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Oscillatory System I-, H2O2, HClO4: The Modified Form of the Bray-Liebhafsky Reaction Anna Olexova´, Marta Mra´kavova´, Milan Melichercˇ´ık,† and L’udovı´t Treindl* Department of Physical and Theoretical Chemistry, Comenius UniVersity, 842 15 BratislaVa, SloVakia ReceiVed: March 17, 2010; ReVised Manuscript ReceiVed: May 19, 2010
The kinetics of iodide ions oxidation with hydrogen peroxide in solutions of perchloric acid at temperature of 60 °C has been studied in detail. We have found conditions under which this reaction proceeds oscillatory. The Bray-Liebhafsky (BL) oscillatory reaction started by the oxidation of iodide ions with hydrogen peroxide is described for the first time. The described results support our assumption (Olexova´, A.; Mra´kavova´, M.; Melichercˇ´ık, M.; Treindl, L. Collect. Czech. Chem. Commun. 2006, 71, 91-106) that singlet oxygen (1O2) is an important intermediate of the BL oscillatory reaction in the sense of the Noyes-Treindl (N-T) skeleton mechanism (Treindl, L.; Noyes, R.M. J. Phys. Chem. 1993, 97, 11354-11362). HOI + H2O2 f I- + H+ + O2(aq) + H2O
1. Introduction Although the Bray-Liebhafsky (BL) oscillatory reaction has been known more than 80 years, its reaction mechanism has not yet been univocally established. It consists of the two stoichiometric processes
I- + H+ + H2O2 f HOI + H2O I2 S 2 I I + O2 S IOO
-
+
2IO3 + 5H2O2 + 2H f I2 + 5O2(aq) + 6H2O
(I) I2 + 5H2O2 f 2IO3- + 2H+ + 4H2O
(II)
resulting in the brutto process
2H2O2 f 2H2O + O2
(III)
Under defined conditions, process I alternates with process II, and one can observe oscillations of iodine, iodide ions, and of oxygen and periodical decrease in hydrogen peroxide.1-9 In 1993, one of us was honored to collaborate with Prof. R. M. Noyes at the University of Oregon, and on the basis of this collaboration, a new explanation of the oscillations in the BL reaction has been proposed.8 According to the proposed skeleton reaction mechanism, H2O2 is oxidized only with HOI, the real oxidant is O2,, and H2O2 itself can oxidize no other iodinecontaining species than iodide ion, even though H2O2 is the only oxidant in the overall stoichiometry for one of the two main processes of the BL reaction. The Noyes-Treindl (N-T) skeleton mechanism8 consists of ten elementary processes + IO3 + I + 2H f HIO2 + HOI
(1)
HIO2 + I- + H+ f 2HOI
(2)
HOI + I- + H+ S I2 + H2O
(3)
IOO S IO2 2IO2 + H2O f IO3- + H+ + HIO2 O2(aq) S O2(g)
(4) (5) (6) (7) (8) (9) (10)
Field et al.10 showed using a mathematical model based on this mechanism8 that [I2] and [O2] oscillate around an unstable steady state. According to them,10 these oscillations can be quite longlived; they take place around a stable or unstable focus and do not involve a true limit cycle. Derrick et al.11 have shown that large amplitude oscillations arise through a homoclinic bifurcation and vanish through a subcritical Hopf bifurcation. La´nˇova´ et al.12 based on their first mass spectrometric study confirmed the assumptions made in the N-T model that the escaping rate of oxygen from the solution influences the oscillations substantially and that it needs to be taken into account in the modeling of oscillations. Just recently, Ren et al.13 have also performed computational simulation of the BL reaction based on the N-T mechanism consisting of the above 10 reaction steps and simulated the oscillations of the concentrations of I2 and of O2. As a result of our kinetic study of the autocatalytic oxidation of iodine with hydrogen peroxide,14 it has been pointed out that during this reaction, singlet oxygen (1O2) is produced, which may play an important role in the oxidation of iodine, in the sense of the N-T mechanism. The aim of this work was the kinetic study of the iodide oxidation by hydrogen peroxide in dilute solutions of perchloric acid at temperature of 60 °C. We have found the conditions under which this reaction is oscillatory. It is for the first time described that the BL oscillations occur in a system composed of the reactants of process II, being started by oxidation of iodide with hydrogen peroxide. 2. Experimental Section
*To whom the correspondence should be sent. † Present address: Faculty of Natural Sciences, Matej Bel University, 974 01 Banska´ Bystrica, Slovakia.
