Copper Cycle Lab
Copper is an interesting metal for many reasons, and we will make use of several of its unusual
properties in the following lab. First of all, copper is a metal that is fairly easy to oxidize given the
correct reagent. In other words, you can strip electrons away from the neutral atoms to convert
elemental copper into copper (II) ions (or more rarely, copper (I) ions). These metal ions then combine
with anions to make different ionic compounds. Second, many copper compounds are colorful and not
very soluble in water, so they can perform interesting precipitation reactions. Finally, copper (II) ions
are fairly easy to reduce back to elemental copper with the addition of a reducing agent that adds
electrons back to the copper. Today’s lab involves converting copper metal into several different copper
compounds, and then converting it back to copper metal. Theoretically, one should be able to recover
the full mass of the original copper sample at the end of the lab. You should take careful notes during
the lab and record all of your qualitative and quantitative observations, as you will need this information
for your lab “report.”
Safety Notice
Read this procedure thoroughly before you begin! We will be using a number of fairly concentrated
acids and bases for this lab. Please exercise great care with these materials, as they can cause serious
burns to skin or eyes. Be scrupulous about wearing your goggles correctly, and insist that your partner
do the same. If you spill any of the strong acids, add NaHCO3 (baking soda) to yourself or to the lab
bench to neutralize them. Excess base can be neutralized with vinegar. Be sure to flush any exposed
body part with water immediately after exposure.
Part A – Oxidizing the elemental Copper to the Nitrate
Copper is impervious to most acids, and will not form hydrogen gas in a single replacement redox
reaction. However it will react with relatively concentrated nitric acid, an “oxidizing acid.” Add 25 mL
of 6.0 M HNO3 to a clean, dry 250-mL beaker, and label it with your name. Then measure out
approximately 1.00 gram of copper metal, and record its exact mass. If you are using 18-gauge copper
wire as your source of copper, try about 14 cm. Take your beaker to the hood and place it on a hot plate.
Turn up the heat to setting 4. Make sure the fan is operating, and carefully add your copper wire to the
acid in the beaker.
As the HNO3 oxidizes the copper, it releases corrosive reddish brown NO2 gas that you will be sure not
to breathe. Make notes in your notebook about all of your observations. When the reaction is finished,
you should see no solids remaining and you should also observe the characteristic color of aqueous
copper (II) nitrate. Write an equation (including state symbols) for the chemical reaction you just
witnessed. There are three total products.
Part B – Converting the Nitrate to the Hydroxide
Let your beaker cool briefly in the hood, and then take it back to your lab station. Add 40 mL of cold
water to cool and dilute your solution. Then add 60 mL of 3.0 M NaOH to convert your nitrate to the
hydroxide. You should observe a gelatinous precipitate. Make notes in your notebook about all of your
observations. Write an equation (including state symbols) for the chemical reaction you just witnessed.
Part C – Converting the Hydroxide to the Oxide
The copper (II) hydroxide is not heat stable. Use a ring stand and Bunsen burner to gently heat the
contents of your beaker just to the boiling point, stirring occasionally. Watch the beaker carefully for
signs of a reaction, and to make sure it does not boil over. Once the reaction is complete, turn off the
flame and let the beaker cool on the ring stand for at least 5 minutes. While you are waiting, record
your observations in your notebook. Also, write an equation (including state symbols) for the chemical
reaction you just saw. There are two products, one of which is a liquid.
Part D – Converting the Oxide to the Sulfate
Allow your beaker to cool and your precipitate to settle. Then decant and discard the supernatant, but
retain the precipitate in the beaker. To decant means to pour off, and the supernatant is the aqueous
layer above a solid precipitate. Try not to lose any of your copper compound. It is okay if your solid
still has a small amount of water in it. Then add 30 mL of 2.0 M H2SO4 and stir gently. Record your
observations in your notebook, and write an equation (including state symbols) for the chemical reaction
you just saw. There are two products, one of which is a liquid.
