Thiazoline and Oxazoline Hydrolyses and Sulfur-Nitrogen and Oxygen

R. Bruce Martin, Regina I. Hedrick, Alice Parcell. J. Org. Chem. , 1964, 29 (11), pp 3197–3206. DOI: 10.1021/jo01034a018. Publication Date: November...
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NOVEMBER, 1964

THIAZOLINE A N D OXAZOLINE HYDROLYSES

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Thiazoline and Oxazoline Hydrolyses and Sulfur-Nitrogen and Oxygen-Nitrogen Acyl Transfer Reactions R. BRUCEMARTIN,REGINAI. HEDRICK, AND ALICEPARCELL Cohh Chemical Laboratory, University of Virginia, Charlottesville, Virginia Received June 8, 1964 Hydrolysis of 2-xnethylthiazoline is general base catalyzed on the basic side of the maximum in the pH-rate profile. On the acid side of the maximum a plot of log kobsd - Ho us. log aHlo is linear, conforming to the result expected for the appearance of an acid-inhibited reaction in acid solutions. Hydrolysis rates of 4-carboxy-2methylthiazoline ethyl ester and 4-carboxy-2,5,5-trimethylthiazolineethyl ester also exhibit maxima in their pH-rate profiles and exhibit acid inhibition of hydrolysis in acid solutions, providing further examples of compounds where formation of a tetrahedral addition intermediate seems necessary to account for the results. The rate of appearance of thiazoline rings from X-acetylmercaptoethylamine and N-acetylpenicillamine is equal to the rate of disappearance of amide and exhibits a maximum rate in weak acids, finally becoming nearly constant in strongly acid solutions. The difficulties in dovetailing the general catalyzed acetyl transfer reaction of Sacetylmercaptoethylamine with other results in the 2-niethylthiazoline system are described. The hydrolysis of 2-methyloxazoline in weakly acid solutions is dependent upon at least the 15th power of the activity of water, substantiating the previous suggestion that an acid-inhibited step occurs in this system also. There is no difficulty in incorporating the results for hydrolysis of 2-methyloxazoline with the general base-catalyzed acetyl transfer from 0 to N in 0-acetylethanolamine. Suggestive evidence is presented for intraniolecular amidate formation when the thiazoline ring formed in concentrated acid solutions of glutathione is exposed to neutral solutions.

Primarily to account for the bell-shaped curve obtained when the initial rate of hydrolysis of 2-methylthiazoline is plotted against pH, the scheme of Fig. 1 was presented.' The half-maximum rate on the basic side of the maximum occurs a t pH = pK1, and is accounted for by reaction of water with protonated species (TH+) in the k , step of Fig. 1. On the acid side of the maximum, the half-maximum rate occurs a t [ H + ] = (kg 1c5)/k2 in the assignments of Fig. 1. Qualitatively, thiazoline hydrolysis is inhibited in*acid by t'he appearance of n proton in only the k z and not in the k3 or IC5 steps of Fig. 1. The scheme of Fig. 1 also accounts for the pH independence of the ratio of formation of S-acetylmercaptoethylamine (S) to formation of S-acetylmercaptoethylaniine (K) from initial solutions containing only thiazoline as well as feat'ures of the disappearance of reactant in acid solutions containing initially either s- or S-acetylmercaptoethylamines. I n order to accommodate all these results kinetically it was necessary to postulate the existence of an unstable intermediate to which the steady-state approximation would apply. On chemical grounds the tetrahedral carbon hydroxythiazolidine (U) addition compound was selected as a likely unstable intermediate. On the basic side of the maximum in the pH-rate profile for thiazoline hydrolysis, formation of the tetrahedral addition intermediate is rate limiting while on the acid side of the maximum the decomposition of tjhe intermediate is rate determining.* One of the properties of the thiazoline systems represented in Fig. 1 that facilitates the analysis of the interrelationships is the possession of unique ultraviolet absorption properties by each compound. Protonated

+

(1) R. B. Martin. S. Lowey, E. L. Elson, and J. T. Edsall, J . Am. Chem. Soc., 81, 5089 (1959).

