摘要:

P R E F A C E S.I. Units Papers are beginning to appear in primary publications in which all units used are SI; most papers, however, continue to be written in a mixture of metric systems. Because of the labour that would be involved at a time when Reporters are already hard pressed, The Chemical Society has not yet requested them to use SI exclusively in articles. Some Reporters have used SI, others follow the traditional mixture. There is therefore lack of consistency between one article and another, and sometimes between an article and the scientific paper being reported. The Society’s policy on units is being announced in ‘Notices to Authors’, published on the inside of the covers of issues of The Journal of The Chemical Society.P R E F A C E S.I. Units Papers are beginning to appear in primary publications in which all units used are SI; most papers, however, continue to be written in a mixture of metric systems. Because of the labour that would be involved at a time when Reporters are already hard pressed, The Chemical Society has not yet requested them to use SI exclusively in articles. Some Reporters have used SI, others follow the traditional mixture. There is therefore lack of consistency between one article and another, and sometimes between an article and the scientific paper being reported. The Society’s policy on units is being announced in ‘Notices to Authors’, published on the inside of the covers of issues of The Journal of The Chemical Society.

ISSN:0069-3022    

DOI:10.1039/GR96865FX001

出版商:RSC    

年代:1968

数据来源: RSC

摘要:

10 INTRODUCTION By E. A. V. Ebsworth (Departinent of Chemistry, University of Edinburgh, West Mains Road, Edinburgh 9) IN the past year, as in 1967, there has been no new discovery in Inorganic Chemistry that seems worth special mention. Among new journals, the Journal of Molecular Structure (Elsevier) (whose first number actually appeared in December 1967) is of particular importance; among new books, there is one about physical methods in inorganic chemistry,l while a collec- tion of essays has appeared as a tribute to Professor Emeleus on his 65th birthday.2 During the year there has been a continued expansion in the size of established journals, most notably for Section A of the Journal of the ChemicaZ Society; this section is to be published more frequently in 1969. Such continued expansion means that the need for a reassessment of the form of publication and retrieval of information becomes more compelling each year. In 1968, too, the question of S.I.units has become a live At present the situation is somewhat confused. It is greatly to be hoped that in the future chemistry will not be faced with ‘American’ and ‘European’ con- ventions for units as well as for signs of electrode potential. H. A. 0. Hill and P. Day, ‘Physical Methods in Advanced Inorganic Chemistry,’ See correspondence in ‘Chemistry in Britain’. Interscience, 1968. ” ‘New Pathways in Inorganic chemistry’, Cambridge University Press, 1968. 21310 INTRODUCTION By E. A. V. Ebsworth (Departinent of Chemistry, University of Edinburgh, West Mains Road, Edinburgh 9) IN the past year, as in 1967, there has been no new discovery in Inorganic Chemistry that seems worth special mention.Among new journals, the Journal of Molecular Structure (Elsevier) (whose first number actually appeared in December 1967) is of particular importance; among new books, there is one about physical methods in inorganic chemistry,l while a collec- tion of essays has appeared as a tribute to Professor Emeleus on his 65th birthday.2 During the year there has been a continued expansion in the size of established journals, most notably for Section A of the Journal of the ChemicaZ Society; this section is to be published more frequently in 1969. Such continued expansion means that the need for a reassessment of the form of publication and retrieval of information becomes more compelling each year. In 1968, too, the question of S.I. units has become a live At present the situation is somewhat confused. It is greatly to be hoped that in the future chemistry will not be faced with ‘American’ and ‘European’ con- ventions for units as well as for signs of electrode potential. H. A. 0. Hill and P. Day, ‘Physical Methods in Advanced Inorganic Chemistry,’ See correspondence in ‘Chemistry in Britain’. Interscience, 1968. ” ‘New Pathways in Inorganic chemistry’, Cambridge University Press, 1968. 213

ISSN:0069-3022    

DOI:10.1039/GR96865BX003

出版商:RSC    

年代:1968

数据来源: RSC

摘要:

3 GASES, LIQUIDS, AND LIQUID MIXTURES By N. G. Parsonage (Imperial College of Science and Technology, London S. W.7) Gases.-For the interpretation of second virial coefficient (B) and molecular beam scattering data Barker and Pompel have proposed the following equation for the potential energy (u) as a function of distance (R): 3 2 i = 0 i = 0 u(R) = s{exp [a(l - r)] CAi(r - - Cc2i+6/(8 + Y " + ~ ) > > (1 where Y = R/Rm, Rm = the position of the minimum in the potential energy curve, c6, c8, and c10 are the dispersion force coefficients, which are here derived from theory rather than chosen to fit the experimental data, and 6, a fairly unimportant parameter, prevents the occurrence of a spurious maximum at small R. Since AO and A1 are chosen to make u(Rm) = E and to ensure that the minimum in u(R) occurs at R = Rm, there are effectively five disposable parameters ( E , Rm, a, A2, A3) apart from 6.Of these A2 and A3 are chosen to make ii and dii/dR at a given small separation equal to the values found from molecular beam scattering. The remaining parameters (E, Rm, R) are then chosen so as to fit the B data. The potential is here fitted to the argon data, and leads to a less-deep potential well (Elk = 147.7 or 148.5") than was found by Rowlinson using his B data with the Kihara spherical-core potential ( ~ / k = 164"). Barker and Pompe have, however, rejected the B values at the lowest temperatures, and it is known that the inclusion of these leads to a sharper and deeper potential. This potential has some of the flexibility of the Guggenheim-McGlashan potential without the disadvantage of having discontinuities in some of the derivatives of the potential.Application of the potential to the calculation of the viscosity coefficient (3) yields values which are greater than those from experiment by ca. 10% at 1500"~. The disagree- ment for the thermal conductivity (A) is less (ca. 4%). It is suggested that the discrepancy for yi arises from errors in the application of the slip correction in the experimental determinations. A similar conclusion about the unreliability of the usually accepted high- temperature viscosity data has been reached by Hanley and Childs.2 They considered both the evidence from thermal conductivity data (as in ref. 1) and also new viscosity data by Kestin and co-workers at 300-570"~~ and Guevara at 1 ~ O O - ~ ~ O O K .* J. A. Barker and A. Pompe, Austral. J . Chem., 1968,21, 1683. H. J. M. Hanley and G. E. Childs, Science, 1968, 159, 1114. J. Kestin and W. Leidenfrost, Physicn, 1959,25, 1033; J. Kestin and J. H. Whitelaw, ibid., 1963, 29, 335. F. A. Guevara, unpublished work. 3334 N. G. Parsonage Klein5 has shown that in order to choose between possible potential func- tions it is necessary to have data outside the reduced temperature range 2 < T* < 10, if only equilibrium properties are used, or 2 < T* < 5 if a transport property is involved, where T* = kT/E6-12. Equation-ofstate data. Staveley et aL6 have reported measurements of B for argon, krypton, and methane, and their binary mixtures from near the normal boiling-points to room temperature.A differential method was used, for which it was necessary to assume values of B at 25"c. The values found were well represented by a Lennard-Jones 6 - 18 potential, whereas the 6 - 12 potential gave results for reduced B us. reduced Twhich were numeric- ally too large by ca. 8%. The former potential also led to values of the co- efficient of R-6 which are in good agreement (ca. 6 x) with the best theoretical values. This result supports that of Weir et aZ.,7 which also showed that low- temperature B data demanded a potential with a deeper but narrower well than is provided by the 6 - 12 potential. Analysis of the mixture data showed a marginal preference for the Hudson and McCoubrey combining rule, derived from the London equation: over the Berthelot rule, ~ 1 2 = ( E I I E ~ ~ ) ~ , and one derived by Fender and Halsey from the Kirkwood-MuHer equation, E l 2 = 2 E i i E 2 2 / ( E i i + E22). (3) Data at fairly low temperature have also been reported by Hoover et aL8 for methane, ethane, and their mixtures.They used the Burnett method with pressures up to 40 atm. The data for B for methane (T 2 191.06"~) agrees with that of ref. 6 to ca. 1 cm.3, although it was necessary here to correct for the appreciable amounts of impurity (a few tenths of 1 % of C02 and N2). For methane, they found the Kihara core potential to be better than the Lennard-Jones 6 - 12; for ethane, on the other hand, the Lennard-Jones 6 - 12 was better than a Kihara potential with either a cylindrical or a triangular prism core.They found B12 to be well reproduced by the usual (Lorentz-Berthelot) combining rules. For H2O B and C, over the range 150-450°c, have been measured by Kell et aL9. Values for B differ from the much quoted ones of Keyes et aZ.1° by as much as ca. 60 cm.3 at 150°c, decreasing to ca 8 cm.3 at 250"c and 5 M. Klein and H. J. M. Hanley, Trans. Faraday Soc., 1968, 64,2927. 6 M. A. Byrne, M. R. Jones, and L. A. K. Staveley, Trans. Faraday SOC., 1968, 64, 1747. 7 R. D. Weir, I. W. Jones, J. S. Rowlinson, and G. Saville, Trans. Faraday SOC., 1967, 63, 1320. 8 A. E. Hoover, J, Nagata, T. W. Leland, jun., and R. Kobayashi, J . Chem. Phys., 1968,48,2633. 9 G. S. Kell, G. E. McLaurin, and E. Whalley, J . Chem. Phys., 1968,48,3805; 1968, 49, 2839.10 F. G. Keyes, L. B. Smith, and H. T. Gerry, Proc. Amer. Acad. Arts Sci., 1936,70,3 19.Gases, Liquids, and Liquid Mixtures 35 0 at 450"c. The Stockmayer potential fits the B data fairly well but is poor for C. Similar experiments have been carried out for D2O. These suggest that the enthalpy of dimerisation is ca. 20 cal.mole-l lower for DzO than for HzO, in contrast to the usual rule that deuterium bonds are stronger than hydrogen bonds. Preliminary results for B of 4He from 22 to 2 0 " ~ have been given by Boyd et ~~1.11 from velocity of sound measurements. A severe disadvantage of this method is that a functional form must be assumed for B as a function of T, and slight alteration of this form causes B to change drastically. Ratzsch and Rasenberger12 have measured B of cc14 + C,H, at 90"c by a differential technique and find evidence of specific interactions.Knobler et al.13 have made direct measurements of the excess B, i.e. BIZ - (&I + B22)/2, by observing the change of P produced (range: 0.05-4.53 torr) when the gases are mixed. The 15 binary mixtures of the n-alkanes from methane to hexane wcre studied from 25 to 100"~. Their results fitted a three-parameter Corresponding States relationship with the Lorentz-Berthelot combining rules for Tc and Vc ; the Hudson and McCoubrey (equation 2) and Fender and Halsey (equation 3) rules were less satisfactory. The Principle of Congruence was obeyed fairly well, the deviations being attributable to non-random orientations of the molecules at these relatively low temperatures.B12 Values have been deter- mined for H20 + Nz, H20 + Ar, and HzO + CH4 by a transpiration technique by Rigby and Prausnitz.14 Temperatures from 25 to 100"c and pressures from 20 to 100 atm. were employed. Transportproperties. As well as the criticism of the high-temperature q data mentioned above, the early work of Johnston et a1.15 on the low-temperature determination of q for simple gases has been questioned. Kestin had previously criticised the theory used with the oscillating disc viscometer data. Now Clarke and Smithl6 have given new values obtained by-a capillary flow method from 114 to 3 7 4 " ~ for argon, krypton, and xenon. Below 1 7 0 " ~ the new values are lower than the earlier ones by up to 3%. Di Pippo and Kestin17 have also examined the consistency of the data for the binary mixtures of helium, neon, argon, nitrogen, and carbon dioxide with various sets of com- bining rules, assuming a Lennard-Jones 6 - 12 potential.They found that the Lorentz-Berthelot rules were incompatible with the data, and, indeed, that only the rule ~ 1 2 0 1 2 ~ = (~11a11~~22a22~)3 was satisfactory. The effect of the presence of chemically reacting species on thermal diffusion has been examined by Paul and de Vries.18 The large heat fluxes which can arise by a combination of dissociation and re-combination of Nz04, for l1 M. E. Boyd, S. Y. Larsen and H. Plumb, J . Res. Nat. Bur. Stand., 1968, 72A, 155. l2 M. Ratzsch and E. Rasenberger, Z. Chem., 1968, 8, 156. l3 E. M. Dantzler, C. M. Knobler, and M. L. Windsor, J .Phys. Chem., 1968,72, 676. l4 M. Rigby and J. M. Prausnitz, J . Phys. Chem., 1968, 72, 330. l5 H. L. Johnston and K. E. McCloskey, J. Phys. Chem., 1940,44,1038; H. L. Johnston l6 A. G . Clarke and E. B. Smith, J . Chem. Phys., 1968,48, 3988. l7 R. Di Pippo and J. Kestin, J . Chem. Phys., 1968,49, 2192. and E. R. Grilly, ibid., 1942, 46, 948. R. Paul and A. E. de Vries, J . Chem. Phys., 1968,48,445, 4867.36 N . G. Pursonage example, with diffusion of the two species lead to a large change in the separa- tion factor for an inert pair of substances also present in the system. Thus, in one experiment In q (q is the separation factor) for Ar + He was changed from 0.051 to 0-074 by the presence of the reacting system 2N02 %N204. This result is in accord with their theoretical treatment, which is an elaboration of earlier work by Brokaw assuming instantaneous equilibrium at each point in the system.Barua et all9 also refer to earlier work by Brokaw20 in discussing their measurements of the diffusion of krypton in the presence of 2N02 % N204. Beenakker et uL21 have measured D12 for the 10 binary mixtures formed from helium, neon, argon, krypton, and xenon over the range 65 to 400"~ at P < 1 atm. There is reasonable agreement with the data of other workers except for the flow-type measurements of Walker and Wcstenberg22 on He + Ar at high temperatures. The results are well described by either the Lennard-Jones 6 - 12 or the 6-exp potential. Of the combining rules tried, only ~ 1 2 ~ 1 2 ~ = (&11~11~~22022~)~/~ was successful; the Lorentz-Berthelot rules failed.De Vries et ~ 1 . ~ ~ have observed a strong temperature dependence of the ratio of the D12 values for carbon dioxide of masses 44 and 46 mixed with argon; for the system He + COZ the ratio was much less sensitive to temp- erature. These temperature effects cannot be explained by the Chapman- Enskog theory. The simulation of rotational relaxation effects by assemblies of loaded spheres has been pursued for 4He + HT and 3He + HD, systems involving molecules of equal total mass.24 After making allowances for the difference .between Lennard-Jones 6 - 12 molecules and hard spheres, the thermal diffusion factors ( M ~ ) for the two systems are well reproduced. Experimental testing of the theory of A for systems involving inelastic collisions due to Mason and M ~ n c h i c k ~ ~ has been carried out by Barua et aL26 Systems studied were hydrogen sulphide (- 78-5-200"c),jhydrogen chloride (0-200°c), and ammonia (-26.2-20.1 "c).The equation was used in inverted form to yield values of Zrot, the mean number of collisions for relaxation. Hydrogen sulphide gave a maximum in the graph of Zrot us. T; a similar effect had been found for sulphur dioxide relaxation by Lambert et aZ.27 and attributed to the molecules being favourably oriented for relaxa- tion at low temperatures. A satisfactory discussion of the results for l9 T. K. R. Dastidar, A. Saran, and A. K. Barua, J . Phys. ( A ) , 1968, 1, 269. 2o R. S. Brokaw, J . Chem. Phys., 1960, 32, 1005; 1961, 35, 1569. 21 R. J. J. Van Heijningen, J.P. Harpe, and J. J. M. Beenakker, Physica, 1968,38, 1 . 22 R. E. Walker and A. A. Westenberg, J . Chem. Phys., 1959, 31, 519. 23 W. A. Oost, J. Los, A. N. Van der Steege, A. J. H. Boerboom, and A. E. de Vries, 24 S. I. Sandler and J. S. Dahler, J . Chem. Phys., 1967, 47, 2621; S. I. Sandler and 25 E. A. Mason and L. Monchick, J . Chem. Phys., 1962, 36, 1622. 26 A. K. Barua, A. Manna, and P. Mukhopadhyay, J . Chem. Phys., 1968,49, 2422. 27 J. D. Lambert and R. Salter, Proc. Roy. SOC., 1957, A, 243, 78; P. G. Corran, J. D. Physica, 1967, 36, 637. E. A. Mason, ibid., 1967, 47, 4653. Lambert, R. Salter, and B. Warburton, ibid, 1958, A , 244, 212.Gases, Liquids, and Liquid Mixtures 37 NH3 + MeNHz-at 66-258"c has been made by Gutweiler and Raw2* in terms of the same theory.A series of reports by Barua et aZ.29 have shown the unsatisfactory nature of any of the current treatments of xT for systems containing C02. A study of -q for C2H2, COZ, CS2, and CsHs at 30-200"~ showed no evidence of effects from quadrupole-quadrupole interactions, but this is probably accounted for by the relatively high temperatures involved.30, 31 Pure Liquids.-Critical region. A major field of study continues to be the behaviour near the critical point. The theoretical and experimental progress in this area have been reviewed recently by Fisher32 and Heller,33 respectively In particular, much effort, both experimental and theoretical, has been directed towards the examination of the indices with which various thermodynamic quantities approach asymptotically their critical values.For example, Buck- ingham et al.34 have determined Cv of Ar at the critical density (pc), using a sample height of only 1 cm. in an attempt to avoid phase separation due to the gravitational field, and have found that near Tc, their results can be represented by : - a = lim [In Cv/ln(T - Tc)] where x = 0.08. These measurements are extremely difficult to make, and the authors say that they can be sure only that a lies in the range 0-0.25. Wheeler and Griffiths35 have proved by a thermodynamic argument a result proposed by that a locus of points with C, = co is incompatible with thermodynamic stability. The analagous rule for multicomponent systems is that C, should not diverge along a continuous line of critical or plait points.Determinations of the isotherms P(p, T ) have been made by Wilcox and Balzarini37 for xenon. Their measurements extended to AT/T = 5 x 10-6, where AT = T - Tc, and were plagued by time-constants which were, in some cases, several days. They confirmed the result of Habgood and Schneider that y, the index for (@/Bu)T, falls abruptly from 1.4 to 1 when AT/Tc falls below 10-3. Roach 38 has reported data on the P-p-Tsurface of 4He near the critical point using the Clausius-Mosotti equation to give p, and has discussed these in terms of the critical indices. Although he finds that the Griffith T-+TG+ 2* J. Gutweiler and C. J. G. Raw, J. Chem. Phys., 1968,48, 2413. 29 A. K. Ghosh, A. K. Batabyal, and A. K. Barua, J. Chem. Phys., 1967, 47, 3704; A. K. Batabyal and A.K. Barua, ibid., 1968,48,2557; A. K. Batabyal, A. K. Ghosh, and A. K. Barua, ibid., 1968,445238; S. K. Deb and A. K. Barua, Trans. Faraday SOC., 1968, 64, 358. 30 A. K. Pal and A. K. Barua, J . Chem. Phys., 1968,48, 872. 31 F. R. McCourt and R. F. Snider, J . Chem. Phys., 1967, 47, 417; W. M. Klein, D. K. Hoffman, and J. S. Dahler, ibid., 1968, 49, 2321. 32 M. E. Fisher, Reports Progr. Phys., 1967, 30, 615. 33 P. Heller, Reports Progr. Phys. 1967, 30, 731. s4 C. Edwards, J. A. Lipa, and M. J. Buckingham, Phys. Reu. Letters, 1968, 20, 496. 35 J. C. Wheeler and R. B. Griffiths, Phys. Rev., 1968,170, 249. 36 0. K. Rice, J . Chem. Phys., 1954, 22, 1535. 37 L. R. Wilcox and D. Balzarini, J . Chem. Phys., 1968,48, 753. 38 P. R. Roach, Phys. Rev., 1968, 170, 213.38 N.G. Parsonage inequality p (6 + 1) + a’ >2, where a’$ and 6 are the indices for Cv at T below Tc, for p, and for P as a function of p, respectively, is not obeyed, attention should be drawn to the letter39 refuting on mathematical grounds a similar result by the same author. Choy40 has presented a lengthy treatment of the probability density pp+ qss ({p}, {q), {s)) ofp + 4 + s molecules where molecules within each subset {p), (q), {s) are close together, but there are large separations between the subsets, a type of distribution which is impor- tant near the critical point. He was able to deduce the relationships cc = 2 - dy/(2 - q) = 2 - dv and y = (2 - q)v, where d is the dimensionality, v is the index for the range of correlation, and q defines the way in which the total correlation function decreases with distance at the critical point.These results had been proposed previously by Kadanoff for the king model. De Paz41 has measured Dll and D12 (for impurities) in argon near the critical point. He has not found any dramatic change in D11 up toAT/T, == 10-5, but D12 falls towards zero. General theory. A number of attempts have been made to widen the range of systems which can be satisfactorily covered by machine calculations. Thus, whereas Monte Carlo (MC) calculations for fluids have used, almost ex- clusively, the canonical ensemble (constant T, V and N ) , Wood42 has given preliminary results for a system of 12 particles in an isothermal-isobaric (T, P , N ) ensemble. It would be hoped that by permitting temporary small expansions such an ensemble would not be so vulnerable to ergodic difficulties when dealing with high-density systems.However, at present, the method adopted does not seem to be readily adaptable to systems with ‘soft’ inter- actions. Hoover and Ree43 have put forward two rather similar ways of determining the entropy S for the solid phase. This quantity is needed for a satisfactory discussion of the solid-fluid transition, since, with the size of ‘saniple’ which can be accommodated in a computer (several hundred particles) the system is almost invariably either completely ‘solid‘ or completely ‘fluid’. In one method, the particles are constrained to remain within their cells whilst p is reduced. Since S for this artificial system at low density can be calculated readily, integration of (aS/a V)T yields S for the solid. Using this method they find that the communal entropy enters approximately linearly with 1 / p.The second method, for which no calculations are reported, involves the imposition of an ‘external field’ to maintain the crystalline order whilst p is reduced. When p is sufficiently low, the ‘field’ is reduced to zero reversibly. McDonald and Singer44 have attempted to obtain data over a range of temperatures in a single run by abandoning the usual importance sampling MC niethod in favour of a simple random-walk procedure. This has led, 39 M. H. Coopersmith, Phys. Reu. Letters, 1968, 20, 432. 40 T. R. Choy, J . Chem. Phys., 1967,47,4296. 41 M de Paz, Phys. Rev. Letters, 1968, 20, 183. 42 W. W.Wood, J . Chem. Phys., 1968,48, 415. 43 W. G . Hoover and F. H. Ree, J . Chem. Phys., 1967, 47, 4873; 1968,49, 3609. 44 I. R. McDonald and R. Singer, J . Chem. Phys., 1967, 47, 4766.Gases, Liquids, aiid Liquid Mixtures 39 however, to poor convergence at some densities. Alder et ~ 1 . ~ ~ have examined the solid phase of hard-disc and hard-sphere systems and found that the Lennard-Jones and Devonshire cell theory gives values for S/Nk which are, respectively, too small by 0.06 and too large by 0.24. The difference in P, and hence S, between the close-packed hexagonal and face-centred cubic struc- tures of spheres could not be detected. Using a development of the lattice-gas model of Yang and Lee, Bellemans et a1.46 have been able to obtain two first-order transitions for a 2-dimensional system, corresponding to the solid-fluid and gas-liquid transitions.They use a square lattice and exclude occupancy of the lst, 2nd, and 3rd neighbour sites, with 4th and 5th neighbours being favoured by potentials -g& and -E, respectively, where g is slightly greater than 1. If g = 0, a phase diagram similar to that of 4He is obtained. This occurs because the system of minimum energy is not the most close-packed, with the result that at O'K an increase of P can cause a transition to the close-packed form. A simpler model in which 1st neighbours were excluded but 2nd neighbours were favoured gave only one transition because it provided only one unique ordered str~cture.~' Barker and Henderson48 have developed a method for fluids which is based on consideration of the real potential as a perturbed hard-sphere potential. The Helmholtz free energy is expanded in a Taylor series in the two parameters which describe the deviation of the potential from the hard- sphere potential. This method has been used for the liquid phase, in con- junction with a cell-model treatment of the solid phase, to yield values for the melting line. For argon, the discrepancy between theory and experi- ment for loglo P, where P is the melting pressure, is ca.0.1 from 100 to 400"~, rising to ca. 0.5 at 8 0 " ~ . belle man^^^ has given preliminary notice of a theory for slightly non-spherical molecules with a hard core based on a perturbation on the simple hard-sphere case. Meeron and Siegert50 have made a develop- ment of the scaled-particle approach to obtain a distribution function for cavities of molecular size.Solutions of the Percus-Yevick(PY) equation have been given for the Lennard-Jones 6 - 12 potential at T < Tc by Watts.51 He truncated the potential at various distances and observed a phase change at a temperature dependent upon the position of truncation. All previous calculations for the Lennard-Jones 6 - 12 potential referred to T > Tc. B a ~ t e r ~ ~ has shown that an analytic solution can be found for the PY equation for 'hard spheres with surface adhesion', i.e. with an infinitely deep, infinitesimally narrow potential 45 B. J. Alder, W. G. Hoover, and D. A. Young, J . Chern. Phys., 1968, 49, 3688. 46 J. Orban and A. Bellemans, J . Chem. Phys., 1968, 49, 363; J.Orban, I. van Craen, 47 L. K. Runnels, L. L. Combs, and J. P. Salvant, J . Chem. Phys., 1967,47, 4015. 48 J. A. Barker and D. Henderson, J . Chem. Phys., 1967,47, 4714; D. Henderson and 49 A. Bellemans, Phys. Rev. Letters, 1968, 21, 527. 50 E. Meeron and A. J. F. Siegert, J . Chem. Phys., 1968,48, 31 39. 51 R. 0. Watts, J . Chenr. Phys., 1968,48, 50. 52 R . J. Baxter, J . Chem. Phys., 1968, 49, 2770. and A. Bellemans, ibid., 1968, 49, 1778. J. A. Barker, Mol. Phys., 1968, 14, 587.40 N. G. Parsonage well at the hard-sphere surface. This leads to a first-order phase transition, as observed iiumerically for the Lennard-Jones 6-1 2 potential. B a ~ t e r ~ ~ has also shown how the PY approximation can be obtained by an application of the variation principle to an expression for the pressure.Cure and Babb54 have solved the PY equation for the 6-exp potential of argon at 55"c. They find little improvement on the Lennard-Jones 6 - 12 potential in reproducing the isotherm of this system. A comparison of experimental and theoretical values of thermodynamic properties has been made by lad^^^ for the PY, Hypernetted Chain (HNC), and the pressure-consistent hybrid of these two approximations (PC) (this hybrid was chosen by Rowlinson so as to give the same value for P whether calculated by the Virial Theorem or by the com- pressibility equation of Ornstein and Zernike). PC was always as good as PY and better than HNC. Because the reason for the success of the PY approximation remains mysterious, work has continued on the more firmly based Yvon-Born-Green equation (YBG).Lee, Ree, and Rees6 have applied the Superposition Approxi- mation to the four-particle distribution function : where each g is the distribution function for the particles indicated by the subscripts. This approximation gives results for g(R) of hard-sphere systems which compare favourably with others up to moderate densities. The modifi- The modification to the Superposition Approximation proposed by Young and Rice57 8123 = g12g13g23 exp [T(123)1, g1234 = gl23g124g134g234/(812g13g14g23g24g34), where 03 ~(123) = 2 pn8n+3(123), n = 1 appears to lead to fair agreement for hard discs and spheres, except near the density at which the phase transition occurs, but is much less satisfactory for attractive potentials.The contention that the Superposition Approximation is not bad for g(R), but that by using it in the YBG equation its errors are magnified, is supported by the work of Chung and E~penscheid.~~ They have used the Approximation in the Kirkwood-Salsburg equation and find correct results for the first four virial coefficients and quite good results for the fifth for the hard-sphere potential. The van der Waals-type theory of Flory et UZ.,~~ which leads to the reduced equation of state: SF/T= 5/(6 - 1) - l/(JV), 53 R. J. Baxter, J . Chem. Phys., 1967,47, 4855. 54 J. C. Cure and S. E. Babb, jun., J . Chem. Phys., 1968,48, 2043. 5 5 F. Lado, J . Chent. Phys., 1967, 47, 4828; 1968,49, 3092. 56 Y-T Lee, F. H. Ree, and T. Ree, J . Chem. Phys., 1968, 48, 3506. 57 D.A. Young and S. A. Rice, J . Chem. Phys., 1967, 47, 4228, 5061. 58 H. S. Chang and W. F. Espenscheid, Mol. Phys., 1968,14, 317. 59 P. J. Flory, R. A. Orwoll, and A. Vrij, J . Amer. Chem. SOC., 1964, 86, 3507, 3515.Gases, Liquids, and Liqirid Mixtiires requires knowledge of 41 New values of u and y have been determined, and with these, previous incon- sistencies in the fit to the n-alkane data have been resolved.60 With the realisation that at liquid densities three-body forces cannot be ignored, attempts are being made to salvage the large amount of work that has been performed assuming additivity of the two-body potentials. Sina- noglu61 has chosen an 'effective pair potential' in such a way that it gives the correct mean energy for a system. The treatment is for dilute solutions and assumes a Kihara spherical-core potential for the two-body potential.The three-body potential assumed is an extension of the Axilrod-Teller expression which it is hoped will remain valid at shorter distances. For a given separation of the molecules under consideration, the three-body interaction with the solvent molecules is calculated by integration over the space accessible to solvent molecules assuming that over this space the distribution is uniform. The 'effective potential' for two solute molecules was found to be reduced by 53-2 % in carbon disulphide but only by 19-6 % in neon. The corresponding potentials for solute-solvent and solvent-solvent interactions were reduced by only 1/2 and 1/3 of these amounts, respectively. Experimental.Grigor and Steele62 have discussed their data on CH4 and CD4 in terms of a small difference between the potentials of the two isotopic species (0-4 % in position of well, 0.9 % in well depth). This difference results from the different mean lengths of the CH and CD bonds. On the other hand, the rival theory of isotopic effects on physical properties, that of Bigslei~en,~~ has been supported by new data on the vapour pressures of D2O and T20 from near the m.p. to 1 1 4 " ~ . ~ ~ Teague and Pings65 have determined the refractive index for argon along the coexistence curve and along eight isotherms from 133 to 1 7 3 " ~ for P from 15 to 100 atm. They found that over three states of matter and from 20 to 300"~ the Lorentz-Lorenz function (= (n2 - l)/[p(n2 + l)]) was con- stant to &1-5%.From the coexistence curve they found the critical index a = 0.364 + 0-007, rather a high value. Crawford and DanielP have deter- mined the melting curve of argon up to 6.3 x lo3 bar and 200"~. They found rough agreement with Longuet-Higgins and Widom's theory67 of melting, based on a van der Waals-type model of the liquid phase. 60 R. A. Orwoll and P. J. Flory, J . Amer. Chem. SOC., 1967, 89, 6814. 0. Sinanoglu, Adv. Chem. Phys., 1967, 12, 283, J. Halicioglu and 0. Sinanoglu, J . Chem. Phys., 1968,49, 996. 62 A. F. Grigor and W. A. Steele, J . Chem. Phys., 1968, 48, 1032, 1038. 63 J. Bigeleisen, J . Chem. Phys., 1961, 34, 1485. 64 W. M. Jones, J. Chem. Phys., 1968, 48, 207. 65 R. K . Teague and C. J. Pings, J . Chem. Phys., 1968,48,4973. 66 R.K . Crawford and W. B. Daniels, Phys. Reu. Letters, 1968, 21, 367. 67 H. C. Longuet-Higgins and B. Widom, Mol. Phys., 1964, 8, 549.42 N. G. Parsonage The Clausius-Mosotti equation has also been used to give p and 1 av [= F(a,)l for 4He from 1.25 to 4 . 2 " ~ and 0.25 to 28 atm.68 In the vicinity of the A-point a was represented by the equations: cc = AI loglo IT- TAI + DI o: = A11 log10 I TA - TI 3. DII T > TA T < TA with AI J. AII and DI # DII. The compressibility and thermal expansion of 3Hesg and the Cv and freezing curves of 3He and 3He + 4He mixtures70 have also been measured. New data on argon include p of solid and liquid (T = 96.41 to 153-94"~, P = 80 to 2000 kg./cm.2) by an unusual method,71 and the velocity of sound (T = 86 to 146"~, P up to 65 atm.).72 Molecular dynamics data on free-path distances and collision rates for hard-sphere and square-well molecules73 provide strong evidence against Eyring's jump theory of diffusion.Thus, whereas Eyring's theory suggests that at V = 1.6 x close-packed volume the probability of a jump of one particle diameter is no such jump was found in these machine experi- ments in lo6 attempts. Two-dimensional systems were also examined. A 'van der Waals' theory of DII, developed by Dymond and Alder,74 defines ceff as the value derived by consideration of the equation-of-state data as that of a hard-sphere fluid. This value is then used in the low-density kinetic- theory equation for Dll for hard spheres modified by a factor suggested by Enskog to extend the equation to high densities.They find satisfactory agreement for liquid CH4. Paul and Mazo75 have tested the equation of Kirkwood for 0 1 2 of solutions of polymers: where 5 is the frictional constant per chain eiement, N is the number of chain elements per molecule, and Ra,- is the distance between the ith and jth chain elements. The term i #i has been computed for n-alkanes and, expressing 4 as &r),,a, a has been determined from the equation. The quantity a is found not to be a constant for a given alkane but to vary from solvent to solvent. Carbon tetrachloride has been one of the most frequently studied liquids 68 D. L. Elwell and H. Meyer, Phys. Rev., 1967, 164,245. 69 G. C. Straty and E. D. Adams, Phys. Rev., 1968,169, 232. 70 R. C. Pandorf, E.M. Ifft, and D. 0. Edwards, Phys. Rev., 1967,163, 175. 71 W. van Witzenburg and J. C. Stryland, Canad. J . Phys., 1968,46, 81 1. 72 D. H. Bowman, C. C. Sim, and R. A. Aziz, Canad. J . Chem., 1968,46, 1175. 73 T. Einwohner and B. J. Alder, J . Chem. Phys., 1968, 49, 1458. 74 J. H. Dymond and B. J. Alder, J . Chem. Phys., 1968,48, 343. 7 5 E. Paul and R. M. Mazo, J . Chem. Phys., 1968, 48, 1405.Gases, Liquids, arid Liquid Mixtures 43 because of its convenience and the symmetry of its molecules. It is, therefore, of interest that Raman data has shown substantial distortion from tetra- hedral symmetry. Liquid Mixtures.-Mixtures of hard spheres have been considered by Henderson and Barker77 by a perturbation method. They expand the Helm- holtz free energy of the mixture in a Taylor series in power of the differences oap - 0, where cap = (oaa + 088)/2 and a and p may be either 1 or 2.The reference system is the pure liquid of molecules of diameter 0, 0 being chosen to make the sum of the first-order terms zero. The Superposition Approxima- tion was used in the evaluation of the second-order terms. Good agreement with MC and MD results was found for c22/011 < 1.67. A similar approach has been adopted for both hard discs and spheres by Bellemans et aZ.78 They found good agreement with the PY results for hard spheres due to Lebowitz and Rowlin~on.~~ Treatments of both binary and ternary mixtures have been presentedSo which utilise a transformation due to Fisher enabling the properties of these systems to be expressed in terms of those of a one-component lattice gas.The mixtures are considered in terms of a simple cubic lattice ‘decorated’ by the addition of secondary sites at the mid-points of all bonds. For a binary mixture, primary and secondary sites are restricted to molecules of com- ponents 1 and 2, respectively. Vacancies may occur in either type of site. The interaction energies of nearest neighbours are EII = ~ 2 2 = 0, ~ 1 2 = +cp. The ternary mixture is represented with a 2-molecule on each secondary site and either a 1- or a 3-molecule on each primary site. The effect on the 3-component phase diagram of various relative values of the interaction energies is discussed. A modification to the Corresponding States treatments of Prigogine and Scott has been put forwardS1 which replaces the evaluations of the mean E and c parameters proposed by those authors by one which is similar to that used for a van der Waals gas : and Calculations for soft spheres (E = aR-n) show good agreement with the PY solution, which can be obtained for this potential by a perturbation process on the hard-sphere model.Applying the above scheme for <E> and <03> <03> = c cxixj 0 i j 3 i j 76 H. S. Gabelnick and H. L. Strauss, J . Chem. Phys., 1968, 49, 2334. 77 D. Henderson and J. A. Barker, J . Chem. Phys., 1968,49, 3377. 78 A. Bellemans, J. Orban, and E. de Vos, Chem. Phys. Letters, 1968, 1, 639; A. Belle- 79 J. L. Lebowitz and J. S. Rowlinson, J . Chem. Phys., 1964, 41, 133. 8o R. K. Clark,J. Chem. Phys., 1968,48,741; R. K. Clark and G.A. Neece, ibid., 1968, 48, 2575. 81 J. W. Leach, P. S. Chappelear, and T. W. Leland, Amer. I t u t . Chem. Engineers J . , 1968,14,568; T. W. Leland, J. S. Rowlinson, and G . A. Sather, Trans. Faraday SOC., 1968, 64, 1447. mans and J. Orban, ibid., 1968, 2, 253.44 N. G . Parsonage to a system of Lennard-Jones 6 - 12 molecules, and using the Lorentz- Berthelot combining rules, the results for nine binary mixtures of simple molecules are in better agreement with experiment than are those of the Average Potential Model of Prigogine, the best of the originally proposed variants. Flory et al. have shown that the experimental results for Ar + Kr and also binary mixtures of n-alkanes can be well reproduced by Flory’s earlier theory employing only one adjustable parameter.82 The same theory has also been successfully applied to HE and VE data on aromatic and alicyclic systems by Benson and Singh,83 the solubility parameter, Barker lattice, and Prigogine Corresponding States theories all having previously failed.Flory’s theory also gave a better fit than the Barker theory for C6H6 + isomeric C6H4Me2 systems, and in addition gave a reasonable value for VE, which Barker’s theory cannot predict. The Frank ‘iceberg’ model for solutions in water has been employed for Ar (solubility, 5 to 2 5 O ~ ) , ~ ~ EtOH and the corresponding deuteriated system (vapour pressure^),^^ and n-alkanes (solubility).86 On the other hand, Miller and Hi1debrands7 have rejected this model in favour of an explanation based on Pople’s theory of flexible hydrogen-bonds in their discussion of the large difference in AS of solution for water and cyclohexane which is found for a large number of inert gases.Other data for the solubilities of gase9 in solvents include He, N2, and Ar in N2H4, MeNHNH2, and Me2NNHzs8 and Ar, CH4, CF4, SF6, H2, and Dz in four paraffin^.^^ In the latter case, an explanation is presented of AS as arising in part from the change in partial molar volume of the solute. The solubilities of C2F6, C3F8, c-C~FS, CClF3, and C3H8 in cyclohexane have been measured by Miller;90 in a plot of energy of vaporisa- tion us. -RTlnx2 (x2 is the solubility expressed as a mole fraction) the first three of these lie on the same line as CF4 and SF6, and fall away from the line through the data for non-fluorine-containing gases.Liquid mixtures containing fluorocarbons and hydrocarbons have been further examined for evidence of weakness of the 1-2 interactions. Gaw and Swintongl from their measurements of the vapour pressures of C6F6 with C-C~HU, C6H6, C6H5Me and pC6H4Me2 from 30 to 70°c, and using the Scatchard-Hildebrand theory in terms of volume-fractions, find a 5 %weaken- ing of the 1-2 energy parameter as compared with the geometric niean 82 H. Hocker and P. J. Flory, Trans. Furuduy SOC., 1968, 64, 1188; R . A. Orwoll and 83 G. C. Benson and J. Singh, J . Phys. Chem., 1968, 72, 1345; J. Singh, H. D. Pflug, 84 A. Ben-naim, J . Phys. Chem., 1967,71,4002. 85 C. V. Linderstrom-Lang and F. Vaslow, J . Phys. Chem., 1968, 72, 2645. 86 H. D. Nelson and C. L. de Ligny, Rec.Truv. chim., 1968, 87, 623. 87 K. W. Miller, and J. H. Hildebrand, J . Amer. Chem. SOC., 1968, 90, 3001. 89 J. Walkley and W. I. Jenkins, Trans. Faraduy Soc., 1968, 64, 19. P. J. Flory, J . Amer. Chem. SOC., 1967, 89, 6822. and G. C. Benson, ibid., 1968,72, 1939. E. T. Chang, N. A. Gokcen, and T. M. Poston, J . Phys. Chem., 1968, 72, 638. K. W. Miller, J . Phys. Chem., 1968, 72, 2248. W. J. Gaw and F. L. Swinton, Trans. Faraduy SOC., 1968, 64, 637, 2023.Gases, Liquids, and Liquid Mixtures 45 prediction. Barker's quasi-lattice theory, which is intended particularly for mixtures involving strong specific interactions, has been used for the hydro- gen-bonded system C7F15H + 1,4-dioxang2 and for CsHsMe + C6HsF and c-C6H11Me + Since the Barker theory offers the possibility of a large number of adjustable parameters, one for each type of interaction in the system, it is difficult to assess whether the agreement with experiment comes from the intrinsic correctness of the theory or from the flexibility arising from the presence of, typically, 3-6 adjustable parameters. The first of these could also be satisfactorily discussed in terms of a Dolezalek treat- ment, using an equilibrium constant for hydrogen-bond formation derived from n.m.r., and assuming that other interactions were as for C7F16 + 1,4- dioxan.Bakx and Knaapg4 have followed up previous work on the H2 + Dz, H2 + HD, and HD + Dz systems, by determining AHm for 0-H2 + p-H2 and 0-Dz + p-Dz, for which they find 2-25 & 015 and 2-35 & 0.15 J. molee1, respectively. Previous determinations of these very small quantities are rather scattered. Heats of mixing, AHm, sound velocities, and adiabatic compressi- bilities have been reported by Deshpande et aLg5 for mixtures of CsH5NH2 with other organic liquids.The related toluidines have been studied in their mixtures with C-C~HI~, CsH6, and 1,4-dioxan ( VE meas~rernents),~~ where the results for mixing correspond to break-up of hydrogen-bonded toluidine aggregates followed by 1-2 specific interactions in the cases of CaH6 and 1,4-dioxan. An amusing solution is that of iodine in Me2S0, where the saturated solution at 25"c contains 88 % iodine!97 This very high solubility is attributed to the similarity of the solubility parameters (Me2S0, 13.0; 12, 14.1). A vapour pressure study of Br2 + Clz mixtures in the liquid and solid states (7' = - 100-+60"c, P up to 2761 mm.Hg) has been made by Chees- man and S~ott.~8 A McGlashan-type calorimeter modified to permit loading from the vapour phase has been used to give AHm data on C6Hs + n-C7H16,~~ but, unfortunately, there are no suitable independent data with which to compare the results.Williamson et a1.100 have used a conventional McGlashan-type calorimeter to find AH, for ccb + EtzO, cc14 + Me&, and cc14 + CsHsN at 25 O C . The results support previous evidence from solid-liquid phase- equilibrium studies that in these systems there are strong specific interactions 92 I. D. Watson, R. J. Knight, I. R. McKinnon, and A. G. Williamson, Trans. Faraday 93 S . N. Bhattacharyya and A.Mukherjee, J . Phys. Chem., 1968, 72, 56; S. N. Bhatta- g4 I. N. Bakx and H. F. P. Knaap, Physica, 1968,39, 1. 95 D. D. Deshpande and M. V. Pandya, Trans. Faraday SOC., 1967, 63, 2346; D. D . 96 M. C. Chowdhary and V. R. Krishnan, Austral. J . Chem., 1967, 20, 2761. 97 T. Soda and J. H. Hildebrand, J . Phys. Chem., 1967,71, 4561. 9* G. H. Cheesman and D. L. Scott, Austral. J . Chem., 1968, 21, 287. 99 H. Watts, E. C. W. Clarke, and D. N. Glew, Canad. J . Chem., 1968,46, 815. loo D. F. Gray, I . D. Watson, and A. G. Williamson, Austral. J . Chem., 1968, 21, 379 SOC., 1968, 64, 1763. charyya, R. C. Mitra, and A. Mukherjee, ibid., 1968, 72, 63. Deshpande and L. G. Bhatgadde, J . Phys. Chem., 1968,72,261.46 N. G. Parsonage between donor 0, S , or N atoms and the CCh.Values for VE at 35"c for 1,4-dioxan with CeHd?, C6HsC1, and CsHsBr, and CsHs f C6HsBr and C-CsHlZ + CeHsBr have all been found to be very small, the largest numeric- ally being for the last system (-0.16 ~m.~rnole-l for x = O*5).lo1 Van't Zelfde et a1.1°2 have studied the solid-liquid phase diagram of Ar + CH4 by thermal analysis and vapour pressure measurements. They found GE = 18.1 cal. mole-l for the liquid mixture at 8 6 " ~ , in good accord with previous data. Study of the effect of isomeric change of the components of a binary mixture has been a popular venture. Values of HE and VE for n-ClsHs4 + hexane isomers, and also some other systems, have been measured.lo3 Isomeric change led to a change in VE from -0-45 to -0.70 ml. mole-I. Geiseler et ul.lo4 have determined total vapour pressure, and thence GE, for n-C7H16 with four isomeric octanols and with four isomeric oximes related to n-octane.Isomeric change caused no apparent change of symmetry of GE us. x , and led to change of G at the peak by only ca. 30 cal. mole-l for GE values of ca. 200 cal. moleu1. Prausnitz et uZ.lo5 found a much larger change in the peak GE (from ca. 150 to ca. 70 cal. mole-l) on going from hex-1-ine to hex-3-ine with acetone as the other component. This is probably because of the possibility of hydrogen-bond formation with the terminal hydrogen atom of the 1-isomer. The molecular complexity of H2O in organic solvents has been investigated by an isopiestic method with the assumption that each molecular species present obeys Henry's Law.Affsprung et u2.lo6 concluded that H2O was monomeric in C-C6H12, CCh, C6H5Me and C6H6, but that both polymers and monomers are present in C2H4C12 and CzHzC14. This result for H2O in C6Hs was also found by Worley.lo6 Bittrich et uZ.lo7 have interpreted their results for GE of CCh with CsHsMe, CsHsEt, C6H5C1, and C5H5N in terms of their donor-acceptor properties. Benson et aZ.lo8 have found substantial deviations from the Principle of Congruence in their study of HE and VE of eleven binary mixtures of n-alcohols. Omission of the data for systems in- volving MeOH lessened but did not eliminate the deviations. The Principle was also found wanting in the density studies at - 165"c of short n-paraffins lol P. R. Naidu, J . Phys. SOC. Japan, 1967, 23, 892.lo2 P. van'T Zelfde, M. H. Omar, H. G. M. Le Pair-Schroten, and 2. Dokoupil, Physicu, lo3 J. G. Fernandez-Garcia, M. Guillemin, and C. G. Boissonas, Helu. Chim. Actu, lo* G. Geiseler, K. Quitsch, H-G. Vogel, D. Pilz, and H. Sachse, 2. phys. Chem. H. G. Harris and J. M. Prausnitz, Amer. Inst. Chem. Engineers J., 1968, 14, 737; lo6 J. R. Johnson, S. D. Christian, and H. E. Affsprung, J . Chem. SOC. ( A ) , 1967,1924; lo7 R. Kind, G. Kahnt, D. Schmidt, J. Schumann, and H.J. Bittrich, 2. phys. Chem. lo8 A. E. Pope, H. D. Pflug, B. Dacre, and G. C. Benson, Canad. J . Chem., 1967, 45, 1968, 38, 241. 1968,51, 1451. (Frankfurt), 1967, 56, 288. L. J. Hirth, H. G. Harris, and J. M. Prausnitz, ibid., 1968, 14, 812. J. D. Worley, Canad. J . Chem., 1967, 45, 2465. (Leipzig), 1968, 238, 277. 2665; H. D. Pflug and G. C. Benson, ibid., 1968,46,287.Gases, Liquids, and Liquid Mixtures 47 and their binary and ternary mixtures.lO9 Very careful determinations of VE and GE from 25 to 60"c have been made by Marsh110 for mixtures of CeH6 and cc14 with octamethyltetrasiloxane, molecules of which are very large and almost spherical. The critical mixing region for binary systems is analogous to the gas- liquid critical region for one-component systems, but has considerable experimental advantages. Relaxation of concentration fluctuations in binary systems has been followed by observation by a microscope with close tern- perature control of the phase.lll The system studied was polystyrene in cyclohexane, and led to a relaxation time of 0.1 sec. at T - Tc = 0.01"~. Chu et aL1l2 have discussed in detail the experimental difficulties which have arisen in their scattering study of the isobutyric acid + water system near to the critical solution temperature. Campbell et al.113 have measured p, VE, vapour pressure, surface tension, and -q of the C6H5NH2 + n-C7H16 system, which was chosen because it is known to have a nearly flat coexistence curve in the vicinity of the critical solution temperature. They observed that *q was anomalously high (ca. 10%) just above Tc. This type of anomaly has been discussed by Fixmanll4 in a series of papers in which he also shows that the behaviour should really be non-Newtonian in this region. Cooperll5 has made calculations of the way in which critical solution temperature should be affected by an applied electric field. In agreement with the experimental results of Debye and Kleboth he found that Tc decreases on application of the field for the type of system which they studied. log M. Y . Shana'a and F. B. Canfield, Trans. Faraday SOC., 1968, 64, 2281. 110 K. N. Marsh, Trans. Faraday SOC., 1968, 64, 883. 111 P. Debye and R. T. Jacobsen, J . Chem. Phys., 1968,48,203. 112 B. Chu, F. J. Schoenes, and W. P. Kao, J . Amer. Chem. SOC., 1968,90, 3042. 113 A. N. Campbell, E. M. Kartzmark, S. C. Anand, Y . Cheng, H. P. Dzikowski, and 114 R. Sallovanti and M. Fixman, J . Chem. Phys., 1968, 48, 5326. 115 M. J. Cooper, J . Chem. Phys., 1968,48, 4272. S. M. Skrynyk, Canad. J . Chem., 1968,46,2399.

