J. Phys. Chem. 1981. 85.2303-2380
(t-Bus) is in keeping with recent measurementslc~ld,lf~lo which give values of 8-12 kcal mol-l. Possible Influences of Weak Collisions. In this and most previous VLPP analyses high-pressure rate constants are derived by assuming "strong" gas-wall collisions.6 However, recent analyses by King and GilberP indicate that this assumption is not always valid in VLPP experiments. After examining several reactions, they found that to account for weak collisional effects experimental collision frequencies in VLPP experiments should be multiplied by an efficiency factor, P, of the order of one-half. Furthermore, more detailed considerations indicate that P decreases with increasing temperature^.'^ At the present time, however, we cannot reliably account for weak collision effects in our VLPP experiments for the following reasons: (1) The dependence of P on surface type is not clear. Walls in these and previous studies2of substituted benzene decomposition are coated with a shiny, carbonaceous deposit. It is not clear whether this type of surface is present in experiments aimed at determining collisional efficiency. (2) The dependence of P on molecular type is not known. Molecules studied in this work are larger and may interact more strongly with the walls than the molecules examined by King. The effect, if any, of aromaticity on ,f3 is also unknown. (3) Reactions reported here are not as far in the falloff as reactions examined in experiments aimed at determining collisional efficiency. At mean reaction temperatures, assuming /3 = 1, k/k, falls in the range 0.3-0.4for the reactions reported here. It is worthwhile, however, to consider the possible influence of weak collisional effects on our results. If, for instance, /3 = 0.5 rather than 1 in the middle of the temperature ranges of the present experiments, then estimated high-pressure rate constants, k,, are raised by -50%. (12) Gilbert, R. G.; King, K. D. Chem. Phys. 1980,49,367. (13)(a) Kelley, D.F.; Barton,B.D.; Zalotai, L.; Rabinovitch, B.S. J. Chem. Phys. 1979, 71,538. (b) Kelley, D. F.; Kasai, T.;Rabinovitch, B. S. J. Phys. Chem. 1981,85,1100.
2383
This uncertainty is comparable to the uncertainty in k, due to possible errors in the choice of Arrhenius parameters (higher Arrhenius parameters lead to higher estimates of kJ. Weak collisions would decrease the slope of the theoretical lines in Figures 1, 3, and 5; this decrease in slope would tend to very slightly worsen the fit to our data. In Figure 2, if @ = 0.5 VLPP-derived k, values would be raised above carrier-derived values by ca. 50% in their region of overlap. However, considering the combined uncertainties in carrier data, VLPP data, and assumed Arrhenius parameters, this change in k, would not be very significant. Weak collisional effects would increase the difference between shock tube and VLPP results for isobutylbenzene homolysis (Figure 4). The difference between rate constants obtained from these type types of experiments is, therefore, not an artifact of weak collisions in VLPP experiments. In any case, since /3 values are not expected to differ significantly from one alkylbenzene homolysis to another, weak collisional effects can have only minor effects on relative rate values derived in this work. Summary Kinetics of propylbenzene dissociation determined in the present VLPP experiments are in excellent agreement with those obtained by carrier methods. For isobutylbenzene dissociation, the present experiments yield rate constants 2.5 times greater than those determined by using the comparative shock tube technique; however, the two techniques do agree on the decomposition rate of cyclohexene which is used as a comparison standard in the shock tube studies. Based on rate parameters derived for propylbenzene and isobutylbenzene dissociation, rate parameters for neopentylbenzene dissociation indicate that the enthalpy of formation of tert-butyl radicals lies in the middle of the range of recently reported values.
Kinetics of Oxidation of Aqueous Bromide Ion by Ozone Kenyu Haruta* and Tetsu Takeyama Product Development Laboratory, Mitsubishi Electric Corporation, Amagasaki, Hyogo 66 1, Japan (Received: January 8, 198 1; In Final Form: April 20, 198 1)
-
The rate of ozone disappearance in the reaction between aqueous bromide ion and ozone to form hypobromite ion, Br- + O3 BrO- + O2 (l), has been measured spectrophotometricallyat temperatures of 5-30 "C, pHs of 1.2-3.6, bromide concentrations of 5 X 10-5-2 X M, and ozone concentrations of 1 X 10-5-5 X M. The rate equation obtained is -d[O,]/dt = (kl0 + kl*[H+])[03][Br-],M s-l, where :k = 4.9 X log exp(-(10000 f 1000)/RTJ, M-ls-', and kl* = 1.7 X 10l2exp(-(11000 f 800)/RT),M-2 s-'. The second-order rate expression applies also in neutral solution. In an alkaline solution, where spontaneous decomposition of ozone occurs, the rate for ozone disappearance is expressed by a linear combination of the above equation and the first-order ozone decomposition rate. The mechanism of the overall reaction is also discussed.
