In the Laboratory
A Copper-Sulfate-Based Inorganic Chemistry Laboratory for First-Year University Students That Teaches Basic Operations and Concepts Emilio Rodríguez and Miguel Angel Vicente* Departamento de Química Inorgánica, Universidad de Salamanca, Plaza de la Merced, S/N, 37008 Salamanca, Spain; *
[email protected] There is a tendency toward increasing specialization in the laboratory work of university chemistry courses; in some cases, students use sophisticated techniques from the very first years. Unfortunately, this means that fundamental laboratory work is sometimes overlooked. In this paper we present a 10hour integrated experiment, based on copper sulfate, which covers all the inorganic chemistry topics in first-year courses in both Chemistry and Chemical Engineering. It introduces the most common laboratory operations and enhances students’ skill in experimental work. The experiment has three steps (Scheme 1): (i) purification of an ore containing copper sulfate and insoluble basic copper sulfates, (ii) determination of the number of water molecules in hydrated copper sulfate, and (iii) recovery of metallic copper from copper sulfate. The main basic operations carried out are weighing, heating, filtration (simple and vacuum-assisted), purification, crystallization, and redox reaction. The main concepts are pure compounds and mixtures; hydrated and anhydrous salts; solubility; unsaturated (dilute and concentrated), saturated, and supersaturated solutions; adsorbed and crystallization water; reversible dehydration; redox reaction, reduction potential; free energy; spontaneity; and catalysis. All the work is carried out on copper sulfate, which can be handled by the students without special precautions. No special equipment is required, and all products are easily recycled.
Hazards Copper sulfate (both hydrated and anhydrous) is harmful if swallowed. A mask may be used when handling ground solids. Sulfuric acid is corrosive and causes severe burns. In case of contact with eyes, rinse immediately with plenty of water and seek medical advice. Never add water to this product. Experimental Procedure and Results All the products used and generated in this experiment may be recycled. Insoluble basic copper sulfates are mixed with the remaining copper sulfate, adding the anhydrous copper sulfate generated in step 2 and an amount of pure copper sulfate equivalent to that used up in step 3. After the mixture is heated in an oven at 80 °C for 5 hours, it may be used again. Metallic copper and FeSO4⭈7H2O can be used in subsequent experiments.
Step 1. Purification of an Impure Copper Sulfate Ore The natural ore used is composed of copper sulfate (CuSO4⭈5H2O) and basic copper sulfates, mainly brochantite [CuSO4⭈3Cu(OH)2] and antlerite [Cu3(SO4)(OH)4]. It has an approximate composition of 70% copper sulfate and 30% copper basic sulfates, in spite of which the solid is pale green in color. The ore we use has a natural origin, but it is easy to prepare by mixing these solids in these approximate ratios.
insoluble copper basic sulfates Step 1
impure copper sulfate ore
Step 2
crystalline CuSO4·5H2O
Step 3
crystalline CuSO4·5H2O
dissolving
copper sulfate solution
heating ≈ 300 °C
anhydrous CuSO4
concentration crystallization
heating ≈ 900 °C
copper sulfate mother liquor crystalline CuSO4·5H2O
CuO
Scheme 1. Step 1: purification of an ore containing copper sulfate and insoluble basic copper sulfates; Step 2: determination of the number of water molecules in hydrated copper sulfate; Step 3: recovery of metallic copper from copper sulfate.