The reaction courses were followed mostly potentiometrically by recording the time dependence of the I-/I2 redox potential
10.1021/jp1024284 2010 American Chemical Society Published on Web 06/16/2010
Oscillatory System I-, H2O2, HClO4
J. Phys. Chem. A, Vol. 114, No. 26, 2010 7027 TABLE 1 c, M a
[HClO4]o [H2O2]ob [I-]oc [H2O2]od [I-]oe
0.055-0.125-0.14 0.055-0.065-0.25 0.003-0.0250.035 0.04-0.05-0.15 0.015-0.02-0.03
IP, sa
PO, sb
2360-120-145 1800 f 300 250-120-435 1000 f 370 1900-125-250 2250 f 280
AO, mVc 100 f 40 35 f 40 50 f 40
520-460-1650 5700 f 2220 90 f 70 1400-12255100 f 2100 90 f 70 1550
a IP: induction period, corresponding to the adequate concentration. b PO: period of oscillation, decreasing step by step. c AO: amplitude of oscillation, decreasing step by step.
Figure 1. Effect of stirring on the E ) f (t) curves [NaI]0 ) 4 × 10 -3 mol dm-3, [H2O2]0 ) 10-1 mol dm-3, and [HClO4]0 ) 0.12 mol dm-3, 60 °C, with dispersed light, (a) without stirring, (second arrow indicates the start of stirring 50 rpm), (b) with stirring of 50 rpm, and (c) with stirring of 50 rpm (first arrow indicates when stirring was switched out).
(Pt electrode) against the potential-saturated Hg/Hg2SO4/K2SO4 reference electrode. The cylindric vessel (height 5 cm, diameter 2.5 cm) was closed by a rubber plug with a small tube (diameter 2 mm) to reduce the pressure difference, with both electrodes and with a glass mantel connected to ultrathermostat. The reactants were inlet into the reaction vessel in the following subsequence: aqueous solution of HClO4, solution of NaI, and finally solution of H2O2, and in such a way, the reaction was started. The time dependences of the electrode potential were monitored with a digital multimeter Metex M 4650CR connected to a PC. The collected data were worked up using the programs Gnuplot and Excel. When the reaction was followed in light, the lab was illuminated by normal neon lights. The reaction in darkness was followed with a reaction vessel coated by aluminum foil. Spectrophotometric measurements were performed using a Perkin-Elmer spectrophotometer. The measurements were based on the monitoring of the time dependence of the iodine concentration at wavelength of 460 nm. A thermostatted 1 cm quartz cell with circle orifice (r ) 2 mm) was placed in the spectrophotometer. The measurements were performed mostly at temperature of 60 °C. The reaction solution was intensively stirred magnetically with the Teflon-coated stirrer 10 s before starting a potentiometric or spectrophotometric registration. Stock solutions of the components were prepared from commercially available chemicals, hydrogen peroxide (p.a., Merck, Germany), HClO4 (p.a. Merck, Germany), NaI (p.a., Centralchem, Slovakia), and D2O (99.86%) (Institute for Research and Utilization of Radioisotopes, Prague, Czech Republic). All other reagents were of analytical grade. Distilled and deionized water was used in all experiments. The H2O2 concentration was determined by iodometric titration, and the HClO4 concentration was determined by alkalimetric titration. The solution of H2O2 was always freshly prepared. 3. Results and Discussion If the oxidation of iodide ions by hydrogen peroxide in dilute H2O solutions of perchloric acid is followed potentiometrically, then the extent of oxidation depends dramatically on the stirring rate of the solution. Even the stirring of as low of a rate as 50 rpm inhibits the oxidation of iodide to IO3- (Figure 1c). If stirring of solution is switched out, then after some induction
Figure 2. Effect of D2O on the E ) f (t) reaction course. [NaI]0 ) 5 × 10-3 mol dm-3, [HClO4]0 ) 0.12 mol dm-3, 60 °C. (a) In light, [H2O2]0 ) 0.15 mol dm-3, volume ratio of D2O and H2O: RV ) 0 and (b) in light, [H2O2]0 ) 0.15 mol dm-3, RV ) 0.78.