ISSN:0069-3022    

DOI:10.1039/GR9686500033

出版商:RSC    

年代:1968

数据来源: RSC

摘要:

4 THERMOCHEMISTRY by H. A. Skinner (Department of Chemistry, The University, Manchester, M13 9PL) THE 23rd Annual Calorimetry Conference (U.S.A.), the 4th All-Union Calorimetry Conference (U.S.S.R.), the 3rd Experimental Thermodynamics Conference (U.K.), and the 4th Calorimetry Conference of Japan all took place during 1968. Each of these National Calorimetry Conferences is by now well-established, and their success augers well for the 1st International Conference on Calorimetry and Thermodynamics, planned to take place at Warsaw, Poland in September 1969. 1968 also saw the 1st International CODATA Conference, held at Frankfurt-am-Main, at which the special problems facing the compilers of standard reference tables of thermodynamic, spectral, and nuclear data were thoroughly examined.Of significant new books and publications, mention should be made of ‘Experimental Thermudynamics’, Vol. I, edited by J. P. McCullough and D. M. Scott, and the announcement from Academic Press of the Journal of Chemical Thermodynamics, to appear in January, 1969. The present status of thermal and thermochemical data for lanthanide metal chalcogenides, borides, pnictides, and halides has been assessed by Westruml, Darby2 has reviewed the thermodynamic properties of transition-metal alloys, and Hubbard, O’Hare, and Feder3 have reviewed recent progress in experimental inorganic thermochemistry. The thermo- dynamic properties of the platinum-group metals and their compounds were reviewed by Goldberg and H e ~ l e r . ~ Continuing development in the design and operation of reaction calori- meters is exemplified by the twin-flow and the batch-type microcalorimeters described by Wadso5 and by Monk and Wadso.6 These are conduction calorimeters using thermopiles made from semi-conductors, and both are available from the L.K.B.Instrument Co. An isothermal titration calorimeter in which the heat of titration is neutralised by a constant cooling Peltier device is described by Christensen, Johnston, and I ~ a t t . ~ Peltier cooling is also employed in the microcalorimeter described by Evans, McCourtney, and Carney.8 Franks and Watsong have reported on a twin differential calorimetric 1 E. F. Westrum jun. in ‘Advances in Chemistry Series No. 71,’ American Chemical Society, Washington D.C., 1967, p. 25. 2 J. B. Darby jun., Znd.and Eng. Chem., 1968, 60, 28. 3 W. N. Hubbard, P. A. G. O’Hare, and H. M. Feder, Ann. Rev. Phys. Chem., 1968, 19, 1 1 1 . * R. N. Goldberg and L. G. Hepler, Chem. Rev., 1968, 68, 229. I. Wadso, Acta Chem. Scnnd., 1968, 22, 927. P. Monk and I. Wadso, Acta Chem. Scand., 1968, 22, 1842. 7 J. J. Christensen, H. D. Johnston, and R. M. Izatt, Reo. Sci. Znstr., 1968, 39, 1356. W. J. Evans, E. J. McCourtney, and W. B. Carney, Analyt. Chern., 1968, 40,262. F. Franks and B. Watson, J . Sci. Znstr., [J. Phys. (E)] 1968, 1, 940. 4950 H. A . Skinner system for the measurement of heats of solution of non-electrolytes at high dilutions (mole fraction, 0-05-2.0 x I. Key Compounds.-Careful measurements by Greenberg and Hubbard'" of the energy of combustion in a bomb calorimeter of graphite in fluorine led to AHf"(CF4, g) = - 223.04 & 0.18 kcal./mole, and have largely settled any remaining doubts on this important thermochemical datum. In conjunction with other seemingly reliable data,ll the sharp value for AHfO(CF4, g) 'fixes' the value for the enthalpy of formation of aqueous hydrofluoric acid at AHf0(HF,20H20) = - 76.75 0.10 kcal./mole.A direct measurement of the latter from the heat of combustion of hydrogen in a fluorine flame calor- imeter has given a value in close agreement.12 New measurements of the enthalpy of hydrolysis of crystalline phosphorus pentachloride by Birley and Skinner,l3 in conjunction with the enthalpy of synthesis of PCI,, gave AHfo(H3P04,100Hz0) = - 309.80 & 0 3 5 kcal./mole, in agreement with the value derived from recent rneas~rementsl~ of the heat of hydrolysis of P4010.The heat of oxidation of phosphorous to phosphoric acid has been measured by solution calorimetry,l5 yielding AHf0(H3P03,c) = - 229.1 kcal./mole. Sousa-Alonso, Chadwick, and Irvingl6 have re-measured the heat of oxidation of iron(rr) ions with dilute hydrogen peroxide. Earlier measurements of this quantity179 18 varied over several kcal./mole. The new measurements, using dilute solutions, are perhaps more relevant to the determination of AHf"(Fe2+, aq) [with respect to AHf"(Fe3+ aq)], and correspond to AH = 10.57 & 0.4 kcal./mole for the process: Fe2+(aq) + H+(aq)+ Fe3+(aq) + &Hz(g) This value is confirmed (AH = 10.6 & 0.9 kcal./mole) by independent rneasurement~l~ of the enthalpy of oxidation of iron(1r) ions by aqueous bromine.The enthalpies of formation of the simpler alkyl radicals (e.g. Me., Et., MezHC., Me&.) are of key importance for the thermochemical evaluation of bond dissociation energies of carbon bonds in organic molecules. Although acceptable values for many of these have been available for several years, they have carried uncertainty limits of at least &l kcal./mole. Moreover, neither the kinetic nor the electron-impact techniques of measurement of dissociation energies in polyatomic molecules have seemed capable of higher 10 E. Greenberg and W. N. Hubbard, J . Phys. Chem., 1968,72,222. l1 J. D. Cox, H. A. Gundry, and A. J. Head, Trans. Faraday SOC., 1965,61, 1594; see l2 R. C . King and G. T. Armstrong, J .Res. Nat. Bur. Stand., 1968,72, A , 113. l3 G . I. Birley and H. A. Skinner, Trans. Faraday SOC., 1968, 64, 3232. l4 R. J. Irving and H. McKerrell, Trans. Faruduy SOC., 1967, 63, 2582. l5 A. Finch, P. J. Gardner, K. S. Hussain, and K. K. Sen Gupta, Chem. Comm., 1968, l6 A. Sousa-Alonso, I. Chadwick, and R. J. Irving, J . Chem. SOC. (A), 1968,2364. l7 D. K. Bewley, Trans. Faraday SOC., 1960, 56, 1629. 18 M. F. Koehler and J. P. Coughlin, J . Phys. Chem., 1959,63,605. l9 J. W. Larson, P. Cerutti, H. K. Garber, and L. G. Hepler, J . Phys. Chem., 1968, also H. A. Skinner, Ann. Reports ( A ) , 1967, 64, 4. 872. 72, 2902.TItermochemistry 51 accuracy than this, so that sharp values have remained elusive. A significant advance is now reported by Chupka20 in the evaluation of threshold appear- ance potentials of parent and fragment ions produced by photo-ionization of molecules and radicals.Improved apparatus21, 22 of much higher sensitivity and photon energy resolution has increased the accuracy of measurement by a factor of 5 or more, and enabled the effect of thermal rotational energy at room temperature in the threshold region to be taken into account. From studies of the photo-ionization of methane, Chupka20 reports Doo(Me-H) = 4.477 & 0.005 ev 103.24 -+ 0.1 kcal./mole; the noteworthy feature here is the reduction (by about a factor of 10) of the uncertainty previously attached to this quantity. The new value corresponds to D(Me-H) = 104.8 kcal./mole at 298"~, and toAHt-O(Me, g) == 34.8 kcal./mole. Chupka and Lifshitz23 have studied the photo-ionization of methyl radicals by the same technique, and its threshold yields Do0(CHz-H) = 107.9 & 1.0 kcal./mole: this corresponds to D(CH2-H) = 109.0 & 1.0 kcal./mole at 298"~ and to AHf"(CH2, g) = 91.7 & 1 kcal./mole.The main source of errar in the threshold from the Me radical arises from uncertainty in the amount of thermal energy in the radicals at the instant of photoionization, since the radicals may have lost thermal energy on leaving the pyrolysis tube. Never- theless, the estimated uncertainty of & 1 kcal./mole represents a five-fold improvement relative to existing uncertainty in AHfO(CH2, g) obtained by other methods. Zmbov, Uy, and Margrave24 have studied the thermal decomposition of tetrafluorethylene over the temperature range 1 127-1244"~, and measured the temperature dependence of the equilibrium constant for the process CzF4(g) + 2CFz(g).The gas C2F4 (produced by pyrolysis of Teflon) was introduced into a heated graphite Knudsen cell, and the effusing gases were ionized by electron impact and analysed by mass spectroscopy. From these measurements, the enthalpy of disssociation was determined, AHr(298 OK) = 76.3 & 3 kcal./mole; in conjunction with AHfO(C2F4, g) = - 157.4 & 1 k~al./mole,~~ this leads to AHf"(CF2, g) = - 403 & 2 kcal./mole, in good agreement with the values obtained by Modica and La GraffZ6 from shock- tube studies on C2F4 and CHF3. Electron impact studies27 on CCl2 radicals (produced by pyrolysis of C13CSiC13) gave 9.76 ev for the vertical ionization potential; in conjunction with measured appearance potentials for the dissociative ionization of CHC13 and CzC14, a value AHf"(CC12, g) = 56-5 & 5 kcal./mole was derived.The new values (i.e. AH,' for CH2, CFZ, and CC12) lead to the following thermo- 2o W. A. Chupka, J . Chem. Phys., 1968,48,2337. 31 J. Berkowitz and W. A. Chupka, J. Chem. Phys., 1966,45, 1287. 22 W. A. Chupka and J. Berkowitz, J . Chem. Phys., 1967, 47, 2921. 23 W. A. Chupka and C. Lifshitz, J . Chem. Phys., 1968, 48, 1109. 24 K. F. Zmbov, 0. M. Uy, and J. L. Margreave, J . Amer. Chem. SOC., 1968,90, 5090. 25 J. R. Lacher and H. A. Skinner, J . Chem. SOC. (A), 1968, 1034. 26 A. P. Modica and J. E. La Graff, J . Chem. Phys., 1965,43, 3383; 1966,44, 3375. 27 J.S. Shapiro and F. P. Lossing, J . Phys. Chem., 1968, 72, 1552.52 H. A . Skinner chemical bond dissociation energies of C= C double bonds (in kcal./mole): D(CH2=CH2)= 170.9 & 2, D(CH2=CF2)= 133.7 & 3, D(CH2=CCl2) = 147.9 5, D(CF2=CF2) = 76.3 & 3, D(CC12=CC12) = 116.4 7, D(CF2=CCl2) = 96.2 & 6. Kinetic studies of thermal bromination of CF3H (361-431 "c) and C2F,H ( 3 2 3 4 5 8 " ~ ) by Amphlett and WhittleZ8 have established the Arr- henius parameters for the broinination reactions, and provide the bond dissociation energy values, D(CF3-H) = 106.7 0.5 and D(C,F,-H) = 102.6 & 1.2 kcal./mole. The value for D(CF3-H) is in good agreement with earlier measurements from independent kinetic studies291 30, and there seems no reason to doubt that it is of the right order of magnitude.In con- junction with the seemingly reliable value for AHf0(CF3H, g) = - 166.2 & 1 k~al./mole,~~, 31 the value AHf"(CF3, g) = - 1 1 1.6 & 1-5 kcal./mole is derived. This in turn with AHfo(C2F~, g) = - 321 -& 1 kcal./mole25 gives D(CF3-CF3) = 97.8 & 3 kcal./niole, in good agreement with the direct determination of this quantity by Tschuikow-Roux32 using a shock tube. However, this background of internal consistency is severely jolted should the conclusions of Modica and Sillers33 prove valid; from measurements of the rate of formation of CF2 radicals from the decomposition of CF,I, C,F,, and CF4 in an excess of argon behind shock waves over a temperature range 170&3000"~, these authors arrived at the value AHfO(CF3, g) = - 104 kcal./mole.The analysis of the experimental results is involved, and requires the enthalpies of formation of the reacting species as input data. The values used (from the JANAF tables) can be questioned. 2. Combustion Calorimetry.--Satisfactory values are now available for the standard enthalpies of formation of the oxides of most of the metallic ele- ments. Many of these were determined from high-precision measurements of their heats of combustion in oxygen by bomb combustion techniques. The thermochemistry group at Los Alamos have contributed notably to this programme over the past decade, and new measurements from this source include the heats of combustion of neptunium34 (to form NpOa), hafnium35 (to form HfOz), cerous oxide36 (Ce203-+ CeOz), and thorium carbide37 (ThC1.91).Solution calorimetric methods were used for obtaining the standard heats of formation of selected non-stoicheiometric oxides of 28 J. C. Amphlett and E. Whittle, Trans. Faraday SOC., 1968, 64, 2130. 29 J. W. Coomber and E. Whittle, Trans. Faraduy SOC., 1966, 62, 2183. 30 C. A. Goy, A. Lord, and H. 0. Pritchard, J . Phys. Chem., 1967, 71, 1086. 31 H. A. Skinner, Ann. Reports (A), 1967, 64, 4. 32 E. Tschuikow-Roux, J . Chem. Phys., 1965,43, 2251. 33 A. P. Modica and S. J. Sillers, J . Chem. Phys., 1968,48, 3283. S4 E. J. Huber jun., and C. E. Holley jun., J . Chem. and Eng. Data, 1968,13, 545. 35 E. J. Huber jun., and C. E. Holley jun., J . Chem. and Eng. Data, 1968,13, 252. 36 F. B. Baker and C. E. Holley jun., J . Chem. and Eng.Data, 1968,13,405. 37 E. J. Huber jun., C. E. Holley jun., and N. H. Krikorian, J . Chem. and Eng. Data, 38 G . C. Fitzgibbon and C. E. Holley jun., J . Chem. and Eng. Data, 1968, 13, 63. 1968, 13, 253.Thermoclzemistry 53 (TbO1.51, Tb01.709, Tb01.817, and Tb01.975), and of hexagonal Nd203.39 The heat of combustion of copper in a bomb calorimeter to form a mixture of CuO and Cu20 was reported by Mah, Pankratz, Weller, and King40, and used in combination with dissociation pressure measurements on CuO to yield more definitive values than hitherto for AHf"(Cu0) and AHf"(Cu20). Gal'chenko, Gedakyan, and Timofeev4I have used a nickel bomb calori- meter fitted with an internal micro-heater to measure the heats of combustion of Zr, Hf, and Ta metals in chlorine to form ZrC14, HfC14, and TaC15. The heat of combustion of B4C in fluorine has been reported,42 and the heat of formation of chlorine trifluoride measured by King and Armstrone3 using fluorine flame calorimetry [AHfo(CIF3, g) = - 38.6 kcal./mole].The heat of combustion of chlorobenzene has been measured by Kolesov, Tomareva, Skuratov, and A l e k h i ~ ~ ~ ~ using a platinum-lined rotating bomb calorimeter, with As203 solution as reductant to remove chlorine from the combustion gases. The value AHf"(CsH5C1, 1) = -2.82 & 0-25 kcal./niole was derived. Measurements by Robb and Zimmer45 of the solubility and heat of solution of COS in aqueous solutions containing As20.5, As203, and HCl provide improved information necessary for the standard state correc- tions in rotating bomb calorimetric studies with arsenic(rI1) oxide as reductant.Kolesov, Shtekher, and M a r t y n ~ v ~ ~ have continued their measurements on the heats of combustion of gaseous fluorinated hydrocarbons and find AHc" = -291.5 & 2 kcal./mole for 1,l-difluoroethane, which corresponds to AHfO(CF2HCH3, g) = -118.4 & 2 kcal./mole; this agrees well with the value25 derived from measurements of the heat of hydr~genation~~ of CH2 = CF2. Bomb combustion studies on trib~tylboron~~ and trihexylb~ron~~ have yielded AHfo[B(Bu")3, 11 = -84.3 & 0.5 and AHf"[B(n-hexyl)s, I] = - 116.8 0-7 kcal./mole. Combustion methods have also been used to obtain the heats of formation of a number of cyclopentanyl titanium compounds.50 39 G . C. Fitzgibbon, D. Pavone, and C.E. Holley jun., J . Chem. and Eng. Data, 1968, 4O A. D. Mah, L. B. Pankratz, W. W. Weller, and E. G. King, U.S. Bur. Mines Report 41 G. L. Gal'chenko, D. A. Gedakyan, and B. I. Timofeev, Russ. J . Inorg. Chem., 42 E. S. Domalski and G. T. Armstrong, J. Res. Nut. Bur. Stand., 1968,12, A , 133. -33 R. C. King and G. T. Armstrong, Nat. Bur. Stand. Report No. 9905, 1968. 44 V. P. Kolesov, E. M. Tomareva, S. M. Skuratov, and S. P. Alekhin, Zhur. fiz. 45 R. A. Robb and M. F. Zimmer, J. Chem. and Eng. Data, 1968,13,200. 46 V. P. Kolesov, S. M. Shtekher, and A. M. Martynov,Zhur.fiz. Khim., 1968,42,1847. 47 J. R. Lacher and P. B. Howard, to be published. 4* G. L. Gal'chenko and N. S. Zaugol'nikova, Zhur. $2. Khim., 1967,41, 1018. 49 G. L. Gal'chenko and N. S. Zaugol'nikova, Zhur.$z.Khim., 1967, 41, 2181. 50 V. I. Tel'noi, 1. B. Rabinovitch, V. D. Tikhonov, V. N. Latyaeva, L. I. Vishinskaya, 13, 547. No. 7026, 1967. 1968, 13, 159. Khim., 1967, 41, 1528. and G . A. Razuvaev, Doklady Akad. Nauk S.S.S.R., 1967, 174, 1374.54 11. A . Skinner The heats of combustion of resorcin01,~~ trans-stilbene and 22,44,66-hexa- nitr~stilbene,~~ several azido-s-triaz~les,~~ selected nitroso-naphthols,54 hydroxylammonium per~hlorate~~ and other explosive compounds56 are reported (the use of diethyl oxalate and diethyl phthalate as desensitisers to control the combustion is of interest in dealing with the latter). 3. Reaction Calorimetry.-Gross, Hayman, and Joel5' have measured the enthalpies of the reactions MF(c) + BF3(g)+ MBF4(c), (M = Li or Na), by passing BF3 gas over the solid alkali fluorides in a calorimeter maintained at 110"c.The enthalpy of formation of KBF4(c) was determined by comparing solution heats in HF(aq) of the mixtures (NaBF4 + KF) and (KBF4 + NaF). Greene, Gross, and Hayman5* reported the enthalpy of formation of LisAIFs(c) from the reaction 3LiF(c) + AlF3(c), measured at room temperature in a calorimeter fitted with an electrically heated reaction zone. A 'hot zone' calorimeter was also used by Vorob'ev, Monaenkova, and Skuratov59 to measure AHf"(BaH2, c) directly from the reaction of the metal with hydrogen gas at 2 atm. pressure. Solution calorimetric methods were used by CubicciottPO. 61 to measure AHf"(BiBr3, c) = - 66.0 and AHr"(Bi13, c) = - 36.0 kcal./mole. From equili- brium studies on Q Bi(1iq) + Q BiXs(g)+ BiX(g), (X = Br or I), and measure- ments of AHsub (BiXs), the dissociation energies D(Bi-Br) = 64 and D(Bi-I) = 52 kcal./mole were obtained; the average bond dissociation energies in BiBrs(g) (5 = 56) and BiIs(g)(E = 43 kcal./mole) are less than in the respective diatomic molecules.Froin measurements of the enthalpies of solution of the metals and their chlorides in hydrochloric acid solutions, Stuve62 has obtained Akl~f"(HoCl3, c) and AHf"(TbCI3, c); Gvelesiani and YashvW3 measured solution heats of the metals and the oxides in hydro- chloric acid to obtain AHf"(La203, c) andAHrO(Smz03, c). Solution calorimetry was also used by Irving and McKerrelP4 to obtain the enthalpies of formation of the crystalline salts Na,P,O,,, NazHzPz07, Na4Pz07, (NaP03)3, and (NaP0314.j1 P. D. Desai, R. C. Wilhoit, and B. J. Zwo!inski, J . Chem. and Eng. Data, 1968, 13, 334. 52 S. Marantz and G. T. Armstrong, J . Chem. and Eng. Data, 1968, 13, 1 1 8. 53 G. C. Denault, P. C. Marx, and H. H. Takimoto, J . Chem. and Eng. Data, 1968,13, 514. S4 J. V. Hamilton and T. F. Fagley, J . Chem. and Eng. Data, 1968, 13, 523. 55 M. F. Zimrner, E. E. Baroody, G . A. Carpenter, and R. A. Robb, J . Chem. and Eng. 56 E. E. Baroody, G. A. Carpenter, R. A. Robb, and M. F. Zimrner, J . Chem. atid Eng. 5' P. Gross, C. Hayrnan, and H. A. Joel, Trans. Faraday SOC., 1968,64, 317. 58 P. D. Greene, P. Gross, and C. Hayman, Trans. Faraday SOC., 1968,64, 633. 59 A. F. Vorob'ev, A. S. Monaenkova, and S. M. Skuratov, Doklady Akad.Nauk 6o D. Cubicciotti, Inorg. Chem., 1968, 7 , 208. D. Cubicciotti, Znorg. Chem., 1968, 7 , 211. 62 J. M. Stuve, U.S. Bur. Mines Report No. 7046, 1967. 63 G. G. Gvelesiani and T. S. Yashvili, Zhur. Neorg. Khinz., 1967, 12, 3233. 64 R. J. Irving and H. McKerrell, Trans. Faraday SOC., 1968, 64, 875, 879. Data, 1968, 13, 212. Data, 1968, 13, 215. S.S.S.R., 1968, 179, 1129.Thermocheniistry 55 Hill and Irving65 have continued their studies on the acetylacetonates of transition metals and have found values for AHf" of Fe(acac)s and Mn(acac)a, using reaction calorimetry (reactions with HCl and with FeCl2-HCl respec- tively). For the average M-0 bond dissociation energies in M(acac)3 molecules, the authors quote Mn-0 = 44, Fe-0 = 47, Cr-0 = 56, and Al-0 = 64 kcal./mole respectively.Clark and Price,66 from measurements of the enthalpy of reaction of indium trimethyl with Br2-CHC13 solution, obtained AHf"(InMe3, c) = 29.5 kcal./mole. FromAHsub = 1 1*6,AHf"(In, g) = 58.2 and AHp"(Me, g) = 34-8 kcal./mole, the average bond dissociation energy, D(In-Me) = 40.5 kcal./mole is derived. From previous kinetic studies, the authors found D(Me2In-Me) + D(In-Me) = 87.9 kcal./mole, which, in conjunction with 5 = 40.5, yields D(Me1n-Me) = 33.6 kcal./mole. Enthalpies of hydrolysis in aqueous HCl of nine trimethylsilane derivatives have been reported by Baldwin, Lappert, Pedley, and Tre~erton,~~ and Hill and Wadso68 have measured the enthalpy of hydrolysis of triacetylammonia. An improved value for AHp"(SOBr2,l) is available from enthalpy of hydrolysis measurement^,^^ and for AH," (sulphamic acid, c) from measurementsT0 of the enthalpy of reaction with sodium nitrite in acidic solution.The heat of explosion of CF2(0F)2 has been measured,71 and AHf"(ClF5,l) was obtained from rnea~urements~~ of its enthalpies of reaction with H2, and with NH3. Izatt, Eatough, Snow, and Chri~tensen~~ have measured the enthalpies of formation of complexes between Ag+ and Cu2+ with pyridine (AH for Ag pyzf = - 11.2, and for Cu py42+ = -21.5 kcal./mole), and the enthalpy of forma- tion of Co(CN)s3- has also been measured.74 Carson, Laye, and Smith75 used a diphenyl ether fusion calorimeter in measurements of the enthalpies of formation of the complexes of the divalent ions of Mg, Ca, Sr, and Ba with diethylene triaminepenta-acetic acid, and of the complexes formed by the trivalent lanthanide group of metals. The enthalpies of formation of pyridine complexes with aluminium halides are reported by Wilson and W0rrall.7~ A sensitive differential calorimeter was used by Brunetti, Lim, and Nancollas77 to measure the heats of complexing of Cu2+ with diglycine and triglycine.65 J. 0. Hill and R. J. Irving, J . Chem. SOC. ( A ) , 1968, 1052, 31 16. W. D. Clark and S. J. W. Price, Canad. J . Chem., 1968,46, 1633. 67 J. C. Baldwin, M. F. Lappert, J. B. Pedley, and J. A. Treverton, J . Chem. Soc. (A), 8* J. 0. Hill and I. Wadso, Acta Chem. Scand., 1968, 22, 1590. :g A. Finch, P. J. Gardner, and K. Radcliffe, J . Chem. and Eng. Data, 1968, 13, 176. 60 G. A. Nash, H. A. Skinner, T.A. Zordan, and L. G. Hepler, J. Chern. and Eng. 71 G. D. Foss and D. A. Pitt, J. Phys. Chem., 1968, 72, 3512. 72 W. R. Bisbee, J. V. Hamilton, J. M. Gerhauser, and R. Rushworth, J. Chem. and 73 R. M. Izatt, D. Eatough, R. L. Snow, and J. J. Christensen, J. Phys. Chem., 1968, 74 R . M. Izatt, G. D. Watt, C. H. Bartholomew, and J. J. Chriitensen, Inorg. Chem., 75 A. S. Carson, P. G. Laye, and P. N. Smith, J . Chem. SOC. (A), 1968, 141,527, 1384. 76 J. W. Wilson and I. J. Worrall, J . Chem. SOC. ( A ) , 1968, 316, 2389. 77 A. P. Brunetti, M. C. Lim, and G. H. Nancollas, J . Amer. Chem. Soc., 1968, 90, 1967, 1980. Data, 1968, 13, 271. Eng. Data, 1968, 13, 382. 72, 1208. 1968,7, 2236. 5 120.56 H. A. Skiitner The available enthalpies of formation of fluoro- and of fluoro-halogenated- hydrocarbons have been re-evaluated by Lacher and Skinner25 using improved values for ANm"(CF4, g) and AHf"(HF, as).The revised data correlate well with the bond additivity and bond interaction scheme of Allen78 in the case of the halogen-substituted methanes and fluorocarbons, but in halogen-sub- stituted ethanes steric repulsion energies of the order of 1 kcal./mole between gauche C-X bonds across the central C-C bond are indicated (X = F or Cl). 4. Equilibrium Studies.-The combination of the mass spectroscope and the Knudsen cell continues to prove fruitful in studies of gaseous equilibria at high temperatures. The vapour species over liquid SnF2, over liquid SnF2 + Sn, over CaFz(s) + Sn(l), and over liquid and solid PbFz have been studied mass spectroscopically by Zmbov, Hastie, and Margrave.79 The presence of dimers and trimers as well as monomers of SnFz was noted.There are no dimers or polymers in PbFz vapour, but the disproportionation 2PbFz(g) -+ PbF4(g) + Pb(g) takes place. The equilibrium constant of the process Ca(g) + SnF(g)-+ CaF(g) + Sn(g) was measured over the range 1304- 1401"~ from analysis of the vapours over CaF2 + Sn, and the derived en- thalpy of reaction, with Do(CaF) = 127-5 -+ 2-5 kcal./mole,*O gives D"(SnF) = 114 5 3 kcal./mole. Values were obtained for Do(Pb-F) = 85 & 2 kcal./mole, and for the enthalpy of dissociation, SnzFe(g)-+ 2SnFz(g), AHo = 39 5 2 kcal./mole. The vapours from NaSnFa and KSnF3 over the range 725-895 OK have been similarly investigated,*l and the species SnF2, NaSnF3, NazSnFe, and NaSnzFs have been identified.Ficalora, Hastie, and Margraves2 have analysed the species present when Al(g) is equilibrated with Sz(g), Sea(g), and Tez(g). A similar study, replacing A1 by In, is reported by Colin and D r o ~ a r t . ~ ~ AHf" Values for the gaseous species MX, M2X, and M2X2 were derived (M = A1 or In; X = S, Se, or Te). The vapour species from Gad%, Ga2Se3, and GazTe3 have been investi- gated.** Dimer molecules were detectedg5 in the vapours over KOH ; AH(dinwriza- tion) = -43 kcal./mole dimer. The vapours over Cs-02 contained Cs20 and Cs202, and the enthalpies of formation of both were determined.s6 The vapours from BeF2 at 9 4 7 " ~ contained some dimer molecules ; AH(dimerisa- tion) = -34.5 kcal./mole dimer.87 Small amounts (ca.1 %) of dimer species 78 T. L. Allen, J . Cfzem. Phys., 1959, 31, 1039. 79 K. Zmbov, J. W. Hastie, and J. L. Margrave, Trans. Faraday SOC., 1968, 64, 861. J. W. Hastie and J. L. Margrave, J . Chem. arid Eng. Data, 1968, 13, 428. J. W. Hastie, K. F. Zmbov, and J. L. Margrave, J . Inorg. Nuclear Chem., 1968,30, 729. 82 P. J. Ficalora, J. W. Hastie, and J. L. Margrave, J . Phys. Chem., 1968, 72, 1660. 83 R. Colin and J. Drowart, Trans. Faraday SOC., 1968, 64, 2611. s4 0. M. Uy, D. W. Muenow, P. J. Ficalora, and J. L. Margrave, Trans. Faraday SOC., 85 A. V. Gusarov and L. N. Gorokhov, Z h u r . 5 ~ . Khim., 1968, 42, 860. 86 A. V. Gusarov, L. N. Gorokhov, and A. G . Efimova, Teplofiz. Vys. Temp., 1967, 87 V. I. Belousov, L. N. Siderov, S. A. Kamarov, and P.A. Akishin, Zhur. fiz. Khim., 1968, 64,2998. 5, 584. 1967,41, 2691.Thermochemistr); 57 were detected88 in the vapours over LaCls(c) and EuCk(c), but none were found over EuCh(1). Mass spectral and microbalance studies of the sublima- tion of CrO3 at 150-250"c have shown the presence of polymers in the vapoursg, and provided values for the enthalpies of the reactions, n[CrO3,(g)] + (CrO&(g), for n = 3, 4, and 5. Hildenbrandgo has studied the high temperature gaseous equilibria in the Mg-Cu-F system by mass spectroscopy, and measured the enthalpy of the reaction MgF(g) + Cu(g)+ Mg(g) + CuF(g), leading to D(CuF) = 103 & 2.2 kcal./mole, so that the second dissociation energy in CuFz(g) is sub- stantially higher than the first [D(FCu-F) = 79 kcal./mole], and in accord with expectation from the valence-state excitation model for s-p-bonding in CuF2.Similar studiesg1 involving the reactions M(g) + MFz(g)+ 2MF(g), (M = Mg, Sr, or Ba), have established that D(FMg-F) > D(MgF), but the reverse order applies in the case of BaFz where D(BaF) > D(FBa-F). There is qualitative accord with the valence-state s-p-bonding model for the lighter Group I1 metal fluorides, but for the heavier metals the electrostatic model (in which MF2 is treated as polarisable ions in contact) is much sup- erior. The high temperature reactions AuMn(g) + Au(g)+ Mn(g) + Auz(g), and AgMn(g) + Ag(g)+ Mn(g) + Agz(g) have been studied, giving AHo" = -9-6 0.692 and -1043 & 4 kcal./moleg3 respectively, and leading to Do"(AuMn) = 43.4 & 3 and Do"(AgMn) = 26-8 & 5 kcal./mole.Molecular beam sampling and mass spectral detection were used by Kohl and Carlsong4 to analyse the equilibria in vapours over liquid solutions of Bi-Sb mixtures, in the range 225-625 "c. Dissociation energies and enthalpies of formation of the gaseous species BiSb3, BizSbz, BisSb, and BiSb were obtained and are additive in relation to the molecules Biz, Sbz, Bi4, and Sb4 [e.g. D(BiSb) = iD(Biz) + +D(Sbz); AHfO(Bi3Sb) = gAHf"(Bi4) + $AHfo(Sb4)]. The temperature and concentration dependence of an absorption band in iodine vapour95 in the region 265 mp has been interpreted in terms of an equilibrium 212 $14; in this event, the equilibrium constant and its temp- erature dependence over the range 150-420"~ require AHo (605"~) = 2-9 & 0.4 kcal./mole, and indicate 1.4 mole "/, T4 in the vapour at 240"c and 2.5 atm.pressure. 5. Bond Dissociation Energies.-Budinikas, Edwards, and WahlbeckS6 have developed a novel double-oven effusion cell, suspended as a single unit, enabling the simultaneous measurement of Knudsen effusion- and torsion 88 J. W. Hastie, P. Ficalora, and J. L. Margrave, J . Less Common Metals, 1968, 14,83. 89 J. D. McDonald and J. L. Margrave, J . Inorg. Nuclear Chem., 1968, 30, 665. 90 D. L. Hildenbrand, J. Chem. Phys., 1968,48, 2457. 91 D. L. Hildenbrand, J . Chem. Phys., 1968, 48, 3657. g2 S. Smoes and J. Drowart, Chem. Comm., 1968, 534. 93 A. Kant, J. Chem. Phys., 1968,48, 523. 94 F. J. Kohl and K. D. Carlson, J . Amer. Chem. SOC., 1968, 90, 4814. 95 A. A.Passchier and N. W. Gregory, J . Phys. Chern., 1968,72, 2697. 96 P. Budinikas, R. K. Edwards, and P. G. Wahlbeck, J . Chem. Phys., 1968, 48, 2859, 2867,2870.58 H. A. Skinner effusion-data. The combined measurements on the dissociation processes SZ-, 2S, Se2+ 2Se, and Te2- 2Te have provided the values 101.7 f. 2-9,75-7 If: 2.5, and 61.3 & 1.1 kcal./mole respectively for DO0(S2), Doo(Sez), and Doo (Te2). The latter is confirmed by photo-ionisation threshold meas~rements,9~ DO0(Te2) = 60.85 & 0.2 kcal./mole. Photo-ionisation yield curves for molecule and fragment ions have been measured by Dibeler and List0n~8~~9 from H2S, S02, BF3, and B2F4, and have led to the dissociation energy values Doo(HS-H) = 89.3 and Doo(S-H) = 83.2 kcal./mole; Do0(FzB-F) = 169.3, Doo(FB-F) = 117.6 and Doo(B-F) = 173.6 kcal./mole; and Doo(F2B-BF2) = 103 kcal./mole.The second bond dissociation energy in BF3, i.e. D(FB-F), is thus decidedly less than the first or third; it is noteworthy that the B-B dissociation energy in B2F4 is larger than the C-C dissociation energy in C2F4. The dissociation energies in TlF, TlCl, and TlBr were measured from photo-ionisation thresholds for the formation of TI+ by Berkowitz and Walter;100 the existence of small barriers in the potential energy curves of 3 ~ 0 and 3 ~ 1 states of TlF is confirmed by the new results. Setser and StedmanlOl have investigated a series of bond-rupture reactions in which metastable excited argon atoms act as rupture agent, Ar* + RCN -> Ar + R + CW. The Ar* carries 270 kcal./mole available energy for transfer.The emission spectra of the products were analysed to evaluate Emax of the highest populated vibrational level of CN", and thus set upper limits to the bond dissociation energies D(R-CN). The values so obtained, uiz: D(H-CN) < 124, D(NC-CN) < 133, D(Me-CN) < 116, D(HzN-CN) < 124, and D(F3C-CN) < 119 kcal./mole, establish that AHf"(CN, g) < 103 kcal./mole, in 'agreement with shock-tube studies1OZ which gave 100 & 4 kcal./mole. A method of assessment of differences between different C-C bond dissociation energies in chemically activated alkyl radicals is described by Larson, Tardy, and Rabinovitch.103 The scheme involves the steps (1) and (2) : (1) H + olefin-+ R* (vibrationally activated) (o1efin)l + Ri 7 L ole fin)^ + R2 Measurement of the yields of the olefins provides the relative rates of decom- position of R* by the alternative routes.Studies have been made in this way g7 J. Berkowitz and W. A. Chupka; referred to in ref. 96. 98 V. H. Dibeler and S . K. Liston, J . Chem. Phys., 1968, 49, 482. V. H. Dibeler and S. K. Liston, Inorg. Chem., 1968, 7 , 1742. loo J. Berkowitz and T. A. Walter, J. Chem. Phys., 1968, 49, 1184. lol D. W. Setser and D. H. Stedman, J . Chem. Phys., 1968,49, 467. lo2 W. Tsang, S. H. Bauer, and M. Cowperthwaite, J. Chern. Phys., 1962, 36, 1768. lo3 C. W. Larson, D. C. Tardy, and B. S. Rabinovitch, J . Chern. Phys., 1968, 49, 299.Thermochemistry 59 on the room-temperature decomposition of the octyl-4, heptyl-3, and 3- methyl-hexyl-2 radicals, and have indicated that Doo(Me-H) > Doo(Et-H) by 5.irLkcal./mole, and DoO(Et-H) > DoO(Pr-H) by 0.8 kcal/mole.A similar method was used by Cadman, Phillips, and Trotman-Dickenson,lO4 in studies of the decomposition of chemically activated CF3CH2Me. The active species, prepared by method (3) or (4); From Me. + *CH2CF3+ MeCH2CF3* From MeCHz * + * CF3+ MeCH2CF3** thereafter may deactivate by collision, or eliminate HF to form CF2=CHMe. The propene :propane yield ratio gives the relative rates of decomposition. The results indicate higher activation in CF3CH2Me** than in CF3CH2Me*, to the extent that D(CF3-CH2Me) > D(CF3CHz-Me) by ca. 4.5 kcal./mole. Tachikawa and Rowlandlo5 have studied recoil tritium reactions with hydrocarbons in a series of moderators (C-Cd?s, Ar, Nz), and have measured the HT:RT yield ratios in the reactions ( 5 ) and ( 6 ) : (3 1 (4) T* + RH+HT + R ( 5 ) T*+RH+RT+H (6) Wide variations in the yield ratio were noted, and a correlation emerged between the bond dissociation energies of the R-H bonds and the HT yields.The correlation proved unsatisfactory with MeCD :CD2, the HT yield being less than expected; the suggestion was made that the relaxation of CHzCD:CD2 to the allylic configuration is too slow to gain the benefit of n-electron delocalisation during the rupture process.lo6 HT Yields from recoil tritium abstraction with H-N < bonds also appear to correlate with N-H bond dissociation energies.lo7 The photolysis of alkyl halides in the presence of HSiC13 initiates a chain reaction in which the rate-determining step is RX + -SiC13+ R - + SiC13X.The relative rates of abstraction RIX + SiC13+, and R2X + SiC13-+, have been measured for mixtures of several alkyl clilorides and bromides by Kerr, Smith, Trotman-Dickenson, and Young,loS and the Arrhenius parameters determined. The activation energy differences, E(MeC1) - E(RCI), and E(MeC1) - E(RBr), increase as R ascends the series Et, Bun, Bui, Pri, Bus, But, and parallel the decreasing bond dissociation energies D(R-CI), D(R-Br), for similar changes in R. The pyrolysis kinetics of MesSi-SiMes have been investigated in a static system over the range 523-555" at pressures 0-2-0.8 mm., by Davidson and lo* P. Cadman, D. C. Phillips, and A. F. Trotman-Dickenson, Chem. Comm., 1968 lo5 E.Tachikawa and F. S. Rowland, J. Amer. Chem. SOC., 1968,90, 4767. lo6 E. Tachikawa, Y-N Tang, and F. S. Rowland, J. Amer. Chem. SOC., 1968,90, 3584. lo7 T. Tominaga and F. S. Rowland, J. Phys. Chem., 1968,72, 1399. lo8 J. A. Kerr, B. J. A. Smith, A. F. Trotman-Dickenson, and J. C. Young, J. Chem. 796. SOC. (A), 1968, 510. C60 H. A. Skinner Stephenson.log First-order kinetics were obeyed, and the activation energy (67.3 & 2.2 kcal./mole) is identified by the authors with the Si-Si bond dissociation energy. Band, Davidson, and Lambert ,11* from electron-impact studies and measured appearance potentials of fragment ions, have evaluated bond dissociation energies D(Me3Si-X) in a series of trimethylsilyl deriva- tives, including D(Me3Si-CI) = 88 3: 2 and D(Me3SiBr) = 78.5 & 2 kcal./mole.Combined with AHf"(Me3SiC1, g) = - 84.6 kcal./m0Ie,~7 the value D(Me3Si-Cl) = 88 yields AHfo(Me3Si-, g) = -25.5 kcal./mole: combined with A.Hf"(MesSiBr, g) = -70.1 k~al./mole,~~ the latter yields D(Me3Si-Br) = 71.3 kcal./mole. There is thus a greater inconsistency be- tween the thermochemical and electron-impact results than is indicated by the attached limits of error. The vertical ionisation potentials of Me3Sn, MezSn, and MeSn radicals, formed by electron impact on hexamethylditin, have been measuredlll by a mass-spectrometric method. In conjunction with appearance-potential measurements112 of Me3Snf from several trimethyltin derivatives, bond dissociation energies, D(MeaSn-R), are obtained for R = Me, Et, Pr, and Me&. The Sn-C bond in SnMee was determined at D(Me3Sn-Me) = 60 kcal./mole ; the average Sn-Me bond strength, from thermochemical measurements, is somewhat lower (E = 54.0 kcal./mole).The kinetics of thermal decomposition of CF30F in the presence of SO3 have been followed by pressure over the range 200-230"~. The kinetic analysis yields D(CF30-F) = 43.5 5 0.5 kcal./mole. 6. Miscellaneous. Kebarle114 has described gas-phase measurements of the relative concentrations of clustered ionic species of the type A+Sn (A+ = Hf, NH4+, S = HzO, NH3, MeOH), using mass spectral detection of A+&, sampled from an ion source containing a known pressure of the 'solvent' species, S. The enthalpies of the gas-phase reactions NH4+(NH3)n + NH3 -+ NH4+(NH3)n+l have been derived,1l5 and are -27 (n = 0), - 17 (n = l), - 16.5 (n = 2), - 14.5 (n = 3), and -7.5 kcal./mole (n = 4).The sharp fall at n = 4 is noteworthy, and contrasts with the steady decrease in- A H for the successive addition of H2O molecules to HC. Measurements of the attach- ment of HzO molecules to F-, C1-, Br-, and I- by Kebarle, Arshadi, and ScarboroughllG indicate that the negative ions solvate more readily and more strongly than do alkali-metal positive ions. The standard absolute free energy, enthalpy, entropy, and heat capacity for hydration of the proton have been evaluated by de Ligny, Alfenaar, and log I. M. T. Davidson and I. L. Stephenson, J . Chern. SOC. (A), 1968, 282. 111 F. W. Lampe and A. Niehaus, J . Chem. Phys., 1968,49, 2949. 112 A. L. Yergey and F. W. Lampe, J . Amer.Chem. SOC., 1965, 87, 4204. 114 P. Kebarle, in 'Advances in Chemistry Series No. 72,' American Chemical Society, 115 S. K. Searles and P. Kebarle, J . Phys. Chem., 1968, 72, 743. 110 P. Kebarle, M. Arshadi, and J. Scarborough, J . Chetn. Phys., 1968, 49, 817. S . J. Band, I. M. T. Davidson, and C. A. Lambert, J . Chem. Sor. ( A ) , 1968, 2068. J. Czarnowski, E. Castellano, and H. J. Schumacher, Chem. Comm., 1968, 1255. Washington D.C., 1968, p. 24.Thennocheniistry 61 Vander Veen,l17 based on a cyclic process involving ferrocene and the ferro- cinium ion : H+(aq) + Ferro(aq) - Ferro+(aq)-+ iHz(g) t .1 H+(g) + Ferro(g) - Ferro+(g) --f H+(g) + e-(g) The assumption was made that AGO (hydration) of the ferrocinium ion differs from AHo (hydration) of neutral ferrocene only by the electrostatic contribu- tion (calculable from the Born model), and was justified on the grounds that the ion is large and that the ion and neutral molecule are virtually equal in size.The absolute values, S"(H+)(aq) = -4.8 & 1-2 cal.deg.-1 mole-1 and AH" = -263 5 3 kcal./mole were derived. The interaction of alkyl groups with water is an important factor in relation to micelle formation, hydrophobic bonding in aqueous protein solutions, and adsorption processes at oil-water interfaces. The large negative entropy of hydration of alkyl-chain segments in water is acknowledged, but consistent results have not been obtained from measurements of the enthalpy of hydra- tion per -CH2- group of an alkyl chain. Aveyard and Mitchell118 have considerably clarified the situation from an analysis of their calorimetric measurements on the heats of solution of a series of alcohols and n-carboxylic acids to high dilution in water at 25"c.It is concluded thatAHsoln per -CHz-- (liquid) group is ca. 350 cal./mole for alkyl chains up to c8 in length. The thermal effects of intramicellar dissolution (solubilisation) of hydro- carbons in aqueous sodium oleate solutions have been estimated from refrac- tometric measurements of the intramicellar solubility over a range of concen- trations and temperature.llg For the process octane (I)+ octane-sodium oleate micelle, the value AH = 2 - 4 kcal./mole was derived, on the basis, however, of a model that may be oversimplified. Measurements of the standard partial molar heat capacity of NaBPh4 in aqueous solution from 0 to 90"c have been made by Subramanian and, Ahluwalia.120 The curves ofACpO and zp2O us. temperature showzd two breaks at the minimum near 50", and the maximum near 70"c. The minimum is ascribed to hydrophobic interactions, and the maximum to a sharp reduction in the structure-making ability of the solvent above 70"c. The denaturation of P-lactoglobulin by urea at pH 2-5-3.5 at several temperatures between 5 and 55"c has been studied by Pace and Tanford.l2I The isothermal denaturation, brought about by exposure to urea, was fol- lowed by the change in optical rotation measured at 365 mp. The monomeric native protein was used in all these studies. The equilibrium constant, defined 117 C. L. de Ligny, M. Alfenaar, and N. G. Vander Veen, Rec. Trav. chim., 1968,87, 585. 11* R. Aveyard and R. W. Mitchell, Trans. Faraday SOC., 1968, 64, 1757. lL9 Z . N. Markina, E. V. Rybakova, A. V. Chinnikova, and P. A. Rebinder, Doklady 120 S. Subramanian and J. C. Ahluwalia, J . Phys. Chem., 1968, 72, 2525. Phys. Chem., 1968, 179, 246. N. C. Pace and C. Tanford, Biochemistry, 1968, 198.62 H. A . Skinner by (fraction of unfolded molecules)/(fraction of native molecules), was found to be strongly dependent on the urea concentration, and to be strongly temperature dependent. From the temperature dependence, AH(denaturation) changes from -40 kcal./mole at 15"c, to zero at 35"c, and to 40 kcal./mole at 52"c. This sharp change in A H indicates a large change in heat capacity associated with the unfolding. The behaviour of P-lactoglobulin resembles that of ribonuclease, chymotrypsinogen, and myoglobin and may be charac- teristic for giobular proteins in general.