Introduction we have demonstrated recently that hypobromite ion, BrO-, is a very effective agent for the removal of odorous gases from sewage treatment plants and that reaction 1 Br-
-
+ O3
BrO-
+ O2
However, few papers - - have been published so far on the kinetics of this reaction. Taube2 studied the reaction of Ozone with hydrogen peroxide in the presence of Br- and H+ at 0 "c by measuring the overall rate for the disappearance of the sum of ozone and hydrogen peroxide by
(1)
can be used advantageously for the in situ production of BrO- in a deodorization process.' 0022-3654/81/2085-2383$01.25/0
(1) K. Azuma, N. Matsunaga, K. Doi, A. Ikeda, and K. Haruta, h o c . Int. clean~ i , congr., . 5th,1980. (2) H.Taube, J.Am. Chem. SOC.,64, 2468 (1942).
0 1981 American Chemical Society
2384
The Journal of Physical Chemistry, Vol. 85, No. 16, 1981
Haruta and Takeyama
iodometry. Assuming an equimolar reaction between ozone and hydrogen peroxide, he found the rate for the disappearance of ozone to be expressed by and evaluated kl,the rate constant for reaction 1, to be 1600 f 100 M-' min-' over a narrow pH range of 2.6-3.0. In the present study, the rate constant for reaction 1 was determined directly by following spectrophotometrically the disappearance of ozone in the reaction of ozonized water with aqueous sodium bromide solution over a temperature range of 5-30 "C and a pH range of 1.2-9.0. Experimental Section Apparatus and Procedure. The experimental apparatus consisted of an ozone saturator and a reaction cell installed in one of the light paths of a double-beam spectrophotometer. The ozone saturator was a cylindrical glass vessel, 15.5 cm long and 9.5 cm in internal diameter (1500 mL), with a thermometer, a pH sensor, a magnetic stirrer, an ozone diffuser, and two tapered joints; one of these joints was connected to the exhaust valve and the other to the reaction cell by a capillary tube through a connecting valve. The reaction cell was of quartz, cylindrical in shape, with a light path of 10 cm and a capacity of 38 mL. Another cell of the same dimension was used as a reference. The air surrounding these cells was replaced with nitrogen during the measurements so that water vapor in the air could not condense on the cell windows. Redistilled water (1000 mL) acidified with sulfuric acid was put into the ozone saturator and stirred, with the exhaust valve open and the connecting valve closed. Ozonized oxygen (30 mg of O3min-l) was bubbled through the stirred solution maintained at a certain temperature and pH chosen for the experiment. After 15 min, with ozonized oxygen still bubbling, the exhaust valve was closed. Thus, the pressure in the ozone saturator increased rapidly to a gauge pressure of 76 mmHg, when the connecting valve was opened to introduce the ozonized water into the reaction cell which contained an aqueous sodium bromide solution (1 mL) of predetermined concentration. Then, the connecting valve was closed manually when the reaction cell was filled with the ozonized water, and the variation in ozone concentration was followed spectrophotometrically by measuring its absorption at 260 nm. Measurements. The extinction coefficient of ozone at 260 nm, the absorption maximum, was measured in a sulfuric acid solution. The arithmetical average of six measurements under the same conditions was (2.95 f 0.12) X lo3 M-' cm-l, which is in a good agreement with literaA remarkable absorption was observed at ture 390 nm after the reaction was completed, which can be attributed to bromine molecule^.^ However, no other absorptions, except the one at 330 nm due to BrO- formed in the reaction in an alkaline medium, were appreciable in the wavelength range of 240-400 nm during the reaction period. The initial concentrations of ozone and bromide ion were calculated from the mixing ratio of the ozonized water and the aqueous sodium bromide solution. The amount of bromide ion consumed in the reaction was calculated from that of ozone consumed on the assumption that 2 mol of bromide ion reacted with each mole of ozone. This assumption held true satisfactorily under the con(3) M. L. Kilpatrick, C. C. Herrick, and M. Kilpatrick, J . Am. Chem. Soc.. 78. 1784 (19.519. ~
(4) M. Griggs, J. Chem. Phys., 49, 857 (1968). Lon(5) R. G. Aickin, N. S.Bayliss, and A. L. G. Rees, Proc. R. SOC. don, Ser. A , 169, 234 (1939).