period (IP) the potential of the Pt electrode increases, and on the E ) f (t) curve, first maximum and then the oscillations can be observed (Figure 1a). On the other side, in the presence of D2O, under the same conditions, the oxidation to IO3- does proceed, and the oscillations do occur even if the solution is stirred permanently with frequency of 70 rpm. When the oscillations already occur in the absence of D2O, stirring leads to the increasing amplitude and to the increasing period of oscillations (POs) (Figure 1b). In the presence of D2O, under the same conditions, no IP is needed for the oscillations to occur. The dependence of oscillatory characteristics on initial reactant concentrations of reactants in H2O solutions is shown in Table 1. One can see the influence of initial concentrations of HClO4, H2O2, and NaI on the length of IP, PO, and amplitude of oscillation (AO). The values of PO and AO are averages of the first five oscillations. One can ask a question of how our studied system relates to the BL reaction because no iodate ion is initially present in the system. Our answer is that after the first cycle, our studied system behaves identically with the classical BL system. The D2O effect and the effect of illumination on the E ) f (t) reaction course can be seen in Figure 2. The most pregnant effect of D2O can be observed in illuminated solutions. The described results do support our assumption14,18,19 that during the autocatalytic oxidation of iodine with hydrogen peroxide very reactive singlet oxygen is produced, and the oxidation of iodine can start in the sense of the N-T reaction scheme.8 In our papers,14,18,19 the production of singlet oxygen as well as its role in the oxidation of iodide ions with hydrogen peroxide has been already accounted for adequately. Although the influence of D2O on the BL reaction is quite complex, but also according to Stanisavljev et al.,22 replacement of H2O by D2O progressively intensifies the
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Olexova´ et al.
Figure 3. Effect of light on the reaction course in H2O solution. [HClO4]0 ) 0.12 mol dm-3, [H2O2]0 ) 0.1 mol dm-3, and [NaI]0 ) 5 × 10-3 mol dm-3, 60 °C, without stirring (a) with dispersed light and (b) in darkness.
reactions of oxidation of the iodine species, and the oxidation pathway is more effectively accelerated than the reduction pathway. According to them,22 it is not possible to explain the observed effects solely as a result of changed acidity, occurring from the isotopic effect. We prefer the following explanation of the kinetic D2O effect on the iodide ions oxidation with hydrogen peroxide: According to Pitt et al.15 and Singh,16 in general, the following reaction steps lead to the generation of singlet oxygen
HO2 f O2- + H+
(a)
2O2- f O22- + 1O2
(b)
2 HO2 f H2O2 + 1O2
(c)
O2- + H2O2 f HO- + HO + 1O2
(d)
The D2O effect on the value of IP as well as the described effect of stirring are in accordance with a very large solvent dependence of the 1O2 lifetime reported by McDermott et al.17 According to them, the variation of lifetime extends over six orders of the magnitude. The shortest lifetime has been observed to be 3.1 µs in H2O and 20 µs in D2O. If one compares the effect of stirring in solutions in the presence of D2O, by stirring of solution at 70 rpm the POs increases, which may be due to the partial expulsion of the 1O2 molecules from solution or to their quenching on the walls of the reactor, which diminishes their concentration. The effect of light or illumination in H2O solutions is also impressive (Figure 3). Illumination increases IP as well as frequency of oscillations. Under the same composition of the H2O solution (5 × 10-3 M NaI, 10-1 M H2O2, and 1.2 × 10-1 M HClO4), IP of the reaction in light is more than twice as long (562 s) as that in darkness (214 s). We can use practically the same explanation as we have already done14 with the inhibiting effect of light on the iodine oxidation with H2O2 by interaction of I atoms with I- ions (whose concentration on the start of reaction is relatively high)
I + I- S I2-
(11)
Figure 4. Dependence of induction period on the HClO4 concentration [NaI]0 ) 2 × 10-2 mol dm-3 and [H2O2]0 ) 0.1 mol dm-3, 60 °C, in darkness without stirring.
