Kinetic Study of the Initial Cyclization of Methyl=bis(p-chloroethy1)amine Hydrochloride Using Chloride Ion-Selective Electrodes Adelbert M. Knevel and P a u l F. Kehr Department of Medicinal Chemistry and Pharmacognosy, School of Pharmacy and Pharmacal Sciences, Purdue University, Lafayette, Ind. 47907 THE METHODS FOR DETERMINATION of cyclization rate constants of acyclic P-chloro amines have been reviewed recently by Hammer and Craig (1). Methyl-bis(p-chloroethy1)amine HC1 (MBA) is a commercial antineoplastic agent (Mustargen Hydrochloride made by Merck Sharp & Dohme). When dissolved in slightly basic medium, MBA (I) cyclizes according to the following reaction :
(1)
(11)
Compound I1 then reacts with an available nucleophile, B, to form the alkylated product I11 as follows: B
/CHzCHzB
fast
\ CH2CHzCI
(11) +CHsN
(2)
(111)
Since MBA contains two chloride ions, the above processes can be repeated with the second chloride ion. However, the second cyclization, according to Cohen ( 2 ) , can only proceed after the first cyclization has been well advanced. The cyclization process is the slow step in the sN1 mechanism and, therefore, is rate limiting and first order. The rate of chloride ions produced is directly proportional to the rate of cyclization of MBA. Therefore, if the rate of chloride ion production is measured, the rate constant, k , can be calculated. Ion-selective electrodes have been reviewed (3). Srivinasan and Rechnitz (4, 5) and Fleet and Rechnitz (6) have used ionselective electrodes in kinetic studies. In light of the wide use of these electrodes in these studies and others, it seemed reasonable that a chloride ion-selective electrode could be used to study the cyclization reaction of MBA. In an alkylation kinetic study of MBA with sulfhydryl, Boatman (7) reported that the rate of alkylation was dependent upon the buffer system used. His work showed that the phosphate buffer reacted with the alkylating agent whereas the borate buffer did not. Consequently, the electrode studies were conducted in both buffer systems to deter(1) C. F. Hammer and J. H. Craig, ANAL.CHEM., 42,1588 (1970). (2) B. Cohen, E. R. Van Artsdalen, and J. Harris, J . Amer. Chem. Soc.,70,281 (1948). ( 3 ) “Ion-Selective Electrodes,” R. A. Durst, Ed., Nat. Bur. Stand. (US.) Spec. Publ. 314,Washington, D.C., Nov. 1969. ( 4 ) K. Srivinasan and G. A. Rechnitz, ANAL.CHEM.,40, 1818 (1968). (5) Ibid., p 1955. (6) B. Fleet and G. A. Rechnitz, ibid.,42,690(1970). (7) J. A. Boatman, M. S. Thesis, Purdue University, Lafayette, Ind.. 1969.
mine if the interference took place in the cyclization or the alkylation step. EXPERIMENTAL
Apparatus. Orion Model 92-17 liquid membrane and Model 94-17A solid membrane ion-selective electrodes (Orion Research Inc.) were used in this study as the indicator electrodes. The reference electrode was a double junction electrode constructed in this laboratory. The electrode consisted of a Beckman saturated calomel electrode (No. 39270) placed inside the sleeve of a Leeds and Northrup calomel electrode (No. 1199-31). The sleeve was filled with 10% KNO3 for solid membrane work and 10% K2S04for liquid membrane studies. This construction was necessary to prevent contamination of the solution with chloride from the reference electrode. The potentials developed between the indicator electrode and the reference electrode were monitored by an Orion Model 407 Ionalyzer Specific Ion Meter which was connected t o a Sargent SRLG Recorder. The recorder was set on 0.2 inch/minute and 2.0 mV full scale. Water was circulated through a 150-ml water-jacketed reaction cell from a water bath maintained at 25 0.25 “C. All pH measurements were made on a Sargent Model LSX pH meter using a Sargent S-30072-15 combination glass electrode. Reagents. The phosphate buffer, pH 7.41, was prepared using KH,PO, and Na,HPO,. Borate buffers, pH 7.40 and 7.45, were prepared using sodium borate and boric acid. The 0.14M pyridine solution was made following standard procedures. For solid membrane work, a stock solution was prepared using 5.844 grams of NaCl and sufficient buffer to make 100.0 ml. Either borate or phosphate buffer was used. For work with the liquid membrane