J . Phys. Chem. 1985,89, 5222-5225
5222
To evaluate the a q ( H ) / a P A term in eq 7, we can use the variation of PA and q ( H ) in the R3N and R3NH+ series. The variation of the PA is obtained from experimental values and the variation of q(H) from a b initio r e ~ u l t s . ~ a4 - = [q(H)(NH4+) - d H ) ( ( C H , ) , N H + ) I /
aPA
[PA((CH3)3N) - PA(NH3)I 0.037 unit charge =-21 kcal mol-' From eq 7-1 1, we obtain the predicted slope of the correlation lines dED/dPA as 0.16 for R3NH+-NH3 and 0.24 for R3NH+--HCN. These predicted slopes are somewhat smaller than the experimental slopes of 0.25 and 0.35, respectively. However, the prediction agrees with the observed trend in that the correlation slope is larger in the R3NH+.HCN series. Indeed, the ratio between the calculated slopes, 1.50, is close to the ratio of the experimental slopes, 1.40. The larger intercept in the NH+-HCN series is also predicted by eq 4. In fact, these equations predict in general that larger intercepts will be associated with larger slopes, and that both will depend strongly on the polarity (gD) of the ligands. W e shall see in the following paper that these relations also apply in comparing XH+--O and XH+-.S series. Therefore, in summary, the variation of q(H) with PA(A) in a series of ions AH+ can account for the major observed trends in the AHDO vs. APA correlations. 3. A Structural Interpretation of A H D O us. APA Correlations. We compare the a b initio 3-21Gi5structures and charge distributions of NH4+-.NCH vs. NH4+-.NCCH3, representing the (15) Deakyne, C. A., et al., to be submitted for publication. (16) Lias, S.G.; Liebman, J. F.; Levin, R. D. J. Phys. Chem. ReJ Data 1984, 13, 695. (17) Raksit, A. B.; Bohme, D. K. In?.J. Mass Spectrom. Ion Processes 1984, 57, 211.
NH+-.NCR series. The energies, geometries, and charge distributions of the related monomer and complex ions are given in Figure 7. In comparing NH4+--NCCH3with NH,+-NCH, we observe in Figure 4 that as APA increases, the H+-.NCR hydrogen bond becomes more extended, from 1.675 to 1.733 A, and the N-H+ bond becomes less extended from 1.061 to 1.051 A. Both changes are consistent with the notion that the decrease in AHDO corresponds to less effective partial proton transfer in the complex as APA increases. Also, as expected, the more polar CH3CN ligand is observed to polarize the NH4+ ion more than does HCN, and correspondingly, the charge on the bonding proton of NH4+ is larger in H3NH+-.NCCH3 than in H3NH+-NCH. This factor contributes to the stronger electrostatic interaction in the former complex. It should be noted that this factor would not be accounted for in simple electrostatic calculations, where polarization upon complex formation is usually not considered. The a b initio results show that the overall charge transfer in these complexes is only 0.1 13e in NH4+-.NCCH3 and 0.124e in NH4+-NCH. Therefore, the contribution of charge transfer, i.e., of covalent factors, to the overall interactions must be small, and the complexes are mainly electrostatic in nature."-13 It is, however, unexpected that charge transfer should be more significant in the more weakly bonded complex. In this respect, it is also interesting that charge transfer in the NH4+-.NH3 complex is calculated as 0.187 (Figure 7). Also, the N-H+-N bond is only 1.484 A, and the N-H+ bond in the ammonia dimer is significantly extended (1.162 A). These results suggest that the similar stabilities of NH4+-.NH3 (AH,' = 24 kcal mol-') and NH4+.-NCH ( A H D O = 2 1 kcal mol-I) result from a compensation of stronger covalent contribution in the former vs. stronger ion-permanent dipole electrostatic interactions in the latter.
Acknowledgment. We thank Dr. C. A. Deakyne for making available results of the ab initio calculations, and Dr. L. W. Sieck for helpful discussions and assistance with the mass spectrometer.
