J. Phys. Chem. 1982, 86, 2626-2632
2828
length b ( b = 30 A in Na-NH3) provided that .$ > b, a unimodal distribution characteristic of critical fluctuations should prevail. This model predicts particular sound propagation properties for e> and a more conventional critical behavior for t < The experimental results indicate, however, that for 6 > the sound propagation in Na-NH3 solutions can be understood on the basis of the theory for critical phenomena also applicable to nonconducting binary mixtures. Distinct features are exhibited for 6 < low2.Due to the
electronic properties of the metal-ammonia system mass diffusion may be coupled to electric polarization modes.70 Such a coupling is, however, not included in the theory discussed above.
Acknowledgment. This work was supported in part by the Deutsche Forschungsgemeinschaft. (70) R. A. Robinson and R. H. Stokes in "Electrolyte Solutions", Butterworths, London, 1959.
Theoretical Study on the Stability and Structure of H+(HCN), and M+(CH,CN), (M+ = H+, LI+, and Na+) Clusters K. Hirao,*t S. Yamabe,' and M. Sanot Deperhnent of Wmktry, C o i m of General Educatbn, Nagoya University, Nagoya, Japan, and Department of Chemistry, Nara University of Education, Nara, Japan (Received: February 9, 1982)
Ab initio molecular orbital calculations have been carried out on the gas-phase ion solvation of H+(HCN), and M+(CH3CN),with a basis set of double-{quality. Protonated HCN and CH3CNform shells of H+(HCN)2and H+(CH3CN),,respectively, and a third solvent molecule attacks not the central proton but a terminal hydrogen of their respective shells. For Li+(CH3CN),and Na+(CH3CN),,the most symmetrical structures are energetically favorable on account of the electrostatic stability. The asymmetry problem of the central proton in H+(HCN)2 is examined by CI method. Energy decomposition and difference density analyses are made to elucidate the nature of the bonding involved. The enthalpy change of the clustering is found to be reproduced well by the present method.
I. Introduction The properties of ion-molecule reactions have been a matter of longstanding chemical interest. This interest has been heightened by new gas-phase kinetic and thermodynamic data due to recent technical advances.'p2 A new type of chemical process such as clustering reactions can provide a great deal of information on the nature of solvation. I t has also been found that ion-molecular clustering is a particularly important process in the ionosphere. Most of the experimental data reported so far have been obtained within 10 years. Work is still continuing in many laboratories. However, there remains a lack of information about the true structure of ion clusters and the origin of the bonding involved. The bonding between ion and molecule is generally found to be weak relative to the normal chemical bond (less than 50 kcal/mol). In order to shed light on these problems, we have carried out a series of ab initio molecular orbital calculations on various types of ion ~ l u s t e r s . ~With these studies we want to derive a general picture of the gas-phase clustering reactions. In this work, we will report theoretical computations for ion clusters of H+(HCN), and M+(CH,CN), (M+ = H+, Li+, and Na+). The equilibria in the gas phase (M+(CH3CN),-, + CH3CN M+(CH3CN),, n = 1-5) were measured for M+ = H+,4Na+, K+, and Cs+p5and thermodynamical data such as the enthalpy change AHo,+, were obtained for the clustering reaction. Also, the temperature dependence of equilibrium constants on the reaction H+(HCN),_' + HCN H+(HCN),, n = 1-5, is determined.4 On the basis of the change in AHon-l,nwith the
-
-
Nagoya University.
* Nara University of Education. 0022-365418212086-2626$0 1.2510
increase of n, some structures of ion clusters were suggested. For instance, the large drop from to AH0293 in H+(HCN),l + HCN H+(HCN), was thought to be attributed to a linear structure, because the third HCN molecule would add to the outer layer in the one-dimensional array of a H-N-C.. . H .N-C-H shell. The similar falloff of AHo,,+, was obtained for H+(CH3CN),+ CH3CN H+(CH,CN),. Different from such a trend in the protonated HCN and CH,CN, the stepwise solvation around alkali ions gave a slow decrease of AHon-1,,. In view of these thermochemical data, it is tempting to make clear the structure of these cluster ions and to analyze the difference of the mechanism of the clustering process. Although the equilibrium of the Li+(CH,CN),-, CH3CN Li+(CH3CN),was not studied experimentally on account of the requirement of high (>700 K) temperature: we determine the structures of the Li+ clusters to compare them with other ion clusters.
