data book: Show all calculations
and clearly label your answers. Ask the teacher to initial your original data.
Procedure: For this experiment, goggles must be worn!
 Weigh a clean dry test tube, and pour 10.0mL of hydrogen peroxide into it. Record the mass of the test tube and sample.
 Have a wooden splint and matches ready. Add a very small amount, the tip of a spatula of MnO2 to the hydrogen peroxide.
 One person should cover the reaction with an inverted test tube, while the other lights the splint and allows it to burn for about 5
seconds.
 Blow out the flame and place the glowing splint halfway down the test tube.
1. List four observations for this reaction.
2. Write the balanced equation for the reaction that takes place when hydrogen peroxide (H2O2) forms
water and oxygen gas.
3. What kind of reaction is this? What kind of compound is the reactant in this reaction?
4. Calculate the mass of your hydrogen peroxide sample. How many moles is this sample of H 2O2?
5. How many moles of oxygen can be formed from this amount of H2O2? How many grams of oxygen is
this?
6. Draw and label a model of this reaction
7. Draw and label an appropriate Energy Diagram for this reaction.
Extension questions:
1. The manganese (IV) oxide was not part of the reaction rather it was a catalyst. Use your references in
class to define a catalyst and re-draw the Energy Diagram to show the effect the catalyst has on the
reaction.
2. How many grams of water were made in this reaction?