THE JOURNAL OF
PHYSICAL CHEMISTRY (Registered i n
U. S. Patent Ofice) (Q Copyright, 1988, by the American Chemical Society)
MAY13, 1963
VOLUME67, NUMBER 5
THE HYDROLYSIS OF ~-DIIETHYLA3lINOETHYLACETATE BYJ. A. ZASLOWSKYAND E. FISHER' Olin Mathieson Chemical Corporation, New Haven, Connecticut, and Niagara University, Niagara Falls, New York Received May 81, 1968 The hydrolysis of P-diethylaminoethyl acetate was investigated over the pH range of 5.5 to 8.4. The data indicate that the hydrolysis proceeds by the interaction of the acid salt of the amine with hydroxyl ion rather than the kinetically equivalent mechanism involving the amine with water. This is explained on the basis of the polarization of the carbonyl bond of the ester by the favorably located positive charge.
Experimental
Introduction The physiological activity of p-aminoethyl carboxylates has prompted .this investigation of the kinetics of
1. Preparation of p-Diethylaminoethyl Acetate Hydrochloride
0
k hydrolysis of a simple prototype a t physiologically significant pH values. Davis and Ross2compared the rates of hydrolysis of ethyl acetate, dimethylaminoethyl acetate, and acetyl-. choline in acidic and basic SOY' aqueous acetone solu-. tions a t 50". The marked increase in rate noted in basic solution for aceitylcholine was attributed to the positive charge on the P-nitrogen atom. The formation of a sterically favorable activated complex was postulated. fX
,('\
0
R--
4t \
CHn
I
CH2 /
OH- 0
The increased rate of hydrolysis of dimethylaminoethyl acetate (relative to ethyl acetate) was attributed to activation of the carbonyl bond by the &nitrogen atom electrons. The present researchL deals with a study of the comparative kinetics of hydrolysis of P-diethylaminoethyl acetate, acetylcholine, and ethyl acetate over the pH range 5.5-8.4 a t 50-65'. (1) In partial fulfillment of the requirements for the Master of Science Degree a t Niagara University, 1960. Presented a t the 137th National Meeting of the American Chemical Society, Cleveland, Ohio, 1960. (2) W.Davis and W.C. J. Ross, J . Chem. Soc., 3056 (1950); 2706 (1951).
(DEA.HCl).-The material was prepared by the method of Gilmana from acetyl chloride and P-diethylaminoethanol in ether. A 5-10% excess of acetyl chloride was employed. The crude product was washed with ether, recrystallized from dry acetone, and driedinvacuo. The n1.p. was 113-114", literature 116-11703; C1- analysis: 18.26%; theor.: 18.12%. 2. Kinetic Studies with Diethylaminoethyl Acetate and its Hydrochloride Salt.-The rates of hydrolysis were determined a t constant pH with a Model K Beckman automatic titrimeter. The pH was effectively controlled to the nearest 0.05 pH unit. The temperature of the reaction was maintained to f0.1' by means of a conventional thermoregulated water bath. A weighed sample of DEA.HC1 (0.3 mmole) was added t o about 100 ml. of distilled water a t the reaction temperature. The pH was rapidly adjusted with 0.05 N NaOH (from the titrimeter). The hydrolyses were conducted a t pH 5.5, 7.5, 7.9, and 8.4 a t 60' and at pH 8.0 a t 50, 55, and 65". The initial value for kinetic evaluation purposes was determined from the terminal base consumption. The data were treated on the basis of pseudo-first-order kinetics. First-order constants were determined by the method of least squares. There was no significant deviation over the entire range of reaction. 3. Determination of the K , of DEA.HCI.-The partial titration curve was determined for weighed quantities of DEA.HCI (0.003 Il.1solutions) a t 50, 55, 60, and 65". From the per cent titrated a t each pH the K , was calculated. The values of K , a t the various temperatures permitted a calculation of the fraction of the total amine present as the free base and as the amine salt a t the experimental pH value^.^ BH+/Bo T,OC. PH K. 0.658 50 8.0 5.2 x 1 0 - 9 .614 55 8.0 6 . 4 x 10-9 ,997 60 5.5 7 . 9 x 10-9 60 7.5 .799 ,613 60 7.9 ,334 60 8.4 ,502 65 8.0 9 . 9 x 10-9 B H + = acid salt. Bo = total ester. 4. Determination of Hydrolysis Rates of DEA.HC1 in Deuterium Oxide.-The same procedure was employed aR in water. (3) H. Gilman, L. C. Heckert, and R. McCracken, J . A m . Chem. Sac., SO, 437 (1928). (4) J. W. Churchill, M. Lapkin, F. Martinez, and J. A. Zaslowsky, ibid., 80, 1944 (1958).