-
+
-
11. Method of Calculation For the SCF calculation of H+, HCN, and CH3CN, the 4-31G basis set is used.6 For Li+ and Na+, we used (1) Kebarle, P. "Ions and Ion Pain in Organic Reactions"; Szwarc, M., Ed.; Interscience: New York, 1972; Vol. 1, p 27; "Ion-Molecule Reactions";Franklin, J. L., Ed.; Plenum Press: New York, 1972; Vol. 1, p 315. (2) Good, A. Chem. Rev. 1975, 75, 561. (3) Yamabe, S.;Hirao, K.; Kitaura, K. Chem. Phys. Lett. 1978,56,546. 1980,102,2268. Yamabe, S.; Osamura, Y.; Minato, T. J. Am. Chem. SOC. Yamabe, S.; Hirao, K. Ibid. 1981,103,2176. Hirao, K.; Yamabe, S. Chem. Phys. Lett. 1981, 79, 279. Yamabe, S.; Hirao, K. Ibid. 1981, 84, 578. Hirao, K.; Yamabe, S. 'Theoretical Study on the Structure and Stability of H,+ and Hn- (n = 3, 5 , 7 , 9 , 11, 13)", to be submitted for publication. (4) Meot-Ner, M. J. Am. Chem. SOC. 1978, 100, 4694. (5) Davidson, W. R.; Kebarle, P. J . Am. Chem. SOC.1976, 98, 6125. (6) Hehre, W. J.; Lathan, W. A.; Ditchfield, R.; Newton, M. D.; Pople, J. A. QCPE, No. 236 (1973), GAUSSIAN'IO. It is extended to GAUSSIAN320 by Kitaura and Morokuma at the Institute for Molecular Science.
0 1982 American Chemical Society
The Journal of Physical Chemistry, Vol. 86, No. 14, 1982 2627
H+(HCN), and M+(CH,CN), Clusters
TABLE I: Total Energy ( E T )and Stabilization Energy ( A E , , - ~ ,for ~ ) the Protonated Cluster H+(HCN), Calculated with SCF of the 4-31G Basis Set AEn-l,n,
n
cluster
geometry
E T , au
1 2
HCN H+(HCN) H+(HCN),
linear linear linear (asym)
-92.731 928 -93.022 351 -185.800 115 -186.170 629' - 185.799 409 -186.170 805' -278.562 017 -278.545 839 -371.316 988 -464.067 086
linear (sym) 3
H+(HCN),
4 5
H+(HCN), H+(HCN),
linear T-shaped linear linear
'
kcal/mol
AHnn-r,nla
kcal/mol
-182.3 -28.8 -30.2' -28.3 -30.3' -18.8 -8.7 -14.5 -11.4
-30.0
-13.8 -11.8 -9.2
a The experimental data are taken from ref 4. Calculated values on the SCF + CI level. CI energy on the HCN and H+(HCN) in their respective equilibrium geometry with intermolecular distance of 250 A is -186.122 533 au.
