2914
HARRY L. CONLEY,JR.,AND R. BRUCEMARTIN
Cupric Ion Catalyzed Hydrolyses of Glycine Ethyl Ester, Glycinamide, and Picolinamide'
by Harry L. Conley, Jr., and R. Bruce Martin Cobb Chemical Laboratory, University of Virginia, Charlottessille, Virginia
(Received February 16, 1966)
Attack of hydroxide ion on protonated glycine ethyl ester at 25.00" and 0.16 ionic strength proceeds 41 times faster than the same reaction with neutral ester. Reaction of hydroxide ion with the cupric complex of glycine ethyl ester provides a further increment of 3200 over hydroxide ion attack on protonated ester, furnishing an example of super-acid catalysis due to chelation near the reactive site. A zinc-mercaptoethylamine complex catalyzes hydrolysis of glycine ethyl ester. Cupric ion also promotes the hydrolysis of glycinamide near neutral pH, but less effectively than in the case of glycine ethyl ester. Cupric ion inhibits the hydrolyses of both glycinamide and picolinamide a t high pH, where ionization of amide hydrogens has taken place. Interpretations of these kinetic observations are made on the basis of equilibrium complexation of cupric ion with reactants and tetrahedral carbon addition intermediates. Results of potentiometric and spectrophotometric studies of cupric and nickel ion complexes of picolinamide are described. A 2: 1 picolinamidenickel ion complex yielding red needles in the solid state has been prepared.
Though it is generally recognized that metal ions are less effective in catalyzing amide than ester hydrolyses, direct comparisons of specific rate constants on analogous compounds of well-defined reacting species have not been possible. An early study of met,al ion activated hydrolysis of amino acid esters2 is compromised by coordination with the metal ions of t,he basic component of the buffer system and has been criticized on other grounds as well.3 Further work on cupric ion promoted hydrolysis of glycine ethyl ester in the presence of glycine buffer, a product of the reaction, was performed a t only one P H . ~ This latter work did establish however, that carbonyl oxygen exchange occurs in the cupric ion hydrolysis, as it does in ordinary ester hydrolysis. Since it was not possible to calculate the concentration of reacting species, only observed rate constants are given. Both of these studies followed the reaction by classical titration methods. A conductivity method has also been used to follow the hydrolysis of glycine ethyl ester in the presence of cupric and nickel ions and excess base. Formation constants were reported, but no variation in hydroxide concentration was a t t e m ~ t e d . ~ Cupric ion catalyzed hydrolysis of glycinamide has been studied as a funcThe Journal of Phpsical Chemistry
tion of pH and temperature by a spectrophotometric method in the presence of high concentrations of carbonate buffers.5 The work reported here aimed to determine specific rate constants for cupric ion catalyzed hydrolyses in terms of well-defined complexes of glycine ethyl ester and glycinamide. Reaction rates were followed on a pH-stat so that buffer-free solutions could be used. Hydroxide ion dependences of the rates with specific complexes are reported. Comparisons are then made among free ester and cupric ion catalyzed hydrolyses and with amide hydrolyses. Similar studies on other amino acid esters are reported elsewheree6 Some of the features of glycinamide-metal ion com(1) This paper is abstracted from the Ph.D. thesis of H. L. Conley, Jr., 1964, from which more details may be obtained. The research
was supported by grants from the National Science Foundation and the National Institutes of Health. (2) H. Kroll, J. Am. Chem. Soc., 74, 2036 (1952). (3) M. L. Bender and B. W. Turnquest, ibid., 79, 1889 (1957). (4) J. M. White, R. A. Manning, and N. C. Li, ibid., 78, 2367 (1956). (5) L. Meriwether and F. H. Westheimer, %%id., 78, 5119 (1956). (6) H. L. Conley, Jr., and R. B. Martin, J . Phys. Chem. 69, 2923 (1965).
CUPRICIONCATALYZED HYDROLYSIS OF GLYCINE ETHYLESTER
plexes appear at relatively high pH values. For this reason investigations were performed on picolinamide with a considerably lower pK,. Picolinamide is also a derivative of an a-amino acid.