ISSN:0069-3022    

DOI:10.1039/GR9686500049

出版商:RSC    

年代:1968

数据来源: RSC

摘要:

5 Part (i) HIGH RESOLUTION NUCLEAR MAGNETIC RESONANCE SPECTROSCOPY By J. Feeney ( Varian Research Laboratory, Klausstrasse 43, Ziirich 8, Switzerland) NUCLEAR magnetic resonance spectroscopy continues to have widespread use in analytical and structural investigations. Further support for such routine n.m.r. applications has been provided recently by extensive compila- tions of data by Brugell and Bovey.2 In this review no attempt will be made to cover the many excellent examples of the use of 1i.ni.r. for molecular structural determinations. Likewise, little consideration will be given to the well established procedures for conformational analysis and kinetic studies of inter- and intra-molecular processes. However, special attention will be paid to any advances which have been made in instrumental procedures and in theoretical aspects of the technique. It is worth noting that the develop- ment of the theory of chemical shifts and spin-spin coupling constants is not being retarded by lack of excellent experimental data of which there is an abundance available.Experimental Techniques.-The most exciting recent advance has been the use of pulse techniques to remove dipolar and quadrupolar broadening in solid samples to an extent where chemical shifts and scalar couplings can be ~bserved.~ -7 Normally to observe high resolution n.m.r. spectra it is necessary to examine samples as gases or liquids, or in solution, to provide the molecules with sufficient tumbling and rotational motion to average out the dipolar interactions. Waugh and his co-workers have shown that by pulsing repeti- tively with brief intense pulse trains, phase shifted by k90" and/or & 180", one can observe effects similar but not equivalent to the motional narrowing observed in liquids.A typical 8-pulse cycle used is (T, Px; 27, P-x; 2, Py; 22, P-y; 2, P-x; 22, Px; 2, P-y; 27, Py). The effect of a fast 90" train is analogous to that of molecular motion, and the transient n.m.r. signal observed is the Fourier transform of the high resolution spectrum revealed by the r.f. analogue of motional narrowing. When the system is subjected to a repetitive cycle of pulses, the spin operators 1 W. Briigel, 'NMR Spectra and Chemical Structure,' Academic Press, New York, 1967. F. A. Bovey, 'NMR Data Tables for Organic Compounds,' Wiley, New York, 1967.P. Mansfield and D. Ware, Phys. Letters, 1968, 27A, 159. 3 J. S. Waugh, L. M. Huber, and U . Haeberlen, Phys. Rev. Letters, 1968,20, 180. 5 J. S. Waugh, L. M. Huber, and E. D. Ostroff, Phys. Letters, 1968,26A, 211. 6 D. Gill, Phys. Letters, 1968, 26A, 544. J. S . Waugh, C . H. Wang, L. M. Huber, and R. L. Vold, J . Clieni. Phys., 1968, 48, 662. 6364 J . Feerzey (rather than the lattice operator in motional averaging) bccome time depen- dent. A series of 90" r.f. pulses of different phase has been applied to a sinall single crystal of calcium fluoride in the presence of some liquid trifluoro- methylbenzene, and the phase detected Bloch decay has been observed. The signal from the solid gave a fast decay followed by the much longer decay of the liquid signal. A compound modulation was obtained because of the difference in chemical shift between the two samples (the dipolar interactions in the solid having been removed).The Fourier transform of the compound modulation gives the slow passage spectrum with all dipolar interactions removed and all chemical shifts reduced by a factor of 4 3 . Linewidths of less than 200 Hz have been observed for solid calcium fluoride by this method and no doubt these linewidths will be further improved. This method promises to be much more convenient than the alternative method of removing dipolar broadening in solids by rapidly rotating the samples around an axis inclined at 54" 44' to the direction of the external magnetic field. It has been shown recently that when a molecule has a spin coupling tensor which has an antisymmetric component this component is not usually removed by the rotation.8 The quest for sensitivity enhancement in n.ni.r. has continued, stimulated largely by the increasing interest in nuclei other than protons where the natural abundance and sensitivity to n.m.r.detection often pose serious sensitivity problems. For example for 13C nuclei the natural abundance is only 1.1 % and the natural sensitivity to detection at constant field strength is only 1.6% that of protons. All known methods of sensitivity enhancement (for example use of large sample tubes, high fields, spectrum accumulation, Fourier transform) benefit by having a high sensitivity in the basic instru- ment: over the last seven years most commercial n.m.r.spectrometers have been able to improve their basic sensitivity specifications by a factor of five as a result of instrumental improvements. During the last year, heteronuclear noise decoupling has become an added powerful means of improving sen- sitivity. Its application to 13C spectroscopy is perhaps the most drama ti^.^ By irradiating at the proton frequency with a noise-modulated r.f. source it is possible to achieve a broad band decoupling effect which decouples simul- taneously all the chemically shifted protons (even those with large C-H coupling constants of 200 Hz). The signal-to-noise ratio of the decoupled 13C spectrum is improved not only by virtue of the fact that the intensity of the collapsed multiplet is now concentrated into a single absorption band but also because of a positive nuclear-nuclear Overhauser effect, which can give signal-to-noise enhancement of a factor of three or more: the two effects combine to improve the sensitivity by a factor of seven or eight, which corresponds to a considerable time saving when spectrum accumulation techniques must be used.By use of this method it is possible to detect all * E. R. Andrew and L. F. Farnell, Mol. Phys., 1968, 15, 157. 9 L. F. Johnson, 9th Experimental Conference, Pittsburgh, Pennsylvania, March 1968.Iiigh Resolution Nuclear Magnetic Resonarice Spectroscopy 65 the carbon atoms in cholesterol after one hundred scans.l*, l1 Broad-band decoupling can be achieved without use of noise modulation techniques, by use of a very high power coherent r.f.source.12 However, one can only use this approach with a time-sharing spectrometer where the measuring r.f., the irradiating r.f., and the receiver are switched on in turn for short non-over- lapping intervals at the rate of a few kHz. Hewittl3 has used a time-sharing spectrometer of this type to obtain beat-free decoupled spectra of bicyclo- butane in a nematic solvent. The need for small dedicated digital computers in nuclear magnetic reso- nance has long been realised, but it is only recently that such facilities have become available. Some obvious uses are data acquisition in digitised form, data reduction to a few significant parameters, data handling, spectrum accumulation spectrum smoothing, resolution enhancement, line-shape and spectral analysis and simulation of multispin systems.Computers will also be used increasingly to control n.m.r. spectrometer operation. An obvious example is the automatic control of magnetic field homogeneity by monitoring some resolution criterion such as peak height and automatically modifying the electric shim coil homogeneity controls to optimise the resolu- tion.14 Of the possible criteria for homogeneity, only the second moment and related criteria can allow one to construct a complete set of non-interacting shim adjustments : however, for a resonance line with sufficient natural linewidth the peak height can be related to the second moment and may also give non-interaxting control. The advent of field/frequency locked spectrometers has made practical a simple method of measuring the r.f.observing field, Hl.l5 One selects a sample with a single narrow absorption line and at a low non-saturating value of r.f. power the spectrometer is adjusted to be exactly on resonance: the r.f. power is then suddenly increased by 40 dB such that a transient wiggle signal is observed (o ca. 2 Hz). This frequency is a measure of yH1 to about 0.5% accuracy. By use of a steady-state method, an electric field effect has been observed on the 14N spectrum of highly purified nitrobenzene when a very high field (60 kv/cm.) is applied.16 As expected, a line-splitting was detected but this could not be described completely by use of either the Lorentz or the Onsager local field model. By use of a low field ( 1 0 ~ ) n.m.r. spectrometer with very large samples (1-25 1.) signals from whole living animals have been detected for the first time.17 lo J.D. Roberts, NATO NMR Summer School, Coimbra, 1968. l1 L. F. Johnson, personal communication. l2 E. Lippmaa and T. Pehk, Kemian Teollisuus, 1967, 24, 1001. l 3 R. C. Hewitt, Rev. Sci. Instr., 1968, 39, 1066. l4 R. R. Ernst, Rev. Sci. Instr., 1968, 39, 998. l5 J. S. Leigh jun., Rev. Sci. Instr., 1968, 39, 1594. C. W. Hilbers and C. MacLean, Chem. Phys. Letters, 1968, 2 , 445. J. A. Jackson and W. H. Langham, Rev. Sci. Instr., 1968, 39, 510.66 J. Feeney Novel Techniques.-Several workers18-22 have shown that it is possible to observe n.m.r. emission and enhanced absorption during rapid free-radical reactions. This is referred to as chemically induced dynamic nuclear polarisa- tion (CIDNP).When free radicals are generated by the breaking of chemical bonds, the unpaired electron states have equal populations initially and tend to reach a Boltzman distribution by spin-lattice relaxation. If fluctuating electron spin-nuclear spin interactions are present, cross relaxation (combined electron spin/nuclear spin) transitions may be induced which disturb the nuclear Boltzman equilibrium. This can result in enhanced nuclear polarisa- tion. When radicals react further the polarisation is transferred to the dia- magnetic reaction products, which give rise to increased n.m.r. signals. Whether or not one obtains emission or enhanced absorption depends on whether there is a dipolar or a scalar coupling involved. CIDNP proton spectra have been reported recently for methyl, ethyl, propyl, isopropyl and t-butyl radicals by Kaptein.22 This author also clarified the conditions con- cerning relaxation times and radical life-times necessary for successful obser- vation of this effect.Cocivera23 has described a method for studying optically excited molecules by n.m.r. Using a 3500 w mercury lamp operating over 3000-4000 8, he irradiated 0.005 M anthraquinone in benzene. Under these conditions one obtains a steady-state concentration of anthraquinone in the triplet state. Because of the low concentration of the triplet state molecules, the n.m.r. spectrum of only the ground state molecules is observed, but it is modified by the presence of the triplet states. Thus we see Overhauser effects which change the spectrum of the ground state dramatically, leading to inversion of most of the lines in the ground state spectrum by a process of stimulated emission induced by the optical excitation of the molecule to its lowest lying triplet state.Multiple-resonance experiments form the basis for several novel applica- tions. Sinevee and Sa1m-1~~ have developed a theory to predict the effects of triple homonuclear resonance by use of weak r.f. fields: their theory predicts not only the line shapes but a new phenomenon, namely, the appearance under certain conditions of a signal with a frequency which is a combination of the basic frequencies. This has been experimentally substantiated by By~trov,~~ who has examined the triple resonance spectra of AMX systems.Freeman and Gestblom26 have described a method of measuring frequency l8 J. Bargon, H. Fischer, and U. Johnsen, Z . Nnrurforsch., 1967, 22a, 1551. l9 J. Bargon and H. Fischer, 2. Nuturforsch., 1967, 22a, 1556. 20 H. R. Ward and R. G. Lawler, J . Amer. Chem. Soc., 1967, 89, 5518. 21 R. G. Lawler, J . Amer. Chem. SOC., 1967, 89, 5519. 22 R. Kaptein, Chem. Phys. Letters, 1968, 2, 261. 33 M. Cocivera, J . Arner. Chem. Soc., 1968, 90, 3261. 24 V. Sinevee and V. Salum, BUN. Acnd. Sci. Estonian S.S.R., Phys. Marh. Sci. Ser., 25 V. F. Bystrov, J . MoI. Spectroscopy, 1968, 28, 81. 26 R. Freeman and B. Gestblom, J . Chem. Phys., 1968, 48, 5008. 1968, 17, N1.High Resolution Nuclear Magnetic Resonance Spectroscopy 67 differences to & 1 mHz. In high resolution n.m.r.the dominant line-broaden- ing mechanism is very often the spatial inhomogeneity of the magnetic field; in such cases it is possible to carry out a double resonance experiment to impose a very weak r.f. field, H z , on a given line, Wa, in the spectrum, such that the saturation is localised to a restricted region of the sample volume (referred to as ‘burning a hole’ in the line). If one then examines a line, Wb, regressively connected to Wa, one observes a similar hole in the exactly cor- responding position on the line profile resulting from changes in the popula- tion of the common energy level. This correlation arises from the intra- molecular nature of the coupling and it forms the basis of a technique for measuring frequency separations with a precision not determined by in- homogeneous line widths.Irradiation with a weak perturbing field in a spin-tickling experiment can be used to locate precisely hidden 13C.CH satellite lines in a proton spectrum and hence to determine the magnitude and signs of CCH coupling constants. The technique is applicable mainly to two or three spin systems but by com- bination with spin decoupling in a triple resonance experiment it can be extended to more complex systems, such as found in ally1 bromide.27 By irradiating at the CHzBr frequency the molecule is effectively reduced to a three-spin system, which is sufficiently simple to carry out the required spin- tickling experiments on the 13C.CH satellite spectra. Govil and BernsteinZ8 have found the same relative signs for HgHF and JtHF in 1,1,2,2-tetrabromo- fluoroethane in the gauche and trans rotameric forms.This was possible because the nuclei in the slowly interconverting forms retain the same nuclear orientations with respect to each other during the exchange process: thus irradiation at part of the multiplet of one nucleus will cause transfer of saturation to only part of the multiplet of the interacting nucleus in the other form involved in the slow intramolecular exchange process. Hoffman and ForsenZ9 were the first to use double-resonance methods to study kinetic effects by observing the transfer of saturation from an irradiated nucleus to a nucleus at another nuclear site involved in exchange with the irradiated nucleus. Fung30 has extended the method to cases where the spin saturation can be propagated along an aliphatic chain to involve protons not participating in the direct exchange process but strongly coupled to the nuclei involved in the exchange.This has been illustrated in experiments with mixtures of 2-phenoxyethanol and t-butyl alcohol. Double quantum transi- tions (DQT) in four-spin systems such as (I) have been used to obtain the relative signs of the coupling constants.31 Strong DQT were observed for low 27 J. Feeney and P. J. S . Pauwels, Mol. Phys., 1968, 14, 209. 28 G. Govil and H. J. Bernstein, Mol. Phys., 1968, 14, 197. 29 R. A. Hoffmann and S . Forsen, J. Chem. Phys., 1963,39, 2892. 3O B. M. Fung, J . Amer. Chem. SOC., 1968,90,219. 31 L. Lunazzi and F. Taddei, J . Mol. Spectroscopy, 1968, 25, 113.68 J. Feeney values of Av/J: for AX3 type systems when hv/J was greater than 30 it was not possible to detect DQT.Two independent methods for determining micelle formation with greater ease have been reported. Micelle formation in solutions of sodium 4,4- dimethyl-4-silapefitane-1 -sulphonate in the presence of paramagnetic ions is accompanied by enhanced chemical shift and linewidth changes compared with experiments in the absence of paramagnetic ions.32 Haque33 has exa- mined fluorinated colloidal electrolytes, sodium perfluorocaprylate and propionate, and observed large changes in 19F chemical shifts on micelle format ion. Primary, secondary, and tertiary alcohols have been characterised by preparing the dichloroacetate esters and observing the characteristic dichloro- acetyl proton resonance.34 The meltifig point of ethane has been determined by observing the motional narrowing of its proton n.m.r.signal at the phase transition; during this investigation it was noted that n.m.r. can be used as an impurity detector by measuring the impurity premelting which takes place below the melting point of the pure substance.35 Herington and Lawren~on~~ have illustrated how to determine the purity of organic samples by n.m.r. The samples were examined from 20" below their melting point up to a temperature at which all the sample is melted. The intensity of the liquid line is measured at dif- ferent temperatures and a linear graph of temperature against reciprocal of fraction melted is obtained. From the slope of the graph the purity of the sample can be estimated, and the intercept on the temperature axis gives the melting point of the pure material.Signal integration techniques, applied to separate proton signals from water molecules within and outside the primary co-ordination sphere of Nil1 com- plexes in aqueous solution at - 30", have enabled the solvation number to be determined as six. This is the first time this technique has been used with paramagnetic ions.37, 38 32 B. R. Donaldson and J. C. P. Schwarz, J. Chem. SOC. (B), 1968, 957. 33 R. Haque, J. Phys. Chem., 1968, 72, 3056. 34 J. S. Babiec, J. R. Barrante, and G. D. Vickers, Analyt. Chem., 1968, 40, 610. 35 L. J. Burnett and B. H. Muller, Nature, 1968, 219, 59. 36 E. F. G. Herington and I. J. Lawrenson, Nature, 1968, 219, 928. 37 T.J. Swift and G. P. Weinberger, J. Amer. Chern. Soc., 1968, 90, 2023. 38 N. A. Matwiyoff and P. E. Darley, J. Phys. Cheni., 1968,72, 2659.High Resolution Nuclear Magnetic Resoriance Spectroscopy 69 The tetrameric nature of alkyl-lithium compounds in solution has been confirmed by enriching methyl-lithium with 13C and observing the 7Li-13C coupling in THF solution (14.5 Hz).~* Lippmaa and his co-workers40 have measured intermolecular Overhauser effects between much larger molecules than had previously been considered possible. Coupling Constants.-The Pople and Santry4I CNDO-2SCF (complete neglect of direct overlap) approximation for calculating coupling constants has been extended by Ditchfield and M~rre11,~~ who have allowed for con- figuration interaction between calculated excited states.The Pople and Santry method is a perturbational method for calculating coupling constants based on a MO approach but which avoids using the mean energy of excited states approximation. They considered only excited states arising from single excitations between molecular orbitals obtained from a minimum basis of valance-shell atomic orbitals. The theory is formulated on the basis of a one- electron Hamiltonian such that the excitation energies are taken as the difference between orbital energies and on the assumption of no configura- tional interaction between the excited states. This approach gave satisfactory results for hydrocarbons and simple heteroatomic molecules. By introducing configurational interaction between calculated excited states, Ditchfield and Murrell have obtained good agreement between observed and calculated directly bonded C-H and C-C coupling constants in hydrocarbons. The calculated non-bonded C-H coupling constants are also in better agreement than those obtained previously.Gil and Teixeira-Dias43 have considered the effects of substitution on directly bonded C-H coupling constants within the framework of the Pople and Santry MO theory. H i r ~ i k e ~ ~ has also considered this problem. P-P coupling constants have been explained45 satisfactorily by use of an extension of the Pople and Santry theory; JPP values can be either positive or negative, as was found with JCP values.46 A linear relationship between directly bonded C-H coupling constants and calculated (SCF method) bond orders has been found for a series of hydrocarbons.47 R o s ~ i ~ ~ has proposed a theoretical explanation of the effects of deuterium substitution of a vicinal proton on directly bonded JCH values on a tetrahedrally hybridised carbon atom.The change in the time- L. D. McKeever, R. Waack, M. A. Doran, and E. B. Baker, J . Amer. Chem. SOC., 40 E. Lippmaa, M. Appa, and A. Sugis, Eesti N.S. V . Teaduste Akudeemia, Toimetised, 41 J. A. Pople and D. P. Santry, Mol. Phys., 1965, 9, 31 I . 42 R. Ditchfield and J. N. Murrell, Mol. Phys., 1968, 14, 481. 4D V. M. S. Gil and J. T. C. Teixeira-Dias, Mol. Phys., 1968, 15,47. 44 E. Iiiroike, J. Phys. SOC. Jupan, 1968,24, 1348. 45 E. G. Finer and R. K. Harris, Chem. Comm., 1968, 1 1 0 . 46 W. McFarlane, Chem. Comm., 1967, 58.47 G . Berthier, H. Faucher, and D. Gagnaire, BUN. Sue. chim. Frnrrce, 1968, 1872. 48 M. Rossi, Chem. Phys. Letters, 1968, 2, 353. 1968,90, 3244. 1967, 3, 385.70 J . Feeney averaged 2s character of the hybrid carbon orbital due to coupling of the bending vibrations is calculated and good values of C-H coupling constants are predicted. has developed a VB description of the contact nuclear spin-spin contribution which does not invoke the mean excitation energy approxima- tion. By use of this approach he has calculated values for long range H-H coupling i-c-electrons which are in reasonable agreement with the experi- mental values. Some of the elegant measurements of coupling constants of fundamental importance are worth mentioning. By examining the AA'BB' lH spectrum from [l ,1,2,2,3,3,4,4-2H8]cyclohexane (II) under conditions of deuterium decoupling at -103" it was possible to extract the following coupling con- Another method of obtaining coupling constants between nuclei in fixed relative positions to each other is to examine rigid bicyclic compounds.In this way the angular dependence of vicinal lH-14N coupling could be determi~~ed,~l by examining molecules such as dibenzobicyclo [2,2,2]octa-2,5- dien-7-yltrimethylammonium bromide (111). The following results were obtained: dihedral angle 0", J N H ~ 2.7 (probably positive); 60°, J N H ~ < 0.3 Hz; 120", JNHd 0.8 Hz. For compounds with dihedral angles of 180" the J N H " ~ ~ stants :50 Jas,astrans 13.12, Je q,e qtrans 2.96, Jaz,e qcis 3-65, and Jgem - 13.05 Hz.value is predicted to be 5-6 Hz. Some coupling constant sign determinations would be useful on this system. JH-S-C-H and JH-S-C-CH have been shown52 to be both positive in sign by double-resonance experiments on the bicyclo-compound 1,2,3,4- tetr achloro-5-exo-mercapt o bicyclo [2,2,1] hept-2-ene. Many other relative sign determinations for homo- and hetero-nuclear coupling constants have 49 M. Barfield, J. Chern. Phys., 1968, 48, 4458, 4463. 51 Y. Terui, K. Aono, and K. Tori, J . Amer. Chem. SOC., 1968,90,1069. b2 V. F. Bystrov and 0. P. Yablonsky, J . Mol. Spectroscopy, 1968,26, 213. E. W. Garbisch and M. G. Griffith, J. Amer. Cliem. Soc., 1968,90, 6543.High Resolution Nidear Magnetic Resonance Spectroscopy 71 been reported, as typified by references 53 to 57.Several empirical correla- tions involving coupling constants have been pointed out. Values of JHC.OH in primary alcohols follow a Karplus relationship similar to that relating JHC. CH vicinal coupling with dihedral angle.5* Long range protoii-methyl coupling constants in propene, mesitylene, and other compounds are linearly related to the square of the mobile bond 0rder.5~ A linear correlation has been found between JCF in trifluoromethyl derivatives with carbon-substituent bond length JCF = -(106rcx + 115) where YCX is bond length in A.6o JCH in a series of methyl derivatives corre- lates linearly with the :product of the electronegativity EX and the bond distance rcx.61 The coupling constants in N-substituted pyridines show a similar dependence on the N-substituent as is found in monosubstituted benezenes.62 An additivity scheme for ortho-, meta-, andpara-proton coupling constants based on changes induced by a single substituent in model com- pounds has been applied to trisubstituted benzenes: only the ortho-coupling constants could be reproduced by this approach.63 Solvent effects on coupling constants, although small, can have important implications.Fineg01d~~ has examined the solvent dependence of vicinal coupling constants in systems which are rotationally invariant in cyclic rings; a gradation of gauche vicitzal coupling constants over a 14 % range with change of solvent points to the errors involved in using medium effects to solve rotational isomerism problems. A large solvent dependence of JPF in phosphorus trifluoride [JPF (gas) 1404 Hz; JPF (Cch solution) 1423 Hz] has been explained qualitatively in terms of electric field effects.65 JBF in silver tetrafluoroborate changes sign on changing solvent from water to aceto- This observation, together with the small magnitude OfJBF (ca.1 Hz), indicates that we have cancellation of large opposing coupling contributions. The reduced X-F coupling constants in the isoelectronic series BeF42-, BF4-, CF4, and NF4+ correlate with atomic number in the same manner as found in the series TeFs, SbFs-, and S ~ F G ~ - . ~ ~ From the analysis of the proton spectrum of formaldehyde [15N]oxime68 (IV) in methyl cyanide solu- 53 R. K. Harris and C. M. Woodman, J . Mol. Spectroscopy, 1968, 26, 432. s4 R. B. Johannesen, F.E. Brinckman, and T. D. Coyle, J . Phys. Chem., 1968,72,660. 55 T. C. Farrar, R. B. Johannesen, and T. D. Coyle, J . Chern. Phys., 1968,49,281. H. Dreeskamp and G. Pfisterer, Mol. Phys., 1968, 14, 295. 57 R . B. Johannesen, J . Chem. Phys., 1968, 48, 1414, 58 E. F. Kiefer, W. Gericke, and S. T. Amimoto, J . Amer. Chem. SOC., 1968,90, 6246. 59 D. J. Blears, S. S. Danyluk, and T. Schaefer, Canad. J . Chem., 1968,46,654. 60 I. Love, Mol. Phys., 1968, 15, 93. G2 S. Castellano and R. Kostelnik, J. Amer. Chem. Suc., 1968,90, 141. 63 T. Schaefer, G. Kotowycz, H. M. Hutton, and J. W. S. Lee, Cunad. J. Chem., 1968, C . P. Yue, Cunad. J . Chem., 1968,46, 2675. 46, 2530. H. Finegold, J. Phys. Chem., 1968, 72, 3244. G5 W. T. Raynes, T. A. Sutherley, H. J. Buttery, and C.M. Fenton, Mol. Phys., 1968, 66 R. J. Gillespie, J. S. Hartman, and M. Parekh, Canad. J . Chem., 1968, 46, 1601. 67 J. Feeney, R. Haque, L. W. Reeves, and C. P. Yue, Cunad. J . Chem., 1968,46,1389. 68 D. Crkpaux and J. M. Lehn, Mol. Phys., 1968,14, 547. 14, 599.72 J. Feeney tion values of JNHA (& 14.2) and JNHB (Zt2.3 Hz) have been obtained. Proto- nated aldehyde [15N]oximes also show similar but smaller effects.68 These results can be explained in terms of the orientation of the lone pair on the nitrogen in the formaldehyde oxime which appears to greatly increase JNHA without appreciably affecting JNHR. In the protonated species where the lone pair is replaced by an N-H bond, JNHA is much smaller than in formalde- hyde oxime. The stereospecific nature of long range H-H coupling is illustrated by the observation that protons on both sp2- and sp3-hybridised benzylic carbon atoms have maximum coupling with the ortho-protons when the benzylic proton is out of plane, while the reverse is true for the coupling with the rneta-protons.69 A large number of 2 J ~ . ~ ~ values have been measured in acyclic organo- phosphorus compounds and values between 0 and 22 Hz obtained. They could not be rationalised theoretically and even their empirical use for struc- ture determination requires caution if very close analogies cannot be found. 70 Chemical Shifts.-Relatively few papers dealing with proton chemical shift calculations have appeared in the past year. It is still not possible to carry out the necessary detailed quantum mechanical calculations to predict chemical shifts and one must resort to specific physical models (such as those based on anisotropic or electric field effects) to explain differences in chemical shifts, Fraenkel and his co-~orkers~~, 72 have postulated that there are significant paramagnetic contributions to the shielding of ortho-protons in aromatic organometallic compounds of lithium, magnesium, and calcium.Magnetic mixing of the ground and low-lying excited states gives paramagnetic con- tributions to the shielding, causing low-field shifts of the ortho-protons. Homer and Callagha~~~~ have re-examined carefully the approach of using neighbour anisotropic effects to predict chemical shifts. The anisotropy of the magnetic susceptibilities of C-C and C-H bonds was investigated and the values of Axcc and AxCH obtained from the n.m.r.measurements were shown to be incompatible with those found from Cotton-Mouton constants. A similar was undertaken for C-F and C-Cl bonds, and doubt has been cast on whether the n.m.r. estimates of AxCF and AxCC1 are meaning- ful. 69 G. P. Newsoroff and S. Sternhell, Austral. J . Chem., 1968, 21, 747. 7O M. J. Gallagher, Austral. J . Chem., 1968,21, 1197. 71 G. Fraenkel, D. G. Adams, and R. R. Dean, J . Phys. Chem., 1968,72, 914. 72 G. Fraenkel, S. Dayagi, and S. Kobayashi, J , Phys. Chem., 1968, 72, 953. 73 J. Homer and D. Callaghan, J . Chem. Soc. (A), 1968, 439. 74 J. Homer and D. Callaghan, J . Chem. SOC. (A), 1968, 518.High Resoliltion Nuclear Magnetic Resoriance Spectroscopy 73 The contribution to proton shielding from steric compression is known to result in a deshielding effect.Cheney75 has correlated this with the com- ponent of the non-bonded H-H repulsive force along the C-H bond axis. The model used postulates that there is induced electronic charge polarisation in the C-H bond as a result of the H-H interaction. An empirical linear relationship was found : = - 105 Z cos 8i exp (-2.671r;) p.p.m. i where 8i is the angle between C-H and the H . . . H internuclear line,and ri is the separation between the sterically interacting protons. The summation is taken over all hydrogen atoms which interact significantly with the hydrogen under study. A similar relationship has been used to predict contributions to 13C chemical shifts resulting from steric crowding. These effects are to be dis- tinguished from electric-field effects resulting from dispersion forces which show no angular dependence.Using the SCF-LCAO-MO method of Pople, E m ~ l e y ~ ~ has computed the 'H, 13C, and 14N chemical shifts in pyridine and pyridinium ion from the charge densities: he found that one cannot neglect c polarisation in such calculations. Tokuhiro and his co-workers calculated the diamagnetic and paramagnetic contributions to lH and 13C shielding in pyridine and found values in good agreement with experiment. Wu78 has used the P ~ p l e ~ ~ theory to calculate 14N chemical shifts. In the absence of low-lying transitions (n --f x*) which give rise to paramagnetic shifts, the 14N shifts were found to depend primarily on (i) the electron density on the nitrogen, (ii) the mean excitation energy, and (iii) the effect of multiple bonds on the nitrogen.Emsley80 has reviewed the relationship between charge densities and 19F chemical shifts in aromatic compounds. For fluorine nuclei para to the substituent a semi-empirical SCF-LCAO-MO method was used to verify the existence of the linear correlation between 19F chemical shifts and x- electron density on the fluorine or the attached carbon atom. 19F Shielding constants calculated by the method of Karplus and Pople, by use of an aver- age energy approach (AE = 8.15 ev) are of the correct order of magnitude, but the calculated changes with substitution are too small. This cannot be explained simply by using different AE values for different substituents and it was concluded that this approach cannot reliably predict lgF chemical shifts.By studying molecules such as (V) and (VI), where the substitution is remote from the fluorine nuclei,s1, 82 the only effects on the fluorine shielding 75 B. V. Cheney, J . Amer. Chem. SOC., 1968,90, 5386. 76 J. W. Emsley, J. Chem. SOC. (A), 1958, 1387. 77 T. Tokuhiro, N. K. Wilson, and G. Fraenkel, J . Amer. Chem. SOC., 1968,90, 3622. 78 T. K. Wu, J . Chem. Phys., 1968,49, 1139. 79 J. A. Pople, J . Chem. Phys., 1962, 37, 53, 60. 8o J. W. Emsley, J. Chem. SOC. (A), 1968, 2018. 8l E. W. Della, Chem. Comm., 1968, 1558. A2 M. J. S. Dewar and T. G. Squires, J . Amer. Chem. SOC., 1968,90,210.74 J, Feeney are those observed when the substitution causes conformational distortion, which is possible in some of the decalin derivatives.82 Molecular distortions in some bicyclic fluorides83 on being substituted have also been cited as the origin of I9F chemical shift changes accompanying substitution.Several empirical correlations between chemical shifts and other molecular parameters have been pointed out. Chemical shifts of benzylic protons in para-substituted toluenes give a linear correlation with Hammett c values of para-substituents.84 Likewise the arnine proton chemical shifts in a large series of substitute-d anilines show a linear correlation with appropriate Hammett cs c0nstants.~5 Proton chemical shields in polyenylic ions (VII) have been correlated with calculated carbon chargs densities.86 There has been an increased interest in the study of isotopic effects on chemical shifts.Isotopic chemical shifts observed on replacing protons by deuterium atoms in the isoelectronic series BH4-, CH4, NH4+ have been measured.87 Although deuteriation results in higher shifts for llB and I3C, lower shifts are observed for 14N, a fact which has been explained in terms of electric-field considerations involving the nitrogen atom, which is more electronegative than carbon or boron. F The calculated shift for 19F shielding between 2HF and IHF has been shown to be in good agreement with experiment.88 83 G. L. Anderson and L. M. Stock, J. Amer. Chem. Soc., 1968,90,212. 84 R. R. Fraser, Gurudata, C. Reyes-Zamora, and R. B. Swingle, Cunud. J. Cheni., 85 B. M. Lynch, B. C. MacDonald, and J.G. K. Webb, Tetrahedron, 1968,24, 3595. 88 D . Oseen, R. B. Flewwelling, and W. G. Laidlaw, J. Amer. Chem. SOC., 1968, 90, 4209. 8? M. Shporer and A. Loewenstein, Mol. Phys., 1968, 15, 9. 88 D. K. Hindermann and C. D. Cornwell, J. Chem. Phys., 1968,48,4148. 1968,46, 1595.High Resolution Nuclear Magnetic Resonarice Spectroscopy 75 When 19F chemical shifts of alkali fluorides are measured in light and heavy water solutions there is a linear variation in shift with isotopic composition of the solvent (total shift of 3 p.p.m. observed):89 increasing the heavy water content increases the 19F shielding. This can be interpreted in terms of changes in the vibrational energy of the fluoride ion F- (H20)z; resulting from isotopic substitution which changes the average excitation energy AE in the paramagnetic shielding term.By examining the IgF shielding of HF in the presence of varying amounts of potassium fluoride and using both light and heavy water as solvents, high field shifts have been found to accompany the substitution of H by D for both the F- (2.96 p.p.m.) and HF (4-7 p.p.m.) resonance frequencies.90 19F isotopic shifts resulting from substitution of 32S by 34s (natural abundance spectra) render fluorine nuclei non-equivalent in the normally symmetrical species S205F2, S206F2, and S308F2; one can thus observe the F-F coupling constants in such systems.g1 Empirical methods of calculating proton shifts in heteroaromatic com- pounds92 and polyhalogenoben~enes~3 have been proposed. Useful charts of proton shifts in oxygenated unsaturated aliphatic compoundsg4 and for acyclic methineg5 protons have been published.Carbon-13 Resonance.-Because of the vast potential of 13C studies it is not surprising that a great deal of effort is now being expended in this area. Improvements in basic instrumental sensitivity and the use of sensitivity enhancement techniques have reached the stage where they have largely overcome the enormous inherent sensitivity problem in the study of carbon-13 nuclei. It is anticipated that there will be an explosive increase in carbon-13 studies when the organic chemist solving molecular structural problems acquires access to suitable facilities for such investigations. The desirability of total heteronuclear decoupling as a method of simplifying and increasing the sensitivity of 13C spectra has already been in~licated.~, 9 6 9 97 This is nor- mally achieved by noise decoupling at the proton frequency but Lippmaa has achieved similar results by using a time-sharing spectrometer where he can use very high powered coherent r.f.decoupling fields to irradiate all protons in a molecule simultaneously.97 Lippmaa and his co-workers have published several papers illustrating the elegant use of 13C rssonance in the study of 89 C. Deverell, K. Schaumburg, and H. J. Bernstein, J . Chem. Phys., 1968, 49, 1276. K. Schaumburg and C. Deverell, J . Amer. Chem. SOC., 1968, 90, 2495. 91 R . A. Stewart, S. Fujiwara, and F. Aubke, J . Chem. Phys., 1968, 49, 965. 92 I. Nicholson, Chem. Comm. 1968, 1028. 93 B. Richardson and T. Schaefer, Canad. J .Chem., 1968,46,2195. 94 N. F. Chamberlain, Analyt. Chem., 1968,40, 1317. 95 0. Yamamoto, T. Suzuki, M. Yanagisawa, and K. Hayamizu, Analyt. Chem., 96 A. Sugis and E. Lippmaa, Eesti N.S. V . Teaduste Akadeernia Toimetised, 1967, 1, 81. 97 E. Lippmaa, T. Pehk, and J. Past, Eesti N.S.V. Teaduste Akadeemia Toimetised, 1968,40, 568. 1967, 3, 345.76 J. Feeney alkanes,98, 99 alkenes,loO alkynes,lol methoxybenzenes,lo2 and strained molecules such as adamantane.97 Other workers have chosen to use the in- direct method of measuring 13C frequencies by using a heteronuclear double- resonance technique, where one observes the effects on the proton spectra when the 13C absorption bands are irradiated.lo3 In this way, the 13C chemical shifts have been measured as a function of pH in sevcral amino-acids and peptides such as glycine, diglycine, triglycine, alanine, and alanylglycine.By calculating charge densities, the 13C chemical shifts of the a-carbon atoms could be predicted. It was also possible to rationalise the 13C shifts changes accompanying protonation (NH2 -+ NH3+) and ionisation (C02H -+ COZ-).~O~ Litchman and Grantlo4 have found that although the 13C chemical shifts in halogen-substituted methanes cannot be explained in terms of a simple additive substituent relationship, when one incorporates a pair-interaction term into a linear expression the chemical shifts can be calculated. In an investigation of methyl derivatives of Group IV and Group IIB rnetalslO5 the M-C coupling constants have been measured and discussed in terms of the contact mechanism.Buccilo6 has pointed out an interesting empirical corre- lation of 13C chemical shifts in methyl and ethyl derivatives with the Pauling electronegativity (E) and numbers of lone pairs (M) of the substituent. 8l3CH3M 8CH313CH2M == 208 - 45E + = 234 - 55E + 13M 7M Correlations between 13C shifts and charge densities have been pointed out for diazoles,l07 triazoleslo7 and some 5-membered nitrogen heterocyclic108 com- pounds. 13C shifts have been reported for tetraiodomethane,lo9 acetyl compounds,110 the triphenylcyclopropenium cation,lll azines and their proto- nated cations,l12 pyrazole and substituted pyrazoles,l13 and 2-substituted pyridines,ll* where it was found that substituent effccts are often different from those observed in monosubstituted benzenes especially at the 2-positions.98 E. Lippmaa and T. Pehk, Eesti N.S. V . Teudusre Akadeemia Toimetised, 1968,3, 210. 99 E. Lippmaa and T. Pehk, Kemian Teollisiiius, 1967, 24, 1. loo E. Lippmaa, S. Rang, 0. Eisen, and T. Pehk, Eesti N.S.V. Tedusre Akadeemia lol S. Rang, T. Pehk, E. Lippmaa, and 0. Eisen, Eesti N . S . Y . Tecrdusre Akadreinia 102 T. Pehk and E. Lippmaa, Eesti N.S. V . Teaduste Alcadeemia Toimetised, 1968,3, 195. lo3 W. J. Horsley and H. Sternlicht, J . Amer. Chem. SOC., 1968, 90, 373s. 104 W. M. Litchman and D. M. Grant, J . Amer. Chem. SOC., 1968, 90, 1400. lo5 F. J. Weigert, M. Winokur, and J. D. Roberts, J . Amer. Chem. SOC., 1965,90, 1566. 106 P. Bucci, J . Amer. Chem. SOC., 1968, 90, 252. lo7 B. M.Lynch, Chem. Comm., 1968, 1337. 108 R. J. Pugmire and D. M. Grant, J . Amer. Chem. SOC., 1968,90,4232. 109 0. W. Howarth and R. J. Lynch, Mol. Phys., 1968, 15, 431. I1O G. A. Gray, P. D. Ellis, D. D. Traficante, and G. E. Maciel, J . Magnetic Resonance, 111 G. J. Ray, A. K. Colter, and R. J. Kurland, Chern. Phys. Letters, 1965, 2 , 324. 112 R. J. Pugmire and D. M. Grant, J . Amer. Chem. SOC., 1965,90, 697. 113 R. G. Rees and M. J. Green, J . Chem. SOC. (B), 1968, 387. 114 H. L. Retcofsky and R. A. Friedel, J. Phys. Chem.. 1968, 72, 2619. Toimetised, 1967, 3, 351. Toimetised, 1967, 4, 346. 1969, 1, 41.High Resohition Nuclear Magrietic Resonance Spectroscopy 77 Carbon-13 spin-lattice relaxation times are often long and their values are of some practical importance.Lippmaa and his co-workers115 have measured 13C spin-lattice relaxation times in several organic compounds. Many publications have appeared dealing with information obtainable from 13CH satellites in the proton spectra of molecules. In the analysis of the 13CH satellite spectra of para-dihalogenobenzenes it was found that long-range isotopic shifts must be introduced to account for the asymmetry in the upper and lower field l3CH satellites.l16 Studies of Other Nuclei.-Only a few of the numerous n.m.r. studies involv- ing other nuclei can be mentioned. Witanowskill7 has found empirical correlations between 14N shifts and molecular structure for sp2-hybridised nitrogen atoms in a large number of compounds. These correlations could be explained in terms of the ground-state molecular orbitals.It was possible to account for the influence of electronegativity of substituents on the nitrogen chemical shifts without reference to excited electronic states. The hetero- nuclear double-resonance technique has been used to measure 14N shifts indirectly in ureas,l18 thioureas,l18 and formamides.llg The 14N chemical shifts in formamide differ by as much as 10 p.p.m. on going from an infinitely dilute solution in acetone to one in methanol, resulting from hydrogen bonding of the methanol to the carbonyl group of the formamides. Oxygen-17 measurements are useful for studying solvation effects of metal ions, as was illustrated in an investigation of CoII in water and methanol solutions120 and also in aqueous hydrochloric acid solutions.121 By using 170-enriched methanol, separate 170 signals for co-ordinated and bulk methanol can be distinguished.To see separate signals the electron spin relaxation time of the paramagnetic ion must be short compared with the hyperfine interaction between unpaired electrons and the relevant nucleus. Also, the ratio of exchange between bulk and co-ordinated molecules must be sufficiently slow not to average out the separate signals. Hydration of API and GaIII ions has also been studied by novel 170 resonance measurements to determine co-ordination numbers122 and the lifetimes123 of the water molecules in the hydration sphere. When cobalt(I1) ions are added to the aqueous solutions, the 1 7 0 signal of the non-co-ordinated water molecules moves to lower field than that of the co-ordinated ones.Liquid Crystal Studies.-Molecular geometries, signs of coupling constants, and chemical shift anisotropies have been measured in several systems 115 A. Olivson, E. Lippmaa, and J. Past, Eesti N.S. V . Teaduste Akadeemia Toimetised, 116 J. M. Read, R. W. Crecely, and J. H. Goldstein, J. Mol. Spectroscopy, 1968,25, 107. 117 M. Witanowski, J . Amer. Chem. SOC., 1968, 90, 5683. 118 P. Hampson and A. Mathias, J . Chem. SOC. (B), 1968, 673. 119 M. Kamei, BUN. Chem. SOC. Japan, 1968, 41, 1030. l2* D. Fiat, 2. Luz, and B. L. Silver, J. Chem. Phys., 1968, 49, 1376. 121 A. H. Zeltmann, N. A. Matwiyoff, and L. 0. Morgan, J . Phys., Chem., 1968,72,121. m2 R. E. Connick and D. Fiat, J . Chem. Phys., 1963, 39, 1349. 123 D. Fiat and R. E.Connick, J . Amer. Chey. SOC., 1968, 90, 608. XVI Koide, 1967, 390.78 J. Feeney (e.g. 3,3,3-trifl~oropropyne,l~~ p-benz~quinone,~~~ thiophen,126 and 1,3,5- trifl~orobenzenel~~) by this elegant procedure.124 -127 When racemic 3,3,3- trichloropropylene oxide was examined in an optically active liquid crystal solvent two different spectra (d and 1) were observed.12* Diehl and his co- w o r k e r ~ ~ ~ ~ have used both direct and moment methods of n.m.r. spectral analysis for oriented molecules containing two or three spin-systems. They have developed a modified LAOCOON I1 computer programme which was able to analyse iteratively the spectra obtained for symmetrical orrho-disub- stituted benzenes dissolved in a nematic phase.130 From a study of the temperature and concentration dependence of liquid crystal spectra131 it is found that the degree of solute orientation increases gradually with decreasing temperature and increases with decreasing solute concentration : in both cases the lines sharpen when the degree of orientation increases.Whereas the concentration dependence is found to be typical of the solute, the temperature dependence depends on the liquid crystal solvent rather than the solute. The degree of orientation also depends on the spinning speed of the sample if the axis of rotation is not along the magnetic field direc- tion.131 For very high magnetic fields (50 kG) it is possible to spin samples at speeds of 250 Hz without destroying the orientation.131 However, such high fields are invariably obtained by use of superconducting magnets, where the axis of rotation is parallel to the direction of the magnetic field so that spinning cannot destroy the orientation regardless of its rate.This has been verified e~perimental1y.l~~ For lower fields, such as 14,000 G, to maintain132 orientation it is necessary to have slow spinner speeds of 3-20 Hz, depending on the liquid ~rysta1.l~~ The use of oriented molecules to measure proton chemical shift aniso- tropies is limited in accuracy because of contributions to the observed shifts from changes in solvent environment in the isotropic and nematic p h a ~ e . ~ ~ ~ ~ 13* This is illustrated by a study of molecular hydrogen in the nematic phase where the measured chemical shift anisotropies are much larger than good theoretical values.Extensions to the experimental technique have involved applying electric 124 A. D. Buckingham, E. E. Burnell, C. A. de Lange, and A. J. Rest, Mol. Phys., 125 P. Diehl and C. L. Khetrapal, Mol. Phys., 1968, 14, 327. 1% P. Diehl, C. L. Khetrapal, and U. Lienhard, Cunud. J . Chem., 1968, 46, 2645. 127 C. T. Yim and D. F. R. Gilson. Personal communication. 12* E. Sackmann, S. Meiboom, and L. C. Snyder, J . Amer. Chem. SOC., 1968,90, 2183. 1z9 P. Diehl, C. L. Khetrapal, and U. Lienhard, Mul. Phys., 1968, 14, 465. I3O P. Diehl and C. L. Khetrapal, Mul. Phys., 1968, 15, 201. 131 P. Diehl and C. L. Khetrapal, Mol. Phys., 1967, 14, 283. 132 A. D. Buckingham, E. E. Burnell, and C. A. de Lange, Mol. Phys., 1968,15,285. 133 A. D. Buckingham, E. E. Burnell, and C.A. de Lange, Chem. Comm., 1968, 1408. 134 A. D. Buckingham, E. E. Burnell, and C. A. de Lange, J . Amer. Chem. Suc., 1968, 1968,14, 105. 90, 2972.High Resolution Nuclear Magnetic Resonance Spectroscopy 79 fields135 (which can rotate the molecular axis of the liquid crystal by 90") and using cholesteric liquid crystal phases.136 There are two cases of cholesteric liquid crystals which must be distinguished. In one case if the long axis of the molecule tends to align parallel to the magnetic field then no macroscopic alignment can take place unless the magnetic field is strong enough to un- wind the helical structure, and only then will a high-resolution spectrum of the solute be obtained. For the second case (which includes cholesterol deriva- tives) the long axis of the molecule tends to align perpendicular to the field, and the axis of the helical structure aligns parallel to the'magnetic field, such that no unwinding takes place: in such cases high resolution n.m.r.spectra of solute molecules can be observed. Biological Studies.-The availability of very high magnetic fields by use of superconducting systems (220 MHz for protons) is proving particularly useful to those interested in biological problems, where one is often dealing with broad, complex, overlapping absorption bands.137 By use of results obtained from studying an extensive series of commonly occurring amino- acids dissolved in trifluoroacetic [lH]- and [2H]-acids it was possible to analyse the 220 MHz spectrum of insulin.138 Bradbury and his co-workers139 have made a study of conformational analysis in polypeptides by use of 220 MHz n.m.r.spectra. Bradbury and Crane-R~binsonl~~ have indicated some of the exciting possibilities available through n.m.r. studies to those interested in biopolymers. There has been increasing use of the technique to study conformational and binding problems in enzymes. Eventually it will be possible to obtain detailed information concerning the nature of the active sites and the types of interactions with the substrates.140 For example, Gerig141 has made quantitative estimates of the bonding of tryptophan to p-chymo- trypsin by measuring line broadening, and Cohen and Jardet~kyl~~ have studied the 1H spectrum of lysozyme. Inter- and intra-molecular interactions in solutions of adenine nucleotides and other mononucleotides have been st~died.l4~ The n.m.r.spectrum of adenosine monophosphate (AMP) shows a similar concentration dependence to that of adenosine and it is concluded that AMP also forms vertical stacks in solution. Conformational studies of other nucleotides14'* and nucleosides,l45 and determinations of helix-coil 135 E. F. Carr, E. A. Hoar, and W. T. MacDonald, J. Chern. Phys., 1968,48,2822. 136 E. Sackmann, S. Meiboom, L. C. Snyder, A. E. Meixner, and R. E. Dietz, J. Amer. 137 E. M. Bradbury and C. Crane-Robinson, Nature, 1968,220, 1079. 138 B. Bak, C. Dambmann, F. Nicolaisen, E. J. Pedersen, and N. S. Bhacca, J. Mol. 139 E. M. Bradbury, B. G. Carpenter, C. Crane-Robinson, and H. W. E. Rattle, Nature, l40 Article 'Enzymes and N.M.R.', Nature, 1968, 218, 1107.141 J. T. Gerig, J. Amer. Chem. SOC., 1968, 90, 2681. 142 J. S. Cohen and 0. Jardetsky, Proc. Nat. Acad. Sci. U.S.A., 1968, 60, 92. 143 M. P. Schweizer, A. D. Broom, P. 0. P. Ts'o, and D. P. Hollis, J. Amer. Chern. Suc., 144 M. Smith and C. D. Jardetzky, J. Mol. Spectroscopy, 1968, 28, 70. 145 R. J. Cushley, J. F. Codington, and J. J. Fox, Cunad. J . Chem., 1968,46, 1131. Chern. SOC., 1968, 90, 3567. Spectroscopy, 1968, 26, 78. 1968, 220, 69. 1968,90, 1042.80 J . Feeney transition temperature~l~~ for different molecular weight polypeptides have been reported. Conformation Studies.-Considerable doubt has been cast on much of the published work on conformation preferences ( A values) of substituents in cyclohexyl systems in which it has been assumed that the rigid model 4-t-butylcycloliexyl derivatives have the same chemical shifts for the axiaI and equatorial methine protons as exist in the simple cyclohexyl derivative.147 By cooling cyclohexyl derivatives down to -80” it was possible in some cases to observe directly the axial and equatorial methine shifts and when these were compared with the 4-t-butyl-derivative values the assumption is found to be invalid. Eliel and Martin148 had earlier arrived at the opposite conclu- sion for cyclohexyl derivatives, but the weight of evidence is against their findings.However, Eliel and Martin14* found that 4-t-butyl derivatives are not good models for fluorine nuclei or for proton shifts in heterocyclic compounds. Of the large selection of conformational studies reported only a few can be mentioned.The barriers to chair-twist interconversion in 4-membered rings containing a trigonal sp2-carbon atom (cyclohe~anone~~~ and methylene- cyclohe~ane~~~) have been reported. Conformational studies of 1,3-dioxans,151 N-methylhydra~ones,~~~ and N N’N”N’”-tetramethylhe~ahydrotetrazine~5~ have also been made. Restricted rotation about C-C bonds in cis- and trans- but-2-ene154 and about P-N bonds155 in dimethylamino(pheny1) phosphine chloride has been observed. Kinetic Studies.-There has been continued investigation into the question of the validity of spin-echo measurements for measuring chemical exchange rates. It was found that this method gives values of rate constants which are too high in the slow-exchange region and too low in the fast-exchange region compared with the rates obtained by the method of complete line- shape ana1y~is.l~~ Delp~echl~~ has considered the coalescence with tem- perature of a complete spectrum arising from exchange between two identical configurations of one molecule in terms of coalescence of simpler partial spectra and he was able to build up a general computer programme for handling such problems.146 E. M. Bradbury, C. Crane-Robinson, H. Goldman, and H. W. E. Rattle, Nature, 147 F. R. Jensen and B. H. Beck, J . Amer. Chem. Sac., 1968,90, 3251. 149 F. R. Jensen and B. H. Beck, J . Amer. Chem. SOC., 1968,90, 1066. 150 J. T. Gerig, J . Amer. Chenz. SOC., 1968, 90, 1065. 151 E. L. Eliel and Sr. M. Carmeline Knoeber, J . Amer. Chem. SOC., 1968, 90, 3444.lS2 C. J. Karabatsos and R. A. Taller, Tetrahedron, 1968, 24, 3557. 153 J. E. Anderson and J. D. Roberts, J . Amer. Chem. SOC., 1968, 90, 4186. 154 H. G. Hecht and B. L. Victor, J . Amer. Chem. SOC., 1968,90, 3333. lS5 A. H. Cowley, M. J. S. Dewar, and W. R. Jackson, J . Amer. Chem. SOC., 1968,90, 156 P. T. Inglefield, €3. Krakower, L. W. Reeves, and R. Stewart, Mol. Phys., 1968, 15, 157 J. J. Delpuech, Mol. Phys., 1968, 14, 567. 1968,217, 812. E. L. Eliel and R. J. L. Martin, J . Amer. Chem. SOC., 1968, 90, 682. 4185. 65.Irish Rmolutiori Nircleas Magnetic Resonance Spectroscopy 81 Miscellaneous Studies.-La~zlo~~~ has written a comprehensive review on solvent effects in n.m.r., and many papers, too numerous to mention, have appeared illustrating the important analytical usage of such effects.The reaction-field model has been used to explain medium effects on frozen-out rotamers of 1,1,2,2-tetrabromofluoroethane at low temperat~es.l5~ Bec- consalPO has derived expressions for the effects of magnetically anisotropic disc- and rod-shaped solvent molecules on the nuclear shielding of a spherical solute molecule. The binary collision gas model used for interpreting pressure dependence of chemical shifts in gases has been extended to liquids consisting of non- polar solutes in magnetically isotropic solvent molecules.161 By observing the temperature and pressure dependence of the proton signal of hydrogen chloride in the gaseous state in the presence of dimethyl ether, hydrogen- bonding parameters for the gas phase interaction could be worked out.lG2 The origin of the line broadening of the n.m.r. signals from protons in the a-position to nitrogen in nitrogen heterocyclic molecules has been considered by Kintzinger and Lehn :163 the broadening is attributed to incomplete removal of N-Ha spin coupling by 14N quadrupolar relaxation. A study has been made of cross-relaxation contributions to relaxation times in systems containing protons in two different environments;164 these contributions arise from interactions between the different protons and are additional to the contributions from interactions between identical protons. Studies of nuclei attached to quadrupolar nuclei continue to provide interest. The observed temperature changes in the lgF spectrum of the hexafluoro- niobate ion have been explained in terms of chemical exchange of the fluorine atoms and also the effects of the quadrupolar induced transitions between the spin states of the g3Nb n~c1eus.l~~ It has been observed166 that the extent of collapse of multiplet structures for protons attached to loB is greater than for the analogous llB compounds, because of the larger quadrupole moment of 1°B. The extent of collapse also increases with field gradients at the boron nuclei in different compounds. When the temperature is increased, the extent of collapse decreases, as a consequence of the decrease in correlation time of the molecular motion.166 Gore and Gutowsky167 have used a density-matrix formulation to study the transient n.m.r. effects occurring when the r.f. field HI is changed dis- continuously for an AB system. undergoing rapid intramolecular exchange. 158 P. Laszlo, Progr. N.M.R. Spectroscopy, 1967, 3, 231. 159 G. Govil and H. J. Bernstein, J. Chem. Phys., 1968, 48, 285. l 6 O J. K. Becconsall, Mol. Phys., 1968, 15, 129. 161 F. H. A. Rummens, W. T. Raynes, and H. J. Bernstein, J. Phys. Chem., 1968, 72, 162 G. Govil, A. D. H. Clague, and H. J. Bernstein, J. Chem. Phys., 1968, 49, 2821. 163 J. P. Kintzinger and J. M. Lehn, Mol. Phys., 1968, 14, 133. 164 A. A. Brooks, J. D. Cutnell, E. 0. Stejskal, and V. W. Weiss, J. Chem. Phys., 1968, 165 D. W. Aksnes, S. M. Hutchison, and K. J. Packer, MoI. Phys., 1968, 14, 307. 166 H. Watanabe, T. Totani, M. Ohtsuru, and M. Kubo, MoI. Phvs., 1968, 14, 367. 167 E. S. Gore and H. S. Gutowsky, J. Chem. Phys., 1968,48, 3260. 2111. 49, 1571.

