Charge-Transfer Association in Enzyme Inhibition Citation Roth, Dana Lincoln (1965) Charge-Transfer Association in Enzyme Inhibition. Master's thesis, California Institute of Technology. doi:10.7907/DPPB-6X75. https://resolver.caltech.edu/CaltechTHESIS:05052010-095701928 Abstract Attempts to determine if charge-transfer bonding forces are involved in enzyme inhibition by gas chromatographic and spectral methods were, in general, unsuccessful. Direct spectral evidence of formation of the β-naphthoquino-linium cation in the presence of the enzyme, α-chymotrypsin, has been shown. Item Type: Thesis (Master's thesis) Subject Keywords: (Chemistry) Degree Grantor: California Institute of Technology Division: Chemistry and Chemical Engineering Major Option: Chemistry Thesis Availability: Public (worldwide access) Research Advisor(s): Niemann, Carl G. Thesis Committee: Unknown, Unknown Defense Date: 1 January 1965 Record Number: CaltechTHESIS:05052010-095701928 Persistent URL: https://resolver.caltech.edu/CaltechTHESIS:05052010-095701928 DOI: 10.7907/DPPB-6X75 Default Usage Policy: No commercial reproduction, distribution, display or performance rights in this work are provided. ID Code: 5772 Collection: CaltechTHESIS Deposited By: Tony Diaz Deposited On: 09 Aug 2010 22:37 Last Modified: 09 Feb 2024 18:39 Thesis Files Preview PDF - Final Version See Usage Policy. 1MB Repository Staff Only: item control page

CHARGE-TRANSFER ASSOCIATION IN ENZYME INHIBITION Thesis by Dana Lincoln Roth In Partial Fulfillment of the Requirements For the Degree of Master of Science in Chemistry California Institute of Technology Pasadena, California 1965 ii DEDICATION To my devoted wife Eileen, whose love and patience lighted the way. iii ACKNOWLEDGMENTS To the late Dr. Carl Niemann, whose untiring dedication to the science of chemistry provided the initial inspiration for this work and who unknowingly awakened my desire for the broader career of librarianship. To Gary Neil, Sang Chul Shim, Leonard Piszkiewicz, Roger Peterson, Al Mukatis and Carole Hamilton for their meaningful contributions to my education. To my parents for their patience and perseverance during my preparative schooling. To the California Institute of Technology for financial support. iv CHARGE-TRANSFER ASSOCIATION IN ENZYME INHIBITION By Dana Lincoln Roth Abstract Attempts to determine if charge-transfer bonding forces are involved in enzyme inhibition by gas chromatographic and spectral methods were, in general, unsuccessful. Direct spectral evidence of formation of the B-naphthoquino- linium cation in the presence of the enzyme, a@-chymotrypsin, has been shown. PART IT. IIl, Iv. TABLE OF CONTENTS INTRODUCTION GAS CHROMATOGRAPHY Results and Discussion Conclusions Experimental SPECTRAL STUDIES Results and Discussion Conclusions Enzymatic Studies Experimental REFERENCES eo © © #8 PAGE 15 15 20 21 23 46 vi LIST OF TABLES TABLE PAGE I. Gas Chromatographic Separation Factors for Equimolar o,m and p-Xylene Mixtures. .... 8 Il. Solvent Effects on the Ultraviolet Spectrum of Pyromellitic Dianhydride ..... 1... eee. 15 Il. Absorption Maxima of Various Aza-~Aromatic TCNE Charge-transfer Complexes ....... 17 IV. Equilibrium Constants and Extinction Coef- ficients of TCNE-Pyridine and TCNE- Isoquinoline Charge-transfer Complexes 18 Vv. Concentration and Spectral Data for the TCNE- Isoquinoline Charge-transfer Complex ..... 18 Vi. Enzyme~-inhibitor and Charge-transfer complex Dissociation Constants for Pyridine and Isoquinoline and the Ratios of the Corre- sponding Free Energies . 1... -- ee ee eee 20 Vil. Solvent Dependence of the Ultraviolet Ab- sorption of B-Naphthoquinoline ......... 