Cu Later! Lab
Jeesun Park
Partner: Dale Walker
B1, IB Chem I
McClung
Objective: Observe some chemical reactions of copper and its compounds while
performing the lab appropriately as to retain the copper as much as possible
Materials:
250mL beaker
400mL beaker
25mL graduated cylinder
100mL graduated cylinder
rubber policeman
stirring rod
watch glass
wash bottle
masking tape
permanent marker
electronic balance
heater
Bunsen burner
Steel wool
Zinc strip
Copper
6M nitric acid
6M sodium hydroxide
6M hydrochloric acid
1M sulfuric acid
Procedures:
Day One
1. Mark a clean, dry 250 mL beaker with a masking tape label indicating you and your
partner’s names. Mass the beaker and record it in your data table. Obtain about 1.5g
of copper and remass the beaker and the copper together to the nearest 0.01g.
Record the mass.
2. In the fume hood, add 20 mL of 6M nitric acid, HNO₃(aq). If necessary, warm the
mixture on a hot plate. Leave the beaker in the fume hood until all of the copper
metal and brown fumes have disappeared. Avoid breathing fumes – they are
nitrogen dioxide, NO₂. Time the reaction and record all observations. The step
takes 5 to 10 minutes.
3. Fill a 400 mL beaker about 1/3 full with cold tap water and ice. Place the smaller
beaker from the fume hood in the ice bath to cool it.
4. Rinse the graduated cylinder. Measure 20mL of 6M sodium hydroxide, NaOH(aq),
into a small beaker. Add it slowly, 2 to 3 mL at a time, into the blue solution in the
ice bath. Stir constantly. Record observations.
5. Add about 100mL of distilled water to the mixture resulting from step 4. Heat the
beaker and its contents gently over a Bunsen burner until it boils. Continue heating
with a reduced flame until all of the solid changes to the same dark color. Allow the
mixture to cool. Decant the liquid, being careful not to lose and solid. Flush liquid
down the sink with plenty of running water. Wash the solid in the beaker with
another 100 mL of distilled water. Wash the solid in the beaker with another 100
mL of distilled water. Allow beaker to stand in your lab drawer.
Day Two
6. Without disturbing the solid on the bottom of the beaker, carefully decant the liquid.
Flush it down the sink with lots of water. Add 40 mL of 1M sulfuric acid, H₂SO₄
(aq). Stir. Record observations, both before and after adding the acid.
7. Use steel wool to scratch off the surface of a Zn strip. Mass the zinc and record to
the nearest 0.01g. Place the zinc in the solution from Step 6, and warm slowly.
When the solution boils, remove it from the heat and allow it to tand and cool.
Don’t use ice here, just room temperature cooling. Record observations, waiting
until the solution becomes colorless and transparent.
8. Wash the accumulated copper solid from the zinc strip into the beaker, using a wash
bottle filled with water. Dry and record the mass of the zinc. Use the same balace as
you used throughout the experiment.
9. Wash the copper and decant, being careful to retain all of the copper in the beaker.
Add about 15mL of 6M HCl(aq) and warm gently to dissolve any remaining piece
of zinc. Wash the copper with distilled water several times, decanting carefully each
time so as not to lose any copper.
10. Mass a clean, dray watch glass or glass plate. Record. Carefully scrape out all the
Cu onto the watch glass, spread it out, and place in your lab drawer to dry.
Day Three
11. Re-mass the glass and determine the mass of Cu recovered.
Data (the original copy attached):
Data Table 1: Masses of Chemicals and Items in the Copper Compound Reactions
Item
Mass (+ 0.01g)
250 mL beaker
114.62
250 mL beaker & copper
116.17
Original Copper
1.55
Zinc strip before reaction
7.08
4.51
Zinc after reaction with CuSO₄
Watch glass
29.18
Watch glass & Copper
30.40
Copper after reactions
1.22
Qualitative Data:
When the original copper solid was put in the transparent nitric acid solution, the
solution turned to blue, a little bit of bubbles in the solution, and some brown gas was
observed near the surface of the solution. When the resulting solution, copper(II) nitrate,
was mixed with the colorless sodium hydroxide solution, the originally blue solution
produced some pale blue colored precipitate, which soon sank to the bottom of the beaker.
When the solution was heated by the Bunsen burner, after the previous reaction was
complete, the pale blue precipitate disappeared and the dark brown precipitate appeared.
On the day two, after the liquid was decanted, nothing was observed when the
sulfuric acid was added to the copper(II) oxide solution. When the solution was heated with
the zinc strip in it, the solution turned to green at first, then to blue and became colorless.
Some bubbles were seen on around the zinc strip in the solution, and soon the rust formed
around the zinc. In the final step, nothing was observed when the hydrochloric acid was
added to dissolve any remaining zinc.
Data Processing:
1. The mass of copper obtained before reactions (in grams)
= (the mass of 250mL beaker & copper) – (the mass of 250mL beaker)
= 116.17 – 114.62
= 1.55
2. The mass of copper after the reactions were completed (in grams)
= (the mass of watch glass & copper) – (the mass of the watch glass)
= 30.40 – 29.18
= 1.22
3. The mass of zinc lost after the reaction with copper(II) sulfate solution (in grams)
= (the original mass of the zinc strip) – (the mass of zinc after the reaction)
= 7.08 – 4.51
= 2.57
4. The percent recovery of Copper metal:
(the mass of recovered copper) * 100 = 1.22g * 100 = 78.71 %
(the mass of original copper)
1.55g
5. The number of moles of copper recovered
1.22g Cu 1 mol Cu
= .0192 moles Cu
63.55 g
6. The number of moles of Zn that was expected to react with Copper(II) sulfate
solution
= 1.55g Cu 1mol Cu 1mol Zn
= .0244 moles Zn
63.55g
1mol Cu
7. The number of moles of Zn that actually reacted with
= 2.57g Zn 1mol Zn = .0393 moles Zn
65.39g
8. The difference between the expected and the actual amount of Zn
= (expected) – (actual) = (.0393 moles) – (.0244 moles) = .0149 moles Zn
Conclusions and Analysis:
The objective of this lab was to observe some chemical reactions of copper
compounds and to do it appropriately so that the original amount of copper could be
retained as much as possible. The original mass of the copper was 1.55g and the copper
after all the reactions was 1.22g(see Data Table 1). We’ve lost .33g of copper in the
experiment, and this loss of copper means that some errors were made in the lab.
In the first reaction when the original copper solid reacted with nitric acid, it was
clear that a new substance was formed as a result because the transparent solutions
turned to blue as they were mixed, which was a clear indication of a chemical reaction.
Also the copper solid disappeared. The chemical formula for this reaction is:
3Cu(s) + 6HNO₃(aq)
copper
nitric acid