Huzinaga's [4s] and [4s2p] basis sets.' The calculated dipole moments of HCN and CH3CN are somewhat too high (HCN: 3.214 vs. 2.885 D exptl.; CH3CN: 4.088 vs. 3.92 D exptl). The present basis set tends to overestimate the polarity of molecules and consequently overemphasizes the contribution of electrostatic interaction. In spite of the limitations imposed by use of the smaller basis set, we are obligated to use it since the extensive basis sets are not practical for studies of larger systems. According to our experience: the calculated stabilization energies with the present basis set are in agreement with the experimental values of AHon-l,n with a few kcal/mol deviation. Thus, the results are quite appropriate for qualitative discussions. Confiiation interaction (CI) using all single and double substitutions has been done for H+(HCN)2to examine the potential of the proton movement. The l a orbitals are constrained to be doubly occupied. For H+(HCN), (n 5 3) and M+(CH3CN)(M+ = H+, Li+, and Na+), all of the geometric parameters are fully optimized. Then, the intermolecular distance in ion clusters, H+(HCN)( and H+(HCN)&,is determined with HCN and H+(HCN)3being kept at their optimum geometry. For M+(CHgCN),(n 1 2) clusters, based on the results of the negligible geometrical distortion of CH3CN, only the intermolecular distance and angle (i.e., orientation) in ion clusters are changed to seek the lowest energy with CH3CN being fixed at its optimum geometry. 111. H+(HCN), Calculated total energies (ET)of H+(HCN), clusters are Their optimcollected in Table I together with ized geometries are shown in Figure 1. As is anticipated, linear structures of H+(HCN) and H+(HCN)2are most stable. The N. *H+distance is enlarged during H+(HCN) + HCN H+(HCN)2,because H+ is shared by two u lone-pair electrons of N atoms in H+(HCN)2. In H+(HCNI2,the position of H+ is asymmetric between two N atoms on the SCF level. However, the energy change is very small for the proton movement. According to the CI calculation, the proton turns out to stay at the central position. An energy difference of this order of magnitude (0.1 kcal/mol) is beyond the accuracy of the present study, and our investigation essentially proves the great mobility of the central proton. A similar change of geometry by the absence or presence of the correlation correction was reported for H5+ (C,, or Dzd).8 Two models, linear and T-shaped, of H+(HCN)3are examined. Which site, the central proton or the terminal H of H+(HCN)2,is better for the third HCN? The linear
-
~
(7) Huzinaga, S. "ApproximateAtomic Functions";The University of Alberta: Alberta, Canada, 1971. (8)Ahlichs, R. Theor. Chim. Acta 1975,39, 149.
1
3
linear
linear
1-shaped
-0 ( 1 6
4
linear
5
linear
+@.@
.____ - - o - - m----. 0Q.o i 110
-
._.__ 2 o c
___
2 0 9
Flgure 1. Optimized geometries of H+(HCN), clusters. Bond lengths are in angstroms and angles in degrees. Values in parenthesis are net atomic charges. An empty circle denotes a hydrogen atom.
Scheme I
-H
-no
moderately positive, electronic cloud which gives the exchange repulsion
more cationic, but n electronic cloud prevents the approach of the third HCN
H+(HCN)3is found to be much more preferable than the T-shaped one? In the latter model, the r-electron density presented by two neighboring N atoms blocks the approach of the third HCN along the axial direction (see Scheme I). This is the so-called exchange repulsion, which arises from the overlap of charge densities in connection with the (9) All of the geometric parameters are fully optimized for the Tshaped structure of H+(HCN)* We find that the NH+N bond of H+(HCN)Z is somewhat bent with a deviation from linearity of about 4.7". However, the minimum energy is found to be stable only by 0.1 kcal/mol with respect to that of the rigid T structure where the NH+N bond is frozen at M O O .
2028
The Journal of phvsical Chemistry, Vol. 86,No. 14, 1982
Hirao et at.
a
C
Fkure 2. 0 nsity difference maps for H+(HCN), (n = 1, 2, and 3). The full lines and the broken lines denote the increase and the decrease of-the charge density relative to the separated nonlnteractlng molecules. The noninteracting species in cases a-c are H+ + HCN, H+(HCN) HCN, and H+(HCN)*
+ HCN, respectively.
+
The unit of the contour lines is bohr3.
Pauli principle. To see the difference of the two models more clearly, we decomposed A E 2 , 3 into five energy components (ES, ESX, PL, CT, and MIX) according to the Kitaura and Morokuma schemelo (see Table 11). ES is an energy term of the electrostatic interaction. ESX is the exchange interaction. P L means a polarization (i.e., inductive) stability. CT is a charge-transfer energy term, and MIX is a coupling term of these various components. ES and ESX are of the first order in the sense of the perturbation. CT and PL are of the second order, and MIX is of the third order. In the table N ES (-18.81 N -19.81 kcal/mol in the linear model and -8.65 N -8.24 kcal/mol in the T-shaped model), which indicates that the other four terms related to the distortion of the electronic cloud cancel one another. Thus, the stability of the cluster comes primarily from the electrostatic attraction. In the T-shaped model, the optimal H+...N intermolecular distance (2.972 A) is much larger than the corresponding distance of the linear model (1.896 A). If the former distance is enforced to 1.896 A , A E 2 , 3 turns out to be positive (repulsive), as shown in the third column of the table. Particularly, ESX, which originates from the exchange repulsion between the N-. .N a-electron density and the lone-pair electrons of HCN, becomes extremely large. The N H charge-transfer (CT) interaction, although relatively small (-3.45 kcal/mol in the linear model), brings
-
(IO) Morokuma, K. Acc. Chem. Res. 1977,10,294. Kitaura, K.; Morokuma, K. Znt. J. Quantum Chem. 1976, IO, 325.
