Chemistry Lab: Stoichiometric Analysis of an Iron-Copper Single Replacement Reaction

Chemistry Laboratory 6A – Stoichiometric Analysis of an Iron-Copper Single Replacement Reaction
Martin Sun
Purpose
The purposes of this experiment were to: determine the number of moles of iron reacted; determine the number of moles of copper produced; and calculate the ratio of moles of copper to moles of iron.
Materials and Methods Materials and methods for this laboratory followed those laid out in Experiment 6A on pages 56-59 of Essential Experiments for Chemistry by Morrison and Scodellaro with the following exceptions: one iron nail was used instead of two;
100 mL of water was used instead of 50 mL; a digital scale was used in place of a centigram balance; a 600mL beaker was used in place of a 250mL beaker; steel wool was used in place of sandpaper to remove the zinc coating on the nails; two tablespoons of copper(II) chloride crystals were used in place of one tablespoon; and the iron nail was not rinsed.
Observations
Prior to the experiment, a 600 mL capacity glass beaker was obtained from the front of the laboratory. The beaker was weighed on a digital scale and its mass was determined to be 213.7 grams. A galvanized iron nail and a lump of steel wool was then collected from the front of the laboratory and the zinc coating on the iron nail was rubbed off using the steel wool. The iron nail was then weighed on a digital scale and its mass was determined to be of 6.5 grams. Two tablespoons of copper(II) chloride (CuCl2) crystals were then added to the beaker. The copper(II) chloride crystals were blue in colour, shown in Figure 1 below.
Figure 1. Photograph of two tablespoons of copper(II) chloride crystals in a beaker.

One hundred millilitres of water was then added into the beaker and mixed using a glass stirring rod to create a copper(II) chloride solution. The solution was observed to be light blue in colour, similar to that of the copper(II) chloride crystals. The iron nail was then placed into beaker, fully submerged in the copper(II) chloride solution to begin the reaction. During the reaction, the glass stirring rod was used to scrape off precipitates that adhered to the nail. A few minutes after the reaction commenced, the solution was observed to have become green and significantly darker in colour. Figure 2 compares the two solutions.
Figure 2. Photographs of solution before (left photo) and after (right photo) the start of the reaction. The iron nail is not submerged in the solution on the left.

During the reaction, very obvious changes in the nail could be seen. The surface of the iron nail was coated with a red substance, similar in colour to that of iron(III) oxide. Although it cannot be observed from the photographs in Figure 3, the length of the iron nail was slightly reduced. Figure 3 below compares the iron nail before and after the reaction.
Figure 3. Photographs of iron nail before (top photo) and after the reaction (bottom).

During the reaction, a red coloured precipitate could be seen collecting on the bottom of the beaker, most of it scraped off from the iron nail. Twenty minutes after the experiment commenced, the solution was decanted and the iron nail was taken out using tongs placed on a paper towel. Once dried, the nail was weighed and the mass was determined to be 4.1 grams. The beaker with the remaining contents was placed on a hot plate for twenty minutes to evaporate the remaining liquid. The red precipitate could be clearly observed, collected at the bottom of the beaker. Figure 4 below shows the contents of the beaker after drying.
Figure 4. Photograph of contents of beaker after drying with the hot plate.

After the contents of the beaker was dried and cooled, the beaker with the contents was weighed and the mass was determined to be 217.2 grams. Table 1 below summarizes all the data collected for mass in laboratory 6A.
Table 1. Summary of all data collected for mass in the laboratory.

Before Reaction
After Reaction
Object
Iron Nail
Dry, Empty Beaker
Iron Nail
Beaker + Copper Precipitate
Mass (g)
6.5
213.7
4.1
217.2

Results In the laboratory, a chemical reaction between copper(II) chloride and iron occurred, resulting in a copper precipitate and an iron(II) chloride solution.
The chemical reaction in this experiment was:

Firstly, seeing the equation, we can deduce that the aqueous, dark solution in Figure 2 is iron(II) chloride. This chemical reaction was a single replacement reaction, in which iron replaced copper in copper(II) chloride, as iron is more reactive, creating iron(II) chloride and a copper precipitate, the red substance which can be observed in Figure 4. Therefore, the moles of iron that reacted should have been equal to the moles of copper produced or a molar ratio of 1:1, under ideal conditions. If not, there was an experimental error. To calculate the amount of each substance in moles, the mass changes need to be first determined, using the data in Table 1. Table 2 below summarizes the calculations.
Table 2. Summary of calculations for the mass changes for iron and copper.

Mass Equation
Calculation
Result (g)
Iron
Mass of Nail: Before – After Reaction
6.5 – 4.1
2.4
Copper
Mass of Beaker: With Copper Precipitate – Empty
217.2 – 213.7
3.5

To calculate the amount in moles of iron reacted and copper produced, the mass needs to be divided by the atomic mass of the corresponding element. Table 3 below summarizes the calculations.
Table 3. Summary of calculations for deriving the moles of iron reacted and copper produced.

