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QUARTERLY REVIEWS THERMODYNAMIC PROPERTIES OF ORGANIC OXYGEN COMPOUNDS By J. H. S. GREEN (CHEMICAL THERMODYNAMICS GROUP NATIONAL CHEMICAL LABORATORY D.S.I.R. TEDDINGTON MIDDLESEX) THE measurement of accurate thermodynamic data is of rapidly increasing importance from two aspects first the data are a basic requirement of modern chemical technology and secondly they contribute to theories of molecular structure. In recent years systematic investigations of certain groups of compounds have resulted in comprehensive compilations of their thermodynamic properties ; an outstanding example is the American Petroleum Institute Research Project 44 on hydr0carbons.l Other im- portant classes of compounds have received much less attention however and the aim of this article is to review the existing state of knowledge of the thermodynamic properties of organic compounds containing only carbon hydrogen and oxygen.Attention is restricted to data on heats entropies and free energies of formation heat capacities and related properties of single substances ; properties of binary and other mixtures thermodynamic dissociation constants and critical data are excluded. After a brief survey of experimental and computational sources the available measurements on the various classes of these compounds are considered in detail and an attempt is made to select the most reliable values. Regularities in the data are then discussed and related to recent methods for the estimation of thermodynamic properties where measure- ments are lacking. Finally some applications of the data to the determina- tion of bond-energy terms and to barrier heights to free rotation are reviewed.Experimental Sources The measurements of heats of combustion by bomb calorimetry as a source of heats of formation is a familiar technique the experimental details of which have been reviewed.2 The accuracy of the method has increased considerably during the last sixty years and in the best experi- Rossini Pitzer Amett Braun and Pimental “Physical and Thermodynamic Properties of Hydrocarbons and Related Compounds,” Carnegie Press Pittsburgh 1953. (a) Rossini (Ed.) “Experimental Thermochemistry,” Interscience Publishers Inc. New York 1956; (b) Skinner “Modern Aspects of Thermochemistry,” Royal Institute of Chemistry London 1958; (c) Coops VanNes Kentie and Dienske Rec. Trav. chim. 1947,66 113 131. 1 125 126 QUARTERLY REVIEWS ments is now about &0.02%.CottrelP has distinguished three classes of combustion results modern work of high precision modern work of relatively low precision and older work for the greater part now over sixty years old. Most of the work on organic oxygen compounds falls in the second and the third category. Rossini4 pointed out in 1937 that most of the old data were subject to uncertainties often as large as several kcal./mole; for many compounds these are still the only values available. In fact recent developments2a of combustion techniques for organic sulphur nitrogen and halogen compounds have yielded values far superior in accuracy to those available for even simple oxygen compounds. The flame calorimeter2a has been used much less than the bomb calori- meter but the few results obtained are of great importance.The heats of formation of some organic oxygen compounds have also been derived from measurements of heats of hydrogenation and of hydrolysis.2b To obtain free energies heats of formation must be combined with entropy values and for organic compounds in general the most accurate results are obtained from low-temperature heat-capacity measurements. If the entropy of a substance is zero at absolute zero then the entropy S at T"' is given for example by where C is the heat capacity at temperature T and at constant pressure and AH' AH" are the enthalpy changes accompanying an isothermal change of state whether of transition fusion or vaporisation. If the terms in this expression were available from O'K then an absolute value of the entropy could be obtained but in practice some extrapolation is needed from the lowest temperature at which the measurements are made down to O'K.If the former is low enough say about 10"~ then a Debye function can be used its parameters being evaluated from the measurements at the lowest temperatures. However in much of the earlier work especially that of Parks and his co-workers the measurements did not go below 9 0 " ~ and an extrapolation procedure5 was used which may give rise to an error in the entropy at 25"c of about one cal./deg./mole. To complete the evaluation of the entropy of a compound in the gas state measurements of vapour heat capacities and of the heat of vaporisa- tion are required. The latter can be obtained by direct measurement and by the use of vapour-pressure data combined with the Clapeyron equation.Variation of vapour pressure with temperature has been expressed by 8 Cottrell "The Strengths of Chemical Bonds," Butterworths Scientific Publications London,2nd. Ed. 1958. Rossmi Ind. Eng. Chem. 1937 29 1424. Kelley Parks and Huffman J. Phys. Chem. 1929,33 1802. GREEN THERMODYNAMIC PROPERTIES OF ORGANIC OXYGEN COMPOUNDS 1 27 many empirical equations,g of which the Antoine equation is currently the most favoured. It has the disadvantage that it causes calculated heats of vaporisation to vary approximately as T-2 and to increase to unreason- ably large values at temperatures much below the observed range. Equa- tions of the Kirchoff-Nernst form imply a more reasonable dependence of heat of vaporisation upon temperature and are more satisfactory; alternatively the equation due to COX’ has been used to represent recent accurate work.However the most accurate values for heats of vaporisation are those obtained by vapour flow calorimetry a method which simultaneously yields accurate values for the heat capacity of the vapour.* In recent years the technique has become of great importance in obtaining complete thermodynamic functions for organic molecules in the vapour state though it has not yet come to be widely used. These accurate measure- ments of vapour heat capacity and heat of vaporisation can be usedg to obtain values for the second virial coefficient; the results can of course be supplemented by those derived from P- V-T measurements. Additional to these thermodynamic quantities derived from heat measurements are those calculated by statistical mechanics from spectro- scopic and molecular-structure data.lo From the molecular dimensions are obtained the moments of inertia required for the calculations of the rotational contribution to the thermodynamic functions.The vibrational contributions can be computed by the standard methods of the harmonic- oscillator rigid-rotator treatment if the complete vibrational assignment of the molecule is available from the interpretation of the infrared and Raman spectra. If the molecule has a group capable of internal rotation then the magnitude of the energy barrier to this rotation can be obtained from the difference between the calculated and the observed specific heat and entropy of the vapour corrected to the ideal gas state.The contribution of this barrier to the thermodynamic functions can then be evaluated. A further difference between observed and calculated specific heats can be attributed in the most accurate work to the effect of anharm~nicity,~ which for polyatomic molecules must be treated empirically.ll In some instances therefore such as acetone12 where an unobserved frequency has Partington “An Advanced Treatise on Physical Chemistry,” Longmans Green and Co. London 1951 Vol. 11 pp. 265-274. Cox Ind. Eng. C‘hem. 1936 28 613. * Sturtevant in Weissberger (Ed.) “Physical Methods of Organic Chemistry,” Interscience Publishers Inc. New York 1959 Vol. I Part I 3rd. Ed. Scott Waddington Smith and Huffman J. Chem. Phys. 1947,15 565. lo (a) Herzberg “Infra-red and Raman Spectra of Polyatomic Molecules,” Van Nostrand Co.Inc. New York 1945; (6) Janz “Estimation of Thermodynamic Proper- ties of Organic Compounds,” Academic Press Inc. New York 1958; (c) Godnev “Calculations of Thermodynamic Functions from Molecular Data,” State Publishing House Moscow 1956. l1 McCullough Finke Hubbard Good Pennington Merserley and Waddington J. Amer. C‘hem. SOC. 1954 76 3661; Pennington and Kobe J. Chem. Phys. 1954 22 1442. l2 Pennington and Kobe J. Amer. Chem. SOC. 1957,79,300. 128 QUARTERLY REVIEWS also to be selected the choice of suitable values for barrier height an- harmonicity and missing frequency to fit the measured properties may be somewhat arbitrary but no large errors are likely if the measurements cover a wide range of temperature. The complete calculations then provide the following thermodynamic functions over a range of temperatures the free energy function (Go - HOo)/T heat content function (H” - H,”)/T as well as specific heat C,” and entropy So.From the measured heat of formation at one temperature and the appropriate values for the elements, AHo” can be found and thence the standard heat AH,” free energy AG,” and logarithm of the equilibrium constant of formation log, K, over the same temperature range. Relatively few organic oxygen compounds have been the subject of a comprehensive study ; for those substances where they are available space does not permit the complete tabulation of all the functions and in the following survey only the values of AH,” So and dG,” at 2 5 ” ~ are given. Thermodynamic Properties No systematic review of the thermodynamic properties of organic oxygen compounds has been made since the pioneer work of Parks and Huffman13 nearly thirty years ago.The heats of formation used by these authors were based essentially on the measurements of heats of combustion recorded in the International Critical Tables; about the same time appeared the extensive compilation by Kharaschl* summarising the com- bustion data on organic oxygen and other compounds. The Landolt- Bornstein tables and their supplements15 give more recent work as does Timmerman’s book,ls whilst the National Bureau of Standards Circular 5001’ gives selected values for compounds containing not more than two carbon atoms. Various authors have listed some heats of combustion and formation in discussing the thermochemistry18 and resonance energieslg of oxygen compounds.To maintain consistency all values given below have been converted on the basis of the atomic weights C = 12.010 H = 1.008 0 = 16.000 and revised to the fundamental constants employed in the tabulations for hydr0carbons.l Heats of combustion have wherever possible been con- l3 Parks and Huffman “The Free Energies of Some Organic Compounds,” A.C.S. Monograph No. 60 The Chemical Catalogue Co. Inc. New York 1932. l4 Kharasch Bur. Stand. J. Res. 1929,2 359. l5 Roth and Scheele (Eds.) Landolt-Bornstein “Physikalisch-Chemische Tabellen,” Julius Springer Berlin 5th Edn. Hauptwerk. 1923 p. 1586; I Eng. Bd. 1927 p.866; I1 Eng. Bd. 1931 p. 1633;.III Eng. Bd. 1936 p. 2896. l6 Timmermans “Physico-Chemical Constants of Pure Organic Compounds,” Elsevier Publishing Co.New York 1950. 1’ Rossini Wagman Evans Levine and Jaffe “Selected Values of Chemical Thermo- dynamic Properties,” Circular 500 National Bureau of Standards Washington 1952. (a) Gray Trans. Faraday SOC. 1956 52 44.4; (b) Gray and Williams ibid. 1959 55 760; (c) Gray and Williams Chem. Rev. 1959 59 239. lo Klages Chem. Ber. 1949 82 358; Wheland “Resonance in Organic Chemistry,” John Wiley & Sons Tnc. New York 1955. GREEN THERMODYNAMIC PROPERTIES OF ORGANIC OXYGEN COMPOUNDS 129 verted to the basis of the modern thermochemical caloric (= 4.1840 abs. joules) and 298.16"~ (25"c); the heats of formation quoted are based on the values17 AH CO (g 25"c) = - 94.0518 and AH H 2 0 (1 25"c) = - 68.3174 kcal./mole. Free energies of formation are derived from the entropy values at 25"c by using the values17 So C (graphite) = 1-361 So H (8) = 31.211 So O2 (g) = 49.003 cal./deg./mole.Limits of error for heats of formation are given in the Tables only when they have been given by the authors but it is not possible to bring them to a uniform basis. In many cases they are not available and it has been preferred not to assign estimated limits of error; for some compounds a comment on accuracy is given in the text. (a) Alcohols Phenols and Polyhydric Alcohols.-The available data for alcohols at 25"c are summarised in Table 1. Although measurements of the low-temperature specific heats20 and of the heat of vaporisation21 of methyl alcohol were made several years ago accurate determination of the complete thermodynamic functions is dependent upon the value of the barrier height to internal rotation.This has been the subject of several investigations.22 Further difficulties arose as to the actual treatment of the internal rotation since at low tem- peratures the molecule fell outside the range of validity of the method usually used and also from the anomalous behaviour of the vapour specific heat which showed an increase with decrease of temperature near the saturation curve. It was shown by Weltner and P i t ~ e r ~ ~ using heat- capacity measurements obtained by flow calorimetry that this behaviour was due to association by hydrogen bonding and they determined the rather large corrections for gas imperfections arising from this. With structural information available from microwave spectroscopy a revised vibrational assignment and a new theoretical treatment of the barrier- height problem Ivash Li and P i t ~ e r ~ calculated the complete thermo- dynamic functions and tabulated them for the range 100-1000"~.Similar difficulties arise with ethyl alcohol studied by Aston and his co-worker~.~~ The values for the barriers to free rotation necessary to give agreement between calculated and observed26,21 entropies were to be very high especially for the hydroxyl group (10,000 cal./ mole). More reasonable values were required by use of a in which the unsymmetrical barrier for this rotation was treated by the use of a potential-energy function with unequal minima that is a trans-form of lower energy than the skew forms and thermodynamic functions cal- 2o Kelley J. Amer. Chem. Soc. 1929,51 181. 21 Fiock Ginnings and Holton J . Res.Nat. Bur. Stand. 1931 6 881. 22 Halford J. Chem. Phys. 1949 17 111; 1950 18 361 1051 and references herein. 23 Weltner and Pitzer J. Amer. C'hem. SOC. 1951 73 2606. 24 Ivash Li and Pitzer J. Chem. Phys. 1955 23 1814. 25 (a) Schumann and Aston J. Chem. Phys. 1938,6,480; (b) Aston Ind. Eng. Chem. 1942 34 514; (c) Aston Szasz and Isserow J. Chem. Phys. 1943 11 532; ( d ) Brick- wedde Moskow and Aston J. Res. Nat. Bur. Stand. 1946,37,263. 26 Kelley J. Amer. C'hem. Soc. 1929 51 779. 130 QUARTERLY REVIEWS culated in this way were tabulated.26d The agreement with the observed entropy was satisfactory but that of heat capacities was much poorer the calculated values being too high. Barrow2' considered the effect of associa- TABLE I. Thermodynamic properties of alcohols phenols and polyhydric alcohols at 25 OC (kcal./mole and cal./deg.lmole) Alcohol Methyl Ethyl Propyl Isopropyl Butyl Isobu tyl S-Butyl t-Butyl Pentyl t-Pentyl Hexyl Heptyl Hexadecyl Cyclopentyl Cyclohexy 1 Benzyl Diphenylmet hy I Triphenylmethyl Phenol Ethylene glycol Glycerol Erythritol Mannitol Dulcitol - AH," 57-02 f 0.05 48.08 f 0.05 66-36 f 0.10 56.24 f 0-12 73.20 f 0-24 61-85 f 0.26 76.18 65.42 79-54 f 0.10 66.92 & 0.22 80.00 f 0.1 1 67.9 f 0.25 81.88 f 0-13 70.1 f 0.25 85-87 f 0.10 74.9 4 0.25 85-65 f 0.40 71.85 f 0-47 96.1 91.75 f 0-48 76.75 f 0-56 97.85 f 0.56 81.65 f 0.66 163.55 151.86 71.77 83-45 38-49 f 0.30 25-16 & 0.50 0.80 f 0-60 23-05 f 0.15 39.46 f 0.08 108.74 95.10 159-80 217.61 319-61 32 1 *90 Ref. 17 17 17 17 28 28 35 28 28 31 31 31 31 31 31 28 28 13 28 28 28 28 35 43 35 35 46 46 46 47 47 51 32 51 51 51 51 (a) S" 30.3 57-29 38.4 67.58 46.1 76- 19 43.0 73.92 54-5 89.42 - - - - 45-3 76.8 60.9 54.8 68.6 77.9 108.0 145.0 49.3 47.7 51.8 57.3 78.7 75.44 35.71 39.9 77-33 48.87 39-9 57.0 56.0 - - - Ref.20 24 26 27 33 32 37 (a) 33 32 33 (b) 42 42 26 43 43 43 43 44 45 13 13 49 49 33 32 53 13 13 13 - AGf" 39.73 38.84 41.77 40-35 41.2 39-13 43-26 41-71 40.33 38.12 - - - - 43.92 42.3 38-63 47.3 37-3 36.5 23.74 23.08 30.6 32.1 6.6 - - - - 26.2 - 65.2 7-89 12-45 77-3 74.82 114.01 152.7 1 225.2 227.2 (a) See text. (6) Calculated from So (l) vapour pressure and heat of vaporisati~n,~~ the effect of gas imperfections being ignored. 27 Barrow J. Chem. Phys. 1952 20 1739. GREEN THERMODYNAMIC PROPERTIES OF ORGANI'C OXYGEN COMPOUNDS 13 1 tion in the vapour by the same method as that used for methyl and also gave a revised vibrational assignment in terms of which only a symmetrical threefold potential function was required for both the methyl and hydroxyl groups to give good agreement with heat capacities and entropies as well as measurements of the equilibrium between ethyl alcohol ethylene and water.Further calculations giving the functions from O-lOOOo~ have been made.27a Data on all other alcohols are more diffuse. The values for heats of formation given in Table 1 for the higher normal aliphatic alcohols are those of a recent revision28 of the combustion measurements previously considered by R o ~ s i n i . ~ ~ A fresh choice of "best" selected values within the experimental errors of the results was made but the greatest change is in the values for the heats of vaporisation at 25"c derived from vapour- pressure measurements.The value given for butyl alcohol is that of Tjebbes30 but a recent measurement by Skinner and Snelson3' yielded at 25"c AH," (1) = - 78.49 & 0.11 AH," (g) = - 66.1 -C 0-25 kcal./ mole. Values of the entropy enthalpy and free-energy functions for propyl and butyl alcohols from 298.16 to 10oO'~ were calculated by D ~ a t k i n a ~ ~ who used a vibrational assignment based on Raman spectra only. The contributions of the possible rotations about the C-C and C-0 bonds were allowed for and the barrier heights were assumed to be 3000 and 2100 cal./mole respectively but there are insufficient experimental data against which to check the calculations. For propyl alcohol however there is satisfactory agreement at 298.16'~ between the calculated entropy 77.25 cal./deg./mole and the value of 77.1 d~ 0.5 cal./deg./mole derivable from the entropy of the the revised heat of vaporisation2s and recent vapour-pressure meas~rements,~~ gas imperfections being ignored.The two most recent measurernent~~~~~~ of the heat of formation of isopropyl alcohol are in good agreement but the position of the tabulated thermodynamic is unsatisfactory. Although reasonably good agreement with the entropy of the vapour derived36a from the value for the was obtained by Aston and his co-workers using a rather schematic vibrational assignment the agreement with heat-capacity 37u Green in preparation. 2s Rossini J. Res. Nat. Bur. Stand. 1934 13 189. 31 Skinner and Snelson Trans. Furuduy Suc. 1960 56 1776.3x Dyatkina Zhur. fiz. Khim. 1954 28 377. 33 Parks Kelley and Huffman J. Amer. Chem. SOC. 1929 51 1969. 34 Copp and Findlay Trans. Furaduy SOC. 1960,56 13. 36 Parks Mosley and Paterson J. Chem. Phys. 1950 18 152. Green Chem. and Ind. 1960 1215. Tjebbes Acfa Chem. Scand. 1960 14 180. (a) Schumann and Aston J. Chem. Phys. 1938 6 485; (b) Aston Isserow Szasz and Kennedy ibid. 1944 12 336; (c) Kobe Harrison and Pennington Petroleum Refiner 1951,30 1 19. 37 (a) Kelley J. Amer. Chem. SOC. 1929 51 1145; (b) Ginnings and Corruccini Ind. Eng. C'hem. 1948,40 1990. 132 QUARTERLY REVIEWS measurements3* was rather poor. A later calculation36c of the enthalpy and heat capacity yielded slightly different values from but no better agreement with either the earlier or subsequent39 observations and it is possible that the measured values are in error.The listed heat of forma- tion of isopropyl alcohol of the gas is the mean of the value (65.56 kcal./ mole) obtained for the liquid together with the heat of vapori~ation,~~ and that deduced from the heat of hydrogenation of acetone discussed later (see p. 139). Similarly the entropy is the mean of a calculated value36c (74.14 cal./deg./mole) and that deduced from measurements of the equilibrium between isopropyl alcohol hydrogen and acetone (see p. 139). For isobutyl s-butyl and t-butyl alcohols there is a paucity of data. The heats of combustion are now accurately known for all three com- p o u n d ~ ~ ~ but an entropy value is available only for t-butyl alcohol the listed value for the vapour being calculated from that of the by use of vapour-pressure and heat of vaporisation data,40 the effect of gas imperfection being ignored.Taft and Riesz41 studied the equilibrium between isobutene water and t-butyl alcohol in dilute acid and obtained results in poor agreement with the earlier13 values for the alcohol; the agreement is very much improved by the new values and we have for the reaction at 25"c with the pure liquid alcohol AH = -12.6 (obs.) -13.5 (calc.); AG = -1.3 (obs.) - 1.1 (calc.) kcal./mole. Measurements of the vapour heat capacities of the butyl alcohols and of pentyl alcohol were made by Sinke and De V r i e ~ . ~ ~ are the only data available. Values for hexadecyl alcohol as solid and hypothetical (at 25"c) liquid were given by Parks and his co-workers;43 the derived heat of formation of the liquid is in satisfactory agreement with the value 152.84 I!= 1-3 kcal./mole obtained by extrapolation from the revised values28 for the lower alcohols.The listed entropies of diphenylmethyl alcohol and triphenylmethyl alcohol are those given by Parks and Huffman13 and are based on unpub- lished measurements by Andrews. The heats of formation of these sub- stances are more reliable and together with the value for benzyl alcohol are taken from the work of Parks Manchester and V a ~ g h a n . ~ ~ However a comparison of their values for phenol quinol and benzoquinone with For all other alcohols entropy measurements at 25"c on the 38 Parks and Shomate J . Chem. Phys. 1940,8,429. 38 Sinke and De Vries J. Amer. Chem. SOC. 1953 75 1815. 40 Parks and Barton J . Amer. Chem.SOC. 1928 50 24. 41 Taft and Riesz J. Amer. Chem. SOC. 1955 77 902. 43 Parks Huffman and Barmore J. Amer. Chem. SOC. 1933,55 2733. 43 Parks Kennedy Gates Mosley Moore and Renquist J. Amer. C'hem. SOC. 1956 44 Kelley J. Amer. Chem. SOC. 1929 51 1400. 45 Parks Todd and Moore J. Amer. Chem. SOC. 1936 58 398. 46 Parks Manchester and Vaughan J. Chem. Phys. 1954,22,2089. 78 56. GREEN THERMODYNAMIC PROPERTIES OF ORGANIC OXYGEN COMPOUNDS 133 later measurement^^^^^^ suggests that their values of - AH may be low by about 0 - 6 - 4 8 kcal./mole. For phenol the values listed are therefore preferred. The thermodynamic functions for phenol have been calculated49 from a complete vibrational assignment together with the barrier height as measured by microwave spectro~copy.~~ These calculations and accurate vapour-pressure and heat of vapori~ation~~ measurements (the correction for gas imperfection is negligible) indicate that the measured value for the entropy of solid phenolq2 is low by 1.6 cal./deg./mole.The listed value for the heat of formation of ethylene glycol is that of Parks and his co-worker~,~~ in preference to that of other and a presumably averaged va1ue.l' Calculations of the thermodynamic functions excluding heat capacity were made by D ~ a t k i n a ~ ~ using the same procedure and assumptions as for propyl and butyl alcohols. A rather more accurate value53 of the entropy of glycerol replaces that derived13 from very early measurements but no improvement is possible on the previously listed values for erythritol mannitol and dulcitol although the heats of formation have been re-meas~red.~' From the heats of solution of erythritol mannitol and dulcitol in water54 the heats of combustion of the hypothetical liquid forms of the substances at 25"c were derived the corresponding heats of formation are -212.04 -314.22 and -3 14.8 1 kcal./mole respectively.The heats of formation of this series of substances C,H,,+,O are reproduced to within 0.5 kcal./mole by the expression - AH," (1,25"c) = 5.82 + 51.46~. There exist only heats of combustion data for all remaining com- pounds ; heats of formation derived from work not previously summarised are given in Table 2. The values for ct- and P - n a p h t h ~ l ~ ~ are rather less accurate than those for the cres01s~~ and xylenols;47 that for cycloheptyl alcohol is from recent work by Skuratov and his co-worker~.~~ Values are available for a number of cyclohexyl alcohol derivatives :57-59 those for the methylcyclohexyl alcohols are due to Skita and F a u ~ t ~ ~ ~ and are listed with the changed assignment of configurations.60 The values listed for various d i o l ~ ~ ~ may be in error by several kcal./mole to judge by these authors' 47 Andon Biddiscombe Cox Handley Harrop Herington and Martin J.1960 5246. 