JOURNAL OF THE AMERICAN CHEMICAL SOCIETY (Registered i n
U. S. Patent Office)
(0Copyrigbt, 1958, by the American
Chemical
Society)
FEBRUARY 12, 1958
VOLUME SO
h
T3
PHYSICAL AND INORGANIC CHEMISTRY [CONTRIBUTION FROM
THE
BIOCHEMICAL RESEARCHDIVISION,CHEMICAL WARFARE LABORATORIES]
Kinetics and Mechanism of Chlorination of Triethylphosphorothiolate in Dilute Aqueous Media at 25" BY NICHOLAS G. LORDIAND JOSEPH EPSTEIN RECEIVED JUNE 25, 1957 The kinetics and mechanism of chlorination of triethylphosphorothiolate (TEPT) in dilute aqueous solution were deterIn acid solutions (PH 4-$, the chlorination involves a reaction brtween molecular chlomined over a range of PH 4-10. rine and TEPT, the rate being controlled by the formation of chlorine from hypochlorous acid. hydronium and chloride ions. I n slightly acid and neutral solutions (pH 6-7), the kinetics indicate that, in addition t o a reaction with molecular chlorine, TEPT is chlorinated by a concerted attack of both hypochlorous acid and hypochlorite ion. The latter reaction is minimized where chloride ion concentration is high. In alkaline solution (PH 8-10), the participating moiety in the reaction is the hypochlorite ion. Kinetic equations are derived for the reactions involved. A discussion of the mechanisms involved is given.
Introduction
0
The discovery that hypochlorite ion behaved as a catalyst in the hydrolysis of isopropyl methylphosphonofluoridatel (Sarin) prompted us to study the reactions of chlorine in dilute aqueous solution with other organophosphorous compounds. The present report is concerned with the interactions between chlorine and triethylphosphorothiolate (TEPT). I n contrast to the single reaction observed between chlorine and Sarin, i.e., the hypochlorite ion catalyzed hydrolysis, the reaction between T E P T and chlorine is much more complex. At PH < 5, a very rapid reaction occurs between T E P T and molecular chlorine in which three moles of C12 is consumed and between four and five moles of acid is produced for each mole of reacting T E P T (Fig. 1). Of the total acid produced, three moles is formed rapidly, roughly paralleling the rate of decline in chlorine titer, the remainder of the acid forming a t a considerably lower speed and only after the uptake of chlorine had ceased. Ethane sulfonyl chloride (as its sulfonamide) was isolated from reaction mixtures in which only three to four moles of acid had formed. The reaction may be expressed as
E t 0 P-SEt
I1 I
rapid + 3C12 + 3Hz0 -+
0 Et
0
/I I
E t 0 P-011
+ FtSO2C1 + 5IICl
(a)
0 Et EtSOZCl
slow + HZO -+ EtSOaH + HC1
(b)
Although a total of eight moles of acid is produced in the reaction (both the sulfonic and phosphonic acids are strong acids), three moles is consumed in the formation of molecular chlorine from hypochlorous acid, leaving a maximum of five which can be measured. In the pH range 5-S two reactions appear to occur: one, an attack by molecular chlorine; the other, a concerted attack of both hypochlorous acid hypochlorite ion upon T E P T . At fiH > 8, the reactive species is hypochlorite ion, which apparently acts as a hydrolytic catalyst, similar to its action upon Sarin. However, the rate of the hypochlorite catalyzed hydrolysis of T E P T is much lower than the similar reaction with Sarin.2
(2) T h e bimolecular rate constant for t h e hypochlorite catalyzed (1) J. Epstein, V. E. Bauer, M. Saxe and A l . M. Demek, THIS hydrolysis of the phosphonofluoridate was found to be 600 1. mole-1 ruin.-' at 25'1; for TEPT, 2.4 1. mole-' min.-'. JOUXNAL, 78, 4068 (1956).
