Honors Chemistry
Name: __________________________
Period: ________
Lab Partners: ____________________________________
Copper Conversions Lab
Introduction
In this multi-day lab you will start with a sample of copper metal and run several successive
reactions that produce different copper compounds. The last reaction will release the copper
atoms as copper metal again. If your technique is perfect, you will recover all the copper atoms
you started with (according to the Law of Conservation of Matter).
Chemical Concepts:
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Chemical Reactions
Law of Conservation of Matter
Materials:
Chemicals
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Copper Turnings, 0.5g
Hydrochloric Acid solution, 6M HCl,
8-10-mL
Nitric acid, 15.8 M HNO3, 2-3 mL
Sodium Hydroxide solution, 8 M
NaOH, 5-8 mL
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Sodium Phosphate, 0.4 M Na3PO4, 15
mL
Sulfuric Acid solution, 2M H2SO4, 1215 mL
Water, distilled or deionized, H2O
Zinc, mossy, Zn, 2g
Apparatus
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Balance, 0.001g readability
Beaker, 250 mL, 2
Beaker, large (for water bath)
Beral-types pipets, 5
Buchner Funnel
Evaporating dish
Filter paper, 2 pieces
Graduated cylinder, 10 mL
Drying oven
Hot plate or Tirrill burner setup
Ice
pH paper 1-12
Ring stand with ring
Stirring rod
Test tube (to store copper product)
Wash bottle (filled with distilled
water)
Watch glass
10 mL Pipette and Bulb
500 mL Filtering Flask
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Name: __________________________
Period: ________
Lab Partners: ____________________________________
Safety Precautions:
Nitric acid is severely corrosive, a strong oxidant, and toxic by ingestion and inhalation.
Hydrochloric acid solution is corrosive to skin and eyes and toxic by ingestion and inhalation.
Sodium hydroxide solution is a corrosive liquid, can cause skin burns, and is very dangerous to
eyes. Sulfuric acid is corrosive to eyes, skin, and other tissue. Avoid contact of all acids and
bases with eyes and all body tissue. Clean up all spills immediately: neutralize any acid spills
with a weak base; neutralize any base spills with a weak acid; wipe up with water. Wear
chemical splash goggles, chemical-resistant gloves, and a chemical-resistant apron.
Background:
The purpose of this laboratory experiment is to:
(a) Illustrate different types of chemical reactions.
(b) Show how a quantity of an element can be carried through a series of chemical
transformations without significant loss of mass, thereby illustrating the law of
conservation of matter, and
(c) Provide experience in fundamental laboratory procedures such as transferring a reagent
from a reagent bottle, transferring a solution or a solid from one vessel to another,
decanting, filtering, washing, and dissolving a precipitate.
A weighed quantity of copper will be carried through the following transformations:
Cu  Cu(NO3)2  Cu(OH)2  CuO  CuCl2  Cu3(PO4)2  CuSO4  Cu
In Part A, copper metal is oxidized by nitric acid to produce a blue solution containing cupric
nitrate, Cu(NO3)2, and a brown poisonous gas of nitrogen dioxide, NO2, as shown in Equation A:
Cu(s) + 4HNO3(aq)  Cu(NO3)2(aq) + 2 NO2(g) + 2H2O(l)
Equation A
In Part B, the blue cupric nitrate solution, Cu(NO3)2, reacts with sodium hydroxide, NaOH, in a
double replacement reaction to produce a blue precipitate of cupric hydroxide, Cu(OH)2,
according to Equation B:
Cu(NO3)2(aq) + 2NaOH(aq)  Cu(OH)2 + 2NaNO3(aq)
Equation B
An acid-base neutralization reaction also occurs in Part B between sodium hydroxide, NaOH,
and the excess nitric acid, HNO3, from Part A according to Equation B2:
NaOH(aq) + HNO3(aq)  NaNO3(aq) + H2O(l)
Equation B2
You will use litmus paper to indicate if your beaker contains acid or base. Litmus paper comes
in two colors, red and blue. They actually have the same dye in them, but that dye is RED if last
exposed to an acid and BLUE if last exposed to a base.
RED BLUE
ACID BASE
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Memory aid: Notice the D's and B's go together.