3 3
3 4
3 5
3.6
37
Flgure 1. Rate constant of spontaneous ozone decomposition in aqueous solution: (A)Stumm (ref 6); ( 0 )Czapski (ref 7); ( 0 )Rizzutl (ref 8); (V)Teramoto (ref 9) at pH ca. 9; (A)Horie (ref 10); (m) Morooka (ref 11); (V)calculated from Hoigne (ref 12) at pH ca. 7; (0 and a)) this study at pH 7 and 9, respectively.
ditions of present experiments with only a few exceptions as will be shown later. The time required to fill the reaction cell with ozonized water and complete the mixing of ozonized water with sodium bromide solution was measured several times, by using redistilled water and sodium hypobromite solution instead of ozonized water and sodium bromide solution, respectively, and carring out the same procedure. The hypobromite concentration, measured by a UV absorption at 330 nm, decreased to a calculated value within a period of 3 s, which is negligibly short compared with the ozone-bromide reaction time of -2 min. The pH value hardly changed during the reaction time at pH below 3, but the temperature of the reaction mixture varied slowly because of the heat conduction from the surrounding atmosphere. Materials. Ozonized oxygen was prepared by silent discharge in 99.99% pure oxygen. Analytical-grade sodium bromide, sulfuric acid, sodium hydroxide, and sodium sulfate were used without further purification. Sodium sulfate was used to make the ionic strength of the reaction mixtures 0.1 M. Results Decomposition of Ozone i n Aqueous Solution. In order to clarify the magnitude of the spontaneous decomposition of ozone, preliminary experiments were carried out at 20 O C and pH 3, 7, and 9 with initial ozone concentrations, 1.8 X and 7.5 X lo4 M, respec[03],,, of 2.5 X tively. The decomposition of ozone obeyed the three-halvesorder law at pH 3 and 7. The rate constants determined from the slopes of straight lines in the plot of ([0,]-1/2 [0,]{1/2) vs. time, where [O,] is the ozone concentration at time t , were kD3 = 4.68 X lo-, M-'f2 s-' and kD7 = 4.48 X M-1/2s-l, respectively. The value of kDa is negligibly small compared to the rate constant for the reaction of ozone with bromide ion, as will be shown later. On the other hand, the decomposition of ozone at pH 9 was found to follow the first-order law and the rate constant, kDe, to be 3.69 X s-l from the straight line obtained in the plot of In ([03],,/[03]) vs. time. The values of k ~ and 7 kDo obtained in the present study are plotted in Figure 1 together with several literature values612 as a function of reciprocal temperature. The full (6) W. Stumm, Helu. Chim. Acta, 37, 773 (1954). (7) G. Czapski, A. Samuni, and R. Yelin, Isr. J. Chem., 6, 969 (1966). (8) L. Rizzuti, V. Augugliaro, and G. Marruci, Chem. Eng. Sci., 31,877 (1976).
The Journal of Physical Chemistry, Vol. 85,No. 16, 198 1 2385
Oxidation of Aqueous Bromide Ion by Ozone
PH
T A B L E I: Stoichiometry of Reaction between Br- and 0, Measured at 27 "C
1 2 3 4 5
3.30 2.58 2.30 1.80 1.50
1.91, 1.87, 1.91, 1.99, 1.95,
1
5.36, 5.02, 6.61, 4.44, 5.33,
5.91, 4.91, 6.80, 4.20, 6.10,
I^
a
I
I
/-
1.5
I
I
0.91 1.02 0.97 1.06 0.87
'~ / i
--
2 0 I
3 0
0
1
I
I
1
1
I
8
16
24
32
40
48
56
64
[Htl X103 ( M I Figure 3. Dependence of second-order rate constants on hydrogen ion; 27 O C , p = 0.1 M, [Br-],/[0310 = 3-5.
