The Kinetics of the Nitrogen Dioxide-Hydrogen Chloride Reaction Ching-Chuan C. KUO,Robert A. Wilkins, Jr., and 1. C. Hisatsune' Department of Chemistry, Davey Laboratory, The Pennsylvania State University, University Park, Pennsylvania 16802
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The reaction between gaseous NOz and HCI, whose stoichiometry had been reported to be 2NOZ 4HCI = 2NOCI 4- 2H20 C12, has been found to consist of three separate sequential reactions: 2N02 HCI = NOCl HN03, HN03 HCI = NOpCl HzO, and NOzCl+ 2HCI = NOCl H20 Clz. These reactions are catalyzed by HN03 and H20 through heterogeneous reaction paths having negative temperature dependences. The active reaction sites in these heterogeneous paths appear to be "03 dimers and polymers which are hydrated in the presence of HzO. Only the first reaction with an equilibrium constant of 645 f 79 (atm units) at 25 O C has an appreciable homogeneous rate that is second order in NOz and first order in HCI. The experimental rate equations for these reactions, their temperature dependences, and reaction mechanisms are presented.
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Introduction Gas phase kinetic data on reactions involving HNO3 particularly in the presence of H20 are not readily available because of experimental difficulties in handling and characterwhich adsorbs on the walls of the reaction izing wet "03 vessels or forms aerosols. England and Corcoran (1974) rewhich are cently reviewed several reactions involving "03 pertinent to the chemistry of the atmosphere and air pollution. They also reported a kinetic study of the gas phase reaction between H20 and NO2 whose equilibrium stoichiometry is given by 3N02
+ H20 = NO + 2HN03
(1)
Nitric acid mists and two phases are formed in this reaction unless the reactant pressures are sufficiently low, and previous studies being carried out under such conditions did not clarify the kinetics. The homogeneous rate-determining step of reaction 1 was found to be 2NO2(or N204)
+ H20 = HONO + "03
+ H20 = HON02 + N20j + HC1= ClNOz + "03
N20j
"03
(3) (4)
Consequently, a similar equivalence was expected between reaction 2 and the reaction 2NO2(or N204)
+ HC1= ClNO + HNO3
(5)
However, in earlier studies of reaction 5 by Harris and Siege1 (1941) and by Talbot and Thomas (1959), H N 0 3 was not identified and the stoichiometry was given instead as 2N02
+ 4HC1= 2NOC1+ 2H2O + Cln
(6)
According to Talbot and Thomas (1959),reaction 6 was heterogeneous and in a glass vessel a t 25 to 55 "C the rate was first order depending only on N204 with a rate coefficient of 108 exp(-12.9 kcallRT) s-l. At 250 to 420 "C, Rosser and Wise (1960) and Gilbert and Thomas (1963) observed this reaction to be homogeneous with NO, Clp, and H 2 0 being formed as products. In this case the rate was first order in both NO2 and HC1 and the second-order rate constant was exp(-23.4 kcallRT) M-l s-l. 236
Ind. Eng. Chem., Fundam., Vol. 15, No. 4, 1976
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Because of potential importance of reaction 6 in the chemistry of air pollution, we have reexamined it by infrared spectroscopy and have found, as we shall describe here, that reaction 5 does indeed take place and can be made to be the sole equilibrium reaction. Moreover, reaction 6 was the consequence of reaction 5 followed by HNO3
+ HCl = N02Cl+ HpO
(7)
and NO2Cl+ 2HC1= NOCl
+ H20 + C12
(8)
The gas phase ' inetics of neither reaction 7 nor reaction 8 has been studied before but Collis et al. (1958) have confirmed reaction 8 in the liquid phase while Bursa and Straszko (1962) have studied extensively the reaction between H N 0 3 and HC1 in aqueous solutions. The latter authors reported the stoichiometry and the rate in the range of 7 to 90 "C to be, respectively
"03
(2)
with "02 decomposing rapidly in a subsequent step into NO, NO2, and H20. Among the reactions of N20j in which "03 is formed, there is the following equivalence between the reaction with H20 and that with HC1 (Wilkins, 1973).
