J . Phys. Chem. 1989, 93, 261-264 While for surface I reagent rotation promotes the reaction and the products are formed with high rotational energy, for surfaces I11 and IV there is a decline in reactivity between j = 0 and j = 6, followed by a significant increase for higher values of j , and the products are formed with low rotational energy. The differences between the rotational energies of the products are especially pronounced for low collision energies and low initial rotational states and can be related to the different geometries at the interaction region, resulting from different features of the potential energy surfaces. The examination of QCT and experimental results for several H LH' reactions (C1 HCl, 0 HC1, 0 HBr, and 0 HR, where H R is a hydrocarbon molecule) indicates very interesting correlations between contrasting features of potential energy surfaces and contrasting dynamical results. One type of surface, COLD surfaces (collinearly directing surfaces), are dominated by attractive forces which tend to direct the reagents toward a nearly collinear configuration at the interaction region, and they lead to the formation of rotationally "cold" products. The second type of surface, H R E P surfaces, are dominated by repulsive forces, and they lead to the formation of highly rotationally excited products. The properties characterizing the COLD surfaces seem to be a strong decline of reactivity withj, at least for low values of j ; formation of products with low rotational energy; strong oscillations of partial cross sections as a function of collision energy (for j = 0); a wide range of & leading to reactive
+
+
+
+
+
261
collisions; and a steep increase in barrier height with the deviation from the collinear configuration. Properties characterizing HREP surfaces seem to be promotion of reactivity with reagent rotation; formation of highly rotational excited products; no significant oscillations in partial cross sections as a function of collision energy; a narrow range of Or leading to reactive collisions; and a smaller change in barrier height with the deviation from collinearity than for COLD surfaces. Our studies for C1+ HCl also indicate that dynamical propensity rules are obeyed much better for COLD surfaces than for H R E P surfaces. The differences between the properties of the two types of potential energy surfaces are especially marked for low rotational states of the reagents and low collision energies. Additional extensive 3D quasi-clessical trajectory studies for several H + LH' systems are under way, aiming to give more insight into the dynamics of such reactions and to test the general nature of the correlations suggested above between the nature of potential energy surfaces and dynamical results. Hopefully, the QCT results presented and discussed here will stimulate similar experimental dynamical studies, as well as theoretical studies, of this important class of reactions.
Acknowledgment. This research was supported by the Fund for Basic Research administered by the Israel Academy of Sciences and Humanities. Registry No. HC1,7647-01-0; DC1,7698-05-7; C12,7782-50-5.
Estimate of the Rate Constant of the Reaction between HOI and 4-[ 2-( Methylamino)propyl]phenol in 1 M Sulfuric Acid Margit Vargat and Peter Ruoff* Department of Chemistry, Rogaland University Center, Ullandhaug, N-4004 Stavanger, Norway (Received: March 9, 1988)
We have studied the kinetics of the reaction between HOI and 4- [2-(methylamino)propyl]phenol (HMP) by following the iodine (I2) absorbance in an excess of the organic substrate (0.1 M) in 1 M sulfuric acid at 25 OC. The goal of our study is to determine the rate constant of the removal of HOI by HMP, which is important for work on a related iodine-catalyzed bromate oscillator. The presence of additional iodate keeps the amount of iodide ion, which otherwise would retard the reaction at a low level. The rate constant of the HMP-HOI reaction is estimated to 8 X lo4 M-' s-I.
Introduction Oscillatory chemical reactions have received an increasing attention in the past two decades.' Among the best understood examples are bromate-driven oscillators2 where an organic substrate is brominated and oxidized by acidic b r ~ m a t e . ~The main mechanistic understanding of bromate ascillators is due to the work of Noyes and c o - ~ o r k e r sthe ; ~ mechanism is generally referred to as the Field-Korb-Noyes (FKN)5 or the OrbBn-Korb-Noyes mechanism (OKN)? depending on whether a catalyzed'~~ or u n c a t a l y ~ e d ~system .'~ is considered, respectively. Varga et al." showed recently that, in uncatalyzed bromate oscillators, where alkyl-substituted phenols are used as organic substrates, added iodide ion acts as a catalyst and increases the number of oscillations considerably.12 The presence of an aliphatic side chain appears necessary for induction of an increased number of oscillations by added iodide ion. Further experiments12indicate that intermediates in the iodide oxidation by bromate are probably responsible for the catalytic effect of iodine. It further appears that the important reaction 'Institute of Inorganic and Analytical Chemistry, L. Eotvos University, H-1443 Budapest, Hungary.
