J . Phys. Chem. 1986, 90, 6366-6370
6366
The pOH Jump: Determination of Deprotonation Rates of Water by 6-Methoxyquinoline and Acridine E. Pines: D. Huppert,*+ M. Gutman,$ N. Nachlie1,t and M. Fishmant Raymond and Beverly Sackler Faculty of Exact Sciences, School of Chemistry, and G . S. Wise Faculty of Life Sciences, Department of Biochemistry, Tel Aviv University, Tel Aviv. Israel 69978 (Received: April 21, 1986)
6-Methoxyquinoline (6MQ) and acridine are known to become more basic when excited to their lowest singlet state. We utilized a direct time-resolved subnanosecond technique and steady-state fluorescence to determine the acid-base kinetics of these excited compounds @N*+ H 2 0 ~2 @N*H++ OH- ( k l ,k-l). The protonation rates ( k , ) were found to be 5.2 X lo6 M-' s-l and 0.2 X lo6 M-' S-I for 6MQ (pK* = 11.8) and acridine (pK* = 10.7), respectively. The neutralization rate (k-J was found to be diffusion controlled (6MQ) k-, = 3.1 X 1O'O M-' s-l. The deuterium isotope effects were found to D 3.5 and k - l H / k - l D = 2.2. The protonation rate increased in the presence of metal cations; for acridine in 2 be k I H / k l = M Mg(N03)2solution kl = 73 X lo6 M-I SKI. Introduction The rapid dissociation of protons from excited aromatic alcohols1q2has been employed for the rapid acidification of aqueous solutions and the determination of diffusion-controlled rate constants of various proton-transfer reaction^.^-^ The primary step of the proton dissociation from the excited proton emitter is of special interest. Its rate is controlled by a free energy relationship; the rate can be slowed by many orders of magnitude by lowering the activity coefficient of the water in the s o l ~ t i o n . ~This , ~ method was employed for determination of the water activity of specific sites on or on the interface of micelles or phospholipid surfaces.s An analogous reaction is the proton abstraction from water by excited heterocyclic compounds that are much stronger bases in their first excited singlet state than in their ground ~ t a t e . ~ . ' ~ Steady-state fluorescence measurements were employed"*'2 to evaluate the rate constant for proton abstraction and reprotonation. More recently kinetic methods based on the time-resolved fluorescence measurements have been applied to determine the rate of excited-state protonation of acridineI3 and na~hthylamine'~ in aqueous solutions. We have demonstrated that compounds like acridine or 6methoxyquinoline (6MQ) can be used for pulse alkalinization of small solutes or proteins in aqueous solution^.^^ Naturally, it will be of interest to evaluate how the initial step, the reaction between the excited base and water, is affected by the state of the water molecules, which function as the proton donor. For this purpose, we carried out the experiments described in this publication. We first investigated, by time-resolved and steady-state fluorescence techniques, the proton-transfer dynamics in pure water. The data compiled by these measurements was taken as a reference for further investigations where we varied the properties of water. We compared the effect of water dilution by methanol with experiments where concentrated solutions of strong electrolytes were employed to reduce the concentration of the free bulk water. As documented below, cations with high solvation energy (Li', Ca2+,or Mg2+) accelerate the rate of the proton transfer, suggesting that the water molecules that are solvating the cations are more acidic than bulk water. There are many enzymes where transition-metal ions (Mn2+, ZnZ+)are an essential active component of the active site. Their interaction with H20molecules or OH moieties in the active site is mandatory for catalysis. On the other hand, enzymes like ATPase act on substrates which provide the metal ion as a complex.I6 These water molecules, activated by the ATP-bound cations may play an essential role in ~atalysis.'~In our opinion the study of metal-induced proton transfer, as described below, will help to elucidate the mechanism of such catalytic reactions. School of Chemistry. *Department of Biochemistry