Experimental Materials. The glycine ethyl ester hydrochloride and glycinamide hydrochloride used were the best commercial products available. Equivalent weights determined by titration were within 2% of the theoretical values. Preparation and standardization of sodium hydroxide and metal ion solutions are described elsewhere.6 Picolinamide was prepared by the reaction of aqueous ammonia with picolinic acid ethyl ester7 and reprecipitated from alcohol solution by addition of benzene, m.p. 104-106'; lit.7105'. Anal. Calcd. for CeHcNzO: C, 59.01; H, 4.95; N, 22.94. Found: C, 59.17; H, 4.89; N, 22.97. Automatic titration of a 0.01 M solution on the pH-stat at 25.00' and 0.16 ionic strength gave pK, = 1.9. This value is subject to uncertainties involved in converting hydrogen ion activities to concentrations in this low pH region, but is in agreement with a values of pK, = 2.1 found at the lower temperature of 20' and ionic strength 0.01. On several occasions high pH solutions of picolinamide and cupric ion produced purple needles while nickel ion gave red needles. The red needles were prepared in larger quantity by mixing amide and nickel ion at concentrations of 0.3 and 0.1 M , respectively. Upon the raising of pH of this solution to about 11.5, red needles began to form. The red needles were collected, washed three times with water, and stored in a desiccator containing Pz06. Results to be presented later suggest that the red needles should be a 2 : 1 picolinamide-nickel ion complex with amide hydrogens ionized on both ligands. Anal. Calcd. for Ni(C6H5N20)2:C, 47.89; H, 3.35; N, 18.62; Nil 19.5. Found: C, 47.63; H, 3.42; N, 18.92; Ni, 19.0 and 19.4 by gravimetric dimethylglyoxime m e t h ~ d . ~This agreement provides strong evidence for amide hydrogen ionization because the analysis cannot accommodate the anions that would otherwise be required for charge neutralization. Methods. Most rate runs were performed on a Radiometer pH-stat equipped with a glass jacket about the reaction vessel which permitted temperature control to *0.02'. Experiments in this paper were performed at 25.00, 30.00, or 35.00 f 0.02'. Other features of the experimental set up are described elsewhere.lV6 Solutions of metal ion nitrates were used in these experiments. Ionic strength was always 0.16 controlled with KN03.
2915
Some rate runs were performed spectrophotometrically on a Cary 14 spectrophotometer with metal ion chlorides and ionic strength controlled to 0.16 with KC1. The temperature was 25.0 f 0.1'. Whether rates were determined on the pH-stat or spectrophotometrically, pH values were read on a pH meter. Hydroxide ion activities were obtained from pH meter readings by utilizing -log K , = 14.00, 13.83, and 13.68 at 25, 30, and 35', respectively.'O Our second-order rate constants are obtained by dividing first-order rate constants by the activity of hydroxide ion and possess units of activity-l see.-'. These rate constants may be converted t,o M-l sec.-l by multiplying them by 0.75, the mean ion activity coefficient for both KOH in KC1 and NaOH in NaCl solution at 25' and 0.16 ionic strength.1° In this paper parentheses denote molar concentrations, and brackets denote molar activities. The pH-stat automatically portrays on a graph the amount of a given concentration of base or acid added from a syringe buret to maintain a constant pH vs. time. By drawing a slope through the initial portion of the rate curve, the initial molar concentration per second, Ratat,of protons liberated or consumed in the reaction may be obtained, upon the appropriate manipulation of units. This initial rate of change of protons is not necessarily equivalent to the initial rate of change of molar concentration of substrate per second, the rate we wish to determine. The latter kinetically important rate can be determined from the former experimentally observed rate by a general method outlined for glycine ethyl ester. The experimentally observed rate of molar proton production, Retat,is related to the molar rate of ester hydrolysis, RE, by Ratst = &RE. We proceed to determine the quantity Q for the substrates studied in this work. Glycine ethyl ester may exist in protonated, E H f , or free base form, E, and the product glycine in the dipolar ion AH* or anionic form A- at pH >4. Without consideration of mechanism and merely as a material balance for protons, the hydrolysis of glycine ethyl ester may be written as EHf Hf
+ E
+ HzO +AH* + H f + CzH50H H+
+
A-
(7) R.Camps, Arch. Pham., 240, 347 (1902). (8) H.H.G.Jellinek and J. R. Urwin, J.Phys. Chem., 58,548 (1954). (9) R.Beloher and A. J. Nutten, "Quantitative Inorganic Analysis," Butterworth and Co., Ltd., London, 1955.