ISSN:0069-3022    

DOI:10.1039/GR9686500063

出版商:RSC    

年代:1968

数据来源: RSC

摘要:

5 Part (ii) INFRARED AND RAMAN SPECTROSCOPY By D. A. Long (School of Chemistry University of Bradford) FACED with the task of reviewing a field in which over three thousand papers were published last year this reporter finds himself somewhat in sympathy with one of Oscar Wilde’s maxims,’ ‘It is a sad thing that nowadays there is so little useless information’. There is of course implicit in this sentiment an acknowledgment of the value of so much that has been published. Perhaps this admission will serve as a general apology to those whose recent contribu-tions to spectroscopy have not survived the whims of personal selection or the restrictions on space. The new-style Annual Report requires in the language of the art phil-osopher a movement from realism towards idealism and expressionism.The varied nature of this particular subject has not lent itself to portrayal entirely in the new style and while for some topics a bold and no doubt subjective selection has been made for others an unashamedly detailed survey seemed more appropriate. The three thousand or so abstracts covering this field for the year 1968 may be compared with the mere eight hundred or so that it was necessary to survey for the two-year period 1962-1963 when preparing an earlier Annual Report2 on this topic. Although the year under review has seen no dramatically new advance in i.r. and Raman spectroscopy the continued expansion of the number of publications in this field is adequate testament to its value and interest for the chemist. The number of papers reporting Ranian spectroscopic studies is growing rapidly reflecting the impact that gas laser sources have made in this field.Lasers continuously tuneable over a limited frequency range are now available and it is an exciting prospect that CW lasers continuously tuneable over the whole visible region may become available for Raman spectroscopy. The potential applications of existing gas lasers to some aspects of Raman spec-troscopy have yet to be exploited. The study of rotational and vibration-rotation Raman spectra under high resolution and the Raman spectra of matrix-isolated species are two examples. The ease with which good quality Raman spectra can now be obtained from a minute amount of sample (which no longer needs to be colourless) will lead to Raman spectroscopy playing a role fully complementary to that at present enjoyed by i.r.spectroscopy in the laboratory of the analytical and preparative chemist particularly the inorganic Oscar Wilde ‘Complete Works,’ Collins London 1966 p. 1203. D. A. Long Atirr. Reports 1963. 8 84 D. A. Long chemist. A demand will soon be created for a laser-Raman instrument of modest price and intermediate performance. It is disappointing to have to record that the intriguing phenomena produced by giant-pulse lasers namely the stimulated inverse and hyper Raman effects have received little attention from chemists although they all open up novel possibilities. Several books and numerous reviews whose subject matter falls in the general field of i.r. and Raman spectroscopy have appeared.Only a limited and necessarily somewhat pcrsonal selection can be mentioned here. Schuler3 has written a review of Raman spectroscopy and lasers which includes a good coverage of the stimulated Raman effect. The excellent bibliography deals Rot only with the normal and stimulated Raman effects but also normal and stimulated Brillouin scattering and self-trapping. Bloembergen4 has written an authoritative review of the stimulated Raman effect and other phenomena which arise from the nonlinear third-order polarisation in the electric field strength. A survey of multiphoton spectroscopy byPetticolas5 includes aspects of the stimulated inverse and hyper Raman effects. A monograph entitled ‘Quantum Electronics’ by Yariv6 deals with the theoretical background to lasers and nonlinear optics.Adams7 has published a valuable critical survey of the i.r. and Raman spectra of metallic and organometallic compounds. Metal-halogen vibrational frequencies,8 far 4.r. spectros~opy,~ and i.r. spectroscopylO of aqueous solutions have all been the subject of reviews. Solid-state Studies.-The solid state is certainly the state of matter most frequently investigated by i.r. and Raman spectroscopy. The gas laser used in conjunction with a double monochromator makes it extremely easy to obtain the Raman spectra of solid compounds and a feature of the past year has been the substantial increase in Raman studies of the solid state. Excellent Raman spectra may be obtained from a minute amount of solid sample contained, for example in an ordinary melting-point tube and placed at the focus of a laser beam.Raman spectroscopy of solids is well on the way to becoming as routine an operation for the chemist as i.r. spectroscopy and will prove no less valuable. Indeed since low-lying frequencies are more easily obtained, Raman spectroscopy may often prove more valuable for the inorganic chemist. Where single crystals are available the laser makes it possible to illuminate along a given crystal axis with a defined state of polarisation and measure the intensity of Raman scattering essentially at right angles for each plane of polarisation. Systematic variation of the directions of illumination and 3 C. J. Schuler ‘Laser-Induced Spontaneous and Stimulated Raman Scattering,’ in Progress in Nidear Energy Series I X ed.H. A. Elion and 0. C. Stewart Pergamon Press, Oxford 1968 vol. 8 part 2. * N. Bloembergen Amer. J . Phys. 1967 25 989. 13 A. Yariv ’Quantum Electronics,’ John Wiley and Sons New Ynrk 1967. W. L. Petticolas Ann. Rev. Phys. Chenr. 1967 18 2 3 3 . D. M. Adams ‘Metal-Ligand and Related Vibrations,’ Arccjld London 1967. R . J. H. Clark Halogen Chenr. 1967 3 8 5 . J. W. Brasch Y. Mikawa and R . J. Jakobsen Appl. Spectroscopic Rep. 1968 1 187. lo F. S. Parker Progr. Infrared Spectroscopy 1967 3 75 Part (ii) Irlfuarcd and Hariiarz Spectroscopy 85 observation and the plane of polarisation of the incident beam enables the scattering tensor to be determined for each vibration. An unambiguous assignment of frequencies to symmetry classes can then usually be made.Measurements of this kind made previously using mercury-arc excitation were not only extremely laborious to obtain but usually inaccurate as a result of convergence errors and the ill-defined state of polarisation of the incident beam. The application of laser-Raman spectroscopy to single-crystal studies was pioneered by Porto and his co-workers.ll Their study of calcite has shown that the anomalies found by many previous workers are attributable to imperfect experimental procedure; it is salutory to note in passing that these anomalies had been given sophisticated theoretical explanations by a number of people. A recent paper12 discusses in detail the chemical applications of singlc-crystal Raman spectroscopy and shows that it can lead to unambiguous assignments whereas polarisation measurements on liquids or solutions (assuming a suitable solvent to be available) usually do not; e.g.polarisation measurements on an octahedral species MX6 in solution will distinguish v1 from v2 and v5 but v2 and v5 cannot themselves be unambiguously dis-tinguished in this way. Beattie and Gilson report single-crystal studies of KzPtC16 KzPdC16 KzSnC16 CSzGeCl6 and (NH4)sTeCb (all with octahedral anions) and K2PdC14 and (NH4)2PtC14 (square-planar anions) which coilfirm previous assignments. Other systems investigated were mercury(1) chloride, cassiterite anatase rutile gallium trichloride aluminium tribromide and molybdenum trioxide. The method relies on the validity of predictions of the Raman scattering tensor resulting from factor group analysis.The procedure is not without its complications and many pitfalls can beset the unwary. These include birefringence crystal imperfection dichroism the resonance Ramaii effect twinning enantiomorphism and internal reflection. The use of isotopic substitution and the examination of spectra at different temperatures can make valuable contributions to Raman studies of oriented single crystals. A recent typical example is the work of Rousseau et on Na14N03 and Na15N03. Earlier investigations of the Raman spectrum of sodium nitrate had revealed apparent anomalies in the scattering tensors for some modes and had left the assignments of the external modes in doubt. Using a laser source these workers found that apart from small intensity contributions in some orientations which are probably due to depolarisation effects arising from the birefringence of the crystal; the polarisability tensors of both internal and external modes are in accord with the selection rules.By comparing the spectra of Nal4N03 and Na15N03 they were able to assign the two external modes unambiguously. One of these is a libration and the other a translation of the NO- ion. A pure libration should show no nitrogen-l 1 S. P. S. Porto J. A. Giordmaine and T. C . Damen Phys. Rev. 1966 147 608: S. P. S. Porto P. A. Fleury and T. C . Damen Phys. Rev. 1967 154 522. I. R . Beattie and T. R. Gilson Proc. Roy. Soc. 1968 A 307 407. l3 D. L. Rousseau R. E. Miller and G. E. Leroi J . Cheni. Phys. 1968 48 3409 86 D. A . Long isotope independence since the centre of gravity of the NO; remains un-changed during the vibration whereas the translation should exhibit a shift.The calculated shift is small and is not observable at normal temperatures since the two lines are broad. Cooling to 35°K produced sufficient sharpening of these lines to reveal a frequency shift of the right order between Nal4N03 and Na15N03 for one frequency (98 cm.-l) and not the other (185 cm.-l). A brief selection of other crystals whose Raman spectra have been studied includes naphthalene anthracene,14 tungstates molybdates,l5 and rhombic sulphur.16 Ferroelectric materials have also been much investigated by both i.r. and Raman spectroscopy. Recent studies include BaTi03,17 NaN02,lS glycine silver nitrate,lg triglycine sulphate,20 potassium dihydrogen phos-phate,21 and SbSI.22 Raman spectroscopy has been proposed as a valuable method of studying transitions in solids,23 e.g.the of the Raman spectrum of a single crystal of SrTiO3 down to 2 5 " ~ . The results for a nominally pure single crystal are in accord with an ideal cubic perovskite structure (point group Oh) at higher temperatures with a transition at 110"~ to a tetragonal phase most probably C 4 h or D 4 h . Both phases show no first-order Raman scattering. An impure crystal undergoes different phase transitions as the temperature is reduced; it is probably tetragonal at 7 8 " ~ with point group S4 C4 or C 4 v and changes to a crystal structure of lower symmetry than tetragonal at lower temperatures. A selection of other spectroscopic studies of phase changes merit mention.Sodium nitrate shows a A transition at 267.7"~ when long-range order dis-appears. In the ordered structure the Raman-active totally symmetric stretching mode of the nitrate is i.r.-inactive but it is permitted for the lower site symmetry of the disordered lattice. The intensity of v1 in the i.r. can there-fore serve as a qualitative measure of short-range order. An i.r. over the temperature range 150"-300"c shows a band in the correct frequency range whose intensity variation with temperature is in accord with the vanish-ing of long-range order but gives no sharp change in short-range order in the transition region. The laser-excited Raman spectrum of crystalline quartz has 14 M. Suzuki T. Yokoyama and M. Ito Spectrochim.Acta 1968 24 A 1091; C. H. Ting J. Chinese Chem. SOC. Formosa 1967 14 1 17. ls R. K. Khanna W. S. Brower B. R. Guscott and E. R. Lippincott J. Res. Nut. Bur. Stand. Sect A. 1968 72 81. 16 A. T. Ward J. Phys. Chem. 1968 72 744. 17 M. DiDomenico jun. S. H. Wemple S. P. S. Porto and R . P. Bauman Phys. Rev., 1968,174,522 ; M. DiDomenico jun. S. P. S. Porto and S. H. Wemple Phys. Rev. Letters, 1967 19 855; D. L. Rousseau and S. P. S. Porto Phys. Rev. Letters 1968 20 1354. 18 E. V. Chisler and M. S. Shur Izvest. Akad. Nauk S.S.S.R. Ser.$z. 1967 31 1098. 19 P. S. Narayanan and A. V. R. Warrier Proc. 1st Internat. Conf. Spectroscopy Bom-20 I. Savatinova and P. Simova Optika i Spektrraskopiya 1968 24 218. 21 M. S . Shur Izvest. Akad. Nauk S.S.S.R. Ser.jiz. 1967 31 1042. 22 R. Bline M. Mali and A. Novak Solid State Comm. 1968 6 327. 23 R. Loudon Adu. Phys. 1964 13 423. 24 W. G. Nilson and J. G. Skinner J. Chem. Phys. 1968 48 2240. 35 W. L. Craft and L. J. Slutsky J. Chem. Phys. 1968 49 638. bay 1967 2 312 Part (ii) Infrared and Raman Spectroscopy 87 been studied26 over the temperature range - 196 to 615"c. A weak mode of A1 symmetry with a room-temperature Raman shift of 147 cm.-l grows in intensity and moves towards zero Raman shift as the a-p phase transition is approached. This mode clearly plays a fundamental role in the phase transi-tion but a full explanation could not be given. 1.r. and Raman ~pectroscopy~~ have been used to study the red to yellow phase transition in mercuric iodide for which the transition temperature at normal pressure is 127"c.Raman spectroscopy28 was employed to show that at normal temperatures and pressures in excess of 13 kilobars the red form of mercury(n) iodide changes to the yellow form with a spectrum identical with that obtained above 127"c at normal pressure. Schutte and Heyns29 have reported a series of low-temp-erature i.r. studies of structural changes with phase transitions in several ammonium salts. For example they find that in (NH4)2Cr207 free rotation of NH4f is indicated above 2 6 8 " ~ but this becomes hindered at lower tempera-tures. The transitions liquid-+solid I+sslid I1 in CF4 have been studied by i.r. and Raman spectros~opy.~~ Phase I appears to be a plastic crystal. The influence of external factors on the spectra of crystals has been the subject of some interesting studies.In their paraelectric phases (Oh) the first-order Raman spectrum is forbidden for SrTi03 and KTa03 but an electric field applied along a (001) direction induces C4v symmetry and renders all the phonons first-order Raman-a~tive.~~ ScotP2 has shown that the Bu modes of Caw04 and CaMo04 which are both i.r. and Raman-inactive in unstrained material become i.r. active and exhibit Eu characteristics under xz or y z shear strain and Au characteristics under xy shear strain. Such effects must be taken into account in interpretation of previously published i.r. spectra from crystals which may not be strain-free. Raman studies of polariton spectra have recently assumed importance. Polaritons arise in polar crystals from photon-phonon interactions.They give rise to Raman shifts in the scattering in the near-forward direction. The frequency shift is zero in the forward scattering direction but increases as the angle between the direction of observation and the forward direction varies over a small range typically 0.6-3.4". Raman scattering from polaritons was first observed by Henry and H ~ p f i e l d ~ ~ in GaP and subsequently in 211034 z6 S. M. Shapiro D. C. O'Shea and H. Z. Cummins Phys. Rev. Letters 1967,19 361. 27 A. J. Melveger R. K. Khanna B. R. Guscott and E. R. Lippincott Inorg. Chem., 1968 7 1630. 28 J. W. Brasch A. J. Melveger and E. R. Lippincott Chem. Phys. Letters 1968 2, 99; C. Postmus V. A. Maroni J. R. Ferraro and S. S. Mitra Inorg. Nuclear Chern.Letters 1968 4 269. 29 C. J. H. Schutte and A. M. Heyns Chem. Phys. Letters 1967,1,487; C. J. H. Schutte and A. M. Heyns ibid. 1968 1 511 515; C. J. H. Schutte ibid. p. 585. 30 R. P. Fournier R. Savoie F. Bessette and A. Cabana J . Chem. Phys. 1968 49, 1159. 31 P. A. Fleury and J. M. Worlock Phys. Rev. 1968 172,613; R. T. Schaufele M. J. Weber and B. D. Silverman Phys. Letters 1967 25 47. 32 J. F. Scott J . Chem. Phys. 1968,48 874. 33 C. H. Henry and J. J. Hopfield Phys. Rev. Letters 1965 15 964. 34 S. P. S. Porto B. Tell and T. C. Damen Phys. Rev. Letters 1966,16,450 88 D. A. Long and quartz35 by Porto and his co-workers. The theory of polariton spectra has been the subject of several recent papers.36 The importance of polariton spectra stems from their potential as tuneable sources of radiation.Such a source could be afforded by the angular dependence of the frequency of spontaneous Raman scattering involving polaritons but it would be of very low intensity. However stimulated Raman scattering involving polaritons would provide a tuneable source which was intense and coherent. It should be remarked that the ability to observe Raman scattering from polaritons, involving as it does observations so close to the forward direction is a striking example of the unique value of the laser for Raman investigations. Such observations could not have been made with conventional sources. The Raman spectra of the semi-conductors G ~ A s ~ ~ GaSe,38 and InSb39 have been studied and interpreted. Phonon dispersion curves for various silicon carbide polytypes have been plotted from Raman spectra data.40 Second-order Raman spectra of the alkali-metal halides are now being reinvestigated with laser excitation.Using a helium-neon laser and a double inonochromator the second-order Raman spectrum (0-400 cm.-l) of NaCl can be directly recorded in 20 min. whereas Toronto mercury arc excitation required a 48 hr. exposure a striking example of the revolution wrought by lasers. Recent studies include NaCl4I and KI.42 A theoretical ~tudy4~ has been made of the second-order Raman spectrum of CsF. Two different assumptions about the polarisability tensor lead to predicted spectra which while they are distinguishable experimentally have in common the unusual feature that they consist almost entirely of very sharp clearly separated lines.1.r. studies include NaC144 and LiF.45 It is interesting to note that two papers46 report that when gas lasers are focussed into LiF or NaCl single crystals lines attributable to multiple phonon processes are observed. The theory of the Raman effect in metals has received attention.47 Raman scattering from Be and AuA12 has been studied.48 In a Raman of 35 S. P. S. Porto Phys. Rev. 1967 162 834; J. F. Scott and S. P. S. Porto ibid. 161, 903. 36 E. Burstein S. Ushioda and A. Pinczuk Solidstate Comm. 1968,6,407; R . Ruppin and R. Englman J . Phys. (C) 1968,1,630; R . Fuchs and K . L. Kliewer J. Opt. SOC. Amer., 1968 58 319. 37 B. Tell and R . M. Martin Phys. Rec. 1968 167 381. 38 J. L. Brebner R. Loudon J. P. Russell and C.T. Sennett Proc. 1st Internat. Conf. Spectroscopy Bombay 1967 2 507. 39 J. Pons-Corbeau and J. Jouffroy Bull. SOP. Frartc. Mineral. Cristallogrphie 1967, 90 498. 40 D. W. Feldman J. H.Parker jun. W. J. Choyke and L. Patrick Phys. Reu. 1968, 173 787. 41 M. Krauzman Compt. rend. 1968 266 B 186. 42 M. Krauzman Compt. rend. 1967 265 B 689. 43 J. R. Hardy and A. M. Karo Phys. Rev. 1968 168 1054. 44 T. Kubota K. Hisano and 0. Matumura J . Phys. SOC. Japan 1968,25,642. 45 Z . G. Akhvlediani Izuest. Akad. Nauk S.S.S.R. Ser. $z. 1968 32 37. 413 G. Heilmann Z . Phys. 1968,215,43 1 ; G. Heilmann Z . Phys. ibid. 214,402. 47 D. L. Mills A. A. Maradudin and E. Burstein Phys. Reil. Letters 1968 21 1178. 48 D. W. Feldman J. H. Parker jun. and M. Ashkin Phys. Rev.Letters 1968,21 607. 49 A. Pinczuk and E. Burstein Phys. Rev. Letters 1968 21 1073 Part (ii) Infrared and Ranian Spectroscopy 89 InSb surfaces surface electric field-induced and resonance-enhanced Raman scattering from LO phonons has been observed and leads the authors to suggest that surface electronic properties may be studied by Raman spectro-scopy. The theory of Raman scattering from solids has been the subject of several papers.50 The interesting field of i.r. and Raman spectra induced or modified by impurities in crystals continues to grow. The introduction of impurities into the lattice of alkali-metal halides will modify the vibrations of the whole crystal lattice and may produce local or quasi-local vibrations associated with changes in the motion of neighbouring atoms.There have been interesting developments in the Raman spectroscopy of such crystals Stekhanov and E1iashberg5I reported the observation of quasi-local vibrations in the Raman spectra of KCl with Br- I- or Li+ impurities (0.1-0-5 mole %) the most characteristic feature being a strong line around 200 cm.-l. Stekhanov and Mak~imova~~ also reported quasi-local vibrations in the Raman spectra of KC1 doped with Na+ Cs+ or Rb+ (0-54-7 mole 7;). Recently Kaiser and MOckeP carried out a careful study of the Raman spectra of extremely pure KC1 and KC1 doped with 0.7 mole % KI. They found no trace of a line around 200 cm.-l and no evidence for an impurity-induced first-order Raman effect. They did find a Raman line at ca. 212 cm.-l in KCl :I- ( 0 3 %) coloured by exposure to a cobalt source.Since Stekhanov et al. had noticed that the mercury arc used in their experiments coloured the crystals it was concluded that the line at 212 cm.-l is associated with a colour centre in a doped crystal. Observations of Raman scattering from colour centres in pure crystals have been reported by Worlock and port^,^^ but there is no evidence for strong sharp localised modes of vibration. However, in KCl-KBr mixed crystals (20-80% KCl) a first-order spectrum has been reported by Porto et al.,55 the main feature of which is a band in the region 120-155 cm.-l of Alg symmetry. A first-order impurity-induced Raman spectrum has also been reported56 for KL with 2 mole % KN02. The theory of first-order Raman scattering by substitutional defects in alkali-metal halides has been treated by Benedik and Nardelli.57 There have been many i.r.studies of impurity-induced spectra. The following selected list of publications representative of the kinds of systems investigated shows how wide ranging these studies are U centres (H-50 A. K. Ganguly and J. L. Birman Phys. Rev. 1967 162 806; C. H. Ting Spectro-rhim. Acta 1968 24 A 1177; M. Hass and H. B. Rosenstock Appl. Optics 1967 6 , 2079; J. R. Hardy ‘Phonons Perfect Lattices Lattices Point Imperfections,’ Scot. Univ. 6th Summer School 1965,245; J. C. Decius J . Chem. Phys. 1968,49 1387. j1 A. I. Stekhanov and M. B. Eliashberg Soviet Phys. Solid State 1965 6 1 1 . 52 A. I. Stekhanov and T. I. Maksimora Soviet Phys. Solid State 1966 8 737. 53 R. Kaiser and P.Moeckel Phys. Letters 1967 25 749. 54 J. M. Worlock and S. P. S. Porto Phys. Rev. Letters 1965 15 697. 56 1. W. Shepherd A. R. Evans and D. B. Fitchen Phys. Letters 1968,27 171. 57 G. Benedek and G. F. Nardelli Phys. Rev. 1967 154 872. J. P. Hurrell S. P. S. Porto T. C. Damen and S. Mascarenhas Phys. Letters 1968, 26 194 90 D. A. Long impurity) in alkali-metal halides;58 rare-earth doped and hydrogenated CaFz showing spectra characteristic of rare-earth +HA ion pairs;5g OH- and OD-in alkali-nietal halides;60 atomic hydrogen in CaF2;61 BH4- and NH4f in alkali-metal halides;62 divalent anion and cation impurities in KCl;63 0 2 - in KBr;64 BO; in alkali-metal halides;65 CO;- NOS NO, and CO in alkali-metal halides;66 ZnS with Cu A1 or Ag A1;G7 silicon doped with boron and lithium;6s NaCl doped with AgCl and CuCl;69 simple mass defect effects resulting from natural 35Cl and 37Cl isotopic abundance in NaCl and LiF;70 the effect of an electric field71 on Li-doped KBr.The theory of impurity-induced i.r. spectra has also attracted attention.72 An interesting application of solid-state Raman spectra has recently been made by Bernstein and co-~orkers.~~ They have observed the vibrational frequency of 0,- in seven different alkali-metal halide hosts. The observed frequency varies with the host lattice. Plots of voo (the 0-0 electronic transi-tion frequency) us. the vibrational frequencyof 0,- are linear for a given cation and the lines for all cations intersect at the same point which corresponds to voo and the vibrational frequency for the free 0,- ion.The i.r. spectra of the various forms of ice have been studied.74 75 High-resolution Studies.-The study under high resolution of rotational and 58 D. Baeuerle and B. Fritz Phys. Status Solidi 1967 24 207; X. X. Nguyen Phys Rev. 1967 163 896; J. B. Page jun. and B. G. Dick Pliys. Rev. 1967 163 910; D Baeuerle and B. Fritz Solid State Comm. 1968 6 453. 59 S. Yatsiv S. Peled S. Rosenwaks and G. D. Jones Opt. Properties Ions Cryst., Conf. 1966 409. 6O A. I. Stekhanov and T. I. Maksimova Fiz. Tverd. Tela 1967 9 3668; M. L. Meist-rich U.S. Clearinghouse Fed. Sci. Tech. Inform. 1967; R. G. Grisar K. P. Reiners, K. F. Renk and L. Genzel Phys. Status Solidi 1967 23 613; M. E. Bauer and W. R. Salzman Phys. Rev. Letters 1967 18 590.61 R. E. Shamu W. M. Hartman and E. L. Yasaitis Plzys. Rev. 1968 170 822. 62 E. H. Coker and B. H. Campbell Proc. S. Dakota Acad. Sci. 1965 44 128; E. H. 63 D. N. Mirlin and I. 1. Reshina Fiz. Tverd. Tela 1968 10 1129. 64 0. Sild Eesti NSV Tead. Akad. Toim. 1968 17 203. 65 T. Mauring Eesti NSV Tead. Akad. Toim. 1968 17 232. 66 M. S. Pidzirailo and I. M. Khalimonova Ukrain. fiz. Zhur. 1967 12 1063; R. Metselaar and J. van der Elsken Phys. Rev. 1968,165 359; W. A. Morgan E. Silberman, and H. W. Morgan Spectrochim. Acta 1967 23 A 2855. 67 H. Kukimoto S. Shinonoya T. Koda and R. Hioki J. Phys. and Chem. Solids, 1968 29 935. 68 M. Balkanski J. Phys. (Paris) Colloq. 1967 2 8 , 69 R. Weber and F. Siebert 2. Phys. 1968 213 273. 70 M. V. Klein and H. F. Macdonald Phys.Rev. Letters 1968 20 1031. 71 R. D. Kirby and A. J. Severs Solid State Comm. 1968 6 613. 72 G. Benedek and A. A. Maradudin J. Phys. and Chem. Solids 1968 29 423; S. Takeno Progr. Theor. Phys. 1967 38 995; V. G . Koval'chuk Ukrain. fiz. Zhur. 1968, 13 437; T. C. Tak U.S. Atomic Energy Comm. 1967 Nuclear Sci. Abstr. 1967 21, 47027; T. P. Martin Phys. Rev. 1968 170 779. 73 J. Rolfe W. Holzer W. F. Murphy and H. J. Bernstein J. Chem. Phys. 1968 49, 963; W. Holzer W. F. Murphy H. J. Bernstein and J. Rolfe J. Mol. Spectroscopy 1968, 26 543. Coker and D. E. Hofer J. Chem. Phys. 1968 48 2713. 74 E. Whalley and J. E. Bertie J. Colloid Interface Sci. 1967 25 161. 75 J. E. Bertie H. J. Labbe and E. Whalley J. Chem. Phys. 1968,49 775; J. E. Bertie. H . J. Labbe and E.Whalley ibid. p. 2141 Part (ii) Infrared and Raman Spectroscopy 91 vibration-rotation spectra is a very important source of structural information yielding rotational and centrifugal distortion constants in the ground and excited vibrational states Coriolis coefficients anharmonicity constants etc. High-resolution i.r. spectroscopy is a well-developed technique but the selection rules are such that some molecules can only be studied by Raman spectroscopy. Even where i.r. spectroscopy is also applicable the selection rules are not identical with those for Raman spectroscopy and by and large, the Raman bands contain more information than the corresponding i.r. bands. Unfortunately experimental difficulties have in the past severely limited the experimental study of Raman spectra under high resolution.There was an important period of activity in this field in the 1950s when development of improved mercury lamps and mirror-type Raman cells with multiple cone collection made possible a series of elegant studies mainly by Welsh and S t ~ i c h e f f ~ ~ and their collaborators. The natural linewidth of the 4398 8, Hg line is ca. 0.2 cm.-l and so it is not possible to study with mercury-arc excitation systems with rotational spacings less than this; indeed the practical limit has been 0.245 cm.-1. Consequently this period of activity came to an end in the main when those relatively few molecules to which this limitation did not apply had all been studied. Other limitations were that the sample must be colourless and available in litre quantities.Thus as one author has recently said ‘High-resolution Raman spectroscopy still represents an as yet relatively untapped source of structural information’. The gas laser is potentially able to provide a sufficiently monochromatic source for Raman excitation to enable high-resolution Raman spectroscopy to be exploited to the full. Already the gas laser in its normally commercially available forms affords narrower line widths than the mercury arc (He-Ne, 6328 8, 0.04 cm.-l; Ar+ 5145 A 0.15 cm.-l) and very much narrower line-widths can be achieved with appropriate precautions albeit with a large drop in power. Despite this engaging prospect of virgin ground there appear to have been no publications dealing with laser-excited high-resolution Raman spectra of previously unstudied systems.Attention has so far been devoted to the prob-lems relating to sample illumination and scattered light collection and detec-tion. A recent series of papers traces the rapid progress in the solution of these problems. In 1965 Weber and port^^^ reported that they had obtained a good photographic record of the pure rotational Raman spectrum of methyl-acetylene at 0.5 atm. with an exposure time of 58 hr. using a cell (actual volume 340 ~ m . ~ ) with Brewster angle windows placed inside the cavity of a helium-neon laser of 20 mw output power. The volume of gas effective in producing scattering was only 0.59 ~ m . ~ compared with 3000 C M . ~ in a conventional mercury arc set-up but the exposure time was about 10 times longer.In 1967, Weber Porto Cheesman and Barrett78 showed that with a multipass Raman 76 B. P. Stoicheff ‘Advances in Spectroscopy,’ Interscience New York 1959 91. 77 A. Weber and S. P. S. Porto J . Opt. SOC. Amer. 1965 55 1033. 78 A. Weber S. P. S. Porto L. E. Cheesman and J. J. Barrett J Opt. SOC. Amer. 1967, 51 19. 92 D. A. Long cell inside the cavity of the helium-neon laser the photographic exposure time for rotational spectra were comparable with the classical mercury arc arrange-ment. They also made experimental tests of the effectiveness of various sample-illumination arrangements using a normal cell either inside or outside the cavity and found that the best arrangement was for the laser beam to be focussed into the cell inside the cavity. The illuminated volume effective in scattering is less than 10-7cm.3 in this arrangement and for a gas at normal pressure only 1011 molecules are involved.This optical arrangement the allied problems of collection of the scattered radiation and the superiority of pulse counting detection systems were fully discussed by Barrett and Adam~7~ in 1968. The two later papers present excellent pure rotational (02,N2,C02, CH3CzCH) and vibrational rotation spectra (02,N2 C02) directly recorded with the sample inside the cavity of an argon-ion laser with an intra-cavity power of 3-5 w at 4880 8 or 5145 A. Satisfactory pure rotation spectra using He-Ne 6328 A excitation (80 mw) and vibration-rotation spectra using Ar+ 4880 A excitation (1200 mw) can be recorded when the sample is placed outside the laser cavity at the focus of the laser beam and an additional con-cave mirror beyond the sample is adjusted to externally resonate the output laser beam through the sample.This sampling system is used on some commercially available Raman spectrometers with laser sources. Clearly then we can look forward to an expansion of activity in high-resolu-tion Raman spectroscopy. Although in the immediate future intensity con-siderations may restrict much of the work to excitation with the argon-ion laser where linewidths are only marginally less than that of the 4358 8 mercury line the need for only micro-samples and the partial removal of the restriction on colour open up interesting possibilities. The pure rotational Raman spectrumso of propyne and [2H4]propyne and the vibration-rotation Raman spectrums1 of [1,1 ,l-2H3]ethane have recently been investigated using con-ventional mercury-arc excitation.In high-resolution i.r. spectroscopy the current trend is the gradual refine-ment of an already well-established technique. Recent improvements in resolving power are leading to re-examination of the rotational and/or vibra-tion-rotation spectra of a number of relatively simple moIecuIes long familiar to those in this field. As a result new or improved values of many molecular constants are becoming available; the fine structure of bands can be more fully interpreted and better values of Coriolis coefficients may be obtained. A typical example is the recent study of the vibration-rotation bands of methylacetylene and [2Hl]methylacetylene by Thomas and Thompson.82 This has yielded values of aA and aB relating the rotational constants A and B in different vibrational states for a number of vibrations and all the Coriolis 79 J.J. Barrett and N. 1. Adams J . Opt. SOC. Amer. 1968 58 31 1. 8o S. I. Subbotin R. Kh. Safiullin V. I . Tyulin and V. M. Tatevskii Oprika i Spek-81 D. E. Shaw and H. L. Welsh Canad. J . Phys. 1967 45 3823. 82 R. K. Thomas and H. W. Thompson Spectrochim. Acta 1968,24 A 1337; R . K. Thomas and H. W. Thompson Spectrochim. Acta 1968 24 A 1353. troskopiya 1968 24 82 Part (ii) Infrared and Raman Spectroscopy 93 coefficients. Other molecules studied during the period under review include : C Z D ~ ~ ~ HF DF HCl and DCl,89 CH3F,So CH3Br,S1 CH3D,92 1°BF3 and 11BF3,93 13CH335C1 and 13CH337C1,94 trans-NzF~,~~ NF3,96 ~1H6]cyclopropane,97 HN3 HNCO HNCS and their deuterium derivative^,^^ allene,g9 and CH3C1.10° Griffiths and ThompsonlOl have considered the problems that arise in the analysis of the rotational spectra of symmetric-top molecules when the K structure is not resolved.Essentially the problem is that if resolution is in-complete and the K structure is not measurable the observed (average) posi-tions of the J lines will be determined by the intensity distribution within the K sub-structure of each J transition and this may vary significantly with J. It emerges that there are differences between prolate and oblate tops. For prolate tops @<A) the plot of v/(J + 1) us. (J + 1)2 will be linear with the intercept (2B - ~ D J K K ~ ) very close to 2B and the slope close to -405.With oblate tops the plot of v/(J + 1) us. (J+ 1)2 will lead to incorrect values of the rotational constants. The authors report new measurements on the rotational i.r. spectra of the prolate tops CD3F CD3Br and CDd CH3CN and CD3CN, and CH3CCH and CH3CCD. The values of B and DJ which they obtain are compared with microwave determinations and explanations are given of any differences. The rotational spectra of the oblate tops CHF3 and trimethyla-mine were also investigated. The microwave values of the rotational constants were used to calculate the apparent shift in the position of J lines due to unresolved K structure as a function of J. This shift increased from 0 to 0.20 cm.-l over the range J = 0 to J = 70.Only when this shift is corrected for 12C180,83 15N160 and 14N160,84 15N2180,85 12C16O2 and 13C1602,86 C2H2,87 83 C. Chackerian jun. and D. F. Eggers jun. J . Mol Spectroscopy 1968 27 59. 84 D. B. Keck and C. D. Hause J . Mol. Spectroscopy 1968,26 163. 85 J. L. Griggs jun. K. Narahari Rao L. H. Jones and R. M. Potter J . Mol. Spec-86 A. e Silva and M. Helena Ann. Phys. (Paris) 1967 2 217. 87 J. F. Scott J . Opt. SOC. Amer. 1968 58 142. 88 S. Ghersetti and K. Narahari Rao J . Mol. Spectroscopy 1968 28 27. 89 D. U. Webb and K. H. Rao J . Mol. Spectroscopy 1968,28 121 ; A. A. Mason and Alvin H. Nielsen J . Opt. SOC. Amer. 1967 57 1464; A. J. Perkins Spectrochim. Acta, 1968,24 A 285. R. Anttila and M. Huhanantti Canad. J . Phys. 1968 46 2025; W.E. Blass and T. H. Edwards J . Mol. Spectroscopy 1968 25 438. 9l R. Azria and C. Joffrin-Graffouillere Compt. rend. 1968 266 B 75. 92 C. Betrencourt-Stirnemann C. Alamichel and G. Graner Compt. rend. 1967, 93 C. W. Brown and J. Overend Canad. J . Phys. 1968,46 977. 94 J. L. Duncan and A. Allan J . Mol. Spectroscopy 1968 25 224. 95 S. Tung King and J. Overend Spectrochim. Acta 1967 23 A 2875. 96 G. W. Chantry H. A. Gebbie R. J. L. Popplewell and H. W. Thompson Proc. Roy. SOC. 1968 A 304 45; R. J. L. Popplewell F. N. Masri and H. W. Thompson, Spectrochim. Acta 1967 23 A 2797. troscopy 1968 25 34. 265 B 549; E. C. Leisegang and D. G. Parkyn J . S . African Chem. Inst. 1968 21 64. 97 J. L. Duncan J. Mol. Spectroscopy 1968 25 449. 98 B. Krakow R. C. Lord and G.0. Neely J . Mol. Spectroscopy 1968 27 148. 99 N. Van Thanh Ann. Phys. (Paris) 1967 2 241. loo M. Morillon-Chapey G. Graner and C. Alamichel Compt. rend. 1968 266 B, lol P. R. Griffiths and H. W. Thompson Spectrochim. Acts 1968 24 A 1325. 240 94 D. A . Long is a satisfactory linear plot of v/(J + 1) us. (J + 1)3 obtained which leads to the microwave values of Bo and DJ. In the case of trimethylamine an un-corrected plot leads to value of BO significantly lower than the microwave value. Inorganic Structural Studies.-The study of vibrational i.r. and Raman spectra to establish assignments of fundamental frequencies and to deduce symmetry or structure is well known and continues to be a major interest of chemical spectroscopists. The principles remain unchanged and the practice remains substantially the same apart from developments resulting from laser excitation of Raman spectra e.g.oriented single-crystal Raman spectroscopy which is discussed in another section. There seemed little point consequently in merely selecting at random a few examples since they would serve no more purpose than to illustrate already well-known principles. An attempt has therefore been made to present a reasonably complete survey of structural spectroscopic studies for smaller inorganic compounds ; most aspects of the extensive field of transition-metal complexes have had to be excluded on space considerations but complex halides have been treated fairly fully. 1.r. and Raman studies102 of XeOzF2 have been interpreted in terms of C2, symmetry with a molecule of pseudo-bipyramidal structure with the two fluorine atoms axial to the xenon and the two oxygen atoms with a lone pair in equatorial positions.The i.r. and Raman spectra of XeF6 have been studied.103 The i.r. spectrum shows more bond-stretching bands than would be expected for Oh symmetry. 1.r. studies of 104 alkali-metal aluminium hydrides MAlH4 (M = Li Na IS, Rb Cs) show that in solution A1H4 has a distorted tetrahedral structure and in the solid state the symmetry is lowered to C2,. A bridge structure is postu-lated. H H Force-constant calculations have been madelo5 for AU2C16 and it is found that the bridge bond stretching force constant is 73% of the terminal bond whereas in Al~Cl6 it is ca. 50%. The Raman and i.r. spectra of poly-crystalline TlAu(CN)2 suggestlo6 that while the Au(CN)2 grouping is linear, strong interaction with the T1 leads to an overall symmetry of C2,.Stammreichl07 has used Raman spectroscopy to determine the Hg-Hg lo2 H. H. Claassen E. L. Gasner H. Kim and J. Huston J . Chem. Phys. 1968,49,253. 1°3 H. Kim H. H. Claassen and E. Pearson Znorg. Chem. 1968,7 616; E. L. Gasner lo* T. G. Adiks V. V. Gavrilenko L. 1. Zakharkin and L. A. Ignateva Zhur. priklad. 105 D. M. Adams and R. G. Churchill J . Chem. SOC. (A) 1968 2141. 106 H. Stammreich B. M. Chadwick and S. G. Frankiss,J. Mol. Structure 1968,1 191. 107 H. Stammreich and R. Teixeira Sans J . Mol. Structure 1967 1 5 5 . and H. H. Claassen Inorg. ,Chem. 1967 6 1937. Spektroskopii 1967 6 806 Part (ii) Infrared and Raman Spectroscopy 95 stretching frequency in a series of mercury(1) salts and has discussed the varia-tions in relation to bond lengths.Raman spectra of mixtures of HgI2 with either molten HgC12 (220") or HgBrz (260") have been interpretedlo8 in terms of formation of mixed halides HgICl and HgIBr. Assignments have been made for yellow Hg212 yellow HgI2 and red HgIz from the Raman spectra of the s01ids.l~~ The far -i.r. spectra of these compounds and also HgC12 HgBr2, Hg(CN)2 HgO (red and black) and HgS (red and black) have been investi-gated.ll0 The formation of Hg(CN)I Hg(CN)312- and Hg(CN)21=- in aqueous solutions of NaCN and NaI to which Hg(II) has been added has been detected by Raman spectroscopy.lll The i.r. and Raman spectra of (H3N)zAgf and (H3N)2Hg2+ have been studied.l12 The metal-N skeletal vibrations give intense Raman scattering but the lines are very broad and only weakly polarised in aqueous solution; this is attributed to strong hydro-gen bonding between the ammine group and the solvent cage.Normal co-ordinate calculations were carried out for these ammines and for the isostructural dimethyl derivatives MezCd MezSn2+ MezHg MezTl+ and Me2Pb2+ and also for (F3C)zHg; variations in force constants are interpreted. 1.r. studies113 are in accord with a structure with twisted azide groups (CZ for Hg(N& and a symmetric trans-structure (C2h) for Hgz(N3)~. The i.r. and Raman spectra of Hg(CD3)2 have been reported for the first time for the liquid state.l14 Addison and coworkers115 have studied the i.r.and Raman spectra of solutions of zinc(II) cadmium(Ir) and mercury(I1) nitrates (and some halides) in acetonitrile. The nitrate spectra show strong perturbation of nitrate ions by solvated cations but only for Hg(I1) was a metal-oxygen frequency observed. In concentrated solutions of the zinc nitrate the principal zinc species is [Zn(MeCN)2][N03]2. 1.r. and Raman studies116 of aqueous solutions of Hg(N03)z and mixtures of Hg(NO&HzO and KN03 are con-sistent with the presence of Hg(N03)f and Hg(N03)~. 1.r. studies117 of various isotopic boroxines H3B303 in the gas phase support a non plan~(C3~) configuration rather than a planar model (D3h). Vibrational assignments have been made for some methyldiboranes.l18 The gas-phase i.r. spectra of MeBFz (and nitromethane) have been analysedllg in detail in terms of a theoretical model with free internal rotation.Cryoscopic and ebullioscopic evidence that Me2A.lF and EtzAlF exist as tetramers is supported by the i.r. 108 J. H. R. Clarke and C. Solomons J . Chem. Phys. 1968,48 528. log R. P. J. Cooney J. R. Hall and M. A. Hooper Austral. J. Chem. 1968 21 2145, ll1 J. Coleman R. A. Penneman L. H. Jones and I. K. Kressin Inorg. Chem. 1968,7, 112 M. G. Miles J. H. Paterson C. W. Hobbs M. J. Hopper J. Overend and R. S . 113 D. Seybold and K. Dehnicke 2. anorg. Chem. 1968 361 277. 114 J. L. Bribes and R. Gaufes Compt. rend. 1968 266 C 584. 115 C. C. Addison D. W. Amos and D. Sutton J . Chem. SOC. (A) 1968 2285. 116 A. R. Davis and D. E. Irish Znorg. Chem. 1968 7 1699. 117 F. A. Grimm L.Barton and R. F. Porter Inorg. Chem. 1968 7 1309. 118 J. H. Carpenter W. Jones R. W. Jotham and L. H. Long Chem. Comm. 1968,881. 119 W. J. Jones and N. Sheppard Proc. Roy. SOC. 1968 A 304 139. E. Decamps and A. Hadni J. Chim. p h j ~ . 1968 65 1030. 