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(26) Makino, K.; Suzuki, N.; Moriya, F.; Rokushika, S.; Hatano, H. Chem. Lett. 1979, 675. Makino, K. J. Phys. Chem. 1980, 84, 1012. (27) Kotake, Y.; Kuwata, K.; Janzen, E. G. J. Phys. Chem. 1979, 83, 3024. (28) Stowell, J. C. J . Ofg. Chem. 1971, 31, 3055. (29) Makino, K.; Suzuki, N.; Moriya, F.;Rokushika, S.;Hatano, H. Anal. Lett. 1980, 13,311. (30) Hogeveen, H.; Gersmann, H. R.; Praat, A. P. Recl. Tfav. Chim. Pays-Bas 1967, 86, 1063.
(31) Hoffman, B. M.; Earns, T. B. J. Am. Chem. Soc. 1989, 91,2169. (32) Kirino, Y. J. Phys. Chem. 1975, 79, 1296. (33) Gilbert, 6 . C.; Trenwith, M. J. Chem. SOC.,Perkin Trans. 21973, 1834. (34) Lagercrank C.; Setaka, M. Acta Chem. Scand., Sect. B 1974,28, 619. (35) Lagercrantz, C.; Forshult, S. Nature (London) 1968, 278, 1247. (36) Smith, P.; Fox, W. M.; McGinty, D. J.; Stevens, R. D. Can. J. Chem. 1970, 48, 480.
Spectrophotometric Determination of N-Methylformamide Autoprotolysis Constant T. Oncescu, A . 4 . Oancea, Polltechnicai Institute of Bucharest, Instltute of Chemlstry, Depattment of Physical Chemistry, Bd. Republicii 13, 7003 1 Bucharest, Romania
and L. De Maeyer" Max-Planck-Institutfuer biophysikaiische Chemie, 0-3400 Qoettingen, Germany (Received:March 17, 1980)
The autoprotolysis constant of N-methylformamide obtained from spectrophotometric measurements on M2harmonizes well with p-nitrophenol and p-nitrophenoxide solutions is reported. The value K,, = 1.8 X the autoprotolysis constants of other related compounds.
The study of proton-transfer reactions between different alkyl amines and nitrophenols in N-methylformamide (NMF) which is in progress in our laboratory indicated that the solvent itself is involved in acid-base equilibria. This participation is not unexpected since the N-monoalkylated amide solvents are known to possess both basic and acidic properties,l as proved primarily by amide hydrochloride precipitation when hydrogen chloride was bubbled into solvent and by hydrogen evolution accompanied by potassium salt formation when metallic potassium is added to solvent. In our experiments the acid-base properties of NMF were revealed by analysis of electronic spectra of p-nitrophenol (PNP) and sodium p-nitrophenoxide (NaPNP) in NMF. Starting from either PNP or NaPNP solution, both phenol and phenoxide absorption bands were present in the spectrum. This behavior we associated with the presence of basic and acidic impurities, but after thorough purification it was found that PNP was sufficiently acidic to protonate and NaPNP sufficiently basic to deprotonate NMF, within a concentration range suitable for spectrophotometric measurements. As the acidic and basic impurities compete with the solvent in acid-base equilibria, special attention was paid to solvent purification. Experimental Section Solvent. NMF (Fluka purum) kept over molecular sieves of 4 A was treated with charcoal (Merck) and thoroughly stirred, filtered, refluxed over BaO until the vapor temperature remained constant (-318-320 K), and distilled under reduced pressure (410
0.400 0.340 0.485 0.735 0.380 0.410 0.520 0.545 0.685 0.345 0.713 0.485 0.620 0.470 0.575 0.495 0.508 0.405 0.170 0.247 0.320 0.380 0,191 0.320 0.377 0.437 0.483 0.500 0.366 0.353 0.302 0.582 0.268 0.342 0.313 0.346 0.366 0.432 0.496 0.557 0.176 0.283 0.373 0.438 0.491 0.268 0.355 0.465 0.180 0.485 0.665 0.880 0.215 0.260 0.330 0.656 0.334 0.551 0.550 0.477 0.532 0.385