(2) I n a footnote B. Zerner and M. L. Bender [ibid., 83, 2267 (196l)I claim t h a t some of the conclusions drawn from t h e scheme of Fig. 1 are invalid because the protonated forms of the intermediate is not included. I t is stated in ref. 1. however, t h a t , "as a consequence of the steady-state approximation the equilibrium involving the conjugate acid of hydroxythiazolidine need not be explicitly considered." Several kinetically indistinguishable mechanisms are often available for describing results and t h e rate constants of one mechanism will include equilibrium and rate constants of another. Furthermore, we show in this paper (see eq. 6) t h a t t h e requirements of general catalysis make the scheme of Fig. 1 the more accurate way of describing the reaction pathways.

thiazolines absorb maximally a t 260, free base thiazoline a t 230-250, and thio esters a t 230 nip; it has heen possible to follow amide absorption a t 200 mp on the slope of an absorption curve with a maximum further in the ultraviolet. Xot only is it possible to nieasure the init,ial rate of disappearance of one of the three compounds, but also simultaneously the rnt>eof appearance of the other two. I n addition, equilibrium studies have been performed to yield values for equilibrium constants defined as follows in the terminology of Fig. 1. Only KST = [SH+]/[TH+] = kiks/kzkeKiz K N T = [N]/[T] = kika/kzkaKi

KNS

=

[H+][N]/[SH+] = k3k6KJkSki

(2) (3)

two of the t8hreeconstants are independent as they are interrelated by the equation KSTKNS= K 1 K N T . The equilibrium ratios determined spectrophotometrically yield equilibrium constants which check well witfh the values obtained by substitution of the rate constant's in eq. 1-3.1,3 A summary of the values of rate and equilibrium constants for the 2-inethylthiazoline system is presented for easy reference in the caption of Fig. 1. Later rate and equilibrium studies indicated that the scheme of Fig. 1 is applicable to hydrolysis of other thiazolines, including 2-ethylthiazoline, 2-methylthiazoline-N-methylperchlorate, 2-(l-acetaniino-2-1nethylpropyl)thiazoline, and 2-methylthiazine, the six-membered ring analog of 2-methylthiazoline.3 Partit,ioning of the tetrahedral addition intermediate to yield either S- or N-acetylmercaptoethylamines k 5 / k z was indicated to be independent of pH and ring size but dependent upon substituents in the 2-positi0n.~ Features of' the hydrolysis of 4-carboxy-2-methylthiazoline4 were also shown to be consistent with the scheme of Fig. 1. In this report results for the hydrolysis of 4-carboxy-2niethylthiazoline ethyl ester and d-carbosy-2,5,5-triniet,hylthiazoilne ethyl ester are presented. Thiazoline ring formation has been reported to occur in concentrated acid solutions of several N-acylmercaptoethylamines, but few definitive quantitative (3) R . R. Martin and A. Parcell. ibid., 83, 4830 (1961). (4) H. S. Smith and G . Gorin. J. O r B . Chem., 46, 820 (1961)

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2.

11 Kl H+

+

SH

I

NH2

(SI c 0

NH I

I

*S

n c=o

I CH3

+ H+

1tK2 n

s

INTERCONVERSIONS Of 2

NH%+

- M E T H Y L - A 2 -T H I A Z O L I N E

Fig. 1.-Scheme for the interrelationships among 2-methylthiazoline and S- and N-acetyl-4-mercaptoethylamines. Values for the rate and equilibrium constants for this system are3 min.-', k3/kZ = 0.05 M, ks/kl = pK1 = 5.22, kl = 1.05 X 0.06 M , ke/kz = 4.5 X min.-' M , pK? = 9.1, ka = 1 X min.-', KST = 11, K N T = 8.5 X lo4,and K N S= 0.045 M .