500
~
~
NICHOLAS G. LORDIAND
510
JOSEPH
EPSTEIN
Vol. 80
30 40 50 60 70 Time (minutes). Fig. 1.-Rates of acid formation ( I ) and chlorine destruction (11) a t PH 5, I.( = 0.1, T = 25 5 0.5’: 0 , [HOC110 = 2.38 X lO-‘M, [Cl-] = 1.16 X lo-* M ; [TEPT] = 4.6 X hl; 0, [HOClIo = 1.6 X lo-‘ M, [el-] = 1.04 X 10-3 M ; [TEPT] = 4.6 X 10-6 hl. 20
10
0
The products of hydrolysis consume the hypochlorite, i.e. 0
//
SEt
€?to-P-
I
slon + OC1- + H?O *
0
0
I’
Et
EtO--P--OCl
1
+ EtSH + O H -
(c)
0 Et 0
II I
EtO-P-OC1
+ HzO
fast
0 I! (EtO)*P-O-
0 Et
EtSH
+ X0C1-
fast
+ OC1- + 2H+
-+EtSOifI - t 4C1-
(dl (e)
These conclusions were reached from a kinetic study of the chlorination reaction a t constant pH over the range 4 to 10 and in the presence of varying amounts of chloride ion and TEPT. At any given pH (maintained constant by addition of alkali) and a t fixed concentrations of chloride ion and T E P T , the rate of disappearance of active halogen (chlorine, hypochlorous acid and/or hypochlorite ion) was found to be first order with respect to the total active halogen concentration, enabling the calculation of a first-order constant (kobs) . 3 Deviations from first-order kinetics were observed only after 50y0 of the reaction had occurred. This could be explained by the fact that the rate of disappearance of active halogen was also (3) First-order rate constants based upon acid production rates when evaluated by Guggenheim’s method (A. A. Frost and R . C,. P a r s o n , “Kinetics and Mechanism,’’ John Wiley and Sons, New York. S.Y , 1953, p. 48) were of the same order of magnitude as k , h .
511
CHLORINATION OF TRIETHYLPHOSPHOROTHIOLATE
Feb. 5, 1958
dependent in some measure upon the T E P T concentration (Fig. 2). At constant $H and TEPT concentration the kobs was directly proportional t o the chloride ion concentration (Fig. 3).
0.07
0.028
0.024
0.05
-
0.020
r: 0.016
3 8' 2
0.012
0.008
0.004
0
0 0.002
0.001
Added [Cl-] , moles/l. of the rate of chlorine destruction on 0 5 10 15 20 25 30 chloride ion concentration: (A) 0 , PH 5, [TEPT] = 4.5 X [TEPT],,, moles/l. X IO-*, 10-6 M , p = 0.1, T = 25 0.5'; ( B ) 0, PH 7, [TEPT] = Fig. 2.-Dependence of the rate of chlorine destruction on 4.5 X 10- M , p = 0.1, T = 25 f 0.5". the initial concentration of T E P T : 0 , pH 7.0, [el-] = 5 X lo-' M , [HOCl] [OCl-] = 5.5 X M , p = 0.1, and 5, there appears to be a ten-fold increase in T = 25.0 =!= 0.5"; 0,PH 4.0, [Cl-] = 9 X M , [HOClIp klobs for ten-fold increase in [H+]; above this pH, = 1.5 X IO-' M,p = 0.1, T = 25 j z 0.5'. the dependence of the rate on [H+] is less than ten
Fig.3.-Dependence
+
At p H 5 (Fig. 3) the straight line relating kobs to chloride ion concentration may be extrapolated t o the origin, suggesting that no reaction ensues in the absence of chloride ion a t this p H (this also implies that there is no reaction between T E P T and hypochlorous acid). Above pH 5, e.g., p H 7, a second reaction independent of chloride ion and resulting in loss of halogen appears t o be involved (intercept, Fig. 3, pH 7). The rate of halogen loss due to the chloride-independent reaction (which was much more rapid than the loss of chlorine in solutions without TEPT) was generally observed to occur in the p H range 5.5 to 8. Thus, the rate of disappearance of halogen could be arbitrarily divided into (a) a chloride ion concentration dependent reaction and (b) a chloride ion independent reaction. At p H < 5, the rate of halogen disappearance can be explained by assuming that only reaction (a) takes place; above PH 5, both reactions must be assumed to explain the observed rate. An analysis of the chloride ion concentration dependent reaction follows. The portion of the observed rate constant which is chloride ion concentration dependent (klobs = kobs - koobs, where koobs is the rate constant at zero [CI-] concentration) divided by the [CI-] is related to pH as shown in Fig. 4; between PH 4
and variable. If, now, the rate is assumed t o be dependent upon the hypochlorous acid concentration4 rather than the total active halogen concentration, the resulting value for k l o b s is found to be independent of the [H+] over the pH range of 6.5 to 8.0 (dotted line, Fig. 4). The observation that, a t constant initial concentration of TEPT, the rate of disappearance of hypochlorous acid in acid solution depended on the concentration of HOC1, chloride ion and hydronium ion raised t o the first power, and that the dependence on T E P T was small, approaching a constant value as [TEPTIowas increased, suggested that the rate controlling step in the over-all reaction of T E P T with active halogen (equation g) was contained in the equilibrium involved in the formation of chlorine (equation f). HOCl
+ HC1
ki Cln
k- 1
+ HzO
(f)
(4) The hypochlorous acid concentration was calculated from the equation HOCl = ( [ H + ] / ( [ H + ] -I- KA)) [HOCllt, where [HOCllt is the total active halogen concentration and K A is the acid dissociation constant of hypochlorous acid under experimental conditions of ionic strength. The value of K A used in this work, wiz.,8.45 X 10-8, was calculated from the equations log K A = log K 2 A d E where K = 4 X lo-, C 0.1 and A = 0.51 a t 25'. (5) S. Glasstone, "Textbook of Physical Chemistry," 2nd Ed., D. Van Nostrand Co., Inc., New York, N. Y., 1946,p. 970.
