Autocatalytic Oxidation of Trithionate Ion by Iodate in a Strongly Acidic

Subscriber access provided by UNIVERSITY OF ADELAIDE LIBRARIES

Article

Autocatalytic Oxidation of Trithionate Ion by Iodate in a Strongly Acidic Medium György Csek#, Changwei Pan, Qingyu Gao, Chen Ji, and Attila K. Horváth J. Phys. Chem. A, Just Accepted Manuscript • DOI: 10.1021/acs.jpca.7b06796 • Publication Date (Web): 12 Oct 2017 Downloaded from http://pubs.acs.org on October 13, 2017

Just Accepted “Just Accepted” manuscripts have been peer-reviewed and accepted for publication. They are posted online prior to technical editing, formatting for publication and author proofing. The American Chemical Society provides “Just Accepted” as a free service to the research community to expedite the dissemination of scientific material as soon as possible after acceptance. “Just Accepted” manuscripts appear in full in PDF format accompanied by an HTML abstract. “Just Accepted” manuscripts have been fully peer reviewed, but should not be considered the official version of record. They are accessible to all readers and citable by the Digital Object Identifier (DOI®). “Just Accepted” is an optional service offered to authors. Therefore, the “Just Accepted” Web site may not include all articles that will be published in the journal. After a manuscript is technically edited and formatted, it will be removed from the “Just Accepted” Web site and published as an ASAP article. Note that technical editing may introduce minor changes to the manuscript text and/or graphics which could affect content, and all legal disclaimers and ethical guidelines that apply to the journal pertain. ACS cannot be held responsible for errors or consequences arising from the use of information contained in these “Just Accepted” manuscripts.

The Journal of Physical Chemistry A is published by the American Chemical Society. 1155 Sixteenth Street N.W., Washington, DC 20036 Published by American Chemical Society. Copyright © American Chemical Society. However, no copyright claim is made to original U.S. Government works, or works produced by employees of any Commonwealth realm Crown government in the course of their duties.

Page 1 of 23

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60

The Journal of Physical Chemistry

Autocatalytic Oxidation of Trithionate Ion by Iodate in a Strongly Acidic Medium György Csek˝o,† Changwei Pan,† Qingyu Gao,∗,† Chen Ji,† and Attila K. Horváth∗,‡ †School of Chemical Engineering, China University of Mining and Technology, Xuzhou 2221111, Jiangsu province, P. R. China ‡Department of Inorganic Chemistry, University of Pécs, Ifj´ uság u. 6, H-7624 Pécs, Hungary E-mail: [email protected]; [email protected]

Abstract The trithionate–iodate reaction has been studied spectrophotometrically in an acidic medium at 25.0±0.1◦ C in phosphoric acid/dihydrogen phosphate buffer monitoring the absorbance at 468 nm at the isosbestic point of iodine–triiodide ion system and at I=0.5 M ionic strength adjusted by sodium perchlorate. Main characteristics of the title system are very reminiscent of those found recently in the pentathionate–iodate and the pentathionate–periodate reactions, the systems paving the way for classifying clock reactions. A thorough analysis revealed that the direct trithionate–iodate reaction plays a subtle role only to produce trace amount of iodide ion via a finite sequence of reactions, and once its concentration reaches a certain level, then the reaction is almost exclusively governed by the trithionate–iodine and the iodide–iodate reactions. The title reaction, as expected, was experimentally proven to be autocatalytic with respect to iodide ion. A simple 3-step Landolt-type kinetic model is proposed to describe

1

ACS Paragon Plus Environment


1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60

adequately the most important kinetic features of the title system that can easily be extended to a feasible sequence of elementary and quasi-elementary reactions.

Introduction Polythionates — especially trithionate, tetrathionate and pentathionate — are crucial intermediates of redox transformation of sulfur-containing oxyanions in several environmentally and industrially important chemical processes 1–4 and their role is also inevitable in the metabolism of microorganisms. 5–7 Despite the fact that these species may play crucial roles in these processes their reactions seem to be poorly studied, and many kinetic aspects of their reactions are still awaiting for clarification. Among the known polythionates, oxidation reactions of tetrathionate are the most studied systems, but besides these ones investigated by our research group 8–11 only a small number of independent studies exists in the literature, such as the detailed study of the tetrathionate–iodine, 12 the tetrathionate–oxygen reaction in presence of iron(III)-catalyst 13 and that of the tetrathionate–hydroxyl radical reaction. 14 About the redox transformation of trithionate and pentathionate even less is known, besides the trithionate–iodine, 15 pentathionate–iodine, 16 pentathionate–iodate 17 and pentathionate–periodate 18 systems the ferrate(VI)–polythionate reactions were studied systematically by Read et al. 19,20 Investigation of the redox transformation of polythionates was recently brought into the focus of interest by clarifying the well-known term clock reaction. A recent dispute 21,22 on whether the stoichiometric constraint is a necessary requirement to classify a system as a clock reaction made it possible to distinguish the autocatalysis-driven and substrate-depletive Landolt-type systems. 23 The core model of Landolt-type systems consists of a slow direct reaction between the reactants to produce the autocatalyst followed by formation of the clock species (generally it is iodine) via the autocatalyst–oxidant reaction. In many real examples it is the Dushman reaction. 24 In the third step the clock species (usually iodine) is removed via its reaction with the substrate molecule to reform the auto-