The Ionic Hydrogen Bond and Ion Solvation. 4. SH+-O and NH'0-S Bonds. Correlatlons wRh Proton Affinity. Mutual Effects of Weak and Strong Ligands in Mixed Clusters Michael Meot-Ner (Mautner)* and L. Wayne Sieck Chemical Kinetics Division, Center for Chemical Physics, National Bureau of Standards, Gaithersburg, Maryland 20899 (Received: January 24, 1985)
Ionic hydrogen bonds of the type SH+-.O, i.e., dissociation energies of the complexes R2SH+-OH2, range from 12 to 19 kcal mol-', and exhibit an inverse linear relationship with proton affinity difference APA = PA(R2S) - PA(H,O), of the form A H D O = 18.6 - 0.16 APA kcal mol-'. Ionic hydrogen bonds of the type NH+.-S are also relatively weak, with values between 11 and 15 kcal mol-'. However, some symmetric dimers of the type SH+--S, e.g., (CH3)2SH+--S(CH3)2,are more strongly bonded, with A H D O = 26 kcal mol-'. Using the weakly bonding sulfur ligand CH,SH and a strongly bonding polar ligand CH3CN in a mixed cluster CH3NH3+.CH3SH.CH3CN,it is observed that the bond-weakening effect of the weak ligand on the interaction of the core ion with the strong ligand is only 13%. In comparison, the effect of the opposite combination, i.e., the effect of the strong ligand CH3CN on the interaction of the core ion with either a second weak ligand (CH3SH) or a second strong ligand (CH3CN), is to cause weakening of the bonds by 26 and 29%, respectively.
Introduction Interactions between protonated species and neutral molecules are important in environments ranging from planetary atmospheres to biological systems. In all of these environments, sulfur cornpounds and, ,,,,der ionizing or acidic conditions, sulfonium ions can occur. It is therefore of interest to examine the participation of sulfonium ions in strong ionic hydrogen bonds of the types
SH+-.O and SH+--S, and of sulfur compounds in bonds of the type NH+*& In an earlier article' we found that ionic hydrogen bonds BH+*-A involving oxygen and nitrogen compounds follow a regular behavior in that an inverse linear relation exists between the bond (1) Meot-Ner (Mautner), M. J . Am. Chem. SOC.1984, 106, 1257.
This article not subject to U S . Copyright. Published 1985 by the American Chemical Society
Ionic Hydrogen Bond and Ion Solvation dissociation energies, AHD', and the difference between the proton affinities of the proton donor B and the proton acceptor A, Le., APA= PA(B) - PA(A). The relation is of the form A H D O = a - bAPA, where a depends on the identities of B and A and b is generally in the range 0.23-0.30. Qualitatively, the correlation results from the dependence of the efficiency of partial proton transfer from B to A on the relative proton affinities of the components. In a practical sense, the correlations generally predict AHDowith the accuracy of f l - 2 kcal mol-', which is approximately as accurate as the usual error limit associated with experimental measurements. This predictive value is useful in predicting ionic hydrogen bond energies that have not been measured, or that occur on inaccessible environments, such as biosystems. Previous work also showed that the relation between the attachment energy of the first ligand molecule to an ion, and of further identical ligand molecules, is consistent for a large variety of clusters. For example, the ratio A H D O ( 1 , 2 ) / M D 0 (0,l) = 0.75 f 0.05 for most clusters; i.e., the attachment of one ligand molecule A to BH+ weakens the interaction with the second ligand molecule by 25%. While this empirical observation applies to the mutual effects of two identical ligands, there is no information on the mutual effects of different ligands, such as attaching a weak ligand A and a strong ligand C to the core ion BH'. Such information is needed in order to generalize our understanding of interactions in ion clustering and solvation. The only related information which is available involves systems in which AH+ is complexed to A and C. However, in this case one of the ligands is identical with the core moiety A, resulting in special stabilization effects involving the symmetric A-.H+-.A bond. The present work shows that sulfur compounds tend to be weak ligands in ionic hydrogenbonding systems. Therefore, we can use the attachment of the combination of a weak sulfur ligand, CH3SH, and a strong polar ligand, CH3CN, to the core ion CH3NH3+,to gain qualitative information on the mutual effects of bonding to substantially different ligands.
