The oxidation of iodine to iodate by hydrogen peroxide

Transcription

1 The oxidation of iodine to iodate by hydrogen peroxide Guy Schmitz Faculte des Sciences Applique es, Universite L ibre de Bruxelles, CP165, Av.F.Roosevelt 50, 1050 Bruxelles, Belgium. gschmitz=ulb.ac.be Received 0th July 001, Accepted rd September 001 First published as an Advance Article on the web rd October 001 The kinetics of the oxidation of iodine by hydrogen peroxide, a part of the BrayÈLiebhafsky reaction, has been studied between 98 and K using a spectrophotometric method. The rate of reaction passes through a maximum when the concentration of hydrogen peroxide increases. At low hydrogen peroxide concentration, the order of reaction with respect to [I ] is roughly one, whereas at high hydrogen peroxide concentration the apparent order is higher and there is no simple kinetic law. This reaction can be observed only with sufficient acid and iodate concentrations. However, its rate does not change much with perchloric acid concentration in the range 0.0 to 0. mol dm~ and iodate concentration in the range to 0.05 mol dm~. The maximum Ðrst order rate constant is 8 ] 10~ s~1 at 98 K with an activation energy of 80 kj mol~1. The results are discussed considering a model proposed previously for the BrayÈLiebhafsky reaction. 1 Introduction The oxidation of iodine to iodate by hydrogen peroxide is a part of the BrayÈLiebhafsky1, (BL) reaction, that is the decomposition (D) of hydrogen peroxide catalysed by iodate and iodine in acidic solutions. H O ] H O ] O (D) As already proposed by Bray,1 it is recognised that this catalysis results from two reactions, the reduction (R) of iodate to iodine and the oxidation (O) of iodine to iodate. IO ~]H` ] 5H O ] I ] 5O ] 6H O (R) I ] 5H O ] IO ~]H` ] H O The sum of reactions (R) and (O) gives reaction (D). These reactions are themselves complex and their mechanisms involve several intermediate species, especially iodide and hypoiodous acid. We have shown, that the kinetics of reaction (R) is closely related to the kinetics of the Dushman reaction5 IO ~]5I~]6H`]I ] H O (1) (O) Here we study the kinetics of reaction (O), the second main component of the BL reaction. First, it is necessary to consider the erimental conditions under which this reaction can be studied. The reactions between hydrogen peroxide and iodine depend on the composition of the solutions. The main parameter is the acidity. In slightly acidic solutions, hydrogen peroxide does not oxidise iodine but reduces it: I ] H O ] I~]H` ] O () It can also oxidise iodide: I~]H` ] H O ] I ] H O () When the rates of these two reactions balance each other we get another catalytic pathway to the decomposition (D). Liebhafsky6 has studied these reactions and concluded that hydrogen peroxide does not react directly with iodine but with hypoiodous acid, a product of the hydrolysis (). I ] H O H HIO ] I~]H` () Thus, the reactions () and () should be written in the form HIO ] H O ] I~]H` ] O ] H O (5) I~]H` ] H O ] HIO ] H O (6) Reaction (5) is important also during the BL reaction. On the other hand, the concentration [I~] is then so low that the reaction (6) can be neglected. The oxidation (O) of iodine to iodate cannot take place unless the acidity is sufficiently high, at least about 0.01 M. At lower acidities we get only reactions (5) and (6). In order to observe reaction (O), it is also necessary to have initially some iodate. Its minimal concentration depends on the acidity and the hydrogen peroxide concentration and can be very small. The reaction can even start without added iodate, but, in this case, it is preceded by an induction period7 long enough to permit some iodate formation. A period of slow iodine concentration decrease8 is also sometimes observed before the start of reaction (O). These observations can be understood noting that the iodide concentration [I~] during reaction (O) is many orders of magnitude lower than during the hydrogen peroxide decomposition following the pathway (5) ] (6). Thus, reaction (O) can be observed only if the acid and iodate concentrations are such that the iodide is quickly removed by the Dushman reaction (1). Liebhafsky has studied the kinetics of the reaction (O) between 0 and 50 C.7,9 He has proposed a Ðrst order rate law [d[i ]/dt \ k [I ], where the erimental rate constant k depends on the hydrogen peroxide, acid and iodate con- centrations. This dependence cannot be ressed by a simple law. It is especially surprising that, although hydrogen peroxide is a reactant, when its concentration increases the rate decreases. Our former results for this reaction10 were in agreement with those of Liebhafsky but uncertainties about some e ects of the concentrations, and their lanation, remained. As understanding of the kinetics of reaction (O) is essential for modelling the BL reaction, we have undertaken a new detailed examination. DOI: /b106505j Phys. Chem. Chem. Phys., 001,, 71È76 71 This journal is ( The Owner Societies 001