electrode, a stock solution was prepared with 8.2 ml of concentrated hydrochloric acid and sufficient 0.14M pyridine to make 100.0 ml. Hydrochloric acid had to be used as the chloride source since NaCl in pyridine gave a solution that was too basic (pH 8.95 for 10-3M s o h tion). Above pH 8.5 hydroxyl ion interferes with chloride measurement by the liquid membrane electrode. Standard solutions of 1.00 X 2.00 X 5.00 X 1.00 X 2.00 X 5.00 X and 1.00 X 10-lMin chloride ions were prepared by serial dilution of a stock solution. These standard solutions were used to calibrate the electrodes. Calibration of the Electrodes. The response of an indicator electrode as compared to a reference electrode is given by the Nernst equation
*
E
=
constant
+ (RT/nF)log a
(3)
where E is the potential measured, the constant is characteristic of the reference electrode, R is the gas constant, T is the temperature, F is the Faraday, n is the charge on the ion being measured, and a is the activity of the ion in the solution being measured. The term, RT/nF, is called the Nernst factor. It is a constant for a particular electrode in a solution at a particular
AN ALYT lCAL CHEMISTRY, VOL. 44, NO. 1 1 , SEPTEMBER 1972
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~
~~~~~~
~~
Table I. Rate Constants from Solid Membrane Electrode in Phosphate Reaction number Rate constant, min-1 MS-MBA-7 0.080 MS-MBA-8 0 074 MS-MBA-9 0.075 Average rate constant: Average deviation: 0 0763 min-’ 0,0024 Table 11. Rate Constants from Solid Membrane Electrode in Borate Buffer Reaction number Rate constant, min-1 MS-MBA-15 0,079 MS-MBA-16 0.082 MS-MBA-18 0.079 Average rate constant: Average deviation: 0.0800 min-’ 0.0013 Table 111. Liquid Membrane Electrode in Pyridine Reaction number Rate constant, min-1 MS-MBA-10 0.071 MS-MBA-11 0.071 MS-MBA-14 0,072 Average rate constant: Average deviation: 0.0713 min-l 0.0004
temperature. The theoretical Nernst factor for a chloride electrode in a solution at 25 “C is 59.16 mV/decade. Electrode and solution characteristics cause variations in the Nernst factor such that for accurate analytical work frequent calibration with standard solutions is necessary. In this research, the Nernst factor was determined by measuring the potential of the seven standard solutions. The data (log a us. potential) were plotted and the slope of the resulting line as determined by linear regression analysis gave the Nernst factor. Procedure. The buffer and standard solutions were freshly prepared on the day of the kinetic studies. The Leeds and Northrup sleeve of the reference electrode was filled with a fresh portion of filling solution on the day that the electrode was to be used to make kinetic measurements. Potassium sulfate was used in the liquid membrane electrode studies because the liquid membrane is less selective for sulfate than for nitrate. The Nernst factor for the electrode was determined by measuring the potential developed by each standard solution. After the standards had been measured, 100.0 ml of buffer was placed into the reaction vessel and the electrodes were immersed in the solution and allowed to equilibrate during constant stirring for 10 minutes. Approximately 6.0 mg of NaCl was then added to the solution. The recorder was turned on and the response was monitored for approximately 30 minutes. This was considered the blank determination. Any drift in response was noted and a suitable correction was made in the calculation of the rate constant. After the blank determination was completed, a fresh portion (100.0 ml) of buffer was placed into the reaction vessel. The electrodes were again allowed to equilibrate for 10 minutes. Then approximately 20 mg of MBA was added to the solution. This quantity of MBA gave an initial concentration of reactant of approximately 1 X 10-3M. The recorder was turned on and the increase in chloride ion was monitored for 45 minutes. The recorder was then turned off after this time and the reaction solution was transferred to a glass prescription bottle. The air was purged with nitrogen and the bottle was tightly capped. The standard solutions of NaCl used ini1864
