The Copper Cycle Laboratory task

Most of the background material for this laboratory will be covered in greater detail in
the lecture course later in the semester. Here is some background information so you
will understand the chemistry behind the reactions you will perform.
Many aspects of our lives involve chemical reactions—from the batteries that power our cars and cell
phones to the thousands of processes occurring within our bodies. Most of these reactions can be
classified into one of three main types of chemical reactions: precipitation reactions, acid-base
neutralization reactions, and oxidation-reduction (also called “redox”) reactions.
Aqueous Solutions(aq)
Many reactions occur in an aqueous environment (i.e., in a solution where ions and compounds are
dissolved in water). When we indicate that a reactant or product has the physical state (aq), we mean
the substance is dissolved in water. When an ionic compound is in aqueous solution, the individual
ions are present in solution; for example, NaCl(aq) exists as Na+ and Cl– ions moving around in water.
Solubility Rules
Many ionic compounds are soluble—i.e., they dissolve in water. Others generally do not dissolve in
water and are considered insoluble. To determine if an ionic compound is soluble—i.e., will
dissolve—in water, we use the Solubility Rules:
Solubility Rules for Ionic Compounds in Water
The compound is SOLUBLE if it has:
1. Li+, Na+, K+, or NH4
+ ion (ALWAYS!)
2. C2H3O2
–, NO3
–, ClO4

3. Cl–, Br–, or I–, except compounds
with Ag+, Pb+2, and Hg2
+2 are
insoluble
4. SO4
2- except compounds with
Ag2SO4, CaSO4, SrSO4, BaSO4,
PbSO4, and Hg2SO4 are insoluble
The compound is INSOLUBLE if it has:
5. CO3
2–, CrO4
2–, PO4
3–, except compounds
with Li+, Na+, K+, NH4
+ are soluble
6. S2–, except compounds with Li+, Na+, K+,
NH4
+, Ca+2, Sr+2, Ba+2 are soluble
7. Hydroxide ion, OH–, except compounds
with Li+, Na+, K+, NH4
+ are soluble
The Solubility Rules indicate which compounds are soluble, and thus are represented as aqueous: e.g.,
KI(aq), BaCl2(aq), NaOH(aq), etc. The Solubility Rules also indicate which compounds are
insoluble—i.e., do not dissolve in water and remain as solids: e.g. BaSO4(s), AgCl(s), CaCO3(s), etc.
Double Replacement/Precipitation Reaction
For example, consider the reaction between aqueous lead(II) nitrate with aqueous potassium bromide,
as shown below:
Pb(NO3)2(aq) + KBr(aq) → PbBr2 + KNO3
Note that the chemical formulas for the products formed are based on their charges, not how
they appear on the reactant side of the chemical equation.
GCC CHM 151LL: The Copper Cycle © GCC, 2013 page 2 of 12
Based on Solubility Rules #4 and #1, we find that PbBr2 is insoluble and KNO3 is soluble. Thus, the
complete, balanced equation is:
Pb(NO3)2(aq) + 2 KBr(aq) → PbBr2(s) + 2 KNO3(aq)
We can cancel the spectator ions from the ionic equation and write the net ionic equation:
Pb2+(aq) + 2 Br -(aq) → PbBr2(s)
This reaction produces a cloudy mixture with small particles of the solid suspended in the solution.
When enough solid has formed, it will begin to settle at the bottom of the beaker. Thus, a clear solution
becoming cloudy when another solution is added is often taken as experimental evidence of a solid or
precipitate forming.
Acids and Bases
Acids can be defined as substances that produce hydronium ions (H3O+) when they are dissolved in
water. A hydronium ion is the product of a hydrogen ion that reacts with a water molecule: H+(aq) +
H2O(l) → H3O+(aq). A hydrated hydrogen ion (H+(aq)) is equivalent to an aqueous hydronium ion. The
two equations below both represent the ionization of hydrochloric acid, HCl(aq), but the second one
shows a particular water molecule explicitly.