48 Pilcher and Sutton J. 1956 2695. 4 9 Green J. 1961 2236. 5 0 Kojima J. Phys. Soc. Japan 1960 15 284. 51 Parks West Naylor Fujii and McClaine J. Amer. Chem. SOC. 1946 68 2524. 52 (a) Moureu and DodC Bull. SOC. chim. France 1937,4,637; (b) Jung and Dahmlos b3 Ahlberg Blanchard and Lundberg J. Chem. Phys. 1937 5 539. 64 Parks and Manchester J. Amer. Chem. SOC. 1952,74 3435.55 Leman and Lepoutre Compt. rend. 1948,226 1976. 66 Skuratov Kozina Shtecher and Varushyenko Thermochem. Bull. I. U.P.A.C. 67 Nicholson J. 1960 2378. 68 Landrieu Baylocq and Johnson Bull. SOC. chim. France 1929 45 36. 5 9 (a) Skita and Faust Ber. 1931 B 64 2878; (6) ibid. 1939 B 72 1127. 6o Eliel and Hober J . Org. Chem. 1958 23 2041. Z. phys. Chem. 1942 A 190,230. 1957 No. 3 25. 134 QUARTERLY REVIEWS TABLE 2. Heats of formation of alcohols phenols and polyhydric alcohols at 25 O c (kcaI./mole) - AH," Octyl alcohol (1) 103.96 f 0.64 (g) 86.56 f 0.75 Nonyl alcohol(1) 110-07 -i 0-72 (8) 91.47 f 0.84 Decyl alcohol (1) 116.18 f 0-80 (g) 96.36 f 0.94 o-Cresol (s) 48.91 & 0-12 (g) 30-74 f 0.22 rn-Cresol (1) 46-38 f 0.07 (g) 31-63 f 0.26 p-Cresol (s) 47-64 i 0.08 (g) 29.97 f 0.36 2,3-Xylenol (s) 57.67 &- 0.1 1 (8) 37.59 f 0.27 2,4-Xylenol (1) 54.69 f 0.11 (g) 38.95 j= 0.18 2,s-Xylenol (s) 58.96 f 0.10 (g) 38.65 f 0.12 2,6-Xylenol (s) 56.75 f 0.12 (g) 38.68 f 0-13 3,4-Xylenol (s) 57.93 f 0.13 (g) 37.44 f 0.14 3,5-Xylenol (s) 58.43 f 0.14 (g) 38-63 f 0.16 a-Naphthol (s) 26-4 @-Naphthol (s) 29.3 Cycloheptanol (I)* 94.0 3,3,5-Trimet h y 1- cyclohexanol (1) 109.2 Cyclohexylme t hanol 102-0 * At 20".t At 17". Ref. 28 28 28 28 28 28 47 47 47 47 47 47 47 47 47 47 47 47 47 47 47 47 47 47 55 55 56 57 58 -AH," cis-2-Methylcyclo- trans-2-Met h ylc yclo- cis-3-Methylcyclo- trans-3- Methylcyclo- cis-4- Methylcyclo- trans-4-Methylcyclo- cis-3 ,cis-S-Dimethyl- trans-3,trans-5-Dimethyl cis-3,trans-5-Dimethyl- hexanol(1) 94-85 hexanol (1) 100-93 hexanol (I) 101.02 hexanol (1) 95.82 hexanol (I) 100.24 hexanol (1) 105.07 cyclohexanol (1) 100- 9 cyclohexanol (1) 1 15-4 cyclohexanol (1) 121.3 Diethylene glycol (l)t 149.3 Triethylene glycol (I)? 191.0 Tetraethylene glycol (I)? 233-1 Propane-1,Zdiol (l)? 118.9 Butane1,2-diol(l)t 124.4 Butane-1,3-diol (I)? 122.3 Butane-2,3-diol (I)? 128.7 Isobutane-1 ,2-diol(l)t 128.2 2-Ethylhexan-1-01 (1) 103.46 Ref.59a 59a 59a 59a 59a 59a 596 59b 59b 52a 52a 52a 52u 52a 5242 52a 52a 30 values for ethylene glycol and for ethylene oxide and propylene oxide (seep. 136). Studies of a number of cis- and trans-isomers of various diols were made by Verkade Coops and their co-workers,61 and the figures in Table 3 are a revision of their values converted into heats of formation at 25"c. For cyclopentanediol more recent gives - AH (cis) = 115.6 (trans) = 116.8 kcal./mole.in satisfactory agreement with the listed values. (b) Ethers Cyclic Ethers and Derivatives.-Thermochemical data on even the simplest aliphatic ethers are extremely scanty and only one modern accurate value of the heat of formation is available that for 61 Verkade Coops Maan and VerkadeSandbergen Annalen 1928,467 217. GREEN THERMODYNAMIC PROPERTIES OF ORGANIC OXYGEN COMPOUNDS 135 TABLE 3. Heats of formation of diols at 25"c Cyclopentane-1 ,Zdiol (s) 1 -Methylcyclopentane-1 ,2-diol (1) 1 -Phenylcyclopentane- 1,2-diol (s) Indane-1 ,2-diol (s) 1,2,3,4-Tetrahydronaphthalene- 1,2-diol (s) 1,2,3,4-Tetrahydronaphthalene-2,3-diol (s) Cyclohexane-l,2-diol (s) l-Methylcyclohexane-l,2-diol (s) l-Phenylcyclohexane-1,2-diol (s) - AH," fkcal./mole) cis trans 115.2 117.2 126.0 128.3 97.5 - 88.7 88.5 98-5 99.8 98.5 98.9 131-7 130.7 143.0 140.5 110.6 109'8 diethyl ether measured by Pilcher," who finds by using a flame calori- meter that AH," (g 24"c) = -60.37 rt 0.18 kcal./mole.(The old measure- ments gave -57.4 and for ethyl methyl ether -52.0 kcal./mole.) The vibrational assignment of dimethyl ether is however well established and several calculations of the thermodynamic functions have been made from it,62 the results of which are in reasonabIe agreement with one another and with the measured entropy 62a and heat The values listed in Table 4 are by Si5ha62c who tabulated the functions from 298.16- 1000"~. Studies of the equilibrium between ethyl alcohol diethyl ether and water have been to derive values for the functions for diethyl ether but the results are in poor agreement with the meagre data available.TABLE 4. Thermodynamic properties of some ethers and cyclic ethers at 25 O c (kcal./mole and cal./deg./mole) Dimethyl ether (8) Di-isopropyl ether (1) Diphenyl ether (s) Ethylene oxide (g) Propylene oxide (g) Trimethylene oxide (g) 1,4-Dioxan (1) 1,3-Dioxan (1) Furan (1) $ 9 (g) Furfuryl alcohol (1) * At 20". - AH^" 44.3 83-94 7.67 12-19 22.02 95.5 94.5 92*1* 14,903 8-293 66-05 - Ref. 62c 65 64 66 186 71 72 56 76 76 35 So 63.74 70.4 55.91 58-13 67- 15 63.40 47-0 42.22 63-86 51.6 - Ref. 62c 42 64 65a 69 69a 70 76 76 43 - AG," 27.27 21.1 34.37 2.79 5.60 - 55.1 - 0*050 - 0.208 36-88 - * Dr. G. Pilcher University of Manchester private communication. (a) Kennedy Sagenkahn and Aston J.Amer. Chem. SOC. 1941 63 3267. (b) Hadzi Cornpt. rend. 1954 239 349; (c) %ha Chem. Listy 1955 49,1569; ( d ) Taylor and Vidale J. Chem. Phys. 1957 26 122; (e) Mashiko and Pitzer J. Phys. Chem. 1958 62 367. 63 Kiastiakowsky and Rice J. Chem. Phys. 1940.8 610. 64 Valentin J. 1950 498 and references herein. 136 QUARTERLY REVIEWS Values of the entropy42 and heat of formations5 of di-isopropyl ether at 25"c are available. Measurements have been made for diphenyl ether leading to a tabulation of the enthalpy entropy and heat capacity in the range 0-570"~ to- gether with the heat of formation.s6 The heats of formation of a number of cyclic ethers have been measured and ethylene oxide has been the subject of a number of more complete investigations and calculations which are in satisfactory agreement.63967 For the heat of formation we follow an earlier compilation17 in listing the value obtained by Crog and Hunt;6a the older work including that of Moureu and D o ~ C ~ ~ ~ gave considerably different results both for this compound and for propylene oxide for which the value obtained by Stulllab is listed in Table 4.The latter value has been used together with a complete vibrational assignment and the barrier height determined by microwave spectroscopy to calculate the complete thermodynamic functions for propylene oxide from 0" to IOOO"K.~~ Calculated functions for trimethylene oxide are also availablesgu but there is no value for the heat of formation of this compound. A measurement of the of 1,4-dioxan at 25"c is available but the heat of formation7' is a rather old one more recent values exists for 1,3-dio~an,~~ but the difference between them is rather large.The remaining values are summarised in Table 5. It may be noted that the results due to Bad~che'~ for phenol and rn-cresol agree to within 0.2 kcal./mole with recent rneasurement~.~~ Other values in the Table are due to S ~ r i n g a l l ~ ~ and S k ~ r a t o v ~ ~ ~ ~ and their co-workers. Their ~ a l u e s ~ ~ ~ * ~ for the heat of combustion of liquid tetrahydrofuran differ by 0.8 kcal./ mole and the mean value -598.4 kcal./mole together with the heat of ~aporisation7~~ was used to derive the listed heat of formation. Larger differences exist between the two sets of results for tetrahydropyran and 1,3-dioxolan both of which are listed. 6s Parks and Manchester Thermuchem. Bull.I.U.P.A.C. 1956 No. 2 8. 68 Furukawa Ginnings McCoskey and Nelson J. Res. Nat. Bur. Stand. 1951 46 (a) Godnev and Morozov Zhur. fiz. Khim. 1948,22 801 ; (b) Gordon and Giague J. Amer. Chem. Suc. 1949 71 2176; (c) Arnold quoted by Kobe and Pennington Petroleum Refiner 1950,29 135; ( d ) Giinthard Mesikommer and Kohler Helv. Chim. Acta 1950 33 1809. 19:; d B Crog and Hunt J. Phys. Chem. 1942,46,1162. 6s Green Chem and Ind. 1961 369. 600 Ziircher and Giinthard Helv. Chim. Acta 1955,38 849; 1957 40 89. l o Jacobs and Parks J. Amer. Chem. Suc. 1934 56 1513. 71 Roth and Meyer 2. Elektrochem. 1933 39 35. Fletcher Martimer and Springall Bull. Chem. Thermodynamics I . U.P.A.C. 1958 No. 1,817. 73 Badoche Bull. SOC. Chim. France 1941 8 212. l4 (a) Springall Mortimer and Fletcher Thermochem.Bull. I. U.P.A.C. 1957 No. 3 12; (b) Cass Fletcher Mortimer Springall and White J. 1958,1406; (c) Cass Fletcher Mortimer Quinay and Springall J. 1958 2595. 76 Skuratov and Kozina DokIady Akad. Nauk S.S.S.R. 1958 122 109. GREEN THERMODYNAMIC PROPERTIES OF ORGANIC OXYGEN COMPOUNDS 137 A detailed study of furan has been made76 and a tabulation given of the complete thermodynamic functions from 273.16" to 1500"~. The sub- sequent change7' in the vibrational assignment does not seriously change these values since the original assignment included four frequencies chosen empirically to give the best fit to the measured entropies and heat capacities and at the same time allowed for the contributions of anharmonicity. The latest value for the heat of formation35 of furfuryl alcohol is in good agreement with an earlier value obtained by Landrieu Baylocq and who also made measurements on some other furan derivatives but unfortunately at an unstated temperature; revised values from their other results are given in Table 5.TABLE 5. Heat of formation of some ethers and cyclic ethers at 25"c (kcal./mole) -AH," Methyl phenyl ether (1) 29.61 Ethyl phenyl ether (1) 37.67 rn-Methoxytoluene (1) 36.69 Tetrahydrofuran (1) 5 1 * 1 9 (g) Tetrahydropyran (1) 2-Methoxytetrahydro- Dihydropyran (1) Dibenzofuran (s) Dibenzopyran (s) 5,5'-Spirobis-rn-dioxan Furfuraldehyde (1) PYran (1)" (s) * At 20". t At 17". 43.5 59.1 61-33 50.7 53.98 104.4 37.6 1.4 15.3 167-9 46.4 Ref. 73 73 73 56 74b 74b 56 74b 74b 56 74b 746 74b 74a 58 - AH," Furylacrylic acid (1) 107.3 Tetrahydrofurfuryl alcohol (1) 102.0 1,3-Dioxolan (l)* 79.5 7 9 9 80.7 4-Methyl-l,3-dioxan (11 20" 102.0 1,3-Dioxepan (I)* 92.5 1,3-Dioxocan (1)* 90.2 o-Dimethoxybenzene (1) 89.5 1,4-Benzodioxan (1) 61 -0 1,3-Benzodioxolan (1) 44.0 2,3-Benzo- 1,4-dioxepan (1) 57.8 Butylene oxide (1)t 39.7 Furoic acid (1) 119.1 Furylethylene (1) 0.3 Ref.58 58 56 72 56 56 56 74c 74c 74c 74c 52a 58 58 (c) Aldehydes.-The vibrational assignment for formaldehyde is well establishedloU and several calculations of the thermodynamic functions have been made. D ~ o r j a n y n ~ ~ tabulated values from 298.16" to 1200"~ which were derived by using the most recent determination by microwave spectroscopy of the moments of inertia of the molecule. No very reliable value for the heat of formation is available however; that listed in Table 6 was derived1' from rather old work.76 Guthrie Scott Hubbard Katz McCullough Gross Williamson and Waddington 77 Bak Brodersen and Hansen Acta Chem. Scand. 1955 9 749. 70 Dworjanyn Austral. J. Chem. 1960 13 175. J. Amer. Chem. SOC. 1952,74 4662. 138 QUARTERLY REVIEWS TABLE 6. Thermodynamic properties of some aldehydes and ketones at 25 OC (kcal./mole and cal.ldeg.lmole) - AH," Ref. Formaldehyde (g) 27.7 17 Acetaldehyde (g) 39-67 79b Butyraldehyde (1) 57-2 30,82 Heptanal (1) 74.5 & 0.9 82 Acetone (1) 59.34 84 $ 9 (8) 5 1 -72 84 Ethyl methyl ketone (1) 66-68 35 Benzophenone (s) 8.1 35,86a Dibenzoylethylene (s) 27.55 f 0-60 46 Dibenzoylethane (s) 61.24 f 0-40 46 S" 52.26 63.15 59.0 83.3 47-9 70.49 57-71 58.6 77.6 76-3 Ref. 78 79b 43 43 84 84 43 13 87 87 - AG," 26.3 31.77 28-6 24.1 37.19 36.30 37.73 - 33-5 - 26.3 - 2.3 Several calculation^^^ have also been made for acetaldehyde that of Pitzer and W e l t n e ~ ~ ~ ~ yielding the most reliable values since the para- meters were chosen to give agreement with the observed heat capacities79c and measurements of the equilibrium between acetaldehyde ethanol and hydrogen.79d (A revisionso of the assignment used by these authors is mainly in the region of higher wave-numbers and therefore has a negligible effect on the calculated values.) By using these calculated functions for acetaldehyde together with those for ethanol25d and hydrogen,17 the listed heat of formation of acetaldehyde can be derived from the heat of hydrogenation measured by Kistiakowsky and his co-workers.81 Recent values for the heat of combustion of b ~ t y r a l d e h y d e ~ ~ ~ ~ and heptanals2 are available; for the former the mean of the two values is listed.Tjebbes30 also gives the heats of formation of the following liquid aldehydes at 25 *c but-2-enal -34.45 0.09 ; 2-ethylhexanal -83.32 31 0.18 and 2-ethylhex-2-ena1 4 2 - 4 6 & 0.17. A heat of combustion of benzaldehyde has been quoted19 and converted into a heat of formation.lsa Revision of the original measurernent~~~ yields AH," (I) = -17.8 kcal./mole the temperature not being stated; with the estimated heat of vaporisation this gives AH," (g) = -6.0 kcal./mole. These values may be subject to fairly large errors as also may the following values (kcal./mole) for a number of substituted benzalde- hydes measured at the temperature stated :83 'Is (a) Smith Trans.Amer. Inst. Chem. Engineers 1946,42,983; (b) Pitzer and Weltner J. Amer. Chem. Suc. 1949 71 2842; (c) Coleman and De Vries ibid. 1949 71 2839; ( d ) Rideal Pruc. Roy. SOC. 1921 A 99 153. Evans and Bernstein Canad. J. Chem. 1956 34 1083. Dolliver Gresham Kistiakowsky Smith and Vaughan J. Amer. Chern. SOC. 1938 60 4-40. 82 Nicholson J. 1360 2377. 83 (a) Bonino Manzoni-Ansidei and Rolla Ricerca sci. 1937 8 5 ; (b) Manzoni- Ansidei and Storto Atti Accad. naz. Lincei Rend. Classe Sci. fis. mat. nat. 1940 1 465. GREEN THERMODYNAMIC PROPERTIES OF ORGANIC OXYGEN COMPOUNDS 139 o-hydroxy (20") -67.0 rn-hydroxy (20") -70-3 p-hydroxy (20") -74.3 o-methoxy (17") -62.6 rn-methoxy (17") -64.9 p-methoxy (1 7 ") -62.8 3-hydroxy-4-methoxy (at 17 ") - 107.3 (4 Ketones.-A complete tabulation of the thermodynamic functions is available only for acetone.The calculations were based on the heat of formation entropy measurements of the liquid heat of vaporisation and heat-capacity measurements by flow calorimetry of sufficient accuracy to permit evaluation of the contribution of anharmonicity in terms of the vibrational as~ignment.~~ The resulting values can therefore be regarded as well established. Several authors have studied the equilibrium between acetone isopropyl alcohol and hydrogen,s5 and the early work was discussed by Parks and Huffman,13 chiefly by using the equilibrium constants measured by Parks and K e l l e ~ . ~ ~ ~ The somewhat different results obtained by later workers were however in mutual agreement. Kistiakowsky and his co-workerss1 found the heat of hydrogenation of acetone at 3 5 5 " ~ to be 13.407 kcal./ mole with which the value 13.69 kcal./mole found by Kolb and Bu1we11~~~ from their equilibrium measurements was considered to be in satisfactory agreement.By using the former value together with heat capacities for acetone,8* isopropyl and hydrogen,17 we find for isopropyl alcohol AH," (g 25"c) = -65.28 kcal./mole. Similarly from the equili- brium measurementsSsb we find for isopropyl alcohol So (g 2 5 " ~ ) = 73.7 1 cal./deg./mole. Vogler and TrumplerssG analysed their measurements over the range 277-327"c and 100-400 atm. to obtain fugacities which were in good agreement with those obtained by estimation. The vapour heat capacities and heats of vaporisation of ethyl methyl ketone and methyl propyl ketone have been measured recently,ss and the thermodynamic functions tabulated from 0" to 1500"~.From these values those for the methyl n-alkyl series have been calculated by the method of increments. Values for the entropy43 and heat of of ethyl methyl ketone are also available the value listed for the latter quantity is preferred to that obtained by Crog and Hunt68 who burnt the substance as a gas but give the result for the liquid having made an unstated "correc- tion" to this state. Skinner and S n e l ~ o n ~ ~ quoting the values6 of 8.31 kcal./mole for the heat of vaporisation of ethyl methyl ketone at 25"c find that there is not a very satisfactory consistency between their value for A H " (g) of s-butyl alcohol (-70-1 kcal./mole) Crog and Hunt's value of AH," (8) of ethyl methyl ketone (-58.9 kcal./mole) and the heat of hydro- 84 Pennington and Kobe J.Amer. Chem. Soc. 1957,79 300. 86 (a) Parks and Kelley J. Phys. Chem. 1928,32 740; (6) Kolb and Burwell J. Amer. Chem. Soc. 1945 67 1084; (c) Vogler and Trumpler Hefv. C'him. Acta 1956 39 757; ( d ) Ciborowski Chim. et Ind. 1958 80 240. 86 Nickerson and McKetta Bull. Chem. Thermodynamics I.U.P.A.C. 1960 No. 3 47; Nickerson Ph.D. thesis University of Texas 1960. 140 QUARTERLY REVIEWS genationso of the latter substance reduced to 25 O c (- 12.95 kcal./mole). The agreement is better by 0.45 kcal./mole if the present values are used but is still not entirely satisfactory. Two recent measurements of the heat of formation of benzophe- none35j86u differ by 0.4 kcal./mole; the mean value is listed here together with the entropy as estimated by Parks and Huffman.13 For dibenzoyl- ethylene and dibenzoylethane the entropy measurements are again due to Parks and Huffman;s7 with the modern values46 for the heats of formation the sign of d Gf O for dibenzoylethane is changed from that given previo~s1y.l~ Recent value^^^-^^ for the heats of formation of a number of ketones are summarised in Table 7.For benzyl methyl ketone the value listed is that from two r n e a s ~ r e r n e n t s ~ ~ ~ ~ ~ in good agreement chosen in preference to a third of 38-96 kcal./mole. The three recent values for diacetyl TABLE 7. Heat of formation of some ketones at 25"c (kcal./mole) - AHfo Acet y lacet one (1) p-Methylbenzophenone 6) 184.0 p-Ethylbenzophenone (1) 243.0 p-Isopropylbenzo- phenone (1) 520-5 p-t-Butylbenzophenone (1) 686.7 Methyl phenyl ketone (1) 34.06 Ethyl phenyl ketone (I) 39.95 Propyl phenyl ketone (1) 45.13 Isobutyl phenyl ketone (1) 52.62 t-Butyl phenyl ketone (1) 49-91 2,4,6-Trimethylaceto- phenone (1) 65.0 101.33 * At 20".Ref. 89 86 86 86 86 88 88 88 88 98 93 - AHf" Ref. 2,4,5-Trimethylaceto- phenone (1) 61.4 93 Cyclopentanone (I)* 57.6 56 Cyclohexanone (1)* 69.5 56 Cycloheptanone (I)* 71.1 56 90 Cyclodecanone (1)" 81.6 90 Dibenzyl ketone (s) 20-3 91 Benzil (s) 42.7 91 Diacetyl (1) 87.5 46 91 92 Benzyl methyl ketone (1) 36-4 91 92 Norcamphor (s) 54-4 94 endo-Ethylene- cyclohexanone (s) 15-1 94 Colomina Cambeiro Perez-Ossorio and Latore Anales real SOC. espaii. Fis. Quim. 1959 6 509. Parks and Huffman J. Amer. Chem. SOC. 1930,52,4387. Colomina Latore and Perez-Ossorio Bull.Chem. Thermodynamics I. U. P.A.C. 1958 No. 1 A19. Nicholson J. 1957 2431. Skuratov Kozina Shtecher Prevalova Kamkina and Zuko Bull. C'hem. Thermo- dvnamics I.U.P.A.C. 1958 N o . 1 A21. 91 Springall and White J. 1954 2764. 92 Nicholson Szwarc and Taylor J. 1954 2767; Therrnochem. Bid/. I.U.P.A.C. 1956 No. 2 16. O3 Baker and Tweed J. 1941 796. 94 Becker and Roth Ber. 1934 B 67 627. GREEN THERMODYNAMIC PROPERTIES OF ORGANIC OXYGEN COMPOUNDS 14 1 -87.60,46 8 6 ~ 4 ~ ~ 87~35~~-are in poor agreement and the choice of 87.5 kcal./mole is listed. (e) Acids and Anhydrides.-With the exception of formic acetic and benzoic acid the available thermodynamic data on organic acids are in- complete and much is of poor accuracy. The experimental results for formic acid were reviewed in detail by Waringg5 and subsequent work comprises an accurate measurement by Sinkegs of the heat of combustion of the liquid substance and a revisiong7 of the vibrational assignment leading to a new value for the entropy of the monomer in the ideal gas state at 25"c.The values listed in Table 8 TABLE 8. Thermodynamic properties of some acids anhydrides and esters at 25 O c (kcal./mole and cal./deg./mole) Acid Formic (1) Formic monomer (g) Formic dimer (g) Acetic (1) Acetic monomer ( g ) Acetic dimer (8) Butyric (1) Palmitic (s) Lactic (s) Lactic (1) Benzoic (s) o-Hydroxybenzoic (s) m-Hydroxybenzoic (s) p-Hydroxybenzoic (s) Oxalic (s) Fumaric (s) Maleic (s) Succinic (s) Phthalic (s) Phthalic anhydride (s) Methyl formate (1) Methyl formate (g) Ethyl acetate (1) Ethyl acetate (g) - AH," Ref.101-52 f 0.06 96 90.49 97a 195.12 97a 115.7 f 0-10 99b 103.8 98 223.0 98 127.2 101 102a 21 1.2 165.89 -k 0.11 161.1 140.0 141.1 142.0 196-7 193.83 188.28 224.77 f 0.06 186.88 91.812 f 0.07 1 10.03 88.6 81.0 106.2 103 -4 101 106 13 lOOd 16 109 109 111 117 117 113 35 35 126 126 102a 130 So 30-82 59.45 82.89 38.2 67.5 96.7 54- 1 104.8 34.0 45.9 40.04 42.6 42.3 42.0 28.7 39.7 38.1 42.0 49.7 42.9 - 84.6 62.0 90.1 Ref. 95 97a 97a 33 98 98 33 105 107 108 lOOd 110 110 110 13 87 87 87 45 45 13 42 130 - AG," 86-39 83.89 171.19 93-1 89.9 183-7 89.9 75.1 125.0 123.7 100.0 101.0 101.8 165.9 150.2 150.2 178.5 59.1 1 61-30 79.1 - 72-0 71.2 76.8 are based on calculations incorporating these new results. 97a For acetic acid Weltnerg8 made a similar detailed study treating the effects of dimer- isation and giving tabulated values of the thermodynamic functions for O6 Waring Chem.Rev. 1952 51 171. '13 Sinke J. Phys. Chent. 1959 63 2063. O7 Mulliken and Pitzer J. Chem. Phys. 1957 27 1305. O* Weltner J. Amer. Chem. SOC. 1955 99 3941. Green J. 1961 2241. 142 QUARTERLY REVIEWS both monomer and dimer. A revision of his values is possible to take account of recent measurementsgg of the heat of formation of the liquid; the more accurate value of Evans and Skinnerggb has been listed. Benzoic acid has been the subject of considerable study2J00 and the thermodynamic properties have been tabulated by Goton and Whalley,lOod whence the listed values are taken. For the higher aliphatic acids all the data are old,lol apart from a few measurements by Schj$nberglo2 and by Hancock Watson and Gilbeylo3 which are not of high accuracy.The values given in the Tables are derived from such revision as is possible of the work cited with correction to 25"c; they may be in error by several kcal./mole. However Coops1o4 reports that the heats of combustion of the fatty acids from valeric to eicosanoic have been remeasured by Adriaanse and these will no doubt replace all previous values. The listed entropy for palmitic acid (at 298.6"~) is presumably more reliable than the older value33 of 113.7 cal./deg./mole. For crystalline lactic acid a recent measurement of the heat of combustion is availablelo6 and a value for the entropy,lo7 whilst for the liquid state there are old values for both quantities.13J08 The values for m- and p-hydroxybenzoic acidslOgJ1o are of low accuracy; they have been con- verted into values for 25"c but the differences between the isomers do not exceed experimental error.For o-hydroxybenzoic acid the heat of com- bustion is rather better established; a summary of the available heats of combustion is given by Timmermans16 and from these the value -723.31 kcal./mole at 25"c has been selected from which the listed value of the heat of formation is derived. The heats of combustion of the homologous series of dicarboxylic acids HO,C.[CH 2]nC02H were measured by Verkade Hartman and Coopslll but by modern standards the results are subject to several uncertainties which it is difficult to allow for satisfactorily. Oxalic acid as was pointed out by Washburn in his classic paper,112 is subject to an extremely large standard-state correction which he calculated for the 99 (a) Stull Thermochem.Bull. I.U.P.A.C. 1956 No. 2 4; (b) Evans and Skinner Trans. Faraday SOC. 1959,55 260. loo (a) Jessup J. Res. Nat. Bur. Stand. 1942 29 247; (b) Challoner Gundry and Meetham Phil. Trans. 1955 247 556; (c) Ginnings and Furukawa J. Amer. Chem. Suc. 1953,75 522; (d) Goton and Whalley Canad. J. Chem. 1956,34 1506. lol Verkade and Coops Rec. Truv. chim. 1928 47 608 and references therein. lo2 Schjiinberg Z. phys. Chem. (a) 1935 A 172 197; (b) 1936 A 175 342; (c) 1937 A 178,274; (d) 1938 A 181,430. lo3 Hancock Watson and Gilbey J. Phys. Chem. 1954 58 127. lo4 Coops Bull. Chem. Thermodynamics I.U.P.A.C. 1959 No. 2 6. lo6 Ward and Singleton J. Phys. Chem. 1952 56 696. lo6 Saville and Gundry Trans. Furuduy Suc. 1959 55 2036. lo' Huffman Ellis and Borsook J.Amer. Chem. Suc. 1940 62 297. lo8 Parks Thomas and Light J. Chem. Phys. 1936 4 64. lo9 Keffler and Guthrie J. Phys. Chem. 1927 31 65. n1 Verkade Hartman and Coops Rec. Trav. chim. 1926,45 373. 