959
J. ,4.ZASLOWSKYASD E. FISHER
960
TABLE I HYDROLYSIS OF DEA. HCl (60')
The p H meter readings were however corrected by the method of Long as described by Bender5
pD
=
pH
+ 0.4
K,(D,o)
=
4.09
x
10-9
5 . The Hydrolysis of Acetylcholine Chloride.-Acetylcholine chloride (Matheson, Coleman and Bell) was recrystallized from acetone. The dry hygroscopic salt had a melting point of 150151" in agreement with Jones and Major.6 The hydrolyses mere conducted in a manner similar to that described for DEA.HC1. The kinetic data, a t BO", are summarized PH
hi, sec.-l (pseudo-first-order)
ki, 1. mole sec. -1
x x x x
26.4 28.0 29.4 29.2
6.0 6.9 6.9 7.8
2.53 2.13 2.24 1.77
10-5 10-5 10-4
Av.
k2:
28.2
6. The Hydrolysis of Ethyl Acetate.-The reaction was studied a ~60" t employing the same general technique which has been described, provision being made to operate in a closed system to prevent evaporation losses. The data are summarized PH
6.0 7.0
*
x x
Discussion Although the rate of hydrolysis over the entire pH range can be expressed by the summation
+
+
+
h(B) b(BH+) h(B)(H+) h(B)(OH-) ks(BH+)(OH-) ks(BH+)(H+) (where B is the concentration of the free base of the ester and BH+ is the acid salt) the relationships
f
=
fl =
=
(kif
ki, sec. -1
6.17 x 6.39 x 6.39 x 5.87 x 5.55 x 5.40 x 5.89 x 6.17 x 6.54 x AV.: 6.04 x
10-3 10-3 10-3 10-3 10-3
10-3 10-3 10-3 10-3 10-3
Least squares value: 6.23 X
presented in Table I. Assuming the important kinetic term to be Izl(B), the autohydrolysis of the free base, the pseudo first-order constant could be converted to a true first-order constant by dividing the experimental yalue by the fraction of free amine present a t each pH value (f). The results are given in columns 3 and 4 of Table I. The constaiicy of the results over the 1000-fold range of hydrogen ion concentration suggest that the assumption was valid. A statistical treatment confirmed the fact that within the precision of the experiments the equation
B
=
fk,
+ hf' + h f ( H + ) + hf(OH-) + k~fl(OH-)+ hfl (H +))Bo
First-order constants were calculated a t four different pH values (5.5-8.4) at 60". The data are (5) M. L.Bender, E. J. Pollock, and M.C. Pieveu, J. Am. Chem. Soc., 84, 595 (1962). (6) L. W.Jones and R. T. Major, ibzd., 69, 200 (1947).
+ H , O ~ B H ++ OH-
Kb
=
(BH+)(OH-) - K _, (B) K,
permits ai1 alternate kinetically equivalent interpretation k,(B)
=
kl(BH+)(OH-) Kb
=
K5(BH+)(OH-)
The value of k5 was calculated directly from kl assuming the latter to be the significant contributor. K , a t 60" was taken as 9.6 X 10-14.' The least square value of kg (at 60") was 516 1. mole-1 sec.-l. This value has the same statistical significance as the firstorder constant k , previously developed. The following alternate kinetic terms were determined a t each temperature T.OC. 50 55 60 65 60 (DzO)
B/B~ BH+/Bo
permit all the terms to be converted to linear functions of Bo (total ester concentration). The following equation results after rearrangement of the terms
V
f 0.666 ,666 ,666 ,387 ,387 ,387 201 ,201 ,0025
10-3 10-3 10-3 10-3 10-3 10-8 10-3 10-3 10-6
represents the data to high degree of confidence (>95%). The equilibrium
7. The Determination of the Presence (or Absence) of the Ethyleneimmonium Ion during the Hydrolysis of DEA.-The rapid reaction of ethyleneimmonium ion with sodium thiosulfate suggested that the presence of thiosulfate ion in the hydrolysis solution of DEA would monitor (at least qualitatively) the formation of the cyclic ion. A known excess of sodium thiosulfate was added to a hydrolysis solution a t pH 8. After approximately 60% reaction, an aliquot of the solution was titrated with standard iodine for thiosulfate assay. There was no consumption of thiosulfate during the hydrolysis.
