In the Laboratory
Formation and Dimerization of NO2 A General Chemistry Experiment April D. Hennis, C. Scott Highberger, and Serge Schreiner* Department of Chemistry, Randolph-Macon College, Ashland, VA 23005 Experiments demonstrating gas laws are found in most general chemistry laboratory manuals. While a number of experiments illustrate Boyle’s law (1), Charles’s law (2), and the ideal gas law (3), experiments demonstrating GayLussac’s law of combining volumes are virtually absent (4). This is in part because most simple gas-phase reactions mentioned in general chemistry textbooks in connection with this law (for example H2 /O2, N2 /H2, or H2/Cl2) are either incomplete, rather slow, or violent, or they require sophisticated gas-handling techniques. One notable exception is the NO/O2 reaction (eq 1), which is rapid at room temperature. 2NO + O2 → 2NO2
(1)
This particular system has the added advantage that it lends itself to the determination of partial pressures of the various gases as well as the evaluation of the equilibrium constant Kp for the dimerization of NO2 (eq 2). 2NO2
N2O4
(2)
In addition, it is an environmentally important system (5, 6) and it can be used to discuss solubilities of chemical compounds. We have incorporated all of these concepts into a simple experiment that can be performed in the introductory chemistry laboratory. Experimental Procedure CAUTION: NO, NO2 and N2 O4 are toxic gases. All procedures should be carried out in a well-ventilated hood. Care should be taken in the handling of hypodermic syringes. All materials were used as received from the manufacturer, with no further purification. Oxygen was supplied by Southern Oxygen & Supply and had a purity of 99.999%. Plastic syringes (5 and 20 mL) were obtained from Thomas Scientific. Luer needle lock ground glass syringes (30 mL) were purchased from PGC Scientific, rinsed with water and acetone, and air-dried before use. Rubber septa (12.5 mm o.d.) were obtained from Aldrich Chemical Co.
Preparation of Gases Prepare two identical gas-containing set ups (Fig. 1). The following steps can be carried out by the instructor well in advance of the laboratory session. Oxygen Setup In the first apparatus, flush the air out of the 250-mL Erlenmeyer flask with O2 gas by introducing a hypodermic needle connected to the oxygen supply line through the rubber septum. After 10 minutes, withdraw the hypodermic needle from the septum and close the oxygen tank. If no oxygen tank is available, O2 can be prepared by the catalytic decomposition of hydrogen peroxide. NO Setup With the second apparatus, add 10 g of NaNO2 and 20 g of FeSO4 to the 250-mL Erlenmeyer flask. Mix the two solids well. Add 10 mL of water to the flask and stopper immediately. The generation of NO starts at once. Since the flask contains air, the NO reacts with oxygen in the air to form NO2, as evidenced by the appearance of a brown gas. After several minutes, all of the oxygen is consumed and the contents of the Erlenmeyer flask slowly become colorless. Once the flask contains pure NO, students are ready to fill their gas reaction chamber.
Procedure Prepare the 30-mL glass syringe (the reaction chamber) by attaching a rubber septum to it and securing it with wire (Fig. 2). Make sure that the plunger is at zero volume before gases are introduced. Withdraw 20 mL of NO with a 20-mL plastic syringe. Transfer the NO to the 30-mL glass syringe. Record the initial NO volume, the atmospheric pressure, and the room temperature. Record the room temperature several times during the course of the experiment in order to obtain an average experimental temperature. With a 5-mL plastic syringe, withdraw 3-mL of oxygen and add it to the reaction chamber. Record the new volume of the glass syringe. Add oxygen in 3-mL increments to the reaction chamber until the total volume of the reaction chamber is just below 30 mL. Add 2 mL of distilled water and record the final volume of the reaction chamber. Results and Discussion Successive 3-mL additions of oxygen to a reaction chamber containing 20 mL of NO result in the formation of NO2 , followed by its dimerization to N2 O4 (eq 3).
Figure 1. Gas generation setup. a: Rubber stopper. b: Bent glass tubing. c: Side-arm flask. d: Rubber septum. e: Tygon tubing. f: Bubbler. *Corresponding author.
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Figure 2. Gas addition syringe and reaction chamber.
