University of Phoenix Copper Cycle Report Lab Report

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Reports consist of the following sections:-Report Sheets – the completed Data and Report Sheets cut out of your lab manual.Calculations – If there are calculations, you must clearly and logically show all of the calculations to getfull credit.Discussion – Explain your results, referring to the chemistry theory where appropriate. Did the dataconform to what was expected? What does it show in relation to the experiment? If the results are notwhat was expected explain why. Possible sources of error (estimate experimental error in numericalresults). Conclusion – briefly state any final result(s), or conclusions that can be drawn from the experiment.References – any sources consulted for the writing of the report. These should be referenced (cited) inthe body of the text. 1
• To observe a sequence of copper reactions that form a cycle.
• To gain skill in recording observations that indicates evidences of chemical change for those reactions.
• To interpret the ‘cycle of copper reactions’ in terms of chemical equations using the simple classification scheme
of grouping chemical equations according to precipitation, acid-base, decomposition and redox reactions.
• To practice quantitative laboratory techniques by determining the percentage yield of copper recovered at the
end of the reaction cycle.
To a chemistry student, one of the most important aspects of a laboratory experiment is the array of physical
and chemical changes encountered. This experiment provides an interesting experience on the reactivity of copper
in aqueous solutions. Here, you will be asked to carry out a ‘cycle of copper’ reactions. It is a ‘cycle’ because the
sequence begins and ends with copper metal. Because no copper is added or removed between the initial and final
steps, and because the reaction is expected to go to completion at each step, you should be able to quantitatively
recover all of the copper you started with provided you are careful and skillful in performing each step of the cycle.
As a multi-step process, it is important that each step in the process is optimized because the percent recovery of
the copper metal in the overall process depends on the percent recovery of the products in each step. Figure 1
shows in an abbreviated form the reaction steps in this cycle.
Figure 1 shows the flow-chart diagram of all the steps in the ‘cycle of copper reactions’. Typically, only the key
reagent(s) is/are presented in a flow-chart diagram. In addition, chemists often times use shorthand abbreviations.
One common abbreviation is the use of Δ for heating a reaction mixture.
It is important to realize that these equations are results of a large number of experiments and it’s easy to lose
sight of this if you just look at them on paper. It is in fact a formidable task to attempt to learn isolated bits of
information that are not reinforced by your personal experience. This is one of the reasons why it is important to
have laboratory experience.
Figure 1. Cycle of Copper Reactions
As you already know, chemistry is pre-eminently an experimental science; therefore it is important that you watch
closely and record what you see as chemical changes happen. Each observation should serve as clue in your mind
to formulate abstract/concrete concepts for interpretation (e.g. chemical formula/equation). Observations and facts
that have not been assimilated into some coherent scheme of interpretation are relatively useless. It would be like
memorizing daily weather reports when you have no knowledge of meteorology.
Chemists look for relationships, trends or patterns of regularity when they organize their observations.
Interpretation is typically carried out by establishing chemical reactions associated in order to explain the said
observation. As you already know, the periodic table, which groups the elements into chemical families, is a product
of this kind of thinking. Each element bears a strong resemblance to other members of the same chemical family
but also has its own unique identity and chemistry. In a similar fashion, it is useful to classify reactions into different
types. In this experiment, we will re-visit aqueous solutions chemistry involving precipitation (ion-combination),
acidbase (proton-transfer) and redox (electron-transfer) reaction(s) that you already learned from CHEM1. These
reactions can be used to interpret the observations in the ‘cycle of copper reactions’.
A. Precipitation
In solution, the combination of positively-charged ions with negatively-charged ions to form an insoluble neutral
compound which precipitates from solution is known as precipitation reaction. For instance, you will recall that
sodium chloride (NaCl) and silver nitrate (AgNO3) salts are generally soluble in water. However, mixing an aqueous
sodium chloride solution, NaCl(aq) with an aqueous silver nitrate solution, AgNO3(aq) results to formation of two
products: (1) a precipitate believed to be silver chloride, AgCl and (2) a water-soluble sodium nitrate, NaNO3
compound. This chemical reaction is expressed in equation 1 which can be simplified further to its net ionic form as
shown in equation 2.