Volum 69, Number 9
September 1966
2916
HARRY L. CONLEY, JR.,AND R. BRUCE MARTIN
where KE and Kz are acid ionization constants. The fraction of the total glycine ethyl ester concentration in the free base form is given by aE = &/([H+] KE), and the fraction of the total concentration of glycine in the free base form is given by a2 = K2/( [H+] Kz). For each molecule of protonated glycine ethyl ester, EH+, that is hydrolyzed, one proton is produced for the fraction of product glycine that is AH*, and two protons are produced for the glycine fraction that is A-. The first fraction is given by 1 - a2 and the second by a2. The observed rate will be reduced by the fraction of ester that is E instead of EH+. Totaling these results yields
+ H2O -%products E H + + OH- " ,products E + H2O 3 products RQ E + OHproducts EH+
+ +
Retat
= RE[(^
-
a2)
+ 2ff2- LYE] = RE[1
+
a2
(2)
-
(3) (4)
The symbols EH+ and E represent protonated and free base ester, respectively. To the rate constants for water attack is affixed a subscript w, while the superscripts refer to the charge on the substrate ester molecule. Since all experiments were performed in dilute aqueous solutions, the constant concentration of water is included in the rate constants k,.
- LYE]
The quantity in brackets is Q for glycine ethyl ester hydrolysis. This same quantity can be derived by other arguments. Only when a2 = a~ is there a 1:1 correspondence between moles of ester hydrolyzed and moles of protons liberated. I n order to determine R E from RBtat,a value of pK2 = 9.66 =k 0.0411112 was used. By potentiometric titration we have determined pKE = 7.79 f 0.02 at 25.0' and 0.16 ionic strength. A valuela of 7.73 determined at 25' and 0.05 ionic strength is nearly identical when the ionic strength difference is taken into account. For glycinamide Q = a2 a3 - LYA, where the successive a-values, respectively, represent the fractions of glycine, ammonia, and glycinamide in the basic forms. Since data were obtained at three temperatures, the acid ionization constants for each of the three compounds must be known at each temperature. Best values of acid ionization constants were found at 0.16 ionic strength and 25.00'. These values were calculated for 30.00 and 35.00' using the integrated form of the van't Hoff equation and appropriate values for AH. The assumptions and results are summarized in Table I. The AH value of 12 kcal./mole assumed for glycinamide is not critical for this study because runs were made in the pH range 10.7 to 11.8, where glycinamide is att least 99.7% in the free base form. For picolinamide, the same equations apply as for glycinamide. Since all hydrolysis experiments were performed at pH >10 and pK, = 1.9 for picolinamide and 1.6 and 5.4 for picolinic acid,14CYA. = 1 = a2,
+
Results Glycine Ethyl Ester. Four reactions must be considered in the mechanisms of glycine ester hydrolysis. They involve water and hydroxide ion attack at protonated and free base ester. The Journal of P h g h l Che?niatrp/
(1)
~~
Table I: Values of AH and pK. for Ammonium Ionizations a t 0.16 Ionic Strength AH, kcal./mole
Glycine Ammonia Glycinamide
10.80" 12.40'
12'
25O
P Ka 30'
350
9.66b 9.27d 8.06'
9.53 9.12 7.92
9.40 8.97 7.77
' E. J. King, J. Am. Chem. SOC.,73, 155 (1951). See ref. 11. Average of three literature values: D. H. Everett and D. A. Landsman, Trans. Faraday Soc., 50, 1221 (1954); D. H. Everett and W. F. K. WynneJones, Proc. Roy. SOC.(London), A169,190 (1938); R. G. Bates and G. D. Pinching, J . Res. Nail. Bur. Std., First 42, 419 (1949); J. Am. Chem. Soc., 72, 1393 (1950). N. C. Li and M. C. M. Chen, reference in c. dA8sumed. J. Am. Chem. SOC.,80, 5678 (1958).
'
The over-all rate of glycine ethyl ester hydrolysis is given by RE = k,+(EH+)
+ k+(EH+)[OH-] + kwo(E)+ kO(E)[OH-] ( 5 )
Dividing through eq. 5 by the total ester concentration, (Etot), and letting a represent the fraction of ester in the free base form, a = KE/[KE [H+]], where KE is the acid ionization constant for ester, we have
+
~
(IO) H. S. Hamed,md B. B. Owen, "The Physical Chemistry of Electrolytic Solutions, 3rd Ed., RRinhold Publishing Corp., New York, N. Y., 1958. (11) An average of five literature values measured at essentially the same conditions of this study: R. M. Keefer, J . Am. Chem. Soc., 68, 2339 (1946): E.J. King, fig.,73,155 (1951); C.Tanford and W. S. Shore, ad.,75,816 (1953); F.Basolo and Y. T. Chen, &id., 76,953 (1954); the fifth value is that of Li and Manning's. (12) N.C. Li and R. A. Manning, ibid., 77, 5225 (1955). (13) 0.H.Emerson and P. L.Kirk, J . BWZ. Chem., 87,597 (1930). (14) F.Holmes and W. R. C. Crimmin, J . Chem. SOC.,1177 (1955).