112 Washburn Bur. Stand. J. Res. 1933 10 525. Parks and Light J. Amer. Chetn. SOC. 1934 56 151 1. GREEN THERMODYNAMIC PROPERTIES OF ORGANIC OXYGEN COMPOUNDS 143 results of Verkade Hartman and Coops. With modern units the heat of combustion becomes -59.7 kcal./mole at 25"c yielding the listed heat of formation. It may be noted that the value for the entropy of this substance is still that derived from the very old measurements discussed by Parks and Huffman.13 The heats of formation derived from the older measure- mentslll on both succinic and pimelic acids are not in good agreement with more recent values for succinic acid the listed value is that selected by Pilcher and Sutton113 after a discussion of their own and other results whilst for pimelic acid the listed value due to Stu11114 is to be compared with the old value of -240-1 kcal./mole.For malonic acid the listed value is derived from a revision of the old measurements,111 the standard- state correction being taken as 3.9 cal./g. The homologous series of monoalkylmalonic acids RCH(C02H)2 from R = ethyl to n-tetradecyl was investigated by Verkade and Coops116 and a more satisfactory revision can be made of this work. Some of the new values thus obtained* are listed. A similar investigation was made by Verkade and Hartman1l6 of the heats of combustion of fifteen methyl- ethyl- and phenyl-substituted succinic acids and their anhydrides.For fumaric and maleic acids the values measured by Schwabe and Wagner117 at 24"c are the most accurate and give a difference of 5.55 rt 0.16 kcal./mole between the two isomers the earlier value1I8 being 5.43 kcal./mole. In Table 9 is summarised the remaining recent work on the heats of formation of organic acids. For the toluic and dimethylbenzoic acids the listed values are those obtained by Colomina Perez-Ossorio and Boned;119 their value for o-toluic acid agrees exactly with that due to Brietenbach and Derkosch,120 but for p-toluic acid the latter authors found - 10 1.0 kcal./mole. The values for three forms of cis-cinnamic acid are due to Eisenlohr and Metzer,121 but the temperature of the measure- ments was not stated. Large discrepancies exist between the various measurements on (+)- (-)- and meso-tartaric acids ; the listed values are those of Dunken and Wolf122 but other authors give results differing from these by as much as 5 kcal./mole.The value for citric acid monohydrate is from a recent The value for crotonic which * Calculated by Dr. A. J. Head (private communication). llS Pilcher and Sutton Phil. Trans. 1955 258 23. 11* Stull Bull. Chem. Thermodynamics I.U.P.A.C. 1959 No. 2 6. 116 Verkade and Coops Rec. Trav. chim. 1933,52 747. n6 Verkade and Hartman Rec. Trav. chim. 1933 52,945. 11' Schwabe and Wagner Chem. Ber. 1958,91 686. 11* Huffman and Fox J. Arner. C'hem. Soc. 1938,60 1400. ll0 Colomina Perez-Ossorio and Boned Bull. Chem. Thermodynamics I.U.P.A.C. 120 Brietenbach and Derkosch Monatsh.1950 81 689; 1951 82 177. 121 Eisenlohr and Metzer 2. phys. Chem. 1937 A 178 339. 122 Dunken and Wolf 2. phys. Chem. 1938 B 38 441 and references therein. 123 Chappel and Hoare Trans. Faraday SOC. 1958,54,367. 124 Clopatt SOC. Sci. Fennica Commentations Phys.-Mat. 1932 6. 1. (a) 1959 No. 2 27; (b) 1960 No. 3 21. 144 QUARTERLY REVIEWS TABLE 9. Heats of formation of some acids and acid anhydrides at 2 5 " ~ (kcal./mole) Acid - AH," Ref. Propionic (1) 121.3 1 02u Valeric (1) 133.6 102c Valeric (I) 131.2 & 1-2 103 a-Methylbutyric (1) 133.1 f 1.2 103 Isovaleric (1) 134.8 i 0.8 103 Pivalic (s) 135.5 0.5 103 Pimelic (s) 242.75 rt 0-25 114 Ethylmalonic (s) 226.3 115 Propylmalonic (s) 232.1 115 Butylmalonic (s) 239.5 115 Citric mono- hydrate (s) 439.2 123 Malonic (s) 212.7 111 (+)-Tartaric (s) 368.7 1 20 (-)-Tartaric (s) 373.8 120 meso-Tartaric (s) 369.3 120 Crotonic (s) 83 1 24 Allylacetic (1) 101.8 & 0.2 102c p-Ethylidene- Acid 2,3-Dime t hyl- benzoic ( s ) 2,4-Dimethyl- benzoic (s) 2,SDimethyl- benzoic 2,6-Dimethyl- benzoic (s) 3,4-Dimethyl benzoic (s) 3,5-Dimethyl- benzoic (s) cis-Cinnamic m.p.42" (s) cis-Cinnamic m.p. 58" (s) cis-Cinnamic m.p. 68" (s) - AH," Ref. 107.65 & 0.19 119b 109.58 f 0.20 1196 109.02 f 0-19 119b 105.33 f 0.19 119b 112.04 & 0.25 119b 111.48 f 0.17 119b 75.3 & 0.2 121 73.9 * 0.2 121 71.5 & 0.2 124 propionic (1) 102-8 & 0-2 102c Acetic anhydride (1) 105.6 i- 0.2 102c Maleic anhydride rn-Toluic (s) 101.90 f 0.25 119a Succinic anhydride p-Toluic (s) 102.58 f 0-24 119a (s) 143.2 116 Propylideneacet ic (1) 149.20 127 0-Toluic (s) 99.5 i 0.22 119a (s) 1 12-23 35 differs by 20 kcal./mole from that tabulated by Kharasch,14 is obviously unreliable.The available data on unsaturated fatty acids are old and satisfactory revision of them is not possible; only the work of Keffle~-l~~ is more recent than that summarised by Kharasch14 and all the work is discussed by Markley.lZ6 Only for phthalic anhydride are there measurements of both the heat (3f formation3s and the entropy45 at 25 OC. Values for the heats of formation of some anhydrides can be derived from the heats of hydrolysis measured by Kistiakowsky and his c o - w ~ r k e r s ~ ~ ~ if the heats of formation of the acids are known; the listed value for acetic anhydride was obtained in this way the new valueggb being used for the heat of combustion of acetic lZ6 Keffler J.Phys. Chem 1930,34 1319; Rec. Trav. chim. 1933,52 945. lZ6 Markley "Fatty Acids:" Interscience Publishers Inc. New York 1947. lZ7 Conn Kistiakowsky Roberts and Smith J. Amer. Chem. Soc. 1942,64 1747. GREEN THERMODYNAMIC PROPERTIES OF ORGANIC OXYGEN COMPOUNDS 145 acid. For maleic anhydride the listed heat of formation of the acid being again used we find AH," (s) = -1 11.63 kcal./mole in reasonably good Lgreement with the value of - 112.23 kcal./mole obtained from the heat of combustion of the anhydride. However for succinic anhydride we find AH," (s) = -145.25 kcal./mole compared with the value of -143-2 kcal./mole derived from a revision of the heat of combustion measured by Verkade and Hartman,llG and the heats of hydrolysis of the methyl- succinic acids are in rather poor agreement with the differences between the heats of combustion of the acids and the anhydrides measured directly.(f) Esters.-Entropy values are available only for methyl formate and ethyl acetate and for the former the value has to be obtained by indirect means. The listed value in Table 8 for methyl formate was derived by a revision of the discussion by Parks and Huffman13 of the equilibria between methyl alcohol hydrogen and methyl formate and between carbon monoxide methyl alcohol and methyl formate the modern value being used (Table 1) for the free energy of formation of gaseous methyl alcohol at 25"c. For the heat of formation of methyl formate the heats of com- bustion measured by Roth and Banse12* have been utilised; they are probably more accurate than the very old measurements used in a previous c0mpi1ation.l~ The entropy of ethyl acetate was measured by Parks Huffman and B a r m ~ r e ~ ~ but no reliable value for the heat of formation is available old sources14J02aJ29 yielding results differing by 2 kcal./mole.A of the equilibrium between ethyl acetate hydrogen and ethyl alcohol at two temperatures (18 1 " and 201 -5 "c) gave for the ester the listed values for the gaseous state but these are not in good agreement with those which can be derived from the liquid-state values and heat of vaporisation and vapour-pressure data cited by the authors. The heats of formation of numerous fatty acid esters are available from the work of Schj&berg,lo2 and the values for the liquid substances corrected to modern units at 2 5 " c are listed in Table 10.Values for the heats of vaporisation partly from old were also given by Schjinberg. Also included in Table 10 are revised values for esters of oleic and elaidic acid from measurements by K e f f l e ~ . ~ ~ ~ Other recent measurements of the heats of combustion of esters are as follows dimethyl tartrate and dimethoxysuccinate,122 a- and /3-naphthyl acetates and ,8-naphthyl benzoate,55 the four ethyl pentynoates,102d and methyl a-t-butylacrylate and a-t-b~ty1propionate.l~~" (g) Peroxides and Miscellaneous Compounds.-Heats of combustion of peroxides are inevitably difficult to measure accurately ; the available 128 Roth and Banse quoted in Landolt-Bornstein I1 Eng. Bd. 1931 p. 1644. 12B Berenger-Calvet J. Chim. phys. 1927 24 325. 130 Vredensky Ivannikov and Nekrasova Zhur.obshchei Khim. 1949 19 1094 lS1 Brown J. 1903 987. l** Keffler J. Phys. Chem. 1937,41 715 182a Crawford and Swift J. 1952,7220. 146 QUARTERLY REVIEWS TABLE 10. Heats of formation (- AHf ") of liquid esters R'C02R" at 2 5 " ~ (kcal. Imo le) R' ... .. . Me Et Prn Bun C,H," C4H C4H7G C4H7d C&f33' C17H33f R" Me 106-2 111.6 117-2 - Et 113-9 119.9 126.3 130-8 Prn 119-3 124.9 132.1 137.3 Pr* 122.0 128.2 133.9 140.0 Bun 125-5 130.1 138-1 145.1 Bu' 1273 133.7 139.8 147-2 Bus - - 136-0 148.3 CBHlln - - - - CSH11' 133.6 139.0 146.0 - 90.6 99- 1 104.6 107.8 1 10.2 1 12-7 112.9 117.6 - - - 102.0 103.2 105.4 107.1 109-2 111.5 113-8 114.3 114.8 116.8 116.8 118.2 - - - 173.7 175.3 105-4 184.7 184.1 109-5 188.5 190.0 113.9 - - 116.3 194.3 196.2 119.2 - - 120.1 - - 202.4 - - (a) MeCH=CH; (b) CH2=CHCH2CH,; (c) MeCH=CH.CH,; ( d ) MeCH,CH=CH; (e) cis-Me.[CH,],CH=CH.[CH,],; (f) rrum-Me.[CH,],CH=CH.[CH,],.TABLE 11. Heats of formation of liquid peroxides and hydroperoxides (kca I./mole) -AH," Ref. Ethyl peroxide 55-6 133 Propyl peroxide 76 f 15 134 t-Butyl hydro- peroxide 63.8 f 0.3 137 Propionyl peroxide 148.2 f 1.6 135 Benzoyl peroxide 93-5 1 20 O-TOlUOyl peroxide 119-4 120 p-Toluoyl peroxide 107.7 120 -AH," Ref. Ethyl hydro- peroxide 58 f 12 134 t-Butyl peroxide 94-0 136 Acetyl peroxide 127.9 f 2.4 135 Butyryl peroxide 161.0 f 1-1 135 Cinnamoyl peroxide 84.7 1 20 Decalin hydro- peroxide 83-2 120 Tetralin hydro- peroxide 44.6 1 20 heats of formation so are summarised in Table 1 1. For ethyl peroxide the value obtained by Zihlman is more reliable than the older value13* but these old values are the only ones available for ethyl and propyl hydroperoxides.The quinol-p-benzoquinone system was studied by Pilcher and S ~ t t o n ~ ~ who give the following values for the solid substances at 25"c (kcal./mole and cal./deg./mole) Q - AH," = 44.65 rt 0.17; So = 38-55; - AG," = 20.49 QH2 - AH," = 87-51 It 0.28; So = 32.77; - dGj" = 52.32 13s Zihlman quoted by Rabbert and Laidler J . Chem. Phys. 1952 20 574. 13* (a) Stathis and Egerton Truns. Furuduy SOC. 1940 36 606; (b) Harris Proc. 136 Jaffe Prosen and Szwarc J. Chem. Phys. 1957 27 416. 137 Bell Dickley Rayley Rust and Vaughan Ind. Eng. Chem. 1949 41 2597. Roy. Soc. 1939 A 173 126. Rayley Rust and Vaughan J. Amer. Chem. SOC. 1948 70 88. GREEN THERMODYNAMIC PROPERTIES OF ORGANIC OXYGEN COMPOUNDS 147 and the thermodynamic properties for the oxidation-reduction process in the solid solution and gaseous states.Rather different values for the heats of combustion were reported by MagnuslW who also gives values for a number of other quinones. Heats of combustion have been reported for paraldehyde,141 and for the condensation product of ethylene glycol with acetone and of cis- cyclopentane- 1 ,Zdiol with acetone.52b Regularities and Estimation Methods Regularities in the thermodynamic data for related compounds and homologous series form the basis of the numerous methods of estimation that have been proposed. Since the previous review,lob the important work of Benson and has provided a unification of procedures involving additivity of atomic values bond values and group values.However the paucity of reliable information on oxygen compounds still prevents the establishment of accurate methods for the estimation of their thermo- dynamic properties. A survey of the simpler and less accurate methods for estimating heats of combustion should be Several empirical correlations between various groups of compounds were pointed out by Parks and Huffman,13 and further observations made in the light of more recent data.43 Similarly a constancy was claimed for the difference between the heats of combustion of furan and benzene and of a number of their derivatives with the same substituent ;58 whilst Bad~che'~ reported regularities in the heats of combustion of benzene derivatives on substitution with methyl and methoxyl groups. More accurate work reveals that such observations are only first approximations however.Empirical equations involving structural parameters only have been developed to express the isomeric variation of the thermo- dynamic properties of hydrocarbons and applied to some physical pro- perties of the saturated aliphatic An alternative approach makes use of comparisons with the compre- hensive and accurate information available for the hydrocarbons. For example increments in the entropies43 and heat capacities39 of normal alcohols have been compared with the corresponding increments for the paraffins. It is now well established that the increment for the methylene group for the normal paraffins mono-olefins normal alkylbenzenes normal alkylcyclopentanes and normal alkylcyclohexanes is -4.926 kcal./mole in the heat of formation or - 157.44 -C 0.05 kcal./mole in the 138 Magnus 2.phys. Chem. (Frankfurt) 1956,9,141. 139 Hubbard Katz Guthrie and Waddington J. Amer. Chem. SOC. 1952 74 4456. 140 Tonaka and Watase Technol. Reports Osaka Univ. 1956 6 367. 141 (a) Cass Springall and White Chem. and Ind. 1955 387; (b) DelCpine and Badoche Conzpt. rend. 1942 214 777. 142 Benson and Buss J. Chem. Phys. 1958 29 546. 113 Handrick Ind. Eng. Chem. 1956,48 1366. 144 Greenshields and Rossini J. Phys. Chem. 1958 62 271. 148 QUARTERLY REVIEWS heat of combustion in the gaseous state provided the number of carbon atoms in the chain is greater than five.lob For the normal paraffins the increment per methylene group is 1.18 kcal./mole in the heat of vaporisa- tion at 25"c. So the increment becomes -156.26 * 0-05 kcal./mole in the heat of combustion for the liquid state.It is noteworthy that a very similar increment is established for the higher members of a number of homolo- gous series of oxygen compounds Increments per methylene group in - AH," (1,25"c) (kcal./mole) Normal alcohols28 156.26 Fatty acids (C5-C9)lo4 156.26 0-17 Methyl esters of fatty acids (CgC16)104 156.26 =t 0.10 Dimethyl esters of oxalic acid series ( 2 0 " ~ ) ~ ~ ~ 156.3 Esters of oleic and elaidic acid 156.2 Esters of fatty acids (C,-C,) 156.3 A 1.0 The identity of the increments for the higher fatty acids and their methyl esters is particularly important in showing that the degree of association of all members of the series in the liquid state at 25"c is the same. For the solid state the increment is more variable as would be expected thus for the alkylmalonic acids it is 155.2 III 0.4 kcal./mole (data115 revised to basis of 25"c) whilst for the oxalic acid serieslll there is an alternation between the values 158.3 =k 0.2 and 155.1 h 0.4 kcal./mole.An expression relating the heat of combustion of alcohols to the number of carbon atoms present has been derived by M a ~ l o v . ~ * ~ ~ The increment is given as -156.246 kcal./mole but the values for individual alcohols are all numerically higher than those recently quoted.28 The same author has also given14sb expressions for the vapour heat capacities of a number of oxygen compounds which can be written as C," = a + bT + cT2 + dT3 + nC,(T) + f(n m) exp(-O-O07T) where n and m are the total number of carbon and hydrogen atoms respectively and C,(T) is the increment per methylene group given by C2(T) = 0.1203 + 21.3 x 10-3T - 116.33 x 10-7T2 + 2.502 x 10-'T3 The values for the constants a b c and d and for the function f(n,m) were given as Alcohols -6,602 17.8 139.85 -12.51 242.2~~1 Aldehydes -6.882 5-98 211.88 -15.01 215(n2+m) Acids -5.692 18.2 122.48 -15-012 127.56 n-lrn-l a 103b 107c 109d f(n,m> Esters - 14-54 24.2 275.8 -25.02 14(1O-~) 145 Verkade Coops and Hartman Rec.Trav. chim. 1926 45 585. 146 Maslov (a)Zhur.jiz. Khim. 1955,29,718 (b) Zhur.prikIad. Khim. 1957,30,736. GREEN THERMODYNAMIC PROPERTIES OF ORGANIC OXYGEN COMPOUNDS 149 In the last expression for f(n,m) z is the number of methylene groups. According to Maslov the differences between the calculated and observed values for the first three members of these series may be about 7% but decrease rapidly with increase of n ; it is not clear however on what experimental data these statements are based.A partial application of the method of group equations to organic oxygen compounds was made by Van Krevelen and Chermin14’ but it can be noted that some of the results are in less satisfactory agreement with more recent experimental data notably for acetone. From the appropriate group equations it would be expected that the differences in the heats of formation of the pairs (butane-isobutane) (pentane-isopentane) (butane- 1 -thiol-butane-2-thiol) (propane-1-thiol-propane-2-thiol) and (propyl alcohol-isopropyl alcohol) would be equal. For the first four pairs the values1** are 2.00 & 0.10 kcal./mole but for (propyl alcohol-isopropyl alcohol) and (butyl alcohol-s-butyl alcohol) are 3.7 1 and 3.2 kcal./mole respectively the data of Table 1 being used.This increase which has been related to the electronegativity of the s ~ b s t i t u e n t ~ ~ ~ indicates an important limitation in the application of the method of group equations. Lovering and Laidler149 have recently given a scheme of thermochemical bond energies for organic oxygen compounds bond contributions to heats of atomisation formation and combustion were obtained for the gas at 2 5 ” ~ . The contributions to the heats of formation for C-C and for primary secondary and tertiary C-H bonds -0.45 3-45,2.63 and 1.78 kcal./mole respectively were taken as those previously found for hydro- carbons. (These are equivalent to an increment per methylene group in the heat of combustion of 157.56 kcal./mole for the gas or 156.40 kcal./mole for the liquid.) It was assumed that the attachment of an oxygen atom to carbon affects the bond strength of hydrogen bound to that carbon atom.Such C-H bonds can be primary secondary or tertiary and make the following contributions to the heat of formation 1-22 0.40 and -0.48 kcal./mole respectively. The C-0 bond is assumed to have no effect on other C-H bonds; its own contribution to the heat of formation is 18.48 kcal./mole that of the 0-H bond is 28.12 kcal./mole and that of the 0-0 bond in peroxides - 1 1-44 kcal./mole. These values reproduce the heats of formation of the compounds from which they were derived to 1 kcal./mole or better but it may be noted that the “observed” values utilised differ in some cases from those tabulated here.Applications of Thermochemical Data The most obvious practical application of thermodynamic data is to the calculation of the equilibrium constants for known or hypothetical 14’ Van Krevelen and Chermin Chem. Eng. Sci. 1951 1 66. 148 Hubbard and Waddington Rec. Trav. chim. 1954 73 910. 14* Lovering and Laidler Canad. J. Chem. 1960 38 2367. 150 QUARTERLY REVIEWS reactions and the estimation of the feasibility of a given process. An example is the calculation by Mazurek15* of the equilibrium constants for numerous reactions involving ethyl alcohol acetic acid and ethyl acetate and for the formation of acetone from these compounds leading to the conclusion that the ketonisation of ethyl alcohol at 250-450" is essentially irreversible.Two other important applications are the calculation of bond energy terms and the determination of barriers to free internal rotation. Bond Energy Terms.-Prediction of heats of combustion and formation can be made by using a self-consistent scheme of thermochemical bond energies. The latter quantities are also important in obtaining values for the resonance energy of substances by comparing the observed heat of forma- tion with that computed for the individual valency bond structures. Springall and his co-workers have recently used such an approach in studying a number of ketoness1 and cyclic ethers,74 obtaining quantitative expression for the molecular stabilisation or strain. The significance of the concept of bond energy has been reviewed by Cottrel13 who lists values for several bond-energy terms; these may be revised slightly in terms of the data presented in this paper.For the heats of the gaseous atoms from the elements in their standard states at 25"c we use the values3 C (graphite) 170.9; iH,(g) 52-09; $0 (g) 59.54 kcal./mole. The value of the bond energy term E(0-H) deduced from these values and the heat of formation of gaseous water17 at 2 5 " ~ (-57.80 kcal./mole) is then 1 10.76 kcal./mole. Following Coates and Sutton,lsl the values for E(C-C) and E(C-H) are obtained from the constant in- crement in the heat of formation of the higher paraffins and are 82-60 kcal./mole and 98.70 kcal./mole respectively. Using the new values28 for the heats of formation of the gaseous alcohols we now find E(C-0)is 80.01 in methyl alcohol 83-25 in ethyl alcohol 83.95 in propyl alcohol and thereafter cqnstant at 84.1 kcal./mole up to decyl alcohol.In some acetals the value of E(C-0) is slightly higher being 86.1 in formaldehyde dimethyl acetal 86.5 in formaldehyde diethyl acetal and 86.0 in acetal- dehyde dimethyl acetal but the heats of c o m b ~ s t i o n ~ ~ J ~ are rather un- reliable. It is likewise difficult to establish accurate values for E(C-0) in ethers owing to the absence of reliable values for their heats of com- bustion. For dimethyl ether ethyl methyl ether and diethyl ether E(C-0) is about 84 kcal./mole [actually 84.8 in diethyl ether the new value (p. 135) being used for the heat of formation]. In tetrahydrofuran we find E(C-0) = 84-1 kcal./mole from the heat of combustion obtained by S ~ r i n g a l l ~ ~ ~ or 85.2 kcal./mole from the value given by S k ~ r a t o v .~ ~ Cottrell has pointed out that E(C=O) is more variable being 164.9 kcal./mole in formalde- hyde 172 kcal./mole in acetaldehyde 179.1 in acetone and very slightly higher in some other ketone~.~l From the heat of combustion of diethyl 150 Mazurek Zhur. obschei Khim. 1952,22 1324. Coates and Sutton J. 1948 1187. GREEN THERMODYNAMIC PROPERTIES OF ORGANIC OXYGEN COMPOUNDS 15 1 E(0-0) is found to be 50.9 kcal./mole E(C-0) being assumed to be 84-1 kcal./mole in this compound. A rather different approach is adopted by Allen152 in recent work since the carbon-hydrogen bond energy has been taken as 99.29 kcal./mole and assumed constant throughout a range of hydrocarbons. The same author has also given a method for including the effects of next-nearest neighbour interactions.Mention must also be made of the work of G1ockle1-l~~ who has established empirical relations between bond energies and bond lengths and extended this to carbon-oxygen bonds. Barrier Heights to Internal Rotation.-It is well known that one of the oldest and most successful methods for determining barrier heights to internal rotation is by a comparison of calculated and observed heat capacities and entropies and the results obtained by this method have been r e ~ i e w e d . l ~ ~ J ~ ~ In recent years microwave spectroscopy has provided two new one of which the frequency method is capable of giving results which are usually more accurate than those obtained from thermo- dynamics whilst the other the intensity method gives results of accuracy TABLE 12.Barrier heights to internal rotation (cal./mole> obtained by thermodynamic and microwave methods Ethane Propane Isobu t ane Neopen t ane Propene Methyl alcohol Ethyl alcohol Phenol Dimethyl ether Propylene oxide Acetaldehyde Acetone Acetic acid Methyl formate Thermodynamic 2875 125 3400 3 620 4300 1950 1070 3300 800 2700 lo00 - - 1000 2500 f 700 Ref. 154a 1 54a 154a 154a 154a 24 27 1546 79b 84 98 Microwave 2875 f 125 3900 I 1978 f 17 1070 3140 2720 2560 f 70 1150 f 30; 1103 f 60 - 760 483 1190 & 40 Ref. 156 157 158 24 50 159 160 161 1 62 163 164 comparable to those obtained from thermodynamics. The thermodynamic procedure involves the difference between two relatively large quantities and so it is very sensitive to errors in both of them-that is in the vibrational assignment and in the measured quantities.It is important 162 Allen J . Chem. Phys. 1959 31 1039. 153 Glockler J . Phys. Chem. 1958 62 1049 and references therein. 164 (a) Pitzer Discuss. Furaday Soc. 1951,10,66; (b) Aston ibid. p. 73. 156 (a) Lin and Swalen Rev. Mod. Phys. 1959 31 841; (6) Wilson “Advances in Chemical Physics,’’ lnterscience Publishers Inc. New York 1959 Vol. 11 p. 367; (c) Proc. Nat. Acud. Sci. U.S.A. 1957 43 816; ( d ) Pauling ibid. 1958 44 21 1. 152 QUARTERLY REVIEWS therefore to compare the final values obtained and these156164 are sum- marised in Table 12 for some hydrocarbons and all organic oxygen com- pounds for which data are available. For methyl alcohol the agreement is exact since the results of several detailed studies by the microwave method were used to obtain calculated thermodynamic functions in agreement with ob~ervation.~~ On the other hand the values for acetic acid are considerably different and further investigation is required.156 Lide J. Chem. Phys. 1958 29 1426. 15' Lide and Mann J. Chem. Phys. 1958,28,914. 158 Lide and Mann J. Chem. Phys. 1957,27 868. 158 Kasai and Myers J. Chem. Phys. 1959 30 1096. 160 Herschbach and Swalen J. Chem. Phys. 1958,29 761. 162 Swalen and Costain J. Chem. Phys. 1959 31 1562. 163 Tabor J. Chem. Phys. 1957 27 974. 164 Curl J Chem. Phys. 1959 30 1529. Kilb Lin and Wilson J. Chem. Phys. 1957 26 1695; Verdier and Wilson ibid. 1958 29 340.