21 INTRODUCTION For over a century it has been recognized that a variety of aromatic compounds can undergo additive combination with acidic aromatic compounds such as: picric acid (1), polynitroaromatics (2, 3,4) and quinones (5), to produce stable, isolable, molecular compounds. The extensive work of Sudborough (2) on the addition compounds of picric acid led to their use in the characterization of polynuclear aromatic hydrocarbons and aromatic ethers (6). Moreover, work with many substituted and unsubstituted aromatic compounds showed that the aromatic nucleus of the donor, or basic addend, was the domi- nant factor in complex formation (2). The early work of Hildebrand on the colors of iodine solutions led to discovery of non-isolable complexes formed by the interaction of iodine with solvent (7). Studies of the ultraviolet spectra of solutions of these complexes have been conducted since 1925 (8). In each case a spectrum resulted which was similar enough to that obtained by adding the spectra of the components that it was erroneously concluded that the complex was completely dissociated in solution. Although the color of the complex was invariably different from those of the separate components and corresponded to longer wavelength absorption, no systematic study of the visible spectra was undertaken. In 1949, the classic paper by Benesi and Hildebrand (9) on the interaction of iodine and aromatic hydrocarbons appeared. Contrary to prior spectrophotometric investigations a strong new ultraviolet absorption band was found in the spectra of iodine solutions of aro- matic compounds which was absent in those of iodine-carbon tetra- chloride solutions. Benesi and Hildebrand attributed the anomalous absorption to complex formation and were able to determine the equi- librium constants and extinction coefficients for complexes of iodine with some of the Cy, - Cy aromatic hydrocarbons. Assuming that the complex was formed by direct interaction of the acidic and basic components and that it existed in equilibrium with them, Benesi and Hildebrand proposed the following analysis of the observed spectra: A+B =~ AB (1) [AB] [A, - ABJ[B, - AB] where AB = concentration of complex at equilibrium & Hn initial concentration of acidic component as) o it initial concentration of basic component Experimentally, the initial concentration of B was much larger than that of A, so that [By - AB] = By. With the assumption that the ab- sorption followed Beer's Law, D = Xb [AB] (3) where D = optical density = molar extinction coefficient b = optical path length equation 4 may be derived, Agb ] 1 = + 4 D K2Bo a (4) A plot of Agb/D vs 1/Bg would give a linear relationship with a slope of 1/K> and an intercept of 1/2. Later, Scott (10) proposed a modi- fication to increase the accuracy of measurements at low concen- trations, by rearrangement of equation 4 to give 5. PoP ob - Bo , 1 (5) A plot of AgBob/D vs Bo offers the advantage of extrapolation through regions of low concentration. Since this time, numerous spectrophotometric investigations have appeared and have been tabulated in an excellent review by Brieglieb (11). The theoretical basis for understanding the spectra of these complexes was advanced by Mulliken (12). In general agreement with Benesi and Hildebrand, Mulliken suggested that the complexes result from "an acid-base interaction, in the electron-donor sense, in which the iodine functions as the acid or electron acceptor.'' Mulliken con- sidered the interaction of a no-bond ground state |, (A, D) and a polar excited state, (D*, A ), to produce a stabilized ground state having a wave function ¥, = aby (D,A) + buy (D', A’) (6) and an excited charge-transfer state with the wave function + _ Yop = ay (DT, AD) + by (D, A) (7) The anomalous ultraviolet or visible absorption is associated with the electronic transition be > Hot . The bonding in the ground state is the sum of the van der Waals forces and the resonance energy of inter- action between J and |,, while the bonding in the charge-transfer state is mainly due to the large electrostatic energy of interaction between D* and A. A good approximation (13) to the energy of the charge-transfer transition is given by hv = 1,7 Ea + © (8) where In = ionization potential of the donor EA = electron affinity of the acceptor = the mutual electrostatic energy of D’, A relative to D,A The approximation involved is that the resonance