studies have been made. Results are reported in this paper for the simultaneous measurement of amine disappearance and thiazoline ring appearance in both Tu'acety1-/3-mercaptoethylaniine and X-acetylpenicillaniine. Transfer of the acetyl group from sulfur to nitrogen in S-acetyl-/3-mercaptoethylaniine is acconimodated by the scheme of Fig. 1 only a t pH 10k6. With the smallest basic form a t each pH, Ka/( [ H + ] Ka) where pKa = (k3 kS)/k2 value of any thiazoline system as well as 3.7 for formic acid, both pH values yield 4.2 X the unique result that k , > k5, appearance of the thin h - ' M-I as the catalytic constant for formate ion. azoline derivative in an attempted synthesis of SThus general base catalysis of thiazoline hydrolysis is acetylpenicillaniine ethyl ester mentioned in the Exquantitatively established. perimental section receives a natural explanation. Acid Inhibition of Thiazoline Hydrolysis.-2-;\IethylGeneral Base Catalysis of Thiazoline Hydrolysis thiazoline exhibits inhibition of hydrolysis in acid The effects of general catalysts on thiazoline hydrolysis solution.' This inhibition sets in a t about pH 2 where were studied by following the initial rate of disappearthe activity of water is still nearly unity. To account ance of absorption a t 264 nip of 2-niethylthiazoline-Nfor the acid inhibition it has been proposed, as shown methyl perchlorate in several concentrations of formic in Fig. 1, that hydrolysis proceeds via attack of water acid-formate buffer a t constant temperature and ionic upon protonated thiazoline to yield a hydrated proton strength. Unlike 2-niethylthiazoline, the X-methyl and a hydroxythiazolidine tetrahedral addition intersalt exhibits a plateau in its pH-rate profile from pH mediate which may react further.' Reactions which 2-5.3 I n this pH range the rate-limiting step is attack liberate protons in acid solutions should exhibit a by water upon the positively charged compound and decrease in rate in acid solution. Bunnett7 has shown the plateau provides a constant base line from which to that plots of the logarithm of the observed first-order study effectsof added catalysts in thiazoline hydrolysis. rate constants us. the logarithm of the activity of water When the observed first-order rate constants ICl' for are often linear with a slope the value of which is chardisappearance of 2-niethylthiazoline-N-methyl peracteristic of the role of water in the reaction mechachlorate a t 25.0" and 0.3 ionic strength maintained with nism. In order to accommodate reactions which also KC1 are plotted against the total concentration of liberate protons in acid solutions, it has been proposed8 formic acid and formate in both forms, straight lines that the logarithm of the observed first-order rate conare obtained to 0.4 M total formate a t pH 4.2 and to 1 stant iiiinus Ho (Hamniett acidity function) us. the M total forniate a t pH 3.2. For both pH values the logarithm of the activity of water is a suitable funcintercept at zero formate concentration is a t kl' = tion to plot. The slope of such linear plots would have 1.6 X niin.-l in agreement with the earlier r e p ~ r t . ~ the same significance as the slopes of Bunnett plots.* A Bunnett plot of log kl' us. log a H I O for the hydroly(10) The Kz/Ki ratio of S m i t h and Gorind is equivalent to our K B T . From their results we also obtain our K N S = 1.6 M and combination w i t h sis of 2-niethylthiazoline is a curve with an initial slope our pKi 3 . 2 yields KNT = 1.3 X IO'. This K N S value is the greatest yet of 30 or niore decreasing as the acidity increases." evaluated. We also calculate t h a t k, = 1 X 10-4 min.-l. These approxiFig. 2 shows the same data plotted according to the mate constants may be considered applicable to the system derived from 4-carboxy-2-methylthiaaoline ethyl ester and coinpared w i t h bhose in other proposal8 for reactions exhibiting inhibition in acid thiazoline systems.3 I n concentrated HCl and " 2 1 0 4 solutions their solutions. Owing to the very slow rates, measurements equilibrium ratio of [ N H + l / [ T H 'I = 2.6, where ["+I represents protonI

+

+

+

ated amide. Then, if the acid ionization constant for protonation a t t h e amide bond is designated as Ka, we have K3 = K N T K1/2.6 3 , consistent w i t h ionization constants a t similar amide bonds presented later. 5

(11) Values of logarithm of activity of uater are taken from ref. 7 and acidity functlon from M. A. Paul and F. A . Long [Chem. Rev., 67, 1 (1957) 1.