absorption measurements at 320 and 410 nm. The experimental results carried out for 11 samples of solvent purified independently, in solutions of different values of ml and m2, are given in Table I. These results were used to calculate K,, Kb, and hence K , from the linear regression of eq 10, giving K , = (6.96 f 0.31) X lo4 M and Kb = (2.60 f 0.12) X lo4 M with a correlation coefficient r = 0.94 for 62 points. Accordingly K , = (1.81 f 0.12) X M2. It is well known that water molecules possess both acidic and basic properties and consequently can be involved in
1,
cm
0.5 1.0 0.5 1.0 1.0 1.0 0.5 0.5 0.5 0.5 0.5 0.5 0.5 0.5 0.5 0.5 1.0 1.0 1.0 1.0 1.0 1.0 1.0 1.0 1.0 1.0 1.0 1.0 0.5 0.5 1.0 1.0 1.0 1.0 1.0 1.0 1.0 1.0 1.0 1.0 1.0 1.0 1.0 1.0 1.0 0.5 0.5 0.5 0.5 0.5 0.5 0.5 0.1 0.1 0.5 0.5 0.5 0.5 0.5 0.5 0.1
0.1
R 5.050 0.779 4.277 2.482 0.274 0.180 0.744 0.664 0.713 0.717 0.849 0.521 0.658 0.363 0.367 0.490 0.796 0.408 0.341 0.759 0.887 1.029 1.519 1.953 2.694 3.176 3.643 4.292 1.161 1.241 0.530 0.560 0.724 0.878 1.620 1.875 2.190 2.291 2.698 3.008 1.279 1.785 2.113 2.474 2.808 1.265 1.182 1.099 1.511 0.322 0.425 0.425 0.451 0.436 0.833 0.844 0.811 0.647 0.650 0.649 0.857 0.601
Ra 25.502 0.607 18.289 6.159 0.075 0.032 0.554 0.441 0.508 0.514 0.721 0.271 0.433 0.132 0.135 0.240 0.633 0.166 0.116 0.576 0.786 1.059 2.308 3.816 7.259 10.088 13.272 18.426 1.349 1.540 0.281 0.313 0.525 0.771 2.625 3.51 8 4.796 5.251 7.279 9.050 1.637 3.185 4.467 6.122 7.887 1.601 1.398 1.210 2.284 0.104 0.181 0.180 0.204 0.190 0.694 0.713 0.658 0.419 0.422 0.421 0.734 0.361
105~(~,m) 16.694 1.095 17.142 7.538 -0.118 - 0.055 -0.523 - 0.329 -1.097 0.324 - 1.048 0.213 -0.576 0.394 0.276 0.241 1.670 0.682 0.240 - 0.587 - 1.040 -1.662 1.199 2.583 4.197 5.736 7.271 8.869 -0.905 - 1.100 0.662 1.346 0.802 1.241 2.095 2.682 3.312 4.091 5.530 6.924 0.930 2.087 3.258 4.478 5.698 -0.439 - 0.967 - 1.409 0.312 0.475 0.184 -0.136 -0.559 -0.833 0.189 -1.755 0.211 - 0.289 - 0.295 - 0.037 2.488 0.435
acid-base equilibria in nonaqueous solvents. Although water is present in a relatively high concentration in NMF (30-70 ppm, i.e., 2 X 109-4 X M), the acid-base interaction between PNP or NaPNP and NMF is not influenced by its presence. This affirmation is supported by the fact that R remained constant when the water to 10-1 M. content was increased from Discussion The autoprotolysis constant of NMF obtained in the present study is large if compared to those of amphiprotic
The Journal of Physlcal Chemistry, Vol. 84, No. 23, 1980 3093
N-Methylformamide A,utoprotolysis Constant
TABLE 11: Some Molecular and Bulk Properties of Simplest Amidesk property mol wt g cmm3 p(298 P , D (ref 1 ) D(298 K), ref 1 ET, cm-’ ref 13 P W 2 9 8 K) pKs*(2139 K) PKc(PNP) pKa*(PNP) PKa,(SH2*)aq
B
NMF 59.07 0.9988 3.82 185.5 54.1 10.74’ 13.205.16’ 6.39 0.15 24
FA 45 1.134 3.68 111.3 56.6 16.8‘ 19.6 8P 9.40 -0.25 14
-
NMA 73.1 0.9492 (313 K ) 3.71 178.9 52.0 -10 (313 K)d 12.2 (313 K ) 8.8g
9.91 -0.11 10
DMF 73.1 0.944 2.82 37.2 43.8 -1V 20.2 10.9h 12.0 -0.19 8.3
AA 59.07
10.5 (367 K)i -12.9 (367 K)
-
-0.40 6.7
H2O 18 0.9982 1.84 78.5 63.1 14.0 17.49 7.15” 8.89 -1.74 1
a Cf. A. J. Parker, Chem. Rev., 69, 1 (1969). G. Wada and T. Takenaka, BUZZ. Chem. SOC.Jpn., 44, 2877 (1971). F. Cf. ref 1, quoted as unpublished data. e M. Teze and R. Schall, Bull. H. Verhock, J,A m . Chem. SOC., 58, 2577 (1936). SOC.Chim. Fr. 13721 (1962). f Cf. G. Charlot and B. Tremillon, “Les Reactions Chimiques dans les Solvents et les Sels B. Fondus”, Gauthier-Villars, Paris, 1963. g M. Gosselet, S. Sibille and J. Perichon, Bull, SOC.Chim. Fr., 249 (1975). W. Clare, D. Cook, E. C. F. KO, Y. C. Mac, and A. J. Parker, J. A m . Chem. SOC.,88, 1911 (1966). Our measured values. NMF = N-methylformamide; FA = formamide; NMA= I’ G. Jander and G. Winkler, J. Znorg. NucZ. Chem., 9, 32 (1959). N-methylacetamide; DMF = dimethylformamide;AA = acetamide.
