In this experiment we will investigate two important concepts. First, we will see an example of an exchange reaction via the oxidation-reduction (redox) reaction (Reaction CL.2) and Reaction CL.3), and secondly, we will calculate a theoretical yield and percent yield to determine the effectiveness of our experimentation.
A redox reaction involves an exchange of electrons. One species must be oxidized, which means it loses electrons, so its oxidation state increases (\( \ce {Al^0 -> Al^3+ + 3 e^-} \)) while the other species must be reduced, which means it gains electrons, and its oxidation state decreases (\( \ce {Cu^2+ + 2e- -> Cu^0} \)).
The molecular equation (Reaction CL.2 for the chemical reaction being carried out in this experiment is: \begin {equation} \ce{3 CuCl2(aq) + 2Al(s) -> 3Cu(s) + 2AlCl3(aq)} \end {equation}
The chloride ions are spectator ions, so the net ionic equation for the reaction is written as shown in Reaction CL.3: \begin {equation} \ce{ \label {Cu2+_and_Al_net} 3Cu^2+(aq) + 2Al^0(s) -> 2Al^3+(aq) + 3Cu^0(s)} \end {equation} where the superscripts indicate the oxidation states of the species, positive for the ions and zero for the elements. The changes in oxidation states of the species are more obvious in the net ionic equation although the student should learn to recognize that the form of Reaction CL.2 represents a redox reaction.
The law of conservation of mass states that matter can neither be created nor destroyed. Thus, if you start your experiment with a certain amount of substance and convert this to another substance, then you should end up with the same total amount of substance as you started with. In this experiment you will determine the mass of copper from the \( \ce {CuCl2} \) that you started with and measure the mass that was produced.
You will obtain a sample of copper II chloride (\( \ce {CuCl2} \)) and record an exact mass (i.e, recording all the digits displayed by the balance before and after the decimal place). You will then dissolve this compound into DI water and convert the copper II ions (\( \ce {Cu^2+} \)) into metallic copper (\( \ce {Cu^0} \)) in a reaction with aluminum metal. In theory, you can calculate the number of moles of copper ions with which you started and you can assume that the same number of moles of metallic copper will be obtained. Therefore, you will be able to calculate the number of grams of metallic copper that you should recover.
To carry out this chemical reaction you will be using aluminum wire. The aluminum atoms in the wire lose electrons and are considered to be oxidized (3 ). The copper (II) ions gain electrons and hence are reduced (4). At the same time that you are producing metallic copper, the metallic aluminum will be changing from the metallic state Al0 (aluminum with the oxidation state of zero) to dissolved aluminum (III) ion (\( \ce {Al^3+} \)). You will assume that you have an excess of aluminum so you need only be concerned with the copper in your calculations for theoretical yield. \begin {align} \ce{ Al^0 (s) -> Al^3+ (aq) + 3e- \label {oxid_Al}}\\\ce{ Cu^2+ (aq) + 2e- -> Cu^0 (s) \label {red_Cu2+}} \end {align}
The mathematical conversion of the mass of \( \ce {CuCl2} \) to moles of \( \ce {CuCl2} \) is performed using the molar mass (MM) of one mole of copper (II) chloride as seen in Equation CL.1. \begin {equation} \label {mass-mol_CuCl2} \frac {\text {mass}\ \ce {CuCl2}}{1}\times \frac {1\ \text {mole\ of}\ \ce {CuCl2}}{\text {MM\ of}\ \ce {CuCl2}}= \text {?? \ moles\ of}\ \ce {CuCl2} \end {equation}
You can obtain the molar mass (MM) of copper (II) chloride (\( \ce {CuCl2} \)) by using the atomic masses given in the periodic table.
For every mole of copper (II) chloride that we start with, we should get the same number of moles of the copper cation, as shown in Reaction CL.6. \begin {equation} \ce{ CuCl2(aq) -> Cu^2+(aq) + 2Cl-(aq)} \end {equation}
Note that for every mole of \( \ce {CuCl2} \) there is one mole of \( \ce {Cu^2+} \). Using 4, which shows the addition of electrons to the copper ion to form the metal, it is seen that the conversion of copper ion to metallic copper is a one to one ratio as expected from the law of conservation of mass (in this reaction, the copper can come from no other source).
You will use the stoichiometry of these equations to help you calculate the theoretical yield of copper metal, which is the mass of copper you expect to produce in the reaction based on the amount of \( \ce {CuCl2} \) that you started with. The calculation of theoretical yield assumes that the reaction goes perfectly to completion with no losses.
In real life chemistry, everything does not always go according to plan. There is no flaw with the chemistry, but perhaps there are various experimental techniques that will prevent you from obtaining the total amount of copper that you predict (the theoretical yield). Most often this is due to an incomplete reaction, which can easily be seen in this experiment if there is a greenish color still persisting in solution. Other errors could be due to mistakes when measuring the mass of the final product. If excess moisture is present, this will cause the mass you obtain to be higher than the amount that you calculated you should produce. One way to evaluate how well your experiment was performed is to look at the ratio of what you recovered compared to what you predicted. This comparison is called percent yield and can be calculated according to Equation CL.2. \begin {equation} \label {perc_yield} \text {percent\ yield}=\frac {\text {actual\ yield}}{\text {theoretical\ yield}}\times 100\% \end {equation}
1.
