1666
E. KLEIN,J. K. BOSARGE AND I. NORMAN
Vol. 64
SPECTROPHOTOMETRIC DETERMINATION OF FAST XANTHATE DECOMPOSITION KINETICS BY ELIASKLEIN,J. KEITHBOSARGE AND IRWIN NORMAN Research Department, Courtaulds, Inc., Mobile, Alabama Received April 28, 1960
The decomposit? kinetics of potassium ethyl xanthate in 0.02 to 4.0 N hydrochloric acid solutions have been detvmined at 0 and 10 . The concentrations of the xanthate ion and xanthic acid species were determined simultaneously b.y monitoring their respective absorbancies a t 302 and 270 mp. From the resulting data, the rate constants, thermodynamic ionization constant, molar extinction coefficients] and the Hammett h-function were determined. This is the first report of the aqueous xanthic acid extinction coefficients known to us. From the temperature dependence of the equilibrium and rate constants,. an estimate has been made of the heat of ionization of xanthic acid, and of the activation energy of tme reaction. The kinetics.are shown to be pseudo-first order under the conditions employed, the rate constant showing a maximum a t 0.4 N HCI. It is shown that alternate mechanisms can be postulated to account for this maximum. The resolution of these explanations lies in the interpretation of activity coefficients appearing in the rate expression. Several experiments at constant ionic strength are presented, and the mechanisms are discussed.
Introduction Since the discovery of cellulose xanthate in 1892 by Cross and Bevan, and the commercial application of this compound in 1900 in the viscose rayon process, the acid-catalyzed decomposition of the substituted dithiocarbonic acid ester has been of practical and theoretical interest. Studies of its model compounds, such as potassium ethyl xanthate, date back to the work of von Halban and Kirsch.'C2 A more modern approach to the mechanism of the reaction was utilized by Lewis3 who studied the non-aqueous kinetics of ethylxanthic acid decomposition. He demonstrated that in the absence of a base, the acid was stable, and postulated an ion-pair as the transition state intermediate. A spectrophotometric experimental approach to the problem, using acid normalities approaching 1.0, was first reported by Iwasaki and Cooke.* A more comprehensive study in both dilute and strong acid solutions was carried out by Ballard, Bamford, Gray and Totmans who showed that the rate constant passed through a maximum as the (H+) increased, and postulated that in strong acid solutions, a new equilibrium involving a protonated xanthic acid species became dominant. Experimental Four samples of potassium ethyl xanthate were used: two were Eastman White Label, repeatedly recrystallized and stored in a vacuum desiccator. The other two were laboratory preparations also recrystallized four times. The agreement in molar extinction coefficient was considered a better measure of punty than the standard iodine titration. The method suggested by Lewis3 was employed in the preparation of stable solutions of ethylxanthic acid. The choice of the solvent was dictated by the type of experiment tJo be performed. Spectral grade isooctane was used for analysis in the ultraviolet and CSZ for infrared analysis. A Gary Model 14 automatic recording spectrophotometer, with the cell carriage replaced by a metal block through which constant temperature water circulated, was employed for the kinetic measurements. The xanthate solution, generally 6.2 ,X M , WZLS introduced first by micropipet (0.100 ml.) into the standard cell, .and allowed to equilibrate thermally. The reaction was initiated by injecting 3.0 ml. of the acid solution as a (1) H. von Halban and A. Kirsch, 2. physik. Chem., 88, 325 (1913) (2) C. V. King and A. Chatenever, J . A m . Chem. Soc., 7 1 , 3587 (1949). (3) G. M. Lewis' Dissertation, New York University, 1947. (4) Iwasaki and Cooke. J . A m . Chen. SOC.,80, 285 (1958). (5) D. G. H. Ballsrd, C. H. Bamford, K. L. Gray and E. D. Totman, Courtsuds Limited, Maidenhead, England: Viscose Research Laboratory, Coventry-private communication.