metallic copper
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a) H+, heating b) iron nails
iron sulfate mother liquor
concentration crystallization
crystalline FeSO4·7H2O
Journal of Chemical Education • Vol. 79 No. 4 April 2002 • JChemEd.chem.wisc.edu
In the Laboratory
About 20 g of the raw solid is added to 200 mL of water. The solid is mainly immiscible with water and floats on the surface. The mixture is heated to about 70 °C, with shaking, and the solution takes on a blue color that becomes more intense over time. At this point, the solubility curve of copper sulfate is explained, emphasizing that in this case solubility increases with temperature, and the concepts of unsaturated, saturated, and supersaturated solutions are discussed. The two curves referring to mass of solution and volume of solution are used (Table 1 and ref 1), making it possible to talk about the density of the solutions. When no more solid is being dissolved (after a few minutes), heating is stopped and the solution is filtered by means of a simple process. The insoluble copper basic sulfates are separated in the filter, and copper sulfate is contained in a dilute solution (ca. 200 mL). The solid is dried, weighed for a final mass balance, and sent to recycling. In the next step the solution is concentrated by heating it to the boiling point for about one hour, adding a few pumice stones to favor boiling. Then the solution (ca. 40 mL) is placed in a porcelain vessel at room temperature (it is filtered if turbidity is observed), and left until the next day. The concentration of copper sulfate in the solution increases from dilute to concentrated at high temperature; the solution becomes supersaturated when cooled to room temperature; and large, well-formed crystals are formed. The crystals are filtered using a filter flask and a Büchner funnel, which allows an explanation of vacuum-assisted filtration. The crystals are dried in an oven at 45–50 °C; this eliminates the adsorbed water but does not affect the crystallization water—concepts that are emphasized in the second part of the work. The dried crystals are weighed in order to calculate the yield of the crystallization process. The mother liquor solution is weighed and its volume is determined, and its content of copper sulfate is calculated from the solubility curves. Again, it is pointed out that the solution is saturated, the solubility of copper sulfate being read from the corresponding curves at the laboratory temperature. The content of copper sulfate in the raw solid is calculated and a mass balance is carried out, which the students may use when looking for possible errors in the process. This step requires about three hours, although the solution must be left at room temperature overnight for complete crystallization. Typical results are as given in Table 2.
Step 2. Determination of Water Content of Copper Sulfate About 2 g of the pure hydrated copper sulfate obtained in step 1, ground and weighed with precision, is heated in a sand bath at about 300 °C. The blue color progressively disappears, until the solid is completely white. This illustrates perfectly the change of properties from the hydrated to the anhydrous salt. The white solid is left to cool for a few minutes, and weighed. It can be demonstrated that this dehydration is reversible: when a few drops of water are added to the anhydrous solid the blue color returns. The students are asked to calculate the number of water molecules in the hydrated salt, on the basis of the mass of the anhydrous compound. The thermal decomposition of the anhydrous copper sulfate can be studied by heating a sample to 900 °C in a
Table 1. Solubility of Copper Sulfate in Water (1) Solubility/g
T/°C
per 100 g of solution
per 100 mL of solution
᎑1.4
19.85
22.75
0.5
20.15
23.15
10
23.40
27.52
20
26.60
31.87
30
30.55
37.64
40
34.78
43.89
50
39.38
51.16
60
44.36
59.58
70
50.13
70.08
80
54.50
78.10
93
61.30
92.20
96
63.25
96.20
100
65.92
102.18
Table 2. Typical Experimental Results Measurement
Result
Step 1. Purification of the Copper Sulfate Weight of solid
20.00 g
Weight of copper basic sulfates
5.60 g
Weight of copper sulfate crystals
8.00 g
Volume of the solution
20.0 mL a
Weight of the mother liquor
24.00 g
Weight of copper sulfate in the mother liquor
6.30 g b
Yield of the crystallization process
40.0%
Content of copper sulfate in the raw solid
71.0%
Step 2. Determination of Water Content of Copper Sulfate Weight of hydrated copper sulfate
2.00 g
Weight of anhydrous copper sulfate
1.30 g
Water loss
0.70 g
Number of moles of anhydrous copper sulfate
8.10 mmol
Number of moles of water
38.90 mmol
Moles of water per mole of copper sulfate
4.8 (≈ 5)
Weight of copper oxide
0.65 g
Weight loss in the second heating
0.65 g c
Step 3. Recover y of Metallic Copper from Copper Sulfate Weight of hydrated copper sulfate
5.00 g
Weight of copper
1.25 g
Weight of iron sulfate heptahydrate
3.50 g
aDuring
crystallization the solution volume decreases by about 20 mL. bAverage value between solubility curves referring to volume of the solution and to weight of the solution. cLoss of a molecule of SO . 3
furnace. A black residue of CuO is thus obtained. In our case, this is not done by the individual students, but by the teacher. The students are informed of the nature and weight of the residue and asked to propose a reaction for the decomposition step that is compatible with the observed weight loss. This step requires about three hours. Typical results are shown in Table 2.