TABLE 11: Energy Decomposition of the Interaction between H+(HCN), and HCN by Kitaura and Morokuma Schemea energy component A E 2 , 3in Table I
ESd ESXe PLf CTP
MIX^
linear* -18.81 -19.81 17.86 -4.58 -3.45 -8.83
T-shaped T(R(H+...N)= shapedb 1.896 A ) C -8.65 -8.24 2.21 -0.95 -0.64 -1.03
14.22 -29.16 111.96 -6.18 -9.57 -52.83
A E , = ES t ESX + PL + CT + MIX. Values in kcal/ mol. b'Models of H+(HCN), shown in Figure 1. Arbitrary model with the same intermolecular distance as in the linear model, Electrostatic attraction. e Exchange repulsion. Polarization stability. g Charge transfer. h Coupling term.
about a partial covalent nature to the N-H bond. The negligible change in the intramolecular distance of HCN by forming the ion cluster is consistent with the fact that the charge migration occurs from the lone-pair orbital which is unrelated to the H-C-N bonding region. In Figure 2, we displayed contour maps of the density differences for H+(HCN), ( n = 1, 2, 3). These densities are obtained by subtracting the superimposed densities of undisturbed H+(HCN),-,. -.HCN from charge densities of H+(HCN)n. The full lines and the broken lines denote the increase and the decrease of electron densities, re-
H+(HCN), and M+(CH3CN), Clusters
spedively. These figures provide a detailed picture of the net reorganization of charge densities through the CT and polarization. We can see from Figure 2a that the proton is trapped moderately by lone-pair electrons with little reorganization of the density of HCN. Figure 2b shows that the negative charge is transferred from the nitrogen lone-pair orbital to the central proton in the new N. ..H+ bond and that the already formed H+.-.N bond at the right is polarized. Consequently, one expects stronger attractive forces (CT and Coulomb interactions). The density reorganization obtained in Figure 2c is caused by the polarization interaction. It must be noted in Figures 1and 2 that the decrease of charges at the terminal hydrogen enhances the cluster extension along the one-dimensional direction.
The Journal of Physical Chetnlstty, Vol. 86, No. 14, 1982 2629 r
'P9
ro 221
1
2
linear
linear
;:E
w-0 2 6 5 '
(H' attack)
linear (CH attack)
Y,*D
(,
,vi.
*c
2101
1.100
IV. M+(CH&N), In Table 111, energies of M+(CH3CN), (M+ = H+, Li+, and Na+) clusters are summarized together with In Figure 3, geometries of H+(CH3CN), (n = 1, 2, 3) are shown. Mol,, of H+(CH,CN), is almost the same as that of H+(HCN)* The similarity of the structures of H+(HCN), and H+(CH,CN), is observed in Figures 1 and 3. We considered another possibility that a second CH,CN attacks a terminal hydrogen of the first CH,CN, because H+ makes the terminal hydrogen of the CH3CN bound to it more positive. Two models, the linear CH attack and CH3 attack, give considerable stabilities but they are less stable than the first model (linear H+ attack). It is interesting to ask whether H+(CH,CN), is of the T-shaped geometry or of the linear type. It appears that linear models are not so preferable as in H+(HCN)3because of the bulky nature of the methyl group of CH3CN. However, calculations show that the linear models (CH attack and CH, attack) are again more favorable than the T-shaped one. The appreciable acidity of the methyl hydrogen in H+(CH,CN), may call for the lone-pair electrons of the nitrogen of the third CH3CN. Although the central proton of H+(CH,CN), with very positive net charge appears to attract the third CH3CN more strongly, the exchange-repulsion block of the r-electronic cloud renders the T-shaped model unfavorable. In the two linear-type models (CH attack and CH3 attack) with a comparable