J. Phys. Chem. 1984,88, 4229-4232
4229
The Influence of Intramolecular Hydrogen Bonding on Proton-Transfer Reactions. A Temperature-Jump Study of Acid-Base Reactions Involving Substituted Salicylic Acids H. Diebler,* F. Secco,? and M. Venturinit Max- Planck-Institut f u r Biophysikalische Chemie, 34 Gottingen-Nikolausberg, West Germany (Received: November 28, 1983; In Final Form: February 8, 1984)
The temperature-jump relaxation technique has been used to study the kinetics of proton transfer between salicylic acid derivatives (3,5-dinitrosalicylic acid, DNSA, and thiosalicylic acid, TSA) and a second donor/acceptor system (ammonia or cacodylic acid). Depending on the system, the investigations have been carried out between pH 6.0 and 9.0. The concentrationdependencies of the observed relaxation time demonstrate that not only the direct proton-transfer process but also the hydrolytic or protolytic pathway, respectively, contribute to the overall reaction under the given conditions. The rate constants for proton transfer from the monoprotonated salicylate anions to the acceptors NH,, cacodylate-, and OH- are far below the diffusion-controlled values, even if these reactions are thermodynamically favored. This decrease in reactivity indicates that the proton to be transferred is engaged in an internal H bond in the mononegative salicylate anions. The strength of the internal H bond is much lower in TSA than in DNSA.
The effect of the presence of intramolecular hydrogen bonds on the rates of acid-base reactions was first recognized and discussed by Eigen et al. in several papers'-3 dealing with proton-transfer kinetics in aqueous solution. For instance, in salicylic acid type molecules where the phenolic proton is involved in
internal H-bond chelation, proton transfer to strong-bases like OHdoes occur only with rate constants well below those of diffusion-controlled processes. 1-3 This observation was accounted for by assuming a two-step mechanism for the overall reaction: the internal H bond has to open first before proton transfer to the base can take place. The magnitude of the deviation from the diffusion-controlled rate is considered to be a measure of the strength of the intramolecular H bond.3 Somewhat later, Eyring and co-workers reported evidence for an alternative mechanism in which the base directly attacks the bridging proton; breaking of the internal H bond and proton transfer to the base are assumed to proceed in form of a concerted (one-step) p r o ~ e s s . ~ Results which have been reported more recently for a few systems apparently indicate that both mechanisms contribute to the overall r e a c t i ~ n whereas ,~ those for several other systems are best accounted for by the two-step process.6 Similarly, a chelating ligand with the binding sites bridged by internal H bonding may also be less reactive toward complexation with metal ions than an analogous non-H-bonded chelate compound. This thesis, first advanced for the complexation of La3+ with pyridylazoresorcinol,' was subsequently used to interpret the anomalously low rates of complexation of Mg(I1) by alizarin yellow G8 and of Ni(I1) by tropaeoline and related compounds? More recently, the low rates of complex formation of Ni(I1) and Co(I1) with substituted salicylic acidslOJlwere explained on the same basis, and the fact that the rate constants were found to be independent of the strength of the phenolic 0-H bond (pKA2)of the salicylate anions was taken as evidence that the ring closure step (assumed to involve loss of the phenolic proton) is not rate determining. However, the details of the reaction mechanism and the nature of the rate-determining step are still under debate. A few recent studies of complexation reactions involving internally H-bonded ligands have been interpreted on the basis of rate-determining ring closure proce~ses.'~-'~ If complex formation reactions with ligands Permanent address: Institute of Analyt. Chemistry and Electrochemistry, University of Pisa, Pisa, Italy.
