In the Laboratory
Gas-Phase Kinetics of O3 + C2H2 and O3 + H2C=CCl2
W
Joel D. Burley* and Alisa M. Roberts Department of Chemistry, Saint Mary’s College of California, Moraga, CA 94575-4527; *
[email protected] Rationale Most courses in physical chemistry present detailed treatments of chemical reaction rates and reaction mechanisms, and gas-phase reactions are frequently used as examples to illustrate these fundamental concepts. Given the important role of gas-phase kinetics within the physical chemistry curriculum, the relative lack of gas-phase experiments for the physical chemistry laboratory is rather surprising. This report describes an experiment that measures reaction rate constants for the ozone–acetylene and ozone–vinylidene chloride reactions and introduces students to some of the techniques and calculations used in gas-phase kinetics. It can be compressed into a single laboratory period or expanded into a multiweek project, depending on the amount of time that is available. One possible explanation for the lack of gas-phase experiments in the physical chemistry curriculum is that most gas-phase reactions require equipment or techniques that are not readily available in undergraduate laboratories. In this regard ozone-plus-organic reactions offer a number of advantages. Ozone is easily produced by the photolysis of O2, and O3 concentrations can be monitored via simple spectrophotometric techniques. In addition, many ozone-plus-organic reactions are important to the chemistry of troposphere (1). Although the two reactions considered here are too slow to be tropospherically important, they are good choices for the experimental approach outlined below.
Figure 1. A schematic diagram of the reaction cell. The cell is constructed of Pyrex; fused silica windows are attached with epoxy to each end of the main body. The main body has an o.d. of 2.54 cm and an optical path length of 10.0 cm; the sidearm has an o.d. of 2.54 cm and an approximate length of 4.5 cm. Both valves are Teflon with Viton O-rings.
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Experimental Procedure Equipment and Chemicals Needed The equipment required for this experiment consists of a UV–vis spectrophotometer, a reaction cell, a vacuum manifold equipped with a 0–500-torr or 0–1000-torr pressure gauge, a low-pressure mercury lamp, a small vial to hold the sample of vinylidene chloride, and a variety of hardware and fittings needed to connect the cell and the chemical samples to the vacuum manifold. The chemicals that are needed are oxygen, acetylene, and vinylidene chloride.
Procedure The rate of reaction between ozone and the organic reactant is measured spectrophotometrically by monitoring the decay of the very strong O3 absorbance at 254 nm. A simple reaction cell, shown schematically in Figure 1, makes it possible to perform the necessary measurements with a standard UV/vis spectrophotometer. (For details regarding the design and construction of the cell, please see the Instructor’s Notes in the Supplemental Materials.W) To conduct a measurement, the entire cell is first evacuated and the sidearm is filled with approximately 300 to 600 torr of the organic reactant. Valve B is closed, and the main body of the cell is re-evacuated and filled with approximately 200 torr of O2 gas. The cell is then detached from the vacuum manifold and placed in the spectrophotometer, which is zeroed at λ = 254 nm. Approximately 0.1 to 0.5 torr of O3 reactant is generated in situ by removing the cell from the spectrophotometer and exposing the O2 in the main body to the output of a low-pressure mercury lamp that is placed directly in front of a cell window for 8–12 minutes. (The sidearm should be wrapped in aluminum foil to shield the organic reactant during the photolytic preparation of the ozone.) After the O3 has been generated, the cell is returned to the spectrophotometer and an initial absorbance value is recorded. Valve B is then cracked open for approximately one second and immediately closed. This results in an injection of organic reactant from the sidearm into the main body of the cell at time t = 0 with no significant back-flow of ozone or reaction products into the sidearm. Absorbance values are recorded at 15- or 30-second intervals for a total elapsed time of three minutes. Depending on the initial conditions, this length of time corresponds to a 40–95% decrease in sample absorbance (i.e., the reaction is between 40% and 95% complete). Absorbance values can be converted into ozone concentrations (molecules/cm3) if the path length of the cell and the photoabsorption cross section for ozone are known. (Reference 2 contains a useful table of temperature-dependent O3 cross sections; σ = 1.41 × 1017 cm2 at λ = 254 nm and T = 298 K.) The concentration of the injected organic reactant can be calculated if the initial pressures of the O2 (main body) and the organic (sidearm) are known, along with the relative volumes of the main body and sidearm and the temperature of the cell.