In Part C, the blue cupric hydroxide solid, Cu(OH)2, is decomposed with heating into a black
cupric oxide solid, CuO, according to the decomposition reaction shown in Equation C:
Cu(OH)2(s) + heat  CuO(s) + H2O(l)
Equation C
In Part D, the black cupric oxide, CuO, undergoes an acid-base neutralization reaction with
hydrochloric acid, HCl, to form a green cupric chloride solution, CuCl2, and water according to
Equation D:
CuO(s) + 2HCl(aq)  CuCl2(aq) + H2O(l)
Equation D
Part E involves a double replacement reaction between the green cupric chloride, CuCl2, and
sodium phosphate, Na3PO4, to produce a blue precipitate of cupric phosphate, Cu3(PO4)2, and
sodium chloride, NaCl, as shown in Equation E:
3CuCl2(aq) + 2Na3PO4(aq)  Cu3(PO4)2(s) + 6NaCl(aq)
Equation E
An acid-base neutralization reaction also occurs in Part E between sodium hydroxide, NaOH,
and the excess hydrochloric acid, HCl, from Part D as shown in Equation E2:
NaOH(aq) + HCl(aq)  NaCl(aq) + H2O(l)
Equation E2
Part F involves another acid-base neutralization reaction as sulfuric acid, H2SO4, dissolves the
blue cupric phosphate, Cu3(PO4)2, to produce a blue cupric sulfate, CuSO4, and phosphoric acid,
H3PO4, as shown in Equation F:
Cu3(PO4)2(s) + 3H2SO4(aq)  3CuSO4(aq) + 2H3PO4(aq)
Equation F
The final reaction in Part G involves the replacement of a less active metal (copper) by a more
reactive metal (zinc). The zinc, Zn, replaces the copper in cupric sulfate, CuSO4, forming zinc
sulfate, ZnSO4, and solid copper, Cu, according to the single replacement redox reaction shown
in Equation G:
Zn(s) + CuSO4(aq)  ZnSO4(aq) + Cu(s)
Equation G
The excess zinc in Part G, over and above that required to replace all of the copper, is dissolved
by sulfuric acid, H2SO4, according to the single replacement redox reaction shown in Equation
G2:
Zn(s) + H2SO4(aq)  ZnSO4(aq) + H2(g)
Equation G2
Pre-Laboratory Assignment
Read the entire lab and then complete the pre-laboratory assignment. This must be done and
checked before beginning work in the laboratory.
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1. Write the correct chemical formula for each substance listed in the first column. Then read
the laboratory procedure to determine where each material is found in the procedure, either
as a product or as a reactant. List the corresponding part(s) of the procedure in the third
column.
Chemical Name
Copper
Chemical Formula
Corresponding Part(s)
Cupric chloride
Cupric hydroxide
Cupric nitrate
Cupric oxide
Cupric phosphate
Cupric sulfate
Hydrochloric acid
Nitric acid
Sodium hydroxide
Sodium phosphate
Sulfuric acid
Zinc
2. Define each of the following terms
a. Reagent
b. Decantation
c. Filtration
d. Precipitation
e. Filtrate
f. Supernatant liquid
3. What does the M in 15.8 M HNO3 and in 6 M HCl stand for?
4. Strong acids and bases are very damaging to the eyes as well as to body tissue in general.
For this reason, what must be worn at all times while in the laboratory (even if you are not
personally working with these chemicals)?
5. NO2 is a red-brown, reactive, poisonous gas. It is evolved in the reaction in Part ___. For
this reason, reaction ___ must be done _________________________.
6. Imagine that two beakers containing colored, but clear solutions were mixed. How could you
determine if precipitation had occurred”? What would you observe?
7. If precipitation occurs, how could you separate the solid from the solution?
8. Copper compounds are frequently colored. Imagine a student is doing Part E of this lab.
This student is filtering the Cu3(PO4)2. The filtrate is slightly blue.
a. Why might the filtrate be blue?
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b. Should the filtrate be discarded? Why or why not?
c. If the filtrate is discarded, how will this affect the final mass of the copper that was
recovered?
d. How might a blue color in the filtrate have been avoided?
Procedure
Part A – Preparation of Cupric Nitrate from Copper by Oxidation with Nitric
Acid.
1. Tare a 250-ml. beaker. Add about 0.5 grams of copper turnings and record the mass to
the nearest thousandth of a gram in the data table.
2. Place the beaker containing the copper in the fume hood. Using a Beral-type pipet, get
about 2 ml. of concentrated nitric acid. Caution: This concentrated nitric acid is strong
and corrosive. Be sure to wear chemical splash goggles, chemical resistance gloves,
and a chemical resistant apron. Add the nitric acid drop-wise. Wait a couple of
minutes as the reaction starts.