2 . Experimental The iodine concentration was followed by spectrophotometry at 60 nm using either a conventional thermostatic quartz cell or a stopped-ñow accessory RX.1000 from Applied Photophysics. The concentration [I~] is so low during reaction (O) that there is no contribution from I ~. The temperature was controlled to ^0.1 C. One drive syringe contained an acidic solution of iodine and iodate and the other a solution of H O with the same acid concentration. Analytical grade KIO, HClO, KI and H O were used without further puri- Ðcation. The H O concentration was controlled by per- manganate titration and the HClO concentration by carbonate titration. The iodine solutions were fresh prepared before each series of eriments. Two di erent methods of preparation were used. Following the Ðrst, solid iodine was sublimed in a nitrogen Ñow and dissolved in deionized water. The concentration of the resulting solution was measured by titration with standard Na S O and controlled spectro- photometrically before each run. Following the second, the iodine was generated by adding KI to an acidic solution of iodate and leaving reaction (1) to take place until completion. Its concentration, deduced from the initial concentration of KI, was controlled spectrophotometrically, as well as the absence of I ~. Both methods gave the same results.. Results and discussion.1. Analysis of the data Fig. 1 shows an example of the e ect of the acidity on the curves ln[i ] vs. time. At the highest acidity we get nearly straight lines consistent with the Ðrst order rate law proposed by Liebhafsky. However, at lower acidities this is no longer true. The value of the apparent order of reaction depends on the erimental conditions. As there is no simple rate law, we have to analyse the data using a di erent method. We deðne an entirely erimental parameter k, a function of all the concentrations, including [I ], by k \ ([d[i ]/dt)/[i ] \ ([da/dt)/a where A is the absorbance at 60 nm. da/dt is calculated by numerical di erentiation of the erimental curves. In order to reduce the well-known e ect of numerical di erentiation on the erimental errors, we have Ðrst to smooth the erimental curves. A convenient method is to take Ðve consecutive equally spaced points and to compute the corresponding best parabola in the least-squares sense. The derivative at the central point A is then given by11 i ([da/dt) \ ([0.A [ 0.1A ] 0.1A ] 0.A )/*t i i~ i~1 i`1 i` where *t is the time interval between two points. The accepted values of k are the mean of at least three measure- ments... Some aspects of the BL reaction The e ects of the di erent concentrations on k are confu- sing and a clear presentation of our results needs to consider Ðrstly some aspects of the BL reaction. Fig. shows examples of the time evolution of iodine in an acidic solution of hydrogen peroxide and iodate. When the initial concentration of iodine is zero or low, we get reaction (R) and the iodine concentration increases until it reaches a constant value. When it is high, we get reaction (O) and the iodine concentration decreases towards the same value. For this value the rates of reactions (R) and (O) are equal and the global reaction is (D). Then we have smooth decomposition of hydrogen peroxide. We call this state the catalytic steady state. The corresponding iodine concentration, denoted by [I ], is a function of the ss temperature, the acidity and the concentrations of iodate and hydrogen peroxide. The catalytic steady state can be stable or unstable. When it is stable the system moves towards it. When it is unstable, the rates of the reactions (O) and (R) cannot become equal and we get a succession of periods R and O. During the periods R the rate of the reaction (R) is larger than the rate of reaction (O) and the concentration of iodine increases. During the periods O the rate of the reaction (O) is larger than the rate of reaction (R) and the concentration of iodine decreases. We get oscillatory decomposition of hydrogen peroxide. The system follows a limit cycle around the unstable catalytic steady state. It is useful to consider the sketch of these behaviours in a phase plane in terms of the concentrations of two intermediate species that can be followed erimentally, the iodine and iodide concentrations. Fig. shows the