NOTES
203 1
or
Decomposition Mechanism of Xanthate in Acid Solution as Determined by a Spectrophotometric Method by I. Iwasaki and S. R. B. Cooke
upon integration
School of Mineral and .Metallurgzcal Engeneering, Universaty of Mannesota, Minneapolis, Minnesota (Received February 21, 1964)
log
(a log (optical density), - _ _( t z - tl)
~
=
7
(optical density),
((32
I n a previous article1 the decomposition of ethyl xanthate (C,HbOCSS-) in acid solutions was studied spectrophotometrically, and it was concluded that the xanthate and xanthic acid were virtually in equilibrium and that the monomolecular decomposition of xanthic acid into carbon disulfide and alcohol was rate determining. Klein, Rosarge, and Sorman, working with highly acid solutions, postulated the existence of an ion-pair activated complex and protonated xanthic acid along with xanthate and xanthic acid to account for the decrease in the decomposition rate in highly acid solutions. It will be shown in the present article that our previous approach to the investigation of the decomposition mechanism, when modified by assuming a stable protonated xanthic acid (eq. I) according to a proposal made by the above authors, will account for the observed rate data; and that the dissociation constants of xanthic acid (HX), and of protonated xanthic acid (HzX+), and the monomolecular decomposition rate constant can be determined as follows. 27%
K:. H+ + X- ; HX + H+ z? H2X+ (1)
J. ICs
ROH
+ CSz'
Since the optical density at a given wave length measures the total concentration of xanthate, xanthic acid, and the protonated xanthic acid, the decomposition of xanthate may be expressed by d(X-) dt
+ d(H,X+) + d(HX) dt dt __-
d(X-
-
+ HX: + H,X+) dt
=
-JGa(HX)
(2)
By using the dissociation constants of the xanthic acid and of the protonated xanthic acid, eq. 2 can be rewritten in terms of the total concentration C.
2.303
(4)
Therefore, the following approximation can now be made.
(A) When (H+) > K 1
Kz Condition (A) indicates that, when log K is plotted against pH, a straight line with a slope of -1 should result. This was demonstrated in the previous article,l from which ka/K1 was determined to be 216 at 23.5'. The value of k , can now be determined from condition (B) thus permitting calculation of Kl and K2.
Experimental and Discussion Potassium ethyl xanthate used in this experiment was prepared in the usual m a n ~ i e r . ~The optical measurements were made with a Beckman DU quartz spectrophotometer a t the characteristic absorption peak occurring a t 270 mp in highly acid solutions.2 The concentration of xanthate solution used in the experiment was normally 1.66 X M , and that of the hydrochloric acid ranged from 0.940 to 10.00 ill. In the stated acid range the decomposition of xanthate is extremely rapid; hence, a rapid mixing apparatus previously described' was used for introducing solutions to the cuvette. The decomposition rate was followed by employing a AZoseley ;\lode1 3 X-Y recorder driven by a Beckman energy recording adapter. A marked increase in temperature of the resulting solution was noted, particularly a t high acidities. Therefore, an alumel-chrome1 thermocouple was in(1) I Iwasaki and S R B. Cooke, J Am Chem Soc , 8 0 , 285 (1958) ( 2 ) E Klein, J K Bosarge, and I Norman, J Phys Chem , 64, 1666 (1960) 3) L S Foster, Utah Eng E x p St. T P , 5 (1929)
Volume 68, Number 7
Julu, 1964
NOTES
2032
serted directly into the cuvette; and the temperature of the mixed solution was measured with a Leeds aiid Xorthrup Mode1 8662 potentiometer. Triplicate tests were made a t three different temperatures; the xanthate aiid hydrochloric acid solutions and the cuvette were either warmed or cooled prior to the run. A linear relationship was obtained upon plotting the logarithm of optical density against time; and thus, the value of K was determined according to eq. 4 for a given condition. Figure 1 presents these values of K as a function of both acidity and temperature. It is evident that K is strongly dependent on the hydrochloric acid concentration and on temperature and that the reproducibility of data for a given acid concentration is satisfactory. The activation energy for the reaction was calculated for each acid concentration, and the average mas determined to be 18.0 + 1.1kcal.,'mole. Table I has been prepared by interpolating the K values from Fig. 1 a t 20'. It was noted that, when 1/~ was plotted against activity of H', a straight line was obtained as shown in Fig. 2, in good agreement with eq. 6. From the slope and the intercept of this line ka was calculated to be 3.41 + 0.01 min.-' and Kz 7.68 + 0.26. The deviation of the experimental points a t an acid concentration of 0.47 M from the
Table I : Experimental Rate Constant ( K ) and Molar Extinction Coefficient ( e ) at 270 mp as a Function of Hydrochloric Acid Concentration a t 20'
on+
K
0.47 1.01 1.99 3.01 4.02 5.00
0.360 0.847 2.16 4.54 8.77 16.2
3.75 3.27 2.61 2.08 1.58 1.10
o.oo
HCI CONCENTRATION
2
4
6
8.78 3.37 0.94 0.29 0.09 0.03
85.58 84.64 77.45 70.11 63.89 58.76
IO
8
z at
270 uw
5.64 11.99 21.61 29.60 36.02 41.21
12
9820 10060 10540 10630 10760 10820
14
16
0"
Figure 2. 10.0
-Relative concn., %-XHX HgXt
HC1 concn.. M
1.',
as a function of aa+ a t different temperatures.