NaCl(aq) + AgNO3(aq) ⟶ AgCl(s) + NaNO3(aq) (1)
Ag+(aq) + Cl-(aq) ⟶ AgCl(s) (2)
Here, the silver chloride precipitate can be easily separated from the solution containing the soluble sodium
nitrate compound. If desired, sodium nitrate salt could also be recovered by evaporating the water from the solution.
B. Acid-Base
In CHEM1, you learned that a hydrogen atom contains one proton and one electron. When it ionizes to form
hydrogen cation (H+), it loses the electron. Therefore, hydrogen cations are also called protons.
An acid can be formally defined as a substance that donates hydrogen cation (proton donor) while a base is a
substance that accepts hydrogen cation (proton acceptor). This definition of acids and bases is also called the
Bronsted-Lowry definition.
In aqueous medium, protons produced by acids are abstracted by the oxygen of water and therefore form the
hydronium ion complex, H3O+. Overall, a substance serves as a strong acid if it completely (100%) dissociates to
produce H+ ion in solution. Therefore strong acids produce higher H+ ion concentration in solution and lower pH.
Acids that partially dissociate to form H+ ions are called weak acids.
A good example of a strong acid in water is hydrogen chloride, HCl gas. Also called hydrochloric acid in its
dissociated form (e.g. aqueous medium), HCl(aq) dissociates 100% into H+ and Cl- in water. As a learner, it is
important to know that the acid (or base) strength depends on the solvent medium. Thus, it HCl may be a strong
acid in water but weak in non-aqueous solvents such as acetone and methanol. The dissociation of hydrogen
chloride in water is shown in equation 3 while that in methanol by equation 4.
HCl(aq) + H2O(l) ⟶ H3O+(aq) + Cl-(aq) (3)
HCl(aq) + CH3OH(l) ⇌ CH3OH2+(MeOH) + Cl-(MeOH) (4)
An example of a weak acid in water is acetic acid, CH3COOH. In acetic acid only one hydrogen atom is
considered acidic. This hydrogen is the one directly attached to oxygen. Its partial ionization in water is given by
equation 5.
CH3COOH(aq) + H2O(l) ⇌ H3O+(aq) + CH3COO-(aq) (5)
On the other hand, a substance serves as a strong base if it binds (or associates) strongly with H+ ions. As a
consequence, bases lower the concentration free H+ ions in solution and give higher pH level. Weak bases bind
weakly to H+ ions.
The Bronsted-Lowry definition may not seem obvious in explaining the basicity of sodium hydroxide. But it follows
that from the Arrhenius definition that NaOH is considered as a strong base in aqueous medium because it ionizes
100% to form HO- and Na+ ions. By adopting the Bronsted-Lowry definition, the hydroxide is a base because it
serves as a proton acceptor in the neutralization reaction (6).
HO-(aq) + H3O+(aq) ⟶ 2H2O(l) (6)
Bronsted-Lowry definition is very useful in explaining the basicity of ammonia, NH3. At first, it looks like it will be
a good acid due to the presence of three hydrogens covalently bonded to an electronegative atom (N). However,
the electron pair around the nitrogen center can abstract a proton from a neighboring acid making ammonia a proton
acceptor rather than a proton donor. Though considered base, ammonia is a weak base.
NH3(aq) + HCl(aq) ⟶ NH4+(aq) + Cl-(aq) (7)
In general, an acid reacts with a base. Strong acids that react with strong bases produce neutral salts while weak
acids that react weak bases to produce weak salts whose strength is relative to their starting materials.