JChemEd.chem.wisc.edu • Vol. 79 No. 4 April 2002 • Journal of Chemical Education
487
In the Laboratory
Step 3. Recovery of Metallic Copper from Copper Sulfate A 5.00-g sample of the copper sulfate obtained in step 1 is dissolved in 50 mL of water acidified with 2 mL of sulfuric acid. The solution is heated, and when the temperature is 60–70 °C, 5 or 6 iron nails are added. Immediately, metallic copper is formed on the surface of the nails. The reaction is Cu2+ + Fe → Cu + Fe2+
(1)
This reaction is spontaneous: ∆G ° = ᎑1.54F J.1 It is a slow reaction and, besides heat, the sulfuric acid catalyst is needed to obtain a good reaction rate. This provides a good opportunity to introduce the concept of catalysis. During the reaction, it is necessary to separate the copper from the nails with a glass rod, to avoid the deposition of a copper sheet that prevents the access of Cu2+ cations to the surface of the nails. The end of the reaction is easily observed by the change in color from blue (Cu2+) to pale green (Fe2+). After this, the formation of hydrogen bubbles is observed. The hydrogen arises by the secondary reaction 2H+ + Fe → H2 + Fe2+
(2)
which is also spontaneous (∆G ° = ᎑ 0.88F J). The concepts of free energy change and speed of reaction are discussed here. In the present case, reaction 1 is at the same time the more exergonic and the faster reaction, but the fact that the free energy change and the speed of a reaction are independent is emphasized. The comparison between eqs 1 and 2 illustrates that the protons, which act as catalyst for reaction 1, become reagents for reaction 2. Once the blue color changes to pale green, the solution is cooled and filtered by vacuum-assisted filtration. The solid, which contains copper and the rest of the iron nails, is washed first with water and then with ethanol and dried in an oven at 60–70 °C. The iron nails are easily separated by using a magnetic stirrer.
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The reaction has a yield close to 100%, and its historical importance (recovery of metallic copper from copper sulfate solutions by sacrifice of scrap iron) is emphasized. Crystals of FeSO4⭈7H2O may be obtained by concentrating the mother liquor solution. An interesting complementary problem for the students is why Fe is oxidized to Fe2+ and not to Fe3+, and why Cu2+ is reduced to Cu and not to Cu+. Solving these problems provides experience with redox calculations. A final experiment is proposed to show that metallic Cu may act as a reducing agent against strong oxidants. A copper wire is introduced into a test tube containing an acidified solution of AgNO3. Metallic silver is formed on the surface of the wire, the reaction being 2Ag+ + Cu → 2Ag + Cu2+
(3)
The solution turns a very pale blue, owing to the Cu2+ formed. This may be verified by taking a small volume of the solution and adding to it a few drops of ammonia; the intensely blue complex [Cu(NH3)4]2+ is formed. This step requires about three hours. Typical results are shown in Table 2. Acknowledgments We thank our colleagues who have participated in the teaching of this subject in different years. Their contributions have considerably improved this experiment. Note 1. Standard electrode potentials: Cu2+/Cu = 0.33 V; Fe2+/Fe = 1 ᎑ 0.44V; H+/ Ú2H2 = 0.00 V; Ag+/Ag = 0.80 V; Fe3+/Fe = ᎑ 0.02 V; Cu+/Cu = 0.52 V.
Literature Cited 1. Seidell, A. Solubilities of Inorganic and Metal Organic Compounds, 3rd ed.; Van Nostrand: New York, 1953.
Journal of Chemical Education • Vol. 79 No. 4 April 2002 • JChemEd.chem.wisc.edu