curred from the nitrogen atom to the oxygen atom. Prosolvents like water and alcohols, and larger than that of tonation at the oxygen atom would not perturb very much +relatedamidic comlp0unds.4~~ Like other amidic solvents, the orbital hybridization in this molecule, especially if the NMF was found to be a powerful dissociating solvent as proved by conductance2 and thermochemical s t u d i e ~ . ~ ? ~proton is shared by two NMF molecules with their dipole moments pointing in opposite directions (11). In this Some of its relevant physical and chemical properties are structure the positive charge is distributed over the whole given in Table I1 together with the corresponding propmolecule. Chainwise association by hydrogen bonding at erties of related compounds and water. the nitrogen pro tons produces further stabilization. One In the gas phase NMF has a large dipole moment ( p = may guess whether a single or double potential well would 3.86 D)ll and seems to exist predominantly as a transexist for the proton in the middle of this structure, but conformer.12 Liquid NMF is a polar solvent as indicated by Dimroth’s E value (representing the lowering of the such considerations, although interesting for the gaseous or crystalline state, are probably futile for the pure liquid. excitation energy as measured by the solvent-induced frequency shift of the electronic absorption maximum of The hypothetical structure I1 allows for anomalous mobetaine)13with a very high dielectric constant (D= 185.5).14 bility of the proton in liquid NMF by a mechanism which An analysis of IR spectra in gaseous and condensed phases involves proton hopping between suitably stacked NMF indicated the presence of important hydrogen bonding.16 molecules (e.g., vertical to the drawing plane in 11)or by The solvation ability is enhanced by hydrogen bonding and successive rotations of the proton carrying molecules in is reflected by large values for the heat of ionic solvation.1o a hydrogen-bonded chain. Kinetic studies of some ion-molecule reactions have alSince these transport mechanisms require the breaking ready shown that NMF acts as a protic solvent.16 Several of hydrogen bonds, it is doubtful whether they can effecNMR studies indicate that the basic properties of amides tively compete with the diffusion of the solvated structure are due particularly to oxygen which in these molecules as a whole. For the solvated proton in DMF, an analogous is easier to protonate than n i t r ~ g e n . l ~ - ~ l structure may be postulated. Here the second mechanism The corrected nondimensional values K,* = K,/[SH],2, for anomalous mobility does not exist. A conjecture about where [SH], is the molar concentration of the pure solvent, the structure of the negatively charged, proton deficient, show no simple correlation with any other property. The solvated anion cannot be made in such a straightforward similarly corrected pKa* (K,* = Ka/ [ SH],) of PNP meamanner. From the charge distribution in I, one would sured in different amidic solvents correlates with the basic conclude that the proton on the nitrogen is more acidic strength of solvents relative to water, calculated as B = than the one on the carbon atom and that, therefore, carbanion formation would be less probable. [SH]s(Ka,(s~2+where Ka,(sH2+), is the acidity constant of SH2+in H20.g@3It is clear that the solvation structure In structure