Experimental Section Materials. 6-Methoxyquinoline and acridine (reagent grade, Merck) were crystallized from ethanol. Deionized water by Micropore 11 with resistivity higher than 10 M a was used in all experiments. D 2 0 (99.75%) was by Merck. The pOH was adjusted by addition of known quantities of 1 M solution of NaOH (Titrizol). Steady-State and Time-Resolved Fluorescence. Time-resolved fluorescence was measured by a system based on a 400-ps light pulse generated by a home-made Blumelin transmission nitrogen laser (337 nm), operating at atmospheric pressure.I8 The laser peak power was 50 kW at a 10-Hz repetition rate. The fluorescence light was passed through a 250-mm Jarrell-Ash monochromator and monitored by a microchannel plate photomultiplier (Hamamatsu 1294U). The photomultiplier output was amplified, digitized, averaged by a Tektronix 79 12AD digitizer with a 7A-19 amplifier, and convoluted by a Tektronix 4051 computer. The overall time resolution of the system was better than 0.5 ns. Steady-state fluorometry was measured with a Shimadzu RF 540 spectrofluorometer. The emission spectra of the heterocyclic compounds were measured while excitation was at the isosbestic point of the absorption spectrum. Thus the two ground states @N and @NH+ had an equal probability of photoexcitation. The spectra were recorded over a wide range of pH (2-12) unless prevented by the metal hydroxyl precipitation. Results and Discussion Steady-State Fluorescence. The fluorescence of 6MQ in (1) Smith, K. K.; Huppert, D.; Gutman, M.; Kaufmann, K. J. Chem. Phys. Lett. 1979, 64, 522. (2) Clark J. H.; Shauiro, S.L.; Campillo, A. J.; Winn, K. R. J . Am. Chem. SOC.1979, 101, 746. (3) Gutman, M.; Nachliel, E.; Huppert, D. Eur. J . Biochem. 1982, 125, 175. (4) Pines, E.; Huppert, D. J . Phys. Chem. 1983, 87, 4471. (5) Gutman, M. Methods Biochem. Anal. 1984, 30, I. (6) Huppert, D.; Kolodney, E.; Gutman, M.; Nachliel, E. J . Am. Chem. SOC.1982, 104, 6949. Politi, M. .I.Chaimovich, ; H. J . Phys. Chem. 1986, 90, 282. (7) Gutman, M.; Huppert, D.; Nachliel, E. Eur. J . Biochem. 1982, 121, 637. (8) Gutman, M.; Nachiel, E.; Fishman, M. In Ion Interaction in Energy Transfer Systems; Papageorgiou et al., Ed.; Plenum: New York, 1986; p 93. (9) Ireland, J. F.; Wyatt, P. A. W. Adu. Phys. Org. Chem. 1976, 12, 131. (10) Weller, A. Prog. React. Kine?. 1961, I , 189. ( 1 1) Mataga, N.; Kaifu, Y.; Koizami, M. Bull. Chem. SOC.Jpn. 1956, 29, 373. (12) Weller, A. Z . Elektrochem. 1957, 61, 956. (13) Gafni, A,; Brand, L.; Chem. Phys. Lett. 1978, 58, 385. (14) Shah, J.; Pant, H. C.; Pant, D. D.; Pant, H. C. Chem. fhys. Lett 1985, 115 192. (15) Gutman, M.; Nachliel, E. FEBS Lett. 1985, 190, 29. (16) Carmeli, C.; Lifshitz, Y.; Gutman, M. Biochemistry 1981, 20, 3490. (17) Gutman, M.; Levy, M. J . B i d . Chem. 1983, 258, 12 132. (18) Von Bergman, H. M.; Pendris, A. J. J . Phys. E . , 1977. 10, 602.
0022-3654/86/2090-6366$01.50/0 0 1986 American Chemical Society
The Journal of Physical Chemistry, Vol. 90, No. 23, 1986 6367
The pOH Jump
l
o
o
X
7
nm
Figure 1. Steady-statefluorescence of 6MQ in its acidic form (A) (0.05 M HC104)and its basic form (B) (0.1 M NaOH); the excitation was at the isosbestic point A,, = 287 nm.
3 TIME
aqueous solution has two emission bands that vary in their intensities with the pH. In alkaline solution of 0.1 M NaOH a single emission band exists with a maximum a t 370 nm. In 0.05 M HC104the emission maximum is shifted to 442 nm (Figure 1). The kinetic scheme of the proton abstraction by the excited 6 M Q is
Figure 2. Time-resolved measurements of 9N* fluorescence at 370 nm in H 2 0 (A), ~~~0= 1.9 ns, and in D20(B), T ~ , O= 3.0 ns.