energy of interaction between % and ¥,, which increases the energy of the charge-transfer transition, is minimal. The validity of this assumption has been demonstrated by McConnell and Dewar (14, 15, 16) for some aromatic systems which are described as weak or mt complexes. As was later pointed out by Dewar (17), since the heats of for- mation of these complexes are at least an order of magnitude lower than their lowest transition energies, the changes in energy of the orbitals in forming the complex are small. Thus, the ultraviolet spectra of the components of the complex approximate those of the separate components. The appearance of a new band of lower energy corresponds to the excitation of an electron from a filled orbital of the donor to a anti-bonding orbital of the acceptor. Since the orbital energy is a function of the basicity of a compound, the anti-bonding orbitals of the acceptor will be associated with lower energy values than those of the donor and the charge-transfer transition will occur at longer wavelengths than intramolecular transitions. An example of strong complexes is given by the interaction of amines and iodine (18). Here, the non-bonding electrons of the amine nitrogen are involved, and the resulting close approach causes a large ground-state interaction between Wp (A, D) and 4, (A , D*), which pre- vents simplified analysis. Strong complex formation is also possible with polynuclear aromatic donors which possess an active ''K region" which is not accompanied by a strong "'L region" (19). The existence of charge-transfer association in biological systems has been shown by Kosower in work on solutions of N-alkyl- pyridinium halides (20). Later work, with DPN and TPN and with the model compound 1-benzyl-3-carboxamidopyridinium chloride, revealed that charge- transfer complex formation occurred with indole derivatives and chymotrypsinogen, the precursor of a-chymotrypsin (21). Studies on the nature of enzymatic inhibition have shown, in the case of a-chymotrypsin, that aromatic compounds are more effective, by an order of magnitude, than their saturated analogs in binding with the enzyme (22). The research discussed in this thesis is predicated on the assumption that charge-transfer bonding forces are involved in enzyme inhibition. The strong acceptor, pyromellitic dianhydride (PMDA), was initially selected as a model for the active site of the enzyme. Comparison of the charge-transfer equilibrium constants of the known inhibitors and the model compound with their enzyme inhibition constants could show if charge-transfer bonding forces play a dominant role in the inhibition of a-chymotrypsin. The determi- nation of charge-transfer equilibrium constants was pursued by two distinct methods. The choice of the Cg - Cg aromatic compounds for the gas chromatographic method of Muhs and Weiss (23) and the un- substituted aza-aromatic compounds for the spectroscopic method of Benesi and Hildebrand (9) was necessitated because of the suitability of analyzing sterically similar compounds and the undesirable experi- mental conditions necessary for gas chromatographic analysis of the latter class of compounds. GAS CHROMATOGRAPHY Results and Discussion As a prelude to the determination of equilibrium constants by gas chromatography, the separation of the isomeric Cg aromatic compounds was required for the requisite experimental refinements. This separation was attempted with four systems: A) squalane - PMDA Fig. la B) diglycerol - PMDA Fig. lb C) a-chloronaphthalene - PMDA Fig. lc D) a-chloronaphthalene Fig. 1d Table I gives the relative separation factor (33), which is defined as follows, for each of the four systems. S,02- (9) T refers to the retention time of a specific compound or mixture and T,>T). In system A, as the column aged, the absolute retention times and the separation factor decreased, which indicates that the added PMDA has no effect. The lack of effect may be ascribed to the lack of solubility in squalane. The very low retention times observed in system B are no doubt due to the limited solubility of the aromatic compounds in diglycerol. The general shape of the recorder response curve is that invariably TABLE I. Gas chromatographic separation factors for equimolar o,m and p-xylene mixtures. System T, Ty Siz A m-xylene p-xylene 0 A o-xylene m, p-xylene 0.176 B m-xylene p-xylene 0 B o-xylene m, p-xylene 0 Cc m-xylene p-xylene 0. 