THOMAS I. CROWELL AND MARIE G. HANKINS
1380
tions, will be characterized by enthalpies of propagation much lower than that for sulfur on account of the lower bond energy of a sulfur-selenium or selenium-selenium linkage. Furthermore, the abundance of "mixed" rings will be relatively large even for small additions of selenium. The effect of these factors on the variation of T+can be demonstrated by considering a TobolskyOwens type of copolymerization involving S S and Sese2 monomer species with the latter characterized by the same entropy of propagation as Ss but with an enthalpy of propagation the same as that for Se8. As indicated in Figure 6, the composition dependence of T+ calculated in this way is in closer agreement with experimental observation. The consumption of S S rings in such a copolymerization should be detectable spectroscopically wit,h sufficiently sensitive techniques. I n the Raman experiment performed here the threshold for removal of Ss rings from the system was obscured by the high optical density of the melt a t the temperatures of interest. Note on Resonance Enhancement of Raman Scattering. The phenomenon of resonance enhancement of Raman scattering occurs when the energy of the exciting radi-
ation approaches the energy of a fundamental electronic transition of the scattering species. The close agreement between the results of the two types of Raman scattering experiment described here implies that the resonance enhancement phenomenon does not play a significant role in this case. Either the principal contribution to the optical absorption must occur in a species which does not contribute greatly to the Raman scattering of interest or the absorbing center must be localized within the vibrating species in such a way as to produce minimal coupling between the electronic and vibrational transitions. The latter situation probably holds in the case of the polymeric chains S, for which the electronic transition of interest is probably associated with the presence of an unpaired electron localized a t each of the chain ends while the vibrational transitions of interest are associated with the normal modes of the entire molecule. Acknowledgment. The authors wish to acknowledge the benefit of helpful discussions with D. Olechna and R. C. Keezer and the assistance of J. O'Neill in the experimental work.
The Hydrolysis of Thiocyanic Acid. I. Dependence of Rate on Acidity Function by Thomas I. Crowell and Marie G. Hankins Department of Chemistry, UnWersitg of Virginia, Charlottesville, Virginia
28901
(Received September 3 0 , 1 9 6 8 )
The rate of formation of ammonia, hydrogen sulfide, and carbon dioxide from potassium thiocyanate, in 0.0511.0 M hydrochloric, phosphoric, and sulfuric acids, measured at 13.5-90.1", shows a linear correlation of log k with Ho. The slope is 1.0 at high acidities where the predominant species is HNCS and about 1.4 below Ho = -2. The pK* of thiocyanic acid is -2.3 to -2.0 at 25", depending on the estimated H- of hydrochloric acid solutions. There is a solvent isotope effect on the absorption spectrum of thiocyanate ion in DgO. The hydrolysis rates are correlated by log IC = [h-/(K* -k h-)] (IC1 k h ) , an equation consistent with concurrent mechanisms involving thiocyanic acid and one of its conjugate acids.
+
Thiocyanic acid is a volatile compound in which the molecular structure HN=C=S predominates over HSC=N.l It readily polymerizes, and forms a variety of heterocyclic compounds.2 Aqueous solutions are stable enough for measurements of the dissociation constant of thiocyanic acid,s which is quite strong, PKAbeing about -2. Hydrolysis occurs, however, in acid solution, according to the equation4
H'
+ HNCS + 2H20
--+
H2S
+ C02 +
4'"
(1)
Because i t is necessary to know the rate of this reaction if the dissociation constant of the acid is to be determined accurately and if any reactions of thiocyaThe Journal of Physical Chemistry
nate ion are to be quantitatively investigated at low pH, we have undertaken a kinetic study of reaction 1.
Results and Discussion Most of the rates reported at 90.1"were measured by the volume of carbon dioxide evolved. Several runs (1) 0. I. Beard and B. P. Dailey, J . Chem. Phys., 18, 1437 (1950). (2) U. Rtick and H . Steinmetz, 2. Anorg. Chem., 7 7 , 51 (1912); W. A. Sherman, "Heterocyclic Compounds," Vol. 7. R . (3. Elderfleld, Ed., John Wiley & Sons, Inc., New York, N. Y., 1961, p 569. (3) (a) T. D. B . Morgan, (3. Stedman, and P. A. E. Whincup. J . Chem. SOC.,4813 (1965); (b) 9. Tribalat and J. M. Caldero, BulZ. SOC.Chim. Fr., 774 (1966). (4) C. J. Hansen, French Patent 661,508, Oct 5, 1928; Chem. Abstr., 24, 473' (1930).
