J . Phys. Chem. 1989, 93, 1689-1691
1689
Complex Dynamical Behavior in the Oxidation of Thiocyanate by Iodate' Reuben H. Simoyi, Department of Chemistry, University of Zimbabwe, Box MP167, Mount Pleasant, Harare, Zimbabwe
Irving R. Epstein, and Kenneth Kustin* Department of Chemistry, Brandeis University, Waltham, Massachusetts 02254-91 10 (Received: December 14, 1988)
Complex dynamical behavior including oligooscillation (multiple extrema in concentration as a function of time) has been observed in the oxidation of thiocyanate by iodate in acidic medium. The stoichiometry of the reaction when thiocyanate is in stoichiometricexcess over iodate is IO3- SCN- + H 2 0 SO4,- + CN- I- 2H'. In excess iodate the stoichiometry I, + 5ICN SO:- + H 2 0 . In high acid concentrations the reaction initially produces iodine, is 710,- + SSCN- + 2H' and then later the iodine is consumed. In excess thiocyanate all the iodine produced is subsequently consumed, while in excess iodate some iodine is left at the end of the reaction. This behavior is explained via a network of nine reactions which are viable in acidic mixtures of iodate and thiocyanate.
-
+
+
The oxidation of thiocyanate by various oxidizing agents has received considerable attention in the past 20 year^.^-^ In all cases, the kinetics and mechanisms of these oxidations have been found to be complex. Sulfate and cyanide are the usual oxidation products, although with a strong enough oxidizing agent, cyanate, ammonia, and carbonate have also been ~ b t a i n e d . ~ , ' The involvement of sulfur-containing compounds in novel chemical oscillators that are pH driven8 has given added importance to the mechanisms of oxidation of these compounds. The reactions of thiocyanate in a continuous stirred tank reactor (CSTR) with chloriteg and with bromatel0 have been found to show oscillatory behavior. No detailed kinetics studies have been performed on the oxidation of thiocyanate by chlorite and by bromate. We have thus been interested in the kinetics and mechanisms of the oxidation of sulfur compounds, especially thiocyanate, by halogen-containing compounds. Our earlier study of the oxidation of thiocyanate by a mild oxidizing agent, iodine," showed the reaction to be extremely complex, characterized by variable stoichiometry and autoinhibition. We report here an even more complex oxidation of thiocyanate, this time by iodate. We describe so far only the dynamical behavior of the reaction in a batch reactor, and we attempt to explain the origin of the observed oligooscillatory'2 behavior, Le., the occurrence of multiple extrema in iodine concentration as a function of time.
Experimental Section Materials. The following analytical grade chemicals were used without further purification: potassium iodate, potassium thiocyanate, silver nitrate, sodium thiosulfate, potassium iodide (free-flowing granular), and perchloric acid, 70% (Fisher). Stock solutions of 5 M sodium perchlorate (Aldrich) were filtered before use. The perchloric acid was standardized with standard N/10 (1) Part 53 in the series Systematic Design of Chemical Oscillators. Part 52: Rlbai, G.; Kustin, K.; Epstein, I. R. J. Am. Chem. SOC.,in press. (2) Radhakrishnamurti, P. S.; Misra, S. A,; Panda, J. K. Indian J. Chem. 1981, 20A, 459. (3) Smith, R. H.; Wilson, I. R. Aust. J . Chem. 1966, 19, 1357. (4) Orszagh, I.; Bazsa, G.; Beck, M. T. Inorg. Chim. 1972, 6, 271. (5) Lewis, C.; Skoog, D. A. J . A m . Chem. SOC.1962, 84, 1101. (6) Wilson, I. R.; Harris, G. M. J. A m . Chem. SOC. 1960, 82, 4515. (7) Wilson, I. R.; Harris, G.M. J . A m . Chem. SOC.1961, 83, 286. (8) Orbln, M.; Epstein, I. R. J . A m . Chem. SOC.1987, 109, 101. (9) Alamgir, M.; Epstein, I. R. J. Phys. Chem. 1985, 89, 3611. (10) Simoyi, R. H. J . Phys. Chem. 1987, 91, 1557. (11) Simoyi, R. H.; Epstein, I. R.; Kustin, K. J . Phys. Chem., in press. (12) Ribai, G.; Bazsa, G.; Beck, M. T. J . A m . Chem. SOC.1979, 101, 6746.