Atomic Mass (g/mol)
Mole Equation
Calculation
Result (mol)
Iron
55.845

0.043
Copper
63.546

0.055

The molar ratio can now be calculated with the data for moles in Table 3. Table 4 below summarizes the calculations.
Table 4. Summary of calculations for deriving the molar ratio between iron and copper.

Moles
Molar Ratio
Iron
0.043 Molar Ratio = 0.78:1
(iron : copper)
Copper
0.055

As calculated in Table 4, the mole ratio quotient is equal to 0.78, therefore the molar ratio is equal to 0.78:1, which is not close or equal to the expected molar ratio of 1:1, under ideal conditions. Therefore, an experimental error was obviously present. An error in the experiment itself is likely, referencing the class data set in Table 5 below.
Table 5. Summary of class data set. Group 3 is the group of the current labwriter. Group 7’s data was omitted from the average as a major, obvious experimental error was present.
Groups
Fe Reacted (g)
Fe Reacted (mol)
Cu Produced (g)
Cu Produced (mol)
1
1.8
0.03223
2.3
0.03619
2
2.3
0.04118
2.4
0.03777
3
2.4
0.04297
3.5
0.05508
4
1.3
0.02328
2.6
0.04092
5
2.4
0.04297
4.5
0.07081
6
2.5
0.04476
3.3
0.05193
7
3.0
0.05372
19.5
0.30686
8
2.5
0.04476
5.6
0.08813
9
2.6
0.04655
4.3
0.06767
With the data in Table 5 above, the average molar ratio can be calculated, summarized in Table 6 below.
Table 6. Summary of average molar ratio for the class data set.

Averages

Moles Reacted/Produced
Molar Ratio
Iron
0.040 Molar Ratio = 0.71:1
(iron : copper)
Copper
0.056

The class average molar ratio has been calculated to be 0.71:1. Comparing this ratio with the individual molar ratio of 0.78:1, the results are similar, both having a molar quotient with a difference of at least 0.2 from the expected quotient of 1.0 (± ~0.1). Therefore, an error within the experiment itself was present. In all the group data sets, the amount of copper was far too high. This could be caused by insufficient drying of the beaker. In a different perspective, the amount of iron could be too low, which may be caused by the zinc coating on the nail, reducing the effects of copper(II) chloride on it, and reducing the amount of iron reacted.
Conclusion
In this laboratory, a reaction between copper(II) chloride and iron occurred. This was done to determine the amount, in moles, of reacted iron and produced copper. Due to this reaction being a single replacement reaction, the amount of both elements should be the same under ideal conditions, yielding a mole ratio of 1:1. However, the molar ratio in the experiment was calculated to be 0.78:1. It is likely than a systematic error in the experiment itself was present, as the class average molar ratio was 0.71:1. The error could be from the remaining zinc on the iron nail after using the steel wool to scrape off most of the zinc, reducing the potency of the copper(II) chloride. In addition, there could still be liquid in the beaker after heating using the hot plate, adding to the weight.

Questions
1. If the tongs used in this experiment were made of pure iron, what might happen if the tongs were allowed to remain in contact with the CuCl2 solution?
Suppose that the copper precipitate is somehow shaken back into the beaker from the tongs. If the iron tongs were placed in the CuCl2 solution, it would essentially act the same as the nail in the experiment. The reaction is represented by this equation:

If iron tongs were placed, this would mean more copper produced without taking into account of the additional iron. The equation would then become:

This equation then becomes unbalanced, and we must double CuCl2 on the left side, and then double the FeCl2 on the right side, then double the Fe on the left side. Since at the end of the experiment, only the nail is weighed, and not the tongs, this means that the more iron that reacted is not being taken account of and the resulting molar ratio would be 1: >1 (iron: copper), not the actual molar ratio which should be 1:1.
2. A student carelessly allows some aluminum tongs to sit in a beaker containing CuCl2 solution for a period of time and a reaction occurs.
Original mass of tongs: 85.1 g
Final mass of tongs: 73.2
a. Write the balanced chemical equation for the reaction.

However, the AlCl2 needs to be changed to AuCl3 as 3 chlorine atoms are needed to balance aluminum’s +3 valence charge, and chlorine has a valence charge of -1.

Seeing this, the equation is now unbalanced as there are three chlorine atoms on the right as opposed to two chlorine atoms on the left.

Now, the aluminum and copper atoms are unbalanced. The balanced chemical equation is:

b. Calculate the moles of atoms of aluminum that reacted.
Firstly, the mass of aluminum must be calculated by subtracting the final mass by the initial mass
85.1 g – 73.2 g = 11.9 grams.
With this, the number of moles can be calculated by dividing this number by the atomic mass of aluminum, 26.982 g/mol.

c. Calculate the number of atoms of aluminum that reacted.
Since 0.441 moles of aluminum reacted, the number of atoms would be 0.441 multiplied by Avogadro’s number, 6.02 × 1023.

d. Calculate the moles of atoms of copper that are produced.
Since this is a single replacement reaction, the moles should be equal to the moles of aluminum that reacted, 0.441 moles.
e. Calculate the mass of copper that is produced.
To calculate the mass, the moles of copper must be multiplied by the atomic mass of copper, 63.546 g/mol.