is probably between an iodine-oxygen intermediate and the organic substrate. One of the possibilities is the attack of iodine (+I) in the form of HOI or H20-I+on the organic substrate H M P (Figure 1). We have studied the kinetics between HOI and H M P under oscillatory with the aim to provide an estimate of (1) Noyes, R.M. In McGraw-Hill Encyclopedia of Chemistry; Parker, S . P., Ed.; McGraw-Hill: New York, 1982; p 716. (2) Noyes, R.M. J. Am. Chem. SOC.1980, 102,4644. (3) See contributions by: Field, R.J.; Tyson, J. J.; Epstein, I. R.;Orbin, M. In Oscillations and Traveling Waves in Chemical Systems; Field, R. J., Burger, M., Eds.; Wiley: New York, 1985. (4) Noyes, R. M.; Field, R. J.; KBros, E. J . Am. Chem. SOC.1972, 94, 1349. (5) Field, R. J.; Koros, E.; Noyes, R. M. J. Am. Chem. SOC.1972, 94, 8649. (6) Orbin, M.; Koros, E.; Noyes, R. M. J . Phys. Chem. 1979,83, 3056. (7) Belousov, B. P. S b . Ref. Radiat. Med. 1959, 145. (8) Zhabotinsky, A. M. Dokl. Akad. Nauk SSSR 1964, 157, 392. (9) Koros, E.; Orbin, M. Nature (London) 1978, 273, 371. (10) Orbin, M.; KBros, E. J. Phys. Chem. 1978, 82, 1672. (11) Varga, M.; PaulB, T.; KBros, E. React. Kinet. Carol. Lett. 1984,26, 363. (12) Ruoff, P.; Varga, M.; Koros, E. J . Phys. Chem. 1987, 5332.
0022-3654/89/2093-0261$01.50/00 1989 American Chemical Society
262
Varga and Ruoff
The Journal of Physical Chemistry, Vol. 93, No. 1, 1989
OH
6.97.0/OI*
6.8.
CHZ
-iog[I2 lo
I
2.7
CH-NH-CH3
3.1
Figure 3. Logarithm of initial rate of iodine absorbance decrease against logarithm of measured initial iodine concentration. The straight line has been obtained by linear regression with a slope of 0.5.
I
1.01
CH3 Figure 1. (HMP).
2.9
The organic substrate 4-[2-(methylamino)propyl]phenol
1.20 1
3.0
0.0
6.0
9.0
12.0
150
TIME (MlN.)
m
Figure 4. Iodine absorbance (460 nm) versus time. Initial concentrations: (A) [IzlO= 2.0 X M; (B) [I2lO= 1.5 X lo-’ M; (C)[I210= 1.0 x M; (D) [1210 = 5.0 X lo4 M. In all runs: [IO3-lO= 5.0 X lo4 M, [HMPIo = 0.1 M, [H2S04]0 = 1.0 M.
m 30
00
60
90
120
150
TIME ( M I N I
60,1-j
Figure 2. Iodine absorbance (460 nm) versus time. Initial concentraM; (B) [I2lO= 1.2 X lO-’M; (C) [IzlO= tions: (A) [IJO = 1.5 X 9.0 X lo-“ M. In all runs: [HMPIo = 0.1 M, [H2S04]0= 1.0 M.
the rate constant of the HOI removal by H M P for subsequent model calculations.