stability, the former is more stable than the latter. H+(HCN)2and H+(CH,CN), may be regarded as inner shells of the proton. At the shell, the most reactive site is not the central proton but the terminal hydrogen for further clustering. Methyl substitution causes a decrease of - h E 2 , 3 by 6.9 kcal/mol (vs. 4.5 kcal/mol exptl). This difference is assigned to that of the acidity (the degree of positive nature) of the terminal hydrogens. In Figures 4 and 5, geometries of Li+(CH3CN), and Na+(CH,CN), are exhibited, respectively. These clusters are found to have quite different structures from those of protonated HCN and CH3CN clusters. That is, symmetrical structures are calculated to be most stable. The gradual decrease of AEn-l,nwith the increase of n reflects this clustering mechanism. We have checked in Li+(CH3CN), the possibility of the structure of Li+... NCCH,.-.NCCH3. But it is calculated to be less stable than the H,CCN. ..Li+. -NCCH3 structure. In Li+(CH3CN), and Na+(CH,CN),, Li+ and Na+ are a t the center between two nitrogen atoms. Li+ and Na+ ions present an essentially spherical static electric field to the solvent molecules. They play a role as a mere electrostatic origin. The net charges in the parentheses in Figures 4 and 5 indicate that there is little electron transfer from the solvent molecule to the central Li+ and Na+. The stability of these clusters comes primarily from the elec-
-
T -shaped
A
Figure 3. optimized geometries of H+(CH,CN), clusters. Bond lengths are in angstroms and angles in degrees. Values in parentheses are net atomic charges.
trostatic attraction and polarization. These effects are not especially sensitive to the angle 0(N...Li+. ..N). So there is a small barrier with respect to the bending. We calculated Li+(CH,CN), at 0 = 180'. With the decrease of 0 from 180° to go', the steric hindrance increases but the exchange repulsion decreases. Thus, in these symmetrical clusters, the distance between the central ion and the solvent molecules is determined mainly by the electrostatic attraction and the orientation of the solvent molecules determined by the steric hindrance and exchange repulsion among the solvent molecules. To examine more precisely the change in the charge distribution on cluster formation, three difference density maps are drawn in Figure 6. They may be classified distinctly into two groups according to the criterion of whether the M+.-.N density increases or decreases relative to the unperturbed ion-molecule. While the H+-..N covalent bond is formed, Li+...N and Na+---Nare antibonding due to the absence of the CT interaction. Li+ and Na+ have inert close shells and so they refuse the acceptance of electron density. The decrease of the density
2890
TABLE 111: ET and AEn-l,n for t h e Reaction M+(CH,CN),, n
cluster
1 2
CH,CN H+(CH,CN) H+(CH,CN),
3
H+(CH,CN),
1 2
Lit Li+(CH,CN) Li+(CH,CN),
3 4
Lit( CH,CN), Li+(CH,CN), Na+ Na+(CH,CN) Na+(CH,CN), Na+(CH,CN), Na+(CH,CN),
1 2 3 4 a
Hirao et ai.
The Jounal of phvsicel Chetnistty, Vol. 86, No. 14, 1982
geometry linear linear (H+ attack) linear (CH attack) linear (CH, attack) linear (CH attack) linear (CH, attack) Tkhaped linear linear (Li+ attack) linear (CH attack) linear (CH, attack) trigonal tetrahedral linear linear (Na+ attack) trigonal tetrahedral
+ CH,CN
-+
M+(CH,CN), Calculabd with SCF Wave Function
ET, au
-131.728 -132.044 -263.821 -263.796 -263.794 -395.668 -395.567 -395.562 -7.236 -139.040 -270.832 -270.785 -270.784 -402.602 -534.358 -161.669 -293.455 -425.231 -556.996 -668.751
262 351 305 635 430 430 085 829 300 757 312 554 508 689 200 699 016 867 486 235
AEn-1,m
AH'n-
kcal/mol
1.