066 D m-xylene p-xylene 0.059 found for gas-solid systems, where the separation is due to adsorptive rather than absorptive processes. The use of a-chloronaphthalene as a liquid support has been previously described by Zlatkis (24) and work in this laboratory con- firms his findings. The addition of PMDA to the a-chloronaphthalene support results in an increased value of the separation factor for m- and p~xylene. The limited effect of the added PMDA is due to charge- transfer complex formation between a~-chloronaphthalene and PMDA, which necessarily dilutes the effect of the added PMDA. The experi- ments using @-chloronaphthalene were further complicated by thermal effects. A column temperature of 50°C was required for reasonable retention times and at this temperature the a-chloronaphthalene was sufficiently volatile to prevent reproducible results. The procedure for determination of charge-transfer equilibrium constants (23) re- quires experimental conditions which remain relatively constant during extensive usage. Conclusions The attempt to determine charge-transfer equilibrium constants by gas chromatography is not feasible for high molecular weight compounds. The experimental necessity of using a liquid support in which the t-acid and aromatic mixture are soluble limits the choice to a m—-base which necessarily dilutes the acid strength. The use of lower molecular weight t-acids may obviate this problem. 10 Experimental Pyromellitic Dianhydride was obtained from du Pont and purified by sublimation at 170° and 704% Hg. The infrared spectrum of the purified material, which melts at 286°, is shown in Fig. 5, The resublimed PMDA develops a yellow color if the sublimation temperature rises above 175°. Hydrolysis of the purified material followed by titration to the phenolphthalein end point gave a molecular weight of 254. 8 compared to the theoretical value of 254.2. Squalane, Diglycerol and a-Chloronaphthalene were obtained commercially and used without further purification. o, m and p-Xylene were Eastman reagent grade and were used without further purification. Chromosorb (80-100 mesh) was obtained from Johns—Mansville and dried at 110° for four hours prior to use. The squalane and diglycerol columns were prepared by the usual technique of mixing with a easily evaporable solvent. The a@- chloronaphthalene was added to the solid support without solvent and mixed by rotation for one week. An Aerograph A-600A chromatographic system was employed using a hydrogen-flame detector and a Hamilton microliter syringe for sample introduction. Fig. 1. la. lb, Ic, ld. Fig. 2. 11 LEGEND TO FIGURES Gas Chromatograms of Equimolar Xylene Mixtures 10% PMDA, 20% Squalane w/w 80-100 Chromosorb 5% PMDA, 10% Diglycerol w/w 80-100 Chromosorb 1.8% PMDA, 15% a@-Chloronaphthalene w/w 80-100 Chromosorb 15% a-Chloronaphthalene w/w 80-100 Chromosorb Infrared Spectrum of Resublimed PMDA (nujol) 12 xylene mixture mown |. { m, p-xylene 840 ' Retention Time-~Seconds 600 _ Response _ _ Recorder TS asuocsey Hen Nae ene DP lh at A Fete aoe auazusqTAéyya q. _ et saat C oreall eueTAx=u q ~, ™~ \ | ec Lapsooey Yt oe) “ a] te oe auazuagTsyyo C we Se . » . aa mR tial ree ae 45 Retention Time-'E nutes a eens | vavelength-—Jicrons 8.0 st Jone cee nena ceeeiee peee Wir ! | F 80 60 40 20 cece sa atetnae emcee ap % Transmittance 15 SPECTRAL STUDIES Results and Discussion Prior to studies investigating charge-transfer complexes of PMDA, the spectrum of PMDA was examined ina variety of solvents. The long wavelength absorption spectrum appeared as a doublet in chlorinated hydrocarbon solvents (Fig. 3) and as a shoulder in oxygenated solvents (Fig. 4). The effects