0022-3654/84/2088-4229$01.50/0
of this type would follow the mechanism which has been proposed in ref 11, their rate constant should be correlated with those of the ligands's proton-transfer reactions. Therefore, we report here kinetic and equilibrium data for acid-base reactions of two salicylates, 3,5-dinitrosalicylic acid and thiosalicylic acid, which presumably differ in the strength of their internal H bond. In a subsequent study we shall examine the behavior of these ligands in their complexation reactions with Ni(II).14
Experimental Section Materials. All chemicals (Fluka and Merck) were of analytical grade. 2-Mercaptobenzoic acid (thiosalicylic acid, TSA) was used without further purification. 3,5-Dinitro-2-hydroxybenzoicacid (3,5-dinitrosalicylic acid, DNSA) was recrystallized from water. Stock solutions of the salicylic acids were prepared by weight and standardized by pH titrations using NaOH. Solutions of known concentration of ammonium perchlorate were prepared by adding ammonia to calibrated amounts of perchloric acid until the equivalence point. Triply distilled water was used to prepare all the solutions. Methods. All measurements were carried out at 25 OC (hO.1 "C). The ionic strength Z was 0.3 M, adjusted with sodium perchlorate. The acidity constants of cacodylic acid ((CH3)2As(O)OH), 3,5-dinitro-2-hydroxybenzoicacid, and 2-mercapto(1) M. Eigen, Angew. Chem., In?. Ed. Engl., 3, 1 (1964). (2) M. Eigen, W. Kruse, G . Maass, and L. de Maeyer, in Prog. React. Kine?., 2, 286 (1964). (3) M. Eigen and W. Kruse, 2.Nuturforsch. B, 18, 857 (1963). (4) E. M. Eyring and J. L. Haslam, J . Phys. Chem., 70, 293 (1966); R. P. Jensen, E. M. Eyring, and W. M. Walsh, ibid., 70, 2264 (1966). ( 5 ) F. Hibbert and G . R. Simpson, J . Am. Chem. Soc., 105, 1063 (1983); B. Perlmutter-Hayman and R. Shinar, Int. J . Chem. Kine?., 10, 407 (1978). (6) A. J. Kresge and M. F. Powell, J . Am. Chem. SOC.,103, 972 (1981); F. Hibbert and A. Awwal, J . Chem. SOC.,Perkin Trans. 2, 939 (1978); T. Fueno, 0. Kajimoto, Y . Nishigaki, and T. Toshioka, ibid., 738 (1973). (7) T. Onodera and M. Fujimoto, Bull. Chem. SOC.Jpn. 44, 2003 (1971). (8) B. Perlmutter-Hayman and R. Shinar, Inorg. Chem., 16, 385 (1977). (9) B. Perlmutter-Hayman and R. Shinar, Inorg. Chem., IS,2932 (1976); 16, 2643 (1977). (10) E. Mentasti, E. Pellitzzetti, F. Secco, and M. Venturini, Inorg. Chem., 18, 2007 (1979). (11) E . Mentasti, F. Secco, and M. Venturini, Inorg. Chem., 19, 3528 (1980). (12) T. M. Che and K. Kustin, Inorg. Chem., 20, 509 (1981). (13) S.Chopra and R. B. Jordan, "Abstracts of the XXII International Conference on Coordination Chemistry, Vol. 1, Budapest, 1982, p 430; Inorg. Chem., 22, 1708 (1983). (14) H. Diebler, F. Secco, and M. Venturini, manuscript in preparation. (15) M. Eigen and L. de Maeyer in "Techniques in Organic Chemistry", 2nd ed, Vol. VIII, Part 11, A. Weissberger, Ed., Wiley, New York, 1963, p 895.
0 1984 American Chemical Society
4230 The Journal of Physical Chemistry, Vol. 88, No. 19, 1984 TABLE I: pK Values of the Acid Dissociation Constants of the Substituted Salicylic Acids, of Cacodylic Acid, and of the Ammonium Ion (25 'C. I = 0.3 M)
DNSA TSA (CHdMOPH NH,'
PKAI 0.28" 3.45 6.19 9.30b
Diebler et al. Scheme I pro toly sis (1)
242
$. H20 k3$hkI3
k23
( 2 ) B t HL t H 2 0
HB t L f H20 (3) k32
k4Skz4
"From ref 17, I = 0 . 2 M. bFrom ref 16b.