shapes of the trajectories of the system estimated from former measurements of these concentrations as functions of time.,1,1 The letters R and O indicate the predominant reaction during the evolution along the branches of the trajectories. When the catalytic steady state is stable the system follows the branch R or O until the catalytic steady state is reached. This is the case in the eriments presented in Fig.. If the initial iodine concentration is less than [I ] we get the reaction (R), if it is ss higher than [I ] we get the reaction (O). When the catalytic ss steady state is unstable (see examples in ref., 1 and 1), beside the R and O branches we have a third unstable branch indicated by a broken line in Fig.. Starting with no iodine, the system follows the branch R until point T is reached. 1 Then the iodide concentration drops suddenly, the system jumps onto the branch O and the iodine concentration decreases. At point T the iodide concentration rises suddenly Fig. 1 ln[i ] vs. time at 9 C; [H O ] \ 0.0 M; [KIO ] \ M; [HClO ] \ 0.0 M ()), 0.10 M (L), 0.0 M (]), 0.0 M ( ); Initial slope for [HClO ] \ 0.0 M (ÈÈ). Fig. [I ] vs. time at 5 C; [H O ] \ 0.0 M; [KIO ] \ 0.10 M; [HClO ] \ 0.08 M; [I ] \ 0 (]), 1 ] 10~ M (K), ] 10~ M (L). o 7 Phys. Chem. Chem. Phys., 001,, 71È76

3 the following semi-quantitative description is sufficient. [I ] ss increases when the temperature increases, it is more or less proportional to [H O ]x/[hclo ]y, where x is a little smaller than one and y is about two and decreases slightly when [KIO ] increases. As oxygen is a product of the reaction, it is not possible to control its concentration and it is difficult to quantify its e ect. The most we can say is that oxygen stabilises the catalytic steady state and decreases [I ] at low ss Fig. Shape of the trajectories in the phase plane [I~] vs. [I ] when the catalytic steady state is stable (a) or unstable (b). and the system jumps onto branch R. This completes the cycle. If the initial concentration of iodine is larger than [I ] at point T, the evolution is the same but starts with a period 1 O.1 In a closed reactor, we do not have a true limit cycle because the hydrogen peroxide concentration decreases during the evolution. To give a more exact representation of the evolution we should consider a space with three dimensions, [I ], [I~] and [H O ]. Such a representation has been given elsewhere.1.. The catalytic steady state and the rate of reaction (O) We have measured the rate of decrease of the iodine concentration along the branch O of the trajectories in a closed reactor for di erent concentrations of hydrogen peroxide, iodate and perchloric acid. However, the above discussion shows that it is important to make a distinction between [d[i ]/dt during the periods O and the rate of reaction (O). There is always a contribution from reaction (R), that becomes increasingly important when we approach the catalytic steady state. In particular, in this state d[i ]/dt \ 0 but the rate of reaction (O) is not zero. It is equal to the rate of reaction (R). When the catalytic steady state is unstable, we can also note that the rate of reaction (O) at point T is not equal to zero. Many complications of the kinetics during the periods O are related to the approach of the steady state or the transition point T. We have found that the value of k increases when the ratio [I ]/[I ] or [I ]/[I ] increases and reaches a constant ss T value only when [I ] is much larger than [I ] or [I ]. ss T Thus the reaction order in [I ] is one when, and only when, we are sufficiently far from the catalytic steady state. This gives a qualitative lanation of the results presented in Fig. 1. We have shown that [I ] decreases when the acidity ss increases. For [HClO ] \ 0.0 M, [I ] is near the initial ss concentration [I ] and the slope of ln [I ] vs. time decreases o quickly with [I ]. When [HClO ] \ 0.0 M, [I ] is much ss lower than [I ] but the slope still decreases with [I ]. It is o only for [HClO ] \ 0.0 M that ln [I ] vs. time is nearly a straight line. [I ] and [I ] depend on the temperature, the acidity and ss T the concentrations of iodate and hydrogen peroxide. They can also depend on the oxygen concentration.,1 We do not have yet a precise description of this dependence but for this work acidity. When the catalytic steady state is unstable, [I ] T seems to change more or less as [I ]. In the following dis- ss cussion [I ] will represent [I ] when the catalytic steady min ss state is stable and [I ] when it is unstable. The reaction (O) T can be studied only if the initial concentration of iodine [I ] o is larger than [I ]. Then, the