of solvent on the long wavelength ultraviolet absorption of PMDA are shown in Table II. TABLE II. Solvent effects on the ultraviolet spectrum of pyro- mellitic dianhydride. Solvent max max a ax Figure 1, 2 dichloroethane 316 306 3a CH,Cl, 316 304 3330 3b CCl, 312 301 3d CHCl, 307 299 3c THF 306 2500 4 DMF 295 2460 4 H,O 295 2310 4 At shorter wavelengths, PMDA exhibits very strong absorption with a maximum at 257 mp in CH,Cl,. This absorption may be as- cribed toa 7 > 17* transition (25). Analysis of the long wavelength absorption is complicated by the multitude of possible electronic ex- citations. 16 The work of Birnbaum (26) on intramolecular charge-transfer and n— 7* transitions in ketones indicate that absorption due to these effects will occur in the observed range. The spectra in non-chlorin-— ated solvents may additionally be complicated by intermolecular charge-transfer interaction with solvent. In general, the long wavelength, low intensity, absorption of carbonyl groups is due to n> 7* transitions, while in anhydride groups the n> o* transitions are possible for singly bonded oxygen atoms. The solvent effects on the transition are directly related to the hydrogen bonding ability of the solvent (25). As shown in Table II, the energy required for electronic ex- citation in PMDA solutions increases with a corresponding increase in hydrogen bonding ability. The anomalous position of the spectral bands in CCl, may be ascribed to the fact that this compound is devoid of hydrogen atoms and thus is not a member of this quasi-homologous series of hydrocarbons. Solutions of PMDA-isoquinoline gave anomalous absorption bands at 331 mp in CH,Cl, and at 326 mp in DMF (Fig. 5). Attempts to determine if this absorption was due to a charge-transfer tran- sition were prevented by the long wavelength absorption of isoquinoline which has a maximum at 319 mp in CH,Cl, (Fig. 6). These solutions had no visible coloration. Owing to the unsuitability of PMDA as an acceptor in the study of the aza-aromatic inhibitors which may in part be ascribed to its absorption above 315 mu and the additional possibility that charge- 17 transfer in these basic compounds involves the non-bonding nitrogen electrons, a new model acceptor was chosen. Tetracyanoethylene (TCNE), which has two electron-deficient carbon atoms, is suscepti- ble to charge-transfer complex formation with donors having loosely held non-bonding electrons. The ultraviolet spectrum of TCNE in CH,Cl, shows a very strong doublet with maxima at 277 mp and 267 mp and extinction coef- ficients of 13,400 and 14, 900 respectively (Fig. 7). Solutions of TCNE-isoquinoline, TCNE-quinoline and TCNE- B-naphthoquinoline in CH,Cl, solvent were prepared and analyzed by the method of Benesi and Hildebrand (9) for charge-transfer complex formation. The results are shown in Table III. TABLE III. Absorption maxima of various aza-~aromatic TCNE charge-transfer complexes. Donor Figure ax max max Color isoquinoline 8a 418 398 yellow quinoline 8b 510 417 398 orange B-naphthoquinoline 8c 508 417 orange pyridine (27) 422 400 yellow The 422 mp absorption of the TCNE-pyridine solution follows a Benesi-Hildebrand relationship (27), as does the 418 mp absorption of the TCNE-isoquinoline solution. The solutions of TCNE and quinoline and B-naphthoquinoline do not reveal a Benesi~Hildebrand relationship. 18 The equilibrium constants and extinction coefficients are given in Table IV. TABLE IV. Equilibrium constants and extinction coefficients of TCNE- pyridine and TCNE-isoquinoline charge-transfer com- plexes. Donor ax Keg max T,°c pyridine (27) 422, 12.0 10, 500 22° isoquinoline 418 512.0 6, 360 25° The Benesi-Hildebrand plot for the TCNE-isoquinoline solution is shown in Fig. 13 and the data is given in Table V. TABLE V. Concentration and spectral data for the TCNE-isoquinoline charge-transfer complex. [ TCNE] isoquinoline Optical Density 8.96 x 1075 1,43 x 1074 0.165 1.45 x 1074 12.62 X 1074 1.035 9.37 x 1075 24, 84 x 1074 0.580 8.74 X 1075 50.96 x 1072 0.565 The anomaly