Experimental Section All chemicals, except the organic substrate, were of commercial analytical quality. The organic substrate, abbreviated as HMP,”*I2 was the sulfate salt of 4- [2-(methylamino)propyl]phenol(Figure 1) with a purity greater than 99%. The HOI-HMP reaction was studied by following the iodine absorbance at 460 nm with a double beam Perkin-Elmer Lambda 15 spectrophotometer. All experiments were performed at 25 f 0.1 OC in 1.0 M sulfuric acid, which was also the reference solution used in the spectrophotometer. The measured iodine absorption coefficient was found to be 0.294 mg-’ cm2, which is in good agreement with the literature value (0.29 mg-’ cm2).I3
Results Reaction of HMP with Iodine. Iodine was produced quantitatively in the solution by using stoichiometric amounts of iodine and iodate ion, according to reaction Rl.I4 The HOI-HMP 103-
+ 51- + 6H+
-
3H20
+ 312
(R1)
reaction was started by adding an acidic iodate solution into a cuvette, containing a mixture of H M P and iodide ion. After rapid mixing of the solutions, the cuvette was placed into the spectrophotometer and the iodine absorbance was followed. The H M P (1 3 ) Gmeliru Handbuch der Anorganischen Chemie; Achte Auflage, Jod,
System-Nummer 8; Verlag Chemie: Berlin, 1933; p 138. (14) Kolthoff, I. M.; Sandell, E. B. Textbook of Quantitatiue Inorganic
Analysis; Macmillan: New York, 1952; p 556.
iodate / iod ide rat io
>o
, 5’9: -8’5.8 5.72’9
31 33 3 5 -log11~1,
3’7
Figure 5. Logarithm of initial rate of iodine absorbance decrease against logarithm of measured initial iodine concentration and applied initial iodate/iodide concentration ratio.
concentration was 0.1 M, while the initial iodine concentration was varied. The decrease in iodine absorbance was very slow and the concentration vs time plot was practically linear. Figure 2 shows runs with different initial iodine concentrations. The initial rate was determined from the numerical output of the spectrophotometer. Figure 3 shows the log uo as a function of log [I&, where uo is the initial rate and [I2lOis the initial iodine concentration. The reaction order in iodine was determined from the slope of Figure 3 by linear regression. The value obtained was 0.5. Because of the very slow rate of the iodine removal, we did not decrease the organic substrate concentration. Different Iodate/ Iodide Concentration Ratios. In the following experiments we used an initial iodate and H M P concentration of 5 X lo4 M and 0.1 M, respectively. The initial iodide concentration was changed such that the iodate/iodide concentration ratio successively varied from 1:4, 1:3, 1:2 to finally 1:l (Figure 4). Figure 5 shows the initial rate of the iodine absorbance change against the measured initial iodine concentration and the iodate/iodide concentration ratio. We observe that the initial rate of the iodine absorbance is largest for the 1:2 iodate/iodide ratio.
:I-/
Reaction Between HOI and 4-[2-(Methylamino)propyl]phenol
The Journal of Physical Chemistry, Vol. 93, No. 1 , 1989 263 Process R5 has an equilibrium constant K5of 2 X 10l2 M-I, and the forward rate constant is 3.1 X 10l2 M-' s-I.l7 During the rapid formation of iodine, HOI also reacts with the organic substrate, which we schematically describe by eq R6. The HMP
5.85
- log [ HM PI,
5.75
1.4 1.6 1.8 Figure 6. Logarithm of initial rate of iodine absorbance decrease against
1.2
1.0
logarithm of initial HMP concentration. The initial iodate concentration has been 5 X loJ M, and the initial iodide concentration has been 1 X lW3 M. The straight line was obtained by linear regression, and the slope was determined to 0.5.