jet into the cell, using an all-glass jacketed syringe capped with a polyethylene needle to avoid metallic contamination. The diverse problems of maintaining kothermal conditions were overcome in the following manner. The cell mount was cooled by circulating constant temperature water from a reservoir. The acid solution, together with the jacketed syringe, was pre-cooled in this reservoir. TO minimize heat transfer due to condensation on the cell faces, compartment was dried with desiccants and a flow of dry nitrogen. To avoid fogging of the cell window from residual moisture, the cell was coated initially with an alcoholic solution of an anionic detergent. The mode of operation was as follows: After the xanthate solution had reached thermal equilibrium (20 minutes afqer loading, as determined by a small thermistor bridge circuit) 3.00 ml. of the acid solution was injected from the synnge. At the same time, the recorder of the spectrophotometer n-as started. As soon as the compartment cover was closed, the pen and wave length sweep of the instrument were activated, and the absorption was scanned from 310 to 260 mp. At the lower wave length the scan direction was reversed, and this cycling procedure continued until the major absorption peak had an absorbance lower than 0.1. After a suitable waiting period, the entire spectrum again was scanned, and the base line values determined for correcting the instrumental zero. From a knowledge of the molar extinction coefficients of the undissociated acid and anion species, it is then possible to convert the absorbances a t 270 and 302 mp into molar concentrations. The sum of these concentrations suitably extrapolated to zero time should then yield the original xanthate concentration, unless some undetected species were also present.
Results The ultraviolet spectrum of xanthate ion (X-> has been reported for some time, but no information is available which characterizes the ultraviolet spectrum of the undissociated acid (HX). During the course of preparation of various samples of ethylxanthic acid and its thioesters for infrared analysis, for a related problem in this Laboratory, it was established that the undissociated acid had an ultraviolet absorption band at 270 mp (the position of the maximum and extinction coefficient being slightly affected by solvent effects). The assignment of the ultraviolet band was facilitated by comparison of a parallel examination in the infrared region. In Fig. 1 are shown the infrared spectra of potassium ethyl xanthate and of ethylxanthic acid-the latter isolated as previouslymentioned, as well as the S-ethyl ester of the latter. In the spectra of the xuithic. ac4d the intense
Nov., 1960
FAST XASTHATE
%DECOMPOSITION KINETICS
absorption region (8-10 p ) shows distinctive bands a t 8.08, 8.96 and 9.37-9.44 j . ~ compared to the xanthate salt (cf. 8.79, 8.91-9.09 and 9.51 p ) . If one assumes that the vibrations in this region arise from the R-0-C part of the molecule, then the frequency shift can be attributed to the effect cf interaction of the SH group on the 0-C bond.6 Figure 2 shows the molar extinction coefficients of ethyl xanthic acid, its anion and its S-ethyl ester in the ultraviolet region. The ionic form has a maximum at 302 mp and the undissociated acid at 268 mp in isooctane. The magnitude of the extinction coefficients is great enough to indicat'e that these bands are of the "K" type7 arising from the C=S chromophore and enhanced by either an S- or SH terminal auxochrome. This is supported by the bathochromic shift in the maximum due to the electron donating S- group, compared to t,he SH group, and by the bathochromic shift encountered with increasing dielectric constant of t'he solvent. A summary of relevent spectral data is shown in 'Z'a,ble I, supporting the conclusion that the 270 mp band is indeed due to the undissociated acid form. TABLE I S U ~ ~ M AOF R YULTRAVIOLET ABSORPTION DATAOF ETHYL? XANTHIC ACID AND RELATED COMPOUNDS e,
Species
Solvent
X. mg
1667
I
I
' 0
1 I I I
I 1 I I I I I I I I I I I 1 0 8 10 12 14 Wave length, p. Fig. 1.-Infrared abso tion spectra of: (a) potassium ethylxanthate; (b) e t h g a n t h i c acid; (c) ethylxanthic S-ethyl ester. I
I
I I
I I
4
2
1
I I I
6
1. mole-' om.-'
CzHjOCS2K
Water 302 17,420 Water 270 660 C2HsOCSzH Isooctane 268 9,772 Isooctane 302 117 C2HsOCS2C2H, Isooctane 278 11,200 C2HsOCSzH Water 270 10,670 Water 302 117b R-OCSzCH2COOCZH~~ 2-Chloroethanol 278 8,600' R-OCS2CH?CON(C2Hj)p5 2-Chloroethanol 280 960d C Z H ~ O C S G H ~Ligroin 280 . . . .B C2HsOCSzH Ligroin 267 approx. ... .e a R = cellulose. bTaken from isooctane data. J. Schurtz, Microchim. Acta, 589 (1955). J. Schurtz, Das Papier, 9, 333 (1955). "9. Hantasch and C. Scharb, Ber., 3570(1913).