k43@34
HB t HL t OH
benzoic acid were measured potentiometrically with a Radiometer PHM 52 pH meter with a Metrohm combined electrode in which the liquid junction was a 3 M NaCl solution. The pH meter was standardized with Merck buffer solutions of pH 4-10. In order to convert Hf activities to H+ concentrations, the electrode was calibrated by measuring the pH of solutions of known concentration of H+ (e.g., 0.01 M) in 0.3 M NaC104. Purified nitrogen was bubbled through the solutions before and during the measurements. Solutions (20 mL, 1 X M) of the acids were titrated with 0.1 M NaOH by using a microsyringe which was accurate to 2 X lom4mL. Under our conditions the ionic product of water is given by K, = [H+][OH-] = 1.7 X M2.16a Studies of the kinetics were carried out by means of the temperature-jump relaxation technique with spectrophotometric det e ~ t i 0 n . lThe ~ cell with the solution was thermostated at 22 OC and then the temperature was raised to 25 OC by discharging a 0.05-~F capacitor of 30 kV. The time constant for the temperature change is about 2 ps under our experimental conditions. The chemical relaxation processes were observed at wavelengths at which the changes in absorption of the salicylates due to the protonation/deprotonation reaction are relatively large, thus providing reasonable amplitudes. The relaxation signal was stored in a transient recorder (Datalab DL 905) and displayed on a two-channel oscilloscope. An exponential curve of variable time constant and amplitude, produced by an electronic device (constructed in this laboratory by C. R. Rabl), was displayed simultaneously via the second channel. The amplitude and the time constant of the synthetic curve were then adjusted to fit the experimental curve. The relaxation times given in this paper are averages of at least four individual values. The maximum spread between repeated runs was usually within 10%.
Results Equilibrium Studies. The protolytic pK values of the substances of interest are listed in Table I. Cacodylic acid ((CH,),As(0)OH) and NH4+/NH3were used as buffers to stabilize the pH of the solutions and to act as donors/acceptors in the protontransfer reactions. The values of the first dissociation constant KAI of the salicylic acids are several orders of magnitude higher than the Hf concentration in the kinetic studies. Therefore, only the second dissociation constants KA2= [H+][L2-]/[HL-] have to be considered. Kinetic Studies. The temperature-jump experiments were carried out with solutions of salicylate of total concentration C, = [HL-] [L2-] and buffer of total concentration CB = [B] [HB], usually under the condition CB>> C,. Following the temperature jump, a fast initial increase of transmittancy was always observed, its time constant being given by the heating time. The amplitude of the initial change of absorption is independent of the nature and concentration of the buffer (the effect is observed also in the absence of buffer) but is strongly dependent on pH. This effect, which has been found to be due to the temperature dependence of the extinction coefficients of the salicylate species (in particular L2-),I8 was well
+
t L
B t H
PKA2 7.02 8.09
+
(16) (a) R. M. Smith and A. E. Martell, "Critical Stability Constants", Vol. 4, Plenum Press, New York, 1976 p 1; (b) ibid., p 40. (17) R. Corigli, F. Secco, and M. Venturini, Inorg. Chem., 21, 2992 (1982). (18) For DNSA, this explanation was confirmed quantitatively by measuring the spectrum of L2-(pH 10; no buffer, no inert salt added) at various temperatures.
(4) hydrolysis
separated from the subsequent chemical relaxation due to the proton-transfer processes occurring in the system. The general reaction scheme for proton exchange between the salicylate and buffer species is shown in Scheme I (charges omitted); it involves a protolytic and a hydrolytic pathway as well as direct exchange, (2) (3).l If the concentrations of H+ and OH- are small compared to those of the other species (L, HL, B, HB), they can be treated as steady states and only one relaxation time is expected for such a system. Its reciprocal is given by the expression'
k34k42
k31k1Z
k12[L1 + ki,[B]
+
k42[HBl
+ k43[HL]
If equilibrium relationships are applied, this expression can also be written in the form 1 / =~ kfF
(2)
with
and where KHB and KHL denote the acid dissociation constants of the buffer and of the salicylic acid (KAZ, Table I), respectively. The results obtained for the various systems are now reported in turn. DNSA. The wavelength of observation was 440 nm in general; a few runs have been followed also at 390 nm. A positive change of transmittancy or light intensity (Le., I , - I,, > 0) was found for the reaction of DNSA with cacodylic acid, whereas a negative change was observed for the reaction with NH3/NH4+. This contrasting behavior obviously reflects differences in A H for the ionization of (CH,),As(O)OH and NH4+. For DNSA-NH,, the reactant concentrations were CL = 2.6 x 10-5-2 x 10-4 M, C, = 1 x 10+-36 x 10-3 M, PH 7.4-8.2. At those pH values (Le., near neutrality) only one time constant was observed, as expected ( T = 90-470 ws). Direct exchange, (2) (3), is not the only reaction path of importance, however, since plots of 1 / T vs. F did not give straight lines passing through the origin. The dependence of 1/(7F) = kf on the concentrations of the various reactant species unambiguously reveals that the additional reaction pathway is not the hydrolytic path but the protolytic path. With the approximation k12= k L 3(both rate constants refer to diffusion-controlled recombination processes) is then obtained (3)
A plot of 1/(7F) vs. 1/([L] + [B]), not shown, confirms the linear dependency. A computer best fit of the experimental data yields
The Journal of Physical Chemistry, Vol. 88, No. 19, 1984 4231
Effect of H Bonding on Proton Transfer
I
I
2
,
L
k,, k13 I (k,, EL1 + k,, [BI 1 ( W 2 S - l 1
+
Figure 1. Dependence of 1/(7F) on kI2kl3/(klZ[L] k13[B]) for the system DNSA-NH3 (25 OC, I = 0.3 M).