I
t '
1
L---+ 325
3 30
335
VT
340
345
x lo3
Figure 3. K,, K?, and k~ as a function of l / T : (1) KI X lo2 after ref. 2 ; ( 2 ) Kl X l o 2 after ref. 1, corrected; ( 3 ) K1 x 102 after ref. 4.
0.3
I/T
x 103
Figure 1. K as a function of 1 / T a t different hydrochloric acid concentrations.
The Journal of Physical Chemistry
straight line may be attributed to the approximation given in condition (B). Similarly, these constants,
NOTES
K z and k3, may be calculated for different temperatures. The results thus obtained are presented in Fig. 3, together with the extrapolated values of K , given by Klein, Bosarge, and Norman2 and by the present authors. Table I also includes the relative percentages of the three species present a t 20' calculated on the basis of the dissociation constants given in Fig. 2. The molar extinction coefficients ( E ) at zero time have been obtained by extrapolating the plot showing the relationships between the logarithm of the optical density and time. Since E was also found to be strongly teniperature dependent, the values given in Table I for 20' and for the respective acidities were obtained by graphical interpolation. When the values of E are compared with the relative percentages of the three species, the molar extinction coefficients of xanthic acid and of protonated xanthic acid at 270 mp are 10, 960, and 9, 750, respectively. KO other absorption peak was observed in the wave length range 210-325 mp in 5 N hydrochloric acid solutions.
2033
Experimental Materials. A sample of dimethylbicyclobutane prepared by the cuprous chloride catalyzed reaction of diazoniethane with dimethylacetylene was furnished through the courtesy of Dr. W. Doering and Dr. J. F. Coburn; it was of 99.7% minimum purity and was used without further treatment. Matheson prepurified nitrogen was used in the runs with added inert gas. Apparatus and Procedures. The vacuum line, air bath thermostat with 500-nil. reaction vessel, teniperature measurement and control, and sample handling procedures were the same as employed in a previously reported study of another system.2 A 6-m. g.1.p.c. column packed with Dow Corning 710 silicone oil on Gas Chrom CL support (Applied Science Laboratories) was operated at room temperature and afforded good product mixture separation. An aluminum column was used with a stainless steel g.1.p.c. detector block. The starting material seemed to decompose and disappear, probably through polynierization, when it was passed through a firebrick packing at 35" in another chromatograph containing copper fittings.
(4) I. Iwasaki and S. R. B. Cooke, J . Phys. Chem., 63, 1321 (1959).
The Kinetics o f the Thermal Ionization of 1,3-Dimethylbicyclo[l.l.0]butance
by John P. Chesick Department of Chemistry, Haverford College, Haverford, Pennsylvania (Receized February 24, 1964)
Coburn and Doering' have synthesized l13-dimethylbicyclo[l.l.O]butane and have studied its reactions. This compound undergoes a thermal isomerization to give 2,3-dimethyl-lJ3-butadieneas the product. This must arise from the simultaneous or successive opening of the two cyclopropane rings without hydrogen niigration. Neither 1,3-dimethylcyclobutene, which would result from a liydrogen migration and bridgehead carbon-carbon bond rupture, nor 2-methyl-1,3-pentadiene, the expected thermal decomposition product of the dimethylcyclobutene, was observed in any significant quantity. The following study was therefore undertaken to determine the activation energy for the homogeneous decomposition reaction of this unusual bicyclic ring system.
Results and Discussion Rate constants were determined a t six temperatures spaced throughout the range 249.6-297.8". The bulk of the runs were a t pressures between 2 and 4 nim. At 237.6" the reactant pressure was varied from 0.2 to 15 mni. with no significant variation in rate constants. Reaction times were varied to give conversions corresponding to one-half, one, and two half-lives a t each temperature, with runs up to three half-lives a t the highest two temperatures. The reaction appeared to be satisfactorily first order. A 13: 1 nitrogen-reactant mixture was prepared and runs at a total pressure between 3.2 and 119 mm. a t 257.6" gave the same rate constants as runs with no inert gas. It is concluded that diffusion to the walls is not rate controlling in any significant part of the reaction at this temperature. Two additional product peaks were observed which eluted soon after the dimethylbutadiene in the g.1.p.c. analyses. These were not identified. They anlounted to less than 1% of the dimethylbutadiene a t the higher temperatures, about 1.3% a t 278.8", and up to 5.4% of the principal product at 249.6'. These peaks did not change significantly relative to the diniethylbutadiene with increase in reaction time, although they (1) J. Coburn and W von E Doering, private communication; J . Coburn, Ph.D. Dmertation, Yale University, 1963. (2) K . M. Klump and J P Chesick. J. Am. Chem. S O C ,85, 130 (1963).
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