Obtain a known amount of copper (II) chloride (slightly < 3 g) and record its
mass to the precision of the instrument (i.e., record all the digits displayed by
the balance before and after the decimal place.) Use a weigh boat so that
you will be able to quantitatively transfer the sample to your reaction
beaker.
2.
Place the sample in a 250-mL beaker.
3.
Rinse any residual copper (II) chloride from the weigh boat into the beaker with
deionized water, adding approximately 60 mL of water to the beaker. Stir until
the copper (II) chloride has completely dissolved. Note the color change
in your lab notebook. The resulting solution should be blue or blue
green and this color is characteristic of \( \ce {Cu^2+} \) ions in
water.
4.
Obtain a piece of aluminum wire with a mass of about 1.5 grams. These are
precut so you may assume that the portion is large enough.
5.
Twist one end of the aluminum wire into a coil that should fit around inside of
the beaker (the more aluminum wire surface area available the better) and
immerse the coiled end in the copper solution.
6.
Observe the reaction that occurs. Note that with time, the blue color of the
solution will fade as the aluminum reacts, and copper metal will accumulate on
the surface of the aluminum. Shake the wire continuously, knocking it against
the sides of the beaker frequently to loosen the copper metal that accumulates
on the surface of the wire. The reaction has reached completion when the blue
color in the solution has disappeared and the aluminum wire has no
more copper forming on it. This process generally will take 20 to 30
minutes.
7.
When the reaction is complete, remove the aluminum wire and shake or tap the
aluminum so that all of the copper falls back into the beaker. Note that if
aluminum accidentally falls or the wire breaks into the beaker, you will be
isolating impure copper (copper mixed with aluminum), so be sure that only the
copper remains in the beaker. Use a wash bottle of deionized water to rinse the
aluminum wire over the beaker to remove any residual copper. You might find it
helpful to scrape the copper off the aluminum with the rubber tip on a glass
stirring rod.
8.
Discard the aluminum wire in the appropriate waste container.
9.
Obtain a piece of filter paper and using a pencil, print your drawer number on it.
Record the mass of this filter paper along with a watch glass to the precision of
the instrument (i.e., electronic balance). The paper must fit into the Büchner
funnel used in the filtration step.
10.
Assemble a vacuum filtration apparatus as described by your instructor
(Figure CL.1). Ask the instructor if there is any doubt about the assembly
before starting this process.
11.
Place the labeled filter paper label side down in the Büchner funnel and isolate
the copper by vacuum filtration. Carefully rinse all of the copper in
the beaker into the Büchner funnel using a wash bottle of deionized
water.
12.
When the copper has been transferred successfully, wash the copper with several
portions of DI water - about 20 mL each time to ensure any non-visible salts are
washed away.
13.
Wash the same filtrate with about 20 mL of acetone. The acetone will wash the
water away and evaporate more quickly than water, thus hastening the drying of
the copper. Allow the vacuum to dry the copper in the funnel for two
minutes.
14.
Carefully transfer the filter paper with copper to the pre-weighed watchglass by
placing the watchglass concave side on top of the funnel and turning the funnel
upside down. The filter paper will be on top showing your drawer number. Place
the watchglass and its contents into the oven to dry for 10 minutes. Allow it to
cool to room temperature for a minimum of 10 minutes. Record its mass to the
precision of the instrument.
15.
Calculate the % yield of copper obtained.
Name: ___________
Section: ________________ Date: ______________________
Mass of starting copper(II) chloride (\( \ce {CuCl2} \))
Show calculations in the boxes below. Include correct significant figures and
units.
Calculation: What is the molar mass of \( \ce {CuCl2} \)?
Calculation: How many moles of \( \ce {CuCl2} \) were dissolved in
solution?
Calculation: Given the number of moles of \(\ce{CuCl2}\) dissolved in solution from
the box above, how many moles of copper ions (\( \ce {Cu^2+} \)) were
initially in solution?
\begin {equation} \ce{ CuCl2(aq) -> Cu^2+(aq) + 2Cl-(aq)} \end {equation}\begin {equation} ? \times \frac {\ce {1 mole Cu^2+}}{\ce {1 mole CuCl2}}= ? \end {equation}
Calculation: How many moles of copper metal (\( \ce {Cu^0} \)) do you
expect to be formed?
\begin {equation} \ce{ 3Cu^2+(aq) + 2Al^0(s) -> 2Al^3+(aq) + 3Cu^0(s)} \end {equation}\begin {equation} ? \times \frac {\ce {3 mole Cu^0}}{\ce {3 mole Cu^2+}}= ? \end {equation}
Calculation: What is the molar mass of \( \ce {Cu^0} \)?
Calculation: What is the theoretical yield of copper metal in grams?
Calculation: Show the calculation for the percent yield of copper metal.
Percent yield of copper metal
Did you collect less copper than expected or apparently more? Explain
why.
What experimental errors occurred during your experiment that may have caused something other than 100% recovery? Use complete sentences and cite more than one example. Explain.