2


Page 2 of 23

Page 3 of 23

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60


catalyst. If this reaction is rapid then the clock species appears only if the substrate molecule is completely consumed providing a straightforward possibility to classify these reactions as substrate-depletive clock reactions. If, however, the rate of this step is commensurable with that of the second one then clock species may form a way in advance to depletion of the substrate molecule. Systems belonging to this group is therefore called autocatalysis-driven clock reactions. In the latter case, thus, substrate–iodine reaction must be a slow process. Since polythionate–iodine reactions are generally iodide-inhibited systems 12,15,16,25 and they proceed relatively slowly, polythionate–iodate and polythionate–periodate systems are good candidates to be classified as real autocatalysis-driven clock reactions. This idea has recently been confirmed by detailed mechanistic investigations of the pentathionate–iodate 17 and the pentathionate–periodate reactions. 18 Taking into consideration the chemical analogy of polythionates trithionate–iodate reaction also seems to be a suitable candidate to expand the group of autocatalysis-driven clock reaction. Herein, we report the main results of the trithionate–iodate reaction. No systematic investigation on this system has yet been available in the literature.

Experimental Section Materials and Buffers Sodium trithionate was first prepared by the reaction of sodium thiosulfate with hydrogen peroxide as described previously. 26,27 The purity of sodium trithionate was found to be better than 99.2%. All the other chemicals (potassium iodate, potassium iodide, phosphoric acid, dihydrogen phosphate and sodium perchlorate) were of the highest purity commercially available and were used without further purification. All solutions were prepared with distilled water (18.2 MΩ−1 cm−1 ) from a Milli-Q purification system. The pH of the solutions was regulated between 1.30 and 2.33 by phosphoric acid/dihydrogen phosphate buffer and the actual pH was calculated by taking the pKa1 of phosphoric acid as 1.8. 28 The total con3



1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60

Page 4 of 23

centration of the buffer components was kept constant at 0.102 M (except when the variation of buffer component concentration was studied) and the ionic strength was always adjusted to 0.5 M by addition of the necessary amount of sodium perchlorate. Temperature of the reaction vessel was controlled at 25.0±0.1◦ C. The initial concentrations of the reactants for all the 99 kinetic runs are depicted in Table 1. Table 1: Initial Concentrations of the Reactants in Different Kinetic Runs No. 1-6 7-12 13-18 19-24 25-30 31-39 40-46 47-52 53-59 60-66 67-73 74-80 81-87 88-94 95-99 *TP O4=[H3 PO4 ]0 +[H2 PO4 – ]0

[S3 O6 2− ]0 /mM 1.0 1.0 1.5 0.5 1.0 0.5 0.5 0.5 1.3–11.6 0.5-3.0 0.3-2.0 0.3-2.0 0.3-2.0 0.15-3.0 0.35-3.0

[IO3 − ]0 /mM 1.0-2.0 0.5-10.0 1.56 1.56 0.5-10.0 1.56 1.56 1.56 0.08-6.93 1.5 1.5 1.5 1.5 1.0 1.0

[I− ]0 /µM 0.0 0.0 0.0 0.0 0.0 0-5.0 0-0.4 0.0 0.0 0.0 0.0 0.0 0.0 0.0 0.0

pH 1.3 1.3 1.3-2.3 1.3-2.3 2.3 1.3 1.9 1.5 1.9 1.9 1.3 1.7 2.3 1.3 2.1

TP O4 */M 0.102 0.102 0.102 0.102 0.102 0.102 0.102 0.0609-0.421 0.102 0.102 0.102 0.102 0.102 0.102 0.102

Methods and Instrumentation The reaction was followed by a Zeiss S600 diode array spectrophotometer in the visible range without using the deuterium lamp. The reaction has been carried out in a standard quartz cuvette equipped with magnetic stirrer and Teflon cap having 1 cm optical path. The buffer components, the reactants trithionate, iodide (if necessary) were delivered from a pipette first. The spectrum of the solution was always recorded before injection of the iodate solution. Then the reaction was started by addition of the necessary amount of iodate solution from a fast delivery pipette. Spectra of the sample at the wavelength range of 400–800 nm were acquired up to approximately 2 days. 4


Page 5 of 23

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60


Data Treatment MRA studies have shown 29 that the only absorbing species in the visible range are iodine and triiodide ion. Therefore evaluation of the kinetic curves was carried out at the isosbestic point 30 of I2 –I− The molar absorbance 3 system at 468 nm by the program package Zita/Chemmech.

of both species was found to be 750 M−1 cm−1 at this wavelength. Originally each kinetic run contained more than 1000 absorbance–time data pairs, therefore it was necessary to reduce the number of time points (80–100) to avoid unnecessary time-consuming calculations. To obtain the kinetic model and the rate coefficients orthogonal fitting procedure has been chosen to minimize the average deviation between the measured and calculated absorbance. Altogether almost 8000 experimental points from the 99 kinetic series were used for the data evaluation.