1174. Tobias Znorg. Chem. 1968 7 1721 96 D. A . Long and Raman spectra120 which also indicate that the AhF4 skeleton in these compounds is a planar 8-membered ring with D4h symmetry. The i.r. spec-trum121 of Ph3AI contains four AI-C frequencies and is consistent with a cen-trosymmetric dimeric structure. In an i.r. study1Z2 of Friedel-Crafts halides, heptane solutions e.g. of TiBr4-AIBr3 mixtures showed no detectable concen-tration of complexes such as TiBr3+ AIBr4- which have been assumed to be active catalysts of isobutene polymerisation.Raman studies123 of InCI2 suggest that in the molten state and probably also in the solid state the structure is In+InC14-. The i.r. spectra124 of benzene solutions of GaCl2 and GaBrz confirm the previously established ion-pair structure GaS(GaX4)- but the tetrahalogenogallate ion has symmetry lower than T d probably CzV. A speculative interpretation of the situation in benzene solution is that the gallium(1) ion has a co-ordination number of five being co-ordinated both to a benzene ring and to two of the chlorines in the anion. Beattie and co-w o r k e r ~ ~ ~ ~ have extended their earlier studies of Group 111 trihalides. The Raman and i.r. spectra of Akdh-6 AlzIs Ga2Br6 Ga2I6 and In& have been recorded in various phases and by extension of normal co-ordinate calcula-tions already developed for GazCIs complete assignments were made for all the compounds studied.Greenwood et aZ.,126 have investigated the i.r. and Raman spectra of the chloride bromide and iodide of indium(I1r). Their spectra support the bridged dimeric structure In216 with four-co-ordinate indium for the iodide by analogy with the spectra of GazBrs. The spectra of the chloride and bromide however indicate that in these two halides the indium is six-co-ordinate in a polymeric layer lattice. W a l t ~ n l ~ ~ has investi-gated the halides of thallium(rrr). A dimeric structure has been postulated for carbon suboxide polymer at room temperature from i.r. studies.12s 1.r. and Raman spectra129 support a CzV structure for the nitromethane anion in NaCH3NOz.An analysis of the group of bands in the Raman spectrum130 of liquid CC14 at ca. 775 cm.-l previously assigned to v3 and v1 + v4 shows that this group of bands arises almost entirely from difference bands and that the fundamental v3 although present is very weak. Silicon difluoride has recently been shown131 to be a relatively long-lived species (ca. 2 min.) and it has been possible to study the lZo J. Weidlein and V. Kreig J . Organonretallic Chem. 1968 11 9. lZ1 H. F. Shurvell Spectrochim. Acta 1967 23 A 2925. 192 P. Schmidt M. Chmelir M. Marek and B. Schneider Coll. Czech. Chem. Comin., lZ3 J. H. R. Clarke and R. E. Hester Chem. Comm. 1968 1042. 124 E. Kinseila J. Chadwick and J. Coward J . Chem. SOC. (A) 1968 969. 125 I.R. Beattie T. Gilson and G. A. Ozin J . Chem. SOC. (A) 1968 813. 126 N. N. Greenwood D. J. Prince and B. P. Straughan J . Chem. Soc. ( A ) 1968 1694. 127 R. A. Walton Inorg. Chem. 1968,7 640; ibid. p. 1927. 128 J. Wojtczak L. Weimann and J. M. Konarski Monntsh. 1968 99 501. 129 M. J. Brookes and N. Jonathan J . Chew?. SOC. (A) 1968 1529. 130 J. T. Kenney and F. X. Powell J. Chetn. Phys. 1967,47 3220. 131 V. M. Khanna R. Hauge R. F. Curl jun and J. L. Margrave J . Chetn. Phys. 1967, 1968,33 1604. 47 5031 Part (ii) Infiaved and Raman Spectroscopy 97 i.r. spectrum of the gas and determine two of the three fundamental frequen-cies (v1 and v3) for this molecule (symmetry CW). Combination of these frequencies with microwave data (v2 and centrifugal stretching constants) enabled a complete quadratic force field to be calculated.The relatively complicated i.r. spectra of gaseous silacyclobutane and [1 l-2H2]silacyclo-butane in the range 24-300 cm.-l have been analy~edl3~ in terms of a double minimum potential for the ring-puckering vibration. The potential barrier is calculated as 440 & 3 cm.-l and the dihedral angle of the puckered ring as 35-9 & 2" for silacyclobutane. Vibrational assignments133 have been made for the organogermanes CsHsGeBr3 and CsDsGeBrs. The far 4.r. spectra of a number of binuclear tin-metal compounds have been studied1Z4 to determine tin-metal frequencies. Whereas the Sn-Sn frequency is weak the tin-metal frequency in the heteronuclear tin-metal compounds is invariably reasonably strong.The vibrational spectra of zirconium tetrachloride indicate a poly-meric structure of ZrzCls units (symmetry D2h) and the spectra of Zr(N03)4 can be interpreted as for Sn(N03)4 and Ti(N03)4 with a zirconium co-ordina-tion number of 8 (symmetry D2d).135 An interesting Raman spectroscopic study136 of the chlorine isotopic structure of the totally symmetric Ti-Cl stretching frequency in Tic14 shows that it could not have originated from a simple tetrahedral molecule ( T d ) . An associated complex of lower symmetry (DM C2h or GV) with a chlorine-bridged dimer structure is suggested. Clark et aZ.13' have studied the MXs2-and MX4Yz2- ions (M = Ti Sn; X = Cl Br I). yield a band at ca. 1 . 5 ~ attributed to the solvated electron. Two Raman and i.r. studies139 of N2F4 agree in their conclusions that both trans(C2h)- and gauche(C2)-isomers are present at both ambient and lower temperatures and that both forms have similar energies.The i.r. spectrum of di-imide N2H2 has been analysed140 in terms of a planar trans-conformation. Provisional assign-ments have been made for the perchlorylimide anion (assumed GV) from i.r. studies141 of Kzf4NC103 and KPNC103 and for the anion HNClO3- (Cs symmetry) from i.r. spectra142 of KH14NC103 KH15NC103 KD14NC103, 1.r. reflection spectra of solutions of alkali metals in liquid 132 J. Laane and R. C. Lord J . Chem. Phys. 1968 48 1508. 133 J. R . Durig B. M. Gibson and C. W. Sink J. Mol. Structure 1968 2 I . 134 N. A. D. Carey and H. C. Clark Chem. Comm. 1967 292. 135 J . Weidlein U. Mueller and K.Dehnicke Spectrochim. Acta 1968 24 A 253. J. E. Griffiths J . Chem. Phys. 1968 49 642. 137 R. J. H. Clark L. Maresca and R. J. Puddephatt Znorg. Chenr. 1968 7 1603. 138 D. F. Burow and J. J. Lagowski J. Phys. Chem. 1968 72 169. 139 D. F. Koster and F. A. Miller Spectrochim. Acta 1968 24 A 1487; J. R. Durig I4O A. Trombetti Canad. J . Phys. 1968 46 1005. *41 A. I. Karelin Yu. Ya. Kharitonov and V. Ya. Rosolovskii Zhur. priklad. Spek-142 A. I. Karelin Yu. Ya. Kharitonov and V. Ya. Rosolovskii Zhur. priklad. Spek-and J. W. Clark J. Chem. Phys. 1968 48 3216. troskopii 1968 8 256. troskopii 1968 8 458 98 D. A. Long and KD15NC103. Goubeau et ~ 1 . ~ 4 ~ state that force-constant calculations144 for NC1032- suggest a double N-Cl bond and a single C1-0 bond.Force-constant calculations based on i.r. frequencies suggest that the geometry of N203 is not unequivocally known. The vibrational assignments for the hyponitrite ion have been revised as a result of further i.r. studies.145 The fundamental frequencies of NFClz have been assigned and c0rrelatedl4~ with those of the isoelectronic molecule CHFC12. Two groups of ~pectroscopists~~~ have intrepidly determined the Raman spectrum of NC13 and agree in finding four frequencies (two polarised) as expected for a pyramidal GV structure. Temperature-intensity i.r. studies14* of PF5 at - 85-25"c show that a band at 301 cm.-l previously assigned as a fundamental is a difference band. Raman and i.r. spectra149 of the solids PC15,SnCh and 2PC15,SnCh support the formulations [PCI4]+[SnC15]- and [PC14+]z[SnC162-] previously advanced on 31P n.m.r.evidence. The vibrational spectra150 of MePC14 suggest that in the solid the structure is MePC13'Cl- whereas in non-ionising solvents the compound is monomeric probably with CzV symmetry. 1.r. and Raman studies151 of polycrystalline PH41 and PDd yield assignments for internal and external modes and barrier heights for phosphonium ion re-orientation in the lattice. The vibrational spectra of solid diph~sphinel~~ are consistent with a trans-structure ( C Z ~ ) . Gas-phase i.r. studies153 of AsF5 confirm a triangular bipyramid structure. The i.r. and Raman spectra of solutions of AsC13 and AsBr3 in tributyl phosphate (TBP) have been interpreted154 in terms of formation of complexes AsC13,2TBP and AsBr3,2TBP.The Raman spectra155 of aqueous solutions containing [0H-]/[As1I1] in the range 3.5-15.0 give evidence for four species As(OH)3 AsO(OHz)- AsO~(OH)~- and AsO~~-. N.m.r. studies show AsC15,PC15 to have the ionic structure [PCh]+[AsCl6]-. The Raman spectrum156 of this compound contains a line attributable to PCl6- formed photochemically. The Raman spectra of solid SbCl5 indicates157 the existence of two forms. The spectrum of the low-temperature form is similar to that of 143 J. Coubeau E. Kilcioglu and E. Jacob Z . anorg. Chem. 1968 357 190. 144 W. A. Yeranos and M. J. Joncick Mol. Phys. 1967 13 263. 145 M. N. Hughes J . Inorg. Nuclear Chem. 1967 29 1376. 146 R. P. Hirschmann L. R. Anderson D. F. Harnish and W. B. Fox Spectrochim. 14' P. J. Hendra and J.R. Mackenzie Chem. Comm. 1968 13 760; M. Delhaye, 14* Sister R. M. Deiters and R. R. Holmes J . Chem. Phys. 1968 48 4796. 149 P. Reich and W. Wieker Z . Naturforsch. 1968 23 737. 150 I. R. Beattie K. Livingston and T. Gilson J . Chem. SOC. (A) 1968 1 . 151 J. R. Durig D. J. Antion and F. G. Baglin J . Chem. Phys. 1968,49 666. 152 S. G. Frankiss Inorg. Chem. 1968 7 1931. 153 S. Blanchard Commis. Energie Atom. (France) Rapport 1967 No. CEA-R 3195. 154 J. E. D. Davies and D. A. Long J . Chem. SOC. (A) 1968 1757. 155 T. M. Loehr and R. A. Plane Inorg. Chem. 1968,7 1708. 156 W. Wieker and A. R. Grimmer Z . Naturforsch. 1947 22 983. 157 K. Olie C. C. Smitskamp and H. Gerding Inorg. Nuclear Chem. Letters 1968 4, Acta 1968 24 A 1267. N. Durrieu-Mercier and M.Migeon Compt. rend. 1968 267 B 135. 129; I. Savatinova and M. Markov Zhur. priklad. Spektroskopii 1967 7 599 Part (ii) Infrared and Rnman Spectroscopy 99 SbC14f in SbC14F and indicates octahedral co-ordination with either a dimeric structure in which two octahedra have one common edge or an ionic structure [SbC14]+[SbCl~]-. The spectrum of the high-temperature form is very similar to that of liquid SbC15 and justifies the assumption that it is composed of trigonal bipyramidal molecules of SbC15. Nitrato-complexes of bismuth(Ir1) have been studied by Raman and i.r. spectroscopy.158 The spectra are con-sistent with a nitrate of CzV symmetry and bidentate co-ordination. Intensity measurements indicate the existence of stepwise complexes containing one to four nitrates per bismuth.A water molecule is also bound to the bismuth. The Raman spectra of liquid NbF5 and TaF5 show no evidence159 for polymerisa-tion. The i.r. spectra in the range 500-200 crn.-l of NbC15 NbBrs TaC15, TaBr5 and wc15 lead to interesting conclusions.160 The dimeric form is retained in solution for NbC15 and TaCl5 but whereas wc15 is a dimer as a solid it may dissolve as a monomer in nonpolar solvents. Some earlier assignments are questioned. The i.r. spectra of a number of compounds con-taining the cluster (Nb&l12)n+(n = 2,3,4) have also been studied.lS1 C2 Symmetry has been proposed for H2S2 and force constants calculated.lSZ The i.r. spectrum of thionitrosyl chloride163 shows that it has the structure NSCl (C,) and not SNCl. The i.r. and Raman spectra of SS have been investi-gatedlS4 and assignments made on the basis of the known puckered ring structure ( D 3 d ) .There have been several spectroscopic studieslS5 of the Group VI monohalides M2Xz (M = S,Se; X = C1,Br); assignments have been clarified and none of the results contradict the gauche structure of CZ sym-metry. Other sulphur compounds investigated include several trimethylsul-phonium and trimethylsulphoxium compounds166 and [(CF3)zCF]zSF2 and CF3SF3 ;IS7 vibrational assignments have been made. Spectroscopic studiesl68 confirm a regular octahedral structure (Oh) for gaseous UFS but in the solid a tetragonal D4h distortion is observed. In solution vibrational spectralS9 of Mo(CN)s4- and W(CN)s4- are consistent with a square Archimedian anti-prism structure of symmetry D4d but in the crystal the symmetry is lower ( D 2 d ) .Assignments have been made170 for WCb. 158 R. P. Oertel and R. A. Plane Inorg. Chem. 1968 7 1192. 159 H. Selig A. Reis and E. L. Gamer J . Inorg. Nuclear Chem. 1968 30 2087. 160 R. A. Walton and B. J. Brisdon Spectrochim. Acta 1967 23 A 2489. 162 Numan Zengin Comm. Fac. Sci. Univ. Ankara 1967 16 9. 163 A. Mueller G . Nagarajan 0. Glemser S. F. Cyvin and J. Wegener Spectrochim. Acta 1967 23 A 2683. 164 J. Berkowitz W. A. Chupka E. Bromels and R. Linn Belford J . Chem. Phys. 1967, 47 4320. 165 E. B. Bradley C . A. Frenzel and M. J. Mathur J . Chem. Phys. 1968 49 2344; P. F. Hendra and P. J. D. Park J . Chem. SOC. (A) 1968,908; E. B. Bradley M. J. Mathur, and C. A. Frenzel J. Chem.Phys. 1967 47 4325. 166 J. A. Creighton J. H. S. Green D. J. Harrison and S . M. Waller Spectrochim Acta 1967 23 A 2973. 167 K. Sathianandan and J. L. Margrave Indian J. Pure Appl. Phys. 1967 5 464. 168 R. Bougon Commis. Energie Atom. (France) Rapport 1967 No. 3235. 169 K. 0. Hartman and F. A. Miller Spectrochim. Acta 1968 24 A 669. 170 J. C. Evans and S. Y . S. Lo J. Mol. Spectroscopy 1968 26 147. R. A. Mackay and R. F. Schneider Inorg. Chem. 1968 7 455 1 00 D. A . Long The i.r. spectroscopy of HF vapour throws doubtlil on the existence of a trimer although higher polymers do contribute to the absorption. The first vibrational spectrum of a planar x2Y6 molecule has been r e p ~ r t e d l ~ ~ for dimeric ICh. Raman studies173 of IF7 and ReF7 in the vapour state confirm D5h symmetry for both species.The vibrational assignments of IF6f have been the subject of conflicting r e ~ 0 r t s . l ~ ~ 1.r. and Raman spectra of the solid CsIF6 indicate175 that IF6- does not have the symmetry Oh. The vibrational spectra of the complexes pyridine-bromine and pyridine-bromine chloride that whereas in nonpolar solvents they are predominently un-ionised, in polar media there exists the equilibria 2PyBrX + (Py2Br)+ + (BrX2)-X = Br or Cl). I-Ili7 and stretching frequencies in iodine and iodine bromide complexes with pyridine and related compounds have been studied by i.r. spectroscopy and the results interpreted in terms of strength of donor-acceptor interaction. Spectroscopic studies179 of the salts [R4NJf[BrHBr]-(R = Me Et Pr Bu or pentyl) in the solid and in various solvents show that the anion may be linear symmetric or linear unsymmetric depending on the environment.Vibrational assignments have been made180 for ICN. The i.r. spectrum of Br2O is in accordlsl with C2 symmetry. The Raman spectrum of the permanganates has been observed1g2 for the first time by use of helium-neon excitation. The vibrational spectra of solid per-rhenic acid183 is consistent with the structure HORe03 (C3,) but in aqueous solutions of 80 % concentra-tion complete dissociation to Re04- occurs. Complex halides have been much investigated and although space pro-hibits a fully comprehensive account they do merit reporting in some depth. 1.r. evidence suggests184 the existence of a flat symmetrical BaCli ion in BaC12-NaC1 systems.Raman studies have been made1s5 of MgCV- CdCI,, CdClF CdBr, and CdI The far 4.r. spectra of anions of general formula ZIIB~,C~~-~~- ZnBrnIn2- and ZnCln14-n2- have been recorded and inter-preted.186 The assignments for AlC1 have been ~1arified.l~~ A Raman study 171 D. F. Smith J . Chem. Phys. 1968 48 1429. 172 H. Stammreich and Y. Kawano Spectrochim. Acta 1968 24 A 899. 174 K. 0. Christe and W. Sawodny Inorg. Chem. 1968 7 1685; J. L. Hardwick and 175 K. 0. Christe J. P. Guertin and W. Sawodny Inorg. Chem. 1968 7 626. 176 S. G. W. Ginn I. Haque and J. L. Wood Spectrochim. Acta 1968 24 A 1531. li7 J. Yarwood and W. B. Person J. Amer. Chem. SOC. 1968,90 594. 178 Y. Yagi A. I. Popov and W. B. Person J . Phys. Chem. 1967 71 2439. lSo S. Hemple and E.R. Nixon J . Chem. Phys. 1967 47 4273. ls1 C. Campbell J. P. M. Jones and J. J. Turner Chem. Comm. 1968 888. lB2 P. J. Hendra Spectrochim. Acta 1968 24 A 125. H. H. Claassen E. L. Gasner and H. Selig J . Chem. Phys. 1968,49 1803. G. E. Leroi ibid. p. 1683. J. C. Evans and G. Y. S. Lo J . Phys. Chem. 1967 71 3942. K. I. Petrov V. A. Bardin and V. G. Kalyuzhnaya Doklady Akad. Nauk. S.S.S.R., la4 J. R. Chadwick P. J. Cranmer and H. C. Marsh J . Inorg. Nuclear Chem. 1967,29, lB5 J. E. D. Davies and D. A. Long J . Chern. SOC. (A) 1968 2054. lS6 G. B. Deacon J. H. S. Green and F. B. Taylor Austral. J . Chem. 1967 20 2069. Is7 D. E. H. Jones and J. L. Wood Spectrochim. Acta 1967 23 A 2695. 1968 178 1097. 1532 Part (ii) Infrared and Raman Spectroscopy 101 of molten cryolitelEE at 1030" suggests the existence of the equilibrium AlFs3- + AlF4- + 2F-.The Raman spectra of chloride melts indicate189 that in the system InC12-LiCl-KC1 Inch- and are present; in Inch-LiCl-KCl InCl.52- and -InCls2- are present ; and in BiC13-LiCl-KC1, BiC14- BiC1s2- and a little BiCls3- are present. Although (Bu2NH2)3BiCls and (BuNH3)2BiCl5 contain the expected complex anion (BuNH3)3SbCls does not contain SbCls3- but has the structure Raman intensity mea~urem.ents~~1 on aqueous solutions containing various ratios of chloride ion to bismuth reveal the existence not only of BiC14-, Dick2- and BiCk3- but also species containing three two and possibly one C1- per BPI; with bromide ions BiBrs3- and not BiBrs5- is the highest species formed.1.r. and Raman studies of complex halides of tin titanium, and tellurium indicate1g2 the trigonal bi-pyramidal structures for the ions SnC15- SnBrs- and TiC15- but a symmetry lower than square pyramidal for Tech-. For (Et4N)Te3C113 the probable structure is (Et4N+)(TeC13+)3(Cl-)4 whereas (Et4N)TizCh appears to contain the chlorine-bridged structure Ti9Clg-. The Raman spectra of the molten complexes SF4SbF5 SeF4SbF5, SeF4AsF5 and TeF4SbF5 indicatelg3 ionic formulations [SF3]+[SbF6]-, [SeF3]+[SbFs]- and [SeF3Ii-[AsF6]- for the first three; the situation is less clear for the fourth but plausible alternatives are ionisation into distorted ions TeF3+ and SbF6- or a fluorine-bridged structure. The Raman spectra of complex anions of formula MXs2- (M = Se Te; X = C1 Br r) have been recorded in the solid phase and in solution;194 evidence was found for SeBr5-in solutions of SeO2 in concentrated HBr.The vibrational spectra of the tetra-halogeno-complexes of MnII have been studied assignments made and force constants cal~ulated.1~5 The intensities of the Raman frequencies of PtCls-have been reinvestigated and previously reported anomalies discussed and ~1arified.l~~ Vibrational assignments based on On symmetry have been madelg7 for Ucls2- and ThC162- from Raman and i.r. spectra. Thallium halides and complex halides have been the subject of a number of studies.198 lS8 C. Solomons J. H. R. Clarke and J. O'm. Bockris J. Chem. Phys. 1968 49 445. lSg J. T. Kenney and F. X. Powell J . Phys. Chem. 1968 72 3094. Ig0 R. A. Walton Spectrochim.Acta. 1968 24 A 1527. 191 R. P. Oertel and R . A. Plane Inorg. Chem. 1967 6 1960. Ig2 J. A. Creighton and J. H. S. Green J . Chem. SOC. ( A ) 1968 808. 193 J. A. Evans and D. A. Long J . Chem. SOC. (A) 1968 1688. lg4 P. J. Hendra and Z. Jovic J . Chern. SOC. ( A ) 1968 600. I95 H. G. M. Edwards M. J. Ware and L. A. Woodward Chem. Comm. 1968 540. Ig6 D. W. James and M. J. Nolan Inorg. Nuclear Chem. Letters 1968 4 97. lg7 L. A. Woodward and M. J. Ware Spectrochim. Acta 1968 24 A 921. lg8 T. Barrowcliffe I. R. Beattie P. Day and K. Livingston J. Chem. SOC. (A) 1967, 1810; D. M. Adams and D. M. Morris J . Chem. Sac. (A) 1968 694; M. J. Taylor ibid., p. 1780; J. E. D. Davies and D. A. Long ibid. p. 2050 102 D. A. Long Matrix-isolated Species.-Vibrational spectroscopy of species containing only a few atoms has always been particularly satisfying in that such systems offer the best chance of detailed and unambiguous analysis.The matrix-isolation technique in which normally unstable species are preserved in matrices of noble gases or nitrogen at very low temperatures has opened up a rich field of new simple species for spectroscopic study. An appreciable number of new species has been identified and studied in the past year by use of this technique. Many of these species are of interest as likely reaction intermediates and they often pose interesting structural problems. So far only i.r. spectroscopy appears to have been used for study of matrix-trapped species but with the availability of laser sources Raman spectroscopy should soon be playing a complementary role.In reviewing this interesting field the reporter has chosen to illustrate the technique by giving some detail of one investigation and to summarise the results for a number of other systems. Noble and Pimentellgg have recently made an i.r. study of the products condensed out at 1 4 " ~ from hydrogen chloride-chlorine-argon mixtures (1 1 160) after passage through a glow discharge. Apart from bands attribxtable to known species (e.g. HCl CIS) and to impurities (nonreproducible intensity) two new i.r. bands were observed. Since the only atoms available in quantity are hydrogen and chlorine these bands must arise from a new species of general formula HnCI,. One of the bands appeared at 956 cm.-l with associated bands at 952 cm.-l and 949 cm.-l.The relative intensities of these three components, approximately 9 :6 1 were those expected for the vibration of two equivalent chlorine atoms of natural isotopic abundance. The simplest molecule con-sistent with this observation would be the symmetric structure ClHCl of symmetry Dmh. On deuteriation the band at 956 cm.-l shifts to 729 cm.-l. The second and stronger i.r. band occurred at 696 cin.-l and showed no chlorine-isotope structure but shifted to 464 cm.-l on deuteriation. The band at 696 cm.-l is assigned to the i.r.-active antisymmetric stretch v3 and the band at 956 cm.-l to v1 + vg. The observed chlorine-isotopic shifts are fully compatible with these assignments. The frequency shifts on deuteriation can only be calculated satisfactorily with a quadratic-quartic function with a high quartic content but this is physically reasonable and a similar instance has been reported for HF2-.These results lead to a value of 259 cm.-l for v1, the symmetric stretch which for Dcoh symmetry should be Raman-active only. A band at this frequency was definitely absent from the i.r. spectrum, further supporting the centrosymmetric structure. The i.r.-active bending mode v2 was not observed although it is expected to lie well above 200 cm.-l, the lower frequency limit of the i.r. spectrometer used but it would be expected to have low intensity. The HCl2 species could play a role in transient phenomena involving chlorine atoms such as in the H2-CI2 explosion and in the HCl chemical laser.The reaction between lithium atoms and methyl halides during condensa-Ig9 P. N. Noble and G . C. Pimentel J . Chem. Phys. 1968 49 3165 Part (ii) Infrared and Raman Spectroscopy 103 tion in solid argon at 1 5 " ~ produces the methyl radical which is shown to be planar (D3h); any equilibrium deviation from planarity in excess of 5" is excluded.200 CF2 is formed when carbon atoms produced by photolysis of cyanogen azide react with molecular fluorine in argon201 at 1 4 " ~ . The com-plete valence force field has been calculated from the observed frequencies for WF2 and WF2 for a nonlinear structure. Photolysis of the CF2 sample produces CF3 which has a GV structure with an estimated 13" deviation from planarity. NCN produced by the photolysis of CNmN3 isolated in an argon matrix at 1 4 " ~ reacts readily with F2 to give NFzCN.If both fluorine atoms and NCN are present in appreciable concentration a new species tentatively identified as the free radical FNCN is produced.202 The reaction of photo-lytically produced F atoms with FCN in argon or nitrogen matrices at 1 4 " ~ yields203 the free radical F2CN. Further fluorine atom attack on FzCN produces inter aha CFzNF and CF3NF2. Simultaneous condensation of beams of lithium atoms and CBr4 in argon gives204 the radical CBr3. A secondary reaction of lithium atoms with CBr3 yields CBr2. Similarly the matrix reaction of lithium atoms and CC14 givesz05 the radical CCh followed by CC12 and the matrix reaction of alkali-metal atoms and perbromochloro-methanes produces206 the trihalogenomethyl radical followed by the tri-halogenomethyl alkali-metal compound.Vacuum-u.v. p h o t o l y ~ i s ~ ~ ~ of SiHzCl2 and SiDzClz in an argon matrix at 1 4 " ~ produces SiC12. Photolysis of OClO in an argon matrix at 4 " ~ gives208 the radical C100. There is some evidence that a structural isomer of this radical also exists. The free radical NCO is producedzo9 by vacuum-u.v. photolysis of HNCO in matrices at 4"-14"~. The action of a microwave discharge on a mixture of a noble gas, chlorine and bromine followed by condensation at 2 0 " ~ produces210 a series of new chlorobromo-compounds for example BrBrClz. These are T-shaped molecules analogous to ClF3. Evidence has been found21l for Li2F2 dimers (possibly linear with Cmv symmetry) when the vapour from solid LiF is deposited in an argon matrix at liquid helium temperature.The spectra of HCN and DCN in argon nitrogen and carbon monoxide matrices show no evidence for rotation of the monomer; the dimer has a linear or near-linear HCN-HCN structure.212 Cyclic dimers of HC1 have been identified213 2oo L. Andrews and G. C. Pimentel J . Chem. Phys. 1967 47 3637. 201 D. E. Milligan and M. E. Jacox J. Chem. Phys. 1968 48 2265. 202 D. E. Milligan and M. E. Jacox J . Chem. Phys. 1968 48 481 1. 203 M. E. Jacox and D. E. Milligan J . Chem. Phys. 1968,48,4040. 204 L. Andrews and T. Granville Carver J . Chem. Phys. 1968 49 896. 205 L. Andrews J . Chem. Phys. 1968 48 979. 206 L. Andrews and T. G. Carver J . Phys. Chem. 1968 72 1743. 207 D. E. Milligan and M. E. Jacox J .Chem. Phys. 1968,49 1938. 208 A. Arkell and I. Schwager J . Amer. Chem. SOC. 1967 89 5999. 209 D. E. Milligan and M. E. Jacox J . Chem. Phys. 1967 47 5157. 210 L. Y. Nelson and G. C. Pimentel Inorg. Chem. 1968 7 1695. 211 S. Abramowitz N. Acquista and I. W. Levin J . Res. Nat. Bur. Stand. Sect. A , 212 C. M. King and E. R. Nixon J . Chem. Phys. 1968 48 1685. 213 B. Katz A. Ron and 0. Schnepp J . Chem. Phys. 1967 47 5303. 1968 72 487 104 D. A . Long in HCI-Xe matrices. The spectrum of the species formed when krypton and chlorine (100 :1) are passed through a microwave discharge and then con-densed at 2 0 " ~ is attributable214 to a normally symmetric Cl3 species per-turbed from Adsorbed Species.-The amount of published work in this field is similar to that in recent years; a substantial portion continues to come from Russian laboratories.Almost all of it deals with i.r. investigations of adsorbed species. The potentialities in this field of Raman spectroscopy particularly with laser sources indicated in the previous report appear scarcely to have been exploited. Only two publications involving Raman spectroscopy were noted. Hendra and Loader continuing earlier have reported a study of acetaldehyde216 adsorbed on silica gel and suggested that condensation of acetaldehyde was catalysed at the surface to produce a physically adsorbed cyclic product. Pershina and Raskin217 investigated the Raman spectra of a variety of substances adsorbed at the surface of microporous glass silica gel, and aluminosilica gel. Perhaps their most interesting observations were that the spectra of SbC13 SbBr3 and dichloroethane show frequency shifts and intensity changes on absorption very similar to those observed for the phase transition liquid- crystal.An extensive review of i.r. spectroscopy of adsorbed molecules which covers 1950-1967 has been published.218 Of papers published during 1968 the reviewer has selected for mention the following as revealing interesting structural information. Studies of the adsorption of formic acid on Vz05 showed219 the existence of three species HC02H HzO and HC02-. Since e.s.r. studies show the formation of V4+ on adsorption it is suggested that the HC02- ions are adsorbed on V4f ions. The adsorption of benzene on highly dehydroxylated Aerosil (surface OH concentrations 1-2 OH/100 A2) occurs at low coverages mainly on OH sites through a 1 :1 interaction involving the benzene x-elec-trons.220 Adsorption on dehydroxylated areas becomes important when a significant number of OH sites have been utilised.The NH2 radical has been identified as forming on the surface of iron dispersed in silica when exposed to ammonia or mixtures of hydrogen and nitrogen in the temperature range 20-500"~. The role of this intermediate in the synthesis of ammonia is dis-cussed.221 Studies of the chemisorption of HCN C2N2 and BrCN on silica-based transition metals (Ni Ir Rh Pt and 0s) showed that chemisorption to Cm by an asymmetric lattice cage. 214 L. Y. Nelson and G. C. Pimentel J . Chem. Phys. 1967 47 3671. 215 P. J. Hendra and E.J. Loader Nature 1967 216 789. 216 P. J. Hendra and E. J. Loader Nature 1968 217 637. 217 E. V. Pershina and Sh. Sh. Raskin Optika Spektroskopiya Aknd. Nnirk. S.S.S.R., 218 M. R. Basila Appl. Spectroscopic Rev. 1968 1 289. 219 M. Adachi T. Imanaka and S. Teranishi J . Chem. SOC. Japan 1968 89 446. 220 A. Zecchina C. Versino A. Apiano and G. Occhiena J . Phys. Chew. 1968 72, 221 T. Nakata and S. Matsushita J. Phys. Chem. 1968,72 458. Otdel Fiz.-Mat. Nauk 1967 3 328. 1471 Part (ii) Infrared and Raman Spectroscopy 105 predominates.222 It suggested that all three gases dissociate to give CN radicals which then give cyanides (M-CN) and cyanates and isocyanates (MOCN MNCO) formation of the latter involving a hydrolysis reaction with OH groups. 1.r. frequencies characteristic of Fe3+ +NO+ Fez+ +NO have been observedz23 for NO adsorbed on Fez03 gel.The richness of the i.r. spectrum of hydrogen (or deuterium) chemisorbed on polycrystalline rhodium substrates is attributedz24 to multiple crystallographic sites available on the polycrystalline substrate. Bands are assigned to hydrogen molecularly chemisorbed in a linear fashion and to H-H and Rh-H stretches in a Rh-H-H-Rh bridged structure. Polarised i.r. studies of thin urea (and thiourea) films several hundred angstroms thick on a steel surface that the c-axis of the urea crystal was parallel to the metal surface and the c-axis of the thiourea crystal nearly parallel; there was also evidence that the urea molecule is deformed from the planar structure. Absorption bands observed in the 17OO-2200 cm.-l region when strongly electronegative gases like oxygen and chlorine are absorbed in ZnO powder have been identified226 as vibrations of bonds between carbon and nitrogen impurity atoms and oxygen in the bulk of the ZnO.It is suggested that adsorption of electron-withdrawing gases brings about this bond formation which appears to be limited to a depletion layer of 10-30 A in depth. Force-constant Studies.-The calculation of force constants continues to to attract the interest of a considerable number of spectroscopists. Significant studies on small molecules continue to be made; e.g. the fundamental harmonic vibrational frequencies and centrifugal distortion data for 16014N35C1 ls015N35C1 18014N35C1 and 1s015N35C1 have been used to calculate227 a unique general quadratic force field; the observed vibra-tional frequencies and rotational constants Bv for six isotopic carbonyl sulphide molecules (135 pieces of information in all) have been usedz28 to calculate 19 potential constants up to the fourth order in OCS.For larger molecules unique quadratic fields cannot be computed but detailed studies of families of related molecules can lead to useful results. Schererzz9 has closely fitted 176 observed out-of-plane vibrational frequencies in a series of chlorinated benzenes with a 23-constant force field calculated ‘stereo’ views of the vibration and shown that his results can be used to predict the fre-quencies of similar vibrations in systems containing other types of substituents. O~erend~~O has continued his work on anharmonicity in polyatomic mole-222 W.Mueller-Litz and H. Hobert 2. phys. Cheno.(Leipzig) 1967 236 84. 223 L. M. Roev and A. V. Alekseev Elem. Fotoprotsessy Mol. Akad. Nauk S.S.S.R., 224 W. H. Smith H. C. Eckstrom and B. Faer J . Phys. Chem. 1968 72 369. 225 Wateru Suetaka Bull. Chem. SOC. Japan 1967 40 2077. 2z6 D. M. Smith and R. P. Eischens J . Phys. and Chem. Solids 1967 28 2135. 227 L. H. Jones R. R. Ryan and L. B. Asprey J . Chem. Phys. 1968 49 581. 228 Y. Morino and T. Nakagawa J . Mol. Spectroscopy 1968,26,496. 229 J. R. Scherer Spectrochitn. Acta 1968 24 A 747. 230 S. Reichman and J. Overend J . Chem. Phys. 1968 48 3095. 1966,346 106 D. A . Long cules; this factor has usually had to be neglected for lack of reliable informa-tion.It is interesting to note that calculations on the force field of Os04 s h o ~ ~ 3 ~ that experimental mean amplitudes of vibration determined with the methods presently available are mostly not accurate enough to use for the precise calculation of interaction force constants. However formulae for the calculation of mean amplitudes of vibration continue to appear. There also continues unabated a passion for converting vibrational fre-quencies into force constants whatever the approximations involved. It must be emphasised that such force constants usually constitute only one of many possible sets of numbers which when fed into some approximate potential energy formula reproduce a set of observed frequencies. Yet such force constants are all too frequently the subject of solemn interpretation.A crude type of calculation employing order of magnitude force constants which are transferred between related molecules has been found to help with assignments in inorganic molecules. that a computer programme automatically cures nonconvergence troubles in a force-constant refinement calculation intrigues this author and no doubt many others who have experienced this difficulty . The Stimulated Hyper and Inverse Raman Effects.-The very large electric field strength associated with the radiation produced by a giant pulse laser has led to the observation of novel spectroscopic phenomena. Woodbury and Ng233 discovered the stimulated Raman effect in 1962. In 1964 Jones and Stoicheff 234 reported observation of the inverse Raman effect (or Raman effect in absorption) and in 1965 Terhune and coworkers235 published the first hyper Raman spectra.None of these new forms of spectroscopy has received much attention from chemical spectroscopists and only the stimu-lated Raman effect has been the subject of any substantial amount of study. The substantial reviews of the stimulated Raman effect and related pheno-mena by Bloembergen and by Schuler have been referred to earlier; they obviate the need for discussion here. It seems worthwhile however to explain briefly the nature of the other two new effects and to indicate their potential for chemical studies. When a scattering medium is irradiated simultaneously with intense monochromatic light of frequency vo (as from a giant pulse laser) and with an intense continuum the molecules are stimulated to emit radiation at vo and at the same time to absorb radiation at vo + Vm or vo - Vm from the con-tinuum where Vm is the frequency associated with a transition between two energy states in the molecule.Absorption at vo + Vm is associated with promotion of the molecule to a higher energy level whereas absorption at The statement in an 231 A. Mueller B. Krebs and S. J. Cyvin Acta Chem. Scatid. 1967 21 2399. 232 L. Nemes and M. Kemenczy Acra Chim. Acad. Sci. Hung. 1967 53 359. 233 W. J. Woodbury and W. K. Ng Proc. Inst. Radio Eng. 1962 50 2367. 234 W. J. Jones and B. P. Stoicheff Phys. Rev. Letters 1964 13 657. 235 R. W. Terhune P. D. Maker and C. M. Savage Phys. Rev. Letters 1965 14 68 1 Part (ii) Infrared and Raman Spectroscopy 107 vo - Vm is associated with a transition to a lower energy level.The resulting induced absorption spectrum is in effect an inverse Raman spectrum. All the Raman-active transitions for the molecule should be observed in the inverse Raman effect unlike the stimulated Raman effect where only one or two characteristic frequencies are observed. Since a complete Raman spec-trum in absorption can in principle be produced in the duration of the giant pulse (ca. 30 x 10-9 sec.) the study of free radicals and short-lived species should be possible. This method of producing Raman spectra may also help to overcome the problems associated with fluorescence. This promising technique has hardly received attention. A key problem appears to be the production of the intense continuum.Vodar et ~ 1 . " ~ have experimented with a continuum produced by passing a giant pulse laser through a cell containing krypton at 10 atm. pressure and reported observation of several lines in the inverse Raman spectra of organic liquids like chloroform and benzene. Otherwise there have been no further publications in this field to date. The dipole pind induced in a molecule by an applied electric field E is given in general by the nonlinear relationship Pind = RE f &$E2 + iyE3 + . . . The scattered radiation known as the Raman effect has frequency depen-dence vo & vm and arises from the linear term in this series. The selection rules intensities and polarisation characteristics are determined by the pro-perties of the second-order polarisability tensor a.The term in E2 is only significant for very large fields. In such cases there will be additional scattered radiation determined by the properties of the third-order hyperpolarisability tensor p. This has a frequency dependence of the form 2v0 & vm and is referred to as the hyper Raman effect. The selection rules are quite different from those for the normal Raman effect. In particular fundamentals inactive in both the i.r. and Raman effects may be hyper Raman active. Examples are the torsional frequency (A,) in ethane ( D M ) and no less than six fre-quencies (2Blu + 2Bzu + 2Ezu) in benzene ( D 6 h ) . The selection rules have been considered in detail by Cyvin Rauch and D e c i ~ s . ~ ~ ~ The effect is very weak and the only experimental work recorded is that of Terhune and co-workers235 who reported hyper Raman spectra of water fused quartz CC14, and MeCN.The hyper Raman effect is a promising new method for the study of molecular vibrational states previously regarded as 'spectroscopically inaccessible' . General.-It has not been possible to allot any substantial space to some topics. However in what follows an attempt has been made to set recent work in some of these fields in perspective and to draw attention to novel developments. 236 S. Dumartin B. Oksengorn and B. Vodar Compt. rend. 1965 261 By 3767 4031 ; 237 S. J. Cyvin J. E. Rauch and J. C. Decius J. Chem. Phys. 1965 43 4083. S. Dumartin B. Oksengorn and B. Vodar J. Phys. (Paris) Colloq. 1967 28 1 108 D.A. Long There has been about the usual number of publications dealing with intensities in i.r. and Raman spectra and the calculation therefrom of bond dipoles and bond polarisabilities and their derivatives. Recent publications of the Minnesota school on COFZ COC12 and on liquid hexa-fluor~benzene~~~ and liquid methyl iodide are broadly representative of i.r. intensity studies which generally have proceeded along well-established lines. The work of Nagarajan and DurigZ4l on intensities in dicyanodiacetylene is illustrative of Raman intensity studies. The gas laser has made possible the direct and accurate measurement of absolute Raman intensities242 and there should be an expansion of interest in this field. One growth point might be the idhence of resonance effects on Raman intensities which are not explicable in terms of existing theories.The CNDO method has been ~ s e d ~ 4 ~ to calculate for a number of molecules dipole derivatives which are in good agreement with values obtained from i.r. intensity measurements. Raman spectroscopy of polymers has been slow to develop although the laser is such a good source for the study of such materials. Reports have been confined to polypropene fibres,244 polymethylene and hexa-gonal and orthorhombic p~lyoxymethylene.~~~ Environmental effects have formed the basis of many spectroscopic studies. Typical of investigations of the effect of high pressures is the work of the Toronto school on pressure-induced i.r. absorption in hydrogen.247 Studies of solvent effects have covered a very wide range of topics e.g.the induced absorption spectra of hydrogen dissolved in Group IV tetra-halides248 which provide evidence for rotational motion of the hydrogen in solution; the i.r. spectra of saturated solutions of HC1 in fused alkali-metal chlorides249 which show that rotation of the HCl is unhindered; the i.r. spectra of solutions of strong electrolytes in D2O which contain bands attributable to hindered rotation and hindered translation of the water molecules;250 the investigation by of Raman spectra of D2O in Hz0 which supports the two-state model of water structure. There have been several interesting general papers relating to Raman spectroscopy. Tang and A l b r e ~ h t ~ ~ ~ have developed a Ranian intensity 23t3 M. J. Hopper J. W.Russell and J. Overend J . Chem. Phys. 1968 48 3765. 239 T. Fuyiyama and B. Crawford jun. J. Phys. Chem. 1968 72 2174. 240 C. E. Favelukes A. A. Clifford and B. Crawford jun.,J. Phys. Cfieni. 1968,72,962. 241 G. Nagarajan and J. R. Durig Bull. SOC. roy. Sci. Lisge 1967 36 552. 242 J. G. Skinner and W. G. Nilsen J . Opt. SOC. Amer. 1968 58 I 1 3. 243 G. A. Segal and M. L. Klein J . Chem. Phys. 1967 47 4236. 2M P. J. Hendra and H. A. Willis Chem. and Znd. 1967,2146. 245 R. F. Schaufele and T. Shimanouchi J . Chem. Phys. 1967,47 3605. 246 G . Zerbi and P. J. Hendra J . Mol. Spectroscopy 1968 27 17. 247 H. L. Welsh Proc. 1st Tnternat. Conf. Spectroscopy Bombay 1967 2 340. 248 M. 0. Bulanin and M. G . Mel’nik Optika Spektroskopiya Akad. Naiik S.S.S.R., 249 L. M. Gurevich,V.N. Devyatkin and K. F. Zhitkov,Zhur.fiz. Khim. 1968,42,1701. 250 D. A. Draegert and D. Williams J. Chem. Phys. 1968,48 401. 251 G. E. Walrafen J . Cheni. Phys. 1968 48 244. 252 J. Tang and A. C. Albrecht J . Chem. Phys. 1968 49 1144. 1967 3 214 Part (ii) Infrared and Raman Spectroscopy 109 theory with the Kramers-Heisenberg dispersion theory as a basis. Theimer253 has discussed the Raman effect in a plasma. Raman scattering from atmos-pheric oxygen and nitrogen has been 0bserved~5~ with a pulsed nitrogen laser and used to measure concentrations as a function of range. Koningstein and co-workers255 have made a series of studies of the electronic Raman effect for rare-earth ions in various crystal environments. Some transitions are associated with an asymmetric scattering tensor. 253 0. Theimer Proc. 1st Internat. Conf. Spectroscopy Bombay 1967 2 516. 254 D. A. Leonard Nature 1967 216 142. 255 J. A. Koningstein and 0. S. Mortensen Nature 1968 217 445. 255 0. S. Mortensen and J. A. Koningstein J. Chem. Phys. 1968 48 3971; J. A. Koningstein and 0. S. Mortensen Phys. Rev. 1968,168 75; J. A. Koningstein and 0. S. Mortensen J. Opt. SOC. Amer. 1968 58 1208; J. A. Koningstein and 0. S. Mortensen, Chem. Phys. Letters 1968 1 693; J. A. Koningstein Phys. Rev. 1968 174 411