values of k are independent min of [I ] but near [I ] the reproductibility becomes poor. As o min there is an upper limit of [I ] given by its solubility and the dilution factor when mixing the reagents, reaction (O) can be observed only if [I ] is lower than this limit. min Why can we observe reaction (O) only if the acidity is sufficiently high and if at least a small concentration of iodate is initially present? The usual answer is that iodide must be removed by reaction (1). Now, we can give another lanation. For a given initial concentration of iodine [I ], we can o observe reaction (O) only if [I ] [ [I ]. If [HClO ] and o min [KIO ] are too small, [I ] is too large and we can get only min reaction (R). With the model discussed at the end of this paper it can be shown that these two lanations are related... The e ects of hydrogen peroxide and acidity For each eriment, we have calculated k vs. [I ] by numerical di erentiation of the curves [I ] vs. time. Fig. shows some results for di erent hydrogen peroxide concentrations. k decreases only slightly with [I ] when [H O ] is low but decreases quickly when [H O ] \ 0.0. The e ect of [H O ] on k is larger when [I ] is smaller. These obser- vations can be lained by the e ects discussed above. The rate [d[i ]/dt is given by the di erence between the rates of the reactions (O) and (R) with an increasing contribution from reaction (R) when [I ] approaches [I ]. Under the condi- min tions of Fig., [I ] is much lower than [I ] when min [H O ] \ 0.01 M. Then, the contribution of reaction (R) remains small. This is no longer true when [H O ] \ 0.0 M because [I ] increases with [H O ]. Then the system min moves closer to the catalytic steady state or the transition point T, and the increasing importance of reaction (R) decreases k. Fig. 5 shows the e ect of hydrogen peroxide on k for a given value of [I ] and di erent acidities. Our results at 5 C are similar. Although hydrogen peroxide is a reactant, the rate can decrease when [H O ] increases, a fact already shown Fig. Variation of k (s~1) with [I ] at 5 C. [KIO ] \ M, [HClO ] \ 0.10 M; [H O ] \ 0.01 M (]), 0.05 M (L), 0.10 M ( ), 0.0 M (K). Phys. Chem. Chem. Phys., 001,, 71È76 7

4 Fig. 5 E ect of [H O ] on k (s~1) at9 C. [KIO ] \ M; [I ] \ ] 10~ M and [HClO ] \ 0.0 M (]), 0.10 M (L), 0.0 M ( ), 0.0 M ()); [I ] \ ] 10~ M and [HClO ] \ 0.0 M (]); [I ] \ ] 10~ or ] 10~ M gives the same values of k when [HClO ] \ 0.0 M. earlier.7,9,10 Of course, the rate must be zero if [H O ] \ 0 and the real situation is that k passes through a maximum when [H O ] increases. If [HClO ] \ 0.0 M, the maximum of k can be seen when [H O ] is between 0.0 and 0.05 M. At lower acidities the maximum is obtained for lower [H O ] values but cannot be located exactly because we cannot measure k for very low [H O ]. The stoichiometry of reac- tion (O) shows that, in order to have an excess of [H O ], we must take [H O ] [ 5[I ]. o o Fig. and, and the above discussion about the e ect of the approach to the catalytic steady state or the transition point give a qualitative lanation of these results. The e ect of [H O ] depends on the acidity because [I ] increases min when the acidity decreases. When [HClO ] \ 0.0 M, [I ] min is much lower than [I ]. The rate of reaction (O) increases with [H O ] and seems to reach a limiting maximum value. The slow decrease observed at high [H O ] can be lained by a small contribution from reaction (R). This e ect is more pronounced for [HClO ] \ 0.0 and 0.10 M. For [HClO ] \ 0.0 M, [I ] is still larger and k decreases min quickly when [H O ] increases. Near the catalytic steady state, the reproducibility becomes poor and the spectrophotometric measurements can be disturbed by oxygen bubbles. Then, the main stoichiometry is no longer (O) but is (D). The e ect of the acidity is di erent at low or high hydrogen peroxide concentrations. In Fig. 5, we see that at low [H O ] k seems to be independent of [HClO ], except for a small diminution at the highest acidity, whereas at high [H O ] k increases with [HClO ]. At low [H O ] we are far from the catalytic steady state or the transition point and [d[i ]/dt is near to the rate of the reaction (O), which seems to be independent of the acidity. At high [H O ], [I ] is higher and min approaches [I ] when [HClO ] decreases. When the value of the ratio [I ]/[I ] decreases to nearly one, k decreases min quickly..5. Other e ects Here we discuss brieñy the e ects of iodate, oxygen and temperature. Although iodate is a product of the reaction, k increases slightly when [KIO ] increases. According to the results of Fig. 6 it tends to a maximum value. If we assume