Results and Discussion Preliminary Observation The absorbance–time profiles of the title reaction at the isosbestic point of the iodine– triiodide system are similar to that observed in the pentathionate–iodate 17 and the pentathionate– periodate 18 reactions. The formation of iodine is delayed and the length of this time lag strongly depends on the initial concentration of the reactants and on the pH as well. Furthermore, the characteristics that absorbance increase can be detected at 800 nm after a fairly long time, may also appear to suggest that the core mechanism of the polythionate–iodate systems must share common features. The only possible end-product having absorption at this wavelength in the present system is the colloidal sulfur. Since light-scattering species may disturb the accurate absorbance measurement, we decided to truncate the experimental curves just before the formation of colloidal sulfur. Such an example can be seen in Figure 1. As shown the time lag for observing colloidal sulfur is always longer than the time needed

5



0.35

0.3

0.25

Absorbance

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60

Page 6 of 23

0.2

0.15

truncation 0.1

0.05

0.0 0

10000

20000

30000

40000

50000

60000

70000

Time (s)

Figure 1: Experimental absorbance–time curves at 468 (blue) and 800 nm (red). Reaction conditions are as follows: pH = 1.3, [S3 O62 – ]0 = 1.5 mM and [IO3 – ]0 = 1.56 mM. for the appearance of iodine, therefore the remaining part of the kinetic curves are suitable to establish a sound kinetic model. This method was successfully tested in cases of the pentathionate–iodate 17 and pentathionate–periodate 18 reactions. In addition to that we also observed, that in excess of trithionate, iodine is a long-lived intermediate due to the relatively slow and iodide-inhibited trithionate–iodine reaction. 15 As expected small amount of iodide ion significantly decreases the time delay of iodine formation making it clear that this species plays a central role as an autocatalyst in this system. Consequently, the trithionate–iodate reaction can also be classified as an autocatalysis-driven clock reaction. The absorbance–time profiles of the measured kinetic curves suggest that initial rate studies are a non-informative way for proper characterization of the concentration dependence of the reactants. To investigate the dependencies of the time lag on the initial concentration of the reactants and that of the pH, we define ti as a time necessary to reach the absorbance of 0.01 absorbance unit at 468 nm, which corresponds to [I2 ] = 1.33×10−5 M, and analyze the concentration dependence of ti . This method has already been applied successfully in cases of the pentathionate–iodate and pentathionate–periodate systems. The main advantage of this definition is that it can exactly be determined experimentally, and no disturbing side reaction (e.g. precipitation of sulfur) can be taken into consideration during the analysis.

6


Page 7 of 23

Concentration Dependence of ti . Figure 2 depicts the dependencies of ti on the concentration of reactants meanwhile keeping the rest of the remaining conditions constant. As we can see, the formal kinetic order of hydrogen ion is clearly below unity, and changes gradually as the pH varies. At lower pHs the formal kinetic order of H+ approaches even zero. It altogether suggests a complex pH-dependence of the governing steps of the reactions. At low pHs the Dushman reaction is relatively fast, consequently in the autocatalytic pathway the pH-independent trithionate–iodine reaction is the rate-determining one. Since at low pHs the formal kinetic order of hydrogen ion approaches zero it straightforwardly means that the initial step of the trithionate–iodate reaction has to be pH-independent. At high pHs when the Dushman reaction becomes the rate-determining step of the autocatalytic pathway, as it is clearly seen, the formal kinetic order of hydrogen ion increases. As Figures 3–4 indicate the formal -4.30 -4.40

log(1/ti)

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60


-4.50 -4.60 -4.70 -4.80 -2.2

-2.0

-1.8

-1.6

-1.4

+

log[H ]0

Figure 2: Dependence of ti on the concentration of hydrogen ion. Reaction conditions are as follows: [S3 O62 – ]0 = 1.5 mM, and [IO3 – ]0 = 1.56 mM (red), [S3 O62 – ]0 = 0.5 mM and [IO3 – ]0 = 1.56 mM (blue), [S3 O62 – ]0 = 2.6 mM and [IO3 – ]0 = 1.5 mM (green). kinetic orders of the IO –3 and the S3 O62 – are also below unity. This result suggests, that the former kinetic order of the reactants have to be unity in the initial step of the reaction. Their lower than unity values can be rationalized by the fact that once the concentration level of iodide reaches its critical point where the reaction is governed by the autocatalytic pathway then both the pH and the iodide concentration determine whether the Dushman- or 7



the trithionate–iodine reaction is the rate-determining one. Since the rate of the Dushman reaction does not depend on [S3 O62 – ] and that of the trithionate–iodine reaction does not depend on [IO3 – ] it consequently means that the formal kinetic orders of both reactants must be lower than one. These results thus altogether confirmed that the initial step of the trithionate–iodate reaction is independent of pH and has to be first order with respect to both reactants. As these arguments hint iodide ion must play a crucial, autocatalytic role -4.20

log(1/ti)

-4.40 -4.60 -4.80 -5.00 -5.20 -3.6

-3.4

-3.2

-3.0

-2.8

-2.6

2-

log[S3O6 ]0

Figure 3: Dependence of ti on the concentration of trithionate ion. Reaction conditions are as follows: [IO3 – ]0 = 1.5 mM for all curves and at pH = 1.30 the slope is 0.34±0.02 (blue); at pH = 1.70 the slope is 0.35±0.09 (green) and at pH = 2.33 the slope is 0.41±0.02 (red). -4.00 -4.20

log(1/ti)