Reaction of Strong Acid with Strong Base:
HCl(aq) + NaOH(aq) ⟶ NaCl(aq) + H2O(l) (8) Reaction
of Weak Acid with Weak Base:
CH3COOH(aq) + NH3(aq) ⟶ CH3COO-(aq) + NH4+(aq) (9)
C. Reduction-Oxidation (Redox)
Reduction-oxidation (also called redox) reactions are reactions that involve the shift or transfer of electrons from
one species to another. Reduction is defined as the process at which a species gains an electron while oxidation is
the process at which a species loses an electron. Consider the reaction between magnesium metal and hydrochloric
acid as illustrated in Equation 10.
Mg(s) + 2HCl(aq) ⟶ MgCl2(aq) + H2(g) (10)
Here, each magnesium atom gives up two electrons to two hydrogen ions forming one magnesium ion and one
hydrogen gas.
Sometimes, the transfer of electrons between atoms is less obvious, as in the reaction:
2 SO2(g) + O2(g) ⟶ 2 SO3(g) (11)
Here the reactants and products are all gases, and no ions are formed. To be able to classify this as a redox
reaction, we use the concept of assigning oxidation numbers (also called oxidation states) to each atom in the
compound. A simple set of rules defines the procedure for assigning the oxidation number (See Tro, Sections 18.2
& 4.9).
In sulfur dioxide, SO2, and sulfur trioxide, SO3, we imagine that all of the electrons in the S-O bonds are assigned
to the O atoms, giving each oxygen atom a full valence shell. This formally gives each oxygen atom a net charge of
-2 which is its oxidation number. So, if oxygen is assigned an oxidation number of -2, the sulfur in SO2 must have
oxidation number of +4 and the oxidation number of sulfur in SO3 would be +6, since the sum of the oxidation
numbers on all the atoms must sum up to the net charge of the molecule (zero in both cases. The oxidation number
of oxygen atoms in O2(g) is assigned as zero (See Tro, Sections 18.2 & 4.9).
Once we have assigned oxidation numbers to each element in the chemical reaction, it can be deduced that the
oxidation number of sulfur goes from +4 to +6 while the oxidation number of oxygen atoms decreases from zero (in
O2) to -2 (in SO3).
From this viewpoint, the change in oxidation number is formally equivalent to the number of electrons that
transferred from sulfur to oxygen (in this case it is 2). Thus we say that the sulfur atom underwent oxidation (oxidized)
while the oxygen atoms in O2 underwent reduction (reduced).
It is important to note, however, that the oxidation numbers we assign do not necessarily represent the real
distribution of electronic charge in the molecule. By assigning the oxidation number according to fixed rules, we
have artificially assigned integer changes in oxidation number to particular atoms (in this case, sulfur and oxygen),
but the changes in electron density around the sulfur and oxygen atoms may not be as large as implied by the
assigned oxidation numbers. Nevertheless, it is reasonable to suppose that the sulfur atom in SO 3 is more positive
than the sulfur atom in SO2 because the added oxygen atom would tend to draw electrons away from the sulfur
D. Decomposition
Although many chemical reactions can be classified into one of the three reaction types described earlier, it is
possible to find examples of chemical reactions that do not neatly fit these types. For instance, when calcium
carbonate is heated at higher temperatures (say, flame from a Bunsen burner), it breaks down into simpler
CaCO3(s) ⟶ CaO(s) + CO2(g) (12)
Chemically, it will be best to describe this reaction in terms of bond-breaking process (C-O bond breakage, in
this case). Decomposition reaction involves the breaking down of a substance into simpler species under an external
stress such as heat, pressure and electricity. Heat is the most common source of decomposition. The temperature
at which this heat is supplied is called decomposition temperature.
Does a comprehensive classification scheme exist?
If you search, you would find examples of other reactions that do not fit into these four categories. Indeed, it is
fair to say that there is NO completely comprehensive scheme that would accommodate all known chemical
reactions described in your general chemistry text will fit into this simple scheme.