3. Warm the beaker gently on a hot plate on low until the copper is completely reacted.
Record detailed observations in the data table. As the reaction slows down, use the
stirring rod to push each piece of copper off the walls of the beaker and into the acid. If
necessary, add a few more drops of 15.8 M HNO3 to react with any remaining copper.
This reaction is done when all the copper metal has reacted. There will be some excess
acid so be careful with this liquid.
Caution: The brown fumes, NO2, released are poisonous; do not remove the beaker
from the fume hood until all of the copper is reacted and no more gas is released.
4. Dilute the solution with 10 mL of distilled or deionized water.
Part B – Preparation of Cupric Hydroxide from Cupric Nitrate
5. Measure the pH of the solution from Part A using 1-12 pH paper. Using your stirring rod
test the results of placing a drop of the solution on the pH paper. Record the pH on the
data table.
6. Measure the pH of the 8 M sodium hydroxide (NaOH) solution using 1-12 pH paper.
Using your stirring rod test the results of placing a drop of the solution on the paper. Be
sure to rinse the stirring rod with water each time you want to test a new solution.
Record the pH of the solution on the data table.
7. To the solution from Part A, add 5 mL of 8 M sodium hydroxide (NaOH) solution.
Caution: Part A solution is dangerous as a strong acid.
8. Allow the precipitate to settle a bit. Test for complete precipitation by adding 1 more
drop of 8 M NaOH to the clear supernatant liquid. If no additional precipitate forms, the
reaction is complete. If additional precipitate forms, add more 8 M NaOH drop wise
until precipitation is complete.
9. Measure the pH of the resulting solution using 1-12 pH paper. Using your stirring rod
test the results of placing a drop of the solution on the pH paper. The solution should be
basic, compare the color of your paper to the scale on the side of the pH paper container.
Since the solution itself is blue, be sure to look at the spread of the liquid as it touches the
paper. If the solution is not basic, add 8 M NaOH drop wise until it is basic. Record the
final pH on the data table.
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Part C – Preparation of Cupric Oxide from Cupric Hydroxide by
Decomposition
10. To the mixture from Part B, add enough distilled water to give a total volume of about
100 ml.
11. Heat the beaker on a hot plate until it just starts to boil, boil for about 3 minutes. While
heating stir occasionally leaving your stirring rod in the beaker between mixings. Take
care to watch the beaker closely during heating. If it starts to jump around, be prepared
to remove it from the hot plate immediately with beaker tongs. Take it off the hot plate
and let it cool.
Figure 1: Buchner Funnel Filtration Setup
12. Prepare a filtration setup with a Buchner Funnel and ensure the rubber funnel stopper fits
firmly in the filtering flask. Connect a 500mL Filtering Flask to the aspirators on the
faucets using a short piece of rubber tubing. Carefully insert a piece of filter paper into a
clean Buchner Funnel. Moisten the filter paper so that it fits snugly into the funnel. Turn
on the water and test the filtration system with a small amount of water to be sure the
water flows freely and rapidly. This will save a great deal of time in the filtering process.
13. Filter the mixture. Pour the mixture into the Buchner funnel using a glass stirring rod to
direct the flow of the mixture. Do not add so much solution into the filter that it raises
over the sides of the filter paper. Wash any traces of solid from the beaker using a stream
of water from a wash bottle.
14. Wash the solid precipitate in the filter paper twice with hot distilled or deionized water.
To do this, add enough water to completely cover the precipitated copper oxide and allow
water to run through.
15. Leave the product in the filter paper for Part D.
Part D – Preparation of Cupric Chloride from Cupric Oxide
16. Discard the filtrate in the 500 mL filtering flask and rinse the filtering flask with distilled
water to ensure it is clean. Place it below the funnel containing the precipitate from Part
C.
17. React the solid by carefully pouring 8 mL of 6 M HCl directly into the funnel. Allow the
resulting solution to run into the filtering flask. Use the aspirator to make the filtering
faster. (See Figure 1). If not all the precipitate on the filter paper is reacted use the pour
the solution in the filtering flask into a beaker and pour it over the precipitate to continue
the reaction until all the precipitate has reacted.
18. If the reaction of the solid seems to have stopped, add more 6 M HCl in 1 mL portions.
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19. When the solid is completely reacted, the filter paper should be washed twice with a
minimum amount of cold distilled or deionized water and the washings allowed to run
into the filtering flask with the solution. Transfer the solution from the filtering flask into
the 250 mL beaker and be sure to rinse the filtering flask with water and add the
washings to the solution in the 250 mL beaker.