The reactions leading to the steady-state composition of the excited populations are @N*
+ HzO
SCHEME I ki
+ N * t H ~ Oc+N*H+
1
-
+
pko=6.2
+N t H ~ O
+ ki, + ki,
= E,,[@N*]
+ EG'[@N*H+]
(1)
where Ex, is the empirical "emission coefficjent" relating the emission of either species with their concentration measured at extreme pH values where only one species is assumed to exist (0.1 M N a O H for @N*' or 0.05 M HC104 for @N*H+). The fluorescence measurements have been carried out at two different wavelengths, AI and A2. A, was chosen at a wavelength where (PN* is emitting almost exclusively (420 nm for acridine and 370 nm for 6MQ). X2 was taken on the red side of the @N*H+emission maximum where the emission of @N*H+is much stronger than that of @N*. Under these conditions [@N*I = ZA,/EA,
& @N*H+ k-z
@N*
+ M+(H*O),,
k3 k-3
@N*H+
(1)
(11)
(1b)
(111)
where in our case the proton donors are metal ion hydrates (M+(HZO),). Under steady-state conditions and in the presence of metal ion hydrates the ratio [@N*H+]/[@N*]is given by the following expression: [@N*H+]
-
[@N*I [H+l(k/ko) + [H+]kz + ki [HzO] + ~ ~ [ M + ( H Z O ) , I (2) kj k-2 k-I[OH-] + k-3[M+(H20),lOH-]
+
+
Time-Resolved Fluorescence of 6-Methoxyquinoline. Timeresolved measurements, carried out at pH 6.0, are shown in Figure 2. In H 2 0 , @N*decays with a time constant rH@= 1.9 ns (curve A). In D 2 0 the decay is significantly longer (curve B) rDZo= 3.0 ns. Similar measurements were carried out at 480 nm to monitor the emission rise and decay of @N*H+ (not shown). The fluorescence rise time was 1.7 ns in HzO and 2.5 ns in D20. Thus we find a good correlation between the appearance of 3N*H+ and the depletion of the @N*population. Analytic Solution of the Rate Constants. In the case of pure water a t p H >5 and with the terms defined in Scheme I and reaction I, the dynamics of @N*H+ and ON* are given by d[ON*H+]/dt = -[ON*H+](kj
+ k-I[OH-]) + [@N*]klw (3)
d[@N*]/dt = -[@N*](kf
Wx,-~x,l[~x,/~x,ll/~,,'
+ M+(Hz0),10H-
(la)
and [@N*H+I =
+ OH-
+NH+ t OH-
k, is the pure radiative rate constant and kicand ki, are the internal conversion (Sl-So) and the intersystem crossing (Sl-Tl) nonradiative rate constants, respectively. We determined the fraction of @N*H+from fluorescence intensity measurements by using the following considerations. The fluorescence intensity (I)at a given wavelength (Xi)is the sum of the emission of @N* and @N*H+ Ih,
@N*H+
In situations where other proton donors (besides bulk water) are present in the solution, a third reaction should also be considered
1kV
where @ N is the unprotonated heterocyclic compound, @N* is the excited compound, and kf and k j are the apparent radiative rate constants of @N*H+and @N* in the absence of proton reactions and are of the form
k f = k,
kl k-1
@N* + H+
OH-
k-1
kt
ns
+ kl,,,) + [@N*H+][OH-]k-l
where klw = kl[H20]. Under conditions where [OH-] C
(4)
6368 The Journal of Physical Chemistry, Vol. 90, No. 23, 1986
Pines et al. I
I
I
I
I
T
'OI-
G
c
>
k
-
I-
v)
z
W
"
I-
-z
n X 0)
W
> I4
0
J W
2.0
a
4.0
6.0
[OH1 m M
Figure 4. Plot of (y2- ki) exp( 1.02Z1/2/(1 + [*I2)) of [OH-] (a) or [OD-] (b) (see text).
U
I
0
10 TIME ;
I 20 ns
Figure 3. Luminescence decay curves of bN*H+ measured at 480 nm as a function of OH- concentration: (-) pure water; 0.1 mM NaOH; (-- -) 2 mM NaOH.
The simultaneous solution of the coupled equations 3 and 4, subjected to boundary conditions [QN*]r=o= [QN*], and [QN*H+]l=o= 0 are given by eq 6 and 719 fQN*l. =
.