of decreasing extinction coefficients with in— discussed by Murrell and McConnell (14-28). creasing equilibrium constants has been reported frequently and is 19 The appearance of a new band in the spectrum of TCNE- quinoline that was not present in the TCNE-isoquinoline spectrum is difficult to explain. Pullman and Pullman (29) report that the differ- ence in the energies of the highest occupied molecular orbitals of the aromatic compounds is only 0.045 8. Moreover, Krishna and Chowdhury (30) in studies on the charge-transfer spectra of iodine and the unsubstituted aza-benzenes, found that the hnax varied by only 11 mp, which would indicate a narrow range of ionization ener- gies for the non-bonding nitrogen electrons. An estimate for the maximum absorption of the n-7 complexes of TCNE-pyridine can be obtained by the method of Batley and Lyons (31); it occurs at 291 mp. Moreoever, the absorption maximum of the TCNE radical anion in tetrahydrofuran has been reported at 430 mp suggesting that this species may be responsible for the 510 mp ab- sorption reported above. 20 Conclusions The determination of the charge-transfer equilibrium constants for a series of aza~aromatic inhibitors by the method of Benesi- Hildebrand using the very strong inhibitors, PMDA and TCNE, is not feasible due to abnormal side reactions. The further study of the complexes which absorb at 510 my might be possible by investigation of their EPR absorption. The significance of charge-transfer bonding in enzyme inhi- bition is evidenced by the similarity of the ratios of the free energies of dissociation of the pyridine- and isoquinoline-~-TCNE charge-transfer complexes to the ratio of the free energies of the corresponding enzyme~inhibitor complexes (22). TABLE VI. Enzyme-inhibitor and charge-transfer complex dissoci- ation constants for pyridine and isoquinoline and the ratios of the corresponding free energies. Compound Ket Ky | (AF/AF) 44 (AF / AF); pyridine 0. 083 27.0 1, 63 1,94 isoquinoline 0.0019 0, 32 21 Enzymatic Studies Following the work of P. R. Hammond (32) on the acridine- acridinium cation charge-transfer complex, in which it was shown that the acridinium cation had an acceptor strength comparable to TCNE, the possibility that, in the special case of interaction with aza- aromatic inhibitors, a-chymotrypsin was in fact acting as the donor in ‘forming a charge-transfer complex was considered. Solutions of a-~chymotrypsin with and without added B-naphtho- quinoline, a very effective inhibitor (22), were prepared using water 0.1 Nin NaCl as the solvent. The spectrum of a-chmyotrypsin showed a shoulder at 290 mu and a peak at 282 mp (Fig. 16). The solvent dependence of the ultraviolet absorption of B- naphthoquinoline is given in Table VII. TABLE VII. Solvent dependence of the ultraviolet absorption of B- naphthoquinoline. Solvent \max Amax Figure H,0 345 330 lla CCl, 346 328 llb The difference spectrum of B-naphthoquinoline and a-chymo- trypsin at pH 7 shows absorption from both S-naphthoquinoline and the B-naphthoquinolinium cation (Fig. 20). Since B-naphthoquinoline does 22 not exist in the protonated form at pH 7 the results here indicate that the enzyme induces formation of significant quantities of the cation. 23 Experimental Pyromellitic Dianhydride was obtained commercially and purified as previously described. Tetracyanoethylene was obtained from Eastman Kodak Company and purified by sublimation at 100° (70 » Hg) or by preparation of the charge-transfer complex with chlorobenzene and subsequent removal of the chlorobenzene at low pressure. Quinoline was obtained from Eastman and redistilled. The material collected boiled at 34-35° (70 » Hg). The ultraviolet spectrum is shown in Fig. 13. Isoquinoline was obtained from Eastman and redistilled. The material collected boiled at 45° (80 » Hg). The ultraviolet spectrum is shown in Fig. 6. B-Naphthoquinoline was obtained from K & K Laboratories and was recrystallized from 0.1 N aqueous acetic acid. Methylene Chloride was Eastman spectroscopic grade, as were the other solvents. The spectra were recorded either on a Cary model 11 or model 14 spectrophotometer. Matched 10-mm silica cells were used. The solutions were prepared in volumetric flasks. The liquids were weighted into the flasks on a Mettler balance and the solids were weighed on an analytical balance. Fig. Fig. Fig. Fig. Fig. Fig. 3. 