MARTIN,HEDRICK, A N D PARCELL

3200

. 4.0

0

-I

.-

-log

Fig. 3.-Plot of logarithm of observed first-order rate constants in min. -1 for hydrolysis of 2-methyloxazoline us. logarithm of activity of water in HC10, (solid circles) and H z S O (open ~ circles) a t 25.0".

were not extended to solutions more concentrated than about 3.5 M HC1O4. The slope of 8.2 is in the range expected' for water functioning a t least8 as a protontransfer agent in the rate-limiting step. The intercept on the ordinate, 1.3 X min.-' M for unit activity of water, corresponds to ICl (k5 k 3 ) / k 2for the constants min.-l, in the scheme of Fig. 1. Since k l = 1.05 x (kb ka)/k% = 0.12 M ,in good agreement with the value of 0.11 M obtained previously from the halfmaxiniuni rate on the acid side of the bell-shaped, pHrate profile for hydrolysis of 2-niethylthiazoline. Acid Inhibition of Oxazoline Hydrolysis -2-Methyloxazoline also exhibits inhibition of hydrolysis in acid solutions.6 In this case, however, since inhibition does not set in until pH 1, it is more difficult to decide whether the decreased hydrolysis rate is due simply to reduction in the activity of water or also to an acid-inhibited step as in the hydrolysis of 2-niethylthiazoline. Fig. 3 shows a Bunnett plot of the logarithm of the observed first-order rate constant us. the logarithm of activity of water for hydrolysis of 2-n~ethyloxazoline. In solutions where -log uHlO = 0.0 to 0.1 the curve of Fig. 3 yields slopes of 17 and 14 in HC104 and HzS04, respectively. In more acidic solutions to -log U H ~ O = 0.6, tangents drawn to the curves give slopes of 5 or so, depending upon where the tangents are drawn. A proton magnetic resonance study'? of 2-methyloxazoline hydrolysis in HC1 solutions yields a slope close to zero when log kl' - log [HCl] - Ho is plotted against log U H ~ OA . zero value in such a plot is equivalent to a slope of about 5 when the ordinate function is that of Fig. 3. Thus the same slope is obtained in nioderately concentrated solutions of three different acids determined by two different methods. At a given water activity the rate in H2S04is greater than that in HC10,. When the same data for 2-methyloxazoline hydrolysis

+

+

(12) R. Greenhalgh, R. M. Heggie. and M. A. Weinberger. Can. J . Chem.. 41, 1662 (1963); R. Greenhalgh, Nafure. 196, 267 (1962).