‘Cycle of Copper Reactions’
Now, we go back to the system that we are interested in this experiment. As you carry out each step of the cycle
of copper reactions, think about what is happening in each reaction in order to predict the products of each step.
Then, you are expected to fit each step into one of the four categories we described.
For additional information on the topics covered in the Background Information, see Sections 18.2, 15.3, 15.11 and
16.5) of your lecture text; Chemistry: A Molecular Approach, 3rd edition, by Nivaldo Tro, Pearson, 2014.
Special Supplies:
Hot plate
Item to check-out
Evaporating dish
copper, Cu wire (22 gauge)
Zn mesh (20 mesh)
Conc. (16M) nitric Acid, HNO3
3M sodium Hydroxide, NaOH
6M sulfuric acid, H2SO4 6M
hydrochloric acid, HCl
methanol, CH3OH
Cu mesh scrubber
Safety Precautions
• Concentrated nitric acid, HNO3, is hazardous. It produces severe burns on the skin, and the vapor is a lung
irritant. When you handle it, you should use a fume hood WHILE WEARING SAFETY GLASSES (as always)
and rubber or polyvinyl chloride gloves (if available). A polyethylene squeeze pipet can be useful for transferring
the HNO3 from a small beaker to your 10-mL graduate cylinder. Rinse your hands with tap water after handling
• The dissolution of the copper wire with concentrated HNO3 should be carried out in a fume hood.
• The brown NO2 gas that evolves is toxic and must be avoided.
• NaOH solutions are corrosive to the skin and especially dangerous if splashed into the eyes – WEAR SAFETY
• Methanol, CH3OH and Acetone, CH3COCH3 are flammable and their vapors are toxic. Use them in the hood to
avoid breathing the vapor, and keep them away from all open flames.
General Notes & Tips for Success
1) This experiment will be performed individually.
2) If using copper wire, make sure to clean the wire with a Cu mesh scrubber provided by the instructor. Handle
the cut Cu wire with forceps. Avoid skin contact with copper wire to prevent contamination by sweat and oil from
your skin. Use forceps in handling copper wire.
3) In the balance room, solid substance will have its own dedicated spatula. Make sure you use the proper
spatula for each solid! Mixing up the spatulas could result in contamination of your samples and cause poor
4) Try to avoid spilling any solids on the balances. In order to do this, always remove your sample container from
the balance when adding or removing material. If any spills do happen, clean off the balance pan immediately
with a brush.
Part 1. Reaction of Cu metal with conc. nitric acid, HNO 3 solution
Cut a pure copper wire that weighs about 0.5g (about a 18-cm length of 22-AWG copper wire). If it is not bright
and shiny, it is likely that the wire is oxidized after prolonged exposure to air thus may be contaminated. To address
this issue, clean it with the Cu mesh scrubber and wipe it with a Kim-Wipe tissue. For better handling of the copper
wire, read section on ‘General Notes & Tips for Success’. Weigh to the nearest milligram. Record its mass in your
reports sheet.
Coil the wire into a flat spiral. If you end up touching the wire, rinse with acetone and wipe it with Kim-Wipe
tissue. Place the coil wire in the bottom of a 250-mL beaker, and – in the fume hood – add 4.0 mL of concentrated
(16 M) nitric acid, HNO3. Observe and record a description of what you see during the course of the reaction in your
report sheet.
Now, swirl the solution around in the beaker until the copper metal has completely reacted (dissolved). What is
in the solution after the reaction is complete? Suggest and write a balanced chemical equation associated with this
reaction in your report sheet.
After the copper has completely reacted (after approximately 40 mins), add deionized water until the beaker is
about half-full. Reserve the solution for Part 2. Steps 2 through 4 can be conducted at your lab bench.
Part 2. Reaction of Cu(NO3)2 with NaOH solution
Collect the solution from Part 1. Stir the solution with a glass stirring rod. Add 30 mL of 3.0 M NaOH solution.