which are consumed at a low concentration of oxygen by the disproportionation17
2 I2- S I3- + I-
(12)
and in the reaction step
I2- + HO2 f H+ + 2I- + O2
(13)
The effect of light on the frequency of oscillations is in accordance with an excessive excitation (indicated by Raman spectroscopy) of hydrogen peroxide embedded in a hydrogenbonded water network during the oscillatory BL reaction described recently by Stanisavljev et al.20 Such an excitation of the symmetric OH-OH vibration could increase the probability of the formation of OH radicals, which can produce HO2 radicals, and finally by their recombination, 1O2 molecules can be formed.15 The value of IP decreases nonlinearly with increasing concentration of HClO4 (Figure 4). This effect may be due to the reaction step16
HO2 + O2- + H+ f H2O2 + 1O2
(14)
In the presence of amid of acrylamide, IP and the POs increase, and if its concentration attains the value of 0.05 mol dm-3, then the oscillations are suppressed (Figure 5). N3- ions exhibit similar effects. Both acrylamide and N3- ions can function as scavengers of intermediary radicals, and they can be active as energy quenchers or charge transfer quenchers. Energy transfer and charge transfer quenching are two major mechanisms of 1 O2 quenching.21 If one follows the reaction course by measurements of time dependence of iodine absorbance, the pregnant effect of volume of gaseous interphase above solution (Vg) on the values of IP, number of oscillations, and POs is observed (Figure 6). At the value of Vg ) 0 cm3 (a), IP and POs are the shortest, and with increasing value of Vg, both values increase (b,c). These observations evidently indicate that the described effect may be due to the fact that the concentration of intermediary singlet oxygen in the reaction system diminishes as Vg increases. The idea that physical and chemical processes might couple in the gas-evolving oscillators and similar systems was first suggested by Bowers and Noyes.23 Sˇevcˇ´ık et al.24 also indicated that chemical processes may couple to both oxygen and iodine
Oscillatory System I-, H2O2, HClO4
J. Phys. Chem. A, Vol. 114, No. 26, 2010 7029 Acknowledgment. The financial support of grant no 1-003909 from the Scientific Grant Agency of Ministry of Education of Slovak Republic is acknowledged. References and Notes (1) Bray, W. C. J. Am. Chem. Soc. 1921, 43, 1262–1267. (2) Bray, W. C.; Liebhafsky, H. A. J. Am. Chem. Soc. 1931, 53, 38– 44.
Figure 5. Influence of acrylamide (ACA) on the E ) f (t) reaction course [NaI]0 ) 2 × 10-2 mol dm-3, [H2O2]0 ) 8 × 10-2 mol dm-3, and [HClO4]0 ) 0.12 mol dm-3, 60 °C, in darkness (a) [ACA]0 ) 0 mol dm-3, (b) [ACA]0 ) 6 × 10-3 mol dm-3, and (c) [ACA]0 ) 5 × 10-2 mol dm-3.