HCl(aq) → H+(aq) + Cl–(aq)
HCl(aq) + H2O(l) → H3O+(aq) + Cl–(aq)
Acids are usually easy to recognize since their formulas start with H and contains nonmetal elements
other than H—e.g. HCl(aq), HNO3(aq), and H2SO4(aq) are all acids. Note that the physical state
aqueous, (aq), must be included to distinguish a compound that is acting like an acid from other forms
of a substance. For example, the formula “HCl” can also be used for hydrogen chloride gas, HCl(g),
so to indicate aqueous hydrochloric acid, one must specify HCl(aq).
One useful definition of bases is that bases are compounds that produce hydroxide ions (OH–) when
dissolved in water. The dissociation of sodium hydroxide, NaOH, is shown below. :
NaOH(s) → NaOH(aq) which is equivalent to Na+(aq) + OH–(aq)
Acid-Base Neutralization Reactions
In an acid-base neutralization reaction, a hydrogen ion-containing acid reacts with a hydroxidecontaining
base to produce water and a salt (an ionic compound):
HCl(aq) + NaOH(aq) → H2O(l) + NaCl(aq)
acid base water salt
Acids can react with bases, regardless of whether the salt is soluble or insoluble. There are other types
of acids and bases that can react without forming water.
If the reactants and products of an acid/base reaction are colorless and soluble, it is impossible to
monitor the progress of an acid-base reaction based solely on the appearance of the solutions. To help
us monitor acid-base reactions, we use litmus paper to determine if a solution is acidic or basic.
Litmus paper changes color depending on the presence of H+ or OH– ions in the substance being tested.
Blue litmus paper turns red in acidic solutions containing H+
ions, and red litmus paper turns blue in
basic solutions containing OH– ions.
GCC CHM 151LL: The Copper Cycle © GCC, 2013 page 3 of 12
Oxidation/Reduction Reactions
In an oxidation/reduction reaction, electrons are transferred from one reactant to the other. In the
simplest form of these reactions, single-displacement reactions (also called single-replacement
reactions), metal ions react with pure metals. If the reaction proceeds, the pure metal gives electrons to
the metal cation. This causes the pure metal to become a cation and the cation to become a pure metal.
The cation must always have an anion partner which is present either in an ionic solid or in a solution.
For example:
Mg(s) + 2 Ag+(aq) → 2 Ag(s) + Mg2+(aq)
metal cation metal cation
If the charge of an element is changing, that is a good indication that an oxidation/reduction reaction is
taking place. Later in the semester you will learn about oxidation numbers which are used to keep
track of more complicated oxidation/reduction reactions.
Step I: Chemistry
The different copper species obtained in each part is shown in Equation 1 below:
Cu(s)
Part I
Cu2+(aq)
Part II
Cu(OH)2(s)
Part III
CuO(s) Part IV Cu2+(aq)
Part V
Cu(s)
blue
I. Oxidizing Copper Metal with Concentrated Nitric Acid, HNO3(aq)
The first step involves transforming Cu metal to copper(II) ions, Cu2+, using concentrated nitric acid,
HNO3(aq). At the same time, the nitrate ions (NO3
–) undergo a series of reactions to form nitrogen
monoxide, NO. This product rapidly reacts with oxygen in the air to form NO2, a brown gas. The
presence of Cu2+(aq) makes the solution blue.
When the reaction mixture is diluted with water, the Cu2+ ions are hydrated (surrounded by water) to
form the octahedral complex ion, [Cu(H2O)6]2+, as shown below. Six water molecules (shown as red O
and white H atoms) are bonded to a Cu2+ ion (shown in gray as the central atom).