ISSN:0069-3022    

DOI:10.1039/GR9686500083

出版商:RSC    

年代:1968

数据来源: RSC

摘要:

5 Part (iii) MICROWAVE SPECTROSCOPY By J. E. Parkin (Department of Chemistry, University College, London) SEVERAL recent developments seem to point to microwave spectroscopy (or molecular rotational resonance as the subject may well become known) being on the threshold of a minor boom. The number of papers considered represents an increase of almost 50 % over those for the comparable period a year ag0.l Not only have more investigations of a traditional nature been reported (determinations of complete or partial molecular structures, barriers to internal rotation, dipole moments, and nuclear quadrupole coupling constants) but the technique has been applied to new kinds of problem. The introduction of commercial instruments with much greater ease of operation than those which have evolved in research laboratories will certainly be reflected in the kind of investigations reported in the next few years.The study of collision processes in gases by microwave spectroscopy has been reported by several authors, and the potentialities as an analytical tool have been pointed to and preliminary work carried out. Microwave Spectrometers.-The most important development of recent years which has become most marked in 1968 is the introduction of the first complete commercial spectrometers2 which incorporate some features absent from all but the most sophisticated research instruments. The backward- wave oscillator source has largely superseded the original klystron source over which it has the great advantages of a much wider frequency-range and tuning at a wide range of speeds.It is an interesting irony that microwave spectroscopy began as a very high-resolution technique and only after about 20 years has a low-resolution scanning instrument become available. The possibility of distinguishing the wood from the trees is of great benefit especially in the spectrum of relatively large molecules where the low- resolution spectrum is an extremely simple linear function of the single molecular parameter proportional to the mean of the reciprocals of the principal moments of inertia Ib and Ic. Since the three principal moments of inertia are sensitive functions of the geometry and atomic masses of the molecule, such a low-resolution spectrum can often distinguish between isotopic or conformational modifications.The spectra then correspond to the band spectra of other regions, but being a property of the molecule as a whole do not have the inherent structural information associated with func- J. E. Parkin, Ann. Reports, 1967, 64, 181. Hewlett-Packard Company, Palo Alto, California, U.S.A. ; Tracerlab Ltd., Wey- bridge, Surrey, England. 111112 J. I? Parkin tional groups in these spectra. Since the relative intensities of a given micro- wave transition, i.e. with the same upper- and lower-state quantum numbers, for different isotopic, vibrational or conformational species in equilibrium are proportional to population (along with certain other normally estimable quantities), this measurement enables relative concentrations or, via a Boltzmann factor, vibrational and conformational energy differences to be obtained.The intensity problem which has plagued microwave spectroscopy has largely been overcome by Harrington3* and is discussed later in more detail. The commercial instruments have been designed with this intensity problem in mind and incorporate many features beyond the scope of this article. These, together with the facility for direct-frequency measurement, rapid scanning and location at a particular frequency, although present individually in earlier instruments, have, as combined here, provided a tool of great versatility and ease of operation comparable to the larger instru- ments designed for other spectral regions. Microwave Spectroscopy in Routine Chemistry.-Since their application to molecular spectroscopy, microwave techniques have been recognised to have many advantages as a routine analytical tool.Dailey5 for instance reviewed such possibilities in 1949 when the inherent drawbacks were also recognised. It is significant that as late as 1966 and 1968 Lide6 and Millen7 make many of the same points. Actual practical application has come only in the last year or so and then only in a preliminary way. The microwave spectrum of a gaseous molecule at high resolution is a very sensitive and unique function of the inertial parameters and therefore of its detailed structure. As such, microwave spectroscopy can be used in a straightforward manner for the empirical characterisation of molecules without the need for detailed spectroscopic analysis. The chief limitations of the technique are firstly that the molecule must be polar, as absorption intensity is proportional to the square of the dipole moment, and secondly that it should have sufficient vapour pressure at not too elevated tempera- tures.In practice, this means dipole moments >0-1 Debye and vapour pressures of a few microns at temperatures up to 3 0 0 " ~ . The advantages as an analytical technique are that very small samples of the order of micro- moles can be investigated, it is non-destructive, and it is rapid and repro- ducible. High resolution is a feature of all microwave spectrometers and it is probable that of the order of half-a-dozen lines and their relative intensities would be sufficient for characterisation. In reality, most molecules have many thousands of such transitions so that the problem is one of documentation rather than of inherent impracticality.Unfortunately no collection of charac- teristic lines has been published since 1950.8 The use of commercial instru- H. W. Harrington, J. Chem. Phys., 1967, 46, 3698. H. W. Harrington, J. Chem. Phys., 1968, 49, 3023. B. P. Dailey, Analyt. Chem., 1949, 21, 541. D. R. Lide, jun., Ado. Analyt. Chem. Instrumen., 1966, 5 , 235. D. J. Millen, Chem. in Britain, 1968, 4, 202. P. Kisliuk and C. H. Townes, J. Res. Nut. Bur. Stand., 1950,44,611.Micro wave Spectroscopy 113 ments with their advantage of speed may render a similar but up-to-date compilation extremely valuable. Potentially the most valuable feature of the newer instruments is the possibility of their use in quantitative analysis of polar gas mixtures. If practical, the method has the advantage of being fast, precise, and unambi- guous.In order to have quantitative significance; intensity data must bear some simple reproducible relation to the concentration of the absorbing species. This condition is met in conventional spectroscopy by Beer’s law in which the absorption intensity is directly proportional to the path-length. Deviations from Beer’s law are well known but comparatively unimportant except in special cases. In microwave spectrometers however, with essentially monochromatic sources, such deviations occur as soon as the microwave power is increased beyond a relatively low level, when saturation of a transi- tion at resonance takes place. At still higher power levels, the absorption intensity decreases and there is a characteristic maximum in the intensity- power level curve. Power saturation has been a troublesome problem since the early days of the subject and spectrometers are operated normally at a level sufficiently low to reduce its effects.Under such conditions they lose much of their sensitivity, especially when the gas is being monitored at low concentrations. A second problem arises when pressures are increased beyond a certain point, typically 20 to 50 microns. The absorption intensity then no longer increases with pressure and the lines merely broaden. Harrington3p has overcome these problems by abandoning the Beer’s law intensity-coefficient, y, defined by AP,/Po = yL, where PO is the incident power level and APg is the loss in power for a waveguide cell of length L.He recognises that a new intensity coefficient I?, defined by AP,/Pd = rL, has the property that its maximum attainable value for a given line is directly proportional to the concentration of the absorbing species and, most importantly, is independent of the broadening collision time T. He further shows that this intensity coefficient I’ is directly proportional to the signal amplitude S for a con- ventional Stark modulated absorption cell.9 In many cases insufficient power is available to fully saturate a transition. Harrington overcomes this difficulty by recognising that the shape of the curve of log S against log PO, which he calls the intensity law, is a function only of the instrument except for a change of origin due to concentration of the absorbing species (proportional to a shift in the log S direction) and to the broadening relaxation time T (proportional to a shift in the log PO direc- tion).Experimental points, obtained well before saturation is maximised, can be translated to this theoretical curve by an amount proportional to the relative concentration of the absorbing species. Calibration can then be effected in a number of ways. This mode of operation applies directly to the case where a variation in molecular concentration is being monitored by a R. H. Hughes and E. B. Wilson, jun., Phys. Rev., 1947, 71, 562.114 J . E. Parkin single transition. In measuring the relative intensities of lines separated in frequency, account must be taken of the instrumental variation in power loss with frequency due to the waveguide itself.This is a common difficulty in microwave investigations where the relative intensities of vibrational satellites are of interest in providing their vibrational energy separation. Harrington shows how large errors can result from neglect of this function, especially when the Stark absorption-cell has not been carefully designed to reduce or eliminate reflections and other frequency-dependent losses. In the first paperlo to appear where the technique has been applied to an analytical problem are demonstrated many of the possibilities of the method. Ways of tackling the intensity measurement problem are investigated by tracing the linearity of three different intensity measures with the pressure of the absorbing gas.These measures were firstly the maximum signal obtained for the transition being monitored (the method suggested by Harrington4), secondly the signal at a relatively low but constant power level insufficient to fully saturate the transition, and thirdly the total peak area under these constant-power-level conditions. They found that all the signals were indeed linear over pressure ranges from 0-ca. 25 microns but at higher total pres- sures the curves become appreciably non-linear. The third method gave linearity over an appreciably longer range of pressure. For instance a linear signal was obtained for a 5 % mixture of acetone in nitrogen up to 50 microns total pressure. Despite these deviations at higher pressures, the curves are quite reproducible and are ideally suited to the analysis of certain kinds of gas mixtures.The example given indicates the analysis to be quite as precise as that obtained mass spectrometrically. Vibrational and Conformational Energy Differences.-The measurement of relative intensities is hindered by several factors, some of which have already been alluded to. Frequency-dependent power losses in the waveguide cell render it imperative that a calibration curve be obtained for each Stark absorption-cell, i.e. a cell background-intensity function. Even when these experimental corrections are made, relative intensities are still not directly comparable with the populations and therefore, from the Boltzmann factors, with the energy differences.Rotational line intensities, y, can be shown (see e.g. ref. 11, p. 343) to be given by : 8x2N f 3ckTAv Y = ___ 1pijp vo2 where the constants c, k, and T have their usual meaning, N is the number of molecules per ~ m . ~ in the absorption cell, f is the population factor, vo is the resonant frequency of absorption, and Av the halfwidth. l p j 1 2 is the square of the dipole-moment matrix element for the transition summed over three lo J. T. Funkhouser, S. Armstrong, and H. W. Harrington, Analyt. Chem., 1968,40,22. l1 C. H. Townes and A. W. Schawlow, 'Microwave Spectroscopy,' McGraw-Hill, New York, 1955.Micro wave Spectroscopy 115 perpendicular directions in space, and is a function of the rotational quantum numbers, the molecular asymmetry parameter, and the square of the total dipole moment.In comparing the intensities of lines of two species in equi- librium, in order to determine the relative populations, these factors must be known. Assuming the same rotational quantum numbers for both lines, the additional factors will include the ratios of the dipole moments and transition frequencies squared, the inverse ratio of the linewidths, and a very compli- cated function of the molecular asymmetry parameters. In the case of dif- ferent vibrational species these ratios are normally sufficiently close to unity to be neglected, and relative populations accurate to a few % are easily obtained. In practice, vibration rotation interaction is sufficient to change the inertial parameters of a molecule, in excited vibrational states, to render the satellite lines easily resolvable in the microwave region, as opposed to the i.r.where sequence structure is often quite unanalysable. This comparison is especially favourable for the microwave method when the molecule has many low-lying vibrational levels arising from hindered torsions or ring-puckering vibrations. Accurate intensity-data for a number of these vibrational satellites gives values for the energy levels with sufficient precision to give quite good molecular potentials. Harrington4 considers this problem in some detail with special reference to the low-lying excited states of the out-of-plane bending vibration of trimethylene sulphide. In an earlier paper12 he obtained the vibrational energies of the first four excited-levels as 0, 141, 156, and 241 cm.-l, with a precision of some 3 cm.-l.There is every reason to suppose that this precision might be improved with more advanced equipment. When the different species are conformational isomers intensity compari- sons are not quite so favourable. The nature of a conformational change is often sufficient to modify greatly such factors as the molecular symmetry and the dipole moment, especially the projection of the latter along a given symmetry axis. This is well illustrated by a preliminary communication13 which reports studies of the two conformational isomers of piperidine, CSHXN, having axial and equatorial imino-hydrogen atoms. The relative intensity data, using admittedly less sophisticated equipment than was employed in the previous example, pointed to a population ratio for the axial and equatorial conformers of 2 : 3 yet to a dipole-moment ratio (actu- ally its pa component) in the ratio 3.2 : 1.Since the latter appears squared in the intensity expression, this gives an observed intensity ratio of some 6.5 : 1 in favour of the less-abundant species. The energy difference between the two isomers then corresponds to a separation of 245 It 150 crn.-l, the equatorial conformer being the more abundant. Collision Processes in Gases.-Microwave transitions are particularly sensitive to the phenomenon of pressure broadening because, as opposed to l2 H. W. Harrington, J . Cliem. Phys., 1966, 44, 3481. l3 P. J. Buckley, C . C. Costain, and J. E. Parkin, Clzem. Cumtn., 1968, 668.116 J. E.Parkirt other spectroscopic techniques, the effective slitwidth is very much smaller than the linewidth. The broadening problem was treated in classical terms14 and later modified15 to give the intensity profile for a line in terms of the variables of equation (l), and the mean collision time T. This standard expres- sion (see ref. 11, p. 342) is the starting point for most treatments of collision- broadening data. A consequence of the theory, borne out in practice, is that for pressures between ca. 0.5 mm. and 10 cm. the peak intensity of a line remains constant but the halfwidth increases linearly with pressure. The broadening parameter can be interpeted in terms of T and hence effective collision-diameters may be determined. These values are considerably larger than those obtained from kinetic theory and it is their interpretation in terms of long- and short-range molecular-interaction potentials which forms the basis of the interest in this kind of study.Several groups have reported work recently on these lines, and it will certainly gain impetus when commercial spectrometers become more readily available. The most fruitful approach so far to the interpretation of pressure-broaden- ing data is based on the work of Anderson16 as modified later by Tsao and Curnutte.17 Interactions considered in their theory include dipole-dipole, dipole-quadrupole, quadrupole-quadrupole, and dispersion forces, where the emitting species is self-broadened or broadened by a foreign species. It is thus one of the few methods of obtaining molecular quadrupole moments, although in practice their evaluation is critically dependent on the validity of the theory.The different methods of determining quadrupole moments, including the microwave method have been reviewed.18 In a series of paper@ a further refinement of the theory is presented in which some of Anderson’s approximations are removed, although this results in very unwieldy expressions. The usefulness of the treatment was demon- strated by application to earlier data on the self-broadening of CH3C1, CH3F, CHF3, and PF3. For these systems the interaction potentials are entirely dominated by the dipole-dipole terms and since these quantities are known experimentally very accurately, the theoretical broadening parameter can be calculated and compared directly with the experimental value.The compari- son was extremely good being within the limits of experimental error, and the new theory appears to represent a great improvement on Anderson’s theory. The treatment is applied to data for the more general cases of line broadening of a number of linear and symmetric rotors by a collision with a range of linear molecules and inert gas atoms. In such cases shorter-range forces such as dipole-qudrupole and quadrupole-quadrupole interactions become important and quadrupole moments may be estimated from the l4 H. A. Lorentz, Proc. Amst. Akad. Sci., 1906, 8, 591. l5 J. H. Van Vleck and V. F. Weisskopf, Rev. Mod. Phys., 1945, 17, 227. l6 P. W. Anderson, Phys. Rec., 1949, 76, 647. 17 C. J. Tsao and B. Curnutte, J . Quant.Spectroscopy Radicitiue Transfer, 1962, 2, 41. l8 Krishnaji and V. Prakash, Reo. Mod. Phys., 1966, 38, 690. l 9 J. S . Murphy and J. E. Boggs, J . Chem. Plzys., 1967, 47, 691, 4142; 1968,49, 3333.Microwave Spectroscopy 117 data. In general, the values obtained are an improvement on previous deter- rninations from microwave data and compare favourably with values ob- tained by other experimental methods, e.g. the value for the N2O quadrupole moment is 6.07 & 0.31 Debye A, compared with values of 6.0 obtained from induced birefringence measurements, 8.50 from nuclear spin relaxation, 5.56 from the second virial coefficient, and theoretical calculations in the region 7-40-7-50. Although this is the most favourable comparison, the agreement for other molecules is quite good and is in general inside the spread of other values.New pressure broadening data have been obtained20 for OCS, CHF3, and N2O in collision with a number of foreign gases by a novel experimental technique. The data are interpreted on the basis of Anderson’s theory and it is concluded, not surprisingly, that it is inadequate and in particular tends to overestimate the dipole-dipole contributions. Further data have been obtained21 for OCS broadened by 0 2 and Nz. A more thorough treatment of this improved data, using the theory of ref. 19, is awaited with interest. The conclusion to be drawn from this work is that although microwave pressure-broadening studies will be a valuable source of molecular quadru- pole moments, much critical work, both theoretical and experimental, needs to be done.Concerned with a different aspect of the collision problem, Oka22 has continued the work reported last year1 on the direct study of collision- induced transitions. He uses a high-power double-resonance technique to pump molecules into a non-Boltzmann distribution and observes changes in the absorption intensity of lines involving quite specific levels other than those being pumped. This change in intensity can only result from preferred collision-induced transitions which appear to obey quite definite selection rules of the dipole or quadrupole type. For instance collisions of ammonia molecules with rare-gas atoms and other ammonia molecules appear to give quite different types of selection rule. Some of the transitions induced by the rare-gas collisions seem to be caused by the octopole moment of ammonia.A large amount of data accumulated in this work points, as did much of the pressure-broadening data, to the inadequacy of current theories of molecular collision phenomena, Double Resonance.-In a 1967 paper received after last year’s report was written, Ronn and Lide23 describe the first successful i.r.-microwave double- resonance experiment. Radiation from a carbon dioxide laser delivering 50 watts in the 10.6 micron region was passed through a conventional Stark- modulated absorption cell containing methyl bromide. The P(20) laser line at 944.1 8 cm.-l was strongly absorbed, most probably by the methyl bromide transition from J,K = 9,l in the ground-state to J,K = 8,0 in the vibration- ally excited-state. They observed simultaneously that microwave lines involv- 2o B.Th. Berendts and A. Dymanus, J . Chem. Phys., 1968, 48, 1361; 1968, 49, 2632. 21 K . Srivastava and S. L. Srivastava, J . Chem. Phys., 1967,47, 1885. 22 T. Oka, J . Chem. Phys., 1967, 47, 4852; 1968,48,4919; 1968,49, 3135. 23 A. M. Ronn and D. R. Lide, jun., J . Chem. Phys.; 1967, 47, 3669.118 J. E. Parkin ing the J = 1 and 2, K = 1 levels in the ground-state decreased in intensity whereas the corresponding K = 0 levels were unaffected. This indicated that the disturbance to the equilibrium Boltzmann distribution was transmitted from J,K = 9,l over a wide range of J values but not to levels of different K. This behaviour is in accord with Oka’s findings on similar systems (see refs.in ref. 22). They were not able to identify increased intensity in transitions of the vibrationally excited state due to the overlapping of stronger lines. If this technique can be extended, possibility of the detailed study of simultaneous vibrational and rotational energy transfer will be opened up. The limitation appears to be that the coincidence of the laser line with the i.r. transition must be almost exact. Tunable lasers are becoming available and there seems to be some hope for new developments in the near future. Microwave-microwave double resonance has now become a standard technique as predicted in last year’s report. F l ~ n n ~ ~ has given a useful review of several quantitative aspects of double resonance spectroscopy including the line shape to be expected.There is no doubt that in analysing congested spectra of large molecules or mixtures of molecules double-resonance tech- niques have considerable advantages, although in general their sensitivity is less than in the conventional spectrometer. Vibration-Rotation Interactions.-Several rather thorough investigations of the ground-states of molecules have been reported recently in which i.r. and microwave data are combined to provide the best molecular force-field. As is well known (ref. 11, p. 105) rotational energy can be expressed as the sum of the rigid-rotor energy and that due to centrifugal distortion. The latter is to first order a function of five inde~endent~~ distortion constants, (four in the case of a planar molecule) which are in turn linearly related to the elements of the inverse force-constant matrix.They can thus be combined with i.r. frequencies to provide a consistent set of molecular force-constants. As obtained from microwave work, distortion constants often have precision sufficient to rank with weights comparable with the frequency data and can render determinable a previously underdetermined problem. A typical recent example is a study of the NSF molecule26 in which earlier data27 are extended to higher quantum numbers to give better determination of the distortion constants. The general force-field contains six harmonic. potential constants f1 andfi associated with S-N and S-F bond stretching, fa the bending constant, and three interaction constants f12, fia, and fia Using the i.r.or microwave data alone, only the three diagonal force-constants could be determined. Combination of the data allowedfia also to be deter- mined, the two remaining constants being insignificantly small. A more interesting way of considering the data is to examine how far the distortion constants could be used to fix the vibrational assignments. Of six possibilities only one is consistent with the microwave data and must be preferred even 24 G. W. Flynn, J. Mol. Spectroscopy, 1968, 28, 1. 25 J. K. G. Watson, J. Chem. Phys., 1966, 45, 1360; 1967, 46, 1935. 26 R. L. Cook and W. H. Kirchhoff, J. Chem. Phys., 1967, 47, 4521. 27 W. H. Kirchhoff and E. B. Wilson, jun., J. Amer. Chem. SOC., 1963, 85, 1726.Microwave Spectroscopy 119 though this results in an abnormally low value for fz of 2.85 mdyne/& nor- mally between 4 and 5 in other molecules containing the S-F bond.Other Investigations.-As mentioned in the introduction, considerably more structure determinations have been reported this year than last. Rather than attempt to document these, a few interesting examples chosen more or less at random will be summarized. Some years ago Costain and SrivastavaZ8 investigated the microwave spectra of gas-phase hydrogen-bonded dimers of trifluoroacetic acid with acetic, monofluoracetic, and formic acids. The spectra consisted of broad evenly spaced bands assigned to successive transitions of the type J + 1 t J . No K structure was observed but some structural information could be inferred. The spectrum of the formic acid dimer has been rein~estigated~~ by measuring the susceptibility dispersion curve for the J = 8 +- 7 transition, and splitting attributable to K structure has been found.The reason why this structure should be apparent in the one technique and not the other is not clear, but a value of 20 D for the dipole moment is obtained from the data and the inertial parameters are not unreasonable. It is significant that these two are the only microwave investigations on hydrogen-bonded dimers reported in the literature although other systems are well known from i.r. studies. It is possible that the broad-band spectra expected for these species have been missed by use of high-resolution instruments ; low-resolution scanning instmments may well alter this situation.The threefold barrier to internal rotation for methyl alcohol changes from 375.6, to 371.8 and 370.3 for the isotopic species CH30H, CD30H, and CH30D. Lees and Baker30 have reported a thorough reinvestigation of the spectrum of this molecule in the millimeter wave region. They were unable to separate the sixfold from the threefold term in the barrier potential, the two moments of inertia about the near-symmetry axis, and two interaction para- meters; they cast some doubt on determinations of this term in other mole- cules. They also obtain a complete geometry for the molecule from the inertial parameters, with significant increase in precision over earlier deter- minations. The bond lengths are YCH 1.094 & 0-003, YOH 0.945 & 0.03, and TCO 1.425 rf 0.002 A, and the methyl tilt is 3’16’ & 11’.The spectrum of nitrosomethane has been investigated31 and a value of 1137 cm.-l found for the barrier to internal rotation. The dipole moment is 2.300 D but the struc- tural parameters are not well determined. Structure determinations have been reported32* 33 of the related compounds SeF4 and SeOFz. The former has an unsymmetrical structure, as implied by the observation of a microwave spectrum, which can be thought of as a 28 C. C. Costain and G. P. Srivastava, J. Chem. Phys., 1964, 41, 1620. 29 G. P. Srivastava and M. L. Goyal, Phys. Rev., 1968, 168, 104. 30 R. M. Lees and J. G. Baker, J. Chem. Phys., 1968.48, 5299. 31 D. Coffey, jun., C. 0. Britt, and J. E. Boggs, J . Chem. Phys., 1968, 49, 591. 32 I. C. Bowater, R. D.Brown, and F. R . Burden, J. Mol. Spectroscopy, 1968,28,454. 33 I. C. Bowater, R . D. Brown, and F. R . Burden, J. Mol. Spectroscopy, 1967, 23,272; 1968, 28, 461.120 J. E. Parkin distorted octahedral structure with selenium lone-pair orbitals occupying two of the positions. The axial and equatorial Se-F bonds have lengths 1.771 and 1.638 A, and are at angles of 169.20 and 100.55" respectively. The dipole moment is 1.779 D. The structure is very similar to that reported34 for SF4. %OF2 has the structural parameters rseo 1.576, rseF 1.730 A, F-Se-F 92*22", and 0-Se-F 104.82". The dipole moment is 2-84 D. The molecules formally have sp3d hybridisation and this data will give insight into the electronic properties of selenium in this ~tate.~5 A spectrum attributed to cyclopropanone has been reported.36 The ob- served moments of inertia are in agreement with a structure with the para- meters, rc(l)o 1-18, rc(qc(2) 1-49, rc(z)c(3) 1.58, YCH 1.085 A, C(3)-C(l)-C(2) 64 and H-C-H 117'35'.The molecule has CZ, symmetry and the only non-zero dipole moment component pa is 2.67 & 0.1 D. These data rule out alternative structures such as allene oxide and the oxyallyl radical, in spite of the indication of Huckel calculations that cyclopropanone should be unstable with respect to the latter. Some of the structural parameters of cyclopropyl bromide have been determined37 as well as the quadrupole coupling constants, which indicate 22% ionic character for the C-Br bond, although this is probably an overestimate. The similarity of the coupling constants with those of vinyl bromide, perhaps indicating a similar C-Br bond, is noted. The ground and low-lying excited vibrational states of the ring-puckering vibration of cyclobutanone and methylene cyclobutane have been investi- gated.3*y39 Barrier heights obtained for the motion are 7.6 and 160 cm.-l respectively indicating considerably more preference for the non-planar structure in the latter compound. The partial structures of cyclohexene have been determined40 and the half-chair conformation and CZ symmetry con- firmed. The dipole moment is 0.331 D. The spectra have been investigated41 of some 1 -halogenoadamantanes, C~OHISX, probably the largest compounds yet to be studied by microwave spectroscopy. The assumption that all the angles are tetrahedral and all the C-H bond-lengths 1.09A leads to the reasonable C-C bond-length of 1.541 & 0.001 A for the chloro-compound. In view of the highly symmetrical cage-structure of adamantane, these assumptions are quite reasonable. The carbon-halogen bond-lengths are 1.370, 1.790, and 1-947 8, for C-F, C-CI, and C-Br, and are in accord with the analogous methyl, ethyl and t-butyl bond-lengths. 34 W. M. Tolles and W. D. Gwinn, J . Chem. Phys., 1962, 36, 1 1 19. 35 R. D. Brown and J. B. Peel, Ausfral. J . Chern., 1968, 21, pp. 2589, 2605, 2617. 3% J. M. Pochan, 3. E. Baldwin, and W. H. Flygare, J . Amer. Chem. SOC., 1968,90, 1072. 37 F. M. K . Lam and B. P. Dailey, J . Chem. Phys., 1968, 49, 1588. 38 L. H. Scharpen and V. W. Laurie, J . Chem. Phys., 1968,49, 221. 39 L. H. Scharpen and V. W. Laurie, J. Chem. Phys., 1968,49, 3041. 40 L. H. Scharpen, J. E. Wallrab, and D. P. Ames, J . Chem. Phys., 1968,49, 2368. 41 D. Chadwick, A. C. Legon, and D. J. Millen, J . Chem. Sot. ( A ) , 1968,1116.

ISSN:0069-3022    

DOI:10.1039/GR9686500111

出版商:RSC    

年代:1968

数据来源: RSC

摘要:

6 CATALYSIS BY METALS By G. C. Bond (Johnson Matthey arid Co., Ltd., Exhibition Grounds, Wembley, Middlesex) THE literature on catalysis has been enriched by the publication during 1968 of books by Sir Eric Rideall and by Thomson and Webb,2 and of one with a strong practical bias edited by Ander~on.~ The regular publication of Aduances in Catalysis4 and of Catalysis Reuiews5 continues. The Fourth International Congress on Catalysis was held in Moscow and Novosibirsk; plans for publication of the papers are uncertain, but it seems likely that only the Rus- sian versions will appear. The three areas which have commanded most attention during the past year are (i) chemisorption of gases on clean metal surfaces, (ii) the structure of supported metal catalysts, and (iii) the mechanism of hydrogenation of unsaturated molecules and related processes.This year’s Report will therefore be devoted exclusively to a review of these three subjects. Chemisorption of Gases on Clean Metal Surfaces.-The study of the chemi- sorption of simple gas molecules on rigorously cleaned metal surfaces con- tinues to attract widespread attention more particularly from physicists than chemists : the increasing complexity and sophistication of the necessary apparatus perhaps makes this inevitable. The process is of great fundamental interest in its own right and knowledge acquired finds application or generates ideas in fields as separated as space research and catalysis, although only its relevance to the latter subject will concern this Report. Clean metal surfaces are highly reactive, and rapidly chemisorb at least a monolayer of oxygen in the presence of this gas: even gold now appears to have some attraction for atomic oxygen.6 To study the reaction of other less strongly adsorbing gases requires the metal surface to be cleaned and then maintained in ultra-high vacuum (u.h.v.) (less than 10-10 ton) for the duration of the experiment.Among the cleaning methods which continue to be applied are thermal desorption, ion bombardment (which severely disturbs the sur- face and has to be followed by annealing’), and thermal sublimation of ‘Concepts in Catalysis,’ E. K. Rideal, Academic Press, London and New York, 1968. ‘Heterogeneous Catalysis,’ S. J. Thomson and G. Webb, Oliver and Boyd, Edinburgh, 3 ‘Experimental Methods in Catalytic Research,’ ed.R. B. Anderson, Academic Press, ‘Advances in Catalysis and Related Subjects,’ eds. D. D. Eley, H. Pines, and P. B. 5 ‘Catalysis Reviews,’ eds. H. Heinemann, Marcel Dekker, New York, 1968, vol. 2(1). 6 R. R. Ford and J. Pritchard, Chem. Comm., 1968, 362; I. G. Murgulescu and M. 1. 7 R. L. Park and H. H. Madden, Surface Sci., 1968, 11, 188. 1968. London and New York. 1968. Weisz, Academic Press, London and New York, 1968, vol. 18. Vass, Rev. Roumaine Chim., 1968, 13, 373. 121122 G. C. Bond surface layers. It is of course important to have available a means of knowjng when a fully cleaned surface is obtained, and low energy electron diffraction (LEED) has proved very useful for this purpose. The problem of keeping the surface clean is not simple, especially since much of this work is now done with quite small areas (e.g.of single crystals), and high surface coverage by the background gas results in a matter of minutes rather than hours after cleaning is stopped even at 10-10 torr. Some cleaning procedures paradoxically have the effect of making the surface dirtier, as for example when they cause dissolved impurities such as carbon to migrate to the surface.* Metal films formed by condensing metal atoms, which have evaporated from an electrically heated wire, on to a cold surface continue to represent the easiest way of obtaining clean surfaces, especially since the films act as getters for residual gas during their formation. Such films when condensed onto glass are polycrystalline and expose no particular crystal plane, but nickel films condensed in u.h.v.on to sodium chloride preferentially expose the (100) plane: enthalpy and entropy of adsorption were obtained for xenon both on this film and on nickel-Pyrex as a function of coverage, and by use of a patch model of a heterogeneous surface it was concluded that no (100) planes were exposed in nickel-Pyrex.9 The chemisorption of hydrogen on a metal surface is formally the most simple conceivable process in this area of study, but for a number of reasons the results exceed expectation in complexity. Almost every study reveals evidence for two or three different states of binding of hydrogen atoms: these have different energies and the relative population of each state can vary with surface coverage.There are also some severe practical difficulties in working with films, not the least of which is to secure uniform contact between the incoming gas and the film, especially with nonsintered, porous films. These points are exemplified in two recent studies of hydrogen on molybdenum10 and on nickel11 films. The sticking probability s of hydrogen on molybdenum films was determined at 78, 195, and 300"~: at the highest temperature the adsorbed &oms are mobile and s on a sintered film was 0-68, falling linearly with increasing coverage. However at 7 8 " ~ s fell rapidly from 0.73 to 0.40, providing evidence for a precursor state.1° The measurement of surface potentials by the static-capacitor method enables transient changes during adsorption to be followed.Hydrogen adsorbs on nickel at 9 0 " ~ forming atoms which are strongly bound and show a negative surface-potential: this is the so-called p-species which is preferentially adsorbed on one group of sites. On unsintered films, later doses are adsorbed partly on the remaining less-active sites and partly as molecules (cry-species) on top of atoms. Sin- tering forms a plane which will adsorb hydrogen in the cry-form but sur- prisingly not in the p-form.ll The adsorption of oxygen and of carbon * E. M. A. Willhoft, Trans. Faraday SOC., 1968, 64, 1925. 9 B. G. Baker and L. A. Bruce, Trans. Faraday SOC., 1968, 64, 2533. 10 D. 0. Hayward and N. Taylor, Trans. Faraday SOC., 1968, 64, 1904. l1 T. A. Delchar and F. C. Tompkins, Trans. Faraday SOC., 1968, 64, 1915.Cntalysis by Metals 123 monoxide has been examined at 7 7 " ~ ; ~ ~ the change of electrical resistance of a platinum film is linear with the number of molecules adsorbed, implying only one form which is thought to be the linear one.l3 There has been much concern in the last year or so concerning the surface structure of finely divided alloy powders, particularly those of nickel-copper.The occurrence of phase separation has been demonstrated, but the point of chief concern is whether the catalytic properties of the surface layer are those of an alloy, each atom having the same behaviour, or are the sum of the properties of the two different kinds of atoms acting independently of each other. A study of the low-temperature adsorption of hydrogen on granular nickel and copper and their alloys has been performed;14 the adsorption is activated on nickel but curiously not on copper, while the independence of the amount adsorbed on bulk composition between 10 and 80 % copper argues that phase separation occurred.One of the most disturbing uncertainties in the early days of the application of LEED and field emission or field-ion microscopy concerned the chemical nature of the surface species responsible for the visual effects. Farnsworth has pioneered the use of mass spectrometry with LEED, whereby the adsorbed species are characterised by their mass spectrum after flash desorption. This combination has been recently applied to confirm earlicr work on the adsorption of carbon monoxide on the (100) plane of nickel.15 Even more elegantly the combination of a time-of-flight mass spectrometry with field-ion microscopy permits the identification of individual surface atoms or species.16 The Structure of Supported Metal Catalysts.-Many of the powerful physical methods for examining surface species mentioned above, and others not referred to, are applicable only to massive metal specimens and not to finely divided metal particles such as are present in supported metal catalysts. Thus, although the study of chemisorption on well-cleaned metal surfaces has been of value in suggesting possible or likely surface structures, its relevance to catalytic mechanisms has not been direct because of (i) the very different natures of the solids and of the purities of their environments and (ii) the fact that (with rare exceptions17) reacting systems have not been studied. A refreshing feature of the scene over the last two or three years has been the increased attention given to the study of practical supported cata- lysts, a study which in the Reporter's opinion has been too long neglected: undoubtedly the accurate description of a porous two-component supported catalyst is altogether more difficult than describing, for example a metal film, but the widespread industrial use of supported catalysts is gradually pricking the conscieiices of academic workers, with results which will now be re- viewed.l2 E. F. W. Thurston, Trans. Furuday SOC., 1968, 64, 2181. 13 T. Sugita, S. Ebisawa, and K. Kawasaki, Surface Sci., 1968, 11, 159. l4 D. A. Cadenhead and N. J. Wagner, J.Phys. Chem., 1968,72,2775. 15 M. Onchi and H. E. Farnsworth, Surface Sci., 1968,11, 203; see also D. Lichtman, F. M. Simon, and R. R. Kirst, ibid., 325. E. W. Miiller, J. A. Panitz, and S. B. McLane, Re!:. Sci. Itutr., 1968,39, 83. 17 R. F. Baddour, M. Modell, and U. K. Heusser, J . Pltys. Cltern., 1968, 72, 3621. E124 G . C. Borrd Experimental studies have had two main objectives; (i) an adequate description of the structure of catalysts, particularly of the mean size or size distribution of the metal particles by means of gas chemisorption, electron microscopy, X-ray diffraction and other methods ; and (ii) to establish whether the specific activity of metal crystallites is a function of their size, i.e. whether sites of different properties exist in concentrations which are dependent on the size of the particle.The above-mentioned methods of deriving information about the metal particle-size distribution are now almost routinely employed. The use of selective gas-chemisorption (usually either hydrogen or carbon monoxide) can, at best, however, give only a mean value of crystallite size, but the sim- plicity of the procedure makes it attractive as a routine means of quality control. The adsorption of hydrogen on Pt-AI203 has been measured as a function of temperature (250-450"~) and pressure (4 x 10-2-120 torr): saturation coverage is obtained at each temperature at 60-70 torr, but the ratio of number of hydrogen atoms adsorbed to total number of platinum atoms falls with increasing temperature as follows: 0.98 at 250°c, 0.71 at 350°c, and 0.57 at 45O"c.l8 A simple and rapid chromatographic method for determining the volume of hydrogen adsorbed by catalysts has been described,lg and differential thermal analysis has been applied to this prob- lem.20 Of course, the application of this method depends, for its success, on the gas not adsorbing on the support or migrating from the metal to the support during the measurement.Therz is a growing body of evidence to show that at sufficiently high temperatures such migration can occur. On heating in hydrogen Pd/A1203 or Pt/Alz03, to which has been added a small amount of ferric nitrate, an e.s.r. signal having g = 2.10 develops and this is attributed to metallic iron: it is suggested that the ferric oxide (formed by calcining the nitrate) is reduced by hydrogen atoms which have migrated from the metal particles.Au/A1203 is ineffective in producing an e.s.r. signal, which incidentally is not seen when iron is absent.21 Several other physical methods have been applied to the study of supported- metal catalysts in the past year, some for the first time. Nickel-chromia catalysts have been examined by using electrical conductivity and contact potential difference : the conductivity is metallic if the catalyst contains more than 96% nickel but it is semiconducting if the nickel content is lower. The contact-potential difference is a maximum at 92-93 nickel regardless of atmosphere.22 Strictly speaking, catalysts of these compositions should be termed 'promoted' rather than supported.Calcium ions in CaY zeolite have been ion-exchanged with Pt(NH3)42+ ions and the resulting material has l8 D. Cismaru and A. Fruma, Rev. Rournaine Chim., 1968, 13, 139, 679. l9 F. Figureas Rosa, L. de Mourgues, and Y. Trambonze, J . Gas Chromatog., 1968, 2O V. Zapletal, K. Kolomaznik, J. Soukup, and V. Ruzicka, Chem. listy, 1968,62, 210. 21 K. M. Sancier and S. H. Inami, J . Catalysis, 1968, 11, 135. 22 D. Tarina, E. Weissman, and D. Barb, J . Catalysis, 1968, 11, 348. 6, 161.Catalysis by Metals 125 been examined by the X-ray absorption-edge method.23 After reduction all the platinum is zero-valent, but 60% of the metal is soluble in cold concen- trated HCl (particle size about lOA) while X-ray diffraction shows the re- mainder to be ca. 60 8, in size. Exposure of the catalyst to hydrogen at 300"c produces a larger change in the absorption edge than does exposure at 100°c, due, it is suggested, to the formation of stronger bonds.Hydrogen adsorbed on Pt-SiO2 or Pt-Al203 gives two i.r. bands, at 2040 and 2110-2120 cm.-l; the intensity of the latter band is much enhanced if the catalyst is pretreated with oxygen, and it is proposed that the band should be assigned to weakly adsorbed hydrogen on areas of platinum oxide.24 One of the most promising of the newer techniques to be applied to the study of adsorption and cata- lysts is Mossbauer spectroscopy;25 its application to examining the prepara- tion of a supported gold catalyst has been described.26 It has sometimes been reported in the past that the ease of reduction of a supported oxide depends on the nature of its support; a detailed study of the reduction of nickel oxide on various supports has now a~peared.~7 Rates were measured manometrically in a closed circulating system; rates of reduc- tion at 400"c decreased in the sequence: SiO2 > Si02-Al203 > A1z03.With nickel oxide on silica, increasing temperature of treatment (300-7OO"c) before reduction decreased the rate of reduction at 400"c, and co-precipitated catalysts were shown to be much more difficult to reduce than impregnated catalysts. A Pt-SiOz catalyst prepared by ion exchange of silica with [Pt(NH3)4](0H)2 and reduced in hydrogen at 500"c has the metal in the form of about 15 A particles as determined by gas adsorption (5-30 A by electron microscopy): this is considerably smaller than the particles in a catalyst prepared by impregnation with H ~ P ~ C ~ C .~ ~ It has again been reported that 'extractable' platinum in Pt-Al203 catalysts shows activity specifically for dehydrocyclisation of straight chain alkanes but not for their dehydro- genat ion .29 Little has appeared during 1968 on the vexed question of whether specific activity (i.e. activity per unit surface area) is a function of crystallite size, although several papers presented at the Fourth International Congress on Catalysis bore on this point.30 The general consensus appears to be that specific activity is remarkably constant over a large range of average particle sizes for a number of different processes, with perhaps the exception of the hydrogznolysis of ethane.31 In an important papeF Boudart and his associates have shown that selectivity for isomerisation of neopentane to isopentane 23 P.H. Lewis, J. Catalysis, 1968, 11, 162. 24 D. D. Eley, D. M. Moran, and C. H. Rochester, Trans. Faraday Soc., 1968,64,2168. 25 W. N. Delgass and M. Boudart, Catalysis Rev., 1968, 2, 129. 26 W. N. Delgass, M. Boudart, and G. Parravano, J . P h p . Chern., 1968, 72, 3563. 27 V. C. F. Holm and A. Clark, J . Catalysis, 1968, 11, 305. 28 H. A. Benesi, R. M. Curtis, and H. P. Studer, J . Catalysis, 1968,10, 328. 29 N. R. Bursian, S. B. Kogan, and Z. A. Davydova, Kinetika i Kataliz, 1968, 9, 661. 30 Anon., Platinum Metals Review, 1968, 12, 136. 31 D. J. C. Yates and J. H. Sinfelt, J . Catalysis, 1967, 8, 348. 32 M. Boudart, A.W. Aldag, L. D. Ptak, and J. E. Benson, J . Catalysis, 1968, 11, 35.126 G. C. Bond (the other reaction being hydrogenolysis to isobutane and methane) on a number of platinum catalysts at 307 O c decreases significantly with increasing dispersion of the platinum, but this is chiefly because activity for hydro- genolysis varies by 300 while activity for isomerisation varies only by fifteen. Another vexed problem relating to supported-metal catalysts has been at least partly resolved in 1968. It is of course well known that the activity of a metal varies greatly depending on the chemical nature of its support and it has often been suggested that some electron transfer from metal to support or vice versa, modifying the electronic structure of the metal, is at least somewhat responsible.33 By comparing the electrical conductivities and activities for methanol oxidation of silver, zinc oxide and Ag-ZnO, it has been conclusively shown that valency induction indeed occurs.34 Mechanism of Hydrogenation of Unsaturated Molecules and Related Processes.-The simplest catalytic process which a molecule can undergo is the substitution of its hydrogen atoms by those of an isotope (e.g.D from deuterium gas or heavy water, T from tritiated hydrogen). Even if the mole- cule is unsaturated, it is possible to find conditions under which substitution can be effected without addition. The exchange of propane with deuterium on a platinum-containing fuel-cell electrode has been studied with a view to establishing how propane is and exchange of neopentane with deuterium has been studied with especial regard to diffusion effects in porous catalyst^.^^ Results concerning the exchange of cycloheptane and cyclo- octane have also been rep0rted.3~ Garnett and his associates continue to examine the exchange of unsaturated molecules with heavy-water catalysed by the Group VIII metals.38 The unresolved problems concerning the mechanism of hydrogenation of alkenes continue to excite much interest, and a number of groups of workers are now employing isotopic-tracer techniques. However, isotopic methods are not essential, and much interesting information comes from looking at double-bond isomerisation during hydrogenation and from measuring comparative rates of reaction.Rates of hydrogenation of a large number of pairs of alkenes over Pt-SiOz have been measured:39 if the ratio of rates for two molecules A and B is RAB elc., it was shown that RAB x RBC = RAC. Ru-C is reported to show very different rates of hydrogenation for alkenes of different structure.40 In the liquid-phase hydrogenation of a number of 33 F. Solymosi, Catalysis Rev., 1967, 1, 233.34 G. M. Schwab and K. Koller, J . Amer. Chem. SOC., 1968,90, 3078. 35 H. J. Barger and A. J. Coleman, J . Phys. Chem., 1968, 72, 2285. 36 F. G. Dwyer, L. C. Eagleton, J. Wei, and J. C. Zahner, Proc. Roy. SOC., 1968, A , 301, 253. 37 B. S. Gudkov, E. P. Savin, and A. A. Balandin, Zzoest. Akad. Nauk S.S.S.R., Ser. khim., 1968, 509. 38 G . E. Calf and J. L. Garnett, Austral. J . Chem., 1968, 21, 1221; G. E. Calf, B. D.Fisher, and J. L. Garnett, ibid., 947; G. E. Calf, J. L. Garnett, and V. A. Pickles, ibid., 961. 39 R. Maurel, J.-M. Elene, J.-F. Mariotti, and J. Tellier, Compt. rend., 1968, 266, C, 599; R. Maurel and J. Tellier, Bull. SOC. chim. France, 1968,4191. 40 L. Kh. Freidlin, E. F. Litvin, and S. K. Tulayev, Neftekhimiya, 1968, 8, 155.Catalysis by Metals I27 cycloalkenes, Pt-Ah03 causes very little isomerisation : dialkylcycloalkenes yield 43-58 % of the corresponding cis-cycloalkane,41 while Rh-C gives slightly more isomerisation and 52-61 % of cis-cy~loalkanes.~~ The be- haviour of all the noble Group VIII metals for isomerisation of branched alkenes has been described,43 the metals falling in the now familiar sequence: ruthenium and osmium are confirmed to be quite active for double-bond migration.44 A revised theory for calculating product distributions from the reaction of ethylene with deuterium has been described.45 The reaction of cyclohexene with deuterium has been discussed,46 and the Horiuti-Polanyi mechanism has been successfully extended to describe the formation of products from the reaction of methyl oleate with de~terium.4~ The reaction of ethylene with tritiated hydrogen has been studied but, unfortunately, only under condi- tions where mass-transport limitation is either partial or complete.4s Anoma- lous effects are observed when ap-unsaturated carbonyl compounds are reduced with tritiated hydrogen.49 Heats of hydrogenation of three- and four-membered rings are reported,50 and the mechanism of reduction of substituted cyclopropanes has been discussed.51 When the spiro-ketone (1) is hydrogenated, no cleavage of the cyclopropane ring occurs with Raney nickel, the product being the saturated spiro-ketone (2), but with Pd-C in a number of solvents cleavage occurs with the intermediate formation of the alkene (3).52 Mann and his co-workers have reported the kinetics of the hydrogenation of propyne, but-1-yne and allene over a number of catalysts.53 Selectivity in the reduction of the three double-bonds in cyclohepta-1,3,5-triene falls in the sequence Pd> Rh> Pt,54 this being the same as with linear diolefins.4L A. S. Hussey, G. P. Nowack, G. W. Keulks, and R. H. Baker, J . Org. Chent., 1968, 33, 610. 42 A. S. Hussey, T. A. Schenach, and R. H. Baker, J . Org.Chem., 1968, 33, 3258. 43 M. Abubaker, I. V. Grostunskaya, and B. A. Kazanskii, Vestnik Moskov. Univ., 1968, 106; M. Abubaker, Z. S. Khrustova, I. V. Gostunskaya, and B. A. Kazanskii, ibid., 148. 44 G. C. Bond, G. Webb, and P. B. Wells, Trans. Faraday SOC., 1968,64,3077. 45 C. Kemball and P. B. Wells, J . Chem. SOC. (A), 1968, 444. 46 D. Moger, G. Mink, and F. Nagy, Magyar Ke'm. Folyoirat, 1968, 74, 315, 318. 47 H. J. Dutton, C. R. Scholfield, E. Selke, and W. K. Rohwedder, J . Catalysis, 1968, 4* L. Guczi and P. TCtknyi, Zeit. phys. Chem. (Leipzig), 1968, 237, 356. 49 H. Simon and 0. Bernguber, Tetrahedron Letters, 1968, 471 1. 50 R. B. Turner, P. Goebel, B. J. Mallon, W. Von E. Doering, J. F. Coburn, and 5 l W. J. Irwin and F. J. McQuillin, Tetrahedron Letters, 1968, 2195.52 M. T. Wuesthoff and B. Rickborn, J . Org. Chem., 1968, 33, 131 1. 53 R. S. Mann and K. C. Khulbe, Canad. J. Chem., 1968, 46, 623 idem, J. Catalysis, 1968,10,401; R. S. Mann and D. E. To, Canad. J . Chem., 1968,46, 161. 54 B. D. Polkovnikov, 0. M. Nefedov, E. P. Mikos, and N. N . Novitskaya, Zzoest. Akad. Nauk S.S.S.R., Ser. khim., 1968, 1240. 10, 316. M. Pomerantz, J. Amer. Chem. SOC., 1968, 90, 4315.128 G. C. Bond Selective poisoning of palladiuni catalysts by cadmium salts alters the position of attack in 6-methylhepta-3,5-dien-2-one from the 5,6 double-bond to the 3,4 d~uble-bond.~~ A number of interesting papers have appeared on the hydrogenation of aromatic rings. The exchange of hydrogen between 13C-labelled benzene and unlabelled cyclohexane occurs in the vapour-phase over Pt-AI203 at 157-182"c and over Au-MgO at 205-235"C.56 Self-exchange of mono- deuteriotoluene takes place at 60-100"c over platinum and nickel : transfer of D to and from the ovtho-position is markedly slower than to and from the nzeta- and para- positions.57 This observation may mean that some earlier work on the exchange of alkylaromatic compounds will have to be recon- sidered.Weitkamp has described a most comprehensive study of the stereo- chemistry of the hydrogenation of substituted naphthalenes;58 1,2-di-t-butyl- benzene is reduced over Rh-C to the corresponding cyclohexane which is almost all the cis-isomer although up to 45% of an intermediate olefin (2,3-di-t-butylcyclohexene) is also formed.59 This is a remarkable demonstra- tion of the effect of substituents on the course of a catalytic reaction. Selectivity aspects of the reduction of p-terphenylGO and of diphenyl ether6I have been described. Several careful studies of the hydrogenolysis of ethane have been re- ported.62 It now seems well established that activity for this reaction at 150-3OO"c is greater for the noble Group VIII metals than for the base metals, and that activity falls on passing from left to right through each row. This is in contrast to most hydrogenations where activity is maximum in the third column of Group V111. 55 L. Kh. Frcidlin, N. V. Boriinova, L. I . Gvintcr, and S. S. Danielova, Zhirr. Jiz. 56 G. Parravano, J . Catalysis, 1968, 11, 269. 5' K. Hirota, T. Veda, T. Kitayama, and M. Itoh, J . Phys. Chern., 1968, 72, 1976. 5* A. W. Weitkamp, Adv. Catalysis, 1968, 18, 1 . 59 B. van der Graff, H. van Bekkum, and B. M. Wepster, Rec. Trav. chim., 1968, 87, 60 Y . Bahurel, G. Descotes, and J. Sabadie, Bull. SOC. chim. France, 1968, 4259. 61 P. N. Rylander and M. Kilroy, Engelhard Industries Technical Bulletin, 1968, 9, 1. 62 J. H. Sinfelt and W. F. Taylor, Trans. Faraday SOC., 1968,64, 3086; J. H. Sinfelt and D. J. C. Yates, J . Catalysis, 1968,10, 362; G. K. Starostenko, T. A. Stovochotova, A. A. Balandin, and K. A. el Chattib, Vestnik Moskoc. Univ., 1968, 52. Khitn., 1968, 42, 98. 777.