1 0
IO
20
30
i
Time ( s e c - ' ) Flgure 2. Reaction of ozone with bromide showing second-order M, [O& = 1.966 X [Br-1, = 4.653 X dependence: (0) M; (0)[Br-1, = 1.012 X IO4 M, [O3lO= 1.932X M; (0) [Br-1, = 1.953 X IO4 M, [03], = 1.153 X M; pH, 3;temperature, 27 O C : p, 0.1 M.
lines drawn by the least-squares method demonstrate good agreement between the present data and literature values. The rate law at pH 7 should be a linear combination of the three-halves-order and the first-order rate laws. Thus, kD7 should be an apparent rate constant. Reaction of Ozone with Bromide Ion in Acidic Solution. Stoichiometry of the reaction was studied in the presence of excess bromide ion over a pH range of 1.5-3.3 by measuring the ratio of the concentration of the bromine produced in the reaction, [BrZIE,to that of ozone consumed, A[03]. The values of [BrZIEwere measured as [BrO-1; the reaction mixture was alkalized after the reaction was completed, and the UV absorption of BrO- was measured at 330 nm. As shown in Table I, the ratio of [Br2]E/A[03]is unity under the present experimental conditions. Therefore, the reaction of BrO- with ozone to form bromous acid and bromate is negligible. Several experiments were carried out to determine the reaction order with respect to ozone and bromide at 27 OC and pH 3 by varying the initial concentration ratio of bromide ion to ozone, [Br-],/[O,],, by -&fold. The results are plotted in Figure 2 according to eq I. Since a good
linearity of second-order plots exists over a wide range of the ratio of [Br-],/ [O3lO,ozone-bromide reaction conforms (9)M. Teramoto, H. Teranishi, and S. Imamura, Preprint of 11th Autumn Meeting of the Society of Chemical Engineers, Japan, Tokyo, A 206 (1977). (10) S. Horie, K. Seino, and H. Takeuchi, Yosui t o Haisui, 15, 345
4.0 3.3
8.4
2
I
3.4
3.5
io3/
T
3.6 (K-I)
Figure 4. Dependence of secondorder rate constants on temperature: p = 0.1 M: (0) pH 1.7; (0) pH 3.1.
closely to a second-order kinetics: the reaction rate is proportional to the first order of both ozone and bromide concentration. The rate constant, kl,obtained from the slope of the straight line was 2.28 X lo2 M-l s-l. To elucidate the effect of hydrogen ion on Itl, hydrogen ion concentration was varied over a range of 2.4 X 104-6.3 X M at 27 "C and [Br-]o/[03]0of ca. 3-6. The values of kl obtained in the same way as above are plotted against the concentration of hydrogen ion in Figure 3, which shows a definite effect of hydrogen ion. The straight line obtained by least-squares method is expressed by kl = klo + kl*[H+] (11) where k10 = 2.11 X lo2 M-l s-l and kl* = 1.74 X lo4 M-2 5-1,
The temperature dependence of k1 was measured over the range 5-30 "C at pH 3.10 and 1.70 and [Br-]o/[03]o of ca. 3-5. The results are given in Figure 4 as Arrhenius plots. At pH 3.10, the second term of eq I1 becomes an order of magnitude smaller than the first and may be neglected safely. Therefore, a plot of k1at pH 3.10 vs. 1 / T should give an activation energy, Elo,for the reaction path independent of hydrogen ion. The activation energy for the hydrogen-ion-dependentpath, El*,can be derived from the value of El0determined as above and the temperature dependence of k1 at pH 1.70. The rate constants, thus determined are klo = 4.9 X lo9 exp(-(10000 f 1000)/RT), M-' s-l (111)
(197.1) \____,_
kl* = 1.7 X 10l2 exp(-(11000 f 800)/RT],M-2 s-l (IV)
(11) S. Morooka, K. Ikemizu, and Y. Kato, Kagaku Kogaku Ronbunshu, 4, 377 (1978). (12) J. Hoigne and H. Bader, Water Res., 10, 1 (1976).
Taking into account the fact that the errors in El0arising from neglecting the second term of eq I1 affected directly
2380
The Journal of Physical Chemlstry, Vol. 85,No.