Cu2+(aq) + 6 H2O(l) → [Cu(H2O)6]2+(aq)
Figure 1
GCC CHM 151LL: The Copper Cycle © GCC, 2013 page 4 of 12
Step II: Chemistry
II. Precipitating Cu(OH)2(s) with NaOH(aq)
In Part II, two reactions are carried out by adding NaOH(aq). In the first reaction, the hydroxide ions
(OH–) from the NaOH(aq) neutralize the excess hydronium ions (H3O+) left over from the previous
part:
H3O+(aq) + OH–(aq) → 2 H2O(l)
Once all the H3O+ ions are neutralized, additional OH– ions react with the Cu2+ ion to form Cu(OH)2
precipitate. Once all the Cu2+ ions have reacted, no more precipitate forms. Adding more OH– ions
makes the solution basic, so it can turn red litmus paper blue. Figure 2 on the next page shows the
step-wise reaction of Cu2+ with NaOH.
Figure 2: Step-wise Illustration of the Precipitation of Cu(OH)2 in Part II – Remember:
[Cu(H2O)]2+ indicates the same substance as Cu2+.
OHOHOH-
[Cu(H2O)6]+2
[Cu(H2O)6]+2
Cu(OH)2
OHOHOHH
3O+
H3O+
H3O+
[Cu(H2O)6]+2
[Cu(H2O)6]+2
OHPart
II Step 1
H3O+
H3O+
H3O+
[Cu(H2O)6]+2
[Cu(H2O)6]+2
At end of Part I
OHOHOHOH-
neutralizes H3O+
Part II Step 2
OH- reacts with [Cu(H2O)6]+2 to form Cu(OH)2(s)
OHOH-
[Cu(H2O)6]+2
Cu(OH)2
Part II Step 3
Excess OH- makes
the solution basic
1st Beaker: At the end of Part I, hydrated copper complex, Cu2+
are present, making the solution blue,
and excess hydronium ions (H3O+) remain from the nitric acid used.
2nd Beaker: Adding NaOH(aq) to the blue solution results in the OH– ions neutralizing the H3O+ ions
to form water: H3O+(aq) + OH–(aq) → 2 H2O(l). The Na+ ions and resulting water
molecules are not shown.
3rd and 4th Beakers: Once all the H3O+ are neutralized, adding more NaOH(aq) results in the OH–
ions reacting with the Cu2+ to form the blue Cu(OH)2(s) precipitate shown at the
bottom of the beaker. Water molecules released from the complex ion are not
shown.
5th Beaker: When all of the Cu2+ ions have been converted to Cu(OH)2(s) precipitate, adding more
NaOH(aq) results in unreacted OH– ions in solution, which makes the solution basic. Red
litmus paper can be used to confirm the solution is basic. Note that the solution is no
longer blue since no Cu2+ ions are present in the solution.
GCC CHM 151LL: The Copper Cycle © GCC, 2013 page 5 of 12
Step III: Chemistry
III. Converting solid Cu(OH)2 to solid CuO
In Part III of the sequence, the reaction mixture is heated. This transforms the Cu(OH)2 precipitate to
CuO precipitate.
The CuO precipitate is separated from the solution, called the supernatant liquid, using a method
called gravity filtration. The mixture is filtered using a filter funnel, and the solid is collected on filter
paper. The supernatant liquid runs through the filter paper and collects in a beaker. This resulting
filtered solution is called the filtrate.
Step IV: Chemistry
IV. Dissolving CuO(s) with sulfuric acid, H2SO4(aq)
In Part IV, the CuO precipitate is dissolved using sulfuric acid, H2SO4(aq). This redox reaction returns
copper to its aqueous phase.
Step V: Chemistry
V. Reducing Cu2+ ions with Zinc Metal
In Part V, zinc metal (Zn) is added to the copper solution to convert the copper ions back to copper
metal, Cu(s). The resulting solution will contain colorless zinc ions, Zn2+(aq) and copper solid. Visible
evidence of this reaction is observed as bubbles of gas being released from the solution. (Since the H3O+
ions do not dissolve the Cu metal, the amount of copper yielded is not affected by excess acid.)
Identify the gas displaced from the acid in this reaction.