+
4.11 x 4.25 x 4.24 x 2.27 x 2.15 x 2.09 x 1.18 x 1.24 x 1.64 x
k exp.
1.04 1.18
10-7 10-6
Av. k2: 1.1
+
8.4 8.4 8.4 7.9 7.9 7.9 7.5 7.5 5.5
kz, 1. mole sec.-I
kr, sec. -1
1.0 0.08 1.1 f 0 . 1
k exp. (pseudo-firstorder constant), sec. -1
PH
The calculated pseudo-first-order reaction rate constants a t 60' and pD 8.4 were 1.57 X sec.-1 and 1.66 X set.-' in duplicate experiments. The K , ( a t 60') of DEA.HC1 in DzO was determined using the relationship between pD and the measured pH.
Vol. 67
kh, 1. mole-' see.-'
kl, set.-'
2.71 x 4 80 x 6 23 X 9 17: x 3 18 x
The data fit the equation: s w - 1 (for water).
10-3 10-3
259 413
10-8 10-3
715 785
516
-15.480 __
k6
=
1.73 X 10IOe
RT
1. mole-'
The value of K,(D,O) necessary for the determination of k,(D,O) was obtained from the data of Abel and co-workers8 and the variation of K,v(D,O) with (7) H. 6. Harned and R. A. Robinson, Trans. Faraday Soe., 36, 978 (1940).
(8) E. Abel, E, Brateau, and 0. Redlioh, 2. phgsilc. Chem., 173, 353 (1936).
May, 1963
HYDROLYSIS OF 6-DIETHYLAMIKOETHYL ACETATE
961
temperature reported by Wynne-Jone~.~KW(D2O) catalyzed hydrolysis as acetylcholine. These data a t GOo is 1.7 X loV1$;K,(D20) a t 25’ is 1.6 X 10-15. indicate that a positive charge on the nitrogen atom has The ratio kl(DBO)/Jcl(HSO) is 0.51; the ratio k h a markedly accelerating effect on the hydrolysis of (D20)/kb(H20) is 1.52. The data of Laughtonlo” @-aminoesters; it is reasonable to expect the effect to and ReitzlObfor a series of solvolysis reactions indicated be more marked when the positive charge is carried a ratio of slightly less than unity for k(D20)/k(H20). by a labile proton rather than by a quaternary nitrogen The value for kg was in accord with the work of Long atom. It appears that the electrically neutral @and co-workersll and Wiberg12on the expected isotope nitrogen atom is not a particularly effective activating effect for OH- catalyzed reactions, e.g., ethyl acetate, group. It should be noted that a t complete hydrolysis of 1.33; methyl acetate, 1.60. Since recent workla has shed considerable doubt as to the utility of the deuDEA the theoretical quantity of acid was not liberated. terium isotope criterion for unequivocal mechanism Although this did not interfere with the kinetic interpreassignment, no attempt will be made to utilize this tation, it was of interest to correlate this “missing” criterion for distinguishing the two possible mechanisms. acid with the base strength of the formed alcohol h cyclic mechanism can participate in hydrolysis (diethylaminoethanol). reactions of p-aminoethyl ester^.^,^^ l5 The question For example a t p H 8.4, 33.4% of the original ester was present as the acid salt; only 54% of the theoretical to be resolved is then that of deciding which of the following is the probable reaction path quantity of acid was liberated a t the conclusion of the reaction. These data indicate that 46% of the formed acid was bound as the acid salt of the alcohol. The calculated K,, a t 60” of diethylaminoethanol is 1.0 X (I