HYDROLYSIS OF THIOCYANIC ACID
1381
were followed by titration for ammonia. First-order plots tapered off with time, often around 30% reaction, probably because thiocyanic acid was lost by vaporization. The rate constants, calculated from the straight, initial portion of the plots, were reproducible only to about &20% but were observed over the acidity range 0.05-6.9 M acid, corresponding to a 3400-fold increase in rate. The initial thiocyanate concentration was 0.192 M . Ultraviolet spectrophotometry, at first avoided because of yellow by-products observed a t high thiocyanate concentrations, was found to give excellent results at low concentrations. The spectrum of thiocyanate ion is remarkably temperature dependent due to charge-transfer complexing with the solvent.6 We found further evidence for the complex in that the spectrum of KSCN is shifted about 1.5 mp toward shorter wavelengths in DzO;a t 240 mp, the absorptivity in H 2 0 is 1.32 times that in D20 between 25 and 60". This solvent isotope effect is comparable with the observed shift for hypophosphite6a and iodide.eb Reactions followed spectrophotometrically with initial thiocyanate 6.88 X M gave pseudo-first-order kinetics, since the strong acid was present in great excess. First-order plots were generally linear to 50% reaction (30 min to 5 months), The rate constants are given in Table I. The reproducibility of the data is shown by the following results: 1.94 M HCl, 10% (sec-l) = 4.89, 4.91; 5.97 M , lO?k = 5.77, 5.82; 6.91 M , 10% = (1,17), 1.41, 1.42; 7.97 M , 10% = (3.01), 3.18, 3.23, 3.27, 3.30; 9.50 M , 106k = 9.45,9.55. If the two values in parentheses are omitted, the average deviation from the means is 0.8%. In the last set of data, the rate constants 3.01 and 3.30 were obtained using 1.37 X M thiocyanate. The invariance of IC with initial thiocyanate concentration shows that the reaction is first order in thiocyanate ion or thiocyanic acid and rules out terms such as [HNCS]2 in the rate Table I: Rate of Thiocyanic Acid Hydrolysis" at 25"
1.94 2.91 3.48 3.82 4.65 5.08 5.42 5.97 6.91 7.97 9.50 10.87 10.99 11.11 6
-0.65 -1.01 -1.21 -1.37 -1.64 -1.79 -1.91 -2.11 -2.46 -2.85 -3.41 -3.91 -3.96 -4.00
-0.52 -0.99 -1.21 -1 -46 -1.81 -2.01 -2.18 -2.50 -3.01 -3.50 -3.96
Initial thiocyanate concentration 6.88 X Calculated as described in text.
4.90 X 1.98 x 4.08 x 6.68 X 1.58 X 2.50 X 3.48 X 5.80 X 1.40 X 3.25 X 9.50 X 3,92 x 4.15 x 4.48 x M.
10-8 10-7 10-7 10-7 10-6
10-6 10-6 10-4 10-4 10-4
Reference 7.
-2
-3
!i
O4
c
5 -5. -6
-7
-8'
-1
0
I
2
3
4
-Hoe Figure 1. Rate of hydrolysis of HNCS us. acidity function, Ho, at 25" (lower curve, calculated by eq 4) and 90.1" (upper curve). Method: 0, uv; 0,gas volume; 0, NHI titration. Medium: 0 , HCl; H.804; 0-,&POI.
6,
equation. Furthermore, there was good agreement between the rate constant at 90.1" and 1-00M hydrochloric acid obtained by this method (1.97 X sec-1) and by carbon dioxide evolution (1.88 X sec-l) ,although the thiocyanate concentrations differed by a factor of 300. The first-order rate constants, L, are plotted logarithmically against the acidity function, Ho, in Figure 1, for 25 and 90.1". Runs at 13.5, 24, and 40", all with 10.7 M hydrochloric acid ( H o = -3.86), yielded an activation energy of 18 kcal/mol and an apparent entropy of activation (from first-order rate constant a t 25") of 20 cal/deg mol. In order to understand the kinetic results, it was necessary to measure the dissociation constant of thiocyanic acid under the experimental conditions. From the optical densities of 6 X lOW4Mpotassium thiocyanate solutions containing hydrochloric acid of various concentrations, the fraction [HNCS]/[SCN-] was calculated, after extrapolation of the thiocyanic acid concentration to zero time if the hydrolysis rate was appreciable. Values of the acidity function, H-, and its negative antilogarithm, h-, were estimated as follows: from the known Ho for a given HC1 s ~ l u t i o n , ~ the concentration of H2S04 yielding the same Ho was obtained. Then the H- value of that H2S04 (5) E . Gusarsky and A. Treinin, J . Phys. Chem.. 69, 3176 (1965). (6) (a) H.Benderly and M. Halmann, ibid.. 71, 1053 (1967): (b) I. Burak and A. Treinin, Trans. Faraday Soc., 59, 1490 (1963);M. Halmann and I. Platzner. Proc. Chem. Soc.. 261 (1964). (7) M.A. Paul and F. A. Long, Chem. Rev., 57, 1 (1957).