ISSN:0069-3022    

DOI:10.1039/GR9686500121

出版商:RSC    

年代:1968

数据来源: RSC

摘要:

7 ELECTROLYTE SOLUTIONS By A. D. Pethybridge and J. E. Prue (Department of Chemistry, The University, Whiteknights, Reading, RG6 2A D) Raman and Infrared Spectroscopy.-The availability of commercial instru- ments and of laser sources of excitation has resulted in increased activity. In our understanding of electrolyte solutions, real progress beyond the assunip- tions of the continuum model for the solvent demands a satisfactory molecular picture of the solvent, and in particular of water, which we do not have at present. One of the best prospects of experimental progress is from spectro- scopic measurements, and an interesting paper by Walrafenl which reports Raman studies of the structure of water will be found in the recently pub- lished proceedings of a symposium on hydrogen-bonded solvent systems.Neither the assignment of frequencies nor the interpretation of intensity variations is easy and such work has not yet led to generally accepted conclu- sions, but there can be no doubt that the spectroscopic study of water itself and of HDO in water2 over wide temperature and pressure ranges will be the subject of much experimental effort in the near future. Rather more clear-cut conclusions have already emerged about ion-ion interactions in particular cases. An interesting paper by Irish and Davis3 reports a detailed study of the vibrational spectrum of the nitrate ion in aqueous alkali-metal nitrate solutions. Raman results are complemented by i.r. spectra. Above a concentration of 1 . 5 ~ , the positions, molar intensities, and especially the halfwidths of bands show pronounced dependence on both the concentration and the identity of the cation.The halfwidth changes are ascribed to the interaction of solvated pairs of ions. The appearance of a new nitrate band at 740 cm.-l in lithium and sodium nitrate solutions more concentrated than 7~ is ascribed to the effect of direct cation-anion contact. It is also of interest that the separation of two of the nitrate frequencies for lithium nitrate shows a cube-root dependence on concentration in a plot which includes a value for the molten salt. Other spectroscopic studies across the entire concentration scale up to and including the molten salt would be worthwhile. The same autliors4 report vibrational assignments for the species HgNO3’ and Hg(N03)2 from studies on aqueous mercury(rr) nitrate solutions.They conclude that the nitrate ion is monodentate in HgN03f. A paper5 characteristic of much recent spectroscopic work on nitrate solutions deals with aqueous cadmium nitrate and contains plentiful ref- erences to similar work. A detailed band assignment is made for the free G. E. Walrafen in, ‘Hydrogen-bonded solvent systems,’ eds. A. K. Covington and G. E. Walrafen, J. Chem. Phys., 1968, 48, 244. D. E. Irish and A. R. Davis, Canad. J. Chem., 1968, 46, 943. A. R. Davis and D. E. Irish, Inarg. Chem., 1968, 7 , 1699. 5 A. R. Davis and R. A. Plane, Inarg. Chetti., 1968, 7, 2565. P. Jones, Taylor and Francis, Ltd., London, 1968, p. 9. 129130 A . D . Pcthybridgc and J . E. Prue nitrate ion and for the inner-sphere species, CdN03+, for which a stability quotient is calculated from intensity measurements. A nitrate band at 740 cm.-l is again ascribed to the inner-sphere cation-anion complex.A study6 of bismuth nitrate solutions provides evidence for bidentate co-ordination by some nitrate ions, and a band at 235 cm.-l assigned to the symmetric metal- oxygen stretch is cited as evidence for covalency. It seems that up to four nitrate ions can attach themselves to a bismuth ion. The spectrum of the nitrate ion in acetonitrile solutions of zinc, cadmium, and mercury(I1) nitrates is strongly perturbed by cations, but only in the case of mercury(r1) is a metal-oxygen frequency dete~ted.~ Useful sources of reference to studies of particular complexes in solution are two recent reviews by Hester.8 Although the vibration of an ionic bond is not expected to produce a Raman line, it will have a large change of dipole moment associated with the vibration and should produce an intense line in the i.r.spectrum. It is of particular interest therefore that French and Woodg report the observation of far i.r. vibrational spectra of some ion pairs in non-aqueous solvents. A band at 175 cm.-l is found for sodium tetraphenylborate in pyridine, 1,4-dioxan, piperidine, and tetrahydrofuran. The wavenumber is independent of the solvent but does depend on the cation. These facts, and the value of the isotope shift found for ammonium tetraphenylborate on deuteriation of the ammonium ion suggest that the ion pairs are of the ‘contact’ rather than the ‘solvent- separated‘ variety.(The opposite conclusion for sodium tetraphenylborate in tetrahydrofuran has been reached from conductivity measurements by polymer chemists.1° However, it is questionable whether the conductivity measurements were of adequate precision or sufficiently rigorously analysed to lead to reliable conclusions.) French and Wood stress that the detailed analysis of ion-pair vibrational spectra will obviously require a more sophisti- cated ion-pair potential function than the hard-sphere model. Precise Raman intensity measurements continue to provide information on the degree of ionisation of strong electrolytes in concentrated solutions. Clarke and Woodwardl1 have measured degrees of ionisation of methyl- mercury( 11) methanesulphonate (1-5~) and sulphate (0.2-0.7~) by measure- ment of the integrated intensities of bands characteristic of the anions. In addition to problems of band overlap, there is no doubt that some earlier workers (but not Woodward) were over-optimistic concerning the independ- ence of molar integrated intensities of effects other than chemical bond formation. In the present paper, and in earlier ones dealing with the ionization of methanesulphonic acidl2a and of methylmercury(1r) nitrate,12b the anion 6 R. P.Oertel and R. A. Plane, Znorg. Chem., 1968,7, 1192. 7 C. C. Addison, D. W. Amos, and D. Sutton, J . Chern. SOC. ( A ) , 1968,2285. 8 R. E. Hester, Co-ordination Chem. Rev., 1967,2,319; Analyt. Chetn., 1968,40,32OR. M. J. French and J. L. Wood, J . Clrem. Phys., 1968, 49, 2358.10 D. Nicholls, C. Sutphen, and M. Szwarc, J . Phys. Chetn., 1968,72, 1021 ; J. Comyn, l1 J. H. R. Clarke and L. A. Woodward, Trans. Faraduy SOC., 1968,64, 1041. l2 (a) J. H. R. Clarke and L. A. Woodward, Trans. Faraday SOC., 1966, 62, 2226; F. S. Dainton, and K. J . Ivin, Electrochim. Acta, 1968, 13, 1851. (b) ibid., 3022.Electrolyte Solutions 131 concentration is calculated from the observed intensity and a molar integrated intensity equal to that of an ammonium salt solution (assumed to be com- pletely ionised) of equal refractive index. The justification for this is discussed, but it is recognised that with, for instance, different metal nitrate solutions (supposedly completely ionized) the molar intensities of solutions of the same refractive index can in some cases differ by as much as 10%. Clarke and Woodward use the word ‘dissociation’ where the present writers prefer ‘ionisation’ because in this method anions present in ion pairs (which, as well as higher aggregates, will undoubtedly occur in such concentrated solutions) are counted as ions.This raises a further general point in connection with questionable attempts to extrapolate ionisation quotients at high concentra- tions to obtain thermodynamic equilibrium constants. Consider a commonly invoked scheme of ionisation followed by dissociation such as K1 Ii 2 A-B +- A+B- $ A+ + B- From a degree of ionization given by the quantity which is conventionally calculated is We find on expansion of the right-hand side of equation (3) and substitution of the equilibrium constants K1 and.Kz, that even if changes of activity coefficients with concentration are ignored, Q = f(Kl,Kz,c); attempts to extrapolate such a function to c = 0 by assuming its concentration dependence to be solely determined by changes in activity coefficients are therefore wrong, and extrapolation on an empirical basis only is unreliable. According to Clarke and Woodward the degree of ionisation of methane- sulphonic acid is about 0.5 in an 1 IM solution. An estimate based on acidity function meas~remeiitsl~ gives about the same value for the degree of ionisa- tion at 3 ~ ! It has been claimed14 that this gross discrepancy can be avoided by changes in the assumptions made about the hydration of the proton in analysing the acidity function data.The polynuclear species reported as hydrolysis products of metallic cations in solution are sometimes surprising. The conclusion that the major species produced from bismuth(n1) is Bis(OH)lP, first reported from e.m.f. measure- ments, is well substantiated by ultracentrifuge and X-ray measurements. The vibrational spectrum of this complex both in solution and in the solid state has now been thoroughly studied.15a The work confirms that the bismuth l3 K. N. Bascombe and R. P. Bell, J. Chem. SOC., 1959, 1096. l4 J. G. Dawber, Chem. Comm., 1968, 5 8 . l5 (a) V. A. Maroni and T. G. Spiro, Inorg. Chern., 1968,7, 183; (6) ibid., 188.132 A. D. Pethybridge and J . E. Priie atoms lie at the corners of a regular octahedron with the hydroxy-groups along the edges.A normal co-ordinate analysis has been done; an anomalously high intensity for a low-frequency Raman band is cited as evidence for bis- muth-bismuth bonding. A similarly thorough analysis has been madel50 of the complex Pb4(0H)44+, which is a distorted cube with the lead and oxygen atoms arranged tetrahedrally. There is again evidence for metal-metal bond- ing. Nuclear Magnetic Resonance.-This has been an area of intense activity during the past year, although much of the work has increased under- standing of n.m.r. phenomena rather than substantially enhancing know- ledge of electrolyte solutions. For some cations with a sufficiently slow exchange rate of water molecules between the first co-ordination sphere and the bulk solvent, separate 1H signals can be observed for water molecules in the two situations and solva- tion numbers calculated from the areas of peaks for solvated ions.Separation of signals is achieved by temperature lowering and/or dilution of aqueous solutions with organic solvents. The ion Mg(H20)s2+ can be 'seen' in con- centrated aqueous solutions of the nitrate or perchlorate at temperatures below -70°c,16 and in solutions in aqueous acetone16, l7 over a much wider range of temperature and concentration. In methanolic acetone the ion Mg(MeOH)62f is formed.16 Complete analysis of line shapes should eventually give rate constants for the solvent exchange. As well as Mg(H20)s2+ the existence of AI(HZO)~~+,~~ Be(H~0)4~+,l~ Ga(HzO)s3-t,18 Co(H20)s2+,19 and N i ( H z o ) ~ ~ + , ~ ~ has been confirmed in similar manner.In mixed aqueous solvents, whilst dimethyl sulphoxide and NN-dimethylformamide compete with water for the inner-sphere of aluminium ions, acetone, tetramethylurea, dioxan, and tetrahydrofuran do not.21 In solutions of aluminium perchlorate in water-methyl cyanide mixtures, a separate signal has been found22 for each of the six species [AI(H20)6-n(MeCN)n]3+ with n = 1-6. When the 1H signal of solvent-OH in the co-ordination sphere of a cation can be separately seen, e.g. as with magnesium perchlorate in methanol, the residual shift of the main solvent signal from that of the pure solvent is due to the anion, and this makes possiblez3 the unambiguous calculation of a molal anion-shift without the usual arbitrary division of shifts between anions and cations. l6 N.A. Matwiyofl'and H. Taube, J . Amer. Chern. Suc., 1968,90, 2796. l7 R. G. Wawro and T. J. Swift, J. Amer. Chem. SOC., 1968, 90, 2792; A. Fratiello, A. Fratiello, R . E. Lee, V. M. Nishida, and R. E. Schuster, J . Chem. Phys., 1968, R. E. Lee, V. M. Nishida, and R. E. Schuster, Chem. Comm., 1968, 173. 48, 3705. l9 N. A. Matwiyoff and P. R. Darley, J . Phys. Chem., 1968,72, 2659. 2o T. J. Swift and G. P. Weinberger, J . Amer. Chem. SOC., 1968, 90, 2023. 21 A. Fratiello, R. E. Lee, V. M. Nishida, and R. E. Schuster, J . Chem. Phys., 1967, 47, 495 1. 32 L. D. Supran and N. Sheppard, Chern. Cumm., 1967, 852. 23 R. N. Butler, E. A. Philpott, and M. C. R. Symons, Clrem. Cutnm., 1968, 371.Electrolyte Solutions 133 Earlier work has been continued in which the addition of a suitable para- magnetic cation (e.g.Co2+) to shift the 1 7 0 signal of non-co-ordinated water molecules enables the signal of water co-ordinated to the diamagnetic cations Be2+, A13+, or Ga3+ to be seen separately.24 The technique has also been used25 to study the hydration of organometallic cations. The complexes of cobalt(r1) formed in concentrated aqueous hydrochloric acid solutions have been studied26 by observations of 1 7 0 and 35Cl line shifts. The relative solvating abilities of solvent molecules in pure solvents and in solvent mixtures is also often studiedz7* 28 by observations of chemical shifts. Sometimes the conclusions are surprising, e.g. that sodium ions interact more with the amine than with water in an ethylenediamine-water mixture.2g In such cases one doubts whether the chemical shift observations are very directly related to the strengths of chemical interactions.A useful review3O of the study of ion-solvent and ion-ion interactions by magnetic resonance tech- niques has appeared recently, as well as a chapter on n.m.r. studies of electro- lyte solutions.31 In the quantitative study of the ionisation of strong acids, new lH studies from 0-65"c of concentrated perchloric acid solutions in water and deuteri- ated water are reported,32a and there is further discussion of this and the corresponding nitric acid system in a subsequent paper.32b In some respects the problems are similar to those in Raman studies. It is clear that the lH shift of H30' is more dependent on its state of solvation and on the specific effects of anions than was initially thought.Measurements of degrees of ionisation of methanesulphonic acid by observation of the IH resonance of ionisable protons and of methyl protons agree well33 with the Raman results of Clarke and Woodwardlzlf at concentrations > 1 0 ~ . (Covington and Lilley33 calculate the chemical-shift contribution from hydroxonium ions from mea- surements on dilute solutions of the acid itself in order to allow for the specific effect of the anions). In solutions more dilute than 7 . 5 ~ , where the degree of ionisation is above 0.8, there are discrepancies between the Raman and n.m.r. work. There is therefore a large discrepancy between the acidity constant values of 73 and 16 mole l.-l calculated by the two pairs of investi- gators.The questionable nature of the necessary extrapolation was com- mented upon earlier. Electron Spin Resonance.-For the rather specialised class of anions which 24 D. Fiat and R. E. Connick, J . Amer. Cl7em. Soc., 1968, 90, 608. 35 G. E. Glass, W. B. Schabacher, and R. S . Tobias, Zmrg. Chem., 1968, 7, 2741. 26 A. H. Zeltmann, N. A. Matwiyoff, and L. 0. Morgan, J . Phys. Chem., 1968,72, 121. 27 J. C . Fanning and R. S. Drago, J . Amer. Chem. SOC., 1968, 90, 3987. 28 A. Fratiello, R. E. Lee, D. P. Miller, and V. M. Nishida, Mol. Phys., 1967, 13, 349. 29 E. G. Bloor a i d R. G. Kidd, Canad. J. Chem., 1968,46,3425. 30 J. Burgess and M. C . R. Symons, Quart. Rev., 1968, 22,276. 31 C. Durell, in 'Progress in Nuclear Resonance Spectroscopy,' eds.J. W. Emsley, J. Feeney, and L. H. Sutcliffe, Pergamon Press, Oxford, 1968, vol. 4. 32 (a) R. W. Duerst, J. Chem. Phys., 1968,48,2275; (b) 0. Redlich, R. W. Duerst, and A. Merbach, ibid., 1968, 49, 2986. 33 A. K. Covington and T. H. Lilley, Tmrs. Foradoy Suc., 1967, 63, 1749.134 A . D . Pethybridge andJ. E. Prue are paramagnetic, e.g. semi-quinones, aromatic nitro-anions, and hydrocarbon radical anions, this method is potentially capable of providing detailed structural and kinetic information about ion pairs in organic solvents, and in particular about the distribution between ‘contact’ and ‘solvent-separated‘ varieties. Three review-type articles30, 34 have recently appeared. Ultraviolet and Visible Spectrophotometry.-Two papers35 report an excel- lent study over a temperature range of the equilibria: $Cr2072- + 3H20 + HCr04- + Cr042- + H+ Rather surprisingly, the outer-sphere association constant of the hydroxy-ion with the tris(ethylenediamine)cobalt(IzI) ion in dioxan-water and in dioxan- deuterium oxide mixtures is four times greater than with the tris(ethy1ene- diamine)chromium(rrr) c0mplex.~6 The ion-pair between tris(ethy1enediamine)- ruthenium(ru) and the iodide ion has a well-defined electron transfer band in the visible at a lower energy than the first cl-d tran~ition.3~ Conventional studies38 of the effect of added thiosulphate and selenite ions on the spectrum of tris(ethylenediamine)cobalt(m) refute the astonishing conclusion39 from optical rotation and circular dichroism measurements that, in 2~-sodium perchlorate, complexes such as [Co en33+ SZO,~-]+ with a stability quotient of 150 1.mole-l are formed as well as further complexes up to [Co en,3+(S2032-)4]5-. Spectroscopic studies40 of the association of copper(1r) with sulphate and of bis(ethylenediamine)copper(II) with thiosulphate ions add little to what is already known.41 Plots against the ionic strength Z of log(K’/yi) where K’ is the association quotient and the mean free-ion activity coefficient yk is given by (4) are found approximately to converge to a common intercept with any reason- able assumption about p. This is simply a consequence of the fact that -log yh = 4AZ*/(1 + pZ6) AI+/(l + ( p + Ap)I’) A I i / ( l + pZ*) - AApI ( 5 ) provided that A p is small. What is crucial is the calculation of K’ values and these are less certain than the authors imply.They are obtained from experi- ments at constant ionic strength. It is assumed that the association quotient does not vary with ionic composition at a fixed ionic strength, which is probably ~njustified.4~ No comment is made on the large variation with ionic strength of the values obtained for the difference between the molar absorp- tivities of free and associated copper ions. More importance attaches to the 34 (a) N. Hirota, J.Phys. Chem., 1967,71, 127; (b) M. C. R. Symons, ibid., 172. 35 (a) H. G. Linge and A. L. Jones, Austral. J. Chem., 1968,21, 1445; (b) ibid., 2189. 36 S. C. Chan, J . Chem. Suc. ( A ) , 1967, 2103. 37 Sr. H. Elsbernd and J. K. Beattie, Znurg. Chem., 1968, 7, 2468. 38 J.Olsen and J. Bjerrum, Acta Chem. Scand., 1967, 21, 1112. 39 R. Larrson, S. F. Mason, and B. J. Norman, J . Chern. SOC. ( A ) , 1966, 301. *O P. Hemmes and S. Petrucci, J . Phys. Chem., 1968, 72, 3986. 41 R. A. Matheson, J . Phys. Chem., 1967, 71, 1302.Electrolyte Solutions 135 ultrasonic absorption results in the same paper.40 The observation of a common absorption maximum for both complexes shows that inner-sphere complex formation occurs by the exchange of the anions with weakly attached water molecules in the axial positions of both CU(H20)s2+ and Cu en,(H~0)2~+. In the spectrophotometric experiments, the inner-sphere complexes probably act as optical 'markers' for ions in the associated class, not all of which are necessarily optically different from free ions.A continuation42 of earlier work on the absorption spectra of cobalt(r1) and nickel(r1) halide solutions up to high temperatures and pressures, leads to the conclusion that in concentrated alkali-metal halide solutions the deep blue colour of both cobalt and nickel solutions above 200-250"c is due to the presence of tetrahedral complexes such as CoC13(H20)-, NiC12(H20)2, or NiC13(H20)-. An interesting paper43 claims that precise differential refractonietry on solutions of the strong trifluoroacetic acid can be used to obtain values for both the acidity constant Ku and the apparent molar volume change AVO on dissociation. The method of analysing the data in which A Vo is an adjustable parameter as well as Ku is such that it is probably a conventional dissociation constant that is measured.The concentration of associated species is [see equation (l)], [A-B] + [A+B-1, and Ku = K1K2/(1 + KI). Conductivity.-A good general review44 of the techniques and results of precise measurements with special reference to non-aqueous solvents has been published. The non-specialist may be surprised to learn that to describe the concentration dependence of molar conductance there is still no agreement about the exact equation which incorporates an ion-size parameter and corresponds to the Debye-Huckel equation. This is particularly irritating as the number of results with an accuracy of about 0.01 % in the molar conduct- ance increases. However, it now seems to be generally agreed45 that retention of terms of order c3j2 is necessary.Although there is some evidence46 that the numerical difference between the predictions of the latest version of the Fuoss-Onsager equation 45c and that of has diminished, a paper by Pitts, Tabor, and which compares the two treatments, emphasises that some physical assumptions implicit in the Fuoss-Onsager treatment as well as mathematical details are open to criticism. A recent paper by Hsia and F u o s ~ ~ ~ which analyses results for aqueous solutions of caesium bromide and iodide begins with the surprising sentence 'Association of 1 :1 electrolytes has not 42 H.-D. Liidemann and E. U. Franck, Ber. Bunsengesellschaft Phys. Chem., 1968, 43 E. Grunwald and J. F. Haley, jun., J. Phys. C/iem., 1968,72, 1944. 44 J. Barthel, Angew. Chem., 1968, 80, 253.45 (a) R. Fernhndez-Prini and J. E. Prue, 2. phys. Chem. (Leipzig), 1965, 228, 373; (6) J.-C. Justice, J. Chim.phys., 1968, 65,353; (c) R. M. Fuoss and Kai-Li Hsia, Proc. Nat. Acad. Sci. U.S.A., 1967, 59, 1550. 72, 514. 46 R. Fernandez-Prini, Trans. Faraday SOC., 1968, 64, 2146. 47 E. Pitts, Proc. Roy. SOC., 1953, A , 217, 43. 48 E. Pitts, B. F. Tabor, and J. Daly, Trans. Faradny SOC., 1969, 65, 849. 49 Kai-Li Hsia and R. M. Fuoss, J . Amer. Chem. SOC., 1968, 90, 3055.136 A. D. Pethyhridge arid J . E. Prrir been seriously considered since the advent of the Debye-Hiickel theory of electrolytes in 1923’. A reference50 to a review of the work of C. W. Davies and others seems appropriate. These workers invoked association in inter- preting data for alkali-metal salts of oxyanions, whilst in the latest contribu- tions of Fuoss and co-worker~4~~ 5l it is invoked for dilute (<O*~M) aqueous solutions of the halides.For instance, for potassium chloride A , = 149.90 C2-l cm.2 mol-l, K A = 0-79 1. mole-l, a = 5-65 A. The ion-size parameter is much larger than the Bjerrum distance of 3.5 8, at which the mutual electrical potential energy of a pair of oppositely charged ions is equal to 2kT. Before accepting the need for a third adjustable parameter in fitting conductivity data for alkali-metal halides in water, the examination of the theoretical basis of the equations used and numerical comparison with the Pitts equation are highly desirable. A further surprising feature of Fuoss’s latest papers is the use of the Debye-Huckel limiting law for the activity coefficient of free ions because52 ‘ions in contact are counted as pairs and long-range interactions between free ions cannot depend on the size of ions but only on their charges’.After re-evaluating the coefficient of the c3I2 term in the Fuoss-Onsager treatment, 53 concludes that precise data for I : 1 electrolytes in a wide range of pure and mixed low-dielectric solvents are satisfactorily described by an association treatment with an ion size parameter in both the conductivity equation and the ionic activity coefficient equation equal to the Bjerrum distance. Furthermore, the variation of the association constants so obtained is quantitatively in accord with the traditional Bjerrum expression. Such a treatment may raise problems in connection with the conductance equation, but a similar procedure54 successfully describes the conductivity behaviour of 2 : 2 sulphates in water and a common description can be given of conductimetric, thermodynamic, and spectrophotometric results.In the sulphate cases, and it would be interesting to examine whether this is also true of Justice’s results, the association distance between free and associated ions could be varied considerably (with appropriate adjustment in K A ) ~ ~ ~ and an acceptable fit of the data still achieved. A further problem in studying ion association by conductivity methods is raised by Onsager and Provencher’s papers6 in the P. J. W. Debye memorial issue of the Journal of the American Chemical Society. If the rate constants for association and dissociation are appropriate, an additional mechanism for relaxation of asymmetric charge distribution is provided and the relaxation effect can be reduced by as much as 23 x.50 C. W. Davies, in ‘The Structure of Electrolytic Solutions,’ ed. W. J. Hamer, Wiley, 61 Ying-Chech Chiu and R. M. Fuoss, J . Phys. Chem., 1968, 72, 4123. 52 R. M. FUOSS, Rev. Pure Appl. Chetrr. (Australia), 1968, 18, 125. 53 J.-C. Justice, R. Bury, and C. Treiner, J . Chim. phys., 1968, 65, 1708; J.-C. Justice, 54 (a) W. G. Davies, R. J. Otter, and J. E. Prue, Discuss. Furaday Soc., 1957, 24, 103, 55 R. A. Matheson, J . Phys. Chem., 1968,72, 3330. 5G L. Onsager and S. W. Provencher, J . Amer. Chern. SOC., 1968,90, 3134. New York, 1959, p. 19. personal communication.123; (b) J. E. Prue, ‘Ionic Equilibria,’ Pergamon Press, Oxford, 1966, p. 64.Electrolyte Solirtions 137 The conductivity of tetra-alkylammonium salts continues to be a favourite topic of study in aqueous, non-aqueous, and mixed solvents. Fernandez- Prini4'j has shown that by treating these salts as incompletely dissociated in water a self-consistent analysis can be given of both conductivity and thermo- dynamic properties. The association constants increase with the size of anion for a given cation, which is also found in ethanol and propan01.~~ Trialkylsulphonium iodides have been investigatedS8 in water, methanol, and acetonitrile over a temperature range. In non-aqueous solvents they seem to be considerably more associated than tetra-alkylanimonium iodides.The limiting conductances of these organic cations, their effects on solvent vis- cosity, and the temperature coefficients of these properties in a variety of solvents have been lengthily discussed59 in terms of so-called solvent-structure theories. These theories, whilst interesting, are as yet of a qualitative nature and although 'explanations' abound, predictions are rare. It is certainly an interesting experimental fact that, whatever the explanation, large organic ions in aqueous solutions move abnormally slowly and the effect is extremely temperature dependent. Tt is also interesting that such effects disappear with the (HOC2H4)4Nf ion. Shedlovsky and co-workers continue60 an interesting series of papers on the behaviour of acetic acid, sodium acetate, hydrochloric acid, and sodium chloride in alcohol-water mixtures at several temperatures.An extensive series of measurements with a lengthy discussion is reported61 on the conductivity and viscosity over a temperature range of aqueous solutions of hydrochloric acid and of potassium hydroxide, fluoride, and chloride. Transport-number measurements have made possible the calculation of ionic conductances in sulpholane.62 Many groups are studying the properties of electrolyte solutions at elevated temperatures and pressures, and a list of these is given in a review by Mar- ~ha11.~3 The behaviour of aqueous solutions at supercritical temperatures and pressures is of interest to geochemists, steam turbine designers, and those concerned with economic methods for producing potable water from sea water.From the theoretical standpoint it is interesting that the dielectric constant of water varies from unity at supercritical temperatures and low pressure to about 100 at 0"c and 1 kbar. Quist and MarshallM have during the past year reported measurements on aqueous solutions of NaCl, HBr, NaBr, and NH3 at 0-100"c at pressures to 4 kbar. The results can be qualitatively explained by the effect of temperature and pressure on density, dielectric constant, and viscosity. As the density falls it becomes necessary to introduce 57 D. F. Evans and P. G. Gardam, J . Phys. Chern., 1968, 72, 3281. 58 D. F. Evans and T. L. Broadwater, J . Phys. Chem., 1968, 72, 1037. 59 R. L. Kay, G. P. Cunningham, and D. F. Evans, ref. 1, p. 249. 6O M. Goffredi and T.Shedlovsky, J . Phys. Chem., 1967,71, 4436. 61 T. Erdey-Gniz, E. Kugler, and L. Majthknyi, Electrochim. Acta, 1968, 13, 947. 62 M. Della Monica, U. Lamanna, and L. Senatore, J . Phys. Chem., 1968,72, 2124. 63 W. L. Marshall, Rev. Pure Appl. Chem. (Australia), 1968,18, 167. 134 A. S. Quist and W. L. Marshall, J . Phys. Chem., 1968,72, 684, 1545, 2100, 3122.138 A . D. Pethybridge and J. E. Prue association constants even for the alkali halides (e.g. at p = 0.75 g. ~ m . - ~ for NaCl). It seems to be generally true that log KA at any temperature is to a good approximation a linear function of log C H ~ O with an almost temperature- independent slope. This suggests equilibria of the type and Quist and MarshaW5 show that the same procedure has considerable success in representing the variation of equilibrium constants with the composition of mixtures of polar and non-polar solvents at ordinary temp- eratures and pressures. Ritzert and FranckG6 report measurements with an inductively heated cell on solutions of salts with divalent cations [BaC12, Ba(OH)2, and MgS041 as well as new measurements on potassium chloride.The maximum conductance values are lower for the salts with divalent ions. Horne and co-workers discuss the effect of pressure on the conductivity of tetra-alkylammonium halides in water67 and of potassium chloride in water- alcohol mixtures.68 Studies on non-aqueous solutions over a temperature and pressure range are reported for tetra-alkylammonium salts in nitr~benzene~~ and in a~etone.~O Two papers71 discuss the design of suitable conductance cells for high-pressure work.Lown, Thirsk, and Wynne-Jone~7~ report measurements on acetic acid, and show that the results at 25"c can be satis- factorily fitted by treating the change in standard partial molal isothermal compressibility as independent of pressure up to 2 kbar. The method is also tested with literature results for the dissociation of water and the ammonium ion. Thermodynamic Properties.-Single electrolytes. Between thermally isolated adjacent drops of solvent and solution in a closed space saturated with solvent vapuur a temperature difference is established. With careful calibration, this effect can provide a method for determining (Po - p)/po where po and p are the vapour pressure of solvent and solution respectively.Instruments which operate on this principle, known as 'vapour pressure osmometers,' have become commercially available and it seems that they are capable of providing values of osmotic coefficients of useful accuracy at lower molalities than the isopiestic method. Only small quantities of solution are required. The method has been used to supplement isopiestic measurements in this way in studies of aqueous solutions of tri-n-alkylsulphonium halides,73 whose behaviour is similar to that of tetra-alkylammonium salts, and of disulphonic acids and 65 A. S . Quist and W. L. Marshall, J . Phys. Chern., 1968, 72, 1536. 66 G. Ritzert and E. U. Franck, Ber. Bunsengesellschaft Phys. Chem., 1968, 72, 799. 67 R. A. Horne and R. P. Young, J . Phys. Chem., 1968,72, 1763.68 R. A. Horne, D. S. Johnson, and R. P. Young, J . Phys. Chem., 1968,72,866. 69 F. Barreira and G. J. Hills, Trans. Faraday SOC., 1968,64, 1359. 70 W. A. Adams and K. J. Laidler, Canad. J . Chem., 1968,46, 1977, 1989, 2005. 71 (a) D. A. Lown and Lord Wynne-Jones, J. Sci. Znstr., 1967, 44, 1037; (b) A. B. 72 D. A. Lown, H. R. Thirsk, and Lord Wynne-Jones, Trans. Faraday SOC., 1968, 64, 73 S. Lindenbaum, J . Phys. Chem., 1968, 72, 212. Gancy and S . l3. Brummer, J . Electrochem. SOC., 1968, 115, 804. 2073.Electrolyte Solutions 139 their salts with divalent cati0ns.7~ An activity coefficient of 1457 for uranyl perchlorate in water at a molality of 5.5 has been quoted75 as an example of extreme behaviour. This value was determined by the isopiestic method, but a recent study76 of the distribution of uranyl perchlorate between aqueous solution and carbon tetrachloride containing trialkyl phosphate gives an approximate value for the activity coefficient of 3.9 at a molality of 5.5.The value was confirmed by sedimentation equilibrium in the ultracentrifuge. It is suggested that the abnormally high isopiestic value is probably due to hydrolysis of the cation which increases the number of solute particles in soh tion. A heating-curve method was used by Garnsey and P r ~ e ~ ~ for the precise cryoscopic determination of osmotic coefficients of alkali-metal salts in dimethyl sulphoxide and in sulpholane (tetrahydrothiophen 1,l-dioxide). Although the dielectric constants of the two solvents are similar, the ionic solvating power of dimethyl sulphoxide is much greater than that of sul- pholane, which results in opposite sequences of osmotic coefficients for the alkali-metal perchlorates, and striking differences in degrees of ion association in the two solvents, particularly for lithium chloride. The results for lithium chloride in dimethyl sulphoxide are in fair agreement with earlier cryoscopic measurements by a different technique, but activity coefficients calculated from both sets of results do not agree with those obtained from e.m.f.mea- surement~.~~ The discrepancy awaits explanation. A recent contribution to the theory of the subject is a preliminary note by Rasaiah and Friedman79 proposing a model which superposes on the Coul- ombic potential between rigid non-polarisable ions a square-topped well or mound over the distance r to r + 2w where r is the sum of the crystallographic radii and w the radius of a water molecule.With a mound of height $kT they claim a surprisingly good description of the activity coefficient behaviour of aqueous sodium chloride. The relationship of the model to hydration effects is obvious. A semi-empirical equation has recently been suggesteds0 which is based on treating a concentrated solution by a so-called ‘two-structure’ model, one part following Debye-Hiickel behaviour, the other that of a randomised fused salt. A three-parameter equation is derived which, like the conventional power series in I s , fits data up to a molality of 2, but has the additional virtue that two of the parameters are functions only of the charge type of the electro- lyte. 74 0.D. Bonner, C. Rushing, and A. L. Torres, J. Phys. Chem., 1968,72, 4291. 75 R. A. Robinson and R. H. Stokes, ‘Electrolyte Solutions,’ 2nd Ed., Butterworths, 76 K. Schwabe, R. Kretschmer, R. Gartner, and R. Rottenbach, 2. phys. Chem. 77 R. Garnsey and J. E. Prue, Trans. Faraday SOC., 1968, 64, 1206. 78 W. H. Smyrl and C. W. Tobias, J . Electrochem. Soc., 1968, 115, 33; J. N. Butler, 79 J. C. Rasaiah and H. L. Friedman, J . Phys. Chem., 1968, 72, 3352. 80 M. H. Lietzke, R. W. Stoughton, and R. M. FUOSS, Proc. Nut. Acad. Sci. U.S.A., 1959, p. 218. (Leipzig), 1968, 238, 391. private communication. 1968, 59, 39.140 A. D . Pethybridge and J. E. Pule Studies of free energies and enthalpies of transfer of electrolytes between solvents continue.81 Those concerned with the resolution of free energies and enthalpies into ionic contributions by some kind of extrathermodynamic assumption will be interested in five recent papers.82 Morriss2e also discusses the present position regarding the ionic radii of alkali-metal and halide ions.There has been marked activity in the determination of apparent molar volumes and their concentration dependence. At sufficiently low concentra- tions the behaviour agrees with Debye-Huckel predictions and the zero- concentration values should eventually further our understanding of ion- solvent interaction. Dunns3 has continued his work with precise results by a dilatometric method for 1 : 1 and 2: 1 salts over a temperature range. Franks and Smiths4 have reported results for sodium and potassium chlorides and tetra-alkylammonium halides obtained with a magnetic float apparatus capable of measuring densities to one part in lo6 (at which point salt impurities become the factor limiting accuracy and reproducibility).Ellis85a, has reported measurements for several electrolytes at temperatures up to 2Oo0c, which have been discussed by Gluecka~f.~~c Other papers report the results of studies of the apparent molar volumes of tetra-alkylammonium halides in water (over a temperature range), g6 in ethanol-water mixtures,87 and in deuterium oxide.8s The differences between the limiting values in deuterium oxide and in water are negative for sodium halides but positive for tetra- alkylammonium bromides. Mixed electrolytes.When studying the thermodynamic properties of solutions of mixed electrolytes there are two basic aims; firstly the prediction of the properties of mixtures from the properties of single electrolytes and secondly the determination of functions suitable for the interpolation of results at intermediate concentrations. Osmotic coefficients (4) of the solvent are usually obtained from isopiestic measurements and activity coefficients(y) of solutes calculated from these by means of the equations of Scatchard or of McKay and Perring. Activity coefficients calculated by the two methods usually agree quite well: the former involves rather cumbrous equations and the latter an unreliable graphical integration from zero concentration. Alternatively, activity coefficients are determined from potentiometric meas- urements on suitably designed cells.81 (a) H. P. Bennetto, D. Feakins, and K. C. Lawrence, J . Chem. Soc. ( A ) , 1968, 1493; (b) J. H. Stern and J. Nobilione, J . Phys. Chem., 1968,72, 1064, 3937. 82 (a) R. G. Bates, ref. I , p. 49; (b) H. P. Bennetto and D. Feakins, ref. 1 , p. 235, and ref. 81(a); (c) C. L. de Ligny, M. Alfenaar, and N. G. Van der Veen, Rec. Trau. chim., 1968, 87, 585; ( d ) W. A. Millen and D. W. Watts, J . Amer. Chem. SOC., 1967, 89, 6051; ( e ) D. F. C. Morris, Structure and Bonding, 1968, 4, 63. 83 L. A. Dunn, Trans. Faraday SOC., 1968, 64, 2951. 84 F. Franks and H. T. Smith, Trans. Faraday Soc., 1967, 63, 2586. 85 (a) A. J. Ellis, J . Chem. SOC. (A), 1968, 1138; (b) A. J.Ellis and I. M. McFadden, 86 (a) F. J. Miller0 and W. Drost-Hansen, J. Phys. Chem., 1968,72, 1758; (b) R. Gopal 87 I. Lee and J. B. Hyne, Canad. J . Chem., 1968,46,2333. 88 B. E. Conway and L. H. LalibertC, J . Phys. Chem., 1968,72,4317. Chem. Comm., 1968, 516; ( c ) E. Glueckauf, Trans. Faraday SOC., 1968, 64,2423. and M. A. Siddiqi, J . Phys. Chem., 1968,72, 1814.Electroly f e Soliltions 141 Isopiestic measurements on mixed electrolytes are made on sets of solutions of selected mole fractions of solutes B and C but whose ultimate total concen- tration is fixed by the criterion that all solutions must at equilibrium have the same vapour pressure. A single-electrolyte solution is included with each set for calibration purposes. The results are fitted to a suitable smoothing function and + and y calculated at various solute mole fractions of B and C and rounded values of the total ionic strength, the upper limit of this usually being fixed by a limit of solubility.It is often difficult to compare different measurements on the same system because individual workers select different solute mole fractions and ionic strengths. Typical of this kind of work are the isopiestic measurements of Wu, Rush, and ScatchardS9 on the four pairwise mixtures containing a common ion of NaCl, Na~S04, MgC12, and MgS04 in water at 25”c. These workers analysed their results by Scatchard’s method and found that all components obey within experimental error the linear relationship (6) known as Harned’s Rule up to a total ionic strength of 6 mole kg.-l.The rule states that log YB(C) = log yS(0) - %BC XC I (6) where YB(C) is the activity coefficient of salt B in a mixed solution of total ionic strength Icontaining a mole fraction x c of salt C, log yB(0) is the activity coefficient of B alone at the same ionic strength and aBC is a constant deter- mined by B and C and the total ionic strength. Platford has confirmed these results for mixtures of NaC1-MgC1zg0 and NaCl-Na2S04.91 He treated his results by both methods outlined above and obtained similar conclusions from each. Butler and co-workers have also obtained confirmatory results for the activity coefficient of one component of these mixtures from potentio- metric measurements on cells of the typeg2 Na(Hg) INaf, M2+, C1-, HzO IAgCl, Ag where M = Mg or Ca, andg3 Pb(Hg), PbSO, INa+, SO,2-, C1-, H20 I Na glass electrode.Unfortunately, the actual e.m.f.’s of this second cell are not reported, but only the calculated activity coefficients. The theoretical background of the subject is discussed in a recent book by Harned and RobinsonY4 and in a chapter of a bookg5 and in a paperg6 by 89 Y . C. Wu, R. M. Rush, and G. Scatchard, J . Phys. Cliei~i., 1968,72,4048. R. F. Platford, J. Phys. Chem., 1968, 72, 4053. 91 R. F. Platford, J . Chem. and Eng. Data, 1968, 13, 46. 92 J. N. Butler and R. Huston, J . Phys. Chem., 1967, 71, 4479. 93 J. C. Synott and J. N. Butler, J . Phys. Chem., 1968,72, 2474. 94 H. S. Harned and R. A. Robinson, ‘Multicomponent Electrolyte Solutions,’ 95 E. A. Guggenheim, ‘Applications o f Statistical Mechanics,’ Clarendon Press, 96 E.A. Guggenheim, Trans. Faraday SOC., 1966, 62, 3446. Pergamon Press, Oxford, 1968. Oxford, 1966, p. 165.142 A . D. Pethybridge and J. E. Prue Guggenheim. Scatchard97 derives equations for the excess free energy of mixing and gives a brief, rather personal, review of the historical setting of the various approximations made in the study of the thermodynamic properties of mixed electrolyte solutions. Pan98 has derived a form of the McKay-Perring equation which avoids graphical integration and is quite closely related to Scatchard’s equations. The latter are clarified in papers by Robinson and co-workers on NaC1-KClg9 and KC1-CaCl2lo0 mixtures, and by Rush and Johnsonlo1 who tabulate corrections of the many misprints in Scatchard’s original paper.Rush and JohnsonlOl report isopiestic measurements on pairs of solutes from the trio HC104, LiC104, and NaC104 as well as an extension of measurements on pure NaC104 up to 16 mole kg.-l. They find that Harned’s rule is not obeyed for either salt in the system HC104-NaC104. More interest- ing is the observationlo2 from potentiometric measurements that in binary mixtures of hydrochloric acid with alkaline-earth perchlorates, Harned’s rule is sometimes obeyed by one component of the mixture but not by the other, e.g. obeyed by Sr(C104)~ or Ba(C104)~ but not by HCl. HE GE Guggenheimg6 has shown that the so-called ‘square-cross’ rule for thermo- dynamic functions implies that triplet interactions are negligible, but does not imply small or negligible interactions between ions of the same sign.The square-cross rule for free energies has been recently verifiedlo3 for the system shown in the diagram (the values shown are for a total molality of one); the 97 G. Scatchard, J . Amer. Chem. SOC., 1968, 90, 3124. 98 C. Pan, J . Phys. Chem., 1968,72, 2548. 99 R. M. Rush and R. A. Robinson, J . Tenn. Acad. Sci., 1958,43, 22. loo R. A. Robinson and A. K. Covington, J . Res. Nat. Bur. Stand., Sect. A , 1968, 72, 101 R. M. Rush and J. S. Johnson, J . Phys. Chern., 1968,72, 767. lo2 I. A. Weeks, Austral. J . Chem., 1967, 20, 2367. lo3 A. K. Covington, T. H. Lilley, and R. A. Robinson, J . Phys. Chern., 1968,72,2759. 239.Electrolyte Solutions 143 rule has been previously demonstrated for enthalpies in the same system.With the same assumptions LilleylO4 has shown that the enthalpy and volume changes on mixing equimolar proportions of 1 :1 electrolytes without a common ion can be predicted from experimental data on only one pair of salts with a common cation and one pair with a common anion in conjunction with data for each individual salt. The square-cross rule has also been demon- strated105 for volume changes in the system LiCl, NaCl, Li2SO4, and NazS04. Acid-base equilibria. In their search for a wide range of well defined buffer solutions Bates et al.lo6 report conventional pa^ values for solutions of piperazine phosphate (PzH3P04) over a wide temperature range. The salt is stable and easily purified and the solutions have a high buffer capacity. The pa^ of a 0-05 molal solution at 25"c which contains the species PzHz2+, PzH+, H2P04-, and HP(h2- is 6.26.This value was obtained by extrapolation of measurements on a cell without transference with various quantities of added potassium chloride. As the operational pH is measured in a cell with a liquid junction, values of pH and pa^ will not usually agree. For buffers from acids of type HA, A pa^ - pH m 0, while for HBf, A w 0.013, and for H2B2+, A is even larger. For the piperazine phosphate buffer, A = 0.03 pH units. The same group has published several papers on buffer standards in deuterium oxide and has now investigated107 the relation between the conventional ~ L Z D scale (from cells without a liquid junction) and the operational pD scales (from cells with glass and calomel electrodes with a liquid junction).To obtain readings on the pD scale when the meter has been standardised with an aqueous buffer solution, 0-45 (molal basis) must be added to the observed meter reading for solutions of buffers from acids of the type HA in deuterium oxide. A proposal for a universal pH scale for solutions at different tempera- tures and in different solvents has been made by de Ligny and Alfenaar.lo8 Precise potentiometric measurements of acidity constants and their temperature dependence have been made for the dibasic piperazinium ion,l09 and for alkylammoniumllOu and hydroxysubstituted alkylammonium ions.llob A spectrophotometric technique has been usedlll to study the dissociation of a large number of substituted phenols and anilines over a temperature range.The references cited contain references to earlier work. Measurements of the acidity constants of molecules in electronically excited states continue; e.g. for the lowest x-x* state of aromatic carboxylic acids the acidity constant is T. H. Lilley, Trans. Faraday SOC., 1968, 64, 2947. Io5 (a) H. E. Wirth and W. L. Mills, J . Chem. and Eng. Data, 1968, 13, 102; (6) H. E. 106 H. B. Hetzer, R. A. Robinson, and R. G. Bates, Analyt. Chem., 1968,40,634. lo7 A. K. Covington, M. Paabo, R. A. Robinson, and R. G. Bates, Analyt. Chem., lo8 C. L. de Ligny and M. Alfenaar, Rec. Trao. chim., 1967, 86, 1182, 1185. log H. B. Hetzer, R. A. Robinson, and R. G. Bates, J . Phys. Chem., 1968,72, 2081. 110 (a) M. C. Cox, D. H. Everett, D. A. Landsman, and R.J. Munn, J. Chem. SOC. (B), 111 P. D. Bolton and F. M. Hall, Austral. J . Chem., 1968,21, 939; P. D. Bolton, F. M. Wirth and A. LoSurdo, ibid., 226. 1968, 40, 700. 1968, 1373; (b) B. A. Timimi and D. H. Everett, ibid., 1380. Hall, and J. Kudrynski, ibid., 1541.144 A. D. Pethybsidge and J. E, Prire decreased by a factor of about 104 compared with the ground-state value.ll3 An extensive review of titrations and acid-base equilibria in non-aqueous solvents has been published,l13 and several accounts of measurements with dimethyl sulphoxide have been reported.l14 An important and scholarly review of the equilibrium properties of acids and bases in amphiprotic mixed solvents has appeared.82a Relevant to the state of solvation of the proton in methanol-water mixtures are studies115n by mass spectrometry on the com- petitive solvation of the hydrogen ion in the gas phase.In mixtures of methanol and water vapour the hydrogen ion is solvated preferentially by MeOH in small clusters and by H2O in the larger ones. For example, when the vapour phase contains only 5 % MeOH, the clusters L4Hf, L5Hf, and L6Hf, where L is ligand, contain 80,65, and 55 % of MeOH respectively. The trend towards preferential solvation by water in large clusters suggests that the same will occur in liquid mixtures, although Wellsllsb does not agree with this interpre- tation. Electrodes.-Two very interesting reviews have been published. ‘Bio- electrodes’l16 consists of a series of articles on various applications of potentio- nietry to biological systems.Many articles are of interest to a physical chemist, particularly those on silver-silver chloride, glass, and ion-selective electrodes. More detailed information on ion-selective electrodes is included in a review by Toren.l17 These electrodes are based on a ‘membrane’ of some kind which may be of glass, ion-exchange resin (solid or liquid), a precipitate embedded in a matrix, or a single crystal. Ideally the membrane selectively transports ions of a single kind, and the successful development of such electrodes will obviously vastly increase the range of potentiometric studies. Much recent work shows that electrodes of this type show a theoretical response over a limited concentration range and are often subject to specific interference by other ions.One hopes that manufacturers will be able to resist the temptation to make excessive claims for their products in advertising literature. One of the best ion-selective electrodes is a single, doped crystal of lanthanum fluoride cemented into a tube containing a reference electrolyte and electrode. Several workers have shown that this electrode shows a theoretical response to fluoride ions down to a concentration of 10-4~. The electrode has been used to rein- vestigate the acidity constant of hydrofluoric acid and the stability constant of the complex HF:,ll8 and to study the fluoride complexes of t i n ( ~ ~ ) . l l ~ 112 E. Vander Donckt and G. Porter, Trans. Faraday SOC., 1968, 64, 3215. 113 G. A. Harlow and D. H. Morman, Analyt. Chem., I968,40,418R. 114 I .M. Kolthoff, M. K. Chantooni, jun., and S. Bhowmik, J . Amer. Chem. SOC., 1968, 90, 23; C. D. Ritchie and R. E. Uschold, ibid., 2821; I. M. Kolthoff and M. K. Chantooni, jun., ibid., 5961. 115 (a) P. Kebarle, R. N. Haynes, and J. G. Collins, J . Amer. Chem. SOC., 1967, 89, 5753; (b) C. F. Wells, ref. 1 , p. 224. 117 E. C. Toren, Analyt. Chem., 1968, 40, 402R. I l 8 K. Srinivasan and G. A. Rechnitz, Analyt. Cliem., 1968, 40, 509; N. E. Vander- Ann. New York Acad. Sci. (‘Bioelectrodes’), 1968, 148, 1-287. borgh, Talanta, 1968, 15, 1009. F. M. Hall and S. J. Slater, Ausfrul. J . Clienr., 1968, 21, 2663.Electrolyte Solutions 145 Electrodes reversible to calcium ions have been used120 in conjunction with a calomel electrode to measure activity coefficients of calcium chloride in mixtures with sodium and magnesium chlorides. The specificity of the elec- trodes breaks down if the concentration of sodium ions rises above 10-2~, but in the absence of other cations the electrodes show a theoretical response over a wide range of concentration.Amalgam electrodes have also been widely investigated. Feakins and co-workers81a discuss the best technique for making these electrodes and measuring their potentials (a digital voltmeter and automatic data recording cquipment were used). The use of lithium,121 s ~ d i u i n , ~ ~ and calcium12z amalgam electrodes in aqueous media has been described. A lead amalgam electrode has been shown123 to respond reversibly to changes in the hydroxide ion concentration in solutions of high pH. A useful review of electrode potentials in non-aqueous solvents has been p~b1ished.l~~ The behaviour of the lithium electrode in propylene carbonate125 and dimethyl sulphoxide126 has been studied.Smyrl and Tobias127 discuss the effect of diffusion of a sparingly soluble salt on the e.m.f. of cells without transference. This effect is more important for electrodes of the same kind, e.g. Ag-AgC1, in non-aqueous solutions, than in aqueous solutions because the ‘insoluble’ salts are usually more soluble than in water. The Ag-AgC1,128a Ag-AgBr,128b and Cd-CdCl2128C electrodes behave reversibly in formamide. Tetra-alkylaxnmonium Salts.-It has been known for forty years that dilute aqueous solutions of these salts show abnormally large deviations from Debye-Huckel behaviour.Recent work (see, for example, several contribu- tions and references therein in ref. 1) has been partly stimulated by interest in the fashionable concepts of ‘hydrophobic bonding’ and ‘hydrophobic hydration’. It is convenient to discuss these two topics separately, as the first is concerned with solute-solute interaction and the second with solute-solvent interaction alone. ‘Hydrophobic bonding’ or better ‘hydrophobic interaction’ describes the tendency of non-polar groups in the same or different molecules to associate in aqueous solution, thereby reducing the interactions between solute and water.12g The attraction is often considered exceptional because it is charac- terized by AH0 > 0 and T A P > 0. In fact the same is true of ion-pair 120 A. Shatkay, J .Phys. Chem., 1967, 71, 3858. lZ1 R. Huston and J. N. Butler, J . Phys. Chem., 1968, 72,4263. 123 G . Schorsch and N. Ingri, Acta Chem. Scand., 1967,21,2727. lz4 H. Strehlow in, ‘The Chemistry of Non-aqueous solvents,’ ed. J. J. Lagowski, 125 B. Burrows and R. Jasinski, J. Electrochem. SOC., 1968,115, 365. lZ6 D. R. Cogley and J. N. Butler, J . Phys. Chem., 1968,72, 1017. 127 W. H. Smyrl and C. W. Tobias, Electrochim. Acta, 1968, 13, 1581. 12* (a) R. W. C . Broadbank, S. Dhabanandana, K. W. Morcom, and B. L. Muju, Trans. Faraday SOC., 1968, 64, 3311; (b) K. W. Morcom and N. L. Muju, Nature, 1968, 217, 1046; (c) R. W. C . Broadbank, B. L. Muju, and K. W. Morcom, Trans. Faraday SOC., 1968, 64, 3318. J. N. Butler, J . Electroanalyt. Chem. Interfacial Electrochem., 1968, 17, 309.Academic Press, New York, London, 1966, vol. 1, p. 129. G. Ndmethy, Angew. Chem., 1967, 6 , 195.146 A. D. Petlij*bsiclgc and J . E. Priie formation in aqueous solution solely due to coulombic attraction.130 The implication in both cases is that ‘bond’ formation releases solvent from constraint. Hydrophobic interaction between two oppositely charged ions will clearly be enhanced by coulombic attraction and could result in associa- tion between ions unlikely to associate for electrostatic reasons alone; indeed, the larger the ions and the smaller their charge, the stronger will be the hydro- phobic interaction. A clear qualitative discussion of the origin and role of hydrophobic attraction in leading to ion pairing was written by Diamond in 1963.131 With the tetra-alkylammonium halides there is strong quantitative evidence from the variation with concentration of conductivities,46 osmotic coefficient^,^^ and apparent molar volumes132 for ion association which increases with size of cation or anion.Such an effect can, if sufficiently small, be formally described by an abnormally small ion-size pararneter,l33 or in the case of osmotic coefficients by a negative specific interaction ~oefficient.l3~ A of the effect of sodium halides on the solubility of tetrabutyl- ammonium perchlorate reveals the same pattern of behaviour. Orthodox members of the ‘hydrophobic hydration’ school object to the concept of ion pairing between (say) tetrabutylammonium and iodide ions, because these two ions belong in different categories, The former is a ‘structure-maker’, whilst ions such as the latter with small surface charge densities but not con- taining organic groups are characterized as ‘structure breakers’ (for a defini- tion of terms see ref.59). The ‘hydrophobic hydration’ schooP36 prefer to invoke, although usually without quantitative discussion, pairs between like- charged cations. However, even in dilute solutions the colligative behaviour shows a marked specific dependence on the anion.134 As the concentration rises higher aggregates involving both cations and anions may become important and eventually hydrophobic interaction will induce micelle forma- t i ~ n . l ~ ~ An attempt to clarify the situation by relaxation studies137a of solu- tions of tetra-alkylammonium salts may be complicated by relaxation involving the alkyl chains alone.137* As would be expected, hydrophobic interaction between tetra-alkylammoniuni ions and solutes such as benzene138 or naphthalene139 increases the solubility of such solutes in water. We turn now to ‘hydrophobic hydration’ which is much more closely associated with controversial views about solvent ‘structure’ and in particular I3O J. E.Prue, J . Chem. Educ., 1969, 46, 12. 131 R. M. Diamond, J. Phys. Chem., 1963, 67, 2513. 132 H. E. Wirth, J . Phys. Chem., 1967, 71, 2922. 133 R. L. Kay and D. F. Evans, J . Phys. Chem., 1966,70, 366. 13* J. E. Prue, A. J. Read, and G. Romeo, ref. 1 , p. 155. 135 J. Steigman and J. Dobrow, J . Phys. Chem., 1968, 72, 3424. 136 H. S. Frank in, ‘Chemical Physics of Tonic Solutions,’ eds. B.E. Conway and R . G. 13’ (a) G . Atkinson, R. Garnsey, and M. J. Tait, ref. 1 , p. 161; (6) M. J. Blandamer, 13* (a) H. E. Wirth and A. LoSurdo, J . Phys. Chem., 1968,72,751; (b) J. E. Desnoyers, 139 J. E. Gordon and R. L. Thorne, J . Phys. Chem., 1967, 71, 4390. Barradas, Wiley, New York, 1966, p. 59. M. J. Foster, N. J. Hidden, and M. C. R . Symons, Trans. Faraday Sac., 1968,64, 3247. G. E. Pelletier, and C. Jolicoeur, Canad. J . Chem., 1965,43, 3232.Electrolyte Solutions 147 the structure of water. It seems140 that phrases such as ‘hydrophobic structure making’ refer particularly to three effects : firstly, an increase in the molecular reorientation time of water molecules in the neighbourhood of a non-polar group, evidence for which is provided by n.m.r. measurements, although this has recently been questioned141 as the result of some measurements on di- quaternary ammonium bromides : secondly, an effect on the transport properties (see p. 009) of large organic ions which puts them into a category different from inorganic ions of either small or large surface charge density: thirdly, the hydrophobic effect is thermodynamically characterised1lo3 140 by a decrease in entropy and a large increase in the heat capacity when alkyl chains are inserted into water. It is readily apprehensible that on the introduction of a non-wetting surface into water, in order to reduce the electric field extending into the non-polar environment, the water molecules immediately adjacent to the non-polar surface should become relatively rigidly orientated and hydrogen-bonded to water molecules in the next shell. The larger the surface of the solute the greater will be the effect which may well reduce both the re-orientation time of the water molecules involved and the solute mobility. What is less evident, however, is that an extensive volume of solution around a tetra-alkylammon- ium ion is converted to some kind of ‘iceberg’ or ‘flickering cluster’ (no estimate seems to have been hazarded as to the number of water molecules in such an ‘iceberg’), or what relationship, if any, such regions have to crystalline clathrates of tetra-alkylammonium salts. Too many papers are insufficiently concerned with deciding the minimum number of adjustable parameters consistent with interpretation of a particular set of results, and too often content to conclude that all is in excellent accord with ill-defined notions of the ‘structure’ of water. As Atkin~onl~~ has stressed, in the absence of quantitative theory, ‘structural’ effects can too easily become a repository of collective ignorance. An interesting observation is that tetra-alkylammonium salts, like other salts, depress the temperature of maximum density of water.143 From the discussion in the paper, it is clear that the authors had expected the ‘structure-making’ tetra-alkylammonium salts would have the opposite effect to ‘structure-breaking’ salts because of a tendency to stabilise strongly hydro- gen-bonded, more ‘structured’ and less dense regions. The explanation that this would happen at 25” but not at 4”c because there are insufficient free water-molecules to form ‘flickering clusters’ at the lower temperature, is, to say the least, not particularly convincing. 140 Ref. 1, p. 221-233. 141 D. D. Eley and M. J. Hey, Trans. Farachy SOC., 1968,64, 1990. 142 G. Atkinson, ref. I , p. 269. 143 A. J. Darnell and J. Greyson, J . Phys. Chenr., 1968, 72, 3021.