nr
kcal/mol
-198.4 -30.6 -15.1 -13.7 -11.9 -11.0
-191.5a -30.2 -9.3
-8.3 -47.8 -39.7 -10.4 -9.7 -26.4 -17.1 -35.8 -30.5 -22.8 -16.6
-24.4 -20.6 -14.9
* 0.3b i 0.5 * 0.2
Taken from ref 4. b Taken from ref 5. 1
linear
1
linear
Nai+0.9511
2
3
trigonal
NI-0.6391 C i i O . 3511
2,106
4
8
linear
(C % attack)
3
linear (t& attack)
0
B
trigonal
Flguro 5. Optimlzed geometrles of Na+(CH,CN), clusters. Bond lengths are in angstroms and angies in degees. Values in parenthese are net atomic charges.
and typical electron redistribution is observed in the HeHe interaction." The figure indicates that the stability of Li+(CH&N), and Na+(CH3CN), comes from electrostatic and polarization interactions.
+
Agw 4. optknired g"mes of U+("), dusters. Bond length are In angstroms and angles In degrees. Values in parentheses are net atomic charges.
-
-
amid Li+. *Nand Na+. .N, the increase around Li+ and Na+ nuclei, and the contraction of the nitrogen lone-pair density are all caused by the exchange repulsion. A similar
V. Discussion on the Shell Surroundings vs. the Most Symmetrical Clusters The calculated of H+(HCN), and M+(CH3CN), is, as a whole, in good agreement with LVion-l,n. A small overestimation is ascribed to the large polarity of HCN and CH3CN as mentioned in section 11. In general, the en(11) Bader, R. F.; Chandra, A. K. Can. J. Chem. 1968,46,953.
H+(HCN), and M+(CH,CN), Clusters
The Journal of Physical Chemistry, Vol. 86, No. 14, 1982 2631
a
b
C
Figure 6. Density difference maps for H+(CH,CN), Li+(CH,CN), and Na+(CH,CN). The full lines and the broken lines denote the increase and the decrease of the charge density, respectively. The unit of the contour lines Is bohr.-,.
thalpy change of the cationic cluster formation may be reproduced well by the SCF wave function with such accuracy as the 4-31Gbasis set. In consideration of the electron correlation correction, the zero-point vibrational energy, and the temperature effect, a few kcal/mol deviation between AE,-l and is a tolerable value on the SCF level. H+(HCN)2and H+(CH,CN), may be re-
garded as inner shells of the proton. These shell complexes are strongly electrostatic in nature but still include an essential charge transfer. That is, the formation of the shell involves relatively tight partially covalent bonds. Outside the shell of H+(HCN)2,the third, fourth, and fiith HCN's are linked successively at each reactive terminal hydrogen, which constitute the linear chain cluster. Sim-
2632
J. Phys. Chem. 1982, 86,2632-2640
ilarly, in the shell of H+(CH,CN),, a methyl hydrogen welcomes the approach of the nitrogen lone-pair electrons of the third CH3CN. Therefore, in these shells the most reactive site for the third body is not the central proton, but the terminal hydrogen on account of the exchangerepulsion block at the former position. In these clusters, the third and more neutral molecules may be regarded as the “surroundings” which are nonequivalent to the first and the second ones. In Li+(CH,CN), and Na+(CH,CN),, the most symmetrical structures are found to be favorable. There is no
tight covalent and directional bonding. This is in contrast to the case of H+(CH,CN),. Different from the proton, Li+ and Na+ have the ionic radii and serve merely as the origin of the electrostatic field due to their inert electronic structure. Acknowledgment. We thank the Institute for Molecular Science for allotment of the CPU time of the HITAC M-200H computer. This study is supported in part by a Grant-in-Aid for Scientific Research from the Japanese Ministry of Education, Science, and Culture.