Utilizing the extinction coefficients of Table I, it is possible to determine molar concentrations of both the acid and anion species by the standard analyses of two-component mixtures. The resulting equations are L,(HX) X = 0.937Ammp - 0.0356A4mmp ( l a ) Lx(X-1 X = 0.57OAmmp - 0.063A27omp (lb) (L= cell length in cm.) (-4 = absorbance)
By the addition of the values obtained from equations l a and lb, we obtain the total xanthate concent'rat'ion (XT),and from the ratio of (lb) t o (la), a function which is related t o t,he ionizat,ioii constant of xanthic acid. X semi-logarithniic plot of (X-), (HX) or (XT) as a function of time is found to be linear for over 50% of the reaction. As is t o be expected, the ( 6 ) F. G. Pearson and R . R. Stasiak, J . d p p l i e d Spectroscopy, 12, 116 (1958). (7) -4. E. Gillam and Z.S.Stern, "Electronic .kbsorption Sgactroscopy." E. Arnold Publishing Ltd., London, 1954, Chap. 8.
2.6
, 210
I
250
I
I
I
I
I
300
Wave length, mp. Fig. 2.-Log cxtinction us. wave length: ( A ) pot.assium ethylxanthate; (B) ethylxanthic acid; (C) ethyl ester of ethylxanthie acid.
slopes of the plots of each species are identical for any given acid concentration, indicating that the acid/anion equilibrium is fast in comparison to the speed of the decomposition reaction. From the extrapolated value of (XT)to zero time, it is also shown that the stoichiometry of the system is established in terms of only the anion and acid concentrations.
E. KLEIN,J. K. BOSARGE AND I. NORMAN
1668
CALCULATED IONIZATION CONSTANTS AND
-100HCl,
X L
N HX f** 0.0199 1.695 f0.014 0.769 .0995 ,401 f ,002 .645 ,298 .I435 f .0018 .585 .496 .0861 f .0003 .596 .696 .0561 f .0006 ,626 .884 ,0426f .0060 .669 1.19 .0293 f .0002 .758
T’ol. 64
TABLE I1 HAMMETTh-FUNGTIONS FOR ETHYLXANTHIC ACID A S CHLORIC ACIDCONCENTRATION
-
A
FUNCTION OF HYDRO-
-loo
with [Cl-] = 0.9-M XXL HX .f*’ Ki hKC/*’ hHY 0.0259 0.0152 2.13 f 0.02 0.672 0.0285 0,0121 1.537 ;t0.007 ,0257 .0643 ,433 f ,003 .672 .0290 .0596 ,366 f .003 .130 f .004 .142 f .005 .672 .0284 .182 .0250 .I80 ,0717i ,0011 .OS15 .0004 ,672 ,0271 .317 .0254 .300 .0478 f .0020 .0245 .460 .0331 f ,0006 ,0410i ,0005 .672 ,0243 .632 .0252 .586 ,0193 ;t .0016 ,0266 .880 Ki
*
00-
-
hKi fa’ 0.877 0.0235 0.0154 0235 .Of342 .803 .765 ,0227 .181 .776 .0215 .328 .492 .798 ,0212 .824 ,0199 .710 .877 .0177 1.218
TABLE 111 PSEUDO-FIRST ORDERRATECONSTANTS AND ACTIVATION ENERGIES FOR THE DECOMPOSITION OF ETHYLXANTHIC ACIDAS A FUNCTION OF ACIDSTRENQTH HCl. N
k, loo,see.-’ (C”) = O.9M
0.0199 ,0995 .298 .496 ,696 .884 1.19 1.59 1.99 4.00 a
0.00755 .0171 .0196 .0203 .0204
kexp
1O0,
kcOD
oo?
880. -1
sec. -1
0.00962 .ON1 .0208 .0213 .0211 .0205 .0190 .0176 .0160 .00907
0.00348 .00627 .00700 .00720 .00688 ,00653 ,00585 .00603 .00502 ,00295
Total xanthate added: 2.00 X 10-4 M .
(XT)Os
0,”
extrapd.
2.00 x 1004 2.02 2.00 2.03 2.00 2.00 1.95 1.99 2.05 2.09
k,,
G T
Ea (cal./mole)
2.76 2.88 2.97 2.96 3.04 3.13 3.24 2.91 3.20 3.07
15,500 16,200 16,700 16 ,700 17 ,200 17,500 18,000 16,400 17,900 17,200
3.02 Av.
16,900 Av.
10“
-
.
In Table 111are also shown the ratio of the rate constants for the two temperatures, and an estimate of the activation energy derived therefrom. The average of the ratio is seen to be 3.02, leading to an estimate of 16.9 kcal./mole. The experimental ratio (X-)/(HX) can be treated in several ways. We have preferred to use a combination of classical thermodynamics, and the Hammett acidity functions h-. The equations
where 1.o
2.0 3.0 4.0 HCl ( N ) . Fig. :%.-Experimental pseudo-first order rate constant for the decomposition of ethylxanthic acid as a functfi of HC1 normality: A, 10”; o, loo, [Cl-] = 0.9 Af; 0,0 .