kz3 = 3.1 X lo6 M-ls-I a nd k,, = 4.5 X lo3 s-l. With k I 2 = k z l / K ~follows ~ k12 = k13 = 4.7 X 10" M-' s-'. If the approximation k12= k13is abandoned, a three-parameter fit based upon eq 4 yields k23 = 1.9 X lo6 M-' s-', kI2= 6.9 X 10" M-' s-I, and (4) k13 = 3.5 X 1Olo M-I s-'. From these values, a plot according to eq 4 is shown in Figure 1. With equilibrium data follows k32 = 1.0 X lo4 M-' s-I, k21 = 6.6 X lo3 s-l, and k31 = 18 s-l. For DNSA-cacodylate the reactant concentrations were CL = M, pH 6.0-7.4. M, CB = 1 X 10-3-6 X 5 X 10-5-3 X Because of the different overall equilibrium (lower K23) the values of F (eq 2) are appreciably larger than in the former case, and the relaxation times are much shorter, T = 5-60 p s . A comparison of the quantities in the rate terms for the protolytic and hydrolytic pathways (eq 2a) clearly indicates that at the given concentration conditions the contribution of the hydrolytic path is much smaller than that of the protolytic path. Furthermore, in all solutions studied the concentration of B (= cacodylic acid anion) was at least 20 times higher than that of L (= DNSA anion), leading to l/(i-F) = k23 (k21/[B]). A corresponding plot gave k23 = 4.5(fl) X lo6 M-Is-l a nd k,, = 4.5(f1.5) X lo3 s-l. The error limits of these rate constants are fairly large since (for reasons unknown) the experimental data showed considerable scatter. With kZ1/kl2= KHL = 9.55 X lo-' M-' and k23/k32 = K23 = KHL/KHB = 0.148 follows kI2= 4.7(f1.5) X 1OloM-l s-l and k32 = 3.o(fi) x 107 M - I S - ] . TSA-NH3. Measurements of the proton exchange between 2-mercaptobenzoic acid (thiosalicyclic acid, TSA) and NH3/NH4+ as second donor/acceptor system have been followed spectrophotometrically at 270 nm, where the spectral differences between the deprotonated (L) and monoprotonated form (HL) of TSA are largest. The concentration conditions were the following: CL = 5 X 10-5-3 X M, CB = 1 X 10-3-8 X M, pH 8.0-9.0. The observed relaxation time varied between 55 and 305 p s . In all experiments, k42[HB]>> k43[HL];with the approximation k12= k I 3 results
+
1 i-F
- = kz3 +
k2l [LI + [BI
k24k43 +
k42[HBl
(5)
The experimental data revealed an approximate linear dependency of l/(i-F) on l/[HB], as shown in Figure 2, but not on 1/([L] [B]), thus indicating that under the given conditions the contribution from the hydrolytic pathway is much more important than that from the protolytic path. If the term k2,/([L] + [B]) is neglected, a two-parameter best-fit evaluation of the experimental data gave k23 = 8.6 X lo6 M-l s-l a nd k24k43/k42= 1.5 X lo4 s-I, Le., kq3= 4.4 X los M-I s-l. With the corresponding equilibrium relationships results k32 = 5.3 X lo5 M-' s-l a nd k34 = 9.2 x 102 s-1.