+
+ 3HC1= NOCl + 2H20 + C12
(9)
and rate = hg(HN03)8 (HCl)
(10)
with hg being 4.4 X lop4exp(-9.80 kcallRT) Mps s-l. Experimental Section The gases NO, Clz, and HCl were obtained from the Matheson Co. and were subjected before use to low-temperature distillations in the vacuum line. The NO2 and NOCl were prepared by reacting NO with Op or C12. Anhydrous H N 0 3was in a soobtained from the distillation of concentrated "03 lution of concentrated H2S04 following the method of was also used to Nightingale et al. (1954). Distilled "03 prepare N02Cl according to the procedure of Schumacher and Sprenger (1929). Chlorosulfonic acid required for the latter preparation was from Matheson Coleman and Bell Co., and it was distilled prior to use. Sample gas transfers and pressure measurements were made in a conventional glass vacuum line having stopcocks lubricated with Fluorolube GR-90 grease from Fisher Scientific Co. Pressures of reactive gases were measured with an oil manometer containing Kel-F No. 3 oil (3M Co.) while a mercury manometer was used for nonreactive gases. A cathetometer was used to read the manometer column heights to a precision of f0.05 mm. Gas samples were analyzed by infrared spectroscopy using absorption bands identified in our earlier studies on NO2 (Hisatsune and Zafonte, 1969), NOCl (Hisatsune and Zafonte,
I
I
I
\
o
NOCl
\.
A
N02CI NOz
- \
A
"OS
I
-'A
-
1
1
4
L
-
u)
6-
\.
W K 0
1
P
W
a TIME
@in)
Figure 1. Changes in concentrations of nitrogen compounds during t h e gas phase reaction of NO2 and HCl. Experimental conditions: reaction temperature = 24 "C, (N02) = 6.4 Torr, (HCI) = 61.5 Torr, (Nz) = 420 Torr.
19691, HNO3 (McGraw et al., 1966), and N0&1 (Bernitt et al., 1967). Our spectrophotometers were Perkin-Elmer instruments Model 112 (calcium fluoride or sodium chloride prism), Model 108 (rapid scanning spectrophotometer with calcium fluoride prism), Model 225 (gratings), and Model 521 (gratings). The operating characteristics of the rapid scanning infrared spectrophotometer is described elsewhere (Hisatsune and Zafonte, 1969; Hisatsune, 1969). The kinetic cell assembly, which was also described before (Hisatsune and Zafonte, 1969), consisted of two 100-ml reactant bulbs connected by a T-stopcock to an absorption cell which served as the reaction vessel. The cell length was limited to 6.20 cm by the size of the compartment in the Barnes Engineering Co. Model 104 variable temperature chamber. Windows of CaF2 or AgCl were heat-sealed to the cell with a mixture of Halocarbon high-temperature grease and Series 15-00 wax (Halocarbon Corp.). Three Pyrex glass cells were employed each with a capacity of about 100 ml. Cell I had a surface-to-volume ( S / V )ratio of 1.2 cm-l while cell I11 had the same S / V but its inner glass wall was coated with Teflon by Chemplast, Inc. (Wayne, N.J.). In cell I1 glass wool (Corning Co. 7220) was packed around the inner wall of the cell leaving the optical path clear; the S/V for this cell was 12.9 cm-l. In some kinetic runs the two reactants were placed separately in the reactant bulbs and then pressurized with nitrogen. The reaction was initiated by opening the T-stopcock which permitted the reactants to mix and flow into the evacuated absorption cell. The spectrum of the reaction system was monitored continuously or intermittently either by repetitive scans of one or more absorption bands or by recording the changes in absorbancy a t one frequency. In other kinetic runs both reactant bulbs were filled with a single reactant and then pressurized with nitrogen. The second reactant was placed a t a lower pressure directly in the absorption cell. Limits in reactant pressures were imposed by absorption coefficients of the infrared bands monitored and by the vapor pressures of the various substances. These vapor pressures, the thermal stabilities of the various gases, and the softening point of the window sealing wax restricted the reaction temperature in our study to a range from about 0 "C to near 60 "C. The experimental procedures and data are presented in detail elsewhere (Wilkins, 1973: Kuo, 1974). Results Reaction Stoichiometries. The infrared spectra of several mixtures of NO2 and HCl revealed immediately that eq 6 is the overall stoichiometry of a sequence of three separate reactions. Figure 1 illustrates a study with HC1 in excess over NO2 where the growth of NOCl was followed first for 30 min
-
I
-.flab./\
*',aC1-f\r 0
IO0
6 h -f-
zoo
300
and then changes in concentrations of all nitrogen compounds present in the reaction mixture were determined by scanning the entire spectrum repeatedly until the reaction was over. The final spectrum of this system showed only NOCl. The stoichiometry during the initial rapid formation of NOCl was followed with the rapid scanning spectrophotometer and it showed the formation of one mole each of HNO:)/dt
was also verified experimentally. In this case 3.88 Torr of NO? was reacted with 107 Torr of HC1 in 354 Torr of N2 at 25 "C. Ind. Eng. Chem., Fundarn., Vol. 15, No. 4, 1976
239
Table I. Arrhenius Frequency Factors ( A ) and Activation Energies ( E )for Eq 13,17, and 18
Rate constant
log A (units)
-0.94 f 0.43 (M-' s-l) 6.37 f 0.47 (M-I s-l) 2.55 f 0.94 (M-l) -2.27 f 0.51 (M-4 s-') c5 ~s(Nz04)' 5.08 0.51 (M-3 s-l) ds -7.77 f 0.11 (M-') -11.13 f 0.73 (M-3 s-') a7 -5.75 f 1.26 (M-5 s-l) b7 -0.29 f 0.63 (M-' s - ~ ) as bsb 0.44 (M-l) N 2 0 4taken as the reactant. From kinetic runs at 25 and 57 "C. as
adNz04)' b F,
I o3
I
Rate constant at 25 "C
-6.60 f 0.59 +6.38 f 0.64 +0.34 f 1.26 -19.7 f 0.7 -6.63 f 0.68 -14.3 f 0.2 -22.3 f 1.0 -25.2 f 1.7 +1.73 f 0.88 -2.9
7.9 X lo3 M-2 s-l 49 M-' s-l 197 M-l 1.43 X 10l2M-4 s-l 8.9 x 109 ~ - s-l3 510 M-' 1.64 X lo5 M-3 s-l 6.0 X 10" M-j s-l 2.9 x 10-'M-1 s-1 400 M-'
i
/
!