4 ,
[QN*Io
((-) Y2
- Y1
exp(--ylt) -
(e) Yz 71
exp(-~~t)\
The terms y1 and yz are the roots of a quadratic equation Yl,Z
=
IO-' as a function
of various concentrations of NaOH (or NaOD), Figure 3. In order to employ eq 10 for the determination of k-,, we had to account for the effect of the ionic strength on the rates of the reactions. Of the terms appearing in eq 10 all but k-, are independent (at first approximation) of ionic screening. The ionic strength correction is given bylo,zo = k-lo exp(-1 .02Z1/2/( 1
(e-)
M (pH > 4(ala4 - aZa3), which leads to the approximations
+
+
+ ai) = kl, + k j + kf + k-l[OH-] (klw + kf)(k-l[OH-] + k;) - k-l[OH-]klw 72 = kl, + kf + k t + k-l[OH-] 71 = -(a4
(8) (9 )
At low OH- concentrations the term k-, [OH-] can be omitted from the denominator of eq 9. Under these conditions and the fact that kf i= 10k; we expect a linear dependence of QN*H+ decay on OH- concentration according to eq 10 72
N
kj
kFk-1 [OH-] + klw + kf + k j
Determination of the Deprotonation Rate of QN*W by OH(k-l). The reaction between OH- and QN*H+ was measured by time-resolved fluorescence of QN*H+at 480 nm in the presence (19) Loken, M. 1972, 11,4719.
As k-, was independently determined and K+N* can be directly measured by steady-state fluorescence titration (pK* = 11.8 f O.l), we obtain kl = 3.5 X lo6 M-' s-l. The effective first-order rate constant of proton abstraction by 6MQ* in aqueous solutions is thus klw= 2 X lo8 s-,, where the concentration of HzO was taken as 55 M. Another method to determine k l is based on the measurement of the quantum yield ratio (QY) of QN*H+ and QN*D+ formation, R = QYH20/QYD20. This ratio was determined by extrapolating to time zero the time-resolved relative fluorescence intensities of the two compounds. The quantum yield for 9N*H+ or QN*D+ formation is given by QY = kiW/(kiw + kd With the assumption,that the initial concentration of QN* and its radiative lifetime (kf) are the same in HzO and DzO,the ratio between the relative fluorescence intensities of QN*H+ and QN*D+ a t time zero is equal to R.
+ +
QYH~o klwH20(klwD20 kf)
R=--
-
QYD~o klwD20(klwH20 kf)
-
klwH20yzD20
(12) klwD20y2H20
R.;Hayes, J. N.; Gohlke, J. R.; Brand, L. Biochemistry
(20) Eigen, M.; Kruse, W.; Maass, G.; De Maeyer, L. Prog. React. Kine?. 19a,2,2a7.
(21) Extracted from measurements by: Lewis, G. N.; Doody, T. C. J . Am. Chem. SOC.1933, 55, 3504. Von Ertl, G.; Gerischer, H. 2.Elektrochem. 1962, 66, 560.
The Journal of Physical Chemistry, Vol. 90, No. 23, 1986 6369
The pOH Jump TABLE I: Kinetic and Themodynami~Constants for the Protonation and Deprotonation of Excited dMethoxyquindine rate const exptl calcd klHZ0 (5.2 f 0.3)X lo6 M-'s-I (3.6 f 1.2) X 106 M-' s-l" klDZ0 (1.5 f 0.2) X lo6 M-Is-I (1.0f 0.2) X lo6 M-' s-lb klHZ0 k i D g 3.5 f 0.3 3.6 f 0.3c k-lH2 (3.1 f 0.6) X 1O'O M-' s-I (4.0f 0.3) X 1O'O M-' K i d k-lDzo (1.4f 0.3) X 1O'O M-' s-l (2.3 f 0.2)X loio M-' s-Ic k-lHzo/ 2.2 f 0.3 1.75 f 0.09 k-lD20
2.5 X lo* 1 1 .9h (1 1.8)g 11.7*(12.3)s
kf pK*H20 pK*Dp
"Based on pK* of 6MQ in H20. bBased on p P of 6MQ and pK shift in D20. cBased on difference in mobility of OH- vs. OD-, the pK* shift in D,O, and the measured klH and kID. dUsing DebyeSmaluchcwski equation for diffusion-controlledsecond-order reaction with reaction radius of 4.5 A. CUsingthe value of k-lH corrected for the mobility of DzO. fusing the ratio of mobilities (ref 21). g Determined by fluorescence titration. Determined from the rate constants. In pure water R was found to be 2.2 f 0.2. The values of yZHzo and y2ware determined experimentally. These are the reciprocal M values of the ON* fluorescence decay time a t [OH-] < (see eq 5). y2Hz0
= klwHzo+ kf = 5.3 X lo8 s-l
+
ylDZ0 = klwDZo kf = 3.3 X lo8 S-I
(13) (14)
It follows that the isotope effect in proton abstraction is given by klwHzo= Ki,
klwDZo
7ZHZ0
= R-
= 3.5 f 0.3
(15)
-fZDZO