Experimental Section All measurements were taken with the N B S pulsed highpressure mass spectrometer, which has been described in detail elsewhere.2 The general experimental procedures used are discussed in the preceding article. The sulfur-containing compounds proved to be very difficult to study in certain cases. W e encountered no experimental problems in those measurements involving trace quantities (mole fractions 10.0001) of mercaptans and sulfides. Similarly, measurements involving N H 3 or CH3NH2 mixtures with H2S, CH3SH, and C2H5SHwere straightforward, as were those involving symmetric proton-bund dimer formation in (CH3)2Sand (C2H&S. However, the criteria which we establish for equilibrium measurements were not satisfied when we attempted to examine the thermochemistry associated with the CH3NH3+.n-C3H7SHand CH3NH3+-C2H5SCH3 ions, as well as clustering of CH3NH3+with higher homologues in the RSH and R2S series. These criteria are (at a given temperature) that (i) the equilibrium constant ( K ) be independent of ligand concentration over a reasonable range at a constant carrier gas pressure; (ii) in a mixture of fixed concentration, K should be independent of total pressure over a t least a factor of 4; and (iii) in a given measurement, the reactant/product ion ratio, from which K is derived, be constant for a sufficient period of time (usually a few ms) to ensure that equilibrium has, in fact, been achieved. The systems mentioned above failed all three criteria within the temperature range of this study (300-600 K). In addition, measurements of symmetric proton-bound dimer formation in (nC3H7)2Sand higher homologues failed criterion (iii) due to reactions involving other sulfides present as impurities. Dimercaptans, such as HSCH,CH,SH, and disulfides, such as CH3SCH2SCH3,as well as their higher homologues, either (2) Meot-Ner (Mautner), M.; Sieck,L. W. J. Am. Chem. Soc. 1983,105, 2956.
The Journal of Physical Chemistry, Vol. 89, No. 24, 1985 5223
InK
I 118
I
I
I
I
I
I
I
1
20
22
24
26
28
30
32
34
1
O ' YT
Figure 1. van't Hoff plots for cluster dissociation equilibria BH+.-A s BH+ + A; BH".-A as indicated.
"t +?
22
9
L
10
20
A PA
30
40
Figure 2. Enthalpy of dissociation (-AHD)vs. proton affinity difference for various RSH+-OH2 hydrates (values in kcal mol-').
fragmented extensively upon protonation or contained unacceptably high levels of impurities.
Results and Discussion I . SH+-O, NH+.-S, and SH+-S Bonds. van't Hoff plots for the clustering equilibria are shown in Figure 1, and the results are summarized in Tables I and 11. The SH+-O type bond is represented by the monohydrates of the sulfonium ions shown in Table I. As with other types of BH+.-A bonds, an inverse correlation is noted between APA = PA(B) - PA(A) and the bond dissociation energy A H D O . The correlation results, qualitatively, from the decreased efficiency of partial proton transfer from BH' to A as the difference in proton affinities increases. This was verified by the theoretical calculations of Desmeules and Allen,3 who found approximately linear correlations over wide ranges of APA for both second-row and third-row proton acceptors and donors, i.e., 2H+-.2, 2H+.-3, 3H+-2, and 3H+-3 combinations. The present results represent the 3H+.-2 case; i.e., SH+.-O, and the 2H+--3 case, Le., NH+--S. Using the present results only, the correlation equation AHDO
(SH+.**O)= 18.6 (f0.3) - 0.164 (f0.009) APA kcal mol-' (1)