Journal of Chemical Education • Vol. 77 No. 9 September 2000 • JChemEd.chem.wisc.edu
In the Laboratory
For this experimental method to be successful, the organic reactant must satisfy the following criteria:
represented by
d O3 = k 1 + k 2 × organic dt
1. It must have a negligible photoabsorption cross section at λ = 254 nm. 2. The products formed in the reaction with O3 must have negligible photoabsorption cross sections at λ = 254 nm.
d ln O3 = k 1 + k 2 × organic dt
4. The rate of reaction with O3 must be significantly faster than the rate at which the O3 is removed by reactions at the cell walls. 5. The rate of reaction with O3 must be significantly slower than the rate at which the reactants initially mix.
A careful review of the literature (1, 3–5) and numerous experiments in our laboratory have indicated that vinylidene chloride (1,1-dichloroethylene, H2C=CCl2) is the best choice for the organic reactant. Acetylene may also be used, but its rate of reaction is fast enough that the observed reaction kinetics can be perturbed by mixing artifacts (see discussion below). Results and Discussion Before the organic reactant is injected into the main body of the cell, O3 decomposes via a slow wall reaction: rate constant = k1
(1)
After the organic reactant is injected, a second removal pathway becomes available: O3 + organic → products
rate constant = k2
(2)
The net rate at which ozone is removed from the reaction cell after the addition of the organic reactant is therefore 37.5
36.0
37.0
35.8
36.5
35.6
36.0
35.4
35.5
35.2
35.0
35.0
34.5
34.8
34.0
(4)
To analyze the data according to eq 4, each individual decay measurement is first plotted as ln [O3] vs time, and the data are subjected to a linear regression analysis. The y intercepts obtained from these regressions correspond to the initial concentrations of O3 (specifically, ln [O3]) present in the cell at time t = 0, and the slopes (i.e., d ln [O3]/dt) correspond directly to the left-hand side of eq 4. A plot of the individual logarithmic decay slopes as a function of [organic] yields a slope of k 2 and a y intercept of k 1. For this scheme to work properly, the cell absorbance at infinite time must be effectively zero (both the organic reactant and the products must have negligible photoabsorption cross sections at λ = 254 nm). If the absorbance does not approach zero, the pseudo-first-order logarithmic decay plots may develop curvature at longer times, distorting the linear regression analysis. In addition, it is usually more accurate to determine k 1 directly from an ozone-only decay measurement than by extrapolation to [organic] = 0. (The ozone-only decay measurement proceeds much more slowly and with much less noise than the ozone + organic measurements, and cannot be perturbed by mixing artifacts.)
ln[O3]
ln[O3]
(3)
If a large stoichiometric excess of organic reactant is present, the concentration of the organic reactant remains effectively constant and the reaction takes place under pseudo-first-order conditions. Equation 3 can be rewritten as
3. It must be a gas or a liquid with a high vapor pressure at room temperature.
O3 + wall → loss of O3
O3
34.6 0
50
100
150
200
Time / s Figure 2. Logarithmic decay plot for O3 + C2H2. The data correspond to initial oxygen (main body) and acetylene (sidearm) pressures of 207 and 504 torr, respectively, and a temperature of ca. 298 K. The ozone concentration is in units of molecules per cubic centimeter. Significant deviations from pseudo-first-order behavior are observed because the mixing rate of the O3 and C2H2 reactants is not sufficiently fast compared to the rate of reaction.
0
50
100
150
200
Time / s Figure 3. Logarithmic decay plot for O3 + H2C=CCl2. The data correspond to initial oxygen (main body) and vinylidene chloride (sidearm) pressures of 205 and 503 torr, respectively, and a temperature of ca. 298 K. The ozone concentration is in units of molecules per cubic centimeter. Deviations from pseudo-first-order behavior are small compared to those in Figure 2 because the O3 + H2C=CCl2 reaction is roughly two to three times slower than O3 + C2H2.
JChemEd.chem.wisc.edu • Vol. 77 No. 9 September 2000 • Journal of Chemical Education
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In the Laboratory 0.008
0.014
0.012
0.006
−d ln[O3]/dt
−d ln[O3]/dt
0.010
0.008
0.006
0.004
0.004
0.002 0.002
0
0 0
1
2
3
4
0
Figure 4. O3 decay rate as a function of acetylene concentration. The squares and triangles correspond to temperatures of ca. 298 and ca. 288 K, respectively. Decay rates appear to be perturbed by mixing artifacts (i.e., for high reactant concentrations the observed decay rates plateau because the rate of reaction is limited by the rate at which the reactants mix).