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60

Page 8 of 23

-4.40 -4.60 -4.80 -5.00 -5.20 -3.4

-3.2

-3.0

-2.8

-2.6

-2.4

-2.2

-2.0

log[IO3 ]0

Figure 4: Dependence of ti on the concentration of iodate ion. Reaction conditions are as follows: [S3 O62 – ]0 = 1.0 mM for both curves and at pH = 1.30 the slope is 0.39±0.05 (red) while at pH = 2.33 the slope is 0.620±0.06 (blue). in the present system, so ti has to be influenced strongly upon adding this species initially to the reacting system. As Figure 5 clearly shows ti significantly decreases by the addition of 8


Page 9 of 23

even trace amount of iodide ion. This fact along with the one indicating a sigmoidal profile of the absorbance–time curves clarifies that the trithionate–iodate reaction is autocatalytic with respect to iodide ion.

0.25

Absorbance at 468 nm

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60


0.2 0.15 0.1 0.05 0 0

20000

40000

60000

80000

time (s)

Figure 5: Measured absorbance–time curves at pH = 2.33, [IO3 – ]0 = 1.56 mM and [S3 O62 – ]0 = 0.5 mM at different [I – ]0 /µM = 0.02 (black), 0.06 (blue), 0.2 (green), 0.35 (cyan), 0.9 (red), 1.4 (magenta), 2.8 (brown), 3.5 (yellow), 5.0 (gray).

Proposed Kinetic Model All these observations appear to suggest that the reaction is a Landolt-type system; therefore the following simplified model is able to describe the main characteristics of the trithionate– iodate reaction. As a next step, therefore, individual evaluation of the kinetic curves has been carried out by the following simplified model: 3S3 O62− + 4IO3− + 6H2 O −→ 9SO42− + 4I− + 12H+

(1)

5I− + IO3− + 6H+ −→ 3I2 + 3H2 O

(2)

S3 O62− + 4I2 + 6H2 O −→ 3SO42− + 8I− + 12H+

(3)

− I2 + I− − ↽− −⇀ − I3

(4)

9



The corresponding rate equations were considered as follows: v1 = ki1 [S3 O62− ][IO3− ]

(5)

v2 = ki2 [I− ]2 [IO3− ]

(6)

[S3 O62− ][I2 ] u + [I− ]

(7)

v3 = ki3

v4 = k4 [I2 ][I− ] − k−4 [I3− ]

(8)

During the calculations k4 and k−4 were fixed 5×109 M−1 s−1 and 8.5×106 s−1 , respectively. These values were directly taken from previous literatures 31,32 to provide the logarithm of formation constant of triiodide as 2.83. 28 For each kinetic curve ki1 , ki2 and ki3 values were determined by nonlinear parameter estimation. The calculated absorbance–time series showed excellent fit of the measured data pairs providing an average deviation of 1.07 % by an orthogonal fitting procedure. Table 2 contains representative results of the individual nonlinear fit providing the apparent rate coefficients of steps (1)–(3). The final results are illustrated in Figures 5 and 6–9.

Careful inspection of the calculated data revealed that

0.25


1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60

Page 10 of 23

0.2 0.15 0.1 0.05 0 0

10000 20000 30000 40000 50000 60000 70000

time (s)

Figure 6: Measured (symbols) and calculated (solid lines) absorbance–time curves at pH = 1.30 and [S3 O62 – ]0 = 1.0 mM. [IO3 – ]0 /mM = 0.5 (black), 1.0 (blue), 2.0 (green), 4.0 (cyan), 5.0 (red), 10.0 (magenta). ki1 is independent of the pH and the base (dihydrogen-phosphate) concentration providing 10


Page 11 of 23

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60


Table 2: Representative Result of the Individual Fit. No. 1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25

105 ki1 M−1 s−1 0.701 1.03 1.34 0.869 1.22 1.51 1.42 1.3 1.15 1.15 1.33 1.21 1.16 1.51 0.842 1.32 1.56 1.46 1.28 1.07 1.25 1.07 1.71 1.03 0.661

ki2 −2 −1 M s 23109 9621 4111 49122 22504 10422 4205 70251 121856 152117 180118 220035 291138 25103 25699 4351 4113 4196 3932 5095 4095 4583 10931 4522 4549

105 ki3 s−1 2.42 2.82 3.22 2.34 2.47 2.13 3.77 2.70 2.64 2.73 2.71 2.63 2.91 2.51 2.04 2.58 2.26 2.48 3.66 2.16 2.31 2.16 2.44 3.22 3.77

[S3 O6 2− ]0 mM 1.5 1.5 1.5 0.5 1.5 1.5 1.5 0.5 0.5 0.5 0.5 0.5 0.5 1.3 1.73 0.3 0.5 0.7 1.05 1.5 1.7 2.0 1.0 1.0 1.0

11

[IO3 – ]0 mM 1.56 1.56 1.56 1.56 1.56 1.56 1.56 1.56 1.56 1.56 1.56 1.56 1.56 0.78 1.04 1.5 1.5 1.5 1.5 1.5 1.5 1.5 1.0 2.0 4.0

[H+ ] mM 12.6 7.94 4.73 19.9 12.6 7.94 4.73 31.6 31.6 31.6 31.6 31.6 31.6 12.6 12.6 4.73 4.73 4.73 4.73 4.73 4.73 4.73 7.94 4.73 4.73