5.9
+ HOI
-
product
(R6)
equilibrium of reaction R5 is rapidly established, while reaction R6 removes the HOI. This leads to an increase of the iodide ion and to a reduction of the steady-state HOI concentration. In principle, a system consisting of process RS in equilibrium and reaction R6 would be sufficient to determine the rate constant of reaction R6, if the rate of iodine (Iz) disappearance and the HOI and iodide concentrations were known. Practically, however, it is difficult to determine the initial HOI and iodide concentration, because iodide ion builds up, while HOI is being consumed. To overcome this problem in an analogous study of the hydroxyiodination of 2-butenoic acid, Furrowt6in an elegant experiment added an excess amount of iodate to the system, in order to keep the iodide concentration low. Furrow considered reactions R3-R6 and applied the steady-state approximation for HOI, HOIO, and I-, which leads to the following rate expression in our system: d[I,lo/dt = (2/ 5 ) ( 5 k3k6/ K5)'/2[H+]o[ HMP] o'/'[ 103-1o'/'[Iz] 0'1' ( 1)
- log [ lOj1,
5.7 3.2
3.4
3.6
3.8
4.0
4.2
Figure 7. Logarithm of initial rate of iodine absorbance decrease against logarithm of initial iodate concentration. The initial iodine concentration was 5 X loJ M. The straight line was obtained by linear regression, and the slope was determined to 0.5.
Using a constant 1:2 iodate/iodide ratio we have varied the amount of organic substrate and determined the order of the reaction with respect to HMP, which was found to be 0.5. The results are shown in Figure 6. Iodate Simultaneously Present with Iodine. In the last type of experiment we varied the amount of iodate in the presence of a constant iodine concentration. The initial rate of the reaction was determined from the slope of the iodine absorbance vs. time plot for the first 15 s. Figure 7 shows the logarithm of the initial rate as a function of the logarithm of the initial iodate concentration. From the slope we determined the order of the reaction with respect to iodate; the value found was 0.5.
Discussion Figure 5 shows that the initial rate of the iodine removal is largest when the iodate/iodide concentration ratio is 1:2. This concentration ratio corresponds formally to the stoichiometry:
+ 21- + 3H+
-
where K5 is the equilibrium constant of process R5 and k3 and k6 are the rate constants of processes R3 and R6, respectively. In our experiments we get the following empirical rate equation -d [121o/dt = k[HMP]oo" [ 103-1 0°.5 [1210O.~
(2)
where k was found to be (7.2 f 0.7) X M4.5 s-l. The [H+] dependence is included into k. Setting eq 1 equal to eq 2, we get the following expression for k6 k6
= ( 5 /4)(k2K5/ k3)
/ [H+10)
(3)
For 1 M sulfuric acid [H+Io= 1.2 M. The iodine concentration was 5 X lo4 M, the iodate concentration varied between 1 X lo4 and 5 X IO4 M, and the H M P concentration was 0.1 M (Figure 5). From this we obtain an average k6 value of 8 X lo4 M-' s-I. The maximum of the initial rate described in Figure 5 can be explained as follows. Due to the rather high acidity used in our experiments, it appears that process R1 is complete within a few seconds, i.e., before any readings on the spectrophotometer are available. If we calculate the "initial" concentrations according to reaction R1, and let A be the initial iodate concentration (5.0 X lo4 M) and nA be the initial bromide ion concentration ( n varies from 1 to 4), we obtain
(R2)
-d[1210/dt = k[HMP]oo.5[A(1- n/5)]0.5[3nA/5]0,5 (4)
However, we believe that the observed maximum in the initial rate is of kinetic and not of stoichiometric origin, although HOI is the intermediate which reacts with HMP. The sequence of reactions which leads to the formation of HOI is the following
The maximum of the initial rate occurs at n = 2.5, independent of the reaction order, so long as reaction orders with respect to iodine and iodate are the same. The rate is symmetrical in n around 2.5, which implies that the rate should be the same at n = 1 and n = 4 with a minimum in log (rate) at n = 2.5, precisely as Figure 5 shows. Implications for the Iodine-Catalyzed Bromate Oscillator. According to an earlier suggestion by Varga et al." the catalytic effect of added iodide ion is due to a rapid reaction of iodine (+I) (Le., HOI) with the organic substrate, which prevents further oxidation of HOI by bromate to iodate. Additional studies by Ruoff et a1.l' support this assumption, showing that iodine-catalyzed oscillations are only obtained when the in situ generated iodine (+I) intermediate is rapidly captured by the organic substrate. This implies that the rate of the HOI-HMP reaction must be considerably faster than the rate of HOI oxidation by bromate. Citri and Epstein'* have recently proposed that the rate constant value of the oxidation of HOI by bromate (R7) is of the order
IO