I11 the planar structure of the parent NMF of the protonated and deprotonated species that may exist in amides plays an important role in determining basicity, acidity, and self-dissociation. Its influence on the kinetics of proton mobility and proton transfer is of particular interest with respect bo acid-base reactions in these media. I11 IV As our measurements show, the presence of water molemolecules has been retained. Chainwise extension of the cules does not change or perturb the solvation equilibrium solvation structure by hydrogen bonding is possible, as well to a noticeable extent. as hopping migration of the proton vacancy by two difIn NMF free rotation about the C-N bond of the planar ferent mechanisms, eventually leading to anomalous mostructure I is hindered by an activation barrier of -59 bility of the lyate anion. Structure IV is a carbanionic structure, which may represent the anionic species in N,N-disubstituted formamides. Its existence in NMF cannot be excluded, however, and mixed structures in which the carbanion acts as a hydrogen bond acceptor to a nitrogen donor must also be considered. In NMF, with I1 R’ = H,structure IV would be an interaction point of two kJ/mol. An appreciable shift of negative charge has ocassociated chains. On the other hand, it cannot exist in
3094
J. Phys. Chem. 1980, 84, 3094-3099
NMA, which also shows marked autoprotolysis. The autoprotolysis constant of NMF measured by a spectrophotometric method harmonizes well with the corresponding values of other amidic solvents and allows the study of protolytic equilibria with reference to SH2+-SH or SH-S- standard pairs. References and Notes (1) L. R. Dawson in "Chemistry in Nonaqueous Ionizing Solvents", Vol. IV, G. Jander, H. Spandau, and C. C. Addison, Eds., AkadembVerlag, East Berlin, 1963, p 259. (2) C. M. French and K. H. Glover, Trans. FaradaySoc., 51, 1418 (1955). (3) R. P. Held and C. M. Crlss, J. Phys. Chem., 8Q,2611 (1965). (4) E. J. Klng in "The International Encyclopedia of Physical Chemlstry and Chemical Physics", E. A. Guggenheim, J. E. Mayer, and F. C. Tompkins, Eds., Pergamon Press, 1965. (5) R. P. Bell, "The Proton in Chemisby", Comell University Press, Ithaca, NY, 1973, Chapter 4. (6) L. Weeda and G. Somsen, Red. TraV. Chim. Pays-Bas, 85, 159 (1966). (7) A. Finch, P. J. Gardner, and C. J. Steadman, J . Phys. Chem., 71, 2946 (1967).
(8) C. M. Criss, R. P. Held, and E. Luksha, J. Phys. Chem., 72, 2970 11968). (9) 6.L. be Ligny, H. J. M. Denessen, and M. Alfenar, Red. Trav. Chim. Pays-Bas, 90, 1265 (1971). (10) D. S.Gill, S. J. P. Singla, R. Ch. Paul, and S.P. Neruia, J. Chem. Soc., Dalton Trans., 522 (1972). (11) R. M. Meighan and R. H. Cole, 2, Phys. Chem., 68, 503 (1964). (12) M. Kitano and K. Kuchitsu, Bull. Chem. SOC.Jpn., 47, 631 (1974). (13) F. W. Fowler, A. R. Katrfky, and R. I.D. Rutherford, J. Chem. Soc. B, 460 (1971). (14) S.J. Bass, W. I. Mathan, R. M. Meighan, and R. H. Cole, J. Phys. Chem., 68, 509 (1964). (15) T . Miyazawa, T. Shimanouchl, and S.Mizushima, J. Chem. Phys., 24, 408 (1956). (16) A. J. Parker, J. Chem. Soc., 1328 (1961). (17) A. Berger, A. Loewensteln, and S.Meiboom, J . Am. Chem. Soc., 81, 62 (1959). (18) 0. Fraenkel and C. Franconi, J. Am. Chem. Soc., 82,4478 (1969). (19) R. J. Gillespie and T. Birchall, Can. J . Chem., 41, 148 (1963). (20) W. E. Stewart and T. H. Siddall, Chem. Rev., 70, 517 (1970). (21) 6. G. Cox, J. Chem. SOC. B, 1780 (1970). (22) Reference 4, Chapter 11. (23) C. D. Ritchie in "Solute-Solvent Interactions", J. E. Coetzee and C. D. Ritchie, Eds., Marceli Dekker, New York, 1969.