CUPRICION CATALYZED HYDROLYSIS OF GLYCINE ETHYLESTER
RE = k,+(l (Etot)
-
a)
+ (k++ (1
-
2917
Kw
.)[OH-]
+ koa[OH-]
(6)
3.3 Three of the four rate constants of eq. 6 appear in terms of different [OH-] dependencies, but k+ and k,O cannot be separately determined experimentally. 3.5 A titration curve for glycine ethyl ester exhibits two unbuffered regions, pH 4.8 to 6.3 and pH 8.8 to 11.2. Rate runs were performed on the pH-stat in each region at ester concentrations of 0.005, 0.01, and 0.02 M . I n the low pH 4.8-5.8 region, a becomes 1 negligible compared to unity, and the last term in eq. 6 is insignificant. A straight line drawn through 12 points in a plot of &/(Etot) us. [OH-] yields k,+ = 5 f 3 X set.-' from the intercept and k + k w o K ~ / K= w 36 f 2 activity-' sec.-l from the slope. This value of kw+ is in excellent agreement with one sec.-l obtained under similar ~0nditions.l~ of 6 X w Since k,+ is surely not less than kwo, k w o K ~ / Kmust 8.8 92 9.6 la0 10.4 10.8 11.2 be less than k w + K ~ / K< w This last number is PH much less than 36, so that k + must equal 36; therefore, Figure 1. Hydrolysis of glycine ethyl ester. Points are k,OKE/K, is negligible compared to k+ and cannot be determined from our studies. A similar conc~usion experimental and smooth curve calculated from eq. 7 with k+ = 28 activity-' sec-1 and ko = 0.78 activity-' sec.-1. applies to other compounds in this paper. In the high pH 8.8-11.2 region, eq. 6 reduces to
-
-
-
Y
I-
-
+
A plot of the logarithm of the left-hand side us. pH for 19 runs yields the curve shown in Figure 1, which becomes linear above pH 10.3 where CY becomes unity. From the linear region kO = 0.78 f 0.02 activity-l sec.-l. By utilizing this value for ko in eq. 7, values of k+ were calculated for each of the 12 experimental points from 8.8 < pH < 10.3, yielding k+ = 28 f 3 activity-' set.-'. This value of k + agrees fairly well with that determined above in the other unbuffered region where the hydroxide ion concentration is only times as great. An average value of k + = 32 4 activity-1 sec.-l is taken for this rate constant. Experiments performed with solutions M in total (Cu2+)and M in total ester concentration at pH 4.1 to 5.3 yield an estimate of the cupric ion catalyzed glycine ethyl ester hydrolysis. By utilizing a value of 3.87 for the logarithm of the first formation constant of cupric ion with glycine ethyl ester4 and reasonable lesser values of the successive formation constants, it may be shown that greater than 96% of complexed cupric ion exists as 1 : l complex in this pH regi0n.l One proton is liberated for each ester molecule hydrolyzed so that Rstat = RE RCUE. I n
*
+
this pH region eq. 2 is the most important for the free ester so that RE = k+(EH+)[OH-1, all quantities of which are known for a given pH. Two terms may contribute to RCG corresponding to the reactions
+ HzO products CUE+^ + OH- +products so that R C ~ E= ~ I , ~ + ( C U E ~++ )k12+(CuE2+)[OH-]. CUE+^
ki'+
Rearrangement then yields (Est,+,
- RE)/(CUE~+)= k i w 2 +
+ ki2+[OH-]
(8) RE was calculated for each of six runs and was only 1-2% of Rstat. (CuE2+) was 2 and 22% of total (CU+~)at pH 4.1 and 5.3, respectively, with inter-. mediate percentages at intermediate pH values. A plot of the left-hand side of eq. 8 us. [OH-] yields, as shown in Figure 2, a straight line, the intercept of which set.-' while the yields klwZf = (4.3 f 0.5) X slope gives kI2+ = (1.01 f 0.05) X lo5 activity-' set.-'. An additional point determined with both total ester and total cupric ion concentrations equal to 0.01 M falls close to the line, lending support to the treatment of this system. (15) V. I. Bolin, 2.anoru. allgent. C h m . , 143, 201 (1925).
Volume 69,Number 9 September 1966
2918
HARRYL. CONLEY,JR.,A N D R. BRUCEMARTIN
8-mercaptoethylamine and zinc ion each at 0.05 M and from pH 6.8 to 7.8. Since log Ki = 9.9 for zinc c
(Zn(mea) +) complex catalyzed hydrolysis of glycine
glycine ethyl ester from p H 4.1 to 5.3. Total ester concentration is 0.01 M , total cupric ion concentration is 0.001 M in circles and 0.01 M in square.