ISSN:0009-2681    

DOI:10.1039/QR9611500125

出版商:RSC    

年代:1961

数据来源: RSC

摘要:

THE HISTORY AND CHEMISTRY OF MUSCARINE By S. WILKINSON (THE WELLCOME RESEARCH LABORATORIES BECKENHAM KENT) THE fungus Amanita muscaria (L. ex Fr.) Qukl. with the unmistakable brilliant orange-red cap flecked with white fragments of its volva is found in birch and pine woods from summer to late autumn. The epithet muscaria refers to its former use as a fly poison and hence the common name “fly agaric”. The genus Amanita is said to be probably responsible for the majority of cases of mushroom poisoning and although the deadly A . phalfoides causes most fatalities occasionally A. muscaria has proved fatal. The fungus contains a number of basic substances including choline acetylcholine and the intriguing alkaloid muscarine. Muscarine is a foundation stone of modern pharmacology being one of the first sub- stances known which reproduced faithfully some of the responses to stimulation of the parasympathetic nervous system.Dixonl was so im- pressed by this activity that he advanced the hypothesis that the vagus nerve liberated a muscarine-like substance which acted as a chemical transmitter of its impulses and this was shown to be substantially correct by Hunt and Taveau2 and by Dale3 in their classical studies of acetyl- choline. The similarity between the actions of acetylcholine and muscarine on smooth muscles and glands gave rise to the definition “muscarinic” actions of acetylcholine to distinguish them from the effects on ganglia and voluntary muscles the so-called “nicotinic” actions. It was obviously of fundamental importance that muscarine should be isolated and charac- terised so that comparative pharmacological studies and absolute assess- ment of its potency could be made.Muscarine however proved extra- ordinarily elusive and only recently almost 150 years after the first approaches were made in 1811 by Braconnot* and S~hrader,~ have its isolation and the determination of its structure been achieved. Although one tends to associate muscarine with A . muscaria it has been isolated from or muscarinic activity demonstrated in a variety of other fungi. It has invariably been fly agaric which is prolific that has been extracted but far higher concentrations of muscarine occur in fungi of Dixon Med. Mag. 1907 16 454. Hunt and Taveau Brit. Med. J. 1906 11 1788. Dale Science 1939 90 393; Dale Feldberg and Vogt J. Physiol. 1936 86 353. Schrader cited by Hamsen Arch.exp. Path. Pharmak. 1903,50 311. * Braconnot Ann. Chim. (France) 1811 79 265; 1813 87 237. 153 154 QUARTERLY REVIEWS other genera. As far as the Reviewer is aware A. pantherina D.C. ex Fr. is the only other species of Ainanita from which the alkaloid has been isolated.6 Activity has been found in extracts of Boletus Zuridus Schaeff. ex Fr.,’ Russula emetica Schaeff ex Fr.,8 Clitocybe rivulosa Pers. ex Fr. and other Russulas and CZitocybe.lo The widest distribution however seems to occur in the genus Inocybe. Wickill and Loup12 have recorded activity in 26 out of 33 species. I. asterospora rimosa cookei and umbrina were stated to have potent muscarinic activity.13 On the basis of pharma- cological assays of I. lateraria Ricken (syn. I. patouillardi Bres.) and I .napipes Lange incredibly high concentrations of m ~ s c a r j n e ~ ~ J ~ J * ~ ~ ~ were recorded probably as a result of the simultaneous assay of acetylcholine. The most reliable figures have recently been obtained by Eugster and his colleagues16 by actually isolating the muscarine. Their following percentage yields show the wide variation in the different species A. nuiscuria G.0002 % ;* I. patouillardi 0.037 % ; I. fastigiata 0.01 % ; I. umbrina 0.003 % ; I . bougardi 0.0 %. The early work has been reviewed previously,17 but in order to present a clear picture it is essential that this phase of the problem be described briefly. In reality the story of muscarine revolves round the presentation of three formulae two erroneous ones C5H,,02N+ and C8HI8OzN+ and the correct one C9H2,,02N+ and this Review is sub-divided to follow these trends.It is noticeable that with the advent of more refined chemical techniques particularly chromatography and ion-exchange many of the principle difficulties associated with the isolation of a minute amount of muscarine from a large mass of material and its separation from closely related compounds were resolved. Many of the doubts and controversies which arose repeatedly were the result of statements which related to chemically impure preparations and based on inadequate analytic21 data. Once pure muscarine had been obtained the elucidation of its structure and its synthesis followed very quickly. * Average yield from batches of A. muscaria collected over a period of 3 years. Inoko Annalen med. fac. Jap. Univ. Tokyo 1887-1889 1 227; J.1892 62 232. Boehm Arch. exp. Path. Pharmak. 1885 19 87. * Kobert St. Petersburg med. Wochenschr. 1891 51 463 cited by Harmsen (see ref. 5); Kobert “Lehrbuch der Intoxikationem.” Vol. II. Euke Stuttgart 1906 p. 1288. Carter Amer. J. Physiol. 1901 5 158. lo Wicki and Loup Trav. Lab. Thzrap. exp. Genzve 1930-1932 13 9; Wicki Schweiz. 2. Pilzk. 1930,8,42; 1931,9,78; Wicki and Roch Rev. mPd. Suisse 1935,55 896. l1 Wicki Bull. SOC. Mycol. GenPve 1928 11 14. la Loup Thesis 114 Geneva 1938. l3 Yasumari Ishida and Kozu J. Appl. Mycol. 1949 3 118; Ford and Sherrick l4 Mecke Arch. exp. Path. Pharmak. 1934 175 23. l5 Fahrig Arch. exp. Path. Pharmak. 1920 88 227. l6 (a) Eugster Helv. Chim. Ada 1957 40 886; (b) Eugster and Muller ibid, 1959 l’Bowden and Mogey J. Pharm. Pharmacoi.1958 10 145; Salemink Pharm. J. Pharmacol. 1911 2 549; 1913 4 321. 42 1189. Weekblad 1960,95 165 197. WILKINSON THE HISTORY AND CHEMISTKY OF MUSCARINE 155 Phase 1. The formula C5H,,02N+ Vauquelinls suspected that the toxicity was associated with the fatty contents of the fungus. LetellieP considered that there was a toxic principle common to species of Arnanila which he referred to as amanitine a name now applied exclusively to toxin isolated from A . phaZZoides.20 Although several subsequent investigators21 realised that toxicity was related to the basic constituents the first classical studies both chemical and pharmacological were made in 1869 by Schmiedeberg and Koppe22 who isolated a syrupy alkaloid having potent physiological activity and arresting the isolated heart in diastole.This preparation was obviously impure and by fractionating the mixture of chloroaurates obtained from it H a r n a ~ k ~ ~ isolated choline and a substance to which he assigned the formula C,H,,AuC14N02. To this material he gave the name muscarine and s~bsequently~~ suggested that it had the structure (l) whose relation to choline (2) is obvious. This structure was apparently substantiated by synthesis. From choline by the action of nitric acid was obtained the so-called “synthetic muscar- ine” to which structure (1) was assigned and which was reputed to be identical with natural m~scarine.~~ Nothnage126 repeated this work with equally convincing proof of identity. However as techniques became more refined pharmacologists began to doubt these observations. Boehm2’ showed that the “synthetic muscar- ine” had a curarising activity not found in the natural preparation and equally subtle differences were described by Meyer28 and This doubt existed until 1914 but in the meantime varying success had attended other attempts to isolate muscarine.Inokos prepared an active la Vauquelin Ann. Chim. (France) 1813 85 25. Letellier Inaug. Thesis Paris 1826 cited by Kobert “Lehrbuch des Intoxika- tionem,” Enke Stuttgart 1906 Vol. 11 p. 1288. 2o Wieland and Hallermayer Annalen 1941 548 1. a1 Sicard and Schoras Compt. rend. 1865,60 847; Letellier and Speneux Ann. hyg. publ. et mid. Iigale 1876,27 71; Apoiger Buchners ReperJJ d. Pharm. 1851,7 289; Kaiser Inaug. diss. Gottingen 1862; Boudier “Des Champignons au Point de Vues de leurs Caractkres Usuelles Chimiques et Toxicologiques,” Paris 1869.p 2 Schmiedeberg and Koppe “Das Muscarin das giftige Alkaloid des Fliegenpilzes,” Vogel Leipzig 1869. 23 Harnack Arch. exp. Path. Pharmak. 1875,4 168. 24 Harnack and Schmiedeberg Zentr. med. Wissensch. 1875,36,598 ; Arch. exp. Path. Pharmak. 1877 6 101. 26 Schmiedeberg Arch. exp. Path. Pharmak. 1881 14 376. 2e Nothnagel Ber. 1893 26 801; Arch. Pharm. 1894 232,261; J. 1893.64 297; 1894 66,437. 27 Boehm Arch. exp. Path. Pharmak. 1885 19 60. 26 Meyer Ber. 1893 26 803. ** Schmidt Annalen 1904 337 37. 156 QUARTERLY REVIEWS fraction from A . pantherim and potent material was obtained by Harmsen30 and Honda,al whilst further extensive pharmacological comparisons were made by H a r m ~ e n ~ ~ S t r a ~ b ~ ~ and Fiihner.33 Finally however E ~ i n s ~ ~ re-examined the method of preparation of “synthetic muscarine” and proved that in fact the product was the choline nitrous ester (3) and not the derivative (1).The pharmacological properties of the nitrous ester were clearly shown by Dale35 and independently confirmed by Wei~~hagen~~ as distinct from those of muscarine and identical with those described earlier for “synthetic muscarine”. Phase 2. The formula C8H1802N+ In 1922 King,37 after critically examining previous procedures and being concerned with the preponderance of choline in the basic fractions drastically modified existing methods of isolation and obtained 90 mg. of a non-crystalline muscarine chloride from 253 kg. of fly agaric. With the possible exception of Honda’s preparati~n,~~ with which it was eqili- active this was undoubtedly the purest sample which had as yet been obtained.King however other than establishing the equivalent weight of 210 for the base by estimating the gold in the chloroaurate and hence eliminating the C,HI40,N+ formula added little to the chemical know- ledge. He could find no evidence that muscarine was quaternary and was therefore in sympathy with those who doubted its formulation as a trimethylammonium base.38 Until 1931 there was a further period of inertia then Kogl and his c011eagues,~~ using Permutit as a means of concentration and fractiona- tion of the reineckates as a method of purification isolated from 1250 kg. of fungus 137 mg. of a crystalline reineckate which they considered to be the pure salt of muscarine. A crystalline chloroaurate was prepared and the analytical figures for these two salts corresponded to the formula C8HI8o2N+.The chloride regenerated from the reineckate by the method of Kapfhammer and Bischoff40 was optically active { [ 0 1 ] ~ ~ + 1-57’ (in water)) and was remarkably stable to alkali but its activity was reduced by one half on prolonged contact with acid. In contrast to Scelba’s earlier observation^,^^ the chloride was found to give reactions 30 Harmsen Arch. exp. Path. Pharmak. 1903 SO 361. 31 Honda Arch. exp. Path. Pharmak. 1911,65,454; Chem. Abs. 1912,6 529. 32 Straub Pfliig. Arch. ges. Physiol. 1907 119 127. 32 Fiihner Arch. exp. Path. Pharmak. 1908 59 179. 34 Ewins Biochem. J. 1914 8 209. 35 Dale J. Pharmacol. 1914 6 147; Dale and Ewins J. Physiol. 1914 48 24. 36 Weinhagen Zphysiol. Chem. 1919,105,249; 1921,112 13; J.Amer. Chem. Soc.. 1920 42 1670. 37 King J. 1922 121 1743. Guth Monatsh. 1925 45 631; Kiing 2. physiol. Chem. 1914 91 241; Heinisch and Zellner Monatsh. 1904 25 537; Zellner ibid. 1911 32 133. 39 Kogl Duisberg and Erxleben Annalen 1931,489 156. O0 Kapfhammer and Bischoff 2. physiol. Chem. 1930 191 182. 41 Scelba Atti Accad. naz. Lincei Rend Clusse Sci. fis. mat. nat. 1922 31 518; Chem. A h . 1923,17 3162. WILKINSON THE HISTORY AND CHEMISTRY OF MUSCARINE 157 characteristic of an aldehyde. The presence of a hydroxyl group was shown by formation of a benzoyl derivative. By Hofmann degradation were isolated the two products which have caused so much difference of opinions namely trimethylamine and the optically active orp-di hydroxy- valeric acid (4). By re-assembling these fragments Kogl et aZ.deduced that muscarine chloride must be a quaternary salt represented by one or other of the two structures (5) and (6) of which (5) was the more probable because of its derivation from serine. The acid (4) was assumed to arise by oxidation with the silver oxide during the degradation. A racemic mixture OH Me CH,.FH*FH CO H Me CH,-CHCH CHO Me CH,. $H *kH - CHO HO OH HC) 'hMe C1- Me3N+ C l - (4) (5) (6) of aldehydes ( 5 ) was synthesised but was 40,000 times less active than mu~carine.~~ Because the separation and the resolution of the two possible racemates of (5) proved impracticable it could not be shown that this low activity was a result of stereo~pecificity.~~ Phase 3. The formula C9H,,02N+ The aldehyde structure deduced by Kogl was not well received.Betaine aldehyde ('7) and a number of related compounds had long been syn- the~ised*~ and found to have the nicotine-curare type of action which is so completely divorced from that of muscarine. + + + Me,N.CH,CHO H,C-CHCH,.NMe I I MeCH .FH .CH .CH,-O.NMe c1- 0 ,o c1- HO OH c1- ( 7 ) CHMe ( 8 ) (9) Fourneau et aZ.45 synthesised the dioxolan derivative (8) with high muscarinic activity but equally as non-specific as many other synthetic compounds. Rogers et ~ 1 . ~ ~ suggested that muscarine might have an alkoxytrimethylammonium structure such as (9) which would account for the formation of the dihydroxy-acid (4) on degradation. However this was a period of speculation and elimination by analogies. Close examination of Kogl's results particularly the inadequate analytical data and absence of any proof as to the homogeneity of the fractions clearly shows on what precarious grounds his formula was based.42 Kogl and Veldstra Annalen 1942 552 1 ; Veldstra Diss. Utrecht 1935. 43 Van der Laan Diss. Utrecht 1942. 44 Berlinerblau Ber. 1884 17 1139; Fischer Ber. 1893,26,464. 45 Fourneau Bovet Bovet and Montezin Bull. SOC. Chim. bid 1944 26 516; 48 Rogers Bovet Longo and Marini-Bettolo Experienlia 1953 9 260. Fourneau and Chantalou Bull. Soc. chim. France 1945,5 10. 2 158 QUARTERLY REVIEWS Several groups of workers realised almost simultaneously that the only solution to the problem lay in isolating a sufficient amount of muscarine for comprehensive chemical and physical analyses and moreover that the recent applications of chromatography might prove invaluable.To Eugster and Waser however must be given the credit of describing the first pure crystalline muscarine chloride. In a preliminary note,*' followed later by full chemical4* and pharmacological details,49 they described the successful applicatiop of partition chromatography on cellulose columns to the separation of the crude bases. The homogeneity of their fractions was evaluated by controlled toxicity tests and by paper chromatography a modified Dragendorff's reagent50 being used to locate the bands. The analyses of their crystalline chloride chloroaurate and reineckate salts of muscarine each corresponded to a new empirical formula CgH2002N+ X-. The optical rotation [a],2o + 6.7" (in water) was considerably higher than that recorded by Kogl et aZ.39 whose material must have been grossly contaminated and whose structures (5) and (6) could no longer be accommodated.Shortly after this announcement BalCnovic et aL51 in Yugoslavia using partition chromatography on cellulose columns and Kuehl I,ebel and in America employing a preliminary fractionation on a resin I.R.C. 50 and purification by chromatography on Super-cel isolated crystalline muscarine chloride and verified the newly established formula. Searching for new techniques for the large-scale preparation of muscarine BalCnovic et found that choline and muscarine could be adsorbed on and successively eluted from the resin Dowex 50-X8. A similar procedure with Amberlite l.R.A.-400 was applied later by Eugster and MullerlGb to the isolation of muscarine from other fungal species. Kogl and his co-workers5* had also resumed work in this field and realis- ing the limitations of their previous method of fractionating the reineckates had by chromatography on Norite succeeded in isolating both pure muscarine chloride and acetylcholine chloride.They withdrew their previous pronouncement in favour of the formula C,H,,ClNO 2. It could now be anticipated that with identical results obtained in four independent laboratories any deductions relating to the structure of rnuscarine would at least be based on the degradations of a single chemical entity. 47 Eugster and Waser Experientia 1954,10,298. ** Eugster Helv. Chim. Acta 1956 39 1002. 49 Waser Experientia 1955 11 452. 6o Thiele and Reuther Naturwiss. 1954 41 230. 61 BalCnovic Cerar Ggspert and Galijan Arhiv Kem. 1955 27 105. 63 Kuehl Lebel and Richter J.Amet. Chem. SOC. 1955 77 6663. 6s BalCnovic and Stefanac Chem. andlnd. 1956,23 ; BaKnovic Bregant and StCfanac s4 Kogl Salemink Schouten and Jellinck Rec. Truv. chim. 1957 76 109; Cox Croat Chem. Acta 1957,29,45. Diss. Utrecht 1958. WILKINSON THE HISTORY AND CHEMISTRY OF MUSCARINE 159 The Structure of Muscarine In order to maintain reasonable chronological order and in spite of considerable overlapping of publications it is felt that because of Eugster’s prior claim to publication of the correct formula his preliminary efforts to determine the structure of muscarine should be discussed first. Eugster4* obtained a yield of 0.0003% of muscarine chloride from fly agaric but much of his work was carried out on material isolated from lrzocybe patouillardi to which previous reference has been made.That the alkaloid was a quaternary salt was shown by its pyrolysis with loss of methyl chloride to the pharmacologically inactive tertiary base normuscar- ine which could be reconverted into the quaternary salt muscarine chloride. The absence of a carbonyl group was proved by non-reactivity in diagnostic tests and confirmed from the ultraviolet and infrared spectra. The presence of a hydroxyl group was indicated by formation of 0-acetyl- muscarine ; and the remarkable stability of muscarine hydroxide eliminated an alkoxytrimethylammonium structure (see 9) such as that suggested by Rogers et al.46 These observations were amplified by Kuehl et aZ.52 The presence of a hydroxyl absorption band (5.75 mp) in the spectrum of muscarine chloride and its absence in that of acetylmuscarine suggested that only one such group was present.Although Zeisel methoxyl determinations were negative they considered that the inertness of the second oxygen atom might be the result of an ether structure. Neither Eugster nor Kuehl admittedly working with micro-amounts could detect either trimethylamine or any acidic product on Hofmann degradation. By a series of micro-scale degradations of muscarine and a number of model substances supported by evidence from infrared spectra E ~ g s t e r ~ ~ proved that the second oxygen atom was located in a tetrahydrofuran ring and that the structures (10) and (1 1) accommodated his available evidence. The nature of the side chain was based on the formation of acetic acid by oxidation and of a little iodoform with hypoiodous acid.However these H2C ,CH 0 FHMe OH ( 1 I> deductions proved to be wrong. In a later publication Eugster and W a ~ e r ~ ~ described an interesting application of catalytic oxidation which was later to prove invaluable in determining the structures of the synthetic stereo- isomers of muscarine. Oxidation of muscarine chloride in the presence of reduced platinum according to the procedure of Sneedon and Turner5’ 66 Eugster Helv. Chim. Acta 1956 39 1023. 66 Eugster and Waser Helv. Chim. Acra 1957 40 888. 67 Sneedon and Turner J. Amer. Chem. Soc. 1955,77 190. 160 QUARTERLY REVIEWS gave the 0x0-derivative muscarone whose infrared spectrum was of the cyclopentanone type veritable proof that the hydroxyl group must be located in the tetrahydrofuran ring and not in the side chain.In the meantime Kogl et aZ.54 had pursued further chemical studies and by carrying out the Hofmann degradation on 100 mg. of muscarine chloride finally settled the vexing question of the nature of the products. Although the major product was normuscarine they isolated and identified trimethylamine by paper chromatography and by X-ray powder analysis of the chloroaurate. In the infrared spectrum of muscarine iodide the characteristic bands of OH (3300 and 1380 cm.-l) quaternary ammonium (1495 cm.-l) and tetrahydrofuran (1079 cm.-l) were located but carbonyl bands (1600 and 1830 cm.-I) were absent. The culminating evidence relating to the disposition of the substituent groups was obtained by opening the tetrahydrofuran ring with hydriodic acid and red phosphorus to give n-hexyltrimethylammonium iodide.(1 2). Such a product could not arise from either of Eugster's structures (10) or (1 1) which would yield 2- and 3-aminohexane derivatives respectively but must be derived from a hydroxy-derivative of the tetrahydrofuran (1 3). Further degradative work was found to be unnecessary for from the possible variants derivable from a hydroxy-derivative of (13) the final selection was made by X-ray crystallographic analysis of muscarine i ~ d i d e . ~ ~ ~ * Differential Fourier synthesis clearly depicted the structure to be that of a quaternary ammo- nium salt of S-aminomethyltetrahydro-3-hydroxy-2-methylfuran and moreover one in which the hydroxyl group was in the trans-position with respect to both the methyl and the CH2.NMe,+ side chain as indicated in formula (14).(14) Muscarine (1 5) Epimuscarine (25 3R 55) (25 35 55) (1 6) Allomuscarine (25 3R 5R) (1 7) Epiallomuscarine (25 35 5 4 58 Jellinck Acta Cryst. 1957 10 277; Diss. Utrecht 1957. WILKINSON THE HISTORY AND CHEMISTRY OF MUSCARINE 161 Four racemic stereoisomeric forms of the basic muscarine molecule are possible each of which can exist in two optically active forms giving a total of eight possible variants. The four racemates have been named and assigned the structures (14)-(17) by Corrodi Hardegger and K0gl;59 the stereochemical nomenclature in parentheses is based on that suggested by Cahn Ingold and Prelog.60 The two enantiomorphs of muscarine and the racemates of epi- allo- and epiallo-muscarine have now been synthesised. Syntheses of Muscarine ( A ) Stereo chern ically Non-spec i j c Syn theses .-( 5) - Muscar ine . After publication of its structure a number of syntheses of (&)-muscarine followed in rapid succession. From the furan derivative (18) obtained by condensing /3-keto-esters with glucose or mannose Eugster et aL61 synthesised by the series of reactions depicted (Synthesis l) the important intermediate 5-dimethylaminomethyl-2,3-dihydro-2-methyl-3-oxo~uran (19) which proved so useful in the preparation of all the stereochemical variants. Synthesis 1. EtO,C EtO,C I M e p C H O Me V C H i NMe2 Mep[&l(OH)] .CH,.OH - Et02C (18) Ph * CH,-CO*NH I ___c CH,,NMe2 - 8 Me@H;NMe2 (20 ' (20) CH,.NMe2 Reagents 1 Pb,O,. 2 Leuckart reaction. 3 N,H,. 4 HNO,. 5 Ph.CH,.OH. 6 ~ N - H C I at 100". 7 Pd-H,. 8 LiAIH,. 9 KBH,. Catalytic hydrogenation of the ketone (19) gave (&)-normuscarone (20) which was further reduced by lithium aluminium hydride to (&)-normus- carine (21).Alternatively one-stage reduction of the ketone (19) with potassium borohydride gave a mixture of racemates of the normuscarines (see later section on isomers of muscarine) from which (5)-normuscarine (21) was isolated by chromatography on deactivated alumina. The infra- red spectrum of the quaternary (-&muscarine iodide obtained from the 5B Corrodi Hardegger and Kogl Helv. Chim. Acta 1957 40 2454. 6o Cahn lngold and Prelog Experientia 1956 12 81. 61 Eugster Helv. Chim. Acta 1957 40 2462; Eugster HMiger Denss and Girod ibid. 1957,40,205. 162 QUARTERLY REVIEWS base (21) was almost identical with that of natural (+)-muscarine iodide and its biological activity one-half of that of the natural isomer.Synthesis 2. + c< /c\H2 + Me-CO-HF /c\H2 $H.CH,Cl Me*CO*HY /c\H2 $HCH2-NMe Me.CO.7 FH.CH2.NMe3 oc-0 CI- - oc-0 c t - L oc-0 (221 c/,C%\ CIL,CH,\ Me.$J+CH yH.CHi&Me A MeeCOCH yH.CHihMe3 OH CL' OH CI- 7 HO\ H/C-CH I/%\ 6 \ + AcO Me.FH-U-1 ~ H w i k M e - MemCH CH.CH;NMe cl' Reagents 1 NMe,. 2 SO,CI,. 3 20% HCI. 4 NaBH,. 5 AgOAc. 6 20% Kogl Cox and Saleminks2 described a straight-forward synthesis from 6-chloro-a-acetyl-y-valerolactone (22) which followed the course indicated in Synthesis 2. The final ring closure carried out according to the procedure described by R e ~ p e ~ ~ gave a mixture of isomers from which about 30% of (-J-)-muscarine chloride was isolated by chromatography on Norite. The product which did not crystallise had an activity of about one-third of that of the natural alkaloid.\O' OH OH Cl- H,SO, then conc. H,SO,. Synthesis 3. Reagents 1 H,SO, then CH,N2. 2 NaBH,. 3 NHMe,. 4 LiAIH, then Mel. A second reputed synthesis by Kogl et aLs4 (see Synthesis 3) was shown later59 to give little or no muscarine but essentially the racemate of allo- muscarine. A mixture of muscarine and allomuscarine was obtained by a method devised by Natsumoto and Maekawa.'j5 62 Kogl Cox and Salemink Experientia 1957 13 137; Annalen 1957 608 81. 63 Reppe Annalen 1955 596 90 118. 64 Corrodi Hardegger Kogl and Zeller Experientia 1957 13 138. 65 Natsumoto and Maekawa Angew. Chern. 1958,70 507. WILKINSON THE HISTORY AND CHEMISTRY OF MUSCARINE 163 (B) Resolution of (3)-Muscarine.-Two rather conflicting procedures By partition chromatography on powdered cellulose Kogl et have been described for the resolution of (&)-muscarine.improved the separation of muscarine from the mixture of racemates obtained in their synthesis 2 and effected its resolution by means of (-)-di-p-toluoyltartaric acid. The salt of (+)-muscarine separated first and the regenerated (+)-muscarhe chloride was found to be identical with natural muscarine chloride. found that the less soluble salt of (-)-di-p-toluoyltartaric acid was that of the unnatural isomer (-)- muscarine. (+)-Muscarhe chloride m.p. 178-179" [aID + 8.1" (in EtOH) was obtained from the less soluble salt of (+bdi-p-toluoyltartaric acid. The (-)-muscarine chloride m.p. 179-180" [a] - 8.4" (in EtO€€) had only 5% of the biological activity (blood pressure in cats) of its enantio- morph.(C) Specijic Stereochemical Syntheses of (+)- and (-)-A4uscarine.- An elegant stereospecific synthesis of (+bmuscarine chloride was devised by Hardegger and Lohse68 (see Synthesis 4). This started from L-glucos- amine (23) and proceeded via L-glucosaminic acid (24) L-chitaric acid (25) In contrast however Eugster et Ts = p-C,H,Me.SO Reagents 1 CH,N, then NHMe,. 2 p-CBH,Me-SO,CI. 3 LiAIH,. and its dimethylam.de (26). Reduction of the tritosyl derivative (27) with lithium aluminium hydride produced a 'low yield of (+)-normuscarine (28) which was quaternised and the (+)-muscarine isolated as the tetra- phenylborate which was converted into the chloride with caesium chloride. Since the absolute configuration and relation of L-glucosamine to L-glyceraldehyde has been established natural muscarine must be 66 Cox Hardegger Kogl Leitchti Lohse and Salemink Helv.Chim. Acta 1958 67 Eugster Hafliger Denss and Girod Helv. Chim. Acta 1958,41,886. 41 229. Hardegger and Lohse Helv. Chim. Acta 1957 40 2383. 1 64 QUARTERLY REVIEWS 2,5-anhydro-l,3,6-trideoxy-~-ribo-hexityltrimethylammonium chloride in which C(5) (marked with an asterisk) is the Rasanov carbon with the L,- configuration and it has been described as L-( +)-muscarhe. CHO CHO (3 3) Two comparable specific syntheses of (-)-muscarine chloride were also carried out. The first66 proceeded from D-glucosamine (29) via D-chitaric acid (30) and in the 2-deoxy-~-ribose (31) was converted into 3-deoxy-~-chitaric acid (32). By analogous series of reactions to those used in the preparation of the (+)-isomer both these acids gave (--)- muscarine chloride (33) which was devoid of activity on the isolated frog heart.Synthesis of Epi- Allo- and Epiallo-muscarine etc. As mentioned previously the ketone (34) proved to be a valuable intermediate. The reduction of this ketone was investigated at considerable length and for various conditions (&)-muscarine and its three racemic stereoisomers being Synthesis 5. H2*NMe2 Nomuscarone Reagents 1 H2-Raney Ni. 2 H2-Pd. 3 KBH,. 4 Pt-0,. 5 LiAIH4. 6 AI(OPr$,. . ." A Normuscarine. B. Epinormuscarine. C Allonormuscarine. D Epiallonormuscarine. 69 Hardegger Furter and Kiss Helv. Chim. Ada 1958 41 2401. 'O Eugster Hafliger Denss and Girod Helv. Chim. Ada 1958 41 205 583 705, WILKINSON THE HISTORY AND CHEMISTRY OF MUSCARINE 165 The various courses of reduction as shown in Synthesis 5 invariably led to a series of mixed racemates of the normuscarines which were separated by chromatography on alumina and subsequently quaternised.The nature of the products was conveniently established as follows. The four isomers fall into two natural groups according as to which keto-derivative (normuscarone) they give on catalytic oxidation thus Nor- and epinor-muscarine -+ Normuscarone Allonor- and epiallonor-muscarine 4 Allonormuscarone In each of the two pairs therefore the individuals differ only with respect to the configuration of the hydroxyl group. That the hydroxyl group of muscarine was trans to both the Me and the CH,*NMe,+ group was known from the X-ray crystallographic work and from the synthesis from L-glucosamine and further proof was derived from the infrared spectra.Epimuscarine iodide gave a band at 3165 cm.-l which was absent from the spectrum of muscarine iodide and is characteristic of hydrogen bonding. This implied that the hydroxyl and the NMe$ group were cis with respect to each other in epimuscarine and trans in muscarine. Similarly the infrared spectra of allo- and epiallo-muscarine showed this discriminating difference associated with hydrogen bonding in the allomuscarine molecule. The pharmacology of the muscarines and muscarones has been studied in great detail. Before reviewing their stereospecificity brief mention must be made of the syntheses by Eugster and his colleagues of two series of derivatives the demethyl- and the dehydro-muscarines which are of considerable importance in such a discussion because in each case the number of asymmetric centres has been reduced from three to two.C0,Et CH.CO,Et I + It CH ,.OH CH * CO Et (361 J2 b 2 N M e 2 + @ ; y e 2 ~ 3 0 - J 0 CH,.NMe (3 8) trans.b.p. 120-140" (37' cis b.p. 90- 100" (40) ?-JCH2AMC3 I - Reagents 1 H,SO, then CH,N,. 2 NHMe, then LiAIH4. 3 KBH,. 4 Mel. Demethylnormuscarone (37) and its quaternary salt demethylmuscarone iodide (40) were prepared'l from the dimethylamide derived from methyl tetrahydro-5-oxofuran-2-carboxylate (36) by first protecting the keto- group and then reducing the amide group with lithium aluminium hydride. 71 Zwicky Waser and Eugster Helv. Chim. Acta 1959 42 1177. 166 QUARTERLY REVIEWS Reduction of demethylnormuscarone with potassium borohydride resulted in a mixture of cis- (38) and trans-demethylnormuscarine (39) which were separated by fractional distillation and chromatography on alumina and were finally quaternised.By further fra~tionation~~ of the mixture derived from the reduction of compound (34) with sodium borohydride (see Synthesis 5 ) were isolated in addition to the normuscarine isomers cis- (4 1) and trans-4,Sdehydro- muscarine (42) as products of partial reduction. Both isomers were dehydrated at 200" to 2-dimethylaminomethyl-5-methylfuran (43) the quaternary salt of which had previously been synthesised and its muscarinic activity examined by Ing Kordik and Tud~r-Williams.~~ OH OH Impure epinormuscarine Normuscarine Reagent 1 H,-Raney Ni. Pharmacological Stereospecificity in the Muscarine and Muscarone Series One remarkable feature has transpired in the course of the very com- prehensive pharmacological studies which have now been carried out with these alkaloids.Although (+)-muscarine does not possess the highest muscarinic activity in contrast both to its stereoisomers and to such derivatives as have so far been examined its action is highly specific and restricted to activity at postganglionic parasympathetic effector sites. Thus W a ~ e r ~ ~ demonstrated that high doses have only peripheral parasympathetic action in the cat and even with toxic doses up to 100 pg./kg. in the atropinised animal no blocking of sympathetic and para- sympathetic ganglia or nerve-muscle transmission could be detected. By inhibiting the effect of cholinesterase he found that the activity of muscarine was enhanced ten-fold and suggested that it might be trans- formed in the body into an active acetyl ester.In vivo in the cat and dog muscarine caused a marked drop in blood pressure. Working with muscarine chloride isolated by the author of this Review F r a ~ e r ~ ~ showed that in the absence of an anticholinesterase the prepara- 72 Denss Girod Hafliger and Eugster Helv. Chim. Acta 1959,42 1191. 73 Ing Kordik and Tudor-Williams Brit. J. Pharmacol. 1952,7 103. 74 Waser Experientia 1955 11 452; Konzett and Waser Helv. Physiol. Pharmacol. 76 Fraser Brit. J. Pharmacol. 1957 12 47. Acta 1956,14,202. WILKINSON THE HISTORY AND CHEMISTRY OF MUSCARME 167 tion was 4-5 times more active than acetylcholine on isolated rabbit auricles but parallel dose-response curves were not obtained.With the same specimen Ambache Perry and Robertson76 described the absence of a pressor reponse in atropinised anzsthetised cats a weak ganglion- stimulating effect readily blocked by atropine and in large doses a stimula- tion of the frog's rectus abdominis. Gyermek and Herr,77 however suggest that there is a lack of ganglionic stimulation in doses 400-1000 times larger than those which stimulate post-ganglionic parasympathetic effector sites. Fra~er'~ found that muscarine was stable to pepsin and that no response could be detected on feeding it orally to monkeys in doses much greater than would be expected to cause poisoning by ingestion of A . muscaria in man. The following brief summary attempts to fix the highlights of the points which have emerged from the spectrum of physiological action but the reader is referred to the relevant publications for the extensive and specialised technical details.Tables 1 and 2 correlate the potencies of muscarine and its isomers in tests on selected pharmacological prepara- tions. (a) The enantiomorphs of (-J-)-muscarine differ greatly in potency which i s almost exclusively associated with the (+)-isomer. (b) Waser7* and Gyermek and Unna79a have clearly illustrated that there is a definite specificity of the action of the muscarines on post-ganglionic effector sites. The racemic forms of epi- allo- and epiallo-muscarine possess only a fraction of the potency of (+)-muscarine. Waser states they have less than one-hundredth of the effect of (+)-muscarine on the blood pressure of cats and on isolated frog hearts.Gyermek and Unna found that all muscarine isomers were devoid of significant action on skeletal muscle. These results (see Table lb) lead to the conclusion that when the hydroxyl group is cis to either or both of the methyl and the CH,-NMe,+ group as in the epi- allo- and epiallo-isomers there is an enormous decrease in mus- carinic activity. (c) In very marked contrast there is a lack of stereospecificity in the series of related ketones (+)- and (-)-muscarone and (&)-allomuscarone which differ only slightly in their action on smooth muscle (Table lc). Although the muscarones have high muscarinic activity they also exert effects at other synapses as well. Thus they exhibit strong nicotinic activity on the frog rectus and block ganglionic and neuromuscular transmissions in the cat.Muscarones and in particular (-)-muscarone were significantly more potent than (+)-muscarhe and acetylcholine in stimulating post- ganglionic parasympathetic sites. In their most recent assessment Gyermek 76 Ambache Perry and Robertson Brit. J. Pharmacol. 1956,11,442. 77 Gyermek and Herr Fed. Proc. 1959 18 399. 7a Waser Experientia 1958 14 356. 7B Gyermek and Unna (a) Proc. SOC. Exp. Biol. Med. 1958,98,882; (b) J. Pharmacol. 1960,12.8,30,37. TABLE 1. (a) ( f)-Muscarine iodide (+)-Muscarine iodide ( +)-Normuscarhe hydrochloride (b) (&I-Epimuscarine iodide ( f )-Allomuscarine iodide ( f )-Epiallomuscarine iodide (c) (&)-Muscarone iodide (-)-Muscarone iodide ( &)-Allomuscarone iodide (d) frans-4,5-Dehydromuscarine iodide cis-4,5-Dehydromuscarine iodide ( f)-4,5-Dehydromuscarone iodide (e) trans-Demethylmuscarine iodide cis-Demethylmuscarine iodide Acetylcholine chloride y/kg.in vivo y/ml. Perfusion of bath-volume Blood pressure in cats Muscarinic (M) contraction. Nicotinic (N) (minimal dose) Frog heart Frog rectus submaximal 0.01 0-032 > 500 0.004 0.01 8 > 500 5.0 5.0 - 3.0 1-7 1.0 0.001 5 0.001 0.003 2.4 15 12 0.0 1 0-0075 0.02 > lo00 >lo00 > 500 2.5 2.5 1.5 0.01 0.2-0.3 50 0.02 0.4 > 100 0.001 5 0.0 1 0.9 1.0 7-0 30 400 0-01 0-002 5 Quotient N/M > 20,000 > 30,00O - > 5 > 65 > 40 250 330 75 250 > 250 90 - - 2500 (a) Concentration of sample (ylkg.) producing fall of blood pressure equal to that given by 0.01 y/kg. of acetylcholine when injected (b) Concentration of sample (ylml.) perfusing through an isolated frog’s heart which produces decreases in amplitude and rate of beat (c) Concentration of sample (y/ml.of bath-volume) giving one-quarter of the maximum contraction of the isolated frog’s rectus abdomink intravenously into an anathetised cat. of both the auricle and ventricle equivalent to those given with 0.002 y/ml. of acetylcholine. muscle. WILKINSON THE HISTORY AND CHEMISTRY OF MUSCARINB 169 and Unna79b found (-)-muscarone was 4-6 times more active than (+)-muscarine. The muscarones therefore resemble acetylcholine in activity and are more closely related to it structurally than is muscarine with its hydroxy-group as illustrated. Muscarone Acetylcholine Although powerful nicotinic activity is introduced by formation of the muscarones that this is a result of the substitution of an 0x0- for a hydroxy-group cannot be concluded because of the simultaneous reduction in the number of asymmetric centres.Steric hindrance of the hydroxyl or keto-group and of the ring-oxygen atom will play important parts in determining the ease of contact with the cholinergic receptors. Moreover since the tertiary base normuscarine is inactive the quaternary nitrogen must be considered to be a third point of contact. According to recent studies by Gyermek and Unna79b all three constituents together with the ring-oxygen atom should be considered as essential pharmacophore groups . (4 The introduction of a double bond with corresponding loss of asymmetry at position 5 as in the 4,5-dehydromuscarines had com- paratively less effect in changing the pharmacological picture than in the muscarine series (Table Id).Even with complete aromatisation of the ring as in the compounds (44) and (45) high muscarinic activity had been ob~erved.~~,*~ (e) Demethylation with resulting loss of asymmetry at position 2 produced a marked decrease in potency (Table le). (f> Witkop Durant and Frjessal have examined the effect of muscarone derivatives on the inhibition of cholinesterase activity. 0-Acetylmuscarine was found to be the most potent inhibitor and as with the other diverse pharmacological preparations the enzyme responded in vitro more markedly to stereochemical differences in the muscarine than in the muscarone series. From models of the isomers of muscarine they concluded that there appeared to be a closer approach in the natural than in the allo-series to coplanarity of the ring with both the quaternary nitrogen atom and the hydroxyl group.This could lead to additional binding strength with the receptor if the ring augmented the effect of the two polar groups. *O Fellows and Livingstone J. Phurmucol. 1940 68 231. 81 Witkop Durant and Friess Experientiu 1959 8 300. 170 QUARTERLY REVIEWS (g) Waser82 recently tested a series of analogues of muscarine in which the tetrahydrofuran ring was replaced by tetrahydrothiophen. There was marked reduction in the diverse biological activities (Table 2). This he considered might be attributed to the inability of sulphur to form a hydrogen bond with the cholinergic receptor and also to the fact that the tetrahydrothiophen ring is known to be larger than the tetrahydrofuran ring83 and could therefore affect the relative position and function of the methyl group.TABLE 2. y/ml. Perfusion of bath-vol. Frog rectus Substance Frog heart submaximal contraction (MI (N) (A) Posn. of OH 0.02 > 500 (33 (B) trans 500 >2o00 cis lo00 > lo00 (C) trans cis (D) trans cis (E) trans cis 30 400 100 200 50 lo00 > lo00 > lo00 (F) Mainly trans >4000 > lo00 In the same publication increasing the length of the alkyl group to n- The positional isomers (46) and (47) have been shown to have little if and iso-propyl was shown to decrease the activities greatly (Table 2). any muscarine W a ~ e r ~ ~ has discussed the use of [14C]muscarone in the elucidation of the nature of the cholinergic receptor site. 88 Waser Experientia 1960 16 347. 83 Marsh Acta Cryst.1955 8 91. 84 E. Gryszkiewicz-Trochimowski 0. Gryszkiewicz-Trochimowski and Zevy Bull. 86 Waser J. Pharm. Phurmacol. 1960 12 577. Suc. chim. Frunce 1958 603 ; Fraser personal communication. WILKINSON THE HISTORY AND CHEMISTRY OF MUSCARINE 171 Other substances isolated from A. rnuscaria Wieland Motzel and Merzss reported the presence of bufotenine in A . rnuscaria; Kogl Salemink and Schullera7 have recently isolated a new alkaloid muscaridine closely related to muscarine which they found to be 4,5-dihydroxyhexyltrimethylammonium chloride (48). Its constitution was established by oxidation with periodate to acetaldehyde and Rjth permanganate to the trimethylammonium derivative of /3-aminobutyric acid 86 Wieland Motzel and Merz Annalen 1953 581 10. *' Kogl Salemink and Schuller Rec. Trav. chim. 1960 79 278.