0022-3654/89/2093-1689$01.50/0
-
+ +
sodium hydroxide (Fisher). In some cases, iodate solutions were standardized by adding excess acidified iodide and titrating the liberated iodine against sodium thiosulfate. This titration was performed once on each fresh batch of iodate to establish consistency. Methods. The reaction was followed by monitoring the absorbance of I2 and 13- at their isosbestic point of 460 nm on Beckman Model 25 and Pye Unicam SP1750 UV/visible spectrophotometers. Cyanide interferes with the quantitative measurement of iodide by a specific-ion electrode, and after a few seconds of use, the electrode surface becomes coated with black sulfide deposits; therefore this technique was not used. The reactions were carried out in 10-mm quartz photocells with Teflon caps to prevent escape of the 1, formed in the reaction. The SP1750 spectrophotometer was equipped with a stirrer to ensure constant agitation of the reaction solution. All reactions were performed at 25 f 0.1 OC and 0.2 M ionic strength (NaC104). In high acid conditions, the reaction became so fast that it was followed on a Hi-Tech SF-3L stopped-flow spectrophotometer. After amplification, the photomultiplier output was digitized via a 12-bit Metrabyte A / D converter and interfaced to a Tandy HD-1200 microcomputer for storage. Stoichiometric Determinations. All reaction solutions used for stoichiometric determinations were stored for at least 48 h before any analyses were performed. These solutions were kept tightly capped, with the reaction vessels wrapped in aluminum foil and stored in the dark.13 When the thiocyanate was in stoichiometric excess over the iodate, the excess thiocyanate was determined by titration with standard silver nitrate using ferric ammonium sulfate as the i n d i ~ a t 0 r . I ~These titrations were performed rapidly, because the nitric acid used to stabilize the indicator eventually oxidizes the thiocyanate in an autocatalytic manner.I5 This technique gave a total titre that also included contributions from iodide and cyanide. In some cases the reagent solution was acidified with perchloric acid before the titration to expel the cyanide as HCN. The sulfate produced was estimated by gravimetric analysis as BaS0,. Iodine formed was determined spectrophotometrically only, since there were several other oxidizing species in solution. In excess iodate conditions, both iodate and sulfate were determined by precipitating simultaneously BaSO, and Ba(103)2. (1 3) Aqueous iodine easily forms 12- when subjected to radiation in the visible region. See: Schwartz, H. A,; Bielski, H. J. J . Phys. Chem. 1986, 90, 1445. (14) Vogel, A. I. Textbook of Quantitative Inorganic Analysis, 3rd ed.; Wiley: New York, 1961; p 265. (15) Stedman, G.;Whincup, P. A. E. J . Chem. SOC.A 1969, 1145.
0 1989 American Chemical Society
1690 The Journal of Physical Chemistry, Vol. 93, No. 5, 1989 I
I--
I
1
3
-- -
I
I
I
I
0 OLM SCN
SCN 0 0 2 M SCN
0 03H
W
u
z 4
-
[ I
- _ _0 0 2 M IO;
m 0 I/)
003M10; 0 0 1 M IO;
m Q
'----______ 1
1
0
3
2
.
1
1
1
I
I
I
I
I
4
5
6
7
0
2 3 4 TIME I minutes 1
1
TIME I minutes)
Figure 1. Absorbance ( A = 460 nm) vs time at different thiocyanate concentrations in excess thiocyanate, showing the transient formation of iodine [I03-lO= 0 001 M, [H+l0= 0.005 M
5
Figure 2. Absorbance traces in excess iodate, showing the oligooscillatory behavior, [SCN-Io = 0.0003 M, [H+]o = 0.02 M. 1
Results Stoichiometry. The stoichiometry of the reaction varied with
1
I
1
the ratio of initial concentrations of iodate and thiocyanate. When thiocyanate was in excess over iodate (by at least a factor of 2), we found the stoichiometry
IO3- + SCN-
+ H20
-
+ CN- + I- + 2H+
(RI)
The silver nitrate titrations gave titres that were lower than expected from ( R l ) mainly because some of the CN- ions were removed from solution as HCN, giving the stcichiometry
IO3- + SCN-
+ H20
-
SO4,-
+ HCN + I- + H+
(Rla)
Acidified reactant solutions gave more accurate results. The 1 :I ratio of iodate to thiocyanate consumed was maintained over a wide range of initial concentration ratios (BaSO, gravimetric analysis). With iodate in excess over thiocyanate (at least 5-fold), the stoichiometry was much more complex:
710,-
+ SSCN- + 2H+
-
I,
+ 5ICN +
+ H 2 0 (R2)