ISSN:0069-3022    

DOI:10.1039/GR9686500129

出版商:RSC    

年代:1968

数据来源: RSC

摘要:

8 THE KINETICS OF REACTIONS IN SOLUTION By D. M. Goodall (Uniuersity of York) Introduction.-THE papers presented at the Nobcl Symposiuni, 1967, in honour of Professors Norrish, Porter, and Eigen, have now been published;1 this reviewer has no hesitation in placing them first among the 1968 references. Norrish on ‘Kinetic Spectroscopy,’ Porter on ‘Flash photolysis and primary processes in the excited state,’ and Eigen on ‘Kinetics of reaction control and information transfer in enzymes and nucleic acids’ are all eminently readable, and lavishly illustrated. Reviews by H. Eyring, Chance, Dainton, and Witt are included and an excellent picture of the foremost problems facing kineti- cists at the end of 1967, with the ideas and instrumentation being developed to meet them, is presented.This report is divided into sections as follows : 1. General theory 2. Diffusion control in photochemical and pulse-radiolytic studies 3. Solvent and ionic charge effects 4, Proton- and hydrogen atom-transfer reactions 5. Kinetic isotope effects 6 . Electrochemical studies of solution reaction kinetics 7. Fast-reaction techniques applied to biochemical problems The uses of kinetics in deduction of organic and inorganic reaction niech- anisms are reviewed in other articles in this Annual Report. Readers un- familiar with the rapid-reaction techniques which are mentioned frequently are referred to the excellent monograph by Caldin,2 and the Interscience compilation ‘Investigation of rates and mechanisms of reactions.3 This review takes up many of the topics considered in the previous Annual Report by Crooks4 and Challi~,~ and the excellent survey of the 1965-1967 literature on fast reactions in solution by Eyring and Bennion,6 which also has a section on biochemical applications.1. General Theory.-Klopman7 sets out to marry organic and inorganic chemistry in a generalised treatment of chemical reactivity; his semi-quantita- tive theory gives simultaneous theoretical foundation to the hard and soft acid ‘Fast Reactions and Primary Processes in Chemical Kinetics,’ ed. S. Claesson, E. F. CaIdin, ‘Fast Reactions in Solution,’ Oxford University Press, Oxford, 1964. ‘Investigation of Rates and Mechanisms of Reactions, Part 2,’ ed. S. L. Friess, J. E. Crooks, Ann. Report Chem. SOC. ( A ) , 1967, 64, 37. B. C. Challis, Ann.Report Chein. Soc. (B), 1967, 64, 99. E. M. Eyring and B. C. Bennion, Ann. Rev. Pliys. Clietn., 1968, 19, 129. G. Klopman, J . Amer. Chern. Soc., 1968, 90, 223. Interscience, New York, 1967. E. S. Lewis, and A. Weissberger, Interscience, New York, 1963. 149150 D. M. Goodall and base concept, to nucleophilic order, and to other reactivity indices. The interactions between a Lewis acid and Lewis base plus their solvating shells are subdivided into : (1) coulomb, (2) electron transfer and delocalisation, and (3) Born solvation effects. A perturbation treatment of the interaction energies suggests classification of reactions as charge- or frontier-controlled [( l), or (2) and (3) predominate respectively]; and reactivity orders can be computed with a knowledge of fundamental atomic and molecular parameters such as ionisation potential, electron affinity, atomic orbital coefficients, charges, and solvent dielectric constant. Dixon8 gives a simple molecular-orbita, method for predicting the direction of alternative proton-abstraction reactions in x-electron systems, based, via a discussion on non-bonding orbital coeffic- ients, on n.m.r.coupling constants between these protons and the remaining atoms in the molecule. The vast accumulation of data on substituent effects in the rates and equilibria of organic reactions has been rationalised by Swain and Lupton:9 they find from statistical analysis of 43 reaction series that any remote substituent constant can be written as: G = fF+ rR where Fand R are field and resonance constants specific for a given substituent and f and r are empiricial weighting factors independent of substituent but different for each reaction series.All parameters are at present given no theoretical content; it would lend weight to the treatment if a correlation with fundamental molecular parameters such as inter-group bond distances and charges could be made (cf. Dewar and Grisedale, ref. 10). A Huckel molec- ular-orbital study suggests that part of the efficiency of bifunctional catalysis, in systems such as the mutarotation of glucose in the presence of pyridine and phenol, arises from delocalization of the transferring charge over the x- electron systems of the two catalysts.ll An important paper by Schwarz12 builds the theoretical framework of chemical relaxation spectroscopy, and gives a general solution to the relaxa- tion problem.Evaluation of complex spectra with many interacting rate constants is best done in terms of the experimentally accessible mean reciprocal relaxation time, and the dispersion about the mean; these quantities can be related to rate constants of individual steps. Applications to the conforma- tional changes of biopolymers are considered later in this report. In another theoretical treatment of relaxation spectra, Hayman13 enquires whether it is possible to distinguish a charge-transfer reaction between aromatics to give a dimer, AB, from a series of steps which utilise both sides of the molecules to give sandwiches ABA, BABABA etc. When all association and dissociation 8 W.T. Dixon, Tetrahedron, 1968, 24, 5509. C. G. Swain and E. C. Lupton, J . Amer. Chem. SOC., 1968,90,4328. lo M. J. S. Dewar and P. J. Grisedale, J. Amer. Chem. SOC., 1962, 84, 3548. l1 H. J. Gold, J . Amer. Chem. SOC., 1968, 90, 3402. l2 G. Schwarz, Rev. Mod. Phys., 1968,40,206. l3 H. J. G. Heyman, J . Chem. Phys., 1968, 48, 3273.The Kinetics of Reactions in Solution 151 rate constants are degenerate, one relaxation time is predicted for both systems, but distinction is in theory possible if this degeneracy is removed (c.J ref. 12). Chal4 has given a simplified treatment for deriving the rate equations in complex interconnected systems where several blocks of reactants are at equilibrium. Shear15 has presented a mathematical proof that a homogeneous chemical reaction system has one, and only one, equilibrium position.Following partial disproval of the classic oscillating chemical reaction between iodate and peroxide, distinction between homogeneous and heterogeneous systems must be made carefully.16 More studies of self-oscillations in open systems involving enzymes have been made,17 and several authors suggest testing of theories with constant source and sink concentrations. Swinkels and Wojcie- chowski describe in detail a matrix approach to the kinetics of such open systems.l* Two papers elaborate the theoretical background for modulation of laser light scattered from a liquid undergoing both thermal diffusion and chemical reaction.19 From measurements of the widths of the central Rayleigh linesz0 (split according to the number of reactions) as a function of scattering angle, rate constants may in principle be evaluated.This technique has yet to be experimentally tested, but holds much promise. The relative merits for kinetics of n.m.r. lineshape analysis and the newer spin-echo technique have been discussed by ReevqZ1 following careful experiments on the determination of hindered rotation rates in amides. 2. Diffusion Control in Photochemical and Pulse Radiolytic Studies.-Using a statistical treatment, Steinberg and Katchalskiz2 have given a new derivation of the diffusion-control kinetic equations; this includes the Forster equation for non-radiative energy transfer, and also accounts for the case where diffusion alters the distance between donor and acceptor during the energy- transfer process.The theory is successfully testedz3 using the excited naphthal- ene (lifetime sec.)-anthranilic acid system in solvents over a viscosity and diffusion coeficient range of lo3, where the Forster equation is inadequate. Wagner and K ~ c h e v a r ~ ~ also find that a simple inverse viscosity dependence for triplet energy transfer probability does not hold for low-viscosity solvents and slow energy transfer, and like Katchalski recognise that at high quencher l4 S. Cha, J . Bio?. Chem., 1968, 243, 820. l5 D. B. Shear, J . Chem. Phys., 1968,48,4144. D. H. Shaw and H. 0. Pritchard, J . Phys. Chem., 1968,72,1403,2693; H. Degn and l7 H. Degn, Nature, 1968,217, 1047; E. E. Sel’kov, European J . Biochem., 1968,4, 79. 18 G. M. Swinkels and B.W. Wojciechowski, Trans. Faraday SOC., 1968, 64, 43. l9 L. Blum and Z . W. Salsburg, J . Chem. Phys., 1968,48,2292; B. J. Berne and H. L. 20 B. Chu, J . Chem. Educ., 1968,45, 224. 21 P. T. Inglefield, E. Krakower, L. W. Reeves, and R. Stewart, Mol. Phys., 1968,15, 22 I. Z . Steinberg and E. Katchalski, J . Chem. Phys., 1968, 48, 2404. 23 Y. Elkana, J. Feitelson, and E. Katchalski, J . Chem. Phys., 1968, 48, 2399. 24 P. J. Wagner and I. Kochevar, J . Amer. Chem. SOC., 1968,90,2232. J. Higgins, J . Phys. Chem., 1968, 72, 2692. Frisch, J. Chem. Phys., 1967,47,3675; ibid., 1968,49,2864. 65.152 D. M. Goodall concentrations the energy-transfer distance is satisfied without any need for diflusion. Protic solvents seem to give an abnormal increase in fluorescence quenching rates with certain aromatic donors containing phenoxide groups, and Umberger25 propounds solvent assistance by proton transfer during the energy donation process.Intermittent photochemical generation of alkyl radicals using rotating sector methods has been used to find rate constants for the reaction of radical pairs in solution. Burkhart26 varied the sector rotation speed and analysed for products after each run, whereas Weiner and H a m r n ~ n d ~ ~ followed the decay of the radical transients in the millisecond dark intervals by use of e.s.r. Both find rate constants in the range 1-4 x lo9 1. mole-l sec.-l, considerably higher than polymer chain termination rates, and in satisfactory agreement with diffusion control theory, when a correction for the three in four anti-bonding triplet collisions is taken into account.The ‘solvent cage’ effect and the primary recombination of photochemically produced radical pairs must always be taken into account in photochemical diffusion control studies. Analogously, some of the most interesting papers on pulse radiolysis are concerned with the search for spurs, i.e. groupings of reactive intermediates, generated by the 7-ray pulse close enough together for there to be a significant probability that they will react with one another. Freeman’s2* recent theoretical analysis of this kinetic problem, predicting a fast, followed by a slow, phase (combination of oppositely charged ions within the spur, then the homogeneous reaction of these ions) has been substantiated by several studies using non-aqueous solvent^.^^^ 30 The two investigations by Thomas and his co-w~rkers~~ use nanosecond pulses so that the fast phase of the decay is not masked. There are now two more theoretical treatments of recombination rates for a pair of oppositely charged ions generated within a spur: Mo~umder~~ uses random walk of the ions in their mutual coulomb field to compute half-lives of neutralisation from a given distance as a function of viscosity, temperature, and other variables, whilst Williams32 concludes that the Nernst-Einstein equation is adequate for a description of recombina- tion from typical initial separation values of 40-120 A.Klein et aZ.33 looked for spurs after 50 nsec. pulses into aqueous solutions of various pH values and ionic strengths, but found no evidence of non-homogeneous distribution of the hydrated electrons in their time frame; an initial fast decay of the hy- J.Q. Umberger, J . Phys. Chem., 1968, 72, 1350. ZG R. D. Burkhart, J . Amer. Chern. SOC., 1968, 90, 273. 27 S. Weiner and G. S. Hammond, J . Amer. Chem. SOC., 1968, 90, 1659. 28 G. R . Freeman, J . Clietn. Phys., 1968,46,2822. 29 K. N. Jha and G. R. Freeman, J . Chern. Phys., 1968, 48, 5480; J. C. Russell and 30 J. K. Thomas, K. Johnson, T. Klippert, and R. Lowers, J . Chern. Phys., 1968, 48, 31 A. Mozumder, J . Chem. Phys., 1968,48, 1659. 32 F. Williams, J . Chem. Phys., 1968, 48, 4077. 33 N. Klein, C. N. Trumbore, J. E. Fanning, and J. W. Warner, J . Phys. Chem., 1968, G. R . Freeman, ibid., p. 90. 1602; R. Cooper and J. K.Thomas, J . Chem. Phys., 1968,48,5103. 72, 880.The Kinetics of Reactions in Sohition 153 drated electron in air-free alkali was attributed to a reaction with excited water. HamilP4 postulates that the hydration of a pulse-generated electron takes sec. in water, and that differences in reactivity of the free and hydrated electrons could account for anomalies in the radiation chemistry of water, usually attributed to spurs ; for instance, solutes in concentrations greater than 1 0 - 2 ~ will scavenge all free electrons before hydration. Dainton’s review, ‘The chemistry of the electr0n,’~5 gives a concise rate constant summary and discussion of the role of diffusion control in reactions of the hydrated electron. The usual activation energy in these reactions is ca.3.5 kcal. mole-1, the diffusional activation energy of the electron; Cer~ek3~ has recently reported measurements of lower activation energies in its reac- tions with the hydronium and nitrite ions, and prefers to discuss these as electron-migration reactions, with anupper limit of 2.3 kcal. mole-l calculated from the known equivalent conductance of the electron and absolute reaction- rate theory 3. Solvent and Ionic Charge Effects.-Parker and his c o - ~ o r k e r s ~ ~ continue to amass data about solvent effects on solubilities and kinetics in a variety of protic and dipolar aprotic solvents. They have analysed the rates of 78 bi molecular nucleophilic substitution and bimolecular elimination react ions in various solvents, and the free energies of transfer of reactants from these solvents to methanol, and make deductions about transfer free energies and solvation of the transition states.Their short cuts in assessing transfer free energies of individual ions by extra-thermodynamic assumptions (e.g. by allotting equal contributions to the tetraphenylboride and tetraphenylar- sonium ions) are rightly criticised by Rodewald, Mahendran, Bear, and Fuchs;38 these workers measure enthalpies of transfer of alkali halides, and give an alternative explanation for some of the S,2 solvent effects. More superficially unrewarding work on the thermodynamics of transfer between solvents is required for a resolution of these problems. K o s ~ w e r ~ ~ has observed a striking solvent effect in the reaction between the 1 -ethyl-4-methoxycarbonylpyridyl radical and 4-nitrobenzoyl chloride.The rate increases by a factor of 104 in changing solvent from 2-methyltetrahydro- furan to dimethylformamide or acetonitrile, and this indicates a charge- transfer transition state ; other benzoyl chlorides react via an atom-transfer transition state. Solvent effects are discussed in terms of Kosower’s 2 value model,40 assuming equilibrium solvation of the transition state. An interesting study of the competing reactions of disproportionation and 34 W. H. Hamill, J . Ctiem. Phys., 1968, 49, 2446. F. S. Dainton, ref. 1 , p. 185. 36 B. Cercek and M. Ebert, J . Phys. Chetn., 1968, 72, 766; B. Cercek, ibid., p. 2279. 37 R. Alexander, E. C. F. KO, J. Parker, and T. J. Broxton, J . Amer. Chem. SOC., 38 R.F. Rodewald, K . Mahendran, J. L. Bear, and R. Fuchs, J . Amer. Cliem. SOC., 39 E. M. Kosower and M. Mohammad, J . Amer. Chem. SOC., 1968,90,3271. 40 E. M. Kosower, ‘An Introduction to Physical Organic Chemistry,’ Wiley, New 1968, 90, 5049. 1968,90, 6698. York, 1968.154 D. M. Goodall recombination of a pair of ethyl radicals has been madc, both in the gas phase and in various non-polar solvents. kc kd c4H10 +-- 2c2H.!je+ C2M4 + CZH6 The logarithm of the rate ratio is a linear function of the square root of the internal pressure of the solvent, and Stefani41 gives a number of other examples which suggest that Hildebrand’s theory is as successful as dielectric constant correlations for a number of electroneutral reactions. Measurements of activation volumes are uniquely suited to the study of solvation changes during a reaction.Browef12 has considerably extended the range of application of this technique by describing a pressure jump apparatus with optical monitoring of equilibrium shifts : positive or negative pressure jumps may be made in 5 msec. from any pressure in the range 50-1400 atm. H i r ~ t a ~ ~ has used e.s.r. to distinguish different degrees of solvation of ion pairs in ethereal solvents. The extent of splitting of the e.s.r. spectrum of the naphthalenide radical anion by its alkali counter-ion depends on the separation between the two, and, in the presence of naphthalene, electron exchange was shown to proceed only via the loosely bound ion pairs; the rate constants in the pre-equilibrium formation of this from the tight ion pair have been evalu- ated.In liquid ammonia, all proton exchange from the ammonium ion occurs via the free ion rather than ion pairs, and the n.m.r. data can be used to derive ion-pair formation constants as well as exchange rates.44 Gigantic ‘primary salt effects’ have been found in a study of anion-anion reactions in the presence of cationic polyelectrolytes,45 and cation-cation reactions with anionic macro-ions. There is, presumably, a high concentration of counter-ions in the neighbourhood of the polyelectrolyte molecule, and electrostatic repulsion between the like ions are reduced. In fact the rate enhancement arises predominantly from a lowering in the activation enthalpy term, as is also observed in the more complex systems where micelles are used to catalyse reactions of their counter-i0ns.~6* 47 Interestingly, one particular reaction between negative ions is not accelerated by use of a cationic micelle;46 if the reactants bind at the same type of site, competitive binding could occur, and neighbouring sites might be too far apart to promote reaction.4. Proton- and Hydrogen Atom-transfer Reactions.-Marcus4s has extended his successful theory for weak- overlap electron transfers to cover proton- and atom-transfers, and strong- overlap electron transfers. The free energy of 41 A. P. Stefani, J . Amer. Chem. SOC., 1968, 90, 1694. 42 K. R. Brower, J . Amer. Chem. SOC., 1968, 90, 5401. 43 N. Hirota, J . Amer. Chem. SOC., 1968, 90, 3603; N. Hirota, R. Carraway, and 44 D. R. Clutter and T. J. Swift, J .Amer. Chem. SOC., 1968, 90, 601. 45 N. Ise and F. Matsui, J . Amer. Chem. SOC., 1968, 90, 4242. 46 C. A. Bunton, E. J. Fendler, L. Sepulveda, and K-U Yang, J . Amer. Chem. SOC., 1968,90, 5512. *? R. B. Dunlap and E. H. Cordes, J . Amer. Chem. SOC., 1968,90,4395; L. R. Romsted and E. H. Cordes, J. Amer. Chem. SOC., 1968,90,4404. 48 R. A. Marcus, J . P/JYS. Chem., 1968, 72, 891 W. Schook, J . Amer. Chem. SOC., 1968,90, 3611.The Kinetics of Reactions in Solution 155 activation for the reaction of AH and B- may be calculated if the energy difference between reactants and products is known, and the rates of the symmetrical exchanges AH + A-, BH + B- have been measured. Predic- tions are made about the magnitudes of the Br~lnsted exponents and kinetic isotope effects and it will be of interest to obtain magnetic resonance data for symmetrical transfers and test these relations.A provocative paper by Pshenichnov and Sok01ov~~ suggests that a reaction with rate-deter- mining proton transfer be treated as a process of consecutive vibrational excitation steps within a hydrogen-bonded complex of donor acid and acceptor base, of lifetime 10-10 sec. Energy transfer would occur by resonant interaction between the randomly tumbling solvent and the complex, and solvents with a vibration frequency near to the A-H fundamental are predicted to be particularly effective. Ritchie and Kingso present detailed LCAO-MO-SCF calculations on potential energy surfaces for proton abstraction from the series HH, HF, HzO, H3N, H4C by the hydride ion. With the exception of the first reaction, shallow energy minima along the reaction co-ordinate are found, rather than activation energy maxima.In the absence of gas-phase data, Ritchie and King add solvation-energy correction factors and use their, by now approximate, calculations to give a rather unsatisfactory account of some solution results. Current problems in the kinetics of proton-transfer reactions in solution have also been discussed by Khristov.51 Ahrens and ma as^^^ have extended Eigen’s survey53 of proton-transfer rates to include S-H . . . N and S-H . . . S transfers. The general trend of Bramsted-plot curvature is still present, but at large ApK values, rates for S-H . . . S transfer level off below the diffusion control limit, and individual deviations from a smooth plot are present, as in proton transfer from carbon acids.An analogous Brcansted plot over the range ApK = - 10 to + 14 has been constructed for proton transfers between fluorenes and bases in dimethyl- sulphoxide solution54. Ritchie and Uschold obtained their latest data from stop-flow studies using benzoate ions as bases. The diffusion control rates are attained within these ApK limits, but diffusion control cannot be obtained when methanol is considered as solvent; for this solvent, considerable free energy is required for de-solvation of the reactants in the pre-equilibrium which precedes proton transfer. Recent papers from the Grunwald group have also highlighted the differences in proton transfers which occur uia a ‘solvent bridge’ Grotthus-type mechanism, and those which necessitate direct contact between acid and base molecules. The transfer from the imidazolium ion to imidazole is diffusion controlled, and has a large reaction diameter: 49 E.A. Pshenichnov and N. D. Sokolov, Znternat. J. Quantum Chem., 1967, 1 , 855. C. D. Ritchie and H. F. King, J. Amer. Chem. SOC., 1968,90, 825, 833, 838. 51 S. G. Khristov, Zhur.fiz. Khim., 1968, 42, 1553. 5 2 M. L. Ahrens and G. Maass, Angew. Chem. Znternat. Edn., 1968, 7 , 818. 53 M. Eigen, ref. 1, p. 245; Angew. Chem. Internat. Edn., 1964, 3, 1. 54 C. D. Ritchie and R. E. UschoId, J. Amer. Chem. Soc., 1968,90, 3415. E. K. Ralph and E. Grunwald, J . Amer. Chem. SOC., 1968,90, 517.156 D. M. Goodall n.ni.r. studies55 of both substrate and 170-enriched water-line broadenings show that there is a mean of 1-4 water molecules mediating in this reaction. Thus, approximately half the proton exchanges here are quadrimolecular, or of higher order.For direct proton transfers between ammonium ions and a m i n e ~ , ~ ~ exchange is relatively slow but does not conform to the Brarnsted rule when the acceptor base is varied. This leads to the conclusion that dehydration of the solvent-separated encounter complex, rather than proton transfer, is the rate-determining step. Long and GoodalF7 studied the rate of proton transfer between acetic acid and the anions of nitroalkanes in mixtures of light and heavy water; here proton transfer is rate determining, and donor and acceptor are in contact. Other studies of proton transfers to carbon bases confirm the lack of solvent bridging in such 59 Kreevoy’s review59 on ‘Developments in the study of A-SE~ reactions in aqueous solution’ starts with identification of the overall mechanism, and progresses to hypotheses on fine details of the transition state; general acid catalysis and isotope effect determinations are the most powerful aids to such studies.The classic way of demonstrating solvent involvement in solution reactions is to look for a dependence on solvent concentration in the rate equation. Bell and co-workers60 have been able to do this in a study of the reversible hydration of ketones, by using mixtures of water with high mole fractions of dioxan or acetonitrile. Their results are striking, and indicate involvement of 0, 1, 2, or 3 water molecules in the transition state according to the catalyst present and whether hydration or dehydration is being studied.Evidence points to a stepwise mechanism with no diffusion apart of reactants: for instance, one water molecule is involved in hydration catalysed by benzoic acid. I1 ir A powerful argument for a stepwise proton-transfer sequence is that a fully concerted mechanism cannot be written for a monofunctional catalyst such as triethylamine. A simple electrostatic model predicts that a stepwise sequence has lower activation energy than a concerted one, but the concerted scheme still has proponents.49, 61 Close analogies can be drawn between this 56 E. Grunwald and A. Y . Ku, J. Amer. Chem. SOC., 1968,90, 29. 57 D. M. Goodall and F.A. Long, J . Amer. Chem. SOC., 1968,90,238. 58 V. Gold and D. C. A. Waterman, J . Chem. SOC. (B), 1968,839,849. 59 J. M. Williams and M. M. Kreevoy, Ado. Phys. Org. Chem., 1968, 6, 63. 6O R. P. Bell, J. P. Millington, and J. M. Pink, Proc. Roy. SOC., 1968, A , 303, 1. 61 W. J. Albery, Progr. Reaction Kinetics, 1967, 4, 353.The Kiiietics of Reactions it1 Solution 157 reaction and the reversible addition of hydrogen peroxide to aldehydes, studied by Jencks: this system shows the interesting feature of a Brransted exponent, a , equal to unity during acid catalysis.62 Swift63 has analysed lH n.m.r. lineshape data to find rate and activation parameters for proton exchange from the hydration spheres of several cations, and these suggest different push-pull mechanisms for weak (1) and strong (2) acid catalysis: in both cases the solvent functions as a base: H H H H M-0-H .. . A-H . . , 0-H -4 M-0- . . . H-A , . . H-0-H 1 1 I I (1) + H M H H M H H-9-H . . . -0-H . . .O-H- H-0 . . . H-0- . . . H-0-H + + 1 I I I I I (2) At low exchange rates the nature and number of solvent molecules bound in the primary hydration sphere may be found from peak area meas~rements.~~ Loewen~tein~~ generated radicals from alcohols, and studied their proton exchanges with the hydronium ion using e.s.r. line-broadening. The rate constants and substituent effects parallel those found with neutral alcohols. Stop-flow with e.s.r. detection was used to follow the hydrogen atom exchange between 2,4,6-tri-t-butylphenol and its ring-deuteriated phenoxy radical in carbon tetrachloride.66 Reaction proceeds in two steps via a symmetrical intermediate [ROHOR].of lifetime sec., and the asymmetry of the actual transition state is demonstrated by the low kinetic isotope effect. Kastening67 has taken account of both the variation in e.s.r. sensitivity, and non-homogeneity in radical concentrations, within a capillary flow tube mounted in the spectrometer waveguide. The theory is then applied to find the rate constant for dismutation of the nitrobenzene negative ion, under conditions of laminar flow in the capillary. An interesting flask photolytic reaction is reported by Weller;68 at acidities where azulene is present predominantly as its protonated cation, a flash generates azulene via the primary formation of the excited azulinium ion, a very strong acid which deprotonates as it is quenched.The restoration of the Ga E. G. Sander and W. P. Jencks, J. Amer. Chem. SOC., 1968, 90, 4377. 63 T. A. Stephenson, T. J. Swift, and J. B. Spencer, J . Amer. Chem. SOC., 1968, 90, 64 R. G. Wawro and T. J. Swift, J. Amer. Chem. SOC., 1968,90,2792; N. A. Matwiyoff 65 R. Pouko and A. Loewenstein, J . Chem. SOC. (A), 1968, 949. 67 B. Kastening, Ber. Bunsengesellschaft Phys. Chem., 1968, 72, 20; B. Kastening and K. H. Grellman, E. Heilbronner, P. Seiler, and A. Weller, J . Amer. Chem. SOC., 4291. and H. Taube, J . Amer. Chem. SOC., 1968, 90, 2796. M. R. Arick and S. I. Weissman, J . Amer. Chem. SOC., 1968, 90, 1654. S. Vaviizka, ibid., 1968, 72, 27. 1968, 90, 4238.158 D. M.Goodall ground-state equilibrium after termination of the flash occurs with a rate constant known from previous flow studies.69 5. Kinetic Isotope Effects.-By use of a sufficiently wide range of donor acid and acceptor base strengths to ensure variations on either side of the sym- metric transition state, maxima in change of primary kinetic isotope effects in a reaction series are now well d~cumented.~O-~~ Kresge71 shows that a smooth correlation of isotope effect with rate constant is found in proton transfer from the hydronium ion to carbon-carbon double bonded systems only when substituents of the type -OR are considered. More diverse substitution in the olefin obscures any correlation. Studies in water suggest maxima near the point of equal pK values of donor and a~ceptor.~O In tetrahydrofuran, with phenyl-lithium as the base, the maximum occurs with donor acids of pK ca.25.72 Benzene has a pK near 36, and the divergence between these two figures is presumed to be due to the fact that the pre- dominant, ion-paired species is a considerably weaker base than the bare phenyl anion. Kreevoy has reported a primary isotope effect which remains constant through large changes in dimethylsulphoxide-water compositions, and it is assumed that the hydrogen ion proton donor remains the H30f unit.59 Competitive abstraction of hydrogen and deuterium from [c~-~H-] toluene by bromine atoms has been studied both in the gas phase73 and in solution, and it is difficult to explain alterations in the observed isotope effects, particularly when similar abstraction studies using chlorine atoms yield concordant gas phase and solution results.Caldin and Tomalin74 have studied the rates of proton and deuteron transfer between 4-nitrobenzyl cyanide (or its [2Hz]-isomer) and ethoxide ion, by a combination of low temperature and stop-flow techniques. Devia- tions from Arrhenius behaviour at the lowest temperatures are attributed to proton tunnelling, and, using the Bell model of a parabolic one-dimensional barrier, concordant results are obtained by independent analyses of the two systems. They then proceed to use the model to analyse recent data where proton tunnelling is indicated, and calculate barrier widths and heights; symmetric transition states are found to show the maximum difference in H and D barrier heights, as expected. Lewis and Robinson75 obtain tritium isotope effects to compare with deuterium isotope effects in cases where tunnelling is suspected; deviations from the Swain relation are small, and a justification is offered.The anomalous isotope effects in the reaction between 89 B. C. Challis and F. A. Long, J . Amer. Chem. SOC., 1965, 87, 1196. 70 R. P. Bell and 4. M. Goodall, Proc. Roy. SOC., 1966, A, 294,273; J. L. Longridge and F. A. Long, J . Amer. Chem. SOC., 1967,89, 1287; A. F. Cockerill, J. Chem. SOC. (B), 1967,964. 71 A. J. Kresge, D. S. Sagatys, and H. L. Chen, J. Amer. Chem. SOC., 1968,90, 4174. 72 Y. Pocker and J. H. Exner, J . Amer. Chem. SOC., 1968,90,6764. 73 R. B. Timmons, J. de Guzman, and R. E. Varnerin, J . Amer. Chem. SOC., 1968,90, 7* E.F. Caldin, M. Kasparian, and G. Tomalin, Trans. Furaday SOC., 1968,64, 2802; 5996. E. F. Caldin and G . Tomalin, ibid., p. 2814.The Kinetics of Reactions in Solution 159 2,6-lutidine and 2-nitropropane continue to haunt experimentalist^,^^ 76 and Davis and Lehman’s76 activation parameters are so widely divergent from those of Lewis that the tunnelling explanation cannot be assumed proved. Willi77 continues his important experimental and theoretical investigations of secondary hydrogen isotope effects in simple SN2 reactions of methyl iodide and [2H3] methyl iodide. Temperature dependence studies are reported, with cyanide ion as the nucleophile. From an accurate knowledge of ground- state force constants, and isotope effects in similar reactions, the best fit of the data is obtained by a set of transition-state force constants which include unchanged C-H stretching contributions.(It is normally assumed that an sp3 to sp2 change increases the stretching vibration frequency.) In this instance alterations in the bending force constant dominate the isotope effect. In the SN1 solvolysis of t-butyl chloride, the secondary isotope effect is constant in solvent mixtures which promote ionization at the same rate,78 and equal to the isotope effect calculated on the assumption of an equilibrium with t-butyl and chloride ions.79 This gives more evidence for the ion-pair nature of the transition state. In interpreting results of remote secondary isotope-effect studies, there seems to be a consensus of opinion in favour of Shiner’s hyperconjugative explanation.8O Servis, BoreiC, and SunkoS1 also suggest hyperconjugation as the reason for their observation of an excellent linear free-energy correlation between the solvolytic rate enhancement on making the change from RR’XCaH to RR’XCCH3, and the secondary isotope effect on solvolysis after the methyl hydrogens have been replaced by deuteriums.Assuming hyperconjugation as the explanation for the first effect, the extension to the second seems logical. 6. Electrochemical Studies of Solution Reaction Kinetics.-The application of electrochemical methods in reaction-rate studies is introduced by Caldin,2 and reviewed rigorously by Strehlow.g2 A 1968 review of developments in alternating current polarography also has references to work in homo- geneous kinetics.83 Most investigations are carried out at constant voltage, and for the case where reaction in solution modifies the diffusion gradient of the electro-active species, explicit solutions for the value of the average current (in a polarographic measurement) or the decay of current with time (at a fixed electrode) are available. These hold only for a first-order reaction 75 E.S. Lewis and J. K. Robinson, J . Anter. Chem. SOC., 1968, 90, 4337. 76 T. A. Lehman, Diss. Abs., 1967, 28, B, 632. 77 A. V. Willi and Chong Min Won, J . Amer. Chem. SOC., 1968,90, 5999. 78 G. J. Frizone and E. R. Thornton, J . Amer. Chem. SOC., 1968,90, 121 1. 79 J. C. Evans ana G. Y-S. Lo., J . Amer. Chem. SOC., 1966, 88,2118. 80 V. J. Shiner, W. E. Buddenbaum, B.L. Murr, and G. Lamarty, J . Amer. Chem. SOC., 1968, 90, 418; J. G. Jewett and R. P. Dunlap, J. Amer. Chem. SOC., 1968, 90, 809; D. S. Noyce and M. D. Schiavelli, J . Amer. Chem. SOC., 1968,90, 1023. 81 K. L. Servis, S. BorEik, and D. E. Sunko, Tetrahedron, 1968, 24, 1247. 83 H. Strehlow, ref. 3, p. 799. 83 W. H. Reinmuth, Analyt. Chem., 1968, 40, 185R.1 60 D. M. Goodull in solution, and Birk and PeroneS4 find that they may be used in the rather special case where the electro-reducible species are ketyl radicals generated by flash photolysis; instead of the normal constancy of concentration in the bulk solution, this concentration follows a second-order decay, and the rate constant for the radical recombination can be found from chronoampero- metric measurements.The majority of papers are concerned with the ECE inechanism, where a chemical reaction is interposed between two charge-transfer steps. These steps may occur at different electrodes, normally a combined ring-disc arrangement spinning in the solution. Theoretical and experimental con- siderations combine to make this the most tractable arrangement.85 It may be illustrated by Bruckenstein’s study of the kinetics of oxidation of arsenic(n1) by iodine in alkaline iodide solutions.86 The iodine is produced at the disc electrode by oxidation of iodide ion, at a rate controlled by the disc current, then moves by connective diffusion towards the ring, where it is rzduced back to iodide; en route it reacts with arsenic(m), and the ring current is diminished according to the rate of this reaction.As before, the method treats first-order reactions only, and reactions with half-lives in the region of lo-’ sec. com- Pete usefully with diffusion between the electrodes. The example above was a new kinetic and mechanistic study, but studies by AlberyE7 and Bruckenstein8* using well-documented reactions confirm the essential validity of the theory. Karp has recently suggested a slight modification of the Albery treatment.8g ECE Processes can be observed using a single electrode, provided both reactant and product are electro-active. One possible scheme is : *nn,e k i n a c A f B - + - C + D If C is reduced at a considerably less cathodic potential than A, the electrode potential may be stepped from an initial high value, where the reaction is proceeding smoothly from left to right, to a lower value where B is being oxidized whilst C is still being reduced.Blount and Hermango have followed the decay of current with time (and also the change in potential with time at constant current) in the sequence where B is p-hydroxyphenylhydroxylarnine and C its dehydration product. The dehydration is found to be catalysed by general acids and bases, and with this conclusion the catalysis scheme sug- gested could be improved by following the arguments of Bell.91 A more complex system is the two-electron oxidation of o-tolidine to the di-i~nine.~~ 84 J. R. Birk and S. P. Perone, Analyt. Cizem., 1968, 40, 496. 85 W. J. Albery and S. Bruckenstein, Trans. Faraday SOC., 1966,62, 1946, W. J. Albery, 86 D. C.Johnson and S. Bruckenstein, J. Amer. Chem. Soc., 1968,90, 6592. 87 W. J. Albery, M. L. Hitchman, and J. Ulstrup, Trans. Faraday SOC., 1968,64,2831. 88 P. Beran and S. Bruckenstein, J. Phys. Chem., 1968, 72, 3630. 89 S . Karp, J . Phys. Chem., 1968, 72, 1082. 90 H. N. Blount and H. B. Herman, J . Phys. Chern., 1968,72, 3006. 91 R. P. Bell, ‘The Proton in Chemistry,’ Methuen, London, 1959, ch. 9. 92 J. W. Strojek, T. Kuwana, and S. W. Feldberg, J. Amer. Chem. SOC., 1968, 90, 1353; T. Kuwana and S. W. Feldberg, Discuss. Faraday SOC., 1968,45, 134. ibid., 1967, 63, 177 1.The Kitretics of Reactions in Solutiora 161 The reaction sequence here is A f B + e : B + C + e ; A + C + 2B. The electrochemical data were supplemented by spectrophotometric concentra- tion-time studies of the reactive species, made possible by the use of a transparent doped tin oxide glass electrode.Careful analysis of the data suggests that 2B is in fact the ion pair AC, though e.s.r. demonstrates that B is a true one-electron oxidised species in some solvents. A paper by Nelson also deals with the distinction between two-electron and ECE processes.93 To acquaint more chemists with these useful electrochemical techniques, Crooks and Bulmer have introduced an undergraduate experiment on kinetics from polarograp hy . 94 7. Fast-reaction Techniques Applies to Biochemical Systems-Rapid- reaction techniques have already made great advances in our understanding of short-lived intermediates in biochemical pathways, enzyme-substrate binding processes, and conformation changes of Many advances in technique are prompted by biochemical challenges, as will be illustrated below.Eigen’s two lucid reviewsg5 describe advances in instrumentation which have gone hand-in-hand with the chemical and biochemical studies in his Institute. These include a pressure-shock wave method with 10 nsec. dead time, combined stop-flow temperature jump, and a square-wave electric-field jump apparatus with 50 nsec. dead time. This last apparatus is used to study haemoglobin relaxations, where unexplained data still remain despite the introduction of a detailed scheme of conformation changes and substrate bindings. The functional unit of haemoglobin here is a tetramer of two weakly interacting ap-dimers. Antoninig6 has been able to study the rate of dissociation of the dimer into cc- and P-monomers, displacing the equilibrium position by flash photolysis, or rapid addition of carbon monoxide in a stop- flow apparatus. Kirschner and Eigen’s studyg7 of allosteric binding to glyceraldehyde-3-phosphate dehydrogenase is a key example of the applica- tion of stop-flow and temperature-jump methods to investigate fundamental operations in enzyme control processes.The observed relaxation times correspond to the Monod model of stepwise substrate binding and ‘all-or- nothing’ conformation change, and only one of the enzyme’s two conforma- tions has catalytic activity. P0h198 has presented an elegant study of the reversible denaturation of chymotrypsin and trypsin in water and aqueous ethanol. Denaturation follows a linear Arrhenius activation energy plot, but the renaturation rate goes through a maximum with temperature.A large heat capacity of activation is indicated, and promises to yield much new infor- mation about solvation changes in protein unfolding. This work was made 93 R. F. Nelson, J. Electroanalyt. Chem. Interfacial Electrochem., 1968,18, 329. g4 J. E. Crooks and R. S. Bulmer, J . Chem. Educ., 1968,45, 725. 95 Ref. 1, p. 333; M. Eigen, Quart. Rev. Biophys., 1968, 1, 3. g6 E. Antonini, M. Brunori, and S. Anderson, J . Biol. Chem., 1968, 243, 1816. 97 K. Kirschner, in ‘Proceedings of the 4th Meeting of the Federation of European 98 F. M. Pohl, European J . Biochern., 1968, 4, 373; 1968, 7, 146. Biochemical Societies, Oslo, 1967.’162 D.M . Goodall possible by development of a simple temperature-jump apparatus ; cooling or heating is achieved with 3 sec. dead time by switching thermostatting water flows. Temperature jump and stop-flow studies by Hammesg9 compare rates of binding of aspartate (the natural substrate) and a-methylaspartate to aspartate aminotransferase, and subsequent relaxations of the enzyme- substrate complex are analysed. Schimmello0 showed that lysozyme binds di- and tri-N-acetylglucosamine with equal rates. Since X-ray data on the mono- and tri-saccharide showed one sugar residue to be identically bound in each case, and the third residue rather loosely bound,lol identity of rates is understandable. Gutfreundlo2 has studied phosphorylation and dephos- phorylation of the enzyme E.coli. phosphatase. Here different phosphory- lating esters react at the same rate, and rearrangement of the active site is presumed to be rate determining. There is currently much theoretical and experimental activity in the fields of conformational changes of proteins, and association of nucleic acid strands. A key experimental paper by Schwarzlo3 describes how transitions between helices and random coils are accompanied by large decreases of dipole moment, and provide a mechanism for dielectric relaxation. This relaxation occurs at higher frequency than rotational relaxation of the polymer; the latter is effective only below lo5 c./sec. With poly-(y-benzyl-L-glutamate) as the polypeptide in solvent mixtures of ethylene dichloride and dichloroacetic acid, the observed mean relaxation time for conformational change is in excellent agreement with theory;lo4 the individual growth step of forming one helical unit of the chain from a coil has a rate constant, k = 1.3 x 1010 sec.-l, which is consistent with diffusion of a peptide group over several A.Poly- (L-lysine) helix-coil changes in water have been investigated by the ultra- sonic-absorption method, but the results are difficult to interpret.lo5 The stability and fast association-dissociation rates of pairs of the different nucleotides A, U, C , G, and some substituted derivatives have been measured in non-aqueous solvents by ultrasonic absorption,lo6 and by dielectric loss in the presence of a high d.c. field gradient.lo7 Differences in stability of the various pairs are alone far too small to account for the invariable biological A-U and C-G pairings.Rates here are very fast, but when the formation of double helices of oligonucleotides containing 3-10 residues of A and U was studied by Eigen and POrschke,lo7 the relaxation time was found in the milli- 99 G. G . Hammes and J. L. Haslam, BiocRemistry, 1968, 7, 1519. loo D. M. Chipman and P. R. Schimmel, J . Biol. Chem., 1968, 243, 3771. lol C. C. F. Blake, G. A. Mair, A. C. T. North, D. C. Phillips, and V. R. Sarma, lo2 Ref. 1 , p. 429. lo3 G. Schwarz and J. Seelig, Biopolymers, 1968, 6, 1263. lo4 G. Schwarz, Biopolymers, 1968, 6, 873; see also ref. 12. lo5 R. C. Parker, L. J. Stulky, and K. R. Applegate, J . Phys. Chem., 1968, 72, 3177. loci G. G. Hammes and A. C. Park, J .Amer. Chem. SOC., 1968, 90, 41 5 1 , lo7 Ref. 1, p. 358. Proc. Roy. SOC., 1967, B, 167, 365.The Kinetics of Reactions in Solution 163 second range. Only one relaxation was found, and from the magnitude of its activation energy the following mechanism was suggested : I I I I I \ I I 1 I I \ I I 1 I \ G i 3 fi I \ I I I \ I I I \ I 1 I I i b 8 I I I I I I I I I 1 I I 1 1 1 1 .i There is a pre-equilibrium, with formation of the third base-pair in a sequence rate-determining for complete helix formation. The renaturation of DNA has been followed using fast electrolytic heating to rapidly perturb the equilibrium position.108 Chance, De Vault, Legallais, Mela, and Yonetanilog provide a comprehen- sive survey of the kinetics of electron transfer in biological systems. Their researches are centred on the role of the cytochromes, essential electron- transfer catalysts in the respiratory chain. Flow studies were carried out with mitochondria as well as with the isolated electron-transport mediators. A pulsed flow apparatus has been developed, with the continuous flow phase (dead time 100 psec.) succeeded by stop-flow observation. Berger has described an optimally designed flow apparatus,l1° and his suggestions for mixer design were incorporated into the Chance system. Chancel09 also demon- strated that laser-activated electron transfer from chlorophyll to cytochrome-c takes less than 2 psec., and Witt has detailed reviews of the use of single and periodic light pulses of lifetime 10-1-10-8 sec. in studies on the photo- synthesis pathway.lll Flash photolytic studies of the role of flavins as photo- oxidizers have been reported by Green and Tollin.l12 The nanosecond flash photolysis apparatus described by Porterll3 and Novak and Windsor114 are of potential interest for biochemists. Postscript--The Landolt ‘iodine clock’ experiment, a classic kinetics demonstration experiment dating from the 1890s,ll5 is now available with lasers and fast recording gear.l16 To the fast-reaction kineticist, nothing is sacred ! lo* J. G. Wetmur and N. Davidson, J . Idol. Biol., 1968, 31, 349. log Ref. 1, p. 437. R. L. Berger, B. Balko, and H. F. Chapman, Rev. Sci. Instr., 1968, 39, 493; R. L. Berger, B. Balko, W. Borcherdt, and W. Friauf, ibid., p. 486. ll1 Ref. 1, pp. 81 and 261. 112 M. Green and G. Tollin, Photochem. and Photobiol., 1968, 7 , 129, 145, 155. 113 Ref. 1, p. 159. 114 J. R. Novak and M. W. Windsor, Proc. Roy. SOC., 1968, A , 308,95. 115 H. Landolt, Chem. Ber., 1885, 18, 249; G. W. A. Fowles, ‘Lecture Experiments in Chemistry,’ Bell, London, 1963, p. 602. J. A. Church and S. A. Dreskin, J. Phys. Chern., 1968, 72, 1387.

ISSN:0069-3022    

DOI:10.1039/GR9686500149

出版商:RSC    

年代:1968

数据来源: RSC