3a. 3b. 3c. 3d. 8a. 8b. 8c. 10. 24 LEGEND TO FIGURES Ultraviolet Spectra of Pyromellitic Dianhydride in Chlorinated Hydrocarbon Solvents. PMDA (saturated) -— 1,2 Dichloromethane PMDA (3.45 X 1074 M) - Methylene chloride PMDA (saturated) - Chloroform PMDA (saturated) ~- Carbon tetrachloride Ultraviolet Spectra of Pyromellitic Dianhydride in Oxygen- ated Solvents. PMDA (6.28 X 1074 M) - H,O PMDA (5.9 Xx 107*M) - DMF PMDA (6.8 X 107*M) - THF Ultraviolet Spectra of PMDA~-Isoquinoline Solutions in Methylene Chloride and Dimethylformamide. PMDA (2.57 X 1074 M) Isoquinoline (1.06 X 107! M) - CH,Cl, PMDA (5.1 X 107+M) Isoquinoline (1.13 X 1077 M) - DMF Ultraviolet Spectrum of Isoquinoline (10~* M) in Methylene Chloride Solvent. Ultraviolet Spectrum of Tetracyanoethylene (7.46 X 1075 M) in Methylene Chloride Solvent. Ultraviolet Spectra of TCNE-Aza Aromatic Hydrocarbon Solutions in Methylene Chloride Solvent. TCNE (9.4 X 1075 M) Isoquinoline (6.4 X 107? M) - CH,Cl, TCNE (9.0 X 1075 M) Quinoline (3.1 X 1074 M) - CH,Cl, TCNE (1.2 X 1074 M) B-Naphthoquinoline (107 M) - CH,Cl, Benesi~Hildebrand Plot for the TCNE-Isoquinoline Complex. Ultraviolet Spectrum of a-Chymotrypsin in 0.1 N NaCl (H,0) Solvent. Fig. Fig. Fig. ll. lla. lib. lie. 12. 13. 25 Ultraviolet Spectra of B-Naphthoquinoline in Various Solvents. B-Naphthoquinoline (1. 16 X 1074 M) - 0.1 N NaCl (H,0) B-Naphthoquinoline (2.65 X 107* M) - CCl, B-Naphthoquinoline (7.95 X 1075 M) - 5% H,SO, (H,O) Ultraviolet Difference Spectrum of B-Naphthoquinoline (10-4 M) and a-Chymotrypsin (10-4 M) - 0.1 N NaCl (H,0) Ultraviolet Spectrum of Quinoline (3.2 X 1074 M) in Methylene Chloride. 26 ce ee ae ee i Fig. 3a. PMDA (3.45x10™i0) 1,2 dichloroethane j i i t WN —" ¢ ot ~ : ot i a Qo gi re i 8 + Ss | { I { 0.5 . i 260 280 300 320 Wavelengthemytvoise ca eet ee cay 27 Figs 3b. PDA (3.45x1074I4) dichloromethane j ie Optical Density fe ete i ete wee et ee ee | _| | 300 325 350 Wavelength-n wr ee + cee cane Reena ame ae em ek 28 Figs 3c. PMDA (saturated) chloroform 1.5 1.0 0.5 300 325 Navelongth~m H- 350. Optical Density 29 Figs 3d. PsDA(saturated) carbon tetrachloride Optical Density 0.5 300 Wavelength=m we 350 30 Figs 4. PDA (6.28x10~m) : HAO 2 PUDA (5-9xL0™it) in DF ot PMDA (6.8x10~é) THF oO . Tes : a : ert , 4 : A : "d . 3 a. O° at ras \ | ee 300 350 Wavelexngth-m Jinn 31 Optical Density (DMF) Figs 5+ PDA (2.57x10™) isog (1.06x10~2i) CHACl, PMDA (5.1x10™) 4s0q (1e13x1L07*i) CHAC1> | “e ci tN Oo So 2 lone > ai 4 Le? S Fat q o or wey § We ra} “ \ To DMF A } 32 360 4,00 Wavelengéh-m ts esas 32 Figs 6+ isoquinoline (107™4u) methylene chloride | nr 1 45 Lo i © ie me: p ont L itp} i S ! | qo i ca ° | TO i Hl 4. i Wavele nether co a ce nterane ene cee ci pe ate ete eee ee ney 33 Fige 7. TONE (7.46x207>it) methylene chloride i t \ i 1o ! -o : Pe] i t i i i : ra ' re 8 | uw ere hd [o) + i \ ee r 280 Navelength-m Optical Density 34 Fig. Sa. TCNE (9.4x107°m) isoquinoline (6.41074) methylene chloride fa ri ° VA LN is 400 450 ~ 500 Wavelorgthos be Optical Density 35 Fig. 8b. TONE (9.0x1072m) quinoline (3.1x1072) methylene chloride Wavelength=m uw Optical Density een ee ee ace my eee en ane ee ete Figs 8c. TCNE (1.2x10~4u) @-napthoquinoline ( methylene chloride 107%) 0.5 400 450 Wavelengthom 1.0 Optical Density wee eee ete ay ee Fig. 9. Scott Modification Z. max (418 mH) Koa (418 m ja) 6.38x10" 512.0 Me meee eee ees 20 4 2 A.B,/D x 109 30 10 By x 10° 38 Aqzsueg Teotqdo o°t S*0 hymotrypsin (4x1077M) ‘igs 10 I ad -c ) ( 0.1M NaCl 1,0 stearate feted ramen nimenpetc sige miata Mine sa oe meen nace occu tem nee ep Ee nee nN 360 250 Navelengsh-ma i @-napthoquinoline (1.16xi0~*x) | Nacl ( 0.1M ) H.0 ian i 7) red “ a : 4 | | { ae \ TA : 4 i