Make sure that you constantly stir the solution while pouring the base into the Cu(NO 3)2 solution. Observe and
record a description of what you see during the course of the reaction in your report sheet. What is the identity of
the precipitate that is formed? Suggest and write a balanced chemical equation associated with this reaction in your
report sheet. Reserve the suspension mixture for Part 3.
Part 3. Decomposition of Cu(OH)2 with heat, Δ
Collect the suspension mixture from Part 2. Stir the mixture gently with a glass rod. Gently heat the solution just
barely to boiling using a Bunsen burner. It is important that you continuously stir the mixture GENTLY to prevent
‘bumping’. Bumping is a phenomenon caused by formation of a large steam bubble in a locally overheated area. If
the solution bumps, you may lose some product, so do not neglect the stirring. Observe and record what you see
during the process in your report sheet.
When the transformation is complete, remove the beaker from the heat source. Do not place the beaker directly
on the lab bench. You can use your blue towel or wire gauze to insulate the heat from the hot beaker to the bench
top. This will give you time to check the progress of your reaction. You will know that the reaction is complete when
the supernatant liquid is clear and colorless on top of fine brown-black product. If the supernatant liquid has any
color or opaque, you will need to continue heating until it is clear and colorless.
After the reaction is complete, return the beaker back on the heat source and continue stirring for a minute or so
and remove the beaker from the heat source and allow the product precipitate to settle until you see a clear
borderline separation of the solid residue from the liquid phase.
Decant (pour off) the supernatant liquid (also known as decantate) into a separate container and dispose it
properly. Be careful not to lose any residual solids into the decantate. To the residual solids, add about 200 mL of
hot deionized water and allow it to settle again. Decant one more time. What is the identity of the precipitate? What
was removed from the precipitate by washing and decantation process? Observe and record what you see in your
report sheet. Also, suggest and write a balanced chemical equation associated with this reaction in your report
sheet. Collect the residue and proceed to Part 4.
Part 4. Reaction of CuO with 6.0 M sulfuric acid, H2SO4 solution
Obtain the precipitate that was collected in Part 3 of the experiment. Add 15 mL of 6.0 M H2SO4 into the beaker.
Stir the mixture constantly during the addition of the acid. What is in the solution? Observe and record what you see
in your report sheet. Suggest and write a balanced chemical equation associated with this reaction in your report
sheet. Reserve the colored solution for Part 5.
Part 5. Reaction of CuSO4 with Zn metal
Transfer operations back in the fume hood. Weight approximately 2.0 g of 20-mesh zinc metal from the balance
room. In the fume hood, add all at once the freshly weighed zinc metal into the colored mixture obtained from Part
4. What happens? What is the gas produced? What is the color and identity of the solid product generated? Observe
and record what you see during the process in your report sheet. Suggest and write a balanced chemical equation
associated with this reaction in your report sheet.
Once the evolution of gas becomes very slow, decant the supernatant liquid and pour it into the waste container
provided. See section on ‘Disposal and Cleanup below. To ensure complete removal of excess zinc solid from the
solid copper product add 10 mL of 6.0 M HCl to the solid, stir, and gently warm the mixture on a hot plate, but do
not allow it to boil. All traces of the silver/gray zinc should react away and gas evolution should cease when the
zinc is consumed, but do not spend more than five minutes on this step. Decant the supernatant liquid (dispose
properly) and transfer the solid product to a clean evaporating dish.
Wash (rinse) the product with 5 mL of deionized water and decant the wash water. Repeat the washing and
decantation at least two more times. If you are using flame as a heating source, make sure that you move away
from it or transfer to the fume hood for the next step. Wash the solid metal further with 5 mL of methanol.
Dispose methanol washings into the proper container. See ‘Disposal and Cleanup’ section below. Dry the metal by
placing the porcelain dish on top of hot plate set at 75°C. Do not crank the hot plate to higher temperatures to
prevent formation of CuO. Using a spatula, transfer the dried metal to a pre-weighed 100 mL beaker and weigh to
the nearest milligram. It is best if you use the same mass balance as in Part 1. Calculate the mass of metal you
recovered and record it in your report sheet.