-
+
NICHOLAS G . LORDIAND JOSEPH EPSTEIN
Vol. SO
where k, = kl [H+] [OH-]. Also
1
+ k ~ and , k,
-- k-1 4- k - 2
[HOCl] - d___ = kl[HOCl][Cl-][H+] + k~[H0Cl][Cl-ldt k-t[Clt] - k2[C12][OH-] = kz[HOCl][Cl-] k,[CLI
(3)
Substituting the equivalence for Clz, equation 2 in equation 3, one obtains
- d[HOCl] -= k,ka[HOCl] [Cl-][TEPT] dt
ku
+ kt[TEPTl
(4)
If total active halogen is experimentally m e a ~ u r e d , ~ then
Transposing
- d[HOCl] --
kxkiK’[Cl-][TEPT] dt ky kr[TEPTl
+
[HOClIt
where
At constant pH, T E P T and C1- concentration^,^ it was experimentally observed that
Hence 5
4
7
6
8
PH. Fig. 4.-Dependence of rate of destruction of active chlorine by T E P T on [ H + ] ; [TEPT] = 5 X 10-6 M,T = 25 & 0.Bo, p = 0.1; 4 , observed values, ----O----, values corrected for ionization of hypochlorous acid.
0.005
0.05
0.004
0.04
A
OEt
where k3 >> kl and k-l. At PH’s between 6.5 and 8.0, where the rate appears to be independent of pH, the rate of chlorine formation may be controlled by a reaction of HOCl with C1- (equation e). This equilibrium has been suggested by Morris6 as the mechanism involved in the hydrolysis of chlorine. HOCl
+ C1-
kz
J _ Clz k-
+ OH-
(e) 0.001
t
If the steady-state approximation is applied to the formation of chlorine, then
dy:l
--
+
= 0 = k,[HOCl][H+][Cl-]
kp[HOCl][Cl-]
- k-i[Clz] - k-~[Clr][OH-l k3[ClzIITEPTl (1)
or [c121 =
+ kZ[HOCl][Cl-] + ks[TEPT]
ki[HOCl][Cl-][H+] k-l k--p[OH-]
+
(6) J. C. Morris, THISJOURNAL, 68, 1692 (1946).
t
1
0’01
0 0
2
4
6
8
10
l / [ T E P T l o X lo-‘. Fig. 5.-Relationship between the reciprocals of the firstorder rate constant (at constant [Cl-] ) and [TEPT] concentration at PH 5 (A, 0 )and PH 4 (B, 0 ) ; T = 25 f 0.5”, p = 0.1. ( 7 ) With respect t o C1-, this condition will prevail for all practica purposes where the added chloride is large as compared t o the initial hypochlorous acid concentration, even though a mole of C1- is formed for each mole of hypochlorous acid decomposed. During the initial periods of the reaction, the concentration of TEPT may be considered approximately constant.