Volume 76,Number 6 May 1989
1382
THOMAS I. CROWELL AND MARIE G. HANKINS
solution8 was assumed to approximate that of the actual HC1 solution. Using the expression, KA = h-[SCN-]/[HNCS], ~ K was A calculated to be -1.91 in 3.70M HC1 (Ho= -1.30, H- = -1.37) and -2.05in4.60MHC1 ( H o = -1.64,H- = -1.82). The average, -2.0 (KA= 100)) is used in this paper. If perchloric acid, instead of sulfuric, is used to relate H- to Ho,-2.31 is obtained. Tables of H- for perchloric acid do not extend to low enough acidities to correlate our results. When H - values are available for hydrochloric acid, a better estimate of PKAcan be made. A PKAvalue of -2.1 for thiocyanic acid in 2.4-5.6 M perchloric acid at 7 M ionic strength may be calculated from the spectrophotometric data of Morgan, Stedman, and W h i n ~ u p ,using ~ ~ H- tables* for perchloric acid which were not available to them. These authors used several methods but apparently did not take hydrolysis into account. A solvent extraction methodaa gave a PKA of -1.83. Tribalat and C a l d e r ~ using , ~ ~ polarography, found the dissociation constant, KO,to be 5.0 at 3 n/l ionic strength in 2.3 M perchloric acid. Since H- = -1.14 for perchloric acid of this molarity,8 neglecting the 0.7 M extra electrolyte, me calculate PKA = -1.5 from their data. The foregoing measurements show that thiocyanato ion is completely protonated at the higher acjdities used and chiefly unprotonated in the more dilute acids. This is probably true at 90.1" as well as at 25", because the ratio [HNCS]/[SCK-] in a given acid solution is practically independent of temperature from 10 to 40". Figure 1 shows that the hydrolysis of thiocyanic acid is strongly catalyzed by acid up to W O= -4.0, the limit of our measurements, which implies that after the protonation of thiocyanate ion, a second proton is involved in the mechanism. However, the lower part of the curves, particularly the 90.1" line extending to 0.05 M HC1, does not show a dependence of k on [Ha0+I2. We therefore postulate a parallel mechanism involving only the neutral HNCS molecule (or a kinetic equivalent) and water
ki
+ HzO products (2) HNCS + H30++ HzO products (3) Noting that the fraction of total thiocyanate which is protonated is ~ - / ( K + A h-), we obtain the rate HNCS
kz
equation (4) for reactions 2 and 3 k = [~-/(KA
+ h-)](ki + kzho)
(4)
The line in Figure 1, 26", is a plot of log k , calculated from eq 4, us. Ho, with k~ = 1.9 X sec-l and kz = 4.5 X lo-* M-' sec-1. While kl would be sensitive to errors in the measurement of [SCN-]/[HSCN] used to determine the dissociation constant of thiocyanic acid, it would not depend on the choice of The Journal of Physical Chemistry
I 9'
I
I
1
I
I
0 rn
0.8
' t
;" f i 0.6
4-
+ (0
0
-1.0
0
-2,o
-
3.0
H, -b log [H) Figure 2. Bunnett plots of rate of HNCS hydrolysis: A, uncorrected (C = 0); B, corrected for partial conversion to SCN- (C = 1).