E , kcal/mol
/
i
i $ 0 .
-2.5
I
IO21
I I
I
I
3.0
3.2
3.4
I
0 085 log(HN03) (Torr)
1
1.0
I
3.6
3.8
1 0 3 1 ~ (OK-')
Figure 9. Arrhenius plots for b5 (closed circles) and d5 of eq 13.
The reactant NO2 and the products NOCl and HNO:l all were followed simultaneously with the rapid scanning spectrophotometer using 6 stands. The experimental data gave second-order rate constants of 15.4, 17.8, and 15.5 M-' s-l from NO*, NOCl, and "03, respectively, which show satisfactory internal consistency considering the experimental uncertainties. A Pyrex reaction cell with S / V = 1.2 cm-' was used most often but four kinetic runs were made with cell I1 ( S / V = 12.9 cm-1) and two with Teflon-coated cell I11 ( S / V = 1.2 cm-'1, all a t room temperature. Cell I1 gave a:, = 2.1 X lo* M-' SKI and c j = 5.9 X 1 O I 2 M-4 s-l while the corresponding rate constants from cell I11 were, respectively, 1.39 X lo4 M-'s-I and 7.2 X 1OI2 MP4 s-l. These values of us may possibly be within the experimental error limits of a5 = 0.79 X lo4 MP2 s-l obtained with cell I but there is little doubt that the "Os-dependent c g term (1.43 X 1012 M-4 s-1 in cell I) is affected by changes in reaction vessel. The Kinetics of Reverse of Reaction 5. The rate of reverse of reaction 5 under excess "03 was initially first order in NOCl but deviated from this order near the end of the reaction due to the equilibrium being established. Correction for this equilibrium using eq 13 in a finite rate form gave data which followed a first-order decay law throughout the entire reaction time but the resulting rate constants were strongly dependent on the amount and dryness of HN03. The order with respect to the acid with excess NOCl was second and the 240
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-0.5
Ind. Eng. Chem., Fundam., Vol. 15, No. 4, 1976
Figure 10. The dependence on " 0 3 of the reverse of reaction 5. See text for the description of the opened circles. The solid and dashed lines have slopes of 2 and 3, respectively.