When the solution becomes colorless, all of the Cu2+ ions have been converted to Cu metal.
All of the excess Zn metal is also converted to Zn2+ ion by the excess H3O+ ions from the sulfuric acid,
H2SO4(aq),used to dissolve the CuO precipitate in Part IV.
Once all the Zn metal is dissolved, the Cu metal can be isolated by decanting, or pouring off, the
supernatant liquid. The Cu will then be rinsed, dried, and weighed as described in the procedure.
GCC CHM 151LL: The Copper Cycle © GCC, 2013 page 6 of 12
The Copper Cycle
In this experiment, you will carry out a series of reactions starting with copper metal. This will give you
practice handling chemical reagents and making observations. It is typical for scientists to observe
materials before they react, what happens during a reaction and how it looks when the reaction has come
to completion. The product of the final reaction will be copper metal and the percent copper that is
recovered will be calculated.
**Lab Notebook**
You should include one table that contains the mass of copper at the beginning and end of the experiment
along with % of copper recovered. This table should include:
• Mass of copper at the start of experiment (in Part I)
• Mass of copper + evaporating dish (from Part V)
• Mass of empty evaporating dish (from Part V)
• Mass of copper recovered (from Part V)
• Percent of copper recovered
Record observations for each of the steps (I-V) of the copper cycle in your lab book. Be sure to label
each step (I-V). The observations for each step should include:
• the appearance of the reactants before the reaction
• the appearance of the reactants during the reaction (for example, bubbles, flames, etc.)
• the appearance of the products after the reaction.
Your observations should include state(s) of matter, color, texture, smell, etc. where applicable. If your
observations are not detailed, you may not receive full credit.
One step also requires a specific chemical test using litmus paper to check for acidity. Be sure to also
record the results of these tests in your lab notebook.
**You will turn in worksheet pages 11-12 along with the duplicate pages from your lab notebook.
Step I: Procedure – Oxidizing Cu with concentrated nitric acid, HNO3(aq)
1. Place a sample of weighing paper in the balance. Tare the balance, so it reads 0.0000 g. Use
forceps to transfer about 0.35-0.40 g of Cu strips onto the weighing paper. Record the mass of the
Cu strips. Transfer the Cu strips into a clean 250-mL beaker labeled with one of your group
member’s initials. Record the appearance of the copper metal in your lab report.
CAUTION: Concentrated nitric acid is highly corrosive, so it can cause severe chemical
burns and damage clothing. Handle with care and avoid breathing the fumes. Any nitric acid
spilled on skin must be rinsed immediately with water for 15 minutes. Any acid spilled on
your work area must be neutralized then the entire area should be washed and dried.
CAUTION: Concentrated nitric acid reacts with copper metal to form brown toxic NO2 gas.
Leave the reaction beaker in the fume hood until all of the brown gas is vented in the hood.
GCC CHM 151LL: The Copper Cycle © GCC, 2013 page 7 of 12
2. In a fume hood, use a 10-mL graduated cylinder to carefully measure about 3 mL of
concentrated nitric acid, HNO3(aq). Slowly pour the nitric acid onto the Cu strips in the beaker,
swirling the beaker to maximize contact between the Cu and nitric acid until all of the solid Cu
has dissolved and the NO2 gas has escaped. Keep the reaction beaker in the hood until all the
toxic brown NO2 gas is gone, and keep your face away from the hood to avoid inhaling nitric
acid fumes and NO2 gas. Describe the reaction between HNO3 and the Cu metal in your lab
report.
3. Dilute the resulting solution with about 10 mL of deionized water. Describe the appearance of the
resulting solution containing Cu2+ in your data table.
Step II: Chemistry – Precipitating Cu(OH)2(s) with NaOH(aq)
In Part II, two reactions are carried out by adding NaOH(aq). In the first reaction, the hydroxide ions
(OH–) from the NaOH(aq) neutralize the excess hydronium ions (H3O+) left over from the previous
part.