that the rate of the isolated reaction (O) is independent of [KIO ] this observation can be lained, remembering that [I ] decreases when [KIO ] increases. The e ect of reac- min tion (R) becomes less important at high [KIO ]. However, we suspect that things are more complicated and we need to know more precisely the e ect of [KIO ] on [I ]. Some min complications come also from the relative weakness of iodic Fig. 6 E ect of [HClO ] and [KIO ] on k (s~1) at 5 C. [I ] \ ] 10~ M; [H O ] \ 0.01 M and [HClO ] \ 0.0 M (]), 0.10 M (L), 0.0 M ( ); [H O ] \ 0.0 M and [HClO ] \ 0.10 M (K). acid and the di erences between [HClO ] and [H`] and between [KIO ] and [IO ~].15 Series of eriments were performed to check the e ect of oxygen. In one group of eriments, the solutions were saturated with oxygen. In the other group, solutions with the same concentrations were prepared under nitrogen. In most cases, they gave identical results. However, oxygen has an e ect at low acidities and high [H O ]. Then it increases the rate of reaction and, when the initial concentration of iodine is close to [I ], solutions saturated with oxygen give reaction (O) min while solutions prepared under nitrogen give reaction (R). As k is a complicated function of the concentrations, there is no real activation energy related to it. However we can estimate a maximum value of k for [I ] \ ] 10~ M. This gives (k ) \ 8 ] 10~ s~1 at 5 C and 0.0 s~1 at max 9 C. Liebhafsky found 5 ] 10~ s~1 at 0 C7 and 0.1 s~1 at 50 C.9 A plot of ln(k ) vs. 1/T gives a straight line with a max slope of 80 kj mol~1. This value is in agreement with that obtained in the same range of acidities by a di erent method (77 kj mol~1).16. A model for the BL reaction.1. The simpliðed model The model we have proposed,17 for the BL reaction can be used for simulations with more or less detail. The main forms are the general model, the simpliðed model and the Braylator. The general model includes three groups of reactions: the reactions of iodine and its compounds constituting the mechanism of the Dushman reaction, the reactions of hydrogen peroxide with these iodine compounds and the reactions of oxygen. It includes many reactions with unknown rate constants and we think it is useless to try to adjust them to a given set of erimental results. It is well known that a good Ðt is not a good proof when the number of adjustable parameters is large. This is why we have proposed a simpliðed model, including only the reactions we consider essential to model the BL oscillations. It gives a semi-quantitative lanation of most erimental facts. Further work should improve the model by adding reactions only when their importance is demonstrated. On the basis of a stoichiometric network analysis,18,19 we have also constructed the Braylator. This is the most simple form of the model keeping its essential features. Although oversimpliðed, it is of theoretical interest because it shows clearly the origin of the BL oscillations.0 The simpliðed model consists of reactions (M1)È(M6). It is derived from the general model keeping only two reactions of hydrogen peroxide, one reduction (M5) and one oxidation (M6) and contracting the Ðrst steps of the Dushman reaction into one pseudo-elementary step (M1). Actually, H I O, or its anhydride I O, is an intermediate compound in this reac- tion.,1 In order to lain qualitatively the e ect of oxygen, 7 Phys. Chem. Chem. Phys., 001,, 71È76

5 we have added to the simpliðed model the global reaction (M7),1 with an empirical rate law r \ k (O )(I~). k is an apparent rate constant, an unknown function of (H`), (IO ~) and (H O ). IO ~]I~]H` H HIO ] HIO (M1) HIO ] I~]H`]I O ] H O (M) I O ] H O H HIO (M) HIO ] I~]H`HI ] H O (M) HIO ] H O ] I~]H` ] O ] H O (M5) I O ] H O ] HIO ] HIO (M6) I~]H` ]1O ] HIO (M7) Numerical simulations have been published showing that the model can lain the main features of the BL oscillations,h5 and of the reaction (O).10 This reaction is the result of the combination (M [ 1) ] (M) ] (M [ ) ] (M [ ) ] 5(M6). Hydrogen peroxide acts as an oxidant in reaction (M6). As it oxidises I O produced by reaction (M [ ), the rate of (M6) is proportional to [HIO]. On the other hand, hydrogen peroxide acts as a reductant in reaction (M5) with a rate proportional to [HIO]. The key of the model is the competition between the oxidation and the reduction of HIO with di erent orders of reaction.1,0 The global result is an oxidation only if [HIO] is sufficiently large. With the quasi-equilibrium (M) we see that it is equivalent to say that [I~] must be sufficiently low, as observed. The direct oxidation of I~ by hydrogen peroxide is negligible