The solution of eq 12, 13, and 14 yields the kinetic rate constants kiwHzo= (5.1 f 0.4) X lo6 M-' s-I klwDzo= (1.4 f 0.2) X IO6 M-' s-I kf = (2.5 f 0.2) X 10' s-' The calculations made so far were based on steady-state fluorescence titration or on time-resolved relative quantum yield measurements. It would have been better if corroborating values derived from direct kinetic study were available. The simplest method to obtain kl would have been to determine the rate of the fluorescence decay in the absence of proton abstraction (kf) by measuring the lifetime of ON* in solvents where ON*H+ is not formed. Unfortunately, although this method was suitable in previous studies with proton emitters,22it is not suitable in our case. It is well-known that the fluorescence intensity of heterocyclic compounds like quinolines of acridines is stronger in water than in aprotic organic ~ o l v e n t s . ~The ~ - ~reason ~ for that is the strong vibronic coupling between the close-lying lowest excited electronic states (n?r* and ?r?r*). This coupling, which is stronger in nonpolar organic solvents than in water, distorts the lowest excited singlet state and increases the Franck-Condon overlap with So or TI. As a result in nonpolar organic solvents the probability of nonradiative decay SI So or triplet formation SI TI is higher than that in water. The short lifetimes of ON* in organic solvents excludes the possibility of an independent estimation of kf. Still, a rough estimation can be drawn from measurements in methanol. In this
-
-
(22) Kolodney, E.; Huppert, D. Chem. Phys. 1981, 63, 801. (23) Lim, E. C. In ExcitedStates; Lim, E. C., Ed.;Academic: New York, 1977; Vol. 3. (24) El-Sayed, M. A.; Kasha, M.Spectrochim. Acta 1959, 15, 758. (25) Noe, L. J.; Degenkolb, E. 0.;Rentzepis, P. M . J. Chem. Phys. 1978, 68, 4835.
[MC104]
M
Figure 5. Dependence of the fluorescence relative quantum yields +(R) on the electrolyte concentration: (0)M = Li; (@) M = Na.
solvent and its deuterated form, CH30D, the lifetime of ON* is 3.04 and 3.07 ns, respectively. These practically identical lifetimes indicate that the replacement of H by D has no effect on the radiative or the nonradiative processes. Thus the shorter lifetime of ON* measured in H 2 0 (1.9 ns) indeed represents the contribution of proton abstraction. However, the lifetime that was extracted from quantum yield measurements (kf)-I = 4 ns is longer than that measured in methanol (3 ns), which implies that these measurements can determine only the upper limit of kf in HzO. Comparison between the Measured Rate Constants and Thermodynamic Predictions. The various rate constants associated with the reversible proton abstraction from water by excited 6MQ are listed in Table I, column 2. The rates estimated by thermodynamic calculations are listed in column 3. As seen from the table the agreement is high, indicating that the reactions are essentially diffusion controlled with no intermediate steps. This adherence allows us to extract from the kinetic values the physical parameters associated with the diffusion of OH- or OD-, namely, their diffusion coefficient and their effective reaction radius. Eflect of Solutes on the Formation of ON*H*. The observed rate of ON*H+ formation decreases when the water in the solution is diluted by organic solvents like methanol. As discussed above, part of the effect is attributed to the increase in the nonradiative decay rate of ON*. Thus organic solvent-water mixtures are not suitable for evaluating the role of water on the rate of proton abstraction. Reduction of the activity coefficient of water by addition of high concentrations of strong electrolytes6led to a marked increase of the relative emission quantum yield of ON*H+. This effect, represented in Figure 5, is specific for the cation and is not a colligative property of the aqueous solution: drawing the results vs. the activity coefficient of the solvent does not agglomerate the effects of the two cations (not shown). The rate of protonation of 6MQ* in water is so high that the absolute incremental protonation in presence of salts is small. Acridine, with a lower pK*, reacts more slowly with water and thus the augmenting effect of the