Once all the H3O+ ions are neutralized, additional OH– ions react with the Cu2+ complex ion to form a
gelatinous blue Cu(OH)2 precipitate.
Once all the Cu2+ ions have reacted, no more precipitate forms. Adding more OH– ions makes the
solution basic, so it can turn red litmus paper blue. The picture sequence on the next page outlines
the step-by-step process that occurs during this step.
Step II: Procedure – Precipitating Cu(OH)2 with NaOH solution
CAUTION: Sodium hydroxide (NaOH) can easily damage eyes. It is corrosive and can cause
chemical burns and damage clothing. Any NaOH splashed into eyes or spilled on skin must
be rinsed immediately with water for 15 minutes. Any base spilled on your work area must
be neutralized then the entire area should be washed and dried.
1. While constantly stirring the Cu solution, slowly add 6M NaOH(aq) from the dropper bottles.
First, the OH– from the NaOH added will neutralize the excess acid left over from Part I.
2. Once all the acid is neutralized, additional OH– ions react with the Cu2+ to form Cu(OH)2(s), a
blue precipitate. Record what you observe in your lab report.
When adding more NaOH does not produce more precipitate, the solution can be tested to determine
if all the Cu2+ has been precipitated and additional OH– has made the solution basic. Use red litmus
paper to test if the solution is basic as follows. Without disturbing any precipitate, use a glass stir rod
to place a drop of solution (NOT the precipitate) on a piece of red litmus paper. If it turns blue, the
solution is basic. Stop adding NaOH when the solution turns red litmus paper blue. Describe your
litmus test in your lab report.
GCC CHM 151LL: The Copper Cycle © GCC, 2013 page 8 of 12
Check solution using red litmus paper (refer to
background handout). Continue adding base until
solution is basic.
Step-wise Illustration of the Precipitation
of Cu(OH)2 in Part II
1st Beaker: At the end of Part I Cu2+ ions are present, making the solution blue, and excess hydronium ions
(H3O+) remain from the nitric acid used.
2nd Beaker: Adding NaOH(aq) to the blue solution results in the OH– ions neutralizing the H3O+ ions to form
water: H3O+(aq) + OH–(aq) → 2 H2O(l). The Na+ ions are not shown.
3rd and 4th Beakers: Once all the H3O+ are neutralized, adding more NaOH(aq) results in the OH– ions
reacting with the Cu2+ to form the blue Cu(OH)2(s) precipitate shown at the bottom of
the beaker.
5th Beaker: When all of the Cu2+ ions have been converted to Cu(OH)2(s) precipitate, adding more
NaOH(aq) results in unreacted OH– ions in solution, which makes the solution basic. Red litmus
paper can be used to confirm the solution is basic. Note that the solution is no longer blue since
no Cu2+ ions are present in the solution. In reality, your solution may still appear blue because
of the dispersion of the Cu(OH)2 in the solution by mixing.
Step III: Procedure – Converting Cu(OH)2(s) to CuO(s)
1. Set up a ring stand as shown in the figure at the right. Set up a
ring clamp, and put a wire gauze on top of it. Above it, attach
another ring clamp with a diameter large enough to go around a
250-mL beaker. You are going to set your 250 mL beaker on the
lower ring and gauze. The upper clamp will hold the beaker in
place so it does not fall.
2. Add about 30-40 mL of deionized water to your reaction beaker
from Part II. Carefully place the beaker on the ring stand inside
the upper ring. CAUTION: Gently heat the beaker over a
medium flame. (Set the inner cone of the Bunsen burner flame
to a height of about 1.5 inch and the lower ring stand about 4 inches above the top of the Bunsen
burner). Constantly stir the solution with the glass end of the stirring rod until all the blue
precipitate turns black, and the solution is clear. If the solution starts to bump or boil, immediately
remove the beaker from the heat and let the solution cool slightly. Describe what happens to the
Cu(OH)2 precipitate upon heating in your lab report.