tially were then taken to determine the Nernst factor as previously described. Two more reactions each followed by electrode calibration were then completed. Each solution was stored as described previously after 45 minutes of reaction. The pH of each solution was measured after a designated time period. In order to prevent contamination of chloride from the combination pH electrode, the initial pH of the reaction solution itself was not measured; but instead, a solution of identical composition was measured. Calculation of the Rate Constant. The potentials at various times were measured from the recorded tracing. Since the second cyclization begins taking place before the first cyclization is completed, no infinite value of potential could be measured associated with the initial cyclization. Therefore, the Guggenheim method was used to determine the first order rate constant. A combination of the Nernst equation with the equation describing first order kinetics adapted to the Guggenheim method gave the following equation:
+
where AE“ and AE’ are the potentials at time (t ‘ A t ) and time t , respectively, as compared to an arbitrarily chosen zero point, b is the initial molar concentration of the reactant, x , is the concentration of the product at the arbitrarily chosen zero point, k is the rate constant, A is the Nernst factor, and At is 2 to 21/2times the half-life of the reaction. The half-life can be approximated by titrimetric or electrode means. A plot of log (lOAE”/A -1OAE’lA) us. t gave a slope of -k12.303.
RESULTS AND DISCUSSION Solid Membrane Electrode in Phosphate Buffer. The initial pH was 7.35 and the pH after two hours was 7.33 for all three determinations. The blank showed negligible drift. Table I shows the rate constants calculated from linear rate plots. Solid Membrane Electrode in Borate Buffer. The initial pH was 7.36 and the pH after two hours was 7.33 for all three determinations. The blank showed a drift of +0.3 mV/hr. Table I1 shows rate constants calculated from linear rate plots. Liquid Membrane Electrode in Pyridine. The initial pH was 7.36 and after 45 minutes, the pH was 7.30 for all three determinations. The blank drifted -0.6 mV/hr. This was probably a result of sulfate contaminatioc from the reference electrode. Table I11 shows rate constants calculated from linear rate plots. Evaluation of Data. Although the liquid membrane electrode gave more reproducible data, the solid electrode proved more satisfactory because it could be used with borate and phosphate buffers in any pH range. The presence of foreign ions interfered with the liquid membrane electrode to such an extent that chloride could not be measured accurately, and a buffer system of pyridine-HCl had to be used instead of the inorganic phosphate and borate systems. The HCl came from the acid salt of MBA. The pH was measured after two hours in the case of the solid electrode studies so that the data could be compared to studies where the reaction had been allowed to run two hours in an effort to measure the second cyclization rate constant. However, upon evaluating the data, it was impossible to measure the second constant accurately and conveniently, because the rate was so slow that only small potential changes
ANALYTICAL CHEMISTRY, VOL. 44, NO. 11, SEPTEMBER 1972
were recorded over long periods of time. The change in potential was of the same order as the drift of the blank. The pH of the solutions used in the liquid membrane studies was taken immediately following the completion of the study because the pH changed so rapidly. Hence no comparison could be made with the solid electrode studies. The results obtained with the solid membrane electrode in phosphate and borate buffers did not differ significantly. Therefore, it was concluded that phosphate must interfere in the alkylation step rather than in the cyclization reaction.
CONCLUSIONS
The ion-selective electrodes used in this study offered a rapid, continuous, and simple method for the measurement of rate constants without a flow-through system. It is suggested that other electrodes could similarly be applied in reaction mechanism studies and in the determination of rate constants. RECEIVED for review February 23, 1972. Accepted May 1,1972.