[H3 PO4 ] mM 45.0 33.9 23.4 56.6 45.0 33.9 23.4 40.6 81.2 122 162 203 274 45.0 45.0 23.4 23.4 23.4 23.4 23.4 23.4 23.4 33.9 23.4 23.4

[H2 PO4 – ] mM 56.6 67.7 78.2 45.0 56.6 67.7 78.2 20.3 40.7 61.0 81.0 102 137 56.6 56.6 78.2 78.2 78.2 78.2 78.2 78.2 78.2 67.7 78.2 78.2


0.25


1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60

Page 12 of 23

0.2

0.15

0.1

0.05

0 0

20000

40000

60000

80000

time (s)

Figure 7: Measured (symbols) and calculated (solid lines) absorbance–time curves at [IO3 – ]0 = 1.56 mM and [S3 O62 – ]0 = 1.0 mM. pH = 1.30 (black), 1.50 (blue), 1.70 (green), 1.90 (cyan), 2.10 (red), 2.33 (magenta). an average value of (1.18±0.29)×10−5 for ki1 . From these data we concluded that the initial step of the trithionate–iodate is probably a rate-determining formation of a weak adduct of the reactants followed by subsequent but rapid reactions to give the overall stoichiometry of eq. (1). The attack of iodate might take place on the β-sulfur of trithionate weakening the sulfur-sulfur bond to be broken. Once the sulfur–sulfur bond breaks up the reaction proceeds rapidly to produce eventually sulfate ion. A conceivable sequence of such reactions may be considered as follows: rate det.

S3 O62− + IO3−

−→

S3 O6 IO33−

f ast

S3 O6 IO3− + 2H2 O −→ 3HSO3− + HIO2 f ast

S3 O62− + HIO2 + 2H2 O −→ 3HSO3− + HOI + H+ f ast

S3 O62− + HOI + 2H2 O −→ 3HSO3− + I− + H+ f ast

3HSO3− + IO3− −→ 3SO42− + I− + 3H+

12


(9)

(10) (11) (12) (13)

Page 13 of 23

0.6


1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60


0.5 0.4 0.3 0.2 0.1 0 0

20000

40000

60000

80000

time (s)

Figure 8: Measured (symbols) and calculated (solid lines) absorbance–time curves at constant [S3 O62 – ]0 /[IO3 – ]0 ≈1.67 ratio and at pH=1.90. [S3 O62 – ]0 /mM = 1.3 (black), 1.73 (blue), 2.6 (green), 3.7 (cyan), 4.3 (red), 8.7 (magenta), 11.6 (yellow). As seen linear combination of (9)+(10)+(11)+(12)+3×(13) establishes the stoichiometry indicated by eq. (1). It is worthwhile to note that the calculated rate coefficient (ki1 ) of the overall reaction (eq. (1)) was found to be (1.13±0.29)×10−5 M−1 s−1 , indicating a vanishingly slow process. This very small value may also question whether this direct thermal process exists at all and the initiation of reaction is rather the consequence of iodide impurity of iodate. Therefore additional calculations were performed where the initiating step was completely neglected from the kinetic model and instead of it trace amount of iodide impurities was fitted in case of each kinetic run. Of course this impurity has to be proportional to the concentration of iodate applied in each run because the same iodate stock solution was used throughout the whole experimental series. The data obtained revealed that there was no correlation at all between the fitted iodide and the initial iodate concentration. The best agreement between the measured and calculated data was found to be at 0.0022 % iodide impurity with respect to the stock iodate concentration but even in this case more than 7 % average deviation was obtained. From this we concluded that the kinetic curves cannot be described adequately by supposing the fact that inherently involved iodide impurity initiates the system with having no direct reaction (eq. (1)) at all. As expected, ki2 values indicate a strong pH-dependence shown in Table 2. It is well13




0.3 0.25 0.2 0.15 0.1 0.05 0 0

10000

20000

30000

40000

time (s)

Figure 9: Measured absorbance–time curves at pH = 1.50 using the following initial concentrations of the reactants of [S3 O62 – ]0 = 0.5 mM and [IO3 – ]0 = 1.6 mM with different [H2 PO4 – ]0 /M = 0.041 (black), 0.082 (blue), 0.122 (green), 0.162 (cyan), 0.203 (red), 0.274 (magenta) concentrations. known that rate of the overall Dushman reaction is proportional to the square of hydrogen ion concentration at moderate pHs. 24,33 Therefore, first the apparent rate coefficients (ki2 ) ′

obtained by the individual nonlinear fit of the kinetic curves are converted into ki2 according to the following expression: ki2 = ki2 /[H+ ]2 . Furthermore, the pH-dependence of ki2 ′

′

is complicated by the fact that ki2 (and of course ki2 values as well) are also affected by [H2 PO4 – ] at constant pH value. This linear relationship is illustrated in Figure 10. This

2.7 2.3 1.9 1.5

′

ki2 ×10−8 (M−4 s−1 )

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60

Page 14 of 23

1.1 0.7 0.02

0.04

0.06

0.08

0.1

0.12

[H2 PO− 4 ]0 (M)