BP \ i ot! \ a 1 S j a ve Oo ord Pe QO S| } i +‘? _— j i 39 250 ~ 300 350 Wavelength=n a Ae Aypsueq Teotydo get OT S*0 a | ne ST 40 apthoquinoline (2.65xL0™i) bon tetrachloride @ -n car peas eee . comet et OS perenne ere cee OT, 1 aeet eine ce a et EE pe ee et, 340 / Wavelength=m ,W 320 300 41 sqyzsueq T30t4d0 St ort $°0 Peele ot PE Sg LETT eh pe ge eee tee A paces ens reeem ae in arn esate ee, {- As ETRE nes Sls EDT SR et es fev consrenceee IRMA Sa ROL PRE TE } aries -(8.0x1079M) 0, (H20) 380 enapthoquinoline Fig. jic. 5% H2S e 360 340 avelength-m ur Vi 42 Fig. 12. @ =napthoquinoline ( 10744) ol-chymotrypsin ( 10-4it) H,0 weteeemt sane eet bm = 250 300 350 Weveleagth-m a. | 5 i val — © Optical Density 43 Aqytsueq [BOT Ado s*t | ort $*0 an eeneeeneenens ste a ate ae —~-¥ ees prereset Sate peage tee erence ma rrereeremnd ACR AUN Te A eS (362x1074y) methylene chloride quinoline eccancttnabenecn ntact nnn eee . verwateg ty etree Geren nett et ee ome Ta -— ~ a aes A a tree eR ay ee ns ema i Hee awe oe ener aa 330 310 290 W avelength=m- 44 Annual Report ~ 1963 A Model for a Portion of the Active Site of a-Chymotrypsin Investigators: D. L. Roth and C. Niemann Support: National Institutes of Health It is known that many aromatic molecules combine with a locus at the active site of Q-chymotrypsin. The nature of this interaction is unknown. In order to determine whether or net a charge-transfer type of interaction is involved a model based upon vapor~phase chromato- graphy has been developed. This model entails use of a solution of a set of x-acids in squalene and determination of the retention co- efficients of a series of xn-bases. The retention coefficients should be dependent net only upon electronic but also upon steric factors. Comparison of retention coefficients so determined with the enzyme~ inhibition constants should determine whether or not a charge-transfer type of interaction is of any significance in the inhibitor--active~ site interaction. 45 Annual Report - 1964 Spectrophotometric Determination of Charge-Transfer Equilibrium Constants Investigators: D. L. Roth and C. Niemann Support: Public Health Service, Division of General Medical Sciences Comparison of charge-transfer equilibrium constants with enzyme inhibition constants should give some information concerning the role of charge- transfer binding in the inhibitor-enzyme complex. The equilibrium con- stant for the isoquinoline-tetracyanoethylene complex has been deter- mined by the method of Benesi and Hildebrand [g. Am. Chem. Soc., 71, 2703 (1949)]. Studies on solutions of the quinoline-tetracyanoethylene and Qnaphthoquinoline-tetracyanoethylene complexes are complicated by significant concentrations of a species believed to be the tetracyano- ethylene radical ion. The difference spectrum of a solution of a-chymo- trypsin and B-naphthoquinoline exhibits a shoulder at 360 mp which can be ascribed to charge-transfer absorption. A Model for a Portion of the Active Site of a-Chymotrypsin Investigators: 0D. L. Roth and C. Niemann Support: Public Health Service, Division of General Medical Sciences Experimental studies on the relationship between retention coefficients for various x-bases in gas chromatography and enzyme inhibition constants, described in the 1963 Annual Report, have been completed. The low tem- peratures required for formation of charge-transfer complexes limited the study to C -C, aromatic hydrocarbons. Addition of pyromellitic dianhydride 6 to the liquid phase of the column produced an increased separation between m- and p-xylene, but the order of elution corresponded to the boiling points of these n~bases, and no conclusions were reached. 10. li. lz. 13. 14, 15. 16, 46 REFERENCES E, K, Anderson, Acta Chem. Scan., 8, 157 (1954). J. J. Sudborough, J. Mme. Chem. Soc., 109, 1339 (1916). J. Fritzsche, Z. Chem., 114, 1 (1869). M. Orchin and E. Woolfolk, J. Am. Chem. Soc., 68, 1727 (1946). ~~ E. Pfeiffer, Ann., 404, 1 (1914). R. L. Shriner, R. C. Fuson and D. Y. 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