is obtained for SH+-O systems (see Figure 2). This equation gives a value of 18.0 for the H2SH+--OH2complex, which unfortunately could not be investigated in the present study due to experimental difficulties. Two A H D O values have been reported for this system: 17.0 kcal mol-' via a direct measurement by Meot-Ner and Field! and 21.2 kcal mol-' by Hiraoka and Ke(3) Desmeules, P. J.;
Allen, L. C. J . Chem. Phys. 1980, 72, 4731
5224
The Journal of Physical Chemistry, Vol. 89, No. 24, 1985
Meot-Ner and Sieck
TABLE I: Thermochemistry' of Cluster Dissociation
BH'
A
AHDo
AGo(T, K)
a D o
H2SH+ CH3SH+ C2H5SH2+ n-C3H7SH2+ i-C,H,SH2' (CH3)2SH+ (CH,)(C,HS)SH+ (n-C3H7)2SH+ (CH,),SH*H20+
HZO HZO H20 HZO HZO H2O H2O HZO H20
SH+-O 17.0,c2Id 15.5 14.7 14.4e 1 3.9e 13.1 12.6' 12.2 11.6
NH4+ CH3NH3+ CH3NH3+ CH3NH3+
HZS H2S CH3SH CZHSSH
11.4 10.8' 13.4 14.6
17.8,' 2Sd (22) 22.8 (23) (23) 22.6 (23) 25.7 23.0
5.20 (476) 4.10 (446) 4 27 (418) 4.63 (347)
APA~
3.7 20.9 24.3 25.1 27.6 34.1 37.0 40.0
SH+-*N
16.7 (20) 22.1 19.6
5.40 (270)
37.1 43.9 26.7 23.3 0 0 0
" A H " , AGO in kcal mol-' (=4.18 kJ mol-'); ASo in cal mol-' K-' (4.18 J mol-' K-'). Error estimates from the standard deviations of slopes and intercepts of van't Hoff plots and from the precision of AGO measurements: AHo f l kcal mol-]; AGO f0.2 kcal mol-'; ASof 2 cal mol-' K-I. bAPA = PA(B) - PA(A). 'Reference 4. dReference 5. eFrom AGO as shown, using estimated ASoas shown.
stituents on the already reduced charge densities in R2SH+ is expected to be less than it is on the higher charge densities in R 2 0 H + . In the preceding article we examined the effect of q(H) on the AHDo/APA relation in a series of BH+-A complexes with varying BH' and a common ligand A.
TABLE 11: Thermochemistrya of Clustering Reactions
E=-
CH3NH:.
yl4Il
H***NCCH3 CH3N -H. -NCCH3 H*.*NCCH3 +/
'
58.2
Top values above arrow, -AH' in kcal mol-' ;bottom values, A S o in cal mol-' K-'. Italicized values are cumulative solvation a
enthalpies. Values in parentheses obtained from thermochemical cycles. From A G " ( 2 7 0 ) =-2.4 kcal mol-';-AS= 20 cal mol-' K - ' . estimated.
'
barle.5 These are shown as error bars on Figure 2. The latter value was obtained by examining the exchange reaction H2SH+.0H2 H 2 0 H30+.H20+ HIS, AHo = -6.7 kcal mol-', using 3 1.6 kcal mol-' for A H D O (H30+.H20)and PA(H2S) - P A ( H 2 0 ) = 3.7 kcal mol-'. It is interesting to note that the predicted value of 18.0 kcal mol-' (from the correlation equation) is identical with the value suggested by Kebarle for this complex in a review article6 which postdated the two experimental determinations. In any event, more data on SH+.-O ions with small APA values is clearly desirable. The depressed slope and intercept of the SH+.-O correlation line, if indeed correct, may be explained by the small charge density on the hydrogens in R2SH+ions. The difference between oxonium ions and sulfonium ions in this respect is illustrated by the difference in charge densities on the hydrogens in H30+' and H3S+,*0.328 and 0.150, respectively. The effect of alkyl sub-
+
-
(4) Meot-Ner, M.; Field, F. H . J . Am. Chem. SOC.1977, 99, 998. (5) Hiraoka, K.; Kebarle, P. Can. J . Chem. 1977, 55, 24. (6) Kebarle, P. Annu. Reu. Phys. Chem. 1977, 28, 445. (7) Almlof, J.; Wahlgren, V. Theor. Chim. Acta 1973, 28, 161. (8) Yamabe, T.; Aoyagi, To; Nagata, S.;Sakai, H.; Fukui, K. Chem. Phys. Lett. 1974, 28. 182.