Figure 10: Plot of ki2 against the base (H2 PO4 – ) concentration at constant pH. Conditions are as follows: [S3 O62 – ]0 = 0.5 mM; [IO3 – ]0 /mM = 1.56 mM and pH = 1.5. Slope of the straight line is (1.82±0.07)×109 and its intercept is (3.87±0.59)×107 . ′

14


Page 15 of 23

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60


connection enables us to determine the rate coefficient of the phosphate base-assisted pathway and that of the buffer-unassisted route of the Dushman reaction. Linear relationship mentioned above suggests that ki2 = k2 [H+ ]2 + k2 [H+ ]2 [H2 PO4 – ] equation holds, where ′

′

′

k2 = (3.87±0.59)×107 M−4 s−1 and k2 = (1.82±0.07)×109 M−5 s−1 are the rate coefficient of the buffer-unassisted and buffer-assisted pathways, respectively. The first experimental evidence that certain buffer components may catalyze the halogenate–halide reactions was noticed by Barton and Wright. 34 In case of acetate buffer and at an 1.0 M ionic strength they obtained the values 3×108 M−4 s−1 and 6.1×109 M−5 s−1 for the buffer-unassisted and the buffer-assisted pathway. Considering the different experimental circumstances the agreement between their data and ours is sound. In terms of elementary reactions eq. (2) can therefore be rationalized according to the following reaction sequence: −− 2H+ + I− + IO3− ↽ −⇀ − I2 O2 + H2 O

(14)

I2 O2 + I−

rate det.

I2 + IO2−

(15)

I2 O2 + I− + B

rate det.

I2 + IO2− + B

(16)

−→

−→

H+ + IO2− − ↽− −⇀ − HIO2 f ast

(17)

HIO2 + I− + H+ −→ 2HOI

(18)

HOI + I− + H+ − ↽− −⇀ − I2 + H2 O,

(19)

where B stands for H2 PO4 – in our case. All the equilibria in this sequence are established rapidly compared to the time scale of the title reaction and the rate determining steps are eqs. (15) and (16). Rest of the reactions, as indicated, are considered as rapid processes. Finally, Table 2 also suggests that ki3 , the apparent rate coefficient of the trithionate– iodine reaction is basically independent of pH in agreement with a previous study. 15 As shown in eq. (7) a constant u parameter is introduced to prevent v3 being infinity at the

15



1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60

Page 16 of 23

beginning phase of the reaction. This parameter was set to a constant value of 10−8 M during the whole course of the calculation process meaning that the iodide inhibitory effect becomes significant when [I – ] substantially greater than 10−8 M. Of course the reason of introducing u is pointing far beyond the fact just to simply avoid the rate of corresponding reaction being infinity. It should also provide a certain connection between the rate coefficients of the elementary reactions driving the trithionate–iodine reaction, thus the source of u strongly depends on the sequence of elementary reactions proposed. Below (see: eqs. (20)–(22)) we present a simply and feasible, but not the only possibility. Since the experiments presented here do not provide a solid basis to unambiguously divide eq. (3) into elementary reaction, the discussion about source of u would merely be a speculation. It is also important to note that the average value ki3 =(2.7±0.5)×10−5 s−1 obtained from our experiments, however, differs by a factor 30 reported in the study mentioned above. 15 Comparison of the experimental circumstances reveals one and only difference between these studies. In studying directly the trithionate–iodine reaction acetate buffer was used instead of phosphate buffer. It appears to suggest that the trithionate–iodine reaction is also subject to buffer-catalysis. In case of the Dushman reaction it is well-understood that catalytic activity of carboxylate ions 34 is more pronounced than that of phosphate ions. Based on our results presented here, we suggest that the same phenomenon may occur as well in case of the trithionate–iodine reaction. Last but not least eq. (3) can also be rationalized in terms of elementary reactions. As it was established by Csek˝o and Horváth 15 strong iodide autoinhibition of this reaction may be explained by a rapidly established iodonium ion transfer from iodine to the β-sulfur of trithionate to produce S3 O6 I – and iodide ion. The intermediate S3 O6 I – is then hydrolyzes in a relatively slow process into bisulfite ion that can react further with iodine in a subsequent fast reaction. 35 − − −− S3 O62− + I2 ↽ −⇀ − S 3 O6 I + I

(20)

S3 O6 I− + 3H2 O −→ 3HSO3− + I− + 3H+

(21)

16


Page 17 of 23

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60


f ast

HSO3− + I2 + H2 O −→ SO42− + 2I− + 3H+

(22)

Finally, a word is also in an order here to clarify that the kinetic model presented is certainly not the complete model of the title system as it is unable to explain the formation of colloidal sulfur. As it was pointed out previously those parts of the kinetic curves where elementary sulfur precipitates — and thus disturbs the measured absorbance–time signals — were truncated. Therefore there is not enough quantitative experimental information left based on which a firmly based sequence of reaction leading to the formation of elementary sulfur may be proposed.