3Cu(NO₃)₂(aq) + 4H₂O(l)
copper(II) nitrate
water
+ 2NO(g)
nitrogen oxide
I think what caused the color change was the combining of copper and nitrate because the
solution’s color is blue. Water and nitric oxide, other resulting solutions, are transparent so
they would not have changed the color of the resulting solution. Some bubbles appeared,
which was the evidence of the nitrogen oxide gas production. Besides the color change,
there was some brown gas observed near the surface of the solution, which was the product
of nitrogen oxide gas and the oxygen in the air, nitric dioxide. The chemical equation for
this reaction is:
2NO(g)
+
nitrogen oxide
O₂(g)
oxygen

2NO₂(g)
nitrogen dioxide
When the sodium hydroxide solution was added to the resulting solution from the
previous steps, which was copper(II) nitrate, the blue solution produced some pale blue
precipitate. This precipitate was copper(II) hydroxide, which is indicated in the following
equation:
Cu(NO₃)₂(aq)
Copper(II) nitrate
+
2NaOH(aq)

Cu(OH)₂
+
sodium hydroxide
copper(II) hydroxide
2NaNO₃(aq)
sodium nitrate
As the solid copper hydroxide was being produced, the blue color of the solution changed
to transparent largely because the blue color cooper(II) nitrate solution was reacting with
sodium hydroxide in the double displacement reaction. The resulting solution, sodium
nitrate, was transparent and thus the color of the produced solution was transparent. When
this product is heated, copper(II) oxide and water are produced as in the following
equation:
Cu(OH)₂(s)
Copper(II) hydroxide

heat
CuO(s)
copper(II) oxide
+
H₂O(l)
water
As the solution is heated, the blue precipitate disappears and the solution turned into dark
brown color. This was not the liquid solution’s color, but that of copper(II) oxide solid that
was produced along with water. This indicates that the chemical reaction had resulted in
new products.
In the next chemical reaction, the sulfuric acid[1M] was added to the copper(II)
oxide solution. The chemical equation for this reaction is:
CuO(s)
+ H₂SO₄ (aq)  CuSO₄(aq) +
copper(II) oxide sulfuric acid
copper(II) sulfate
H₂O(l)
water
In this reaction, no change in particular was observed.
When the zinc strip was added to the solution and was heated, many changes
occurred. First, the dark brown solution turned into transparent green, and then to
transparent blue. Also some bubbles and the forming of the rust on the surface of the zinc
strip in the solution was seen at the same time. The chemical equations are:
Zn(s)
Zinc
+
Zn(s)
Zinc
+
CuSO₄(aq)