Feb. 5, 1955
CHLORINATION OF TRIETHYLPHOSPHOROTHIOLATE
513
TABLE I ANALYSISOF CHLORIDE DEPENDENT RATEOF DESTRUCTION OF CHLORINE AS A FUNCTION OF PH AT T = 25 f 0.5',
1
= 0.1
Av.
fiH
4.0
4.3 4.7 5.6
5.5
6.0
6.5
7.0
[HOCIlo,
k'obs
[TEPTh X 18, M
X 106, M
l/Kj
0.98 1.96 4.6 4.9 4.6 4.6 51 4.5 4.6 2.1 9.2 10.6 51.0 4.8 5.3 8.5 42.5 5.3 5.4 20.0 20.8 27.0 49.5 5.4 27 49.5 4.8 49.5
14 15 15 15 4.6 4.5 5.8 3.5 8.4-23.8 14.5 8.4 14.6 5.6 4.8 5.7 7.7 7.3 5.9 8.0 13.3 12.2 7.5 9.0 7.3 7.2 6.7 4.5 6.6
1.0
1.0 1.0 1.008
1.026
1.OS5
1.266
1.845
where [TEPT] and [Cl-] are assumed to be equal to their initial value, [TEPTIo and [Cl-lo. Equation 5 may be rewritten
C I
216 & 7 325 f 8 457 459 f 11 218 98 127 & 8 57.6 f 2 . 2 60 f 2 . 5 43.5 70.1 71.0 82.1 30.2 30.7 f 0 . 8 30.0 f 1 . 9 38.1 f 0.6
xky'k108 '
kx exp.
kx calcd.
1.91
633
651
*.
312 131
335 145
1.56 1.87
83.5
81.7
1.19
37.6
38.4
18.1 f 1 19.6 f 0 . 1 22.5 f 0 . 4 23.5 f 3 . 7 23.6 f 1 . 4 27.0 f 3 . 3 14.4 f 1 . 2 18.6 f 1 . 2 18.2 f 0 . 1
1.97
25.8
24.7
1.78
19.2
20.4
12.5 17.4 f 1 . 6
2.18
18.4
19
TABLE I1 CHLORIDE-INDEPENDENT REACTIONAS A FUNCTION OP [TEPTjoA N D pH AT T = 25 f 0.5" AND p = 0.1 Av. k%bl
A plot of [Cl-]K'/k',b, vs. l/[TEPT]o should be linear at constant pH with a slope equal to k,/k,kr and intercept equal to 1/k, from which k, and k,/k3 can be evaluated. This was generally observed (e.g., see Fig. 5). A compilation of the results is given in Table I. The value of kl may be estimated from data a t low PH levels, where k, = kl[H+] (since kl[H+] >> kz). Similarly, a t high pH levels where k2 >> kl [H+], k, can be set equal to k2. From data a t pH 4 and 7, kl was found to be 6.33 X lo6 and k2 to be 18.4 1. mole-' min.-I. The line of best fit (used for calculation of k,/k3 and k, exp.) was constructed using the method of least squares. In some cases k, exp. was determined by direct calculation from equation 5, assuming kJk3 to be 2 X 10-6. For purposes of comparison, k, also was calculated from the above values of kl, k2 and the kz. These values are equation k, = kI[H+] shown in Table I as k, calcd. Data relevant to the chloride independent reaction is shown in Table 11. It is clear by comparing the values of koobs for a given and low concentration of T E P T (e.g. 5 X lob5) a t various PH levels that the relationship between koobs and PH is bellshaped in character, approaching a maximum around pH 6 and zero a t p H 5 and 9. The data also
+
9H
[TEPTlo X 10s
5.5
8.5 42.5 5.4 27 5.4 27 2.6 5.2 10.4 26 8 40
6.0 6.5 7.0
8.0
(hob
- k'obd)
X 10'
4.2 f 0.9 11.8 f . 2 7 . 3 f .1 11.9 f . 4 5.4 f . 6 1 l . b f .000 3.5 4.6 6.8 11.9 3.5 11.0
reveal that the rate is fractionally dependent upon [TEPT], the dependence varying with PH. Other experiments a t pH 7.0, in which the concentration of active halogen was varied while the [TEPT] was kept constant, indicated an order of dependence upon the total halogen concentration of greater than one. For example, it was found that a fourfold increase in active halogen content only doubled KO&,. The data appear to indicate that the chloride-independent reaction resulting in active halogen destruction does not occur unless significant concentrations of both hypochlorous acid and hypochlorite ion are present. At pH 9 and 10, the disappearance of active chlorine is first order with respect to both T E P T and hypochlorite ion and independent of both chloride and hydronium
NICHOLAS G. LORDIAND JOSEPH EPSTEIN
514
ions (Table 111). As in the reaction under acid conditions, three moles of hypochlorite ion are consumed per mole of reacting T E P T (equations c, d and e). TABLE I11 HYPOCHLORITE CATALYZED HYDROLYSIS OF T E P T I* = 0.1
AT
25O,
kboci- 1. [OCl-lo x 104
[TEPTlo
f i H C
9.0 9.0 10.0 10.0
1.52 1.34 1.06 1.24
11.6 5.8 11.4 5.25
x
k"
mole-' min. - 1
0.00274 .00148 .00253 ,00126
2.36 2.55 2.22 2.40
104
*
a Pseudo first-order rate constant. kocl- = k/[TEPTl0. At these pH levels, the hydrolysis of T E P T by OH- (bimolecular rate constant = 0.17 1. mole-' min.-l a t 25') can be neglected.
