Chemical Reactions of Copper
Introduction
Chemistry is the study of matter and the changes it undergoes. In this experiment, you will carry
out a series of reactions involving the element copper. You will observe the four classes of
chemical reactions--synthesis, decomposition, single and double replacement. You will be
expected to write balanced equations for these various reactions. You will see evidence of
chemical change and the conservation of matter. You will learn common lab practices and lab
techniques. Finally, the concept of recycling is introduced.
Procedure
Day 1 Conversion 1 -- Changing Cu to Cu(NO3)2
Label a 125 mL flask with your name and your lab partner's name. Weigh and record its mass.
Take approximately 1.5 grams of copper wire; lightly sand with steel wool, weigh it and record
its mass. Loosely fold the wire so that it is flat and place the wire in the flask and carefully add
20 mL of 6.0 M nitric acid (HNO3). (Caution: The nitric acid is concentrated and dangerous to
skin and clothing.) Stopper the flask and briefly observe the reaction that takes place. It may
help to place the flask on a sheet of white paper to observe the color of new substances being
produced. Unstopper the flask and set the flask in the fume hood.
The brown gas produced by the reaction in nitrogen dioxide, a common component of
atmospheric pollution. The blue color of the solution is characteristic of many copper
compounds dissolved in water. Expressed in words, the reaction is
copper + nitric acid --> copper (II) nitrate + nitrogen dioxide + water
Day 2 Conversion 2 -- Changing Cu(NO3)2 to Cu(OH)2
Put a laboratory apron on. Test your solution from Conversion 1 with pH paper, by dipping the
end of a stirring rod into the solution and then touching a piece of pH paper. It should indicate
that the solution is acidic due to a little excess HNO3. (pH paper should indicate a pH less
than 7.) Obtain 35 mL of 6.0 M sodium hydroxide. Test it in the same manner. Record both test
paper results.
Read this entire paragraph before starting!
Add about 150 mL of ice water to a 400 mL beaker. (This is called a cold water bath.) Place the
flask into the beaker so that the water comes up almost to the neck of the flask. (Why do you
think this is necessary?) Add, slowly and cautiously, 5 mL portions of the 35 mL of sodium
hydroxide to your copper (II) nitrate in the flask. Continuously mix the solution with a stirring
rod. Proceed until all 35 mL of NaOH have been added. The sodium hydroxide, a base,
neutralizes the excess nitric acid and then reacts with the copper ions in the solution. The
neutralization reaction produces considerable heat, so stir well and add the sodium hydroxide
solution slowly.
Test the resulting solution with the acid/base test paper and record the result. If the solution is
still acidic, add more NaOH (as before) and test again. The solution must be basic. (pH paper
should indicate a pH more than 7.)
Observe the solution. It should be royal blue (not green) and thick, the result of the production of
copper (II) hydroxide, a solid. (A solid which results from a chemical reaction is called a
precipitate.)
copper (II) nitrate + sodium hydroxide --> copper (II) hydroxide + sodium nitrate
Day 2 Conversion 3 -- Changing Cu(OH)2 to CuO
Add approximately 40 - 50 mL of distilled water to the flask containing the
copper (II) hydroxide precipitate. Place the flask higher than normal over your Bunsen burner
and as an added precaution, clamp the flask to the ring stand so it cannot be knocked off the
ring clamp. Heat the precipitate gently and stir constantly, until all the material
is converted to a black substance, copper (II) oxide. Then turn off the Bunsen
burner. Continue stirring for 30 seconds after the heat is turned off.
Remove the stirring rod. Use your wash bottle to rinse off any black particles of CuO that are on
the stirring rod back into the flask. Let the solution cool for three minutes. Then decant the
extra liquid, using your stirring rod to direct the flow of liquid into the sink, as shown by your
teacher. Don't lose any solid! You don't need to pour off every drop of liquid--just most of it.