salts is more pronounced. In Figure 6 we draw the fraction of ON*H+of the total excited population (calculated from eq l a and lb) as a function of the pH, when measured in the presence of concentrated solutions of various metal cations. The inflection point at p H 5.5 is due to the deprotonation of the ground state while that at high pH is due to the deprotonation of ON*H+ by OH-. These results were analyzed with eq 2. The denominator of this expression is dominated by kf = 3.2 X lo7s-I. All other terms are much smaller; k2,the rate of @N*H+dissociation, is less than 1 SKI.The product k-l [OH-] will be smaller than kf' even a t p H 10,and the product k-3[M+(H20),0H-] will be small due to the
6370 The Journal of Physical Chemistry, Vol. 90, No. 23, 1986
I
1
0 2
3
4
I
I 5
6
7
I
I 8
9
1
given in Figure 5 were analyzed according to this approximation, and the rate constants are listed in Table 11. The validity of these values was corroborated by measuring the effect of Li+ on the decay time of the excited acridine. The lifetime of 9 N * measured at 427 nm was found to be a linear function of the cation concentration and the protonation rate of $N* by Li+(H20) was 14 X lo6 M-I s-l. Similarly the rate for Mg2+(H20)was k3 = 79 X lo6 M-' s-I, which is identical with the rate calculated from the data published by Weller.'O The tendency of metal hydrates to donate a proton (see Table 11) is in accord with the observation of Zundel et a1.26that these cations polarize the 0-H bond in the cation hydration sphere with the subsequent appearance of an absorption continuum in the infrared spectrum of these aqueous solutions.
I
I 0
1
Pines et al.
1
1
2
PH Figure 6. Population fraction of the excited protonated form ON*H+/ [ON* + ON*H+] (obtained from steady-state fluorescence intensities) as a function of the pH measured in the presence of a fixed high concentration of strong electrolytes: (@) in the presence of 1.0 M Mg(N03)2;(0) in the presence of 2.0 M Na(N0,); (B) in pure water.
TABLE Ik Rate Constants of Protonation of Excited (SI)Acridine by Cation Hydrates. cation 106k M-' s-l method" H20( bulk) 0.22 f 0.04 SSF, TRM SSF Cs+(Hzo)n 0.8 f 0.2 2.9 f 0.4 SSF K+(HzO)n 6.9 f 1.4 SSF Na+(HzO), SSF, TRM Li+(H20), 15.8 f 1.4 CaZ+(Hz0). 43.0 5.5 SSF, TRM 12.9 f 6.1 SSF Mg2+(Hzo)n a Measured by steady-state fluorometry (SSF) or by time-resolved measurements (TRM).
*
low concentration of the solvated metal-hydroxy complex. Thus eq 2 can be rewritten as 9N*H+ -9N*
i.e., the variable @N*H+/ON*will be a linear function of [H+] with an intercept a t [ k l ( H 2 0 ) + k3[M+(H20),]/ki. The data
Summary In aqueous solutions certain heterocyclic compounds when in their lowest electronically excited state (@N*) react with the solvent very rapidly to form the equivalent protonated species (@N*H+). When initiated by short intense laser pulse, this reaction creates a transient increase in the solution pH, (the "pOH jump"). In the presence of a stronger base, like N a O H (at pH > lo), the protonated species neutralize the base with a diffusion-controlled rate. The primary step of the process is a proton transfer from nearby water molecules to a 9 N * molecule. Its rate was found to depend on the pK* value of the heterocyclic compound, which implies that the excited-state lifetimes are sufficiently long to permit the establishment of the hydrolytic equilibrium. The sensitivity of the hydrolytic process to the water acidity was demonstrated by adding metal cations to the solution. These cations alter the acidic properties of their solvating water molecules, increasing their acidity. As a result these water molecules protonate basic species much more radpily than bulk water. Acknowledgment. This study was supported by The American-Israel Binational Science Foundation (84/00 100) grant to M.G. Registry No. 6-Methoxyquinoline,5263-87-6; acridine, 260-94-6; deuterium, 7782-39-0. (26) Zundel, G. In The Hydrogen Bond in Theory and Experiments, Schuster, P.;Zundel, G.; Sandorfy, C., Eds.; North-Holland, Amsterdam, 1976; p 687.