VOL. 29

are treated as for 2-methylthiazoline hydrolysis in Fig. 2, the plot exhibits an initial negative slope to -log U H ~ O= 0.3, then tends to level off to zero slope. Two experiments of oxazoline hydrolysis were conducted in 88% DzO a t p H meter readings of 1.6 and 2.1 in this solvent mixture. At both pH values the hydrolysis rate is half that observed in ordinary water on the broad niaxiniuni from about 1 < p H < 1 in the bell-shaped pH-rate profile. Disappearance of N-AcetylmercaptoethylaminesThe initial rate of disappearance of N-acetyl-p-mercaptoethylamine as well as the appearance of 2-methylthiazoline has been followed in acid solutions a t 200 and 260 nip, respectively. Both rates increase with increasing acidity, attaining a niaxinium a t about pH 0.5, and then decline until about Ho = -2 where the two rates are identical and remain approximately constant a t about low5min.-', as the acidity is increased further. The observed first-order rate constant for disappearance a t the maximum is about 0.8 X lo-* niin.-l. Definitive interpretation of .these results on either an equilibrium or rate basis is complicated by the large number of factors to be sorted out. In these acid solutions Y-acetyl-p-mercaptoethylaniine may undergo transfer to S-acetyl-P-mercaptoethylaniine and both reactant and this product may hydrolyze. As the acidity increases and the ratio k 6 / k z[H+] becomes small, there is no kinetic pathway for the transfer reaction even though S-acetylmercaptoethylamine is favored over reactant amide and 2-methylthiazoline by the values of the equilibrium constants given in the caption of Fig. 1. Formation of 2-methylthiazoline also takes place but not until pH ' kz[HB] and formahydrolysis of 2-methylthiamline, addition of a carboxyl tion of the tetrahedral addition intermediate i s rate ethyl ester in the 4-position increases k l , the rate of limiting. formation of a tetrahedral addition intermediate, At the half-maximum value on the acid side of the over 30-fold. This substitution does not markedly pH-rate profile for the hydrolysis of 2-methylthiazoline, affect partitioning of the tetrahedral intermediate to the pH = 0.95 corresponds to k,/kz = 0.11. At pH give products or reactants, (ka kS)/kZ, nor the ratio values less than unity the kz[HB] product, where [HB] of S-acetyl to N-acetyl derivative in the products, is hydronium ion, exceeds k,, and hydrolysis is markedly k5/k3. Further substitution of two methyl groups in the 5-position, to give 4-carboxy-2,5,5-trimethylthiazo- inhibited. When kz[HB] exceeds k,, eq. 8 reduces to line ethyl ester more than halves the rate of formation of tetrahedral addition intermediate, kl, compatible with fii = kDki[HzO]K~/kz[H+l (9) an increased value of pK1, increases the probability where K B is the acid ionization constant for the HB-B of the intermediate partitioning to reactant thiazoline catalytic system, in this case hydronium ion and water. over products, and most markedly favors formation of Since we have used the scale of unit activity for pure S-acetylpenicillamine ethyl ester a t the expense of Swater, Kg = 1. Equation 9 predicts that a plot of acetylpenicillamine ethyl ester. The low values of log kl' - Ho us. log a H z O should be linear with a slope (k, k 5 ) / k 2 for these two 4-substituted thiazolines corresponding to water acting as a nucleophile.8 This provide the fourth and fifth thiazoline derivatives in predicted slope of S l . 2 to +3.3 is much less than the which acid inhibition of hydrolysis cannot be accounted observed slope of +8.2 in Fig. 2 which is in the range for by the decrease in activity of water. expected for water functioning as a proton transfer Thiazoline Hydrolysis.-Two independent lines of agent.7 If water were added as a catalyst on the leftevidence indicate that hydration of thiazoline cation hand side of eq. 7, the high slope could be acconmioto give a tetrahedral intermediate is general base dated. To do so, however, yields a result inconsistent catalyzed so that the scheme of Fig. 1 should be modiwith the results of S-N transfer discussed below. fied to give Results reported for the appearance of thiazoline ki derivatives from N-acetylmercaptoethylamines are TH+ + H*O + B + HB (6) consistent with eq. 6 and the scheme in Fig. 1. As I n the plateau region from 2 < pH < 5 where interstated in the results section, acid catalysis of the k , mediate forniation is rate limiting, S-methylthiazolinestep does not occur a t pH >0, and eq. 5 is consistent N-methylperchlorate hydrolysis is general base catawith this result. In so far as the N-acetylmercaptolyzed as described above, and the hydrolysis rate is ethylamine-thiazoline interconversion is concerned, halved in 80% D20.3 Thus eq. 6 seems established the reaction D B Ft S HB is also acceptable for the hydration step in thiazoline hydrolysis. and in fact preferred if the argument of the previous Comparison of the catalytic constants for water and paragraph is to be accommodated. This last reacformate ion catalyzed hydrolysis of 2-niethylthiazolinetion seems quite unacceptable, however, from the point N-methyl perchlorate yields a Brqinsted exponent of of view of the S-N acyl transfer reaction because this about 0.4. Since the reaction of the same compound reaction is general base catalyzed in ternis of protonated with hydroxide k~ = 13 X lo3 illin.-' M-1, is amine, rather than free amine as required by the last more than 100 times faster than the Br@nstedexponent equation. would predict, direct nucleophilic attack of hydroxide S-N Acetyl Transfer.-Two lines of evidence suggest ion on the ring is evidently much more rapid than its the occurrence of a two-step mechanism and the forfunction as a general base. Thus the previous intermation of a tetrahedral intermediate in S-N transfer in S-acetyl-P-niercaptoethylamine (S). This evidence (19) K. Linderstrorn-Lanp and C. F. Jacobsen, J . B i d . Chem., 187, 4 4 3 (1041). is independent of that given above for a tetrahedral