+
,
500
6
1000
l / [ H B I (M-'1 Figure 2. Dependence of ~ / ( T F )on l/[HB] for the system TSA-NH, (25 OC,I = 0.3 M). TABLE II: Rate Constants (25 'C, I = 0.3 M) for Proton-Transfer Reactions k,
HL
+ B F=kb L + HB
HL B DNSA- NH3 (CH3)2As(0)OHzOa HzOb TSANH, OHH,O+ NH3
k,, M-'
s-I
kbrM-I
SKI
1.9(f0.4) X lo6 4.5(f1.5) X lo6 6.6(*1) X 4.5(f1.5) X 10''
l.O(f0.2) X lo4 3.0(fl) X lo7 6 . 9 ( f l ) X 1O'O 4.7(f1.5) X 10"
8.6(&0.8) X lo6 4.4(f0.4) X lo8
5.3(f0.5) X lo5 9.2(+0.9) X I O z c
3.5(f0.5) X 1O'O
18(f3)C
From the system DNSA-NH,. *From the system DNSA-cacodylate. First-order rate constant (s-I).
Actually, if the recombination of the anionic thio ligand L2with H+ occurs approximately diffusion controlled (rate constant k12, ApK = 9.8), then kZ1will be close to 5 X 10' SKIand the contribution of the second term of the right-hand side of eq 5 is not negligible but will amount up to 27% of the total value of kf, depending on the concentration conditions. Assuming kzl = 5 X lo2 s-l, a two-parameter fit yields kz3 = 8.2 X lo6 M-l s-l a nd k24k43/k42= 1.2 X lo4 s-l. These values are not very different from those obtained before with kzl = 0, but the fit is not far less satisfactory than before. If appears likely, therefore, that the recombination of H+ with the -S- group of the TSA anion occurs slower than diffusion controlled.
Discussion The observation that in the system DNSA-NH3 the hydrolytic pathway does not contribute significantly to the overall reaction is readily rationalized: In the kinetic experiments it was always [HB] >> [HL], and thus the last term of eq 2a simplifies to B ] .value of this expression k24k43/k42[HB]= K w ~ ~ ~ / K H B [ HThe is small compared to that for the protolytic term under our concentration conditions if the value of k4, is at least one order of magnitude below its diffusion-controlledlimit. The rate constant k43refers to the reaction of the monoprotonated DNSA with OHas acceptor. Normally, such a thermodynamically strongly favored process would be diffusion controlled.'S2 In case of salicylic acid-type compounds, however, the presence of an internal H bond in HL- reduces the rate of proton transfer by several orders of magnitude.'$* This point will be discussed in more detail below. A summary of the rate constants obtained in this study is presented in Table 11. The two values for the rate of recombination of H+ with the deprotonated DNSA (L2-) which have been evaluated from the systems DNSA-NH, and DNSA-cacodylate agree fairly well (4.7(fl.5) X 1Oloand 6.9 X 1OloM-I s-]) and are of the expected magnitude. Similarly, the rate constant 3.5 X 1Olo M-I for H+ + N H 3 NH4+is in good agreement with the value 4.3 X
-
4232 The Journal of Physical Chemistry, Vol. 88, No. 19, 1984 1O1O M-’ s-l reported from N M R studies.19 The decrease in proton-transfer rate due to internal H bonding (3) from is clearly expressed by the direct transfer process (2) DNSA (= HL-) to NH3. Although the equilibrium between (2) and (3) is in favor of the latter, with an equilibrium constant K23 = KHL/KHB= 191 (or ApK = +2.3), the experimental value of k23 (1.9 X lo6 M-I s-I) is smaller than the diffusion-controlled limit by about a factor lo3. Normally (Le., without internal H bond), proton transfer from an 0 donor to N H 3 as acceptor does occur diffusion controlled if ApK 1 2.0 (see Figure 9 f of 2 X lo9 M-l s-l. ref l ) , with k The decrease in rate is most straightforwardly rationalized in terms of the two-step mechanism, Le., by assuming that a small fraction of H L is present in a form in which the internal H bond is not closed, and that this open form (= HL’) reacts diffusion controlled (k2) with the acceptor:’
-
-
-
HL
+ K
-+
HL~
kz
L
NH~+
”3