The evaluation of the “pseudo” first-order rate constants was carried out from the slope of log XT vs. time graphs. The XT values were used once it had been established that -dXT/dt = -d(HX)/dt = -d(X-)/dt, and -dXT/dt a XT. A summary of the calculated experimental rate constants is given in Table I11 and Fig. 3 for the two temperatures employed. In addition, data are shown at 10” for experiments at a constant (Cl-) of 0.9 M , obtained by adding appropriate quantities of NaCl to aqueous hydrochloric acid solution of 0.02 to 0.90 N . Because of the small salt effect on the rate (less than lo%), an additional experiment at 0.1 N HC1 and 2.0 M NaCl was performed. The decrease in rate again was less than 10%.
define tlie relationship betwcen the IIainmelt function and the classical ionization constant, K i . Since the experimental data are at 0.02 and higher ionic strengths, the preliminary assumption was made that the mean activity coefficient of the HC1 solutions, f*, was equal to the mean activity coefficient of HX, i.e., (f*)2HC1 = f ~ . f x - and , that fax = 1. Using Harned and Owen9 data for (f&)2HC1, the K i values of ethylxanthic acid were calculated, and plotted against acid normality. The data are shown in Table 11. Extrapolation of the 0 and 10’ data to zero acid normality gave Xi values of 0.0235 and 0.0258, respectively. The extrapolations show that the original assumptions were in error by a second-order term, the latter being (8) L. P. Hemmett, “Physical Organic Chemistry,” hlcGraw-Ilill Book Go.. New York, N. Y..1940. (9) Harned and Owen, “The Physical Chemistry of Electrolytic Bolutions,” Reinhold Publ. Corp., New York. N. Y., Chap. 11, 1950.
FASTXANTHATE DECOMPOSITION KINETICS
Nov., 1960
inversely related to temperature. The Ki values for a constant ionic strength of 0.9, arising from mixed NaCl/HCl solute concentrations show an activity coefficient non-linearity typical of mixed salt solutions.lo Having extrapolated a value for Ki a t the two temperatures, we can then calculate h- as a function of (H+); these data are shown in Table 11. The temperature dependence of Ki also leads to an estimate of the thermodynamic values for the reaction (HX)
(H+)
+ (X-)
AH = 1400 cal. mole-’ Ah’ = -2.2 e.u.
Discussion Any explanation of the mechanism of potassium ethylxanthate decomposition in HC1 solution must explain both the observed pseudo-first order dependence on total xanthate concentration, and the maximum in the rate constant vs. acid concentration curve. The explanations postulated in the literature involve direct decomposition via ethylxanthic acid1s2.* in equilibrium with the xanthate ion. However, rate equations derived from this postulate cannot explain the maximum in the experimental rate constant curve. Two basic alternate mechanisms can be devised which are consistent with present knowledge. In 1954, Ballard, Bamford, Gray and Totmans postulated equilibria and reaction schemes, based in part on their own experiments and in part on the work of Lewis. a Ki
( X - ) + (H+)
I r (HX)
(2)
For comparison, the experimental and calculated values, based on the above solution, are shown in Table IV. While the agreement is excellent, the numerical values lead to improbable predictions. For example, combining the values of (4) with the material balance of (2)) one arrives at an expression relating the total xanthate concentration to xanthic acid concentration, all measured at zero time (XT) = (HX)[1.059
where (X-) (HX)
= concn. of C2HkOCSa-, xanthate anion = concn. of C2H60CSzH,xanthic acid
+
(HX*) = concn. of CnH60CSz; ion-pair activated comH plex (HtX)+ = concn. of CnHsOC&H, protonated xanthic acid H+
and Kl, Kz and Ka the appropriate equilibrium constants. The protonated xanthic acid was introduced to account for the maximum in the rate curve. The relationship between kexp and k is given by
HCI,N 0.0199 .0995 .2985 .4975 .6965
kD, sec.-l 0.00348 .00627
.00700 .OO720 .00688 .00653 .(MI585
+
= 45.06; K ~ K I 16.26 K2 = 2.51; Ki = 6.48
kKz = 0.3845; K1 (10)
Ref. 9, Chapter 14.