rate showed a negative temperature dependence in the range of 5 to 25 "C. Near 55 "C however, the rate was much faster than that expected from the lower temperature results and, furthermore, was now first order in "03 even though the reaction stoichiometry remained unchanged. Small amounts of N0&1 were also formed but they were expected and in agreement with the amounts calculated from the known equilibrium (Ray and Ogg, 1959) NOCl
+ NO2 = NO + N02C1
(14)
Traces of NOZC1 were also generated when "03 was in large excess but they were attributed to reaction 7. Thus, the rate of reverse of reaction 5 was represented by rate = -d(NOCl)/dt = [a-j(HN03) + b-5(HN03)'](NOCl)
(15)
where the second term was dominant a t room temperature and below. Just as the apparent order with respect to "03 in reaction 5 increased when the acid pressure became appreciable as shown in Figure 6, the reverse of reaction 5 also showed a variable order with respect to "03. Figure 10 illustrates a log-log plot of rate/(NOCl) as defined by eq 15 vs. ("03) for kinetic runs made a t 25 "C. Only the first half of the data in each run was used making corrections due to contributions of the forward reaction unnecessary. For "03 above about 3 Torr, the reaction order with respect to this acid is closer to
Table 11. Rate Constants and Arrhenius Parameters for Eq 15 Temperature, "C
a-5 (M-' s - ~ ) b-5 X (M-2 s-l) a-5
5
25
56
10.0 f 0.1
0.22 i 0.03 1.76 f 0.41
1.20 f 0.28
= 1.3 X lo7 exp(-10.6 kcal/RT) M-l
b-5 = 8.1 X
-
s-l
exp(+15.4 kcallRT) M-2 s-l
third than to second but the data points represented by opened circles are on a line with a slope of 2. The latter data = 7.08 Torr but this acid are from a kinetic run with ("OB) had been distilled twice from fresh concentrated HzSO4 shortly before. The remaining data in this figure are from runs in which acid samples stored over a period of time had been used. Evidently, the dryness of the acid affected the reaction rate and the presence of moisture in the acid caused an increase in order from second to third. The experimental conditions under which a-5 and b-5 were determined are described elsewhere (Wilkins, 1973). The reaction temperatures ranged from 5 to 56 "C and the gas pressures were 1.41 to 192 Torr and 1.48 to 10.2 Torr, reThe numerical values of the spectively, for NOCl and "03. rate constants and the Arrhenius parameters are summarized in Table 11. As in Figures 7-9, the main cause of scatter of the experimental data in this table was the strong dependence of the rates on HNO:3 and on its dryness. T h e Kinetics of Reaction 7. As we described in the stoichiometry section, reaction 7 was isolated from reaction 8 by in excess. In a kinetic run with (HN03) = 27.6 using "03 Torr, (HC1) = 4.59 Torr, and (N2) = 447 Torr a t 24.8 "C, the rate was observed to accelerate at the initial stage of the reaction. However, when 2.10 Torr of H20 was added to essentially the same reation mixture, no acceleration in rate was observed although the rates were now higher. Neither was this was decreased initial increase in rate evident when "03 below 20 Torr. The reaction order with respect to HC1 was determined from several runs in which H N 0 3 was kept constant at 19.0 Torr and HC1 was varied from 1.87 to 19.0 Torr. The rates in each of these kinetic runs were constant during the first few minutes and showed no dependence on HCl over the pressure range used. Thus, the rate of reaction 7 was zero order in HC1. The order with respect to H N 0 3 on the other hand was deduced from a log-log plot of the initial rates vs. "0.7 which gave a least-square slope of 3.98 & 0.26. Cell I ( S / V = 1.2 cm-') was used in these runs but in a single experiment with cell I1 ( S / V = 12.9 cm-l) the "03 order was still fourth. The dependence of the rate on H2O was deduced from kinetic runs in which HNO,I and HCl were kept essentially constant at 19.0 and 2.0 Torr, respectively, and H20 was varied from 0.44 to 2.68 Torr. A plot of initial rates against H 2 0 indicated a rate term first order in H20 as well as a term fourth order in "03. Addition of 2.57 Torr of N02Cl to the reaction mixture initially or increasing HCl to 7.02 Torr did not affect this H2O dependence. In all of these kinetic runs, HzO and HNO:j were placed directly in the cell while HC1 pressurized with N2 was placed in the reactant bulbs. Interestingly, when He0 was added to anhydrous "03, the pressure of the acid dropped as evidenced by decreases in intensity of H N 0 3 infrared bands at 891 and 1720 cm-'. The relationship among the initial pressure of the anhydrous acid ("Os),,, the actual observed spectroscopically, and pressure of the acid ("03) the water pressure (H20) was found experimentally to be
t
*
i
/
7
/" 3 3,O
I
3.2 103/T
,
I
3,4
3.6
(OK-')
Figure 12. Arrhenius plot for a7 of eq 17. where a was 0.032 T o r r 1 at 25 "C and 0.22 Torr-' at 0 "C for H20 pressures up to about 6 Torr. The Pyrex cell I was used in these studies. Plots of initial rates vs. (H20) described above gave apparent first-order rate constants b*; with respect to HLOfor reaction 7. However, these constants still contained the HNO 1 dependence which was found to be fifth order in the acid as illustrated in Figure 11. The rate of reaction 7 therefore was
+
rate = d(NO,Cl)/dt = a ~ ( H N 0 3 ) ~b;(H20)(HNOj)i (17) and the observed initial acceleration in rate was due to the formation of water and to the second term in this equation. Reaction 7 was studied (Kuo, 1974) over the temperature range of 0 to 59 "C with (HCl) = 1.87 to 19.6 Torr, (HNO 9) = 5.62 to 31.0 Torr, and (H20) up to 5.79 Torr. One run was made with NOYCl added initially but the kinetics showed no changes. The rate constants a: and b7 are summarized as Arrhenius plots in Figures 12 and 13 and the least-square Ind. Eng. Chern., Fundarn., Vol. 15, No. 4, 1976
241
14
1
i/l
i
/
1
n
5
0
1
3.2
3.0
3