c
Experimental Apparatus .-A Precision Scientific Company constant temperature bath maintained a jacketed cylindrical reacThe Beckman Model K Automatic tion vessel a t ztO.5'. Titrator was utilized to maintain constant pH and measure rate of acid production. p H measurements were made coincidentally with the Beckman Model G p H meter. Absorbance measurements were made on the Beckman DU spectrophotometer. Materials.-The triethylphosphorothiolate employed in this study had the following characteristics: C, found 36.3 (calcd. 36.35); H, 7.5 (7.63); P , 15.8 (15.63); S, 16.2 (18.2); and molecular weight 226 (198). Its purity was >9570. Hypochlorous acid was prepared by the addition of nitric acid to an aqueous suspension of High Test Hypochlorite, a product of Mathieson Chemical Corporation, containing 70% calcium hypochlorite, and subsequent distillation. The distillate was treated with freshly precipitated HgO to remove hydrochloric acid and then redistilled. A 0.01 M solution of HOCl stored in a refrigerator was found to be relatively stable over a period of a month. The HOCl was standardized by titration according to established procedures.* Residual chloride was estimated by titration for free acidity. All other chemicals were of ACS reagent grade. Procedure.-Approximately 250 ml. of 0.1 M KNOa containing known quantities of chloride ion were placed into a reaction vessel. A volume of hypochlorous acid solution, usually from 1 to 5 ml., so as to result in a final concentraand 2 X lo-* M was added. tion of HOCl between 5 X (The exact concentration was estimated spectrophotometrically as described below.) The PH was adjusted to the desired level by the addition of small increments of either NaOH or HN03 and maintained by the automatic addition of 0.01 N NaOH (Beckman Autotitrator). Then an accurately measured volume (1 to 2 ml.) of freshly prepared stock solution of T E P T was added. Zero time was taken as the time when one half of the T E P T solution had been added. The rate of acid formation was followed by noting the volume of 0.01 N NaOH needed to maintain constant pH per unit time. At regular intervals aliquots of the reaction mixture were withdrawn and analyzed for total active halogen content by the following method: 1 or 2 ml. aliquots of the reaction mixture were added to a solution of 2 ml. of 0.05 N nitric acid and 3 ml. of 270 potassium iodide solution. The optical density was measured a t 363 mp, using distilled water as the reference sample and correcting for the absorbance due to the presence of excess iodine ion. The first-order rate constants, koba, were calculated directly from the slopes of plots of the logarithm of corrected optical density as a function of time.g (8) I. M. Kolthoff and E. B. Sandell, "Textbook of Quantitative Inorganic Analysis," T h e Macmillan Co., New York, N . Y., 1948, p. 307. (9) T h e data b y A. D. Awtrey and R. E. Connick, THISJOURNAL, 73, 1341, 1842 (1951), on t h e extinction coefficients indicate that the extinction coefficients of 1, (19) and IzC1- (550) were small compared to that of 1,-(26,400)a t 353 mp. T h e measured optical density when corrected for t h e presence of I - could be assumed t o be due to only the
Vol. 80
Volume changes due to the addition of titrant and/or removal of aliquots during the course of the reaction were assumed to be negligible in all calculations, since they never amounted to more than 5Q/, a t the end of the reaction. Most measurements were made only during the early stages (first 30 minutes) where volume changes were never greater than 2%. A t pH