Figure 6. Volume effect of gaseous phase over solution on reaction course [NaI]0 ) 5 × 10-3 mol dm-3, [H2O2]0 ) 0.1 mol dm-3, and [HClO4]0 ) 9 × 10-2 mol dm-3, 60 °C, no stirring, λ ) 460 nm. (a) Cuvette completely filled in, (b) cuvette filled in 0.5 cm under aperture, and (c) cuvette filled in 1 cm under aperture.
nucleation and evolution to cause, enhance, or suppress oscillations in the BL systems. We realize that the assumptions made in this work also need simulation calculations, and we hope to be able to do so in the near future.
(3) Matsuzaki, I.; Nakajima, T.; Liebhafsky, H. A. Faraday Symp. Chem. Soc. 1974, 9, 55–65. (4) Schmitz, G. J. Chem. Phys. 1974, 71, 689–692. (5) Sharma, K. R.; Noyes, R. M. J. Am. Chem. Soc. 1976, 98, 4345– 4361. (6) (a) Furrow, S. D. J. Phys. Chem. 1987, 91, 1707–1709. (b) Chemical Oscillators Based on Iodate Ion and Hydrogen Peroxide, Chapter 5. In Oscillations and TraVel WaVes in Chemical Systems; Field, R. J., Burger, M. J., Ed.; Wiley: New York, 1985. (7) Kolar-Anic, L.; Schmitz, G. J. Chem. Soc., Faraday Trans. 1992, 88, 2343–2349. (8) Treindl, L.; Noyes, R. M. J. Phys. Chem. 1993, 97, 11354–11362. (9) Laurenzy, G.; Beck, M. T. J. Phys. Chem. 1994, 98, 5188–5189. (10) Noyes, R. M.; Kalachev, L. V.; Field, R. J. J. Phys. Chem. 1995, 99, 3514–3520. (11) Derrick, W. R.; Kalachev, L. V. J. Nonlinear Sci. 2000, 10, 133– 144. (12) La´nˇova´, B.; Vøesˇa´l, J. J. Phys. Chem. A 2002, 106, 1228–1232. (13) Ren, J.; Gao, J. Zh.; Yang, W. Port. Electrochim. Acta 2008, 26/4, 349–360. (14) Olexova´, A.; Mra´kavova´, M.; Melichercˇ´ık, M.; Treindl, L. Collect. Czech. Chem. Commun. 2006, 71, 91–106. (15) Pitt, E.; Scharmann, A.; Suprihadi, T. Z. Z. Naturforscher 1992, 47, 463–468. (16) Singh, A. Photochem. Photobiol. 1978, 28, 429–433. (17) McDermott, W. E.; Pchelkin, N. R.; Benard, D. J.; Bousek, R. R. Appl. Phys. Lett. 1978, 32, 469. (18) Melichercˇ´ık, M.; Olexova´, A.; Treindl, L. J. Mol. Catal. A: Chem. 1997, 127, 43–47. (19) Genigova´, J.; Melichercˇ´ık, M.; Olexova´, A.; Treindl, L. J. Phys. Chem. A 1999, 103, 4690–4692. (20) Stanisavljev, D. R.; Dramicˇanin, M. D. J. Phys. Chem. A 2007, 111, 7703–7706. (21) Zhao, L. Singlet Oxygen; Free Radical and Radiation Biology Graduate Program, The University of Iowa: Iowa City, IA, 2001. (22) Stanisavljev, D. R.; Vukojevicˇ, V. B. J. Phys. Chem. A 2002, 106, 5618–5625. (23) Bowers, P. G.; Noyes, R. M. In Oscillations and TraVeling WaVes in Chemical Systems; Field, R. I., Burger, M., Eds.; Wiley-Interscience: New York, 1985. (24) Sˇevcˇ´ık, P.; Kissimonova´, K.; Adamcˇ´ıkova´, L’. J. Phys. Chem. A 2000, 104, 3958–3963.
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