Electron Exchange between Ferrocene and Ferrocenium Ion. Effects of Solvent and of Ring Substitution on the Rateli2 Edward Shih Yang, Man-Sheung Chan, and Arthur C. Wahl" Department of Chemistry, Washington University, St. Louis, Missouri 63 130 (Received: June 2, 1980)
The rates of electron exchange between bis(cyclopentadienyl)iron(II) and -(III) (ferrocene and ferrocenium ion) and between oxidized and reduced forms of several derivatives of ferrocene have been measured in a number of different solvents by the NMR line broadening method over a temperature range of 0-30 O C . It was found that the rates did not vary with the dielectric properties of the solvents as predicted by the Marcus theoretical model for electron exchange between neutral and singly charged spherical reactants with similar structures, reactions for which solvent reorganization is the principal deterent to exchange. Also, the product of the collision number (2)and the transmission coefficient ( K ) was found to be an order of magnitude smaller than the generally assumed value of KZ= lo1' M-l s-l. Addition of NaPFBor NaC104to acetonitrile solutions of ferrocene and ferrocenium ion reduced the exchange rate by about a factor of 2 at high (0.1-0.5 M) salt concentrations. The presence of substituents on the cyclopentadienyl rings affected the rate of electron exchange only moderately; the largest effect of about a 10-fold increase in rate was observed for the decamethyl derivative. The presence of the methylenetrimethylamine group on one ring, resulting in 1+ and 2+ charged reactants, caused a reduction in rate by a factor of -5, much less than the factor of -60 estimated for Coulombic repulsion between uniformly charged spheres 7 A in diameter, an indication that the charged quaternary amine groups are widely separated (> 10 A) in the transition state. Electrical conductivity measurements of cobaltocenium hexafluorophosphate in a variety of solvents indicated that it and very probably the similar ferrocenium salt are essentially completely dissociated at the low concentrations used for measurement of electron-exchange rates in most solvents investigated.
-
Introduction The rates of electron exchange between ferrocene (Fe(CP)~)and ferrocenium ion (Fe(Cp),+) and between their deviatives in various solvents are of interest because the main deterent to electron transfer is probably the necessity for solvent re~rganization.~ There is no Coulombic repulsion between the reactants, so the work required to bring them together is minimal, and their structures are similar, the distance between rings being 3.3 f 0.1 A for both Fe(Cp)z6and Fe(Cp)z13,eso the energy required for internal rearrangement should be small and essentially the same for electron exchange in different solvents. Therefore, investigation of the rate of electron exchange in different solvents should increase our understanding of the requirements for solvent reorganization. 0022-3654/80/2084-3094$01 .OO/O
and examination of Consideration of bond space-filling models indicates that the iron is buried between the cyclopentadienyl rings, which nearly touch and thus shield the iron from solvent molecules. The main solvent interactions, therefore, are probably with the two rings, which together form the surface of a molecule or ion having a cylindrical shape, the length and diameter being approximately equal. Thus, the exchange systems investigated resemble, to a reasonable approximation, the simple model proposed by Marcus' of spheres in a continuous, unsaturated dielectric medium, and comparison of experimental results with theoretical predictions should be informative. A number of solvents with a considerable range of dielectric constants were used, and, to investigate the pos0 1980 American Chemical Soclety