Electrochemistry in Ordered Systems. 2. Electrochemical and Spectroscopic Examination of the Interactions between Nitrobenzene and Anlonic, Cationic, and Nonionic Micelles Gregory L. McIntire,+ Donna Mark Chiappardi, Robert L. Casselberry, and Henry N. Blount’ Center for Catalyiic Science and Technolcgy and Brown Chemical Laboratoty, The Universi!y of Delaware, Newark, Delaware 1971 1 (Received: August 13, 1981; In Final Form: February 8, 1982)
The reductive electrochemistryof nitrobenzene (NB) to the corresponding anion radical (NB-.) and dianion has been examined in anionic, cationic, and nonionic micelles. The observed voltammetry of NB in these systems has been related to the nature of the interactions of NB and NB-with the respective micelles. These interactions have been interpreted in terms of models of various micelle/substrate interactions which were tested by using values of the diffusion coefficient of NB in each micelle system over a range of NB concentrations. Results indicate that the stability of NB-. in anionic micelles arises from a relatively strong surface interaction. The stability of the nitrobenzene anion radical in nonionic micelles suggested by voltammetry and supported by ESR also derives from a surface interaction. Cationic micelles, however, provide a pseudophase into which NB partitions such that NB-. resulting from reduction of NB is concentrated within the small volume of the micelle and therein undergoes homogeneous chemical reactions to yield phenylhydroxylamine. These results point to the importance of understanding potential interactions between substrates and micelles in describing the behavior of micelle-solubilized substrates.
Introduction Micelles, dynamic aggregates of amphiphilic molecules, possess regions of hydrophilic and hydrophobic character in which normally water-insoluble species may be solubi1ized.l This unique property of micellar solutions, as well as the short-range order afforded to the reaction environment by micelles, has given rise to the growing use of these media in a broad spectrum of applications. The microscopic order of the micelle system is known to be of significance in biological,2synthetic,, and energy-transfer systems4 wherein the solubilized species can, under appropriate conditions, serve either as an electron acceptor or as a n electron donor. There are a variety of interactions which may be operative between solubilized substrates and host micelles including electrostatic a t t r a ~ t i o nsurface ,~ adsorption: pseudophase extraction,7 and substratelamphiphile coassembly to yield a unique micelle composit i ~ n .Understanding ~ ~ ~ the specific nature of the interaction between a solubilized species and the host micelle is critical to describing the impact of the micellar microenvironment on the functional behavior of the substrate. Work in these laboratories has addressed the effects of micelle solubilization on the redox behavior (formal pot Research Laboratories, Eastman Kodak co., Rochester, NY 14650.
tentials) and reactivities of the one-electron oxidation and reduction products (radical ions) of solubilized sub~trates.~JOA previous report has detailed the redox behavior of sodium dodecyl sulfate (SDS) micelle-solubilized 10-methylphenothiazine (MPTH).g The product of the monoelectronic oxidation of MPTH, the cation radical, associates both with free dodecyl sulfate anions (DS-) and with the micellar phase.g Observation of an interaction between the cation radical and the anionic dodecyl sulfate suggested that negatiuely charged species, arising from the reduction of solubilized neutral substrates, might prefer(1) Fendler, J. H.; Fendler, E. J. ‘Catalysis in Micellar and Macromolecular Systems”; Academic Press: New York, 1975. (2) For example: (a) Helenius, A.; Simons, K. Biochim. Biophys. Acta 1975,415,29. (b) Mukherjee, T.; Sapre, A. V.; Mittal, J. P. Photochem. Photobiol. 1978, 28, 95. (3) Cordes, E. H. Pure Appl. Chem. 1978,50, 617. ( 4 ) Brugger, P.-A.; Infelta, P. P.; Braun, A. M.; Gratzel, M. J. Am. Chem. SOC.1981,103, 320. (5) Quina, F. H.; Politi, M. J.; Cuccovia, J. M.; Baumgarten, E.; Martins-Franchetti,S.M.; Chaimovich, H. J. Phys. Chem. 1980,84,361. (6) Eriksson, J. C.; Gillberg, G. Acta Chem. Scand. 1966, 20, 2019. (7) Bunton, C.A.; Romsted, L. S.;Savelli, G. J. Am. Chem. SOC.1979, 101, 1253. (8) Funasuki, N.; Hada, S. J. Phys. Chem. 1980,84, 736. (9) McIntire, G. L.;Blount, H. N. J . Am. Chem. SOC.1979,101,7720. (10) McIntire, G. L.; Blount, H. N. In “Solution Behavior of Surfactants: Theoretical and Applied Aspects”; Mittal, K. L, Fendler, E. J., Eds.; Plenum Press: New York, 1982.
0022-3654/82/2086-2632$01.25/00 1982 American Chemical Society