Percent Recovery
Calculate the percent recovery of your metal using the equation below:
%𝑅𝑒𝑐𝑜𝑣𝑒𝑟𝑦= !
“##$%&’($)’&’*!’+”, 𝑥100%
If you are careful at every step, you will recover nearly 100% of the copper you started.
Disposal and Cleanup
Recovered copper metal must be collected in a container provided by the instructor. The supernatant solution
decanted in Step 5 contains zinc sulfate, ZnSO4 and zinc chloride, ZnCl2 which can be disposed of down the sink
drains. Methanol can be poured out directly into the sink with running water. Decantates with excess concentrated
acids or bases must be diluted in a beaker with overflowing running water while in the sink.
1. Initial mass of copper wire
2. Mass of watch glass + dry, recovered copper
3. Mass of watch glass (enter 0 if beaker mass is teared)
4. Mass of recovered dry copper product
Percent Recovery
Calculate the percent recovery of the metal at the end of the complete reaction cycle. Show your calculation
in the space below.
Observations and Chemical Equations
Record your observations in ink for each step in the cycle as indicated on the following pages. Observations
should be descriptive but concise, and should include descriptions of the starting materials, changes that
occur during the reaction (for example “…a red/brown gas was emitted and the temperature of the mixture
increased”), and of the resulting product. You will also be asked to write the balanced, molecular form of the
chemical equation that describes each reaction, and to classify each reaction using the four item scheme
described in the experiment instructions. The four choices are:
A. Precipitation
B. Acid-Base
C. Reduction-Oxidation
D. Decomposition
You may use pencil in writing balance chemical equations. You can work on your chemical equations,
classifications, and answer to other questions while you are waiting for the reactions to complete. However,
you should never leave a reaction unattended.
Part 1: Reaction of solid copper with concentrated nitric acid
Observations: (partial example given)
Reactants: The copper wire was bright/shiny after cleaning with steel wool. The nitric acid solution
was nearly clear/colorless although a very faint reddish brown discoloration was evident.
Reaction: _________________________________________________________________________________
Products: _________________________________________________________________________________
Balanced molecular chemical equation (this first reaction equation is given as an example)
What ions are in solution after the reaction is complete? ________________________________________
What is the oxidation state of the copper reactant?
What is the oxidation state of the copper product?
Reaction Classification: ____________________________________________
Part 2: Reaction of copper(II) nitrate with sodium hydroxide solution
Balanced molecular chemical equation for the copper(II) nitrate + sodium hydroxide reaction
What ions are in solution after the reaction is complete? ________________________________________
What is the oxidation state of the copper reactant?
What is the oxidation state of the copper product?
Primary Reaction Classification:
In addition to reacting with the copper nitrate, the sodium hydroxide will also react with excess nitric acid left over
from the Part 1 reaction. Write the balanced molecular chemical equation for this reaction and classify this
Balanced molecular chemical equation for the nitric acid + sodium hydroxide reaction
Secondary Reaction Classification: __________________________________ Part
3: Decomposition of copper(II) hydroxide with heat
Balanced molecular chemical equation for this reaction
What is the chemical name and formula of the precipitate? ___________________________________________
What ions are removed from the precipitate by washing? ___________________________________________
(Hint: considered the ions that are in excess or otherwise left over from the previous steps.)
What is the oxidation state of the copper reactant?
What is the oxidation state of the copper product?
Reaction Classification: ____________________________________________
Part 4: Reaction of copper(II) oxide with sulfuric acid
Balanced molecular chemical equation for the copper nitrate + sodium hydroxide reaction
What ions are in solution after the reaction is complete? ________________________________________
What is the oxidation state of the copper reactant?
What is the oxidation state of the copper product?