ISSN:0009-2681    

DOI:10.1039/QR9611500153

出版商:RSC    

年代:1961

数据来源: RSC

摘要:

HALIDES OF THE PHOSPHORUS GROUP ELEMENTS (P As Sb Bi) By D. S. PAYNE (THE UNIVERSITY GLASGOW) IT is a commonplace to observe that the past decade has seen a renaissance of inorganic chemistry. During this period certain topics notably the chemistry of uranium and the transuranic elements the hydrides of boron and their derivatives and non-aqueous solvent systems have been in- vestigated with great intensity to give results of considerable interest. However during this period a steady stream of work has appeared dealing with the chemistry of many if not all of the superficially well-known elements. This new knowledge whilst perhaps less spectacular is important in that it leads to a more complete picture of the potentialities and a better understanding of the general underlying principles of the “well- known” elements.Amongst the groups of elements whose chemistry is accepted as being well known one of the best examples is the phosphorus group. There is a particular interest in this group of elements both in its juxtaposition to the sulphur and the silicon group and because of the occurrence of at least two distinct series of compounds (MI11 and Mv). Recent work has added considerably to our knowledge of these elements and it is the purpose of this Review to draw attention to this consideration being limited to the halides only. Nitrogen whilst having a certain simi- larity to the phosphorus group elements arising from the ability to form a simple octet leading to tervalency has substantial differences because of its inability to make use of orbitals beyond the 2s and 2p levels.The extent to which the phosphorus group utilises the orbitals available for bonding varies markedly with both the element and the nature of the atom or group involved in bond formation; it is most marked in the case of the 3d orbitals of phosphorus and least obvious in the case of the 6d orbitals of bismuth. Before discussing the detailed chemistry of the halides it is appropriate to mention certain relevant features of the elements concerned. The following electronegativity data including values for the halogens are taken from Pritchard and Skinner’s review? N = 3.0 P = 2-1 As = 2.0 Sb = 1.9 Bi = 1.8 F = 3.9 C1 = 3.0 Br = 2.8 I = 2.5. However the values for the phosphorus group are probably of varying reliability. Although the values P = 2.19 As = 1.97 and Sb = 2.28 have recently been put forward,2 their validity has already been ~hallenged.~ The 1st and 2nd ionisation potentials decrease from phosphorus to bismuth.Pritchard and Skinner Chern. Rev. 1955 55 745. Fineman and Daignault J. Inorg. Nuclear Chem. 1959 10 205. Drago J. Inorg. Nuclear Chem. 1960 15 237. 173 174 QUARTERLY REVIEWS This trend is however broken in the 3rd 4th and 5th ionisation potentials where a very slight increase occurs at bismuth. This coupled with the greater size of the Bi3+ ion than of the other ions in the series is a factor in the diminished stability of the higher valency states of bismuth compared to the other members of the group. Values for bismuth (8 16.6 and 25.4 ev) are comparable with those for gallium (8.1 15.9 and 30.6 ev); thus it might be expected that solvation effects being comparable chemical similarities between bismuth and gallium would exist.For example obvious cationic tendencies might be observed particularly in aqueous solution. Van Wazer4 has drawn attention to the principles of bonding in phos- phorus compounds and these may quite appropriately be extended to the whole group. Setting aside the very small number of less-common configu- rations four main types of situation are found. (a) Three-connected atoms. In these situations the atom is at the apex of a triangular pyramid. This is a common arrangement and many examples have been investigated in detail. Evidence based on bond-length shortening and on enthalpy data4 suggests that in the case of phosphorus there are three pure a-bonds with little or no r-character and this might well apply to the remainder of the group.The bond angle drops from 104" & 4" in phosphorus trifluoride to 100" 2" in the tribromide; for the correspond- ing arsenic trihalides the values are 102" & 2" and 100" & 2" respectively with a value of 98.5" for the tri-iodide. However the angles in antimony trichloride and tribromide and in bismuth trichloride and tribromide are very similar; 96" &- 4" and 96" -j= 2" 100" -)= 6" and 100" & 4" respect- ively.6 The values of these angles fall between those expected for bonds involving pure p orbitals (90") and that for sp3-hybrid orbitals (109" 28'). The shapes of these molecules can be interpreted satisfactorily in terms of a greater repulsion between non-bond pairs and bond pairs than between bond pairs and bond pairs.' (b) Five-connected atoms.The atom is at the centre of five atoms the bonds usually being directed in a trigonal bipyramidal fashion. Molecules containing this arrangement are not common. Evidence suggests that for phosphorus these bonds are almost wholly of a-type.* The general dis- position of the five bonded pairs is that to be expected in terms of the simple repulsive forces between bond and non-bond pairs. Alternatively the five bonds can be represented as arising from ~p~d-hybridised orbitals. In a few cases five-connected atoms can be encountered in anions of the type MX,2-. The structure here is that of a distorted octahedron (tetra- gonal pyramid) in which one of the octahedral positions is taken by a Van Wazer "Phosphorus and its compounds," Interscience New York 1958 Vol.I. Chem. SOC. Special Publ. No. 11. Gillespie and Nyholm Quart. Rev. 1957 11 369. ti Van Wazer J. Amer. Chem. Soc. 1956,78 5709. PAYNE HALIDES OF THE PHOSPHORUS GROUP ELEMENTS 175 non-bonded pair such as occurs in (NHd,SbCl,8 and K2SbF6.9 Similar related structures are found in KSbF4 lo and NaSbFp.ll (c) Six-connected atoms. For a co-ordination number of six the arrangement is octahedral as follows directly from consideration of repulsions of six equivalent electron pairs or from the form of sp3d2- hybridisation. This form as a complex anion is encountered more often than the five-co-ordinated situation. For phosphorus the ligands in such a system are largely ~ b o n d e d . ~ (d) Four-connected atoms. A feature of the halide chemistry of this group is the frequent occurrence of four co-ordinated atoms usually as the cation MX,+.Van Wazer4 has shown that in the case of phosphorus this co-ordination involves four a-type bonds with an average of one n-bond per phosphorus atom the bond order being 1.25. The v-bond appears necessary to reduce the positive charge which would otherwise accumulate on the phosphorus atom as a result of electronegativity differences. The need for v-bonding is eliminated in three co-ordinate compounds with strongly electronegative ligands the lone pair being responsible for balancing charge differences. In fike- and six-co-ordinate phosphorus the lowest-energy d-orbital(s) are already in use and are not available for n-bonding. The general picture based on electronegativity differences of the bonds formed in the halides is one of appreciable ionic character particularly for the chlorides and fluorides.However this ionic character is somewhat offset by T-bonding. Since d-orbitals can only be used in bonding where the ligand possesses a high relative electronegativity,12 five- and six-co- ordinate compounds might be restricted to chlorides and fluorides. The effect of ligand size might also be to favour the chlorides and fluorides rather than the bromides and iodides. An additional factor in determining stability is the increasing tendency of the halogens on passing from fluorine to ipdine to form p,-d bonds. As in the neighbouring groups the mixed halides of the phosphorus group are apparently less stable than the simple halides although there are certain exceptions e.g. PF3C12. The presence of fluoride in the mixed trihalides appears to increase the stability to re-organisation processes.13 The stereochemical consequences of the various co-ordination polyhedra in the mixed haiides have still to be realised for example the compound PF,Cl could occur in three stereochemical forms (l) (2) and (3) [(I) appears to be the form present in the gaseous phase at roomtemperature] and the anion PF4C12- in two forms (4) and (5).So far it has not been Edstrand Inge and Ingri Acta Gem. Scand. 1955 9 122. * Bystrom and Wilhelmi Arkiv Kemi 1951,3 461. lo Bystrom Backlund and Wilhelmi Arkiv Kemi 1952 4 175. l1 Bystrom Backlund and Wilhelmi Arkiv Kemi 1953 6 77. l2 Craig Maccoll Nyholm Orgel and Sutton J. 1954 332. lS Delwaulle Cras Bridoux and Migeon XVIIth I.U.P.A.C. Conference Munieh 1959.176 QUARTERLY REVIEWS possible to investigate these possibilities fully in the halides although they have been recognised in the trifluoromethyl-derivatives where (CF,),PCl is in the trans-form whereas (CF&PC12 is a mixture of various stereochemi- cal forms.l* A common feature of halides generally is the existence of halogen bridges arising from donor bonds from halogen to acceptor orbitals on adjacent atoms; these bridges are often of the form shown at (A) incor- porating a four-membered ring. Examples are more common amongst chlorides and bromides than amongst fluorides and iodides possibly owing to steric and electronegativity factors. So far the possibility of such bridges in the halides of the phosphorus group elements appears limited to P2C1, l5 and certain complex anions.The importance of the tetrahedral configuration is shown in the existence of many halides particularly the pentahalides in ionic four-co-ordinate states in preference to molecular states with higher or lower co-ordination numbers. There are however a few pentahalides which occur as such and give simple molecular lattices e.g. SbCl,. Two forms of each of the halides might be expected corresponding to molecular and ionic lattices but in fact evidence of this is limited to PCl,F SbCl,F and possibly to PCI,. The ionic forms of the halides might be expected to be particularly stable in suitable solvent environments (high dielectric constant). The anionic form is often but apparently not invariably six-co-ordinate. The trihalides Certain of the trihalides have a long and interesting history.BoylelG was the first to prepare bismuth trichloride (1 664) whilst antimony trichloride was first correctly identified by Glauber in 1648. The simple trihalides are obtained by the direct interaction of the elements in stoicheio- metric proportions. Only in the case of bismuth and fluorine is difficulty experienced owing to the formation of a protective halide film. The reactions are all exothermic to varying degrees. Various metathetical reactions are available for preparation in some fashion depending for their success on the volatility or low solubility of the products. Metathesis is particularly l4 Harris Ph.D. Dissertation Cambridge 1958. l6 Kennedy and Payne J. 1960 4126; Kennedy Payne Reed and Snedden Proc. l6 Boyle Chem. Soc.,'!959 133. Experiments and Consideration concerning Colour," London 1664.PAYNE HALIDES OF THE PHOSPHORUS GROUP ELEMENTS 177 applicable to the preparation of the fluorides e.g. PCl + AsF -+ PF3 + AsCl,. For the preparation of iodides it is convenient to utilise the reaction between the corresponding chloride and for example potassium iodide in a suitable solvent such as acetone. Various halides in particular the fluorides can be prepared from the oxide or sulphide by the action of the hydrogen halide. In certain cases it is possible to achieve reaction between the oxide or more favourably the sulphide and elemental halogen. Phosphorus appears to be excluded from these reactions. The halides of bismuth and antimony can be obtained from aqueous solutions in the presence of excess of acid but this is impossible for arsenic or phosphorus halides (other than arsenic tri-iodide) owing to their ready hydrolysis.The number of possible mixed trihalides is large but so far only a few have been prepared. The methods used for their preparation have involved either the application of re-organisation reactions in which mixtures of different trihalides are heated together until an equilibrium mixture has been obtained or the partial exchange of halogens with another halide. Those mixed halides which have been isolated and whose individual properties have been recognised are PF,C1 PFC12,17 PF,Br PFBr,,l* and SbBrI,.lS Others have so far only been identified in the course of ebullioscopic,20 Raman spectra,21*22 and nuclear magnetic resonance studies ;23 for example a mixture of phosphorus trichloride and tribromide produces in a few minutes at room temperature both PC1,Br and PClBr in equilibrium with the original halides,23 whilst PC1,F and PBr,F give PFClBr in the equilibrium mixture.Mixtures of PC1,F with PBr and PBr,F with PCl yield mixtures containing a large number of molecular species; the isolation of the individual species has not been attempted. In no case is equilbrium established instantaneously the reaction being considerably slower when fluoride is present in one of the molecules inv01ved.~~~~~ The reaction rates are appreciably faster than the re-organisa- tion reactions of four-co-ordinated phosphorus. To date attempts to prepare arsenic fluorochlorides by partial fluorination have not been succe~sfu1.~~ The trihalides exhibit many features of chemical interest.Reference has already been made to the re-organisation reactions of the phosphorus series and it appears likely that this type of reaction is general to the group. The composition of any equilibrium mixture will not however necessarily include appreciable amounts of mixed halides unless there is some particular feature favouring stability. l7 Booth and Bozarth J. Amer. Chem. SOC. 1933 55 3890. Booth and Frary J. Amer. Chem. SOC. 1939,61,2934. l8 Clark J. 1930 2737. 2o Raeder Z . anorg. Chem. 1933 210 145. Delwaulle and Francois J. Chim. phys. 1949 46 80. ** Delwaulle Compt. rend. 1947 224 389. 2s Fluck Van Wazer and Groenweghe J. Amer. Chem. SOC. 1959,81 6363. 24 Wilkins J. 1951 2726. 178 QUARTERLY REVIEWS The trihalides in the liquid phase are generally poor electrical conduc- tors.Arsenic trifluoride has a specific conductance of 2-4 x 10-50hm-l cm.-l at 25” which is of the same order as values for bromine trifluoride iodine pentafluoride and hydrogen As with these other fluor- ides the presence of potassium fluoride greatly increases the conductance and the compound KF,AsF (K+AsF,-) can be isolated. Likewise the addition of antimony pentafluoride results in an increase in conductance and the compound SbF,,AsF is isolated. The self-ionisation of AsF into AsF,+ and AsF,- is therefore postulated.26 Phosphorus trifluoride with a specific conductance of loe9 ohm-1 cm.-l at -1 13”c shows no tendency to form a complex with potassium fluoride.27 Previous reports on the formation of PF4- by this reaction thus appear incorrect.28 In contrast fused antimony trifluoride has a relatively high conductance comparable with that of fused zinc Arsenic trichloride with a dielectric constant of 12.8 at 20°c is a poor conductor but a good solvent for a wide range of halides a feature made use of in the preparation of halide complexes.29 Similar solvent properties are found for antimony tri~hloride~~ and antimony tribr~mide.~~ A common feature is the ability to promote ionisation of dissolved compounds by halide-ion transfer leading to the formation of complex ions.The high transport number (0.88-0-97) for chloride ion in a solution of tetramethylammonium chloride in arsenic trichloride arises from the ease of the chloride exchange AsCl + AsC1,- -+AsCl,- + AsCI, by which electrolytic conductance occurs.32 Arsenic trichloride is a strongly polar molecule and numerous examples of its addition compounds are known.Most are best regarded as solvates involving truly ionic species only in favourable situations. The system tetramethylammonium chloride-arsenic trichloride contains the com- pounds (CH,),NCl,AsCl and (CH,),NC1,3AsC13 whereas in the cor- responding tetraethyl system (C2H5),NC1,2AsC1 and 3(C,H,)4NCl,5AsCl are recogni~ed.~~ This tendency towards solvate formation is also shown by the existence of the compounds 2PC1,,5AsCI3 and 2PCl5,4SbC1, which have been shown by cryoscopic and conductometric methods to contain the arsenic and antimony trihalide molecules as solvates only.34 Although there is a considerable number of addition compounds of trichlorides and tribromides of arsenic and antimony the simple MX,- unit has been recognised in only a few.One example apparently well 25 Woolf and Greenwood J. 1950,2200. m Gutmann Quart. Rev. 1956 10 451; Angew Chem. 1959 71 57. 27 Woolf J. 1955 279. 28 Lange and Stein unpublished work reported by Simon “Fluorine Chemistry,” 2s Lindqvist Acta Chem. Scand. 1955 9 73. so Porter and Baughan J. 1958 744. 31 Jander and Weiss Z . Elektrochem. 1957 61 1275; ibid. 1958 62 350. a2 Gutmann Svensk Kem. Tidskr. 1956 68 1 . 33 Agerman Anderson Lindqvist and Zackrisson Acta Chern. Scand. 1958 12 84 Kolditz 2. anorg. Chem. 1957 289 118. Vol. I. Academic Press Inc. New York 1950 p. 139. 477. PAYNE HALIDES OF THE PHOSPHORUS GROUP ELEMENTS 179 substantiated is [Me,N]+ [AsCI,] -.35 Potentiometric studies in arsenic trichloride should help to determine the true nature of the trihalide complexes.36 Although the existence of salts of the form MIMII*X (along with more complex forms) has long been recognised in the chemistry of arsenic antimony and bismuth and the examples are very numerous details of the structure are known in only a few instances.Despite the obvious interest in phosphorus analogues of these complexes no convincing evidence has so far been presented for their exi~tence.~’~~~ Investigation of the tetrafluoroarsenites by nuclear magnetic resonance shows that these salts undergo rapid exchange with the arsenic trifluoride used as solvent.37 The numerous complexes of tervalent antimony with fluorine formed from antimony trioxide potassium carbonate and hydrofluoric acid have been investigated in some detail. In K,SbF a co-ordination number of five is exhibited and the structure is in keeping with an octahedral arrangement with a non-bonding pair occupying the sixth ~ o r n e r ~ leading to a square- pyramidal structure in which the antimony atom is in fact displaced outside the basal plane as a result of bond-pair-non-bond-pair repulsion^.^^ KSb,F, however contains four-co-ordinate antimony and can be seen as two trigonal bipyramids linked through a bridging fluorine with non- bonding pairs at the pyramidal apices.39 Four-co-ordination might be (6) expected in the compounds RSbF (R = K Rb Cs NH, or TI) however none has so far been shown to involve an [SbF4]- ion.KSbF4 for example contains the [Sb4FlsI4- complex ion (6) in which a five-co-ordinate antimony is encountered two of the fluorines in the SbF polyhedron being shared.1° A more complex structure is found for KSb4F, and for several isomorphous compounds.Here four independent antimony tri- fluoride groups are Ioosely joined to a fluorine ion (the thirteenth).40 This structure could be described as involving a solvated fluoride ion and 36 Gutmann 2. anorg. Chem. 1951,266 331. 36 Andersson and Lindqvist Acta Chern. Scand. 1955 12 79. a8 GrdeniC and SCarniCar Proc. Chern. Soc. 1960 147. as Bystrom and Wilhelmi Arkiv Kemi 1951 3 373. 40 Bystrom and Wilhelmi Arkiv Kerni 1951 3 17. Muetterties and Phillips J. Amer. Chem. Soc. 1957 79 3686. 180 QUARTERLY REVlEWS classified along with the arsenic and antimony trichloride solvates men- tioned earlier. Various bismuth compounds of the form M1BiX4 have been recognised but none investigated structurally.It is tempting to assign a simple ionic form to compounds such as BiCl,,NOCl (e.g. NO+BiC14-) but there is no supporting evidence. The conductance of arsenic trifluoride and arsenic trichloride and the existence of AsF,,SbF and AsF,,BiF suggest an AsX,+ cation. However no evidence direct or indirect is available in support. One feature of the trihalides of particular interest is their ability to act as electron-pair donors and to a lesser extent as acceptors. In the role of donor the high electronegativities of the halogens would lower the bonding capacity but this effect should diminish on passing from fluorine to iodine Little is known however of the complexing tendencies of the tribromides and tri-iodides so that discussion must be confined to the trichlorides and trifluorides.The withdrawal of electrons by the halogens would favour acceptor n-bonding by the d-orbitals. Investigation has so far been con- fined to the examination of the stability of complexes with obvious accep- tor molecules such as the boron halides. Recent work41 reports the failure of the trichlorides or tribromides of arsenic and antimony to form complexes with halides of boron aluminium or gallium; indications are that previous reports of addition compounds are incorrect. Weak complexes between phosphorus trichloride and these three halides are however found. The gallium trichloride-phosphorus trichloride system contains the complex C1,P-+GaCl3 in the solid but it is not stable A number of adducts between the trihalides and organic amines and oxides are known with varying molecular Examples of 1 1 compounds are Me,As,PCl, Me,N,AsCl, Me,N,SbCl, and Me,N,PBr, of 1 :2 compounds Et3N,2SbC1, of 2:l compounds 2Me3P,PCl and ~ M ~ P A s C ~ .~ ~ * * ~ The formation of 1 :1 complexes appears similar to the formation of [R,P,AsMe,]+Cl- by the reaction of trialkylphosphines and dimethyl- chl0roarsine,~6 however although the formulation [R3X,M*1Cl2]+Cl- might hold for the 1 :1 complexes it is difficult to extend this to the other complexes. In view of the probable acceptor characteristics of MX molecules other formulations involving p,-d bonding might be expected. A decrease in P-N bond energy from 6.4 kcal. in Me,N,PCl to 3.1 kcal. in Me,N,PBr is either the result of the lower electronegativity of the bromine (leading to a lower tendency for the phosphorus to act as ac- ceptor) or the consequence of the greater steric effect of the bromine.41 Holmes J. Inorg. Nuclear Chem. 1960 12 266. Greenwood Perkins and Wade J. 1957 4345. 4s Trost Canad. J. Chem. 1954 32 356. 44 Holmes and Bertaut J. Amer. Chem. SOC. 1958 80 2980 2983. 4b Holmes J. Amer. Chem. SOC. 1960 82 5285. I6 Coates and Livingstone Chem. and Znd. 1958 1366. PAYNE HALIDE$ OF THE! PHOSPHORUS GROUP ELEMENTS 181 Although phosphoryl chloride functions as a donor and arsenic and antimony trichloride as acceptors complexes between these compounds appear to be best described in terms of purely dipole-dipole interaction (comparable with the solvates referred to earlier). The compounds AsCl,,POCl and SbC1,,2POC13 33 are shown by Raman and infrared spectra not to involve ionic species.*' Similar observations cover the systems BiCl,-POCl, AsC1,-Me,CO and SbC1,-Me,CO in which the compounds BiC1,,2POC13 AsCl,,Me,CO and SbC1,,2Me2C0 O C C U ~ .~ * * ~ ~ A 1 1 compound between antimony trifluoride and dioxan is reported.50 The adducts of ammonia primary and secondary amines and the tri- halides cover a wide range of molecular ratios. No systematic investigation has been undertaken. However in the case of phosphorus trichloride and ammonia at -80" the triamide P(NHa, is obtained and at -20" the compound [.NH.P(NH,)-],. No evidence of simple addition compounds of the type PCl,,xNH is found.51 Phosphorus trichloride with an excess of methylamine yields the compound (7) P4N6Me6 52 and with a n h e the compound [.NPh-P(NHPh).],.53 Reaction of the trifluoride and tri- chloride of arsenic with secondary amines gives simple dialkylamido- dihalogenoarsenites R,N.AsX ; with primary amines hydrogen halide is readily eliminated from the products to give alkylimidohalogeno- arsenites (-NR.ASX.),.