+ H 2 0 kla_+ products Zn(mea)(E)+ + OH- ki:products Zn(mea) (E)+
We then obtain Cupric ion also catalyzes the hydrolysis of glycine ethyl ester in the pH 8.6 to 9.8 region. The distribution of ester among four cupric complexes complicates the analysis so that no quantitative treatment of the data obtained is included.’ Comparison of the values for water and hydroxide ion attack on glycine ethyl ester, IC,+ and k+, respectively, and on the 1 : l glycine ethyl ester-cupric ion complex, klw2+ and k12+, indicates that at pH >5 hydroxide ion attack is the more important in both cases. Thus water should be replaced by hydroxide ion in a previous formulation3 of cupric ion catalyzed hydrolysis of glycine ethyl ester where the experiments were performed a t pH 7.3. Since hydroxide ion catalysis of hydrolysis of neutral esters extends into regions where pH 13.27 Titration of solutions containing 0.02 M picolinamide and either 0.01 M Zn2+ or Cd2+ precipitates the metal ion hydroxide a t about pH 8, indicating a relatively weak association of these metal ions with picolinamide. Solutions containing two moles of picolinamide per mole of either Cu2+ or Ni2+ exhibit absorptions at lower wave lengths and of greater intensity than solutions of the aqueous metal ions, indicating binding to the nitrogen containing ligand. For example, aqueous cupric ion absorbs maximally near 800 mp with a molar extinction coefficient E 13, while a 2 : 1 picolinamidecupric ion solution exhibits maximum absorption at 700 mp with E 31. Titration of solutions containing two moles of picolinamide per mole of Cu2+ or Ni*+ requires the addition of 2 equiv. of base per mole of metal ion to reach pH 10. The picolinamide-cupric ion solutions reversibly change color from blue at the beginning of the titration to violet a t the end, with an absorption maximum a t 560 mp and E 53. This last color does not fade on prolonged standing. Further binding to nitrogen atoms is indicated by the spectral shift to shorter wave lengths on titration. The original blue picolinamide-nickel ion solutions change to yelloworange during the course of the titration. These color changes are typical of those found in other amides and (26) Evidence for the view that binding occurs through the carbonyl oxygen before and the amide nitrogen after ionization of amide hydrogens has been reviewed: A. S. Brill, R. B. Martin, and R. J. P. Williams in "Electronic Aspects of Biochemistry," B. Pullman, Ed., Academic Press, New York, N. Y., 1964, p. 540. (27) For the amide hydrogen ionization from positively charged NLmethylnicotinamide, pK, = 13.2: R. B. Martin and J. G. Hull, J . Biol. Chem., 239, 1237 (1964).
CUPRICION CATALYZED HYDROLYSIS OF GLYCINE ETHYLESTER
peptides where amide hydrogens are undergoing cupric28or nickel29ion promoted ionizations. Preparation of a solid, neutral nickel(picolinamide)2 complex is described in the Experimental section. The overlapping ionizations that occur when solutions 0.02 or 0.04 M in picolinamide and 0.01 or 0.02 M, respectively, in either Cu2+ or Ni2+ are titrated with base were analyzed by the projection strip method.30 For the 2 : l cupric ion complex, pK1 = 5.09 and pK2 = 6.51, while for the 2 : l nickel ion complex, pK1 = 7.99 and pK2 = 9.44. These values may be considered reliable to *0.05 log unit. The difference between pK2 and pK1 for the cupric ion complex is similar to that observed for cupric ion promoted amide hydrogen ionizations from other amidesa1and peptides.28 The 2 : 1 picolinamide-nickel ion complex exhibits a difference of 1.45 log units between the two amide hydrogen ionizations. I n other instances of nickel ion promoted amide hydrogen ionizations, the difference in pK, values is 0.6 unit or less.2e These small differences are equal to or less than those based on purely statistical grounds. The difference of about 1.45 units for the 2 :1 picolinamide-nickel ion complex is larger than the statistical difference of 0.6 unit but still much less than that expected for a dibasic acid with the acidic groups so closely spaced. The cooperative nature of the transition from the blue, paramagnetic, octahedral nickel complex to the yellow, diamagnetic, planar complex upon ionization of the amide hydrogens is again indicated. Evidence has been presented that cupric ion promotes amide hydrogen ionization in picolinamide as well as glycinamide. For both amides, hydrolysis at pH 11 to 12 is inhibited by the addition of cupric ion. On the other hand, cupric ion promotes the hydrolysis of glycinamide at pH