Haruta and Takeyama
16, 1981
-d[Br-]/dt = k,[Br-][O,]
0.0
were solved by use of series analysis employing the Stirling form of Taylor’s expansion, in which kDd was varied as a parameter to obtain the best matching between the calculated and observed curves of [O,] change. The results are presented in Figure 5. The circles represent the experimental data obtained at pH 9 and 20 “ c and the full lines numerical calculations for a k p ’ of s-l. The numerical calculations simulate the 3.45 X experimental observation to a considerably good extent over a wide range of the [Br-],/[03], ratio of 1.5-6. In s-l agrees well addition, the value for kDe’ of 3.45 X with kDa of 3.69 X s-l determined in the absence of bromide ion.
’
0
30
60
90
120
Time (sec)
Flgure 5. Disappearance of ozone in alkaline solution in the presence of bromide; circle, experimentalvalues; full line, simulation curve: (0) [Br-]o/[03]o = 1.5; (0)[Br-]o/[03]o = 6.0; ( 0 )[Br-]o/[O,]o= 4.0; temperature, 20 OC;pH, 9; p , 0.1 M.
the value of El*, one may conclude that the activation energies for the two reaction paths are approximately equal and that the catalysis of reaction 1by hydrogen ion leads to an increase in frequency factor rather than to a decrease in activation energy. On the basis of the activated complex theory, the entropies of activation at 25 “C were calculated to be -16.29 cal deg-l mol-I for klo and -4.59 cal deg-I mol-’ for kl*. These results are similar to those for the reaction of ozone with chloride,13in which the entropy of activation increased from -4.98 cal deg-’ mol-l for the hydrogen-ion-independent path to -0.68 cal deg-l mol-I for the hydrogen-ion-dependent path, while the activation energies of both paths were nearly equal to each other. Reaction of Ozone with Bromide Ion in Neutral and Alkaline Solution. Several experiments were done at pH 7 and 9 in much the same way; the amount of bromide ion consumed in the reaction was calculated from that of ozone consumed on the assumption that 1 mol of bromide ion reacted with each mole of ozone, since the main product was HBrO. The experimental results obtained at pH 7 are given in Table 11. The spontaneous decomposition rate of ozone, kDT[03]2/2,is ca. 4 X lo4 M s-l, which is less than 5% of the rate of reaction 1,i.e., ca. 1 X M s-l. Thus, the rate constant of reaction 1 at pH 7 was calculated, as in the preceeding section, from the slope of In ([Br-1,. [03]/[03]o[Br-])vs. t plot. The average rate constant of two experiments was 164 M-l s-l, which agreed approximately with the first term, kI0, of eq 11. The spontaneous decomposition of ozone becomes marked at pH 9 and should be considered to be competitive with reaction l. In the present study, only the concentration of ozone was followed spectrophotometrically as a function of time. Therefore, it is very difficult to estimate the concentration of bromide ion straightforwardly from the amount of ozone consumed. Hence, numerical calculations were made to simulate the time variation of ozone concentration assuming that reaction 1 and spontaneous decomposition of ozone occur independently, that the rate constant of reaction 1 at pH 9 is equal to 164 M-l s-l obtained at pH 7, and that the net rate of spontaneous decomposition of ozone follows the first-order law, d[O,]/dt = k ~ g ’ [ O , ] Thus, . eq V and VI -d[03]/dt = kl[Br-][03] + kD9’[03] (VI (13)L.R. (1949).