Conclusion In this work kinetics and mechanism of the trithionate–iodate reaction was presented for the first time. All the experiments suggest that a simple, three-step Landolt-type kinetic model is working consistently well to describe the characteristics of the title reaction. One of the main kinetic feature of this system was found to be the iodide-driven autocatalysis and since appearance of iodine is not concurrent with the depletion of substrate molecule the title system is a prototype of autocatalysis-driven clock reactions. We have also shown that the direct trithionate–iodate reaction plays only a subtle role to produce iodide ion via a finite sequence of reactions, and once iodide concentration reaches a certain level, then the reaction is almost exclusively governed by the trithionate–iodine and the iodide–iodate reactions. This study also demonstrates that the family of autocatalysis-driven clock reaction may easily be extended in the near future based on the common features to be shared by the polythionate–iodate and polythionate–periodate reactions. Furthermore, the application of the autocatalysis-driven clock reactions in chemically driven hydrodynamic patterns could as well display new spatiotemporal dynamics such as fronts and pulses of iodine which would be different from the pH fronts observed in the chlorite–polythionate systems. 36,37

17



1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60

Page 18 of 23

Acknowledgment This work was supported by the National Natural Science Foundation of China (Grant No. 21573282), the Fundamental Research Funds for the Central Universities (Grant No. 2015XKZD09), the Natural Science Foundation of Jiangsu Province (Grant No. BK20160240), the Hungarian Research Fund NKFIH-OTKA grant no. K116591 and the GINOP-2.3.2-152016-00049 grant. We also thank the valuable comments of an anonymous reviewer that helped us to improve this manuscript.

References (1) Steudel, R.; Holdt, G.; Göbel, T.; Hazeu, W. Chromatographic Separation of Higher Polythionates and Their Detection in Thiobacillus Ferrooxidants. Angew. Chem. Int. Ed. 1987, 26, 151–153. (2) Schippers, A.; Sand, W. Bacterial Leaching of Metal Sulfides Proceeds by Two Indirect Mechanisms via Thiosulfate or via Polysulfides and Sulfur. Appl. Environ. Microbiol. 1999, 65, 319–321. (3) Takano, B.; Zheng, Q.; Ohsawa, S. A Telemetring System for Monitoring Aqueous Polythionates in an Active Crater Lake. J. Volcanol. Geotherm. Res. 2000, 97, 397– 406. (4) Xu, Y.; Schoonen, M. A. A.; Nordstrom, D. K.; Cunningham, K. M.; Ball, J. W. Sulfur Geochemistry of Hydrothermal Waters in Yellowstone National Park, Wyoming, USA. II. Formation and Decomposition of Thiosulfate in Cinder Pool. J. Volcanol. Geotherm. Res. 2000, 97, 407–423. (5) Chambers, L. A.; Trudinger, P. A. Microbiological Fractionation of Stable Sulfur Isotopes: A Review and Critique. Geomicrobiol. J. 1979, 1, 249–293.

18


Page 19 of 23

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60


(6) Suzuki, I. Oxidation of Inorganic Sulfur Compounds: Chemical and Enzymatic Reactions. Can. J. Microbiol. 1998, 45, 97–105. (7) Kelly, D. P. Thermodynamic Aspects of Energy Conservation by Chemolithotrophic Sulfur Bacteria in Relation to the Sulfur Oxidation Pathways. Arch. Microbiol. 1999, 171, 219–229. (8) Horváth, A. K.; Nagypál, I. Kinetics and Mechanism of the Reaction Between Hypochlorous Acid and Tetrathionate Ion. Int. J. Chem. Kinet. 2000, 32, 395–402. (9) Horváth, A. K.; Nagypál, I.; Epstein, I. R. Kinetics and Mechanism of the Chlorine Dioxide–Tetrathionate Reaction. J. Phys. Chem. A 2003, 107, 10063–10068. (10) Varga, D.; Horváth, A. K. Simultaneous Evaluation of Different Types of Kinetic Traces of a Complex System: Kinetics and Mechanism of the Tetrathionate–Bromine Reaction. J. Phys. Chem. A 2009, 113, 9988–9996. (11) Filáry, A.; Horváth, A. K. Photochemically Induced Catalysis of Iodide Ion and Iodine in the Tetrathionate–Periodate Reaction. Phys. Chem. Chem. Phys. 2010, 12, 6742– 6749. (12) Awtrey, A. D.; Connick, R. E. Rate Law and Mechanism of the Reaction of Iodine with Tetrathionate Ion. J. Am. Chem. Soc. 1951, 73, 4546–4549. (13) Druschel, G. K.; Hamers, R. J.; Banfield, J. F. Kinetics and Mechanism of Polythionate Oxidation to Sulfate at Low pH by O2 and Fe3+ . Geochim. Cosmochim. Acta 2003, 67, 4457–4469. (14) Druschel, G. K.; Hamers, R. J.; Luther, G. W.; Banfield, J. F. Kinetics and Mechanism of Trithionate and Tetrathionate Oxidation at Low pH by Hydroxyl Radical. Aquatic Geochem. 2004, 9, 145–164.