Part E – Converting the Sulfate to the Ammonia Complex
Add a couple of mL of your solution to a clean, dry test tube with a pipette (about 1 cm or less of
solution in a standard test tube from your drawer). SAVE the rest of your solution in the beaker and set
it aside just for the moment. Place the test tube in a small Erlenmeyer flask to support it, and take it to
the hood. Activate the fan, and then add a small squirt of concentrated ammonia. Concentrated
ammonia has a powerful smell, so try not to directly breathe the vapors. Add just enough ammonia to
observe a dramatic change in the solution, indicating the formation of an aqueous metal complex called
tetraammine copper (II), [Cu(NH3)4]+2. A metal complex, also known as a coordination complex, results
when a metal cation is surrounded and stabilized by a number of other molecules called ligands. The
ligands often have lone pairs, and form weak bonds (actually better described as intermolecular forces)
by donating some of their electron density to the positively charged metal ion. A kind of cage is created
by the ligands around the central metal cation. Note that the entire complex is usually water soluble, not
a solid precipitate. Because the bonds are weak, the complex can often be easily destroyed.
Return to your lab station and dump the contents of your test tube back into the rest of your copper
solution. The excess of H2SO4 present should neutralize the ammonia and break apart the complex.
Record all of your observations from these steps, and write an equation (including state symbols) for the
chemical reaction you just saw. It is easiest to write one reactant as simply the copper (II) ion and
ignore the sulfate, which is just a spectator ion in this reaction.
Part F – Converting the Sulfate to elemental Copper
Assuming you have done the lab carefully to this point, you should have all of the original copper in the
form of aqueous copper sulfate. We will use granulated zinc to get the elemental copper back.
Calculate the stoichiometric amount of zinc you will need for this reaction (remember that zinc usually
forms +2 ions in compounds), and then add 2.5 times the calculated mass. The excess is to help the
reaction go all the way to completion in a reasonable amount of time. Stir the mixture periodically, and
observe changes in the solution color – how do you know if the reaction is complete? If the solution still
shows evidence of aqueous copper (II) ions, add additional zinc until the reaction is complete (or almost
complete – it may be hard to drive the reaction all the way to completion). Record all of your
observations from these steps, and write two equations (including state symbols) for the chemical
reactions you just saw. One reaction involved the reduction of copper metal. The other reaction was a
“side reaction” which created some bubbles (of what?) as zinc reacted with the excess acid in solution.
You should now have solid copper as well as possibly some excess solid zinc in your beaker. We will
use 6.0 M HCl (an acid, but NOT an oxidizing agent!) to dissolve the excess zinc. About 12 mL of 6.0
M HCl ought to be enough – check for traces of zinc before moving on. Record all of your observations
from this step, and write an equation (including state symbols) for this final chemical reaction.
Finally, filter your solution to reclaim the isolated, solid copper. Wash your copper with 20 mL of
water, twice. Then wash it with 10 mL of ethanol if some is available. The washes will help to remove
impurities, and the ethanol will speed the drying process. Allow your sample to dry completely in the
oven if there is time left in class, or in the hood overnight if there is not. Find the mass of your final
copper sample, and compare this to the original mass.
Lab Report
Instead of writing a traditional lab report, I want you to summarize your learning about the lab on the
front of one piece of 11 x 17 inch (or similar sized) paper. Your report should consist of:
 Some sort of hand-drawn visual representation of the copper cycle, something akin to what
Edward Tufte calls a “supergraphic.” You can draw it as a circular cycle (sort of like the Kreb’s
Cycle) with new reactants and products flowing in and out. You can also use a flow chart or a
map or any other visual form that you think is suitable. Focus on showing all of the steps in a
clear and organized way! Annotate your diagram with words or sentences to make the process
clear. I expect that this diagram is your own work, not taken from an outside source or copied
from your lab partner!
 A list of all of the balanced chemical reactions (with state symbols!). You can incorporate these
directly into your diagram if there is a neat and well organized way to do this.
 A small, neatly organized data table showing the physical characteristics of each coppercontaining substance, including its color, appearance, and solubility.
 Information on the mass of copper you started and ended with, and a corresponding percentage
recovered.