Reaction Classification: ____________________________________________
Part 5: Reaction of copper(II) sulfate with zinc solid
Balanced molecular chemical equation for the copper(II) sulfate + zinc reaction before HCl(aq) is added
Note: the gas produced during this reaction is due to a side reaction between zinc and the excess sulfuric acid
from the previous step – it is not a product of this reaction.
What is the oxidation state of the copper reactant?
What is the oxidation state of the copper product?
Primary Reaction Classification:
In addition to reacting with the copper(II) sulfate, the zinc will also react with excess sulfuric acid left over from the
Part 4 reaction as noted above. Write the balanced molecular chemical equation for this reaction.
Balanced molecular chemical equation for the zinc + sulfuric acid reaction
Give the name and formula of the gas evolved?
What is the oxidation state of the zinc reactant?
What is the oxidation state of the zinc product?
Secondary Reaction Classification:
Hydrochloric acid was added to remove excess zinc from the desired solid product. Write the balanced molecular
chemical equation for this reaction and classify this reaction.
Balanced molecular chemical equation for the zinc + hydrochloric acid reaction
Give the name and formula of the gas evolved?
What is the oxidation state of the zinc reactant?
What is the oxidation state of the zinc product?
Secondary Reaction Classification:
Write the net ionic form of the chemical equations for the two zinc + acid reactions. It may be helpful to write the
ionic form of these equations on a separate sheet to guide you when writing the net ionic form.
Zn(s) + H2SO4(aq) rxn:
Zn(s) + HCl(aq) rxn:
Notice any similarities? (This is a rhetorical question.)
Reflection Questions
1. The following scenarios represent mistakes a student could make during this experiment. Indicate how each
error would affect the final copper metal recovery – that is would it be lower than expected, would there be no
difference, or would it be higher than expected. Also give a brief explanation supporting your selection.
a. In Part 1 after the copper + nitric acid reaction is complete some of the solution is spilled.
The Cu recovery would be low because the copper is in the solution phase as the Cu2+ ion at the
end of the Part 1 reaction, and so less copper will be recovered at the end of the cycle.
b. In Part 2 an insufficient amount of sodium hydroxide solution is added.
c. In Part 3 the mixture “bumps” when heated and some liquid and solid is lost.
d. In Part 3, while decanting the clear liquid from the solid, some of the liquid is spilled.
e. In Part 4, a student inadvertently adds 20 mL of the 6 M H2SO4(aq) instead of the indicated 15 mL.
f. In Part 5 a student forgets to add HCl(aq) and some unreacted zinc remains in the final product.
g. In Part 5 the final product is not completely dried when it is weighed.
2. In Part 1 we used nitric acid to oxidize the solid copper and drive it into solution as the Cu2+ ion. Show that the
4.0 mL of nitric acid used was an excessive amount by calculating the volume (in mL) of 16 mol/L nitric acid
solution needed to completely react 0.525 g of solid copper. Start this problem by copying the balanced
chemical equation for this reaction to ensure you use the correct stoichiometry in this calculation.
Balanced molecular form of the chemical equation for copper solid reacting with nitric acid solution
3. In Part 2 sodium hydroxide solution is added to the copper(II) nitrate solution to precipitate copper(II)
hydroxide. Not only do we need to add enough sodium hydroxide to precipitate all of the copper ion, but, as
you should have shown in Problem 2, there is excess nitric acid present and we must add sufficient sodium
hydroxide to neutralize this as well. Show that NaOH(aq) was added in excess by completing the following
a. Calculate the volume (in mL) of 3.0 mol/L NaOH(aq) needed to completely precipitate the 0.525 g of Cu2+ in
b. Calculate the volume (in mL) of 3.0 mol/L NaOH(aq) needed to completely neutralize the excess nitric acid.
Assume that the concentration of nitric acid, after it has reacted with copper, is 8.1 mol/L and the volume of
the solution is 4.0 mL.

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