Journal of Chemical Education • Vol. 74 No. 11 November 1997
In the Laboratory x = 0.47, or 53% of the NO2 dimerizes
To solve for the product volumes: V(NO2): 2(3.0)(0.47) = 2.8 mL V(N2O4): [2(3.0)(1 – 0.47)/2] = 1.6 mL
To calculate the partial pressures of the product gases (Patm = 0.987 atm): 2NO + O2 → 2NO 2
N2O4
P (NO2): [2.8 mL /18.4 mL] × Patm = 0.15 atm P (N2O4): [1.6 mL /18.4 mL] × Patm = 0.086 atm
(3)
To calculate the equilibrium constant Kp:
While the reaction of NO with oxygen goes to completion, the dimerization of nitrogen dioxide to dinitrogen tetroxide does not. When the reaction is completed, the mixture contains both NO2 and N2O4 . Typical student data and calculations are shown in Table 1. We tell students that it is best to do the calculations using a spreadsheet. By setting up columns and formulas for each of the quantities to be calculated, one can enter two quantities (volume of oxygen added and volume of syringe in Table 1) for each experimental point and have the spreadsheet calculate all others. The experimental data used to do all of the calculations are the total volume of oxygen added and the observed total volume of gas in the reaction syringe after each addition (second and third columns in Table 1). In setting up this spreadsheet, students need to realize that between the third and fourth additions of oxygen, the NO is exhausted. Finally, most students have difficulty deciding how to allot the experimental volume that is not due to oxygen or nitric oxide to the product substances, nitrogen dioxide and dinitrogen tetroxide. To help them set up that part of the spreadsheet, we tell them that they have to determine the extent of dimerization of NO2 to N2 O4 . Once they have established that quantity, it is fairly easy to determine the experimental volumes of the two product gases. Sample calculations for the second data entry in Table 1 for the extent of dimerization, the volumes of the product gases, the partial pressures of product gases and the equilibrium constant are shown below (x = the fraction of NO2 that does not dimerize). 2NO + O2 → 2NO 2 Stoich. V (mL) 2(3.0) excess V (mL) 14.0
3.0 0.0
2(3.0)( x)
Kp = P (N2O4)/ P (NO2)2 = (0.086)/ (0.15)2 = 3.8 atm{1
In the student data shown in Table 1, the average dimerization was 61%, which yielded an average equilibrium constant of 2.7 ± 1.0 atm {1 at 295 K. The value for the equilibrium constant compares reasonably well with Kp as calculated from thermodynamic data (8.6 atm{1 at 295 K). Upon addition of 2.0 mL of water to the reaction chamber, the syringe volume decreases significantly. For the data shown in Table 1, the volume of the syringe after addition of water decreased to 16.2 mL. Since oxygen is “insoluble” in water and the nitrogen oxides are very soluble, the drop in volume can be attributed to the dissolution of NO 2 and N2O 4. Students can visualize and analyze their data and analysis by generating two graphs using their spreadsheet columns. Figure 3 depicts a plot of the calculated pressures of nitrogen dioxide and dinitrogen tetroxide versus the total volume of oxygen added. Figure 4 shows a plot of the equilibrium constant as determined after each addition of O2 versus total volume of oxygen added. Students should be able to deduce from Figure 3 that as long as NO is in excess, the partial pressures of the two product gases increase as a function of oxygen added; but once oxygen is in excess, the partial pressures of the two reactant gases decrease as a function of oxygen added. Figure 4 illustrates that the equilibrium constant is indeed a constant (at least within experimental error) regardless of the amount of oxygen added.
N2O4
Conclusion
[2(3)(1 – x )] /2
An experiment has been developed that allows students to determine partial pressures and equilibrium constants for the formation and dimerization of NO2 . The experiment described in this paper can be carried out in about 45 minutes with students working in groups of two. For departments with PCs equipped with spreadsheet software,
Therefore, the fraction of NO2 that dimerizes to N2O 4 is equal to (1–x). To solve for x: V (syringe) = V (NO) + V (O2) + V (NO 2) + V (N2O4) 18.4 = 14.0 + 0.0 + [2(3.0)(x)] + [2(3.0)(1–x )/2]
Table 1. Typical Data Table of Gas Volumes, Pressures, and Equilibrium Constants Volume (mL) Addition Oxygen Total Unreacted Unreacted NO2 No. Initial Added Syringe NO Oxygen
P (atm) NO2 Final
N2O4 Final
NO2
N2O4
Kp (atm{1)
0
0
20.0
20
0
0.0
0.0
0.0
0
0
—
1
3
18.4
14
0
6.0
2.8
1.6
0.15
0.086
3.8
2
6
16.6
8
0
12.0
5.2
3.4
0.31
0.20
2.1
3
9
14.3
2
0
18.0
6.5
5.8
0.45
0.40
2.0
4
12
15.0
0
2
20.0
6.0
7.0
0.40
0.46
2.9
5
15
18.1
0
5
20.0
6.2
6.9
0.34
0.38
3.3
6
18
21.0
0
8
20.0
6.0
7.0
0.28
0.33
4.2
7
21
25.6
0
11
20.0
9.2
5.4
0.35
0.21
1.7
8
24
29.0
0
14
20.0
10.0
5.0
0.34
0.17
1.5
Vol. 74 No. 11 November 1997 • Journal of Chemical Education
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Equilibrium constant (atm{1)
Pressure of NO2 or N2O4 (atm)
In the Laboratory
Volume of oxygen added (mL) Volume of oxygen added (mL)
Figure 3. Calculated pressures of product gases after each addition of oxygen vs. total volume of oxygen added. –d – NO2 ; –j – N2O 4.
Figure 4. Equilibrium constant after each addition of oxygen vs. volume of oxygen added.
this experiment readily provides data that students can enter and manipulate. All calculations and plotting of results can easily be accomplished with packages such as Lotus, QuattroPro, or Excel.
Acknowledgment This experiment was developed as a result of one of us (S.S.) attending an NSF-sponsored workshop on Computer Based Laboratories for Chemistry and Physics at Evergreen State College. We wish to thank the many introductory chemistry students who have performed this experiment over the last five years. Literature Cited 1. Hall, J. F. Experimental Chemistry; Heath: Lexington, MA, 1989; p 149. 2. Hall, J. F. op. cit.; p 153. 3. Milio, F. R.; Debye, N. W. G.; Metz, C. Experiments in General Chemistry; Saunders; Philadelphia, 1991; p 157. 4. Everett, K. G. J. Chem. Educ. 1982, 59, 802. 5. Shooter, D. J. Chem. Educ. 1993, 70, A133. 6. Hollenberg, J. L.; Brice, L. K. J. Chem. Educ. 1987, 64, 893.
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Journal of Chemical Education • Vol. 74 No. 11 November 1997