19



1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60

(15) Csek˝o, G.; Horváth, A. K. Non-Triiodide Based Autoinhibition by Iodide Ion in the Trithionate–Iodine Reaction. J. Phys. Chem. A 2010, 114, 6521–6526. (16) Xu, L.; Csek˝o, G.; Kégl, T.; Horváth, A. K. General Pathway of Sulfur-Chain Breakage of Polythionates by Iodine Confirmed by the Kinetics and Mechanism of the Pentathionate–Iodine Reaction. Inorg. Chem. 2012, 51, 7837–7843. (17) Xu, L.; Horváth, A. K. A Possible Candidate to Be Classified as an AutocatalysisDriven Clock Reaction: Kinetics of the Pentathionate–Iodate Reaction. J. Phys. Chem. A 2014, 118, 6171–6180. (18) Xu, L.; Horváth, A. K. Autocatalysis-Driven Clock Reaction II: Kinetics of the Pentathionate–Periodate Reaction. J. Phys. Chem. A 2014, 118, 9811–9819. (19) Read, J. F.; Bewick, S. A.; Donaher, S. C.; Eelman, M. D.; Oakey, J.; Schaubel, C.; Tam, N. C.; Theriault, A.; Watson, K. J. The Kinetics and Mechanism of the Oxidation of Inorganic Oxysulfur Compounds by Potassium Ferrate. Part III–Trithionate and Pentathionate Ions. Inorg. React. Mech. 2005, 5, 281–304. (20) Read, J. F.; Bewick, S. A. The Kinetics and Mechanism of the Oxidation of Inorganic Oxysulfur Compounds by Potassium Ferrate. Part IV–A Theoretical Analysis of Four Models Proposed to Explain Some of the Unusual Results for Trithionate, Tetrathionate and Pentathionate Ions. Inorg. React. Mech. 2005, 5, 305–330. (21) Lente, G.; Bazsa, G.; Fábián, I. What is and what isn’t a Clock Reaction? New J. Chem. 2007, 31, 1707. (22) Sant’Anna, R. T. P.; Monteiro, E. V.; Pereira, J. R. T.; Faria, R. B. The Ozone–Iodine– Chlorate Clock Reaction. Plos One 2013, 8, e83706. (23) Horváth, A. K.; Nagypál, I. Classification of Clock Reactions. Chem. Phys. Chem. 2015, 16, 588–594. 20


Page 20 of 23

Page 21 of 23

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60


(24) Dushman, S. The Rate of the Reaction Between Iodic and Hydroiodic Acids. J. Phys. Chem. 1904, 8, 453–482. (25) Kerek, A.; Horváth, A. K. Kinetics and Mechanism of the Oxidation of Tetrathionate by Iodine in a Slightly Acidic Medium. J. Phys. Chem. A 2007, 111, 4235–4241. (26) Kelly, D. P.; Wood, P. A. [35] Synthesis and Determination of Thiosulfate and Polythionates. Methods in Enzymology 1994, 234, 475–501. (27) Csek˝o, G.; Horváth, A. K. Non-Triiodide Based Autoinhibition by Iodide Ion in the Trithionate–Iodine Reaction. J. Phys. Chem. A 2010, 114, 6521–6526. (28) IUPAC Stability Constant Database. Royal Society of Chemistry: London, 1992-1997. (29) Peintler, G.; Nagypál, I.; Jancsó, A.; Epstein, I. R.; Kustin, K. Extracting Experimental Information from Large Matrixes. 1. A New Algorithm for the Application of Matrix Rank Analysis. J. Phys. Chem. A 1997, 103, 8013–8020. (30) Peintler, G. ZiTa, version 5.0; A Comprehensive Program Package for Fitting Parameters of Chemical Reaction Mechanism. Attila József University: Szeged, Hungary, 1989-1998. (31) Turner, D. H.; Flynn, G. W.; Sutin, N.; Beitz, J. V. Laser Raman Temperature-jump Study of the Kinetics of the Triiodide Equilibrium. Relaxation Times in the 10−8 –10−7 Second Range. J. Am. Chem. Soc. 1972, 94, 1554–1559. (32) Ruasse, M.; Aubard, J.; Galland, B.; Adenier, A. Kinetic Study of the Fast Halogen– Trihalide Ion Equilibria in Protic Media by the Raman-Laser Temperature-Jump Technique. A Non-Diffusion-Controlled Ion–Molecule Reaction. J. Phys. Chem. 1986, 90, 4382–4388. (33) Schmitz, G. Kinetics and Mechanism of the Iodate–Iodide Reaction and Other Related Reactions. Phys. Chem. Chem. Phys. 1999, 1, 1909–1914. 21



1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60

(34) Barton, A. F. M.; Wright, G. A. Kinetics of the Bromate–Iodide Reaction: Catalysis by Acetate and Other Carboxylate Ions. J. Chem. Soc 1968, 1747–1753. (35) Yiin, B. S.; Margerum, D. W. Nonmetal Redox Kinetics: Reactions of Iodine and Triiodide with Sulfite and Hydrogen Sulfite and the Hydrolysis of Iodosulfate. Inorg. Chem. 1990, 29, 1559–1564. (36) Horváth, D.; Bánsági, T.; Tóth, A. Orientation-Dependent Density Fingering in an Acidity Front. J. Chem. Phys. 2002, 117, 4399–4402. (37) Liu, Y.; Zhou, W.; Zheng, T.; Zhao, Y.; Gao, Q.; Pan, C.; Horváth, A. K. ConvectionInduced Fingering Fronts in the Chlorite–Trithionate Reaction. J. Phys. Chem. A 2016, 120, 2514–2520.

22


Page 22 of 23

Page 23 of 23

TOC Graphic

Autocatalysis-driven Clock Reaction

Absorbance

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60


Time

23


Autocatalytic Oxidation of Trithionate Ion by Iodate in a Strongly Acidic

Recommend Documents