~* The list of addition compounds of the trihalides is long and contains many compounds of unexpected composition whose structures are un- known.The application of modern techniques is of importance in this field ; already certain misconceptions have been revealed for example the reported addition compounds of arsenic trichloride and tribromide with copper and silver metal e.g. AsBr3,3Ag 2AsC13,7Cu have now been shown by X-ray diffraction to consist of a mixture of silvedi) or copper(1) halides with amorphous arsenic.55 The addition compound 2AsF3,3S0 has been shown by nuclear magnetic resonance to have the complex structure (8).56 47 Kinnell Lindqvist and Zackrisson Acta Chem.Scund. 1959 13 1159. Zackrisson and Alden Acta Chem. Scand. 1960,14 994. I s Lindqvist and Einarsson Acta Chem. Scand. 1959 13 420. Haendeler Glazier and Breck J. Amer. Chem. SOC. 1953 75 3845. 61 Becke-Goehring and Schulze Chem. Ber. 1958 91 1188. 63 Holmes and Forstner J. Amer. Chem. SOC. 1960 82 5509. 63 Michaelis and Schroeter Ber. 1894 27 491 ; Grimmel Guenther and Morgan J. Amer. Chem. SOC. 1946 68 539; Goldschmidt and Lautenschlager Annulen 1953 580 68. 64 Olah and Oswald Canad. J. Chem. 1960 38 1428 1431. 8K Rudorff and Gelinek Chem. Ber. 1957 90 2654. K8 Gillespie and Oubridge Proc. Chem. SOC. 1960 308. 182 QUARTERLY REVIEWS One aspect of the trihalides of the phosphorus group which has been recently exploited is their ability to form d,-d bonds leading to com- pounds such as Ni(PF,), and Ni(CO),(SbCl,).The properties of the trihalides are considerably modified in these compounds so that for F*A” 40 %S-O-S’/ O\ / qs\? F / I &‘F F fs (8) example Ni(PF,) can be vaporised in the presence of steam without undue hydrolysis The phosphorus trihalides are particularly favourable ligands for the platinum metals and many of the adducts have been known for a considerable time. Phosphorus trifluoride and plati- num(@ chloride give two volatile compounds dichlorobis(trifluoro- phosphine)-pp‘-dichlorodiplatinum (PF,),Pt,Cl and dichlorobis(tri- fluorophosphine)platinum (PF,),PtC12.5e Phosphorus trifluoride exhibits a strong trans-effect and in this respect is similar to carbon monoxide.Both of these ligands bind by means of simple a-bonds and also by a strong 71-bond formed by the overlap of filled dp-orbitals of the metal with vacant orbitals on the ligand. The pentahalides and related compounds Certain of the simple pentahalides of this group have been known for a long time; one in particular phosphorus pentachloride has played an important part in the development of chemical ideas first as an example of equilibrium in thermal dissociation then as a valency problem and of recent times as an introduction to the extensive ionic chemistry of com- pounds of this type. Whilst it is common to talk of specific pentahalides this is in many cases merely a reference to the stoicheiometric composi- tion and in no way reflects the structural state.It will be convenient to discuss initially our knowledge of those molecules that are of an un- ambiguous MX or MX,,Y form and then to survey the various other derived compounds. There are no examples of pentahalides of bismuth other than the fluoride. Molecular Species.-There is little doubt concerning the molecular existence of the compounds PF, PCl, PCI,F, PCI,F AsF, SbF, SbCI, SbCl,F, SbCI,F, and SbC1,F. That of others such as PCl,F, PBr,F, PBr,F and SbBr,F, is less certain. The number of possible species even excluding stereoisomers which might be obtained by allowing for simple and mixed fluorides chlorides bromides and iodides of phos- phorus arsenic and antimony is over 150 only about a tenth having so far been recognised.ti7 Wilkinson J. Amer. Chem. SOC. 1951 73 5502. 68 Chatt and Williams. J. 1951. 3061. PAYNE HALIDES OF THE PHOSPHORUS GROUP ELEMENTS 183 The simple halides are prepared by direct reaction of an excess of the halogen with the element. Fluorides are conveniently prepared by fluorina- tion of the chlorides. The mixed halides are prepared by the addition of halogen to the appropriate trihalide e.g. PC12F3,59*60 PC13F2,61 PC14F,61 PBr3F,,lS PBr,F,1g*63 SbC12F3,64*65 SbC1,F2,64*66 SbBr2F3.67 Molecular PC14F is one of the products of the pyrolysis of [PCl4]+[PF6]- in arsenic trichloride suspension6* and molecular SbC1,F can be obtained by the action of arsenic trifluoride on antimony penta~hloride.~~ The direct addition of fluorine to a trihalide has not yet been accomplished. The reaction of fluorine and phosphorus trichloride in the vapour phase leads to appreciable yields of [pCl,]+ [PF,] - together with molecular PC14F.70 The reaction of halogen with trihalide to give a pentahalide has not been closely studied.It is generally an exothermic reaction which occurs readily on direct mixing. In the case of phosphorus trifluoride and chlorine the reaction being homogeneous proceeds very slowly in the gaseous phase at room temperat~re.~~ However in the liquid phase the reaction is rapid.59 In a stainless steel vessel reaction in the gas phase leads to only phosphorus pentafluoride and a large amount of solid probably [PC14]+ [PF,]-.60 For most of the compounds above evidence of the molecular form is largely indirect. Examination of the molecular structure of PF5,72 PCl,,73 PF3C12,72 AsF,,~~ and SbC1,,73 in the vapour phase shows all to possess the expected general features of the trigonal bipyramid.In the case of phosphorus pentachloride and antimony pentachloride the apical bonds are longer than the equatorial bonds (PCl, apical 2-19 5 0.02 equatorial 2-04 5 0.06; SbCl, apical 2-43 0-06A). The extra repulsion exerted on the bonding pairs in the apical positions relative to those in equatorial positions can explain this elongation ; alternatively the explanation has been given in terms of incomplete hybridisation leading to a set of three sp2 equatorial bonds and a further set of two dp apical bonds. Antimony pentachloride forms a molecular lattice in the solid 0.06 equatorial 2.31 6D Kennedy and Payne J. 1959 1228. 6o Muetterties Bither Farlow and Coffman J.Inorg. Nuclear Chem. 1960 16 52. 62 Moissan Compt. rend. 1885 100 1348. 63 Kolditz and Bauer 2. anorg. Chem. 1959 302 241. 64 Ruff Zedner Knoch and Graf Ber. 1909 42 4021. 6s Henne and Trott J. Amer. C'hent. SOC. 1947 69 1820. 66 Swarts Buff. Class Sci. Acad. roy. Befg. 1895 (3) 29 874. 67 Henne U. S. P. 1,984,480 (1931). 68 Kolditz 2. anorg. Chem. 1956 286 307. * 6D Kolditz Z. anorg. Chem. 1957 289 125. 70 Payne unpublished results. 71 Wilson J. Amer. Chem. SOC. 1958 80 1338. 72 Braune and Pinnow 2. ph-vs. Chem. 1937 B 35 239; Brockway and Beach J. Amer. Chem SOC. 1938 60 1836; Gutowsky and Hoffman J. Chem. Phys. 1951 19 1259. Booth and Bozarth J. Amer. Chem. SOC. 1939,61,2927. Roualt Ann. Phys. 1940 14 78. 74 Akers Diss. A h . 1955 1638. 184 QUARTERLY REVIEWS and here Sb-Cl apical = 2-34 A Sb-Cl equatorial = 2.29 the elonga- tion appears to be appreciably reduced by crystal-lattice forces.In PCl,F in accord with the ideas of repulsion of bonded and non-bonded pairs the chlorine atoms are in the apical positions and the fluorines in the equatorial positions. The dissociation of the phosphorus pentachloride in the vapour to give phosphorus trichloride and chlorine is familiar and has been studied on many occasions. Antimony pentachloride behaves in a similar dissociation in each case at 1 atmosphere amounts to a few per cent at 100”c. PF, AsF, and SbF have not been observed to undergo such reversible thermal dissociation. PCl,F shows no dissociation up to 1 50° above this temperature disproportionation (partially reversible) occurs.Indirect evidence suggests that PC1,F dissociates reversibly to phosphorus dichlorofluoride and chlorine.’ Ready thermal dissociation appears to be restricted to PCl, SbCl and PCl,F in which chlorine atoms are in adjacent positions and hence intramolecular elimination of chlorine is easy. In PC12F3 in which the chlorine atoms are in the apical position the dissocia- tion reactions does not occur. There is no evidence for the existence of phosphorus pentabromide in the v a p o ~ r . ~ ~ For many of the pentahalides reliable molecular-weight data in the vapour phase are missing. In the case of antimony pentafluoride preliminary results on the molecular weight suggest that appreciable association occurs; values as high as 1230 (SbF5 = 217) just above the boiling point have been P,Cll0 species have been recognised mass spectrometrically in the vapour.15 Little information on the liquid state is available.All pentahalides so far examined appear to have low or very low electrical conductance and with the exception of antimony pentafluoride are non-associated mobile liquids. Antimony pentafluoride is a viscous liquid shown by nuclear magnetic resonance studies to consist of chains of SbF groups each sharing two fluorines with two neighb0u1-s.~~ Ionic and other less well-defined Species.-There are now several penta- halides which are recognised as possessing ionic lattices. In only a few cases has the evidence for an ionic lattice been obtained directly from X-ray examinations more often the evidence has been obtained indirectly from conductance or other properties of solutions in solvents known to favour break down of ionic lattices.Phosphorus pentachloride was the first compound in which an ionic lattice ([Pel,]+ [PcI6]-) was recognised.aO Solutions of phosphorus pentachloride in various polar solvents are electrolytic conductors and in methyl cyanide the ions [PCI,]+ and pcl6]- 75 Ohlberg J. Amer. Chem. SOC. 1959 81 811. 76 Braune and Tiedje Z. anorg. Chem. 1926 152 39. 77 Harris and Payne J. 1958 3732. Hubb Peacock and Robinson unpublished results 1951. See also Dodd and 7B Hoffman Holder and Jolly J. Phys Chem 1958 62 364. 8o Clarke Powell and Wells J. 1942 642. Robinson “Experimental Inorganic Chemistry,” Elsevier New York 1954 p. 21 5. PAYNE HALIDES OF THE PHOSPHORIJS GROUP ELEMENTS 185 were identified by transport experiments.81 The exchange of radioactive chloride ion and the phosphorus pentachloride in methyl cyanide solution shows clearly that the [Pel,]- anion is less resistant to attack by chloride ion than the [PCl,]+ cation.82 Likewise the ready formation of [PCl,]+ [PF,]- by the action of arsenic trifluoride on a solution of phosphorus pentachloride in arsenic trichloride is evidence for the relative ease of attack on the anion.83 The P-Cl distances in cation and anion of 1-98 A and 2.06 A respectively suggest that in the former (tetrahedral form) the binding is stronger.In carbon tetrachloride solution the molecular weight suggests that association may occur (M = 259 at a concn. of 0.0241~ and 473 at 0.0959~. PCl, M = 208).82 Exchange experiments involving molecular chlorine and phosphorus pentachloride in carbon tetrachloride solutions show that the five chlorines attached to phosphorus are not equally reactive.The results can be interpreted to show that three chlorine atoms react more rapidly than the other two. The exchange may involve as a transition state the addition of the halogen molecule to the equatorial Although solid phosphorus pentabromide consists of [PBr,]+ and Br- ions,85 in methyl cyanide it is an electrolytic conductor in which the ions as shown by transport experiments are [PBr4]+ and [PBr,]-. The anion is presumably on the limit of stability being stabilised only by solvation. Complexes involving the [PBr,]- anion have not been reported yet.86 Conductance of a lower order is also observed in solutions in arsenic tri~hloride,~' sulphur dioxide,87 bromine,88 and nitroben~ene.~~ In nitro- benzene the equilibrium between phosphorus tribromide and bromine appears to involve both undissociated PBr and PBr,,86 but there is no indication as to the nature of these species.Fluorination by arsenic trifluoride of phosphorus pentabromide suspended in carbon tetrachloride or carbon disulphide gives [PBr,]+ [PF,]- the reaction being more complex than the fluorination of [PCI,]+ [PC&]- in arsenic trichloride but again evidence of the resistance of the cation to attack is Pyrolysis of [PCl,]+[PF,]- leads to both PC1,F and [PCl,]+F-.68 The change from molecular PC14F to ionic [PCl,]+F- has been observed to occur by a first-order reaction.g1 Pyrolysis of [PBr,]+ [PF,]- does not yield [PBr,]+F- which must be prepared by the action of bromine on phosphorus dibromofluoride.In this reaction the molecular form appears to exist at Payne J. 1953 1052. 82 Kolditz and Hass 2. anorg. Chem. 1958 294 191. 83 Kolditz 2. anorg. Chem. 1956 284 144. 84 Downs and Johnson J. Amer. Chem. SOC. 1955,77 2098. 85 van Driel and MacGillavry Rec. Trav. chim. 1943 62 167; Powell and Clark 86 Harris and Payne J. 1956 4617. Walden 2. phys. Chem. 1903,43 434. 88 Plotnikov Z. phys. Chem. 1904 48 230. 8 9 Finkelstein Z. phys. Chem. 1925 115 306. O D Kolditz and Feltz 2. anorg. Chem. 1957 293 155. 91 Kolditz 2. anorg. Chem. 1957 293 147. Nature 1940 145 971. 186 QUARTERLY REVIEWS -30" and below but is transformed by heat into the stable ionic f0m1.6~ The compound [PCl,]+ [PCl,F]- contaminated with IpCI,] +F- is the pro- duct of the controlled pyrolysis of [PCl,]+ [PF6]- suspended in carbon tetrachloride.Above 110" it disproportionates to PCl and PC1,F.l5 The analogous compound [PCl,]+ [PCl,Br]- is the product of the addition of bromine to phosphorus trichloride in arsenic trichloride solution. The nature of the anion was established in addition to evidence from the analytical and conductance data by direct fluorination with arsenic trifluoride to give [PCI,]+ [PF6]-.92 Despite the considerable amount of work reported aboke relatively few of the numerous cations and anions derived from single and mixed pentahalides have been recognised. The [PCI,]+ cation is clearly a particularly stable species but with a suitable synthetical route substituted [PX,]+ cations should be obtainable. Amongst the chlorofluoro-anions [PCl,F,]- [PC13F3]- [PC12F4]- and [PClF,]- remain to be recognised.No iodide species has yet been encountered. In addition to the large number of species made possible by varying the nature of the attached halogen there is the added feature of stereoisomerism which might be encountered in certain of the octahedral anions to increase still further the total number of individual compounds. The phosphorus trichloride-bromine system is complicated and many groups of workers have investigated it without complete clarification. Since it was first observed in 1847 that phosphorus trichloride and bromine reacted together exothermically only one compound of certain constitution has been produced [PCl,]+ [PCI,Br]-. 92 Recent Russian workg3 on the system has led to the recognition of two maxima in the phase diagram at the compositions PCl,Br and PC13Br18 and compounds corresponding to these compositions have been reported.A further composition PC13Br8 appeared of significance from viscosity data. In nitro- benzene the compounds behaved as a series of quasiphosphonium com- pounds [PCI,Br]f[BrBr,,]- with n = 1 3 and 8. In American work no simple solid phase was found. Instead an unstable aggregate ofcomposition PC13.0Br5. which broke down to a compound PCI,. 67Br0. 33 was A cryoscopic study of phosphorus trichloride and bromine in nitrobenzene shows the existence of PX and PX species.95 In all the compounds so far reported only the [PCl,]+ [PCI,Br]- and the more complex PCI,. 67Br0.33 appear to possess any degree of stability. Even the former breaks down in non-polar solvents to give PCI and PBr,.The existence of solids contain- ing such a high proportion of bromine as PC1,BrI8 may in fact arise from solvation by bromine molecules or the formation of a type of inclusion compound. Several adducts with carbon tetrachloride have been found e.g. PBr,F 2CC14 PBr5,2CC1,. With the existence of such an extensive group of ionic compounds in the 82 Kolditz and Feltz Z . anorg. Chern. 1957 293 286. 83 Kialkov and Kuz'menko Zhur obshchei Khim. 1951,21,33; 1952,22 1290 1335. g4 Popov Geske and Baenziger J. Amer. Chem. SOC. 1956 78 1793. 86 Harris and Payne J. 1956 4613. PAYNE HALIDES OF THE PHOSPHORUS GROUP ELEMENTS 187 case of phosphorus equally extensive ranges of compounds should be expected for arsenic and antimony if not bismuth. However the results to date are somewhat limited being centred on the [AsCl,]+ cation.The failure of numerous attempts to prepare arsenic pentachloride even in an ionic form is significant. The AsCI molecule and the [AsC16]- ion are presumably inherently less stable than those involving phosphorus or antimony. The core of phosphorus in the quinquevalent state is s2p6 whereas in arsenic and antimony it is dlO; promotion energy data suggest that the 4s and 4p and the 4d orbitals in arsenic will be separated by an appreciably larger energy difference than the 3 4 3p and 3d in phosphorus or the 5s 5p and 5d orbitals in antimony with a consequent lowering of the stability. The extent of the effect will however depend on the electro- negativity of the bonding atom being less for fluorine than for chlorine.In the tetrahedral state other factors notably that of r-bonding are involved and the [AsCl,]+ cation is comparable to [PCI,]+ in stability. Compounds such as AsCI,,PCI 96 have been known for a long time and can now be recognised as [AsCl,]+ [PC16]-,97 others are [AsCl,]+ [SbCl,]-,98 and [AsCI,]+ [ A S F ~ ] - . ~ ~ * ' ~ ~ An attempt to prepare the compound AsCl,+F- by the pyrolysis of [AsCl,]+[AsF,]- in an analogous way to the dis- proportionation of [PCl,]+F- failed the products being only arsenic trifluoride and chlorine.lol The series of [AsCl,]+ complexes is extensive all involving anionic metal halide species e.g. [GaCl,]- [AlCl,]- etc. Many similar [PCl,]+ complex compounds are known.lo2 Only a few derivatives of the pentafluoride and pentachloride of anti- mony are known.The phase diagram of the SbF,-SbC1 systems is complex with at least six individual phases. (SbF,),(SbC1,)2 and SbF,CI were isolated as crystalline solids but no further detail is available.64 Much of this work has centred not on the isolation of the mixed pentachlorides but rather on the preparation and use of these compounds as fluorination catalysts and is reported only in the patent literature. The compound [SbCl,]+F- 69 is prepared by the fluorination of antimony pentachloride by arsenic trifluoride. Bismuth pentafluoride is the product of the action of fluorine on bismuth trifluoride at 500" and is a black solid with a high melting point. A series of addition compounds with LiF NaF KF and AgF are reported.lo3 The complexes of the pentahalides are numerous; some are clearly of the type [MVX4]+Y- or Y+[MVX6]- but many are obviously more com- plicated e.g.(SbF,),Te or SbF5N0,.lo4 No further reference will be made 86 Cronander Bull. Soc. chim. France 1873 [2] 19,499. g7 Gutmann Z . anorg. Chem. 1951 264 151. Gutmann Monatsh. 1951 86 473. Kolditz 2. anorg. Chem. 1955 280 313. loo Dess Parry and Vidale J. Amer. Chem. Soc. 1956 78 5730. lol Kolditz 2. anorg. Chem. 1957 289 128. lo2 Groeneveld Rec. Trav. chim. 1952 71 1152. lo3 Fischer and Rudzitis J. Amer. Chem. Soc. 1959 81 6375. lo' Aynsley Peacock and Robinson Chem. and Ind.. 1951 11 17. 188 QUARTERLY REVIEWS to the large number of obviously ionic complexes which have been investigated largely by conductometric methods. Only one structure has been fully determined namely that of the complex SbCl, POCl, in which the oxygen of the POCl occupies one of the octahedral positions around the antimony :lo5 similar complexes are formed with sulphoxides and ~ulphones.~~ The tendency of phosphorus pentafluoride to form simple molecular complexes apart from the [IT6]- anion is exemplified by the com- pounds PF,,NMe and PF,,CH,.CN.Phosphorus pentafluoride thus exhibits the general acceptor properties of a Lewis acid with a wide range of bases such as ethers sulphoxides amines amides and esters.60*106 This strong acceptor tendency makes phosphorus pentafluoride comparable with boron trifluoride in catalytic activity e.g. in polymerisations.lo6 Adducts of the pentahalides with trimethylamine and trimethylphosphine are (Me3P),,PCl, (Me,N),SbCl, Me,P,SbCl, and (Me,P),,SbCl,. Reaction of the pentachlorides with trimethylarsine and trimethylstibine leads to appreciable reduction often as far as the elements.Phosphorus pentachloride appears to be more readily reduced than antimony penta- chloride in these reactions.44 The reaction of the pentahalides with ammonia and primary and secondary amines has been reported but no simple products or a clear picture of the reaction has r e s ~ l t e d . ~ ~ ~ * ~ ~ * The reaction of phosphorus pentachloride and ammonium chloride leading to the phosphonitrilic chlorides has been widely examined.log Other halides of the phosphorus group elements Halides containing bonds between the elements of the group are so far limited to compounds of the form M2X4 for phosphorus and arsenic and to the compound Bi4C14. P214 is readily prepared by the reaction of stoicheiometric amounts of iodine with white phosphorus dissolved in carbon disulphide.Its orange crystals dissociate irreversibly when heated and are hydrolysed to a variety of products among them hypophosphoric acid containing a P-P bond. Solid P214 consists of a molecular lattice containing non-planar units with LIP1 = 102.3" and LIPP = 93-9".110 Its reactions have been very little studied. It is reported to give an adduct with boron tribromide P2I,,2BBr3.ll1 P,Cl is formed when phosphorus trichloride and hydrogen are together submitted to an electrical discharge,l12 or when phosphorus trichloride alone is subjected to a high-voltage discharge between mercury loB Lindqvist and Branden Acta Cryst. 1959 12 642. lo6 Woolf J. Inorg. Nuclear Chem. 1956 3 285. lo' Becke-Goehring and Niedenzu Chem.Ber. 1957 90 2072. lo8 Audrieth and Sowerby Chem. and lnd. 1959 748. log Paddock and Searle Adv. Znorg. Chem. Radiochem. 1959 1 347. 110 Leung and Waser J. Phys. Chem. 1956 60 539. ll1 Tarible C'omp. rend. 1907 132 204. Besson and Fournier Compt. rend. 1910 150 102. PAYNE HALIDES OF THE PHOSPHORUS GROUP ELEMENTS 189 e1ectr0des.l~~ Fluorination of PJ by antimony pentafluoride led mainly to phosphorus trifluoride and pentafluoride however a small amount of a less volatile material of molecular weight approaching that of P2F4 was obtained.l14 The compound AsJ is obtained by the reaction of stoicheio- metric quantities of arsenic and iodine and also by heating a mixture of arsenic and arsenic tri-iodide in carbon di~u1phide.l~~ It is soluble in carbon disulphide in which its molecular weight was determined ebullio- scopically.l16 No antimony compounds of this type have been reported.Bismuth has been reported to give a dichloride a dibromide and a di- iodide and also a monochloride. There has always been much doubt as to the dibromide and di-iodide and discussion must await further investiga- tion. The bismuth-chlorine system has been thoroughly investigated. No evidence of bismuth dichloride was obtained but the existence of a mono- chloride was confirmed. The monochloride forms shining black diamag- netic crystals recognised as containing a tetramer Bi4C14.117 The extent of the contributions to this relatively small field has been enormous especially within the past decade even so the subject is very far from complete and the general features are only just beginning to emerge.Considerable technical skill is involved in handling these halides most of which are readily hydrolysed. That progress has been made in the past is the result of a determination to overcome many difficult problems of technique and the future must equally depend on such careful skilled and determined experimentalists. 113 Finch Canad. J. Chem. 1959 37 1793. Harris. Dersonal communication. lI5 Bambeker and Philipp Ber. 1881,14,2644; Karantassis Bull. SOC. chim. France lI6 Hewitt and Winmill J. 1907 91 962. 11' Corbett J. Amer. Chem. SOC. 1958 80,4757; J. Phys. Chem. 1958,62 1149. 1925 [4] 37 853.