GCC CHM 151LL: The Copper Cycle © GCC, 2013 page 9 of 12
3. Allow the beaker and contents to cool. While they are cooling, set up the gravity filtration
apparatus. Obtain a second ring stand, and attach a ring clamp that is small enough to hold the
plastic funnel. Prepare the filter paper as shown below:
Step 1: Fold filter paper in half
and crease lightly.
Step 2: Fold again into
quarters.
Step 3: Lift up one layer of the filter paper, leaving 3 layers below. Place
the filter paper cone into the funnel. Press the edges of the filter
paper against the sides of the funnel, and wet the single-sided
edge with deionized water, so the paper sticks to the funnel.
Finally, place the plastic funnel in the small ring clamp, and place a 400-mL beaker beneath it to
collect the filtrate (the liquid that goes through the filter paper). The funnel’s stem should be just
inside the beaker to prevent splashing.
4. Use the markings on a clean 150-mL beaker to measure out about 15 mL of deionized water. Boil
the water on a hotplate to wash the precipitate in step 6.
5. When the 250-mL reaction beaker has cooled to room temperature, pour the CuO precipitate into
the funnel to filter the contents. Transfer the last traces of the solid from the reaction beaker into
the funnel, using a stream of deionized water.
6. Use a disposable pipet to wash the precipitate on the filter paper using the hot deionized water
heated in the 150-mL beaker. Allow each portion of hot water to drain through the filter paper into
the beaker below before adding the next portion.
7. Wash the 250-mL beaker, and rinse with deionized water. Replace the 400-mL beaker under the
filter funnel with the clean 250-mL beaker. Discard the filtrate (solution) collected in the 400-mL
beaker into the properly labeled waste container. Clean and dry the 400 mL beaker for use in
Part V. Keep the CuO solid in the filter paper for use in Part IV.
Step IV: Procedure – Dissolving CuO(s) with sulfuric acid, H2SO4(aq)
CAUTION: Sulfuric acid, H2SO4(aq), is corrosive, so it can cause severe chemical burns and
damage clothing. Handle with care and avoid breathing the fumes. Any sulfuric acid spilled
on skin must be rinsed immediately with water for 15 minutes. Any acid spilled on your work
area must be neutralized, and the entire area should be washed and dried.
1. Add about 10 mL of 3M sulfuric acid, H2SO4 (check the label before pouring), to the funnel to
dissolve the CuO precipitate. You may use a disposable pipet to direct your stream of sulfuric acid
to dissolve the precipitate. Allow the solution to drain through the funnel to the rinsed 250-mL
beaker. Repeat the procedure until all of the CuO solid dissolves to Cu2+ ions. Use as little of the
sulfuric acid as possible in this step. Describe the reaction between the CuO precipitate and the
H2SO4 in your lab report.
2. Use your water bottle to wash the last traces of solution from the empty funnel into the 250-mL
beaker which now contains the acid solution and aqueous Cu2+ ions. Keep this resulting solution
for use in the Part V.
GCC CHM 151LL: The Copper Cycle © GCC, 2013 page 10 of 12
Step V: Procedure – Reducing Cu2+ ions with Zn Metal
1. Use a weighing boat to measure and transfer about 1 g of Zn mesh to the Cu2+ solution in the 250-
mL beaker. Constantly stir the mixture with a glass stirring rod. Do not use any metal object that
will react with the acid to stir the solution. Continue stirring until all the Cu2+ ions have been
reduced to Cu metal as indicated by the solution becoming colorless. Dissolve any excess Zn by
adding a few drops of 3M H2SO4(aq). Describe your observations of the reduction of Cu2+ to Cu
metal.
2. Allow the Cu metal to settle at the bottom of the beaker. Without losing any of the solid, carefully
decant (pour off) as much of the supernatant liquid as possible into a 400-mL beaker. Some liquid
will remain in the first beaker with the Cu metal. Wash the Cu metal 3 times using 20-mL portions
of deionized water by stirring and then allowing the solid to re-settle. Again, decant the liquid into
the 400-mL beaker each time.