K2
(4)
k,
(calcd.), sec.-1
0.00348 .00623 .00717 .00720 .00698 .00662 .00585
It should be noted that, assuming that the protonated acid is in equilibrium with the undissociated acid rather than the ion-pair as shown the derivation results in an expression mathematically equivalent to equation 3. The second mechanism to be considered is based on a generalized reaction scheme using the basic postulates of the transition state theory. One assumes here that, in addition to the experimentally verified equilibrium involving the xanthic acid and its anion, there is another equilibrium between the reactants (X-) and (H+), and some activated complex of indeterminate nature (M), the latter decomposing to form the products, CzHsOHand CS2. The system is described by
(3)
and differentiation of this equation with respect to (H+) indicates that the maximum should occur when (H+)z = l/KZKa. Since neither of these constants is tractable experimentally, they must be calculated. A good fit to the experimental data is obtained from a solution having these values for the constants.
(5)
TABLEIV COMPABISON OF CALCULATED AND EXPERIMENTAL SOLUTIONS TO EQUATION 6
.w
+ CSa
+ 6.48(H+)+ 0.0235/(Hf)]
For acid normalities of 1.0, this would predict that only a fraction (13%) of the total xanthate exists in the xanthic acid form. Yet experimentally, the extrapolation data (Table 111) show that the xanthate concentration is accounted for on the basis of the sum of equation la and lb. The material balance could only be reconciled if the rather tenuous assumption were made that the ion-pair, the protonated acid, and the undissociated xanthic acid all had identical spectra. However, even if this assumption were accepted, the magnitude of the equilibrium constant, K z , (2.51) would predict that the majority of the xanthate in dilute to intermediate acid levels were in the form of a stable ion-pair, which again leads to an improbability.
1.194 ROH
1669
(X-)
K1 + W+)I r(HX) (M)
+Products
where The rate expression is given by RT (M) = k’(X-)(H+)’rate = -d(XT)/dt =
fu
(61
E. KLEIN,J. K. BOSARGE A N D I. KOHMAN
1670
Vol. 64
mechanism on this basis therefore reduce to the question of what must the structure of (M) be to predict that the activity ratio coeEcient should be independent of ionic strength. One can begin by assuming that (M) is identical to (HX). There are several objections t o this position, beginning with the observation that the AH value for ionization and the activation energy, Ea,differs by an order of magnitude in contrast to the expectation that they would be closely related if the initial hypothesis were valid. If, however, it is further assumed that ( A I ) is a collision energized form of (HX), the contrary argument must take the following rationale: the activation of (HX) to (M) would not be expected to result in a change in electrolyte character and, consequently, ~ H X= fM, which when combined with equation 7, leads to
0.020
0.015
&I e .
2 0.010
0.005
kex, =
k'Kih-
K m
18)
The right side of equation 8 does not show a maximum as a function of (Hf), and thereby contradicts the experimental facts. Considerations of several other mechanistic I possibilities using the criteria of pseudo-first 0.5 1.0 1.5 order kinetics, and maximum in the rate constant HCl ( N ) . curve, have led to the tentative conclusion that the Fig. 4.-Plot of - _ _ ~kXT-krKi &-&+us. HC1 most probable form of (M) is that postulated by H+((h- + K i ) fM Lewis,3 the ion-pair complex normality: 0,10 ; o, lo", [Cl-] = 0.9 M ; A, 0'. I
Introduction of the materia1 balance (XT) = (X-) (HX), integration, and equating with the experimental rate constant, ICexp, leads t o
+
(7)
The right side of equation 7 shows a maximum when plotted against (H+), without requiring the postulate of a protonated acid species; it results from the divergence of (H+) and h- values at high acid concentration. In Fig. 4,the ratio kexp/(H+)/hKi is plotted against N H C ~ ;the linearity is found to be good at NHCI > 0.1 and remarkably independent of the acid strength. One concludes that the neglect of the activity coefficient ratio does not lead to gross error at the higher ionic strength and, consequently, this ratio is not highly sensitive to the latter (cf. log f~ = log fx - log f~ t). The criteria for the
+
+
S
H H I /
//
I 1
\
H-4-C-4-C
SThis hypothesis demands that the equation log f ion pair = log fx- + log f ~ + (9) be substantiated by theory, as it is in experimental fact. A more obvious proof of the mechanism must await isotope studies to elucidate the exact nature of the transition complex. Acknowledgment.-We wish to express our appreciation to Drs. Bamford, Ballard, Gray and Totman for making their work available to us, and for informative discussions and correspondence. We also wish to thank Mr. John Wharton, Director of Research, and members of the Board of Courtaulds, Inc., for permission to publish. H H Hf