(VI)
B.Yeatts, Jr., and H. Taube, J.Am. Chem. SOC.,71,4100
Discussion The reaction scheme can be expressed by reactions 1-3 Br-
+ O3
-
BrO-
+ O2
+ H+ HBrO HBrO + H+ + Br- e Br2 + H 2 0 BrO-
(1) (2)
(3)
according to our stoichiometric results, where reaction 1 is the rate-determining step. Williams14reported further oxidation of BrO- to bromate when ozone was bubbled through a bromide solution. Under the present conditions, however, the ozone concentration is too low compared to the bromide concentration to initiate further oxidations. Taube2 considered that chain reactions initiated by reaction 4 cannot be neglected in the reaction of ozone with hydrogen peroxide catalyzed by bromide ion. Br2
+ O3+ H 2 0
or Br2
+ O3 + H20
-
-
H02
OH
+ Br + BrOH + O2
+ Br + Br- + H+ + O2
(4a) (4b)
If reaction 4 takes place appreciably, either OH or Br radical produced will propagate chain reactions which consume ozone independently of reaction 1without extra consumption of bromide thus making the [Br2IE/A[O3]ratio less than unity. However, this is not the case with the present reaction system, as shown in Table I. Therefore, reaction 4, if it occurs, is too slow compared with reaction 1 to be of any great importance in the acidic solution. As was described in the preceeding section, eq I1 for the rate constant of reaction 1was obtained on the assumption that bromine is the predominant product in the reaction of ozone with bromide ion at pH below 3. This assumption can be justified as follows. Since the equilibrium constant of reaction 3, K3, is calculated to be 9.2 X M2 at 25 O C and at an ionic strength, F, of 0.10 M from the data reported by Liebhafsky,’*and bromide ion concentrations used in the rate measurements are 8 X 10-5-104, the ratio of [HBrO]/[Br,] is calculated to be 0.01 at pH 2 and 0.10 at pH 3. If the ratio of [HBrO]/[Br2]is denoted by a,the (14)P.M.Williams, R. J. Baldwin, and K. J. Robertson, Water Res., 12, 385 (1978). (15)J. A. Zeevalkink, D. C. Visser, P. Arnoldy, and C. Boelhouwer, Water Res., 14, 1375 (1980). (16)L. M.Dorfman and G. E. Adams, “Reactivity of the Hydroxyl Radical in Aqueous Solutions”, Natl. Stand. Ref. Data Ser. (US.,Natl. Bur. Stand.), No. 46 (1973). (17)R. F. Hampson, Jr., and D. Garvin, “Chemical Kinetics and Photochemical Data for Modelling Atmospheric Chemistry”, NBS Tech. Note (U.S.), No. 866 (1975). (18)H. A. Liebhafsky, J. Am. Chem. SOC.,61, 3513 (1939).
Oxidation of Aqueous Bromide Ion by Ozone
TABLE 11: Reaction Rate for BrNeutral Solutiona 105[0,],, M 2.39, 2.02,
105[Br-],, M 2.65, 3.38,
+
The Journal of Physical Chemistty, Vol. 85,No. 16, 1981 2387
0, in k l , b M-' s-' 154., 173., av 163.,
a Temperature, 20 'C; initial pH, 7.0; and ionic strength, Determined from the data obtained in the 0.10 M. early stages of the reaction, where ozone consumption was less than 20% of its initial concentration and where pH change was n o t appreciable yet.
difference between the true value of [Br-] and the one calculated on the above assumption is expressed by eq VII. [Br-lt,,, - [Br-lcalcd = ( a / ( l + a)I([O3Io- [Od) (VII) From eq I and VII, the magnitude of errors involved in the ordinate values of data points in Figure 2 is represented by eq VIII. By assuming a to be 0.10 and substituting
[Br-lcalcd + (+--)([O~IO - [OJ) [Br-l0)-' In
(VIII) [Br-lcalcd into eq VI11 the extreme values of concentrations used in the experiment at pH 3, i.e., [Br-lo = M, [Br-] = 8 X M, and [O,] = M, we have M, [O3lO= 2 X determined Ay to be -1.8 X lo2. The experimental point corresponding to these concentrations is given in Figure 2 as a half-black circle at the right-hand end whose ordinate value is 6.6 X lo3 M-'. Therefore, an error of 10% in the calculated value of bromide concentration would lower the experimental point by -3%, which is within the experimental error. However, with a few exceptional runs, in which [Br-lo was in only slight excess of what is stoichiometrically required or in which the experiments were carried out above pH 3, corrections must be made according to eq VIII. The former exceptions are shown after due corrections by open circles in Figure 2 and the latter by the extreme left point in Figure 3. According to eq II-IV, the rate constant k1 at 0 "C and pH 3 is estimated to be 57.49 M-' s", which is in good agreement with kl = 26.67 M-' s-' reported by Taube,2 considering the uncertainty of 10% xn activation-energy values. Taube reported that hydrogen ion has no effect on the rate of reaction 1 under his experimental conditions, where [H+] was varied over a very limited range, i.e., 1.16 X 10-3-2.6 X lo-, M. His conclusion seems reasonable since the value of kl changes at most 6% over this range of [H+], as shown in Figure 3. However, to evaluate the rate constant for the reaction H202+ HBrO --, Br- + H+ + O2 + H20, he assumed k1 to be independent of [H+] even at [H+]> lo-' M, which would have resulted in a considerable error in the value of its rate constant. We now consider the mechanism for the hydrogen-ionindependent path of reaction 1. Apparently, neither Br nor 0 radicals could be the products of the rate-determining step, since they have large enthalpies of formation which should result in an enthalpy change of reaction, AH, much larger than the observed value of the activation energy. Thus, the reaction product of the rate-determining step must be BrO- and Oz, The ground states of bromide, ozone, and BrO- are X'S, XIAl, and XIZ+, respectively. Therefore, the oxygen produced in reaction 1 may not be in the triplet ground state (X3Z -). Otherwise, reaction 1violates the selection rule and skould be very slow, since
-
the transmission coefficient of a reaction involving a multiplicity change is usually less than lo4. The value of the frequency factor obtained in this study (4.9 X lo9 M-' s-') is of an ordinary order of magnitude. The oxygen, therefore, may well be in a singlet excited state (A'A, or BIZ +), T i e enthalpy changes, AH, for the reactions forming oxygen in A'A, and B'Z,+ were estimated from the thermodynamic datalg to be ca. -4 and ca. 9 kcal mol-', respectively, both of which are not inconsistent with the activation-energy values obtained experimentally (ca. 10 kcal mol-'). The entropies of oxygen in the A'A, and BIZ,+ levels may be considered nearly equal to that of oxygen in the ground state, since the rotational constants for these excited states are hardly different from that of the ground state and the influence of vibrational and electronic constants on the value of the entropy is quite small compared to that of rotational constants. Thus, the entropy changes, AS,of these reactions are estimated from thermodynamic datalgto be ---18 cal deg-' mol-', which is comparable to the entropy of activation for the hydrogen-ion-independent path, -16 cal deg-' mol-', obtained in this study. Therefore, the structure of the activated complex of reaction 1 is considered to be close to that of one of the products, i.e., BrO-. In the reactions of ozone with bromide ion and with chloride ion, hydrogen ion enhances the frequency factor, as described in the preceding section. It may be of interest to note that the rate laws for the reactions of hydrogen peroxide with bromide ion and with chloride ion have a similar form as that for these reactions.20 In those cases, however, hydrogen ion does not seem to enhance the frequency factor but to lower the activation energy by 3-4 kcal/mol. In an alkaline solution, a chain mechanism involving OH and HOz radicals has been suggested for the spontaneous decomposition of ozone: 0, + OH-
--
HOZ+ Oz(5) 0, + HOz OH + 2 0 2 (6) 0, + OH HOz + O2 (7) OH + HO2 H2O + 0 2 (8) The first-order equation is expressed by eq IX.I5 Since d[ O,] /dt = -2k5[0,] [ OH-] (IX) the rate of reaction between OH (or HOJ and bromide ion is greater than that between OH (or H02) and ozone,16 bromide ion acts as an inhibitor for ozone decomposition. In other words, bromide ion reacts with OH to produce another radical, which reacts with still another radical or ion to terminate the chain reaction. For example: OH + BrBr + OH(9) 4
Br
+ HOz
- + Br-
O2 + H+
(10)
If BrO- should react with OH to produce BrO radical, BrO radical may react with OH as in the gas phase, playing the same role as Br. On the basis of a simple mechanism involving only reactions 5-10, it follows that, as the bromide concentration is increased, the spontaneous decomposition rate of ozone (19)J. C. Bailar, Jr., H. J. EmelBus, Sir Ronald Nyholm, and A. F. Trotman-Dickenson, "Comprehensive Inorganic Chemistry", Vol. 2, Pergamon Press, Elsford, NY, 1973. (20) A.Mohammad and H. A. Liebhafsky, J. Am. Chem. SOC.,56,1680 (1934).
2388
J. Phys. Chem. 1081, 85, 2388-2392
should decrease and approach the rate determined by reaction 5. However, the results obtained in this study show that the spontaneous decomposition rate of ozone in the presence of bromide is nearly equal to that in the absence of bromide as described in the preceding section. It may, therefore, be concluded that Br (or BrO) radical reacts with ozone and generates a different chain path which should be analogous to reactions 6-8. Since it is well-known that atomic bromine reacts with ozone in the gas phase ( k =
7.23 X lo8 M-' s-' at 27 "C)," it is not unreasonable to consider that same reactions take place in solution to accelerate the spontaneous decomposition of ozone. As a result of this study, the reaction of bromide ion with ozone was found to be fast enough for practical application to a deodorization process for seawage plants. A pilot scale test is in progress in our laboratory.