Now repeat this washing process with another 40 - 50 mL of deionized water. Stir, let the solid
CuO settle and decant off the extra liquid. Repeat a third time. This process is called washing
the precipitate.
copper (II) hydroxide --> copper (II) oxide + water
Day 3 Conversion 4 -- Changing CuO to CuCl2
Add 30 mL of 3.0 M hydrochloric acid to the flask containing the black copper (II) oxide. Swirl
the flask gently, while leaving it flat on the counter. The oxide will soon dissolve, leaving a clear
aqua-blue solution of copper (II) chloride. If the oxide has not all dissolved after 3 minutes, add
more HCl, a few mL at a time, until it does.
copper (II) oxide + hydrochloric acid --> copper (II) chloride + water
Day 3 Conversion 5 -- Changing CuCl2 to Cu
Cut about 0.7 grams of aluminum wire. Bend it half, and in half again. Then bend it to form an
"L" shape, so that the bottom of the "L" will just fit in the bottom the flask. Put the wire in the
flask and watch what happens; the aluminum wire seems to "grow copper hair." Occasionally
shake the wire to dislodge the copper. Since the reaction produces hydrogen gas, the absence of
bubbles will indicate when the reaction has stopped. The solution will gradually turn grayish as
gas is produced; later you will find that it has cleared when the reaction is complete. Let the
flask stand overnight.
There are two reactions going on simultaneously. One is the reaction that results in the recovery
of the metallic copper. The other is the reaction of the aluminum with a left-over hydrochloric
acid; this produces the hydrogen gas.
(1) copper (II) chloride + aluminum --> aluminum chloride + copper
(2) hydrochloric acid + aluminum --> aluminum chloride + hydrogen gas
Day 4 -- Washing and drying the recovered copper
The next day you will find that the copper has settled. There should be no blue color at all left in
the solution. Remove any remaining aluminum wire from the flask. Decant and discard the clear
liquid. Wash the copper three times with 50 - 80 mL portions of warm distilled water, decanting
the liquid each time. Don't lose any copper! After the final decanting, the flask should contain
only pure, but wet, copper. Place it in an oven for drying overnight.
Day 5 -- Weighing the recovered copper
Weigh the flask containing the dry copper on the same balance you used on Day 1. Also make
observations about your copper. Does it appear pure? Is there any discoloration? (Remember
that compounds containing copper are usually blue or green in color.)
Scrape as much copper into a garbage can as possible. Wash the flask with soap and water.
Return all equipment.
Analysis: Each question should be answered in complete sentences and typed.
1. Write balanced equations for each reaction that occurred. Include states of matter for each
substance involved (Refer to your observations and lab procedure to help determine the states of
matter.
a. Categorize each reaction as single replacement, double replacement, decomposition,
synthesis or none of the above.
b. For each reaction, state one piece of specific evidence that a chemical change was
taking place.
4. Answer the following questions for the reaction performed in conversion 1:
a. Where did the copper atoms from the wire go?
b. What elements were in the gas molecules released by the reaction?
c. Where did these atoms come from?
d. Where did the atoms in the water produced come from?
5. What is a precipitate? What is the purpose of washing a precipitate? What is the purpose of
decanting?
6. Calculate the percent of copper recovered. Show all work.
7. Was your percent recovery over or under 100%? Why do you think this is? Explain two
specific places in the lab where a mistake could have been made causing your percent to be
above or below 100%. (This answer should pertain to YOUR percent, so it should be answered
based on your answer to question #6.)
7. What was the purpose of the cold water bath in Conversion 2 on Day 2?
8. How is the concept of recycling related to this experiment?
9. Draw a flow chart that illustrates the transformations of copper throughout this lab.
10. Use the balanced chemical equation from Day 1 to answer the following question: How
many grams of copper(II) nitrate did you produce from the amount of copper(grams) you started
with on Day 1? Show ALL YOUR WORK!!!!
11. Use the first balanced chemical equation from Day 3 to answer the following question: How
many grams of copper were produced from the amount of aluminum(grams) you started with on
Day 3? Show ALL YOUR WORK!!!!