+-qHZa 2r4
(2)
52. 6
2CH3CN 44. 8
q(H)pD r2
When the OH+-OH2 and SH+-OH2 correlations are compared, the dE/dq(H) terms must be similar since this is a function of the dipole moment and polarizability of H 2 0 which is identical in the two series. Also, even though the radius of S is larger than that of 0, the SH-OH2 distance is not necessarily substantially larger than the OH+-.0H2 distance. Therefore the principal difference probably lies in the dq(H)/dPA term. In view of the reduced charge density on R2SH*+,its variation with PA is also expected to be smaller. The electrostatic arguments in this and in the preceding paper may be summarized as follows. The ionic hydrogen bond is mostly electrostatic and, with polar ligands, E is dominated by the qpD/$ term and dE/dPA by the product (pD/?) (dq/dPA). When series with common ions and different ligands are compared, such as BH+--OH2 vs. BH+--N=CH, the more polar ligand should exhibit a higher correlation intercept (Le., a = A H D O for APA = 0) and a steeper slope 6. This trend is verified, for example, when one compares the NH+.-OH2 series ( a = 30.0, 6 = 0.26) with the NH+-.N=CH series ( a = 35.3, 6 = 0.34). Considering series with a common ligand but with a variety of ions, and assuming that larger q(H) also yields larger dq(H)/dPA values, eq 2 and 3 also suggest a larger slope and higher intercept. This is observed in comparing NH+-OH2 ( a = 30.0, 6 = 0.26) and OH+-OH2 ( a = 30.4, 6 = 0.30), both of which have similar intercepts and slopes, with SH+.-OH2 ( a = 18.6, 6 = 0.16). The present data does not afford a reliable correlation line for NH+-.S bonds. We note that for a given APA in the range of 25-45 kcal mol-', the A H D O values of NH,+-S and SH+-.O bonds are similar. This is in accord with the theoretical results of Desmeules and Allen, which suggest similar correlation lines for 2H+-.3 and 3Hf--2 type bonds. Our present data include A H D O for two symmetric SH+.-S dimers, Le., (CH3)2SH+--S(CH3)2 and (CH3SC2H5)H+(CH,SC,H,). Both give surprisingly high values of 26 kcal mol-'
J . Phys. Chem. 1985, 89, 5225-5235 when compared with AHDO for H2SH+-SH2 (14.1 f 1.3 kcal mol-', see Table I). The difference is unexpected, since in both symmetric R20H+-OR2 and R3NH+.-NR3 dimers with APA = 0 the values of AHDO are constant, Le., 30 and 24 kcal mol-', respectively, regardless of the identity of R. The reason for the different behavior of the sulfur dimer is not obvious. 2. Mutual Effects of Weak and Strong Ligands; The Clusters CH3NH3+.CH3SH.nCH3CN.The examination of a large series of clusters of the type BH+.nA showed that the mutual effects of ligand molecules A weaken the bonding to each subsequent A molecule added to the cluster. For all of the hydrogen-bonded clusters, the relative weakening effect of consecutive ligands turned out to be unexpectedly constant. Thus, for a large variety of clusters, the attachment energy of the second A molecule to the cluster is smaller by a factor of 0.75 f 0.05 than the attachment energy of the first ligand; Le., the first ligand molecule decreases the attachment energy of the second ligand molecule to BH+ by 25%. In all cases examined experimentally, the first and second ligand molecules were identical. It is of interest to examine the mutual effects of ligands when the two molecules are different, especially when one is a highly polar, strongly bonding ligand and the other is a weakly polar, weakly bonding ligand. The weak attachment energies of sulfur ligands to NH+- groups affords such a test. For this purpose we examined the mutual effects of attaching CH3SH and CH3CN ligands to CH3NH3+. The results are shown in Table 11. The difference between the bonding of the ligands is illustrated by the difference between the first attachment energies, 13.4 vs. 26.2 kcal mol-'. The data for the second clustering steps can be used to examine the mutual effects of the ligands. For example, to find the effect of CH3SH on the attachment energy of CH3CN to the cluster, we compare AHo for reactions 4 and 5 . The results show that
CH3NH3'
+ CH$N
-
5225 C H ~ N H ~ + S C H ~ AHo C N = -26.2 (4)
CH3NH3+*CH$H + CH3CN CH3NH3+.CH3SH*CH3CN AHo = -22.7 ( 5 ) -+