reference acid, sulfuric or perchloric, used to determine H - and ~ K as A described above; the evaluation of ICz requires only a rough value of ~ K A .The applicability of eq 4 at acidities much lower than those at which PICA was determined ( H o near -1.4) of course requires the existence of an H- acidity function in these media, as assumed. Bunnett plotss of log k HOus. log ~ H rise ~ Osteeply with increasing acidity to a maximum of 0.69 for H o = -2.46 and then decrease to 0.65, except for a value of 0.57 at Ho = -3.41, a point which appears slightly low in Figure 1. The more recently suggested plotlo of log k HOvs. Ho log [H+] shows similar behavior (Figure 2, curve A ) . In view of the known dissociation of thiocyanic acid, a more rational application of the plots is to consider HNCS a weakly basic reactant in an acid-catalyzed reaction,l0 and also to correct its concentration for the decrease due to dissociation. The function log k - log [h-/(hK A ) ] Ho is then plotted against H o log [H+] (Figure 2, curve B). The cause of the nonlinear function which increases at low acidities is, according to our interpretation, the "uncatalyzed" reaction of the species HKCS, with rate constant kl. The slopes of all these plots seem to be approximately zero at high acidities where le1 is negligible in comparison with kzho. This is necessarily so if the slope of Figure 1 approaches 1. Little more information bearing on the hydrolysis mechanism is available. The fact that the reaction is faster in D20-DC1 ( k ~ / = k 0.28) ~ argues against ratecontrolling protonation. When ammonium thiocyanate
+
+
+
+
+
+
(8) R. H. Boyd, J. Amer. Chem. Soc., 83, 4288 (1961). (9) J. F.Bunnett, tbid., 83, 4958, 4968, 4973, 4978 (1961). (10) J. F. Bunnett. Can. J. Chem., 44, 1899, 1917 (1966).
HYDROLYSIS OF THIOCYANIC ACID is used instead of the potassium salt (5 M salt, 0.15 M HCl, 90.1"), the rate is almost unchanged. This is evidence against the mechanistic importance of thiourea or other ammonia derivatives. Nucleophilic attack of water on the carbon atom of the protonated imino group in CH2N - C S]+ would have seemed the obvious mechanism, but is contrary to the Zucker-Hammett hypothesis and the Bunnett slopes of zero. However, the ambident, unsaturated thiocyanic acid molecule suggests more than one plausible mechanism; for example, in these highly acidic solutions, it is possible that the HaNCS+ might be protonated once again to H3NCS2+ followed by hydrolysis to NH3 and COS. The last step would show H30+ dependence which, superimposed on the very much larger Ho dependence, would not be detectable.
Experimental Section Materials. Reagent grade concentrated acids and thiocyanate salts were used, except that constantboiling hydrochloric acid was prepared for the spectroscopic runs and found to have the same absorptivity as the commercial material a t 220-260 mp. Deuterium chloride (20% solution in deuterium oxide) was obtained from Diaprep, Inc. Procedure. The volume of carbon dioxide evolved was measured over mercury in a gas buret surrounded by a 25" water jacket. The vapors passed from the reaction vessel, a 25-ml round-bottomed flask in the oil bath, through a 25" reflux condenser and three U tubes successively containing Drierite, cupric sulfate on
1383 pumice, and Drierite. The gas reaching the buret was shown to contain neither sulfide nor thiocyanate. The reaction vessel was not shaken during the run; shaking was shown not to affect the rate constant. Reactions to be analyzed for ammonia were run in volumetric flasks, the carbon dioxide and hydrogen sulfide escaping into the atmosphere. The reaction mixture was made alkaline after the desired reaction time and the ammonia distilled off and titrated with standard acid. In kinetic runs followed by ultraviolet spectrophotometry, the initial thiocyanate concentration was iM though the acid was present in the only 6.88 X particular high concentration desired; stoppered flasks could, therefore, be used. The absorptivity of thiocyanate ion in aqueous solution is given by the following values of a t 240 mp: 2.5 (lo"), 3.3 (25"), 4.0 (40"), 5.0 (59"); a t 245 mp: 0.90 (lo"), 1.15 (25"), 1.38 (39"). The absorptivity of HNCS, 9.1 X lo2 at 240 mp (maximum) and 8.0 X lo2 a t 245 mp, was temperature independent. These last values were obtained by extrapolating the optical density of a solution of potassium thiocyanate in 10.7 M hydrochloric acid back to zero time as reaction 1 proceeded. Spectra were determined on a Beckman DU spectrophotometer with thermostated cell compartment and on a Hitachi recording spectrophotometer. The kinetic runs were followed at 240 mp. Acknowledgment. We are grateful for the support of the U. S. Army Research Office (Durham).
Volume 78,Number 6 May 1969