Part E – Preparation of Cupric Phosphate from Cupric Chloride
20. Using a hot plate, heat 250 mL of distilled water for use in step 29.
21. Cool the solution from Part D to about room temperature by placing the beaker
containing the solution into a large beaker of cold water.
22. Neutralize the solution by first adding 8 M NaOH, drop by drop with constant stirring,
until the liquid in the beaker acquires a slight murkiness or cloudiness.
23. The add 6 M HCl, drop-by-drop with constant stirring, until the cloudiness just
disappears. The solution should be neutral at this point and have a pH of about 7. Test
the pH by touching a stirring rod first to the solution and then to a piece of 1-12 pH
paper. Add more NaOH or HCl as necessary to adjust the pH to 7. The solution may still
be cloudy.
24. Add 15 mL of 0.4 M Na3PO4, slowly and with constant stirring. Observe precipitate
formation.
25. Allow the precipitate to settle a bit. Test for complete precipitation by adding 1 more
drop of 0.4 M Na3PO4 to the liquid. If no additional precipitate forms, the reaction is
complete. If additional precipitation does occur, add more 0.4 M Na3PO4 drop wise until
precipitation is complete. If the quantity of precipitate is so large that a thick slurry is
formed, add a small amount of water to give a suspension in which the precipitate is free
to settle.
26. Using a hot plate, very carefully heat the solution until it just begins to show signs of
boiling. This will make filtration easier by helping to consolidate the precipitate. Stir the
mixture so that it does not boil vigorously or burn. Do not allow the mixture to boil
vigorously.
27. Prepare a filtration setup as described in step 10 and illustrated in Figure 1.
28. Filter the mixture. Do not add so much solution into the filter that is rises over the sides
of the filter paper. Wash any traces of solid from the beaker using a stream of water from
a wash bottle.
29. Wash the solid precipitate in the filter paper twice with hot distilled or deionized water.
To do this, add enough water to completely cover the precipitate and allow it to run
through.
30. Discard the colorless filtrate; leave the product in the filter paper for Part F.
Part F – Preparation of Cupric Sulfate from Cupric Phosphate
31. Discard the filtrate from the 500 mL filtering flask and rinse it with deionized water to
ensure it is clean. Place it below the funnel containing the precipitate from Part E.
32. React the solid by carefully pouring 12 mL of 2 M H2SO4 directly into the funnel. Allow
the resulting solution to run into the filtering flask.
33. If the solid does not completely react when the H2SO4 runs through the filter the first
time, the liquid that has run through should be poured back through the filter. Be sure to
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use a second clean beaker to transfer the filtrate in the filtering flask and pour this into the
funnel. This should be repeated until the solid completely reacts.
34. If the reaction of the solid seems to have stopped, add more 2 M H2SO4 in 1 mL portions.
35. When the solid is completely reacted, the filter should be washed twice a minimum
amount of cold distilled or deionized water and the washings allowed to run into the
beaker with the solution. Be sure to rinse the second beaker with water and add the
washings to the solution.
Part G – Preparation of Copper from Cupric Sulfate
36. To the solution form Part F, add about 2 grams of mossy zinc. Use a weighing boat to
mass the amount of zinc.
37. Allow the solution to stand, with occasional stirring, until it is entirely colorless. This
indicates that all of the copper has been removed from solution. If absolutely necessary,
add a little more zinc. The more zinc you add in excess the more sulfuric acid will be
required in the next step.
38. Once the solution is colorless, the excess zinc must be dissolved. To do this, add 2 M
H2SO4 in 1 mL portions with stirring until the excess zinc is completely reacted. Warm
the solution to speed the rate of solution. Complete dissolution of zinc will be indicated
by the fact that the precipitate is uniformly copper colored and no more hydrogen gas
bubbles are produced.
39. Allow the copper to settle. Decant off and discard the supernatant liquid. In the process
of decanting, the liquid is carefully poured off while the solid remains behind. To avoid
the loss of solid, the liquid is not drained off completely; some is allowed to remain
behind in the solid.
40. Wash the copper in the beaker three times with cold distilled or deionized water. To do
this, add about 20 mL of water to the precipitate in the beaker. Stir thoroughly with a
stirring rod, allow the precipitate to settle, and decant off the supernatant liquid. Repeat
this two more times.
41. Weigh a clean, dry evaporating dish and record the mass in the Data Table.
42. Use a small amount of water to flush the copper from the beaker into the evaporating
dish. Decant the supernatant clear liquid from the dish.