copper(II) sulfate
H₂SO₄(aq)
sulfuric acid

Cu(s)
copper
+
ZnSO₄(aq)
zinc sulfate
ZnSO₄(aq) + H₂(g)
zinc sulfate
hydrogen gas
The rust on the zinc strip was the reproduced copper from the previous solution, copper(II)
oxide. Some zinc was lost in this reaction by 2.57g (see Data Processing) because some
formed zinc sulfate aqueous solution. The second chemical equation above is to show the
reaction in case some sulfuric acid remained from the previous step, which indicates no
precipitate or color change but some hydrogen gas. The production of the hydrogen gas
probably was the bubbling in the solution, and some of them could have been trapped
around the zinc strip between the rust, which might be why it looked like the zinc strip was
producing gas. Thus the bubbling prove that some sulfuric acid was present when the zinc
strip was put into the solution, which was not accounted for in the Zinc and copper(II)
sulfate reaction. Since the actual take up of Zn was greater than expected (see Processing
Data) by .0149 moles, it also supports that some sulfuric acid was present in the copper(II)
sulfate solution.
The final mass of copper turned out to be 1.22g, which was .33g less than the
original copper of 1.55g. The percent recovery of the copper turned out to be 78.71% (see
Processing Data: 4). The number of moles of the recovered copper is .0192 moles (see
Processing Data: 5)
Evaluations on Procedures and Results and Improving Investigation:
Overall, the procedures were written well in a clear language to avoid any
confusion. Also, it instructed the experimenter when to record qualitative data and
quantitative data, indicating the uncertainties also. Appropriate lab terms were used, such as
“colorless”, not “clear”, when referring to the transparent solutions. Also to avoid any
systematic errors that might result arise from using different kinds of balance during the lab,
the procedures recommended that the same kind of electronic balance should be used
throughout the lab.
There is one evidence regarding a random error, which we do not have control over.
We can never be sure if a chemical reaction is complete or not up to 100%, and the bubbles
and the more loss of Zn than expected in its reaction with Copper(II) sulfate indicate that
the previous chemical reaction between copper(II) oxide and the sulfuric acid was not at all
complete. This remaining sulfuric acid probably affected the lab by reacting with Zn on the
surface of Zn strip. The point is, the surface area of Zn strip is very limited, and thus more
sulfuric acid reacts with Zn, less Zn is left to react with copper(II) sulfate. This could result
in incomplete reaction, which would reduce the amount of copper to be recovered at the
end. This, in turn, will reduce the percent recovered. Just the fact that the chemical
reactions are not always complete leaves the room for errors because the procedures were
written upon the assumption that all of the reactions would be complete.
Another error involving the possibility of an incomplete chemical reaction would
be when the copper(II) hydroxide is heated. This was to decompose the compound into
copper(II) oxide and water. The procedures instruct that the experimenter should heat it
until it boils. However, there is no method or time for the experimenter to test if the
decomposition was complete after the heating. It might be the case that the copper
compound did not decompose completely, and thus resulted in less amount of copper(II)
oxide. This will not necessarily reduce the amount of copper to be available for the next
chemical reaction, because of the copper left out does not disappear into the thin air, but it
leaves an argument that there is no way to figure out if the decomposition was complete.
One systematic error could have been the use of rubber policeman. When we used
the rubber policeman to get Cu rust off the Zn strip physically, it was impossible to get
every little tiny particles of Cu off the strip. This would affect the amount of Cu to be
recovered at the end because if the rubber policeman is not good enough to get all the rust
into the beaker, the rust remaining on Zn will be forever lost and will not be added to the
recovered copper’s mass. This will reduce the percent recovered of copper. This is an
inherent problem because it would cost too much to get a intricate system to regain all of
the copper rust on Zn strip in a high school setting. We can always get close to being
perfect, but in reality it is very hard to achieve.
There seems to be no experimenter error. My data matches with my partners so we
have communicated properly, and we’ve written down any and all qualitative observations
and took them into consideration.
Personal Statement
I confirm that this lab report is completely my own work and that no other student’s work
was turned in that was (even in part) taken from this.
Signed:__________________________
Date:____________________________