Logarithmic decay plots for acetylene and vinylidene chloride are presented in Figures 2 and 3, and decay data for a range of acetylene and vinylidene chloride concentrations are summarized in Figures 4 and 5. The logarithmic decay plots for the acetylene reactant (Fig. 2) show pronounced deviations from pseudo-first-order behavior that become more severe as the acetylene concentration is increased. These deviations probably result from inadequate initial mixing of the reactants (i.e., the mixing rate of the reactants is too slow compared to the rate of reaction) and yield logarithmic O3 decay rates that plateau for acetylene concentrations above 1.7 × 1018 cm3 (Fig. 4). The logarithmic decay plots for the vinylidene chloride reactant (Fig. 3) deviate only slightly from the expected pseudo-first-order behavior, suggesting that the reaction kinetics are not seriously perturbed by mixing artifacts. This conclusion is consistent with the logarithmic O3 decay data in Figure 5, which are nearly linear with respect to the concentration of the vinylidene chloride reactant. Regression analysis of the Figure 5 data for T ≈ 298 K yields k 2 = 2.7 × 1021 cm3 molecule1 s1, which is roughly consistent with the literature value of 3.7 × 1021 cm3 molecule1 s1 (5). Cautionary Notes In order for these types of kinetics measurements to be quantitatively accurate, initial reactant concentrations must be known to a reasonable degree of accuracy, the temperature of the reaction cell must be carefully controlled, and the reactants must be thoroughly mixed at a rate that is rapid compared to the rate of reaction. In the present experiment the latter two requirements are probably not met, since the reaction cell is not thermostated and mixing artifacts appear to perturb the acetylene data. While these problems can be addressed by modifying the design of the reaction cell, such modifications are likely to increase the cost and complexity of the experiment. The version of the experiment described 1212
0.5
1.0
1.5
2.0
2.5
3.0
[vinylidene chloride] / (1018 molecule/cm3)
[acetylene] / (1018 molecule/cm3)
Figure 5. O3 decay rate as a function of vinylidene chloride concentration. The squares correspond to T ≈ 298 K, and the triangles and circles correspond to two separate sets of measurements at T ≈ 286 K. Decay rates are nearly linear with respect to the concentration of the vinylidene chloride reactant, which suggests that the reaction kinetics are not seriously perturbed by mixing artifacts. Regression analysis yields k2 (298 K) = 2.7 × 1021 and k2 (286 K) = 1.8 × 1021 cm3 molecule1 s1.
here utilizes a very simple cell design to minimize the cost and complexity and to encourage students to think about how the apparatus or procedure can be improved to yield more quantitative results. Acknowledgments This experiment was initially conceived at the workshop “Developing Your Physical Chemistry Laboratory” sponsored by the National Science Foundation and hosted by Richard Schwenz of the University of Northern Colorado in the summer of 1995. We wish to thank John Thoemke for many useful suggestions and the reviewers for their helpful comments. Additional financial support was provided by the Saint Mary’s College Faculty Development Fund. Supplemental Material The complete description of this experiment and supplemental materials are available in this issue of JCE Online. W
Literature Cited 1. Atkinson, R.; Carter, W. P. Chem. Rev. 1984, 84, 437–470, and references therein. 2. Finlayson-Pitts, B. J.; Pitts, J. N. Atmospheric Chemistry: Fundamentals and Experimental Techniques; Wiley: New York, 1986; p 143. 3. NASA Panel for Data Evaluation. Chemical Kinetics and Photochemical Data for Use in Stratospheric Modeling; Evaluation No. 11, JPL Publication 94-26; Jet Propulsion Laboratory/ California Institute of Technology: Pasadena, 1994; p 49. 4. Atkinson, R.; Aschmann, S. M. Int. J. Chem. Kinet. 1984, 16, 259–268. 5. Hull, L. A.; Hisatsune, I. C.; Heicklen, J. Can. J. Chem. 1973, 51, 1504–1510.
Journal of Chemical Education • Vol. 77 No. 9 September 2000 • JChemEd.chem.wisc.edu