43. Dry the copper by heating the evaporating dish on top of a boiling water bath over a
Tirrill burner. The heat of the escaping steam causes the water in the dish to evaporate
and dries the contents of the dish without overheating the copper.
44. When the copper in the evaporating dish is dry, wipe the outside of the dish dry with a
clean towel. Allow the dish to cool to room temperature and weigh the dish and the
copper. Record this mass in the Data Table.
45. Save your recovered copper. Place it in a test tube labeled with your name and turn it in
to your instructor.
Disposal
Dispose of the remaining copper in the beaker provided by the instructor.
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Data Table
Data for Part A
Balanced Equation:
Observations:
Descriptions before the reaction:
Copper:
Nitric acid:
Description during reaction:
Description after reaction:
Data: Mass of Beaker
____________g
Mass of Beaker and Copper ____________g
Initial mass of copper
____________g
Data for Part B
Balanced Equations:
Observations:
Description of beaker before the reaction:
Description of beaker during reaction:
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Description after reaction:
Data: pH of solution from Part A: _________
pH of 8 M NaOH solution: _________
pH of final solution:
_________
Data for Part C
Balanced Equation:
Observations:
Description of beaker before heating:
Description of beaker after heating:
Description of filtrate:
Data for Part D
Balanced Equation:
Observations:
Description of filter paper before reaction:
Description of filter paper and beaker after reaction:
Data for Part E
Balanced Equations:
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Observations:
Description of beaker before reaction:
Description of filter paper after reaction:
Description of filtrate:
Data for Part F
Balanced Equation:
Observations:
Description of filter paper before reaction:
Description of beaker after reaction:
Data for Part G
Balanced Equations:
Observations:
Data: Mass of evaporating dish:
Mass of dish plus dry copper:
Mass of copper actually produced:
____________________g
____________________g
____________________g
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Calculations:
Final copper mass:
Initial copper mass:
Percent recovery:
_________g
_________g
_________%
Questions and calculations:
Please answer all questions on a separate sheet of paper. Show all calculations clearly.
1. For each reaction:
a. Write a balanced stoichiometric equation and the net ionic equation. Note:
Parts B, E, and G include a second reaction in which excess reagents are used up.
Include these as well for only the balanced stoichiometric equation. This can be
done on your data table space.
b. Include all physical states, using the following abbreviations: s = solid, g = gas, l
= liquid, and aq = aqueous. (include in net ionic equation)
c. Write the color of each material under the formula in the balanced equations.
d. Label the reaction types for each equation (whether it is a double replacement,
single replacement, decomposition, or synthesis reaction).
2. List the observations you made that gave evidence that a chemical reaction had taken
place.
3. What is the maximum temperature to which the evaporating dish containing copper will
be heated in Part G? How is this known?
4. What will happen if the copper is overheated during the drying process? How will this
affect the percent recovery?
5. Comment on the physical appearance of the reclaimed copper. How could it be made to
look more like the original copper?
6. Calculate the percent recovery of the copper by means of the formula for percent yield.
Percent Yield = mass of copper recovered/mass of copper used x 100%
7. Discuss your percent recovery. Was it reasonable? Why or why not? Support your
answer.
8. Calculate the amount of copper lost (or gained) during the reaction process.
9. List some possible errors that might lead to a mass of reclaimed copper less than that
originally used.
10. List some possible errors that might lead to a mass of reclaimed copper more than that
originally used.
11. Answer the following questions specific to each part:
Part A
a. You know how many grams of copper you used. How many grams of blue
copper nitrate are dissolved in the liquid in Part A?
b. How many grams of nitrogen dioxide gas escaped up the vent in Part A?
Part B
c. How many grams of copper (II) hydroxide did you produce in Part B.
Part C
d. Now calculate how much copper (II) oxide you produced in Part C.
Part D
e. Now calculate how much copper (II) chloride is floating around in your new blue
solution.
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Part E
f. Now calculate how much copper (II) phosphate is precipitated.
Part F
g. Now calculate how much copper (II) sulfate is floating around in your new blue
solution.
Conclusion
With this lab include a short summary describing what happened to the copper atoms as they
went through each procedure. Conclude with a paragraph that tries to account for the
discrepancy between the amount of copper you should have gotten, and the amount you actually
produced.
What does this experiment mean to you in terms of recycling? Why is recycling a good method
for waste management? What are the limits of recycling? Relate your response to this
experiment.
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