Acknowledgment. We thank Dr. K. Azuma for useful suggestions.
Kinetics of the Reactions of Hydrogen, Nitrogen, and Hydrogen/Nltrogen Mixtures with Molten Lithlum Michael P. Gardner* and M. M. Nlshlna Chemistiy Research and Applications Department, TRW Defense and Space Systems Group, Redondo Beach, California 90278 (Received: January 19, 198 1; In Flnal Form: March 30, I98 1)
Rate measurements were conducted in constant-volume apparatus under isothermal conditions (400-500 "C) to determine the reaction rate for hydrogen, nitrogen, and hydrogen/nitrogen mixtures at the surface of molten lithium. All reactions were found to be pseudo first order in reactant gas pressure. The apparent activation energies for the reactions were EA = 45.4 kJ mol-' for the hydrogen reaction, while the nitrogen reaction gave EA = 62.2 k J mol-'. Reaction of lithium vapor was found not to affect the observed kinetic results; however, reaction product solubility could possibly influence the reaction rate determinations during the early stages of product layer formation. Examination of the product of nitrogen/hydrogen reactions with lithium by X-ray diffraction demonstrated probable product formation of Li3N and LiH and absence of formation of H/N mixed products.
Introduction Interest in the chemical and physical properties of molten lithium has grown considerably over the past decade for several reasons. Its possible utilization as a combined tritium breeder and possible heat transfer media is responsible for current interest. In addition, potential gas-lithium metal reactions have been proposed for heat sources in dynamic power systems. Information is currently needed on the preparation and purification of this metal, including the prevention of contamination. Our interest in gas-lithium metal chemistry arises from the potential utilization of molten lithium as a getter for spent reactants from a bench-scale HF laser. Past studies have attempted to measure the surface reactivity of molten lithium with hydrogen's2 and with n i t r ~ g e n with ~ - ~ a wide variety of results. Considerable (1) R. J. Pulham, P. F. Adams, P. Hubberatey, G. Parry, and A. E. Thunder, Radiat. Eft. Tritium Technol. Fwion React., Proc. Znt. Conf., 4, 144 (1975). (2) G. Parry and R. J. Pulham, J.Chem. SOC.,Dalton Trans., 19,1915 (1975). ( 3 ) V. I. Cheburkov and A. N. Rozanov, Metall. Metalloued. Chist. Met., 1 168 (1968). (4) T. E. Little, "Reactivity of Nitrogen, Oxygen, and Halogenated Gases with Molten Lithium Metal", Technical Memorandum prepared for Naval Ordance Systems Command, TM 72-232, NTIS AD 759378. (5) I. Besson and W. Muller, Ann. Uniu. Sarau., Sci., 4, 322 (1955). (6) E. F. McFarlane and F. C. Tompkins, Trans. Faraday SOC.,58 997 (1962). (7) P. A. Longton, U.K.A.E.A. Report, IGR-TN/C276, 1955. 0022-3654/81/2085-2388$01.25IO
attention has been given to separating reaction kinetics occurring at a fresh metal surface from those dominated by the formation of a surface product film. Addison and Davies8compared stirred and unstirred lithium pools for the nitrogen reaction, while Little4 utilized rapid kinetic measurement techniques to assess nitrogen-lithium surface reactivity. No literature was found containing studies of mixtures of hydrogen/nitrogen reactant. In this study rapid pressure decay techniques were employed to study the reactivity of molten lithium with H2, N2, and H2/Nz mixtures to assess the chemical kinetics of these reactants prior to and during the formation of surface product layers. Experimental Section The lithium employed in this study was purchsed from Lithium Corp. of America and was 99.9+% pure. Hydrogen and nitrogen gases were obtained from Union Carbide Corp., Linde Division, and were 99.99% pure. Hydrogen/nitrogen mixtures were prepared in the laboratory. Experiments were conducted in a stainless-steel, constant-volume reactor. The reactor is constructed of 304 stainless steel and is fitted with four access ports, a small thermocouple inlet port, and a large flange sealed loading port. The reactor volume is 1044 cm3, excluding the pressurization plenum. Of the four access ports, two are (8) C. C. Addison and B. M Davies, J. Chem. SOC.A, 1822 (1969). (9) M. S. Chandrasekharaiahand J. L. Margrave,J.Electrochem. SOC., 108, 1008 (1961).
0 1981 American Chemical Society