in the conditions of the BL reaction.6,7 I~ can be removed by reaction (M1) at the start of reaction (O) but, during this reaction, we have another pathway for the iodide removal, the sum of reactions (M) and (M6) giving H O ] I~]H`]HIO ] H O... The catalytic steady state At the catalytic steady state, the only global reaction is (D) and the rates of production and consumption of all the iodine compounds are equal. This gives the following relations between the rates r of the steps (Mi). i r \ r \ r \ 0 r \ r r \ r ] r Introducing the ressions for these rates, we have obtained the ression C [I ] \ k k k 5 DC k k [H O ] D1@ 1 ~1 5 ss k k k k k [IO ~][H`] a5 ~ ~ where a is a parameter function of the oxygen concentration. 7 The simpliðed model lains the increase of [I ] when ss [H O ] increases and when [IO ~] or [H`] decrease... The limiting rate law The simplest lanation of the rate law [d[i ]/dt \ k [I ], where k is a complex function of the concentrations but reaching a maximum value, is based on the rate of reaction (M [ ). With [d[i ]/dt \ k [I ] ~ [ k [HIO][I~][H`], this maximum would be reached when k [I ] A k [HIO][I~][H`] with k \ k. However the ~ ~ value of k has been measured directly8,9 and is larger ~ than k. Thus this lanation cannot be correct. The sim- pliðed model gives another lanation and an analytical analysis of it is useful. Let us consider the model under conditions very far from the catalytic steady state when [I~] is very low and [HIO] very large. Then, we can neglect the component from left to right of reaction (M1), neglect the rate of reaction (M5) compared to (M6) and write the model as follows. (7) I ] H O H HIO ] I~]H` (M [ ) HIO H I O ] H O (M[ ) I O ] H O ] HIO ] HIO (M6) HIO ] I~]H`]I O ] H O (M) HIO ] HIO ] IO ~]I~]H` (M [ 1) The quasi-steady state approximation for the intermediate species gives the relations 1 k ~ [HIO][1 k [I O] \1 5 k 6 [I O][H O ] \1k [HIO ][I~][H`] \1k [HIO][HIO ] ~1 Reaction (M) is at quasi-equilibrium during our eriments: k [I ] \ k [HIO][I~][H`] ~ From these relations and [d[i ]/dt \1d[IO ~]/dt \1r, ~1 we get [ d[i ] \ k k ~ k k [H O ] ~ 6 dt k k 5k ] k [H O ] [I ] ~1 6 This is exactly the form of the rate law observed very far from the catalytic steady state, Ðrst order in [I ] with a maximum value when [H O ] increases. Closer to the catalytic steady state, the rate is lower because the rate of reaction (M5) and the reversibility of reaction (M [ 1) cannot be neglected. Then, there is no simple rate law. Conclusions The kinetics of the oxidation of iodine by hydrogen peroxide is complex and is the result of two kinds of e ects: the kinetics of reaction (O) very far from the catalytic steady state and the approach to this state. Very far from the catalytic steady state, when hydrogen peroxide acts only as an oxidant, the kinetics is relatively simple. The reaction is order one with respect to iodine and its rate increases with the hydrogen peroxide concentration up to a limiting value. The acidity and the iodate concentration have no, or a very small, e ect. The complexity comes out in conditions where hydrogen peroxide acts also as a reductant, especially when we cannot neglect the well known reaction HIO ] H O ] I~]H` ] O ] H O. In these conditions, there is no simple relation between the rate and the ratio [I ]/[I ] but all the factors that decrease this min ratio, decrease the rate. The erimental parameter k decreases when [I ] decreases and when [H O ] increases. It increases slightly with the iodate concentration and increases or decreases with the acidity, depending on the other concentrations. Oxygen has no e ect on the rate of reaction except near the catalytic steady state, where it promotes reaction (O) and increases its rate. We have discussed these results in terms of our model for the BrayÈLiebhafsky reaction. It lains the results very far from the catalytic steady state. The simpliðed model (M1)È (M7) gives also a semi-quantative lanation of the other results but appears to be an oversimpliðcation. On the other hand, the general model includes too many reactions with unknown rate constants. Our aim is to Ðnd a model including all the important reactions, but no more, based on di erent kinds of eriments, not only on the shape of the BL oscillations. Before improving our earlier numerical simulations, we also need more independent information about the rate constants. This work gives such information. Acknowledgement The author thanks Professor Ljiljana Kolar-Anic for very useful discussions during the writing of this paper. Phys. Chem. Chem. Phys., 001,, 71È76 75

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