+

+

SD

+

+

MARTIN,HEDRICK, A X D PARCELL

3204

intermediate in thiazohne hydrolysis. The three-part pH-rate profile for S-?J transfer shown by l?ig. 1 of ref. 3 and described in the Results section of this paper is difficult to explain on any other basis. I n addition the leveling off of the transfer rate at high concentrations of formate buffer strongly suggests that the ratelimiting step has been altered froni one containing catalytic components to one that is uncatalyzed. Such a leveling off is presumptive evidence for a change in the rate-limiting step. Whenever a change in the rate-limiting step is suspected or is being sought, the possibility of a leveling off as occurs in Fig. 3 should be investigated. A niechanisni which can account for the details of S-N acetyl transfer in S-acetyl-6-niercaptoethylaniine written with the same synibols used in Fig. 1 follows.

zS + H + Kz [SI [H’]/[SH+l S + HB zD H + + B K S [S][H+]/[DH+l SH+

=

ks

=

k7

DHC

D

+ H+

KD = [D][ H + ] / [ D H + ] ka

D+N

Assuming a steady-state condition for the tetrahedral intermediate we obtain -1

v = - - - -

[SH+]

d(SH+] Kzh[HBl /[H+1 dt 1 [H+]ki[Bl/k,Ko

+

where k , [HB] and k , [B] represent the sum of the products of the catalytic coefficients and each general acid and general base concentration, respectively. I n the most acid solutions an equilibriuni prevails in the catalyzed step and eq. 10 becomes vH[HL] = K2kaKD/Ka = 2.0 X 10-6min.-’.W

(11)

This numerical value is obtained from the line of unit slope a t 2.5 < pH < 3 of the logarithm of observed firstorder rate constant us. pH plot.5 In the plateau region, 3 < pH < 4, of the pH-rate profile, where the observed rate is independent of p H , eq. 10 gives v H I= ~ KZk8‘

=

2.4 X 10-3 min.-‘

(12)

where k8’ is the catalytic coefficient for hydronium ion acting as a general acid, and the numerical value corresponds to that previously r e ~ o r d e d . ~ At pH >4 eq. 10 yields zlo~-[H+]= K 2 k 8 ” [ H 2 0 ]= 1 7 X

min.-’ M

(13)

where kg“ is the catalytic coefficientfor water acting as a general acid and the nunierical value is taken from the unit slope a t pH >4 of the logarithm of the observed first-order rate constant us. pH This mechanism portrayed in ternis of general acid catalysis of free amine is kinetically indistinguishable from general base catalysis of protonated aniine as the data were described in the Results section. For the situation where low concentrations of formate buffer are added in the plateau region froin 3 < pH < 4 of the pH-rate profile, application of eq. 10 yields ZF =

Kz(h.3‘

+ ks”’[HF]/[H+])

where [ H F ] represents the molar concentration of the conjugate acid form of formic acid and I