the addition of the weak ligand CH3SH decreases the attachment energy of C H 3 C N by a factor of 22.7/26.2 = 0.87, Le., AHo is lowered by 13%, while the prior addition of the strong ligand CH3CN lowers the attachment energy of CH3SH by a factor of 9.9/13.4 = 0.74, Le., by 26%. By comparison, the addition of one CH3CN molecule lowers the attachment energy of a second CH3CN molecule by a factor of 18.6/26.2 = 0.71, Le., by 29%. Thus, in the present cases the effect of the strong ligand on the attachment of a second ligand is proportionally similar whether the second ligand is weak or strong. However, if the first ligand is weak, its effect on attachment energy of a second, strong ligand is significantly smaller. This may be expected since the effect of a weak ligand on the charge distribution in CH3NH3+should be minimal. The differences between the effects of weak and strong ligands are seen to extend to higher clustering steps. Thus, -AHo for the addition of CH,CN to CH3NH3+CH3SHis larger by 4.1 kcal mol-' than the addition to CH3NH3+.CH3CN. By comparison, -AHo for the addition of C H 3 C N to CH3NH3+. CH3CN.CH3SH is still larger by 3.1 kcal mol-' than the addition to CH3NH3+.2CH3CN.Therefore, the bond-weakening effects of CH3CN and CH3SH in the mixed clusters are roughly additive. The present data seem to be the only available for protons solvated by more than two different ligands. Further study of mixed clusters is desirable, since many association ions occurring in nature may involve a variety of components.
Acknowledgment. This research was supported by the Office of Basic Energy Sciences, United States Department of Energy.
Intramolecular Electron Transfer in Donor-Acceptor Systems. Porphyrins Bearing Trinitroaryl Acceptor Group G. Bhaskar Maiya and V. Krishnan* Department of Inorganic and Physical Chemistry, Indian Institute of Science, Bangalore 560 01 2, India (Received: February 20, 1985)
Porphyrins bearing picryl group in ortho, meta, and para positions of one of the mesoaryl groups of tetraphenylporphyrin (TPP) have been synthesized in the free-base form. The metal [Cu(II) and Zn(II)] derivatives of the free-base picrylporphyrins (PPc) have been prepared. The broadened Soret absorption, the decreased optical absorbance values of Q bands, and the reduced singlet emission quantum yields of PPc indicate the existence of intramolecular interaction. The extent of this interaction is found to be greater than those observed for the intermolecular systems and varies with the position at which the picryl moiety is attached to the porphyrin as ortho > meta > para. The energies of the redox states, E(CT) of P'Pc-, calculated from the electrochemical redox potentials depend on the nature of the metal ion as free-base PPc > CuPPc > ZnPPc. The nature of intramolecular interaction between the picryl moiety and porphyrin unit is essentially TT (CT). Conclusive evidence for light-induced electron transfer in ZnPPc is presented from EPR studies. The decay profiles of the EPR signals vary with the position of the picryl moiety as ortho < para < meta. Computer simulation of structures substantiated by the 'H NMR results point out the restricted conformational freedom of the picryl moiety in ZnPPc. Arguments based on symmetry considerations of HOMO of the excited singlet state PPc and LUMO of picryl group indicate plane-to-plane orientation of the donor and acceptor in the picrylporphyrins.
Introduction Intramolecular systems comprising of an electron donor and an acceptor are important in the study of excited-state energytransfer and electron-transfer mechanisms. Investigations of systems containing porphyrins and metalloporphyrins as electron donors bear a direct relevance to photosynthetic research because of their close structural resemblance to the plant pigments. A 0022-3654/85/2089-5225$01.50/0
study of the model compounds comprising a donor porphyrin and an acceptor at prefixed distances and orientations is paramount to elucidate the influence of these parameters in the intramolecular photochemical electron transfer. Several interesting studies to mimic the reaction-center complex through the synthesis of covalently linked porphyrin-quinone moieties are reported in literature.' In all these systems, evidence has been presented for 0 1985 American Chemical Society
