Experiment 7: Virtual Copper Carousel Lab


Experiment 7: Virtual Copper Carousel Lab

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Starting with elemental copper, this experiment uses a series of chemical reactions to end with the recovery of the elemental copper. Since you start and end with metallic copper and can perform the cycle of reactions again and again, this has been referred to as a ‘copper carousel’. Just like going around and around on the amusement park ride, you can run this reaction cycle again and again. You will be asked to observe the changes associated with these reactions, write balanced reaction for each step, identify which reactions are oxidation-reduction reactions, and calculate the percent recovery using hypothetical data.

Early alchemists were intrigued by the nature of matter and how it changed. However, their investigations were hindered by a lack of instrumentation (except for balances to measure mass), and they did not have the benefit of modern atomic theory or the periodic table to help interpret their results. They relied on descriptive chemistry, or simply describing the physical changes that accompany a chemical reaction.

Chemical changes (i.e., reactions) also produce observable changes in the substances involved in the reaction. Some reactions, such as combustion, will generate heat and light. Other reactions, such as the reaction of carbonates with acids, will produce bubbles of gas. Still other reactions will result in a change in the color or appearance of the reacting materials. These changes can be used to determine whether or not a chemical reaction has occurred. Modern chemists describe these chemical changes (i.e., reactions) using chemical formulas to describe the starting materials and products, and using chemical equations to indicate how these substances are changed during the reaction.

There are many common types of chemical reactions. Oxidation-reduction (redox) reactions involve the transfer of electrons. Not all reactions are redox reactions. Simple precipitation reactions most often do not involve the loss or gain of electrons – they are not redox reactions. Acid-base neutralization reactions are another example of reactions that are not typically redox reactions. In this experiment, only some of the reactions are redox reactions.

Metals occupy a large portion of the periodic table. Properties and reactivity of metals have been studied throughout human history; indeed, the ability to use metals has had profound effects on civilization. The coinage metals — gold, silver, and copper — have always been valued. Metals can be classified by their location in the periodic table: active metals (Groups 1 and 2 or Groups 1A and 2A), transition metals (Groups 3 to 12 or Groups 1B to 8B), and p-block metals (Groups 13 to 16 or Groups 3A to 8A). Metals in their elemental form share many properties: they are generally shiny, malleable, ductile, and good conductors of heat and electricity. All but three are solids near room temperature.

The chemistry of metals, particularly the transition metals, is quite rich and a wide range of chemical compounds can be prepared from them. An excellent example of this is the element copper. Copper (Cu from the Latin cuprum) is element #29 on the periodic table. It has been known to mankind since antiquity; an entire age of human development is named after bronze, an alloy of this versatile metal. While most people are familiar with copper in its metallic state as wires or sheets, the solution chemistry of the element is equally noteworthy. Copper is found most often in aqueous chemistry as the Cu2+ cation, although compounds of copper(I) are also frequently encountered.

Starting from the elemental metal, you will watch a series of reactions to explore the chemistry of copper by preparing a series of compounds which move around a ‘carousel’, ultimately returning to the initial metallic state:

For this experiment, we’ll be using a number of YouTube videos that demonstrate the procedure.


This video demonstrates one way to do the copper cycle procedure: https://youtu.be/ylnMuFpDqUA Watch this video and answer the questions that are asked below. Record observations for each step, as requested.

Reaction I:
1. Weigh an empty evaporating dish and record the mass to four places to the right of the decimal point and include the units.
Mass of empty evaporating dish:

2. Obtain a piece of pure copper wire weighing about 0.5 g. Weigh the copper in the evaporating dish and record the mass with the appropriate units and number of significant figures.
Mass of empty evaporating dish:
Appearance of copper wire:

3. Transfer the wire to a 250-mL beaker. Be sure that the copper wire lies flat on the bottom of the beaker. THE NEXT STEP SHOULD BE PERFORMED IN THE HOOD! Add about 4-5 mL of concentrated nitric acid (16 M HNO3) to the beaker. Keep the beaker in the hood until the reaction is complete. Record your observations. The chemical equation for this reaction is complex, and is:
Reaction I: Cu(s) + 4 HNO3(aq) → Cu(NO3)2(aq) + 2 H2O(l) + 2 NO2(g)
List at least three observations of reaction of copper with nitric acid (include appearance of copper wire, the liquid, and any gas formation):

4. When the reaction is complete add ~100 mL distilled water to your beaker.
List your observations:

In solution, the copper ion does not exist as a ‘naked’ cation, but is coordinated by six neutral water molecules in an octahedral arrangement, [Cu(H2O)6]2+. This leads to a slightly different (but no less correct) rendering of the reaction as:
Reaction Ia: Cu(s) + 4 H3O+ (aq) + 2 NO3 – (aq) → [Cu(H2O)6]2+(aq) + 2 NO2(g)

Reaction II:
5. Add 15 mL of 6 M NaOH slowly while stirring with a glass rod. Be sure to keep the glass rod in the beaker to prevent any loss of your copper compounds during this step.
List at least two observations of this reaction:

Complete and balance the chemical equation that describes what happened.
Reaction II: [Cu(H2O)6](aq) + NaOH(aq) →

6. Let the mixture settle for a while.
List at your observations (include the color of the liquid and the solid):

Reaction III:
7. Assemble a ring stand, ring, wire gauze, and Bunsen burner, as illustrated in the video. Heat the solution in the beaker gently using a Bunsen burner. You do not need to heat the solution to boiling; a gently heating will be sufficient. After 5 to 10 minute the reaction should be complete. Complete reaction will be indicated when the solid product is uniform in color and the solution is clear and colorless.
List your observations of this reaction:

Complete and balance the chemical equation that describes what happened.
Reaction III: Cu(OH)2(s) + (heat) →

8. Remove the beaker from the wire gauze and replace it with a clean 600-mL beaker with 300 mL of distilled water (no stirring rod). Heat the water until it is close to boiling.
9. Allow the mixture from Reaction III to settle. Decant the supernatant solution (i.e., the liquid above the solid), and discard the supernatant down the drain with plenty of water. Be careful not to lose any of the solid CuO that was formed.
10. When the water in step 8 is hot, turn off the Bunsen burner and pour ~100 mL of the hot water into the beaker with the solid CuO. Stir briskly to wash any unreacted NaOH from the solid. Allow the solid to settle, and decant the supernatant as you did in step 9. Be careful not to lose any of the solid CuO.
11. Repeat step 10 two more times, being careful not to lose any of the solid CuO.
List your observations of this process:

Reaction IV:
12. Slowly add 30 mL of 3 M H2SO4 to the solid CuO while stirring with a glass rod.
List at least two observations of this reaction:

Complete and balance the molecular and net ionic chemical equations that describes what happened.
Reaction IV (molecular eq.): CuO(s) + H2SO4(aq) →
net ionic eq.:

Reaction V:
13. RETURN TO THE HOOD TO PERFORM THIS NEXT STEP. Weigh 1.30 g of 30-mesh zinc metal in a weigh boat or small beaker. Add the zinc to the solution all at once. Stir until the solution is colorless and the evolution of gas ceases, and then continue stirring for 2 to 5 more minutes.
List at least four observations of this reaction:

What is the formula of the gas that is evolving in this step?
Complete and balance the molecular and net ionic chemical equations that describes what happened.
Reaction V (molecular eq.): CuSO4(aq) + Zn(s) →
net ionic eq.:
This reaction has excess H2SO4(aq). Complete and balance the chemical equation that describes the reaction of the excess zinc.
Reaction Va: Zn(s) + H2SO4(aq) →

14. Once the reaction is complete, carefully decant the liquid. Rinse the precipitate two times with ~50 mL distilled water.
15. Transfer the precipitate to the evaporating dish.
16. Dry the precipitate over a steam bath as shown in the video.
17. Carefully remove the evaporating dish from the steam bath using tongs. Dry the bottom of the dish with a paper towel and reweigh the dish with your product. The balance reading in the video is very hard to read. Record this value as 44.7565 g.
18. You can evaluate the experimenter’s lab technique by how much of the original copper was recovered after all the reactions were complete. Percent recovery is calculated as:


What is the percent recovery for this experiment? (Show all steps of your calculation.):

This video demonstrates another way to do the copper cycle procedure: https://youtu.be/tyrne4AFOvY
Watch the first 4.5 minutes of this video and answer the following questions.
1. What was the purpose of the litmus paper?

2. What were the observations of the litmus paper?

3. What technique was used to recover the products from Reactions III and V?

4. Which acid is used in place of the sulfuric acid in Reaction IV?

This video discusses other reactions of copper: https://youtu.be/1I7lzHy0jKE Watch this video and write 50 – 150 words describing the most interesting things that you learned.

Additional Questions:

1. Many of the reaction in this experiment are oxidation-reduction (or redox) reactions. Define the following terms:
a) oxidation –
b) reduction –

Which of the five reactions in the copper cycle are redox reactions?

2. Write a balanced chemical equation for the reaction of iron with oxygen to form iron oxide (Fe2O3).

Then write a description of what you would observe as a piece of iron rusts. What was the sign of the chemical reaction in which iron rusts into iron(III) oxide?

For this reaction, which reactant is oxidized and which is reduced?

3. In the first reaction in this lab you dissolve the copper metal in nitric acid. Why must this reaction be performed in the hood until the reaction is complete?

4. In the first video, the products were recovered using a technique called decanting? Describe what is meant by decanting.

5. The mass of recovered copper may be measured greater or less than the amount at the beginning.
a) Give at least two reasons why the mass of recovered copper might be greater than the initial mass of copper.

b) Give at least two reasons why the mass of recovered copper might be less than the initial mass.

c) What steps could be taken to minimize these sources of error in the experiment?

6. Using your balanced chemical equation for Reaction V, calculate the minimum amount of zinc that is necessary to completely react with 0.4564 g of copper ions. How does this value compare with the amount of zinc added in the first video?

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