3. Weigh a clean, dry evaporating dish. Transfer the Cu metal and any remaining water into the
evaporating dish using a stream of deionized water. Decant most of the water from the evaporating
dish. Use a disposable pipet to remove as much remaining water from the evaporating dish without
losing any solid. You may use the corner of a paper towel to “wick” out any remaining water after
decanting.
4. Write one of your group members’ names on a folded piece of paper towel. Place your group’s
evaporating dish on the paper towel in the oven (between the hoods) to let the Cu completely dry.
Check it after about 10 minutes. If the copper pieces are loose then it is dry. If it appears black in
color, then the copper has been heated too much and has turned to copper (II) oxide.
5. When the Cu appears completely dry, let the evaporating dish cool to room temperature, and weigh
the evaporating dish with the Cu. Record the final mass in your lab report.
Wash and dry all of your glassware, equipment, and your lab area to prevent chemical
contamination and potential hazards.
Calculating Percent Copper Recovered
Theoretically, the mass of Cu recovered should be equal to the mass of the original Cu sample. The
overall efficiency of the experiment is measured by calculating the percentage of copper recovered:
100%
mass of initial sample
percent recovered = mass of final product ×
Ideally, the percent recovered should be close to 100%, which indicates that most (if not all) of the
copper was successfully transformed through all five parts of the experiment.
GCC CHM 151LL: The Copper Cycle © GCC, 2013 page 11 of 12
The Copper Cycle
Name: ________________________
Partner: _______________________
Section Number: ________________
***Turn in pages 11-12 along with your lab notebook copies***
For each equation shown below (1 for each step of the copper cycle), indicate
A. the products of the equation, include phases of all products,
B. balance the equation, and
C. tell what type of equation it is: combination, decomposition, single replacement, double
replacement (precipitation), or acid/base neutralization (Hint: Many metal oxides can function as
bases in aqueous solutions).
Step I. ___ Cu (s) + ____HNO3 (aq) à ____NO2 (g) + ____H2O(l) + __ Cu(NO3)2 (aq)
Type of reaction: This is a complicated reaction. NO+ in the solution gets reduced to NO gas, which
bubbles out of the solution and reacts with O2 in the air to form NO2. Balance the above equation even
though it does not exactly describe the reaction you perform. The reaction has characteristics of
decomposition, acid-base and oxidation-reduction reactions (single replacement).
Step II. ___Cu(NO3)2 (aq) + ____NaOH (aq) à
type of reaction: ______________________________________
Step III. ____Cu(OH)2 (s) ⎯⎯→ Δ ______________ + ____H2O (l)
type of reaction: ______________________________________
Step IV. ____CuO (s) +____ H2SO4 (aq) à _____________ + _____________
type of reaction: ______________________________________
Step V. ____CuSO4 (aq) + ____Zn (s) à _____________ + ______________
type of reaction: ______________________________________
GCC CHM 151LL: The Copper Cycle © GCC, 2013 page 12 of 12
POST-LAB QUESTIONS:
1) Percent Copper Recovered
A student performing this experiment started with a 0.3769 g sample of copper turnings, which was
dissolved in concentrated nitric acid. After completing the series of reactions, the student isolated
0.3492 g of copper. Calculate the percent copper recovered by the student. Show your work.
2) Indicate whether the following procedural errors would result in an incorrectly high or incorrectly
low percent recovery. Circle and explain your answer.
a. The solution was not basic before being heated in Part III.
High Low
b. In Part III, the solution was poured into the funnel until it went above the top of the filter paper, and
some black solid was disposed of with the filtrate.
High Low
c. The solution decanted in Part V was slightly blue in color.
High Low
d. After all the copper metal was obtained in Part V, it took too long for the excess Zn granules to
dissolve, so a student added concentrated nitric acid to the solution, resulting in a brown gas.
High Low
e. In Part V, the copper metal was weighed while it was still damp.
High Low

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