At [H+IoI 5 X M, the ratio of thiocyanate consumed to iodine produced (A = 460 nm) was as in (R2), 5:1. As the initial acid concentration was increased, however, the iodine produced increased to approach a limiting ratio of 5:3 (it never attained this ratio), and the stoichiometry was nearly 6103- + 5SCN-
+ 2Hz0
-
312 + 5s042+ 5CN-
+ 4H'
(R3) In such high acid concentrations, the cyanide ions exist overwhelmingly as HCN. Generally, also, ICN is unstable in the presence of iodide and acid:
+ IICN + I- + H+ e HCN + I, H + + CNHCN CN-
+ 12 P $
ICN
(R4) (ref 16)
(R5) (R6)
Gravimetric analysis of precipitates of BaS04 + Ba(103)zgave values that were about 95% of those expected from stoichiometry R2. In excess thiocyanate, however, the gravimetric analysis using BaS0, was much more accurate, at 98%. Reaction Dynamics. The reaction dynamics were also very complex, being dependent both on the ratio of the initial concentrations of iodate and thiocyanate and on the initial acid concentrations (16) Griffith, R. 0.;McKeown, A. Trans. Faraday SOC.1935, 31, 868
40 60 TIME (minutes) Figure 3. Excess iodate conditions in low acid concentrations. There is a monotonic increase in I2 absorbance up to the stoichiometric value. [IO [I2],,. In fact, when thiocyanate is in sufficient (about 10-fold) excess over iodine, the reaction is quite fast, and pseudo-first-order kinetics can be observed.I8 Reaction M3 is significant only after SCN- is oxidized and the sulfurcarbon bond has been cleaved. Reaction M5 is necessary to achieve an accurate depiction of the iodine species present. In excess iodate conditions (Figure 2), reaction also commences with ( M l ) and (M2). The difference now is that thiocyanate is depleted instead of iodate. The drop in [I2] next observed results from reaction M3. The further rise in iodine concentration at the end of the reaction is caused by (M6), which depletes ICN and yields the stoichiometric amount of iodine in highly acidic environments. Oxidation of thiocyanate goes through several intermediates" before formation of CN-, which will then consume the iodine formed by (M2). When iodate is in excess, the overall oxidation of thiocyanate, which produces iodide, is controlled by the [H+] dependence of (M2). If [H'] is sufficiently high, iodine is produced rapidly enough to accumulate; lower [H'] places (M2) at a kinetic disadvantage compared with ( M l ) and (M4). The reaction network we have proposed qualitatively explains the observed oligooscillatory behavior. The mechanisms of reactions M1 and M7 (or M8) have not, however, been fully explored.
Acknowledgment. We acknowledge the University of Zimbabwe for granting leave of absence to R.H.S. during the summer of 1987. We thank Momingstar Manyonda and Jonathan Masere, who performed some of these experiments. This work was supported by Research Grant CHE-8800169 from the National Science Foundation and Grant 2.9999.10:2789 from the University of Zimbabwe Research Board.
Effects of Intramolecular Hydrogen Bonding on the Rates of Complex Formation of Co2+ and Zn2+ with Substltuted Salicylic Acids H. Diebler,* Max- Planck- Institut fur Biophysikalische Chemie, Gottingen- Nikolausberg, West Germany
F. Secco, and M. Venturini Department of Chemistry, University of Pisa, Pisa, Italy (Received: April 8, 1988;
In Final Form: January 10, 1989) The kinetics of the complex formation of Co2+and Zn2+with two substituted salicylates have been investigated and compared to those of Ni2+(reported previously). The rate constants for the unprotonated salicylates are consistent with a rate-determining first substitution step at the respective aquometal ion. Because of strong intramolecular hydrogen bonding, the monoprotonated ligands are much less reactive (80-4000-fold). The metal ion dependence of the rate constants for these species leads to the conclusion that in the reactions involving NiZ+and Co2+ ring closure is rate determining, whereas in case of the very labile Zn(I1) the opening of the internal H bond in the initial complex becomes rate determining. In a previous publication we have reported on the kinetics of complex formation of Ni2+ with a series of substituted salicylic 0022-3654/89/2093- 169 1$01.50/0
acids.' The unprotonated form of 3,5-dinitrosalicylic acid (DNSA) reacts with Ni2' with a rate constant ( k , = 3.1 X lo4 0 1989 American Chemical Society