Detection of Nanogram Quantities of Hexachlorophene by Ultraviolet Liquid Chromatography Peter J. Porcaro and Peter Shubiak Research Department, Givaudan Corporation, Clifton, N.J., 07014
GASCHROMATOGRAPHY with electron capture has been extensively applied to the detection of subnanogram amounts of hexachlorophene [2,2'-methylenebis(3,4,6-trichlorophenol)]. The detection is usually made after derivatizing, and the most common derivatives are the trimethylsilyl ether ( I , 2), diacetate ( 3 ) ,and dimethyl ether ( 4 , 5 ) . Each has met with particular favor for various reasons. The problem of derivatizing, however, is not a major one for the eventual detection at the low level capabilities of electron capture detectors. Whether these detectors use radioactive sources (3H, e3Ni, 226Ra)or an activated rare gas (He, Ar), they are all prone to easy contamination, sensitivity losses, and a variety of other peculiar behaviors. A new approach has been made for the separation and detection of this industrially important material using liquid chromatographic techniques and ultraviolet detection. The method has shown itself to be sensitive, reliable, and troublefree compared to the use of electron capture detection using gas chromatography. Ultraviolet detection has thus far been used only for hexachlorophene in the macro region in applications such as drugs and cosmetics (6, 7). EXPERIMENTAL
Reagents. A 55/45 V/V mixture of hexane and n-butylchloride is used. Spectro or reagent grade material is adequate with no further purification necessary. Procedure. The UV detector is used at its maximum sensitivity, 0.02 absorbance full scale. A 2-ft stainless steel 2.3-mm i.d. column is packed with Sil-X silica 36-40 p particle size (Nester-Faust) with no prior conditioning or activation. A constant flow rate of 0.7 ml/min is maintained, generating a pressure of approximately 200 psi for the system which has two restrictors and pulsation dampers in-line as supplied. Septum used is Viton with no pre-leaching required. Buna-N and EPR are not suitable. Derivative Preparation. The derivative used is the di-pmethoxy benzoate or dianisate ester of hexachlorophene (HCP-DA). CI
0
-
Ee
OCH,
CI
Apparatus. A Waters Model ALC 202/R-401 Liquid Chromatograph was used. It incorporates a differential refractometer and an ultraviolet detector. The UV detector is made by Laboratory Data Control (Riviera Beach, Fla.) which employs a low pressure mercury lamp source emitting its strongest radiation at 254 nm. (An alternate optional bandpass is available at 280 nm, if desired.) The recorder used was a Honeywell Electronic Model 194. UV spectra were recorded on a Beckman ACTA I11 spectrophotometer. (1) J. Wisniewski, Facts Metlzods, 8, 10 (1967). (2) P. J. Porcaro and P. Shubiak, ANAL.CHEM., 40, 1232 (1968). ( 3 ) R. S. Browning, Jr., J. Grego, and H. P. Warrington, Jr., J. Pharm. Sci., 57, 2165 (1968). (4) A. Curley, R. E. Hawk, and R. Kimbrough, Lancet, 2, (7719)
296 (1971). ( 5 ) W. H. Gutenmann and D. J. Lisk, J . Ass. Ofic.Anal. Chern.,
53, 522 (1970). (6) D. A. Elvidge and P. Peatrell, J. Pharm. Pharmacol., 13, l l l T (1961). (7) R. W . Daisley and C . J. Olliff, ibid., 22, 202 (1970).
2,2 '-Methylenebis(3,4,6-trichlorophenol)di-p-methoxybenzoate
Dissolve 2.0 grams of HCP in 60 ml of 10% NaOH contained in a 125-ml erlenmeyer. Add 8.0 ml of anisoyl chloride (Aldrich Chemical Co.) and mix contents for 1 hr. Collect crystals on a coarse sintered glass funnel using slight vacuum. Wash the crystals with 100 ml of distilled water and allow vacuum to air dry, Transfer to 50 ml of hexane, and add diethyl ether until solid dissolves. Filter and heat solution gently on a steam bath until it starts to boil. Remove and place in an ice bath to induce crystallization. Recrystallize two additional times. Dry at room temperature. The yield is approximately 500 mg of pure diester, mp 209.5-10 "C. Structure was verified by NMR, IR, and MS. The UV spectrum of the diester (HCP-DA) is shown in Figure 1, and the UV scan of underivatized HCP is contrasted as shown in Figure 2. Calibration Curve Using Pure Diester. Ten milligrams of prepared HCP-DA are dissolved in exactly 100 ml of n-butylchloride. Aliquots of 0.4, 0.8, 1.2, 1.6, 2.0, 2.4, and 2.8 ml
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