If K1 is a rapid preequilibrium (Le., k-l >> k2[NH3]),then k23 = Klk2. Since k-l, the rate of closing the internal H bond (HL’ HL), is 3 X lo7 s-l for salicylaldehyde,20 it can be safely assumed that this condition is fulfilled for the NH3 concentrations used in the present study. With k23 = 1.9 X lo6 M-’ s-I a nd k2 2 X lo9 M-’ s-I, this interpretation yields K1 1X We may compare these results with those obtained for DNSA (= HL) cacodylate- (acceptor, cac-). Although for this system the equilibrium (2) + (3) is in favor of state (2), K23 = 0.15, the rate constant k32 (3.0 X 10’) is far below the diffusion-controlled limit; and the value of the rate constant for H L cac- (4.5(fl) X lo6 M-l s-l) is close to that for H L NH3, despite the difference in the energetics. This behavior, too, can be accounted for by eq 6: Since K23 = KlK2 = 0.15 and K1 > 1, Le., k2 is again diffusion controlled (the reverse rate is correspondingly lower). With k23 = Klk2, where kZ3 4.5 X lo6 M-’ s-I and k2 2 X lo9 M-I s-I, one obtains K1 2 X lC3, a value very similar to that derived before from studies of the system DNSA-NH3. While these results for the deprotonation of HL- by N H 3 and cac- are fully consistent with the assumed two-step mechanism, the same type
-
-
+
+
-
+
-
(19) M. T. Emerson, E. Grunwald, and R. A. Kromhout, J. Chem. Phys., 33. 547 (1960). ’(20) T . Yasunaga, N. Tatsumoto, H. Inoue, and M. Muira, J . Phys. Chem., 73,477 (1969).
Diebler et al. of second-order rate law applies also to Eyring’s concerted (one-step) mechanisms4 A true distinction between the two mechanisms can therefore not be made from experiments done under our conditions. Similar considerations apply to the reactions involving thiosalicylic acid. The value of the rate constant for proton transfer from the monoprotonated acid to OH- (k = 4.4 X lo8 M-’ S-I ) is by a factor 8 higher than the one quoted in Table IV of ref 2, but it is still about one order of magnitude below the diffusioncontrolled limit. Since this process is strongly favored (ApK = +7.6), the observed retardation in rate may again indicate internal H bonding in the monoprotonated thiosalicylic acid anion (HL-). In this case, however, the H bond is rather weak; applying eq 6 gives K1 N 0.1. It is generally assumed that donor or acceptor atoms other than 0 and N (and F) do indeed form only weak H bonds. In the reaction of thiosalicylic acid with N H 3 as acceptor the rate constant k23 = 8.6 X lo6 M-’ S-I is about 200-fold lower than the diffusion-controlled limit and the reaction is only weakly “downhill” (ApK = 1.2). Under such conditions, reactions of other -SH donors with N acceptors proceed also almost 2 orders of magnitude slower than diffusion controlled (e.g., thioglycol, see Figure 9c of ref 1).Iq2l This comparison reveals that also the rate of proton transfer from TSA to NH3 does not provide evidence for a drastic additional inhibition due to strong intramolecular H bonding. From the kinetic studies it is to be concluded, therefore, that in monoprotonated TSA intramolecular H bonding is at best rather weak. This conclucion is in line with the general views on H bonding involving SZ2As in proton-transfer reactions of other thio compounds, the low reactivity of thiosalicylic acid is mainly due to weak hydrogen bond formation with the acceptor group (because of the low polarity of the -S-H bond), and perhaps ro a requirement for appreciable electron-density redistribution within the TSA molecule during the protonation/deprotonation process.
+
Acknowledgment. Thanks are due to M. Jung of this institute for help with the computer evaluation. F.S. and M.V. express their gratitude to the Max-Planck-Gesellschaft for a short-term fellowship. Registry No. (CH3)2As(0)OH, 75-60-5; N H 3 , 7664-41-7; D N S A , 609-99-4; TSA, 147-93-3. (21) M. L. Ahrens and G. Maass, Angew. Chem., 20, 848 (1968). (22) G. C. Pimentel and A. L. McClellan, “The Hydrogen Bond”, W. H. Freeman, San Francisco, 1960, Chapter 6.