ISSN:0009-2681    

DOI:10.1039/QR9611500173

出版商:RSC    

年代:1961

数据来源: RSC

摘要:

THE THEORY OF CHARGETRANSFER SPECTRA By J. N. MURRELL WHEN solutions of chloranil (yellow) and hexamethylbenzene (colourless) are mixed together an intensely red solution is obtained. The optical density of this solution at 5000 A is proportional to the product of the concentrations of the two components and it is therefore deduced that the red colour is due to the absorption of light by a 1 :1 complex.1 In crystals of these “quinhydrone-like” complexes the two components are stacked one upon the other in parallel and Nakamoto3 has shown that the red colour is due primarily to the absorption of light which is polarised in a plane perpendicular to the planes of the aromatic rings. This is in contrast to the strong absorption of the separate components which takes place when the light is polarised in the plane of the aromatic rings.Many molecular complexes of the above type can be formed between two components one of which is an electron donor (D) and the other an electron acceptor (A). It is now generally accepted that the intense colour of these complexes is to be associated with the transfer of an electron from the donor to the acceptor. Mulliken was the first to put this idea into the language of quantum chemi~try.~ He considered the interaction of a no- bond ground state $(D,A) and a polar excited state #(D+-A-) to produce a stabilised ground state having a wave function and an excited state the charge-transfer state (CHEMISTRY DEPARTMENT THE UNIVERSITY SHEFFIELD) $0 = $(D,A) + h*(D+-A-) $1 = W+-A-) + p#(D,A). X and p will in most cases be small compared with unity.The charge- transfer band is associated with the electronic transition ~,,+$1 and occurs at the frequency u = (El - E,)/h where h is Planck’s constant. The term charge-transfer absorption may be used whenever there is a large electron displacement in going from the ground to the excited state. Thus as well as for the quinhydrone complexes charge-transfer bands have been postulated for hydrated inorganic ions of the type C1-(H20) -f C1(H,O)n- Fe2+(H20) -+ Fe3+(H,0),- Fe3+OH- -+ Fe2+0H Michaelis and Grannick J. Amer. Chem. SOC. 1944,66 1023. Anderson Nature 1937 140 583. Nakamoto. J. Amer. Chem. Suc.. 1952. 74. 1739. Mulliken J. Amer. Chem. Soc. 1950,72 600; 1952,74 811; J. Phys. Chem. 1952 56 801. 191 192 QUARTERLY REVIEWS and for the alkali halides in both the gaseous and the crystalline state e.g.Na+Cl-+NaCl. In addition one can observe intramolecular charge- transfer absorption in substituted organic molecules of the types (1) and (2) and in transition metal complexes.*a Earlier reviews5 have discussed charge-transfer absorption for molecular complexes and inorganic ions. It is the present purpose to expound the theoretical interpretation of charge-transfer spectra in general both inter- and intra-molecular including the most recent work in this field. The emphasis will be on the nature of the charge-transfer absorption rather than on the charge-transfer stabilisation of the ground state. Only the spin-allowed transitions which give rise to the intense absorption bands will be considered. The energy and intensity of the absorption will first be discussed separately in general terms and then the mathematics will be examined in detail.The Energy of the Charge-transfer Band.-The energy of the electronic transition (D,A) + (D+-A-) is given to a first approximation by the expression ID - EA - C where ID is the ionisation potential of the donor EA the electron affinity of the acceptor and C the mutual electrostatic energy of D+ and A- relative to that of D and A. For example HasseP has shown that in crystals of the benzene-Br complex the Br lies per- pendicular to the plane of the benzene ring the complex having CgV symmetry. Benzene is the donor in this case its vertical ionisation potential is 9.24 ev (1 ev = 8068 cm.3.’ The distance from the nearest bromine atom to the centre of the benzene ring is 3.36 A; the Br-Br bond length is 2.28 A.Assuming that the donated electron goes into an orbital centred equally on the two bromine atoms and that the hole it leaves behind is distributed over the six carbon atoms of the benzene ring we calculate a coulombic energy of 3.22 ev for the complex. Unfortunately the electron affinity of Br is unknown although it is said to be positive.* What we require is in fact the vertical electron affinity that is the electron affinity calculated for the same internuclear distance as in the neutral molecule. It probably lies between 0 and 2 ev. From the formula given above we then calculate that the charge-transfer band should occur somewhere 4a Jnrrgensen MoZ. Phys. 1959 2 309. Rabinowitch Rev. Mod. Phys. 1942,14 112; Andrews Chern. Rev. 1954,54,713; Hassel Mol. Phys. 1958 1 241.Watanabe J. Chem. Phys. 1954,22 1564. Orgel Quart. Rev. 1954,8,422; McGlynn Chern. Rev. 1958,58,1113. * Massey “Negative Ions,” Cambridge Univ. Press 1950. MURRELL THE THEORY OF CHARGE-TRANSFER SPECTRA 193 between 4 and 6 ev depending on the value for E(Br&. It is in fact observeds at 4.24 ev. This value would be obtained by using an electron affinity for Br of 1-78 ev if more direct methods fail electron affinities can perhaps be obtained from charge-transfer spectra. It can be seen from the above calculation that the electrostatic contribu- tion to the charge-transfer energy is by no means unimportant. Thus an electron and a positive hole separated by 1 A have a mutual energy of 14.4 ev. If the donor and the acceptor approach near enough to each other then the charge-transfer state (D+-A-) could have a lower energy than (D,A) even though ZD-E~ is positive.This situation arises for the alkali halides. For the states NaCl and Na+Cl- in the gas phase we have I(Na) = 5.14 E(C1) = 3.82 and C(2.36 A) = 6.10 ev. It follows that the ionic state has the lower energy and the charge-transfer band arising from the transition to the covalent state should occur at 4.78 ev it appears,lO in fact at 5-14 ev. If only one component of the complex carries a charge and if this charge is just redistributed in going from the ground to the excited state then the coulombic term will be zero. This will be the case for the hydrated inorganic anions. For Cl-(H;O), there will be some small contribution to the charge-transfer energy arising from the different polarisation energy of the H20 in the field of the C1- compared with that of Cl in the field of H,O-.This can probably be neglected. Daintonll has shown that the energy of the charge-transfer band of the bivalent transition-metal ions is linearly related to the redox potential of the system M2+ + M3+ + e. This suggests that the charge-transfer is to be associated with the transition M2+(H,0) -f M3+(H20),- rather than to M2+(H20) -f M+(H,O),+. Again it is probably the coulombic energy which swings the balance between these two processes being about + 7 ev when the metal acts as donor and - 7 ev when the water acts as donor. For intramolecular charge-transfer the quantities I D and EA of the two parts of the molecule have to be estimated from the values obtained for related molecules. For example in aniline we might take I(NH2-) equal to the ionisation potential of ammonia1 (10-15 ev) and E(Ph-) equal to the electron affinity of benzene13 (-1.63 ev).There will be a small error here because the NH2- group is more or less planar but NH is not. The cou- lombic energies associated with intramolecular charge transfer are often greater than those in the intermolecular case since the donor and the acceptor are closer together being joined by a covalent bond. In the case of aniline there are two degenerate orbitals of the phenyl lo Franck Kuhn and Rollefson 2. Physik 1927 43 155; Miller and Wynne-Jones l1 Dainton J. 1952 1533. l2 Watanabe J. Chem. Phys. 1957 26 543. l3 Hedges and Matsen J. Chem. Ph-vs. 1958 28 950. Andrews and Keefer J. Amer. Chem. SOC. 1950,72,4677. J. 1959 1375.194 QUARTERLY REVIEWS group which can accept the electron (the orbitals being taken to be the same as those of benzene). They have the symmetries shown in (3) and (4). The coefficients represent orbitals to the molecular the contribution of the six carbon 2pn atomic orbital. It is now seen that taking an electron from the non-bonding .;rr-orbital of the NH2 group ($) and putting it into the vacant benzene orbitals gives rise to two charge-transfer states which we might represent1* by +-‘#a and $-‘#b. However although the orbitals #a and #b are degenerate the two charge-transfer states are not. They have different coulombic energies associated with them. Thus $-‘#a has a charge density as in ( 5 ) (obtained by squaring the coefficients) for which the attraction of the positive and negative charges gives an energy of 4-84 ev whilst $-l& has a charge density as in (6) and an associated coulombic energy of 6.03 ev.From the formula I - E - C we now calcu- late that +-‘$a has an energy 1-94 and &‘#b an energy of 5.75 ev. The lowest excited state of benzene is at 4.71 ev and we therefore expect to find this same state in aniline but somewhat perturbed by the NH group it is found at 4-31 ev. However the second excited state of aniline should correspond to the charge-transfer state +-l#b it appears at 5.27 ev. The second charge-transfer state +-l$a cannot be clearly distinguished as it would appear in the same region as the higher excited states of benzene. The charge densities of the excited states of aniline which have been discussed above are those obtained by using the molecular orbitals of the benzene ring.From the valence-bond theory rather different charge- transfer states are considered “resonance” structures (7-9). (7) namely those associated with the familiar If these are used we predict that the donated electron can appear either in the ortho- or the para- but not in the meth- position. This might be a good approximation to describe the electron density in the ground state but evidence from the effect of acid on the spectra of aminoazo-compounds suggests that the electron densities in l4 Murrell Proc. Phys. SOC. 1955 68 A 969. MURRELL THE THEORY OF CHARGE-TRANSFER SPECTRA 195 e0.000 70,000 the excited states are better represented by the molecular-orbital struc- t u r e ~ . ~ ~ The importance of the coulombic contribution to intramolecular charge-transfer absorption is illustrated by the aromatic nitro-amines.For these molecules the compounds in which the amino- and the nitro- groups are in ortho-relation absorb at considerably longer wavelengths than do the other isomers. It is tempting to attribute this to some subtle internal hydrogen-bond effect but probably it is just that in the first excited state where the nitro-group has gained an electron and the amino- group has lost an electron the two charges are closer together for the ortho- compounds than for their isomers. For a series of molecular complexes involving different electron donors and the same electron acceptor it has been shown16 that there is a reason- able linear correlation between the ionisation potential of the donor and the frequency of the charge-transfer band.This can be seen from Fig. 1 - - 30,000 40,000 v (cm?) FIG. 1 . 1 Naphthalene. 7 Cyclohexane. 13 Propene. 2 Mesitylene. 8 trans-But-2-ene. 14 cis-Dichloroethylene. 3 o-Xylene. 9 Benzene. 15 trans-Dichloroethylene. 4 p-Xylene. 10 cis-But-2-ene. 16 Diethyl ether. 5 2-Methylbutadiene. 1 1 Chlorobenzene. 17 Cyclopropane. 6 Toluene. 12 Butadiene. 18 t-Butyl alcohol. which records the charge-transfer band for a number of Z complexes. The expression Z - E - C is only a crude approximation to the energy of the charge-transfer state the same order of approximation as saying that +(D,A) and $(D+-A-) represent the wave functions of the ground and the charge-transfer state of the molecule or complex. If there is some overlap of donor and acceptor orbitals there is some “mixing” of these McConnell Ham and Platt J.Chem. Phys. 1953 21 66; Hastings Franklin The relation between the ionisation potential of the donor and the frequency (v) of the charge-transfer band for some iodine complexes.a l5 Murrell J. 1959 296. Schiller and Matsen J. Arner. Chem. SOC. 1953 75 2900. 1 9 6 QUARTERLY REVIEWS two states with resulting stabilisation of the ground state and destabilisa- tion of the excited state. The energy of the charge-transfer band is increased by both effects. A necessary criterion for observing a charge-transfer band in the accessible region of the spectrum appears to be that the inter- action between $(D,A) and $(D+A-) be not too large. Thus one might expect at first sight to observe a charge-transfer band for BF,-NH [ I(NH,) = 10.5 E(BF,)17 = 2 C(1-6 A) = 9 ev] somewhere in the visible or near infrared region of the spectrum.In fact one gets no absorption below 2300 A at which point the NH group itself starts to absorb. The reason would appear to be that t,h(DA) and #(D+A-) interact too strongly owing to the large overlap between the donor and the acceptor orbitals. The stabilisation of the complex is experimentally about 2 ev. A theoretical expression has been proposed by Hastings et a1.16 to cover these strongly interacting cases. Since the ground and the charge-transfer state have very different electron distributions it is to be expected that changing the dielectric constant or polarity of the surrounding medium should have an effect on the energy of the charge-transfer band. A polar solvent will stabilise D+-A- relative to DA.There appears to have been no experimental work carried out to examine this point for molecular complexes but the effect is observed among Brooker's dyes.ls*lg The merocyanine dye shown below has two low-energy structures (10) and (1 1). In pure pyridine this dye has its first absorption band at Amax. = 540 mp. When successive amounts of water are added there are shifts first to longer wavelengths and then back Me Me \ / Me Me a > C = = C H - C H G O - a L > C - C H = C H O O - Ph (I I) to shorter wavelengths (see Table 1). The interpretati~nl~ of this type of TABLE 1. The efect of polarity of the solvent on the wavelength of a typical merocyanine dye. Pyridine (vol. %) 100 95 90 80 70 60 ~ m a x . ( f l l p ) 540 548 552 554 554 556 Pyridine (vol.%) 50 40 30 20 10 5 ma x . ( m ~ ) 554 554 553 552 548 545 behaviour is that in pyridine the quinonoid is more stable than the ionic form but that on addition of water the ionic form is relatively more stabilised. The two structures have about the same energy with 1 :1 solvent mixtures; in water-rich solvents the ionic form is the more stable. l7 Savard and Simons J. Chem. Phys. 1939,7,2. l9 Platt J. Chem. Phys. 1956 25. 80. Brooker and Keyes J. Amer. Chem SOC. 1951 73 5356. MURRELL THE THEORY OF CHARGE-TRANSFER SPECTRA 197 The interaction between the ionic and the quinonoid forms prevents the frequency of the charge-transfer band from becoming zero at the crossing point. The Intensity of Charge-transfer Bands.-The integrated intensity of an absorption band the oscillator strength is given theoretically by f= (1.085 x loll) Mo?uo1 where uol is the frequency in cm.-l and Mol the transition moment is given by n Mol = e 1.. . 1 # o ( x r i ) # l d q . . . d7,. i= 1 where e is the electron charge and ri is the position vector of the ith elec- tron. If #o and lCri are one electron functions Mol is just the dipole moment of the charge density a,bo $i* Since t,h0 and t,bl must both be eigenfunctions of the donor-acceptor pair they must be orthogonal. That is s . . . s #o#l dT1 . . . dr = 0. Now if D and A are so far apart that none of the donor orbitals has any region of overlap with any of the acceptor orbitals then the ground and the excited state of (D,A) are orthogonal to all the charge-transfer states (D+-A-) and transitions from the ground state to any of the charge- transfer states are forbidden.In order to observe charge-transfer bands it is necessary that there must be some region of overlap between the orbitals of D and those of A. Let us consider for simplicity a one-electron system such that the ground state has an electron in a donor orbital #d and in the charge-transfer state this has been transferred to an acceptor orbital#,. That is $(D,A) = $d and #(D+-A-) = +a. Under the influence of the Hamiltonian these interact to give a new ground stated +’d and a new excited state +fa and the intensity of the charge-transfer band is proportional to the square. of the dipole moment of the electron density +’d$la. Fig. 2 shows this type of behaviour for two 1s-orbitals. The wave functions and density are plotted along the line joining the two nuclei.The dipole moment of the electron density $’d$’a is directed along the internuclear axis so that it is the component of the light whose electric vector is parallel to this axis that is absorbed. The charge-transfer band is said to be polarised along the inter- nuclear axis. Nakamoto’s observation that the first absorption band of the “quin- hydrones” is polarised in a plane perpendicular to the planes of the two components is evidence that these bands have charge-transfer character. However to find that a band was polarised in a plane parallel to the aroma- 198 QUARTERLY REVIEWS FIG. 2. Charge-transfer between orbitals of the same symmetry. tic rings would not mean that it was not a charge-transfer band. The transfer of charge between D and A is not always associated with a transi- tion polarised in the direction of charge transfer.For example in Fig. 3 FIG. 3. Charge-transfer between orbitals of different symmetry. we see that if +d is an s-orbital and +a a p-orbital the charge-transfer band will be polarised at right angles to the internuclear axis. One can observe a charge-transfer band of this type even though the charge-transfer state cannot by symmetry lead to any stabilisation of the ground state. In aniline the charge-transfer state +-It,hb will give rise to a band polarised along the two-fold axis and this state can stabilise the ground state and give rise to the band at 5.27 ev; on the other hand does not stabilise the ground state and transition to this state would be polar- ised at right angles to the two-fold axis.As stated earlier there must be some region of overlap between the donor and the acceptor orbitals for a charge-transfer band to appear and the more the overlap the stronger the bands. It follows that the intensity of charge-transfer bands is sensitive to steric effects which inhibit this over- MURRELL THE THEORY OF CHARGE-TRANSFER SPECTRA 199 lap. For example ortho-substituents in NN-dimethylaniline or nitro- benzene turn the dimethylamino- or the nitro-group out of the plane of the aromatic ring reducing the overlap between the donor and the accep- tor orbitals Wepster and his co-workers20 have estimated the angle by which the substituent groups are turned out of the plane from the reduc- tion in intensity of the intramolecular charge-transfer band. Some of their results are given in Table 2.TABLE 2. The angle (8) by which the substituent group is twisted out of the aromatic plane by ortho-substituents.20 x E e Nitrobenzene 25 1 8900 0 o-Nitrotoluene 25 1 6070 34" 2-Nitro-m-xylene 250 1500 66" NN-Dimethylaniline 25 1 15500 0 NN-Dimethyl-o-toluidine 248 6360 50" N,N,2,6-Tetramethylaniline 262* 2240 68 " * This maximum is almost certainly not associated with the charge-transfer band but belongs to the band which in dimethylaniline occurs at 2950 A and is shifted to the blue region by ortho-substituents. Further discussion of the intensity of charge-transfer bands will follow the more detailed investigation in the next section. Mathematics of Charge-transfer Spectra.-In this section we consider the details of the interaction between the ground state and the charge- transfer state and its effect on the intensity of the charge-transfer band.We again consider just a one-electron system in which we write the ground and the charge-transfer state wave functions as We now specify that +d shall be an eigenfunction of the neutral donor and (ba an eigenfunction of the acceptor negative ion. If we write the total Hamiltonian for the complex as H = V(D+) + V(A) + T (2) where V(D+) is the eIectrostatic field of Df V(A) the field of A and T the kinetic energy operator then from our definition it follows that Now if the perturbed ground and the charge-transfer state which result from the interaction of the donor and acceptor are written 2o Burgers Hoefnagel Verkade Visser and Wepster Rec. Trav. chim. 1958 77 491. 200 QUARTERLY REVIEWS then by perturbation theory we have ( 5 ) Had - SadHdd.Had - SadHaa had = Ada = E d - E a ' E a - E d where Had = J4aH4dd-r and Sad = J+a&dT. expressions (3) we obtain coefficients By substituting for H as given by eqn. (2) and making use of the xad(Ed - Ea) = Vad(A) - SadVdd(A) and h ( E a - Ed) = vad(D+) - SadVaa(D'). (6) Introducing the normalised charge density $d&/Sad we have We notice the lack of symmetry in these expressions had#-hda. The integrals W involve the interaction of the electrostatic field of the acceptor or the donor positive ion with a charge density which integrates to zero. The overlap density&+ is concentrated in the region between D and A if the orbitals $bd and +a are about the same size. However if one of these orbitals is much smaller than the other then the overlap density will be concentrated near the smaller of the two.If +d is small compared with +a then $d$a is centred on D and it follows that Wda(D+) will be large but Wad(A) small. On the other hand if $d is large compared with +a then +d$a is centred on A Wad(A) will be large and wda(D+j small. This can be seen from Fig. 4 which shows the electron density [ (+d+a)/Sda -+a2] 0 A FIG. 4. The overlap density between Is-orbitals of different size. MURRELL THE THEORY OF CHARGE-TRANSFER SPECTRA 201 for two 1s-orbitals of the type + = ((3/7r)te5r plotted for various values of (a and <d along the internuclear (The larger ( the smaller is the orbital.) In general we expect +a to be larger than #)d since it is an eigenfunction of a negative ion.We therefore predict that Wda.(D+) 9 Wad(A). (8) This prediction is supported not only by an examination of the overlap density but also by the fact that V(D+) falls off much more slowly than V(A) the latter being zero outside the electron cloud of A. From expression (8) it follows that IAda I> IAadI- (9) This result can also be seen from the condition of orthogonality between +dl and +a' one deduces Ada + Aad + S a d == 0. (10) If S a d is positive and Ea> Ed then Aad will be positive and it again follows that lhda I > IAa 1. In other words there is more ground state introduced into the charge-transfer state than vice versa. If Wad(A) is approximately zero then the ground state is and the charge-transfer state +dt = +d (1 1) (12) (1 3) $d = +a - Sad$d. - A E d 1 had2(Ea - Ed) The stabilisation of the ground state is given by and the elevation of the charge-transfer state by A E a 2 hda2(Ea - Ed).Again it follows that the excited state is generally destabilised more than the ground state is stabilised. So much for the energy of the charge-transfer band. Its intensity is determined by the transition moment between $dl and #at Mda' = J$d'M&ldT = M d a + AdaMdd + AadMaa + a term in A' Use of the orthogonality relationship (10) converts this into (15) This type of expression was first obtained by M~lliken.~ It shows that there are two contributions to the transition moment. The first is pro- portional to the dipole moment of the transferred electron and the hole 21 Murrell J. Amer. Chern. Soc. 1958 81 5037. 202 QUARTERLY REVIEWS it leaves behind (this is itself proportional to the distance through which the charge is carried) and is related to the stabilisation of the ground state through the coefficient had.We have shown that in general had is small so that the first term in eqn. (16) may perhaps be neglected. However Mulliken has pointed out that even if had is zero and there is no stabilisa- tion of the ground state one could still observe a charge-transfer band through the influence of the second term in eqn. (16). Orgel and Mulliken22 attribute the intensity of contact charge-transfer absorption to just such a term the expression “contact charge-transfer absorption” is used to describe the absorption which appears for example in I,- heptane mixtures around 2600 A although there appears to be no stable complex formed.23 It is now necessary to consider if this interpretation of contact charge- transfer absorption is reasonable.The second term in expression (16) can be written in a different form if we again introduce the normalised overlap density By comparing expressions (7) and (17) we see that the reason for expect- ing that had will be small namely that $6d is small compared with will here make the second term in (16) small also. It is not possible to say definitely that the second term in expression (16) will be zero if the first is zero but we can say that factors which make one large or small make the other large or small at the same time. For Orgel and Mulliken’s interpreta- tion of contact charge-transfer absorption to be correct it is necessary that the overlap density lies so far away from A that Wad(A) is zero but not so close to D such that the dipole.moment of [(+d+a)/Sad - +d2] is zero. There is one special case when contact charge-transfer absorption as interpreted above can certainly be observed. If +a and +d have different symmetries as in Fig. 3 then S a d and had will be zero but M a d need not be zero. A charge-transfer band polarised in a plane perpendicular to the direction of charge transfer can be observed without any stabilisation of the ground state. There are other sources of charge-transfer intensity which were not considered in the early work on charge-transfer complexes but have been taken to be important in the field of intramolecular charge-transfer spectra these sources are the excited states of the donor and acceptor. If the donor has a transition +d-++d’ which gives rise to an intense absorp- tion band and if the charge-transfer state $a interacts with the donor excited state to give the effect will be that the charge-transfer band arising from the transition 22 Orgel and Mulliken J.Amer. Chern. Soc. 1957 79,4839. 23 Evans J. Chem. Phys. 1954 23 1436, MURRELL THE THEORY OF CHARGE-TRANSFER SPECTRA 203 &++a’ will now have borrowed some intensity from the donor absorption band. The amount borrowed will be proportional to and to the intensity of the donor band. The coefficient &*a is given by an expression similar to that for &a namely Since $bd* is an excited-state orbital we expect that it will be “blown up” relative to C$d. In general then we expect Sad* to be greater than Sad. Even if the donor and the acceptor do not approach close enough for 4 d and +a to overlap there may still be some overlapping of +d* and 4%.This mechanism therefore provides another interpretation of contact charge-transfer absorption. The borrowing of charge-transfer intensity from excited states of the acceptor can be shown to involve the overlap of the ground-state donor orbital +d and a ground-state orbital of the neutral acceptor ($8 is an orbital of A-). This will probably be unimportant in the case of inter- molecular charge-transfer spectra since exchange repulsion will tend to prevent ground-state orbitals of D and A from overlapping to any large extent. It may however be important for intramolecular charge-transfer absorption as in aniline where the donor and acceptor are held rather closely together by direct covalent bonding.Although we have reason to believe that it is in general the excited states of the donor that contribute the largest part of the charge-transfer intensity there may be particular cases where for reasons of symmetry this is not so. For example if in the quinhydrone complex the quinone and the quinol sit symmetrically one upon the other in parallel planes then the charge-transfer bands which are polarised perpendicular to the aromatic rings cannot by symmetry have picked up any intensity from the low- lying excited states of the quinone or quinol since these bands are polarised in the planes of the aromatic rings. It is probably true to say that in general these rigid symmetrical structures are the exception rather than the rule in solution.In solution we may expect to find a large number of different configurations of the complex in which the least symmetrical although it is the least stable may give rise to the largest charge-transfer intensity. In the series of complexes between iodine and the methylbenzenes we find that as the complex becomes more stable the charge-transfer intensity falls.24 However for the chloranil-methylbenzene complexes the reverse situation arises an increase in stability being accompanied by an increase in intensity of the charge-transfer band (see Table 3).25 Theoretically we expect that in such a series the relative behaviour of stability and charge-transfer intensity should depend on the variation of the difference in energy between the most stable configuration and the configuration giving the largest charge-transfer intensity.These are not 24 Andrews and Keefer J. Amer. G e m . SOC. 1952,74,4500. 25 N. Smith Ph.D. Thesis University of Chicago Ill. 204 QUARTERLY REVIEWS necessarily the same. For instance Orgel and Mulliken interpret the behaviour of the iodine-benzene series as being due to the chance contacts giving rise to the greatest charge-transfer intensity the relative number of TABLE 3. Charge-transfer bands the variation of the extinction coeficient with the equilibrium constant of the complex. Benzene or deriv. Benzene Toluene o-Xylene m-Xylene p-Xylene 1,2,4-Trirnethyl 1,3,5-Trirnethyl 1,2,3,6Tetramethyl 1,2,3 &Tetramethyl 1,2,4,5-Tetramethyl Pentamethyl Hexamethyl Complexes20 with h 292 302 3 16 318 304 332 - - - 332 357 375 I E 16,400 16,700 12,500 12,500 10,100 8850 - - - 9000 9260 8200 K 0.15 0.16 0.27 0.3 1 0.3 1 0.82 - - - 0.63 0-88 1.35 Complexes21 with chloranil x E K 340 2180 0.30 365 1920 0.50 385 2090 1.05 390 2000 0.84 410 1960 0.89 420 1985 1.02 410 2250 1.17 445 2585 2.65 450 2495 2.47 470 2320 3.02 480 2680 5.32 505 2880 9.08 these contacts being reduced as the complex becomes more stable.The chloranil-benzene complexes are more stable and by reason of their geometry it is difficult to obtain much overlap of the donating and the accepting orbitals except when they lie one upon the other in parallel planes. The charge-transfer intensity will therefore arise mainly from the interaction between the ground and the charge-transfer state and we get the expected increase of intensity with an increase of stability.It is interest- ing that the charge-transfer intensity of the chloranil-benzene complexes is much less than for the iodine complexes. This is in agreement with our proposal that much greater intensity can be obtained when the charge- transfer state can borrow from the donor excited states. If our proposal for the benzene-iodine complexes is correct that is that some of the least stable structures contribute most to the charge- transfer intensity then as the temperature is raised the measured extinc- tion coefficient of the charge-transfer band should increase even though the net absorption decreases as less molecules are involved in complex-forma- tion. This prediction has not been tested for the I,-benzene complex but Ross has observed this type of behaviour for the trinitrobenzene-aniline and -naphthalene complexes (reported in ref.22). We have only used one-electron wave functions to describe the electronic states involved in the charge-transfer absorption. If many-electron wave functions are used the matrix elements become more complicated but the most important terms have the form given above.21 We have also based our arguments for the magnitude of the charge-transfer interactions on the MURRELL THE THEORY OF CHARGE-TRANSFER SPECTRA 205 assumption that D and A are both neutral molecules. If one or both is an ion then some of the statements made above will not be applicable. Conclusion.-By my original definition of charge-transfer absorption as one involving a large charge displacement I have excluded the possibility of charge-transfer absorption when the donor and acceptor are identical as in the N-ethylphenazyl dimer26 or biphenyl.27 In fact it has been found convenient to interpret the electronic states of such systems by using as a basic set the wave functions AA A+A- A-A+ etc.However any state of the system must contain an equal contribution from A+A- and A-A+ so there cannot be any actual charge displacement associated with a transi- tion. A state of the type A+A- & A-A+ could perhaps be called a charge- resonance state and a transition from the ground state AA to such a state could be called a charge-resonance transition. The term two-way charge transfer has also been used for this type of molecule and also for example in molecules such as the Ag+-benzene complex where the structures AgC6H,+ and Ag2+CGH6- are probably both important in stabilising the ground state.My definition of a charge-transfer transition also excludes cases when both the ground and the excited state are equal mixtures of DA and D+A- since again there is no net charge transferred from one part of the molecule to the other in the excitation. This situation may occur in the molecules of the type BF,-NH and amongst Brooker’s dyes. However an exactly equal mixture will occur so infrequently that it is not worth making it a special case. In this Review I have restricted my attention to the case in which both D and A have closed-shell electronic structures. Other cases are however important a notable example being the charge-transfer bands associated with complexes between oxygen and aromatic molecule^.^^^^^ It has recently been shown that these charge-transfer states play an important role in the intensification of the singlet-triplet absorption bands of aromatic molecules induced by oxygen and nitric o ~ i d e .~ ~ * ~ O In complexes involving heavy atoms one expects the spin-selection rules to be broken down by spin-orbit coupling. This must certainly occur in the 1,-benzene complex since spin-orbit coupling is known to be important for the iodine molecule (I,) itself. Platt31 has pointed out that the charge-transfer state of this complex lies very close to one of the expected triplet states of benzene and there must be considerable inter- action between the two. Charge-transfer states together with spin-orbit interaction provide a mechanism for the intensification of the singlet 26 Hausser and Murrell f.Chem. Phys. 1957 27 500. 27 Longuet-Higgins and Murrell Proc. Phys. SOC. 1955 68 A 601. 28 Evans f. 1953 345. 2 9 Mulliken and Tsubomura f. Amer. Chenz. SOC. in the press. 30 Murrell Mol. Phys. 1960 3 319. 31 Platt personal communication. 4 206 QUARTERLY REVIEWS triplet absorption bands of aromatic molecules dissolved in methyl iodide.32 In conclusion I must emphasise that in general it will not be possible to separate the electronic states of the system into those which are purely charge-transfer and those which have no charge-transfer character. This will be particularly true for intramolecular absorption. For example it has been calculated14 that the first excited state of aniline has 17% of charge-transfer character and the second excited state 68%. As has been pointed out above the second state has sufficient charge-transfer character to show the expected dependence of its intensity on steric effects the first state has not. In a similar vein the charge-transfer states of Cl-(H20) will resemble highly excited or Rydberg states of C1- so that the distinction between charge-transfer states and excited states of the donor becomes meaningless. 32 Kasha J. Chem. Phys. 1952 20 71.

ISSN:0009-2681    

DOI:10.1039/QR9611500191

出版商:RSC    

年代:1961

数据来源: RSC

摘要:

MECHANISMS OF ELECTRON TRANSFER AND RELATED PROCESSES IN SOLUTION By J. HALPERN* (DEPARTMENT OF THEORETICAL CHEMISTRY UNIVERSITY CHEMICAL LABORATORY LENSFIELD ROAD CAMBRIDGE) THIS Review is concerned with the mechanisms of simple redox reactions in solution notably those involving the transfer of one or more electrons between metal ions or complexes. While many such reactions e.g. Crll + Colll -+ Crlll + Coil . . . . . . (1 1 U I V + 2Felll -+ UVl + 2Fe" . . . . . . . (2) Snll + TI111 3 Snlv + TI1 . . . . . . . (3) are among the most familiar of inorganic and analytical chemistry interest in them until relatively recently appears to have been largely confined to considerations of stoicheiometry and thermodynamics and it is only during the last decade that their kinetics and mechanisms have received serious attention.In addition to the stimulation provided by the general expansion of interest in inorganic chemistry during this period the considerable progress which has been achieved in the study of electron-transfer mechan- isms has been made possible in part by a number of specific experimental developments. Perhaps foremost among these is the introduction into common laboratory practice of a wide range of isotopic tracers. These have found application in detailed mechanistic studies and have opened up the possibility for the investigation not only of ordinary redox reac- tions but also of isotopic exchange reactions involving electron transfer between different oxidation states of the same element,t e.g. Fell + *Fell1 -f Fell1 + *Fell . . . . . . (4) Mn0,2-+ *MnO,- -f Mn0,-+ *Mn0,2- .. . (5) Such reactions have the advantage in many instances of being slower and hence more readily accessible to kinetic measurement than those involving net chemical change and of providing simpler models on which to base theoretical calculations. Their study has therefore received par- ticular attention. The range of electron-transfer reactions whose kinetics *Nuffield Foundation Travelling Fellow; on leave from the Department of Chemistry University of British Columbia Vancouver Canada. ?The earliest such study of the isotopic exchange between Pbll and Pblv was reported by Hevesy and Zechmeister in 1920 (Ber. 1920 53 410). The limited availa- bility of isotopic tracers however precluded extensive work of this type until relatively recently. It is now also possible although as yet only in rather special cases with tech- niques such as electron spin resonance and nuclear magnetic resonance spectroscopy to measure rates of electron-transfer reactions in which there is no chemical change without resorting to isotopic labelling.207 208 QUARTERLY REVIEWS can be studied has also been greatly extended by the development of various techniques for measuring the rates of fast reactions among these are flow techniques,lS2 nuclear magneticre~onance~ and electron spin resonance4 spec- troscopy and relaxation method^,^ which have yielded measurements of rate constants up to lo9 1. mole-1 sec.-l. It is noteworthy however that by no means all the recent progress in this field is due to these newer develop- ments. Along with them and stimulated by them the study of familiar reactions such as (1)-(3) by conventional kinetic methods has played and continues to play an important role.The elucidation of the mechanisms of electron-transfer reactions presents problems of considerably greater difficulty and complexity than are sug- gested by the simplified equations by which we usually represent such reactions. Some of the questions which arise and with which this Review will be concerned relate to the following themes (i) The nature and sequence if more than one of the elementary steps which comprise the overall reaction. Connected with this is the question whether multiequivalent redox processes such as (2) and (3) occur in a single step or through successive 1 -electron steps. (ii) The detailed mechanisms of the elementary steps themselves e.g.the composition and configuration of the activated complex the roles of ligands and solvent etc. (iii) The significance of atomic rearrangements accompanying the electron transfer. In this connexion it should be recognised that all electron- transfer processes in solution involve some atomic rearrangement. In some reactions for example those involving the Mn0,2-Mn0,- or Feaq2+-Feaq3+ couples these are relatively subtle i.e. small differences in the metal-ligand bond lengths and in the polarisation of the surrounding solvent corresponding to the two oxidation states. In others for example reactions involving the U4+-U022+ or PtC1,2-PtC1,2- couple gross changes in the geometry and composition of the co-ordination shell occur. (iv) The significance of the large variations (typical examples of which are to be found in Table 1) in rate and in AH and ASS which are observed in series of related reactions in which the metal ions or ligands are varied.In so far as the two themes can be conveniently separated the experimental evidence relating to the kinetics and mechanisms of electron-transfer reactions will be considered first being followed by a discussion of some of the theoretical aspects of the subject. No attempt has been made at Sheppard and Wahl J. Amer. Chem. SOC. 1957 79 1020; Gjertsen and Wahl 2 Gordon and Wahl ibid. 1958,80,273. 3 McConnell and Weaver J. Chem Phys. 1956,25 307; Bruce Norberg and Weiss- ‘(a) Ward and Weissman J. Amer. Chem. SOC. 1957 79 2086; (b) Adam and 5 Diebler 2. Elektrochem. 1960 64 128. ibid. 1959,81 1572.man ibid. 1956,24,473. Weissman ibid. 1958 80 1518. HALPERN ELECTRON TRANSFER IN SOLUTION 209 complete coverage particularly of the experimental subject matter ; emphasis is placed rather on the discussion of typical systems notably those which have been most extensively investigated and are best under- stood. More detailed treatments of certain aspects of the subject are to be found in earlier reviews by Zwolinski R. J. Marcus and Eyring,6 Amphlett,’ Basolo and Pearson,* T a ~ b e ~ and Stranks.l” Mechanisms of Direct Electron Transfer Although some electron-transfer reactions to be considered in a sub- sequent section occur through indirect mechanisms the kinetic evidence in most cases i.e. a rate-law of the first order in each of the reactants points to a direct reaction between the oxidant and the reductant through an activated complex involving both.Kinetic measurements in such cases serve to define the composition of the activated complex (apart from solvent participation) but not its detailed configuration. However particularly in cases where the reactants contain no solvent ligands or where the reactants and/or products are substitution-inert at least some features of the mechanism and of the structure of the activated complex may be inferred through studies of the type to be described. These have resulted in the recognition of at least two broad classes of mechanism the so-called “outer-sphere” and “inner-sphere” (or bridged) types.* In the first of these electron transfer occurs through the intact co-ordination shells of both metal ions; and in the second through a bridged inter- mediate in which the two metal ions are linked by a bridging ligand common to the co-ordination shells of both.The distinction between the two is not always sharp and many reactions cannot at this stage be assigned with certainty to either class. Outer-sphere Reactions.-In reactions of this type examples of which are listed in Table 1 ( ~ ) electron transfer occurs through an “extended” activated complex in which the first co-ordination shell of each metal ion is presumably intact. The evidence for this type of mechanism is usually (a) a rate-law corresponding to an activated complex containing all the ligands in the first co-ordination shells of both metal ions e.g. k [Co en32+] [Co ens3+] and/or (b) the demonstration that electron *Various designations have been employed to distinguish these two classes of reaction.For the second class the designation “inner-sphere” is preferred to “bridged” because bridging may also occur in outer-sphere mechanisms. The terms “extended” and “compact” have also been used to designate the two types of activated complex. Zwolinski Marcus and Eyring G e m . Rev. 1955,55 157. Amphlett Quart Rev. 1954 8,219. * Basolo and Pearson “Mechanisms of Inorganic Reactions” John Wiley and Sons Inc. New York 1958 Chapter 7. (a) Taube in EmelCus and Sharpe “Advances in Inorganic Chemistry and Radio- chemistry,” Academic Press Inc. New York 1959 p. 1 ; (6) Taube Canad. J. Chern. 1959 37 129. lo Stranks in Lewis and Wilkins “Modern Coordination Chemistry” Interscience Publishers Inc. New York 1960 p.78. 210 QUARTERLY REVIEWS transfer is faster than substitution into the co-ordination shell of either metal ion e.g. the very rapid electron exchange between Fe(CN)G4- and Fe(CN)e3- both of which undergo substitution only slowly. In such a case the fact that the co-ordination shells of the metal ions remain intact during electron transfer may be confirmed by simultaneous isotopic labelling of both the metal ion and the ligands but this is usually considered unnecessary. TABLE 1. Kinetic data 1 -electron-transfer reactionsa Reaction (A) Outer-sphere reactions Co phenP-Co hen,^+ Co en32+-Co ens3+ Co(NH3),2+-Co(NH3),3+ CO(NH),),~+-CO(NH,)~OH~+ Cr2+:ko(NH3)63+ Cr di~y,~+-Co(NH~)63+ Co2+-Co3+ b , -Co(NH3),3fOH- -C0(NH3)G3+C1- 9 -Co(NH3)50H,3+ , -Co(NH3)5C12+ -Co(NH3),Br2+ Mn02-MnO4- I rC1,3-IrClG2 - Fe(CN):-Fe(CN):- Fe ~hen,~+-Fe phen? Ferrocene-ferrocinium (B) Inner-sphere reactions Cr2+-Cr3+ b , -CrOH2fb , -CrF2+ , -cis-CrF,+ , -CrC12+ , -CrBr2+ , -CrNQs+ kbi AH,' AS,' Temp.(1. mole-l (kcal. (e.u.) Ref. sec.-l) mole-l) 3.2 -1 0 1.1 64.5 <lop8 64.5 5-4 x 64-5 0.33 64.5 4-4 x low2 25 5 x 10-5 25 9.0 x 10-5 4 1.0 x 104 4 > 1.6 x 104 4 4 0 3 o >lo5 4 7.1 4 6.5 x lo2 0 7 x lo2 1 > 3 x lo2 - 75 9 x 105 24.5 B 2 x 10-5 0 2.6 x 10-3 0 1.2 x 10-3 0 9 0 >60 0 > 1-2 24.5 0.7 (12.6) (-13) 13.7 -33 12.7 -33 12.2 -35 14.7 -30 13.0 -10 17 + 4 - - - - - - - - - - 10.0 - 9 - - - - - - - - (21) (- 8) - - 13.7 -20 13 -24 - - - - - - 11 12 13 14 14 14 14 156 15a 15a 15a 15a 1 16 17 18 14 19 19 20 21 20 20 20 aData are for aqueous solution (various ionic strengths) except for C~(phen),~+~~+ bThis reaction is of uncertain mechanism but is listed here for purposes of comparison.l1 Bonner and Hunt J. Amer. Chem. SOC. 1952,74 1886. l2 Baker Basolo and Neumann J. Phys. Chem. 1959,63,371. l3 Lewis Coryell and Irvine J. 1949 5386. l4 Stranks Discuss. Faraday SOC. 1960 29 73. l5 Zwickel and Taube (a) Discuss Faraday Soc. 1960 29 73; (b) J. Amer. Chem. l6 Sloth and Garner J. Amer. Chem. SOC. 1955 77 1440. l9 Anderson and Bonner ibid. 1954 76 3826. 2o Ball and King ibid. 1958 80 1091. 21 Chia and King Discuss. Faraday SOC. 1960,29 109. (aqueous acetone) and ferrocene-ferrocinium (methanol). SOC. 1961 83 793; (c) ibid. 1959 81 1288. Deck and Wahl ibid. 1954 76 4054. Eichler and Wahl ibid. 1958 80 4145. HALPERN ELECTRON TRANSFER IN SOLUTION 21 1 Reactton kbi AH AS+ Temp.(1. mole-l (kcal. (em) Ref. sec,-l) mole-l) Cr2+-CrNCS2+ 27 1.8 x 10-4 - 7 9 -CrfNH3)5F2+ 25 2.7 x 13-4 , -CI-(NH~)~C~~+ 25 5-1 x 10-9 1 1 . 1 , -Cr(NH3),Br2+ 25 0.32 8 -5 , -CO(NH~)~OH~~+ 20 0.5 2-9 , -Co(NH3)5OH2+ 20 1.5 x lo6 4.6 - , -Cr(NH3),12+ 25 5.5 , -Co(NH&C12+ 20 >lo3 - , -Co(NH3)5OAc2+ 25.1 0.18 - , -Co(NH,),(H-~uccinate)~+ 14.1 0.17 - , -Co(NH,),(H-f~marate)~+ 14.3 0.75 7.5 , -C~(NH,),(H-terephthalate)~+ 16.6 36 - , -C~(NH,),(Me-fumarate)~+ 2.5 0.43 - , -C~(NH,),(H-phthalate)~+ 14.2 0.055 5-1 , -C~(NH,),(H-isophthalate)~+ 14.2 0.1 1 2-6 (c) Reactions of uncertain mechanism Fe2+-Fe3+ 0 , -FeOH2+ 0 , -FeF2+ 0 , -FeF2+ 0 , -FeCI2+ 0 , -FeBr2+ 0 , -FeSCN2+ 0 , -FeN32+ 0 , -FeC,O,+ 0 , -Fe phen:+ 25 V2+ -CO(NH3)63+ 25 , -CO(NH&OH~~+ 25 , -Co(NH3)5Cl2+ 25 V2f -V3+ 25 Pu3+-Pu4+ 25 -PuOH3+ 25 V3 t-Fe3+ 25 Nb02+-NpOz+ 4.5 0.87 9.7 2.5 9.7 4.9 12.2 7 x lo2 1 x 103 1.8 x 103 3.7 x 104 3.7 x 10-3 -0.5 5.7 1.0 x 20 74 2.0 x 104 5 x 10-3 9.4 6'9 8-6 9.0 8.3 8.0 7.4 13.2 8.6 - 0.2 9-1 - 12.6 7.2 2.2 11.0 16.7 - - 30 - 23 - 33 - 52 - 18 - - - 33 - 47 - 56 - 25 - 18 -21 - 22 - 24 25 - 27 +6 - 14 - 37 - 40 - - 25 -31 - 32 - 12 - 15 20 22 22 22 22 15c 15c 23 24 24 24 24 24 24 24 25 25 26 26 25 27 29 30 27 28 15b 15b 15b 31 32 32 33 34 22 Ogard and Taube J.Amer. Chem. SOC. 1958 80 1084. 23 Taube Myers and Rich ibid. 1953 75 4118; Taube and Myers ibid. 1954 76 24 Sebera and Taube ibid. 1961 83 1785. 25 Silverman and Dodson J. Phys. Chem. 1952,56 846. 26 Hudis and Wahl J . Amer. Chem. SOC 1953,75,4153. 27 Home Microfilm Diss.Abstr. 1957 17 1673; J. Phys. Chem. 1960 64 1512. 28 Sutin and Gordon J. Amer. Chem. SOC. 1961 83 70. 29 Laurence Trans. Faraday Soc. 1957 53 1326. 30 Bunn Dainton and Duckworth ibid. 1959 55 1267. 31 Krishnamurty and Wahl J. Amer. Chem. SOC. 1958,80 5921. 32 Keenan J. Phys. Chem. 1957,61 11 17. 33 Cohen Sullivan and Hindman J. Amer. Chem. SOC. 1954,76 352; 1957,79 3672. 34 Higginson Rosseinsky Stead and Sykes Discuss. Faraduy SOC. 1960,29,49. 2103. 212 QUARTERLY REVIEWS The first of these criteria is applicable both to inert and to labile com- plexes. The second is applicable only to reactions in which both metal complexes are substitution-inert but since it does not depend on a knowledge of the rate-law it can be employed even when the rate is too fast to be measured.Unfortunately neither criterion is readily applicable to most reactions of aquo-ions which thus can rarely be proved to be of this type. The reactions in this class [Table 1 ( ~ ) ] cover a wide range of rates fromk < lo-* 1. mole-l sec.-l for the CO(NH,),~+-CO(NH,)~~+ exchange to k > lo5 for many reactions of the type Fephen,2+-Fephen,3+.1*~10 In general metal ions surrounded by unsaturated or large polarisable ligands such as o-phenthroline bipyridyl cyanide or chloride exchange electrons rapidly usually much faster than the corresponding aquo-ions or ammine complexes. The significance of this and of some of the other trends which appear in Table 1 will be considered below. The demonstration of an outer-sphere mechanism still leaves open the question of how closely the co-ordination shells of the two ions approach each other in the activated complex particularly with reference to the possibility of intervening solvent or electrolyte layers.Only in a few instances is information about this available. Thus from the observation that the OH-catalysed isotopic exchange between CO(NH,),~+ and Co(NH,),,+ (corresponding to the rate law k [CO(NH,),~+] [CO(NH,),~+-OH-]) is accompanied by OH- substitution on the newly created ColI1 ion Stranks14 concludes that electron transfer occurs predominantly through the intermediate [ (NH 3) ,CO~~I*NH,.OHCO~~( NH ,),I rather than through [(NH,) 5C~111*N H,*OH*NH3CoI1( NH 3)n-1]. It seems reasonable that the Cl-catalysed exchange proceeds through an analogous intermediate [(NH,) ,ColI~.NH,CICo~~( N H3)?&] but in this case the rapid hydrolysis of the Co1I1 chloroammine precludes demonstra- tion of net C1- transfer.The markedly C1-catalysed oxidation of Cr2+ by Co(NH,),,+ which yields CrC12+ presumably proceeds through a similar intermediate.35 The rates of certain outer-sphere reactions exhibit marked sensitivity to ions of opposite sign e.g. the Mn042-Mn04- exchange' which is accelerated by cations in the order Cs+>K+ Na+>Li+. This may be due simply to salt effects or alternatively to the operation of electron transfer paths involving outer-sphere bridged intermediates such as [O,MnO-Cs-OMnO,] 2-. Closely related to the reactions of metal complexes in this class in that they involve electron transfer between extensively delocalised orbitals with little accompanying structural rearrangement are certain electron- 85 Taube Chem.Soc. Special Publ. No. 13,1959 p. 57. HALPERN ELECTRON TRANSFER IN SOLUTION 213 transfer reactions between organic molecules and anions (e.g. naphthalene and its mononegative ion) whose rates have been measured by electron spin resonance spectroscopy.* For the exchange between naphthalene and its anion in tetrahydrofuran a cation-bridged intermediate having a sandwich structure with a solvated cation (e.g. Na+) lying between the two hydrocarbon systems has been proposed.36 Since the anion and the cation are strongly associated in this medium it is to be expected that electron transfer will be accompanied by transfer of the cation to the newly formed anion and will thus appear as a net atom transfer. Electron spin resonance measurements yield direct evidence for this in the case of electron transfer to benzophenone from its sodium k e t ~ l .~ ~ Inner- sphere Reactions.-In reactions of this class which have been studied particularly by Taube and his co-workers electron transfer is preceded by substitution into the co-ordination shell of one of the ions with the formation of a bridged intermediate in which the two metal ions are linked by a common ligand. A typical reaction of this type is where X may be any of a large number of molecules or anions e.g. H20 OH- Cl- OAc- etc. [Table 1(~)]. In each case it is found that X appears in the co-ordination shell of the newly formed CrIII ion. (In the cases of H20 and OH- oxygen transfer has been demonstrated by 1 * 0 labelling of the oxidant.37) Since both the CoI1l complex and the Crrrr product are substitution-inert this implies that electron transfer occurs through the bridged intermediate (I) which is formed by an initial sub- stitutional step Cra+ + (NH,),ColllX + 5Hf -+ CrlllX + Coil + 5NH4+ .. . . . (6) (NH,),ColllX + Cr(H,0),2+ -+ [(NH3),Co1~~-X-Cr~~(H,O),] + H20 . (7) (9 Similar mechanisms have been demonstrated for electron transfer between Cr2+ and various chromic complexes21,22 of the type (H,O),CrII*X and (NH,),CrIIIX ; also for various Pt1I-PtIV exchange reactions for example,38 Pt en22+-Pt en,C1,2+. The last reaction and its analogues are catalysed by Cl- according to the rate-law k[Pt en22+] [Pt en2C122+] [Cl-1 suggesting that electron transfer occurs through the symmetrical bridged intermediate (11). [*CI-Pt en2-CI Pt en2-C1I3+ + Pt en,*C122+ + Pt en,a+ + CI- .. . . . . (8) The accompanying exchange with isotopically labelled C1- in the solution (which may be used to follow the reaction Pt en2C122+ being substitution- inert) confirms the mechanism. *CI- + Pt enZa+ + Pt en,C1,2+ +. (11) 36 Aten Dieleman and Hoijtink Discuss. Faraduy SOC. 1960,29 182 37 Kruse and Taube J. Amer. Chem. SOC. 1960,82,526. 38 Basolo Morris and Pearson Discuss. Faraday SOC. 1960,29 80. 214 QUARTERLY REVIEWS In all these reactions the bridged intermediate following electron transfer contains one labile bond (e.g. Co"-X or PtIl-X) and is thus decomposed too rapidly to be detected. In one case however the oxidation of Co(CN),,- by Fe(CN),& where both nietal ligand bonds are substitution-inert a product believed to have the bridged structure [(NC),CO~~~.NC-F~~~(CN),]~- has actually been Although inner-sphere bridged mechanisms of this type are probably fairly common they can be demonstrated unequivocally only in certain cases the minimum requirement being that the reactant complex of one of the metal ions and the product complex of the other be substitution- inert.Unfortunately this condition like that for the demonstration of the outer-sphere mechanism is rarely fulfilled for reactions between aquo- ions. The observation of catalysis by anions even when accompanied by their incorporation into the co-ordination shell of one of the products does not necessarily reflect the participation of the anion as a bridging ligand. Thus S042- and P2074- accelerate the oxidation of Cr2+ by (NH,),CoOH~+ and each is incorporated into the product Cr"I complex.40 Similarly during the oxidation of Cr2+ by (NH,),CoC12+ in the presence of P2074- both Cl- and P20,4- are incorporated into the product CrIII complex.The structure of the activated complex in these cases (where X is the bridging and Y the non-bridging anion) presumably is (NH3),Co~1~-X-Cr~1(H,O),Y. Not unexpectedly it is found that the catalytic effect of anions as non-bridging ligands is much smaller than as bridging ligands. Although the possibility of doubly bridged intermediates with structures analogous for example to that of A12C16 has been considered in certain reactions,41 there is no direct evidence for them. Chia and King21 have shown that electron transfer between Cr2+ and cis-CrF,+ proceeds predominantly through the singly bridged intermediate [(H ,O),F-Cr-F-Cr(H20) ,I3+ rather than the symmetrical species Similarly apparently only one oxygen is involved37 in bridging (and is transferred) in the oxidation of Cr2+ by c~~-(NH,),CO(H,O),~+ or cis-en CO(H~O)~~+.Although the transfer of the bridging ligand from the oxidant to reduct- ant constitutes the evidence for the bridged mechanism in the cases cited it is not clear that such a transfer is an essential feature of this mechanism 39 Hain and Wilmarth J. Amer. Chem. SOC. 1961 83 509. 4 0 Taube ibid. 1955 77 4481. 41 Carpenter Ford-Smith Bell and Dodson Discuss. Faraday Soc. 1960,29 92. HALPERN ELECTRON TRANSFER IN SOLUTION 215 since in these instances it follows simply from a consideration of the relative substitution-lability of the two metal-ligand bonds after electron transfer.The oxidation of Cr2+ by IrClO2- is also believed23 to proceed through a Cl-bridged mechanism but in this case the products are IrCl,3- and Cr(H20),3+. The description of inner-sphere mechanisms as atom transferg2 (as distinct from electron transfer) processes may thus be some- what misleading. This point also is emphasised by the C1-bridged PtII-PtIv exchanges (equation 8) which involve the transfer of two electrons and thus are not equivalent simply to the transfer of a C1 atom from PtIv to Pt” (Cl+ transfer seems an even less realistic concept). Possibly a more meaningful distinction between concepts of electron and atom (or group) transfer in the present context could be made in terms of the actual mechanism of the electron-transfer process the former being used to denote “conduction” mechanisms which do not involve actual oxidation or reduction of the bridging group and the latter “chemical” mechanisms in which an electron is transferred by a process of successive oxidation and reduction of the ligand (or vice versa) by the two metal ions.The two mechanisms are at least in principle quite di~tinguishable.~~ Not surprisingly the rates of inner-sphere redox reactions are very sensitive to the nature of the bridging group reflecting the essential role of the latter in the electron-transfer process. Thus the data in Table 1 reveal a 106-fold variation in the rate of oxidation of Cr2+ by CoIII(NH,),X when the bridging ligand X is varied. The high rates for terephthal and fumarate relative to the other carboxylic acids in this series are of special interest and have been interpreted in terms of bridged inter- mediates such as O-Cr(H,O) *+ 1 (N H3)6Co-0 // C-CH = CH-C \ 0 // \OH in which the two metal ions are co-ordinated to different carboxyl groups the electron being transferred between them by “conduction” through the conjugated .rr-electron system.Strong support for this is provided by the observation that in the corresponding oxidation of Cr2+ by (NH,),CO~~~.O.OC-CH :CH-CO-OMe electron transfer is ac- companied by hydrolysis of the ester and incorporation of the methyl alcohol into the co-ordination shell of the CrlI1 product. With hydrogen maleate or methyl maleate (but not the corresponding fumarates) as the bridging ligand and Cr2+ or V2+ as reductant cis-trans-isomerisation and hydrogen exchange with the solvent (D,O) also accompany electron transfer and this has been construed as evidence that an electron passes into the maleate group during the reaction.44 42 Stewart Experientiu 1959 15 401.43 Halpern and Orgel Discuss. Faruduy Soc. 1960,29 92. 44 Fraser Sebera and Taube J. Amer. Chem. SOC. 1959,81,1906; Fraser and Taube ibid. 1959 81 5000 5514; ibid. 1961 in press. 216 QUARTERLY REVIEWS Electron transfer between Cr2+ and CrN32+ which is accompanied by transfer of N3- is also believed to proceed through a polyatomic bridge (Cr-N=N=N-Cr); the much lower rate observed with CrNCS2+ is attributed to the unsymmetrical structure of the corresponding NCS- bridged intermediate in which electron transfer leads to the formation of an unstable CrSCN2+ complex.2o The possibilities afforded by the use of such conjugated bridging groups for elucidating details of the electron- transfer mechanism and for systematic variation of structural and elec- tronic parameters make these systems extremely valuable and their study promises to play an important role in the further development of the subject.Hydrogen Transfer and Bridging Mechanisms (D 2O Isotope Effects).- A possible mechanism of electron transfer between metal aquo-ions first suggested by Dodson and David~on,~~ is through transfer of a hydrogen atom between the hydration shells. For the Fe2+-Fe3+ reaction this path may be depicted as Fe2+0H + HO*Fe3+ -f Fe2+O-H . . .0*Fe3+ -+ FeOH2+ + H,O*Fe3+ Fe3+ + OH- *Fez+ + H,O+ H H [ H - t H 1 5 + It 11 . . (9) or for the Fe2+-FeOH2+ exchange Fe2+0H + HO-*Fe3+ -+ Fe2+0-H .. .-O*Fe3+ -f FeS+OH- + H20*Fe2+ . (10) H [ H - + H 1 4 + the intermediate in the latter case being symmetrical. Among the evidence which has been advanced in support of such a mechanism is the following (1) The activation energies of a surprisingly large number of diverse redox reactions involving metal aquo-ions are close to 10 kcal./mole and their activation entropies close to -25 e.u. suggesting that they proceed by a common mechanism which probably involves water.46 (2) In certain redox reactions involving metal complexes there seems to be a requirement that at least one of the inner shell ligands be a water rnoIecule,*6 e.g. one of the CN- ions must be replaced by a water molecule before Fe(CN) sp is oxidjsed by hydroperoxide. Similarly the electrolytic reduction of Cd(CN)42- proceeds through the aquotricyano-complex.(3) The rates of the Fe2+-Fe3+ and the Fe2+-FeOH2+ reaction are lowered by a factor of 2 in passing from H20 to D20 as While this isotope effect is consistent with a mechanism involving breaking of an 0-H bond the support it provides for it is not compelling in view of uncertainties concerning the differences in solvent characteristics of H20 and D20 and of the D20 isotope effects observed in other redox reactions (Table 2). Thus even larger deuterium isotope effects are found in the 45 Dodson and Davidson J. Phys. Chem. 1952 56 866. 46 Reynolds and Lumry J. Chem. Phys. 1955 23 2460. 47 (a) Hudis and Dodson J. Amer. Chem. Soc. 1956,78,911; (b) Sutin and Dodson personal communication. HALPERN ELECTRON TRANSFER IN SOLUTION 21 7 TABLE 2.D20 Isotope effects in redox reactions. Reaction kH2O/kD2O Ref. Fe2+-Fe3+ 2 47a Fe2+-FeOH2+ 2 47a Fe2+-FeC12+ 2.5 47b Fe2+-FeN32+ 1.5 30 Cr2+-Co(NH3),3+ 1-3 15b Cr2+-Co(NH3)50H23+ 3.8 15c Cr2+-Co(NH3),0H2 + 2-6 1% Cr2+-Co(NH3)5 (H-fumarate) -1 24 Cr(dipy)32+-Co(NHd,3+ 1 -0a 15a Cr(dipy),2+-Co(NH3) 50H 23+ 2.6 15a T~+-TP+ -1-5 48 Co2+-Co3+ 2 1 1 V2+-Co(NH3)? 1-7 156 V2+-Co(NH3)5C12+ 2.2 15b OThere is however a 30% reduction in the rate of this reaction in going from CO(NH~)~~+ to C O ( N D ~ ) ~ ~ + as oxidant;15 this is probably related in origin to some of the D20 isotope effects for aquo-ions. oxidations of Cr2+ by CO(NH~),OH~~+ and Co(NH3),0H2+ which are known3' to proceed through inner-sphere oxygen-bridged mechanisms and in the oxidation of Cr dip^,^+ by Co(NH3),0H2+.The latter reaction presumably proceeds by an outer-sphere mechanism but in this case water is absent from the co-ordination shell of the reductant. An appreciable although somewhat smaller D20 isotope effect is observed even for electron transfer between Cr2+ and Cr(NH3),C12+ which is known to proceed by a C1-bridged mechanism.* On energetic grounds there are some objections to the suggestion of a net hydrogen transfer in a reaction such as (9). The endothermicity of this process is expected to approach that of the self-ionisation of water (-13 kcal./mole) and this is difficult to reconcile with an observed activa- tion energy of 10 kcal./mole. A somewhat modified view10*27 of the role of water in these reactions which probably has greater validity is that coupling of the hydration shells of the two ions by hydrogen bonding lowers the energy of the activated complex and by increasing the overlap between the exchanging orbitals provides a more effective conducting path for electron transfer.In this context transfer of hydrogen (as with other bridging groups) is incidental to its bridging role and whether or not it occurs depends on the relative proton affinities of the two hydration shells after electron transfer. Thus hydrogen transfer would be expected to accompany reaction (10) but not (9). *Other studies of kinetic isotope effects in redox reactions might also be mentioned in this context. Murmann Taube and P a u ~ e y ~ ~ found an appreciable oxygen isotope effect [If(160)/k('*O) = 1.0351 in the inner-sphere oxidation of Cr2+ by Co(NH3),0H,3+ suggesting appreciable weakening of the Co-0 bond in the oxygen-bridged activated complex.This reaction also exhibits a smaller but measurable nitrogen isotope effect [k(14N)/k(15N) = 1.003] indicative of some weakening of the Co-N bonds. N~Oz+-Np02~ + 1.4 33 Cr2+-Co( NH3),C12+ 1.3 22 48 Gilks and Waind Discuss. Faraday SOC. 1960 29 102. 4S Murmann Taube and Pausey J. Amer. G e m . SOC. 1957,79,262. 218 QUARTERLY REVIEWS Reactions of Uncertain Mechanism.-The uncertainties concerning the detailed mechanism of the Fe2+-Fe3+ exchange extend also to the redox reactions of most other aquo-ions and labile complexes. For these reactions the conditions required for unequivocal demonstration of either the outer- or inner-sphere mechanism are generally not fulfilled and at least at this stage their mechanisms must be inferred indirectly.Only very limited progress has been made in this direction and it is in this area that some of the most important and challenging problems connected with the study of electron-transfer mechanisms are at present to be found. Attempts so far to infer the mechanisms of such reactions have invoked some of the following criteria. (1) Comparisons with reactions of known mechanism. Such comparisons involve particularly the study of ligand effects. The special role of the bridging ligand in the inner-sphere mechanism might be expected to give rise to a different pattern of dependence of the rate on the nature of the ligand from that observed for outer-sphere reactions thus providing a diagnostic tool for distinguishing between the two types.Thus the observa- tion that the relative rates of oxidation of V2+ by the series of ColI1 com- plexes Co(NH,),3+ CO(NH~),OH~~+ Co(NH,),0H2+ and Co(NH3),C12+ (Table 3) parallels more closely the relative rates of oxidation of Cr(bi~y),~+ than of Cr2+ has led Zwickel and T a ~ b e l ~ to conclude that the oxida- tion of V2+ in these cases unlike that of Cr2+ involves an outer-sphere mechanism. Similary it would appear although contrary arguments also have been advancedg that the small dependence of the rate of Fe2+-FeX2+ exchange (Table 1) on the identity of the halide X- compared with thelarge variationsfound in the Cr2+-CrX2+ and Cr2+-Co(NH3),X2+reactions argues against an inner-sphere bridging role for the halide in the case of Fe2+-FeX2+ ; also the 103-fold difference between the Fe2+-Fe3+ and Fe2+-FeOH2+ rates is considerably smaller than the corresponding difference between the inner- sphere Cr2+-Co(NH3)50H23+ and Cr2+-Co(NH3),0H2+ reactions but does not appear unreasonable for an outer-sphere hydrogen-bridged mechanism.In this connexion it is of interest that in the outer-sphere oxidation of Cr dipy,2+ where owing to the absence of water from the co-ordination shell of the reductant hydrogen bridging of this type is not possible the rate for Co(NH3),0H2+ as oxidant is actually slightly lower than for CO(NH~),OH~~+.~~ On the other hand the large effect of N3- on the Fe2+-Fe3+ exchange and the somewhat different AH$ and AS$ values for the Fe2+-FeN32+ reaction are suggestive50 of a different mechanism in this case possibly an inner-sphere mechanism similar to that for Cr2+-CrN32+.20 Unfortunately the patterns of ligand effects for the few reactions of known mechanism so far studied are not sufficiently distinctive or well understood for conclusions such as these to be drawn with great confidence.The accumulation of more results of this type for reactions of 5 0 Dainton Discuss. Furuduy Soc. 1960,29 125. HALPERN ELECTRON TRANSFER IN SOLUTION 21 9 TABLE 3. Relative* rates of oxidation of V" and Cr" by various Coili complexes (after Zwickel and T a ~ b e l ~ ~ ) . Oxidant Reductant Cr2+aq V2+aq Cr dipyz+ Co(NH,):+ 1 1 1 Co( NH3)5C12+ 105 1.6 x 103 1.5 x 103 CO(NH~),OH~~+ 6 x lo3 135 91 CO(NH,),OH~~ 2 x 1O1O 107 50 *Relative to Co(NH,),,+ in each case. known mechanism however will undoubtedly enhance the value of this approach.In a few cases the mechanistic inferences to be drawn from the study of ligand effects seem more convincing. For example the hydrolysis of the bridging ester which accompanies the inner-sphere oxidation of Cr2+ (but not the outer-sphere oxidation of Cr dip^,^+) by (NH,),Co*II(Me fumarate) is also observed with V2+ as reductant strongly implying an inner-sphere mechanism also for the latter r e a ~ t i o n . ~ ~ ' ~ ~ (2) AS and related kinetic parameters. The measurement of kinetic parameters particularly of AS$ which is related to the structure of the activated complex might also be expected to provide information about the mechanism of electron-transfer reactions. Although some attempts in this direction have been made,51 the relatively few results now available (Table 1) for reactions of known mechanism do not appear to yield sufficiently distinctive patterns for the mechanism of say the Fe2+-Fe3+ exchange to be assigned with confidence on this basis.The potential usefulness of the method thus rests on the accumulation of more extensive and more accurate data and on the achievement of a better understanding of the influence on ASS of other specific factors notably the charges of the reacting ions the ionic strength and (in the case of reactions involving a net chemical change) of the overall entropy of reaction. The dependence of AS$ on the charges of the reacting ions has been discussed by several author^.^^,^^ Newton and R a b i d e a ~ ~ ~ showed that for a large number of redox reactions of the actinide elements (including simple electron transfers and reactions involving hydrolytic changes) the entropy of the activated complex S$ (given by S = AS$ + CSoreactant) is largely a function of its charge.For activated complexes of charge +3 + 4 + 5 and + 6 values of S$ (presumably reflecting largely the hydra- tion of the activated complex) were found to lie in the ranges 29-40 67-81 72-106 and 102-128 e.u. respectively. Higginson et aZ.,34 on the other hand have drawn attention to a correlation (shown graphic- ally in Fig. 1) between -d S and the charge of the activated complex for a number of isotopic exchange reactions and to the fact that many electron- transfer reactions involving a net chemical change exhibit deviations from 51 Higginson Discuss. Faraday SOC. 1960 29 123. 52 Newton and Rabideau J . Phys. Chem. 1959,63 365.220 QUARTERLY REVIEWS w 2-,‘ - ’i *i+’2+ 0 3 4 5 6 7 Charge on activated complex FIG. 1. Dependence of ASS of isotopic exchange reactions on the charge of the activated this correlation in the direction of the entropy of reaction AS” (e.g. TI+-CO~+ ASS = + 5 AS” = + 30; Ag+-Ce4+ ASS = + 2 AS” = + 20; Fe2+-T13+ ASS = - 5 AS” = + 10). When fitted to a relation of the type (where ASf.(, is the entropy of activation based on the correla- tion in Fig. 1 expected for a reaction with AS” = 0) complex (based on the data cited by Higginson et aLS4). (ASS - ASS(0)) == 01 AS” . . . . . . . . . (1 1) values of 01 for these reactions were found to range from 0.5 to 1 i.e. the additional contribution to the entropy of activation corresponds to a large fraction of overall entropy of reaction.Since the entropy of reaction arises largely from changes in the hydration of the ions this implies that the configuration of the activated complex approaches rather closely to that of the final product; this is more readily reconciled with an “extended” (i.e. outer-sphere) than with a “compact” structure for the activated complex for the reactions in question. Such interpretations are clearly subject to considerable uncertainty but serve to illustrate the possibilities of this approach. For electron-transfer reactions where there are accompanying hydrolytic changes e.g. Co3+ + V3+ + H20-+Co2+ + V02+ + 2H+ analogous entropy correlations suggest that at least partial hydrolysis occurs in the activated complex.34 Taubeg has drawn attention to differences in ASS between certain bridged electron-transfer reactions and related non-redox reactions which HALPERN ELECTRON TRANSFER IN SOLUTION 22 1 proceed through structurally similar activated complexes e.g.(NH3),CrBr2+ + Cr2+ ASS = - 33 e.u.; (NH3),CrBr2+ + Hg2+ dS$ = - 16 e.u. The difference was attributed to more stringent requirements in the way of simultaneous bond rearrangement in the case of the electron- transfer reaction. A relation between AH and the heat of reaction AH" paralleling that noted above for entropies was found by Newton and R a b i d e a ~ ~ ~ to hold for a large number of redox reactions of the actinide elements. Within this correlation A HS for purely "electron-transfer" reactions (e.g. Pu3+ + PuOz+ -+ Pu4+ + PuO,+) appeared to be somewhat lower than for reactions involving hydrolytic or structural changes (e.g.Np4+ + NpO,,+ + 2H20+2Np0,+ + 4H+) presumably reflecting an additional con- tribution to AH from bond rearrangement in the latter cases. Another kinetic parameter which is potentially capable of providing information about the structure of the activated complex and helping to distinguish between inner- and outer-sphere mechanisms is the volume of activation A V$ determined from the pressure-dependence of the rate. Measurements of this type have not as yet been reported on electron- transfer reactions but a recent study of the mechanism of a substitution reaction53 serves to illustrate their possible value. (3) Solvent eflects. One of the chief uncertainties in the mechanisms of electron-transfer reactions relates to the participation of the solvent (e.g.in the case of aquo-ions to the number cf co-ordinated water molecules in the activated complex). One approach to gaining more in- formation about this is through the study of solvent effects. The study of . redox reactions in D,O has been vigorously pursued from this standpoint but in relatively few instances has the study of redox reactions between metal ions been extended to non-aqueous solvents. The Fe"-FeI*' ex- change has been examined for reaction in n i t r ~ m e t h a n e ~ ~ ~ and in various alcohols28 and found in all these cases to be much slower than in water; in alcohol the reaction is markedly accelerated according to a first-order rate law by small amounts of water supporting the view that the latter plays a specific role in the mechanism. Other redox studies in non-aqueous media include those on the CeI1*-PbIv and Co1I-PbIV reactions in acetic and the UIV-Uvr exchange reaction in methanol-water mixed The latter reaction exhibits marked changes in kinetics and apparently in mechanism with solvent composition.In all these cases interpretation of the results and inferences concerning details of mechanism are severely limited by an insufficient understanding of the properties of ionic species in non-aqueous solvents. Considerably more experience with systems of this type will be required before they can be interpreted with confidence. 53 Hunt and Taube J. Amer. Chem. SOC. 1958,80 2642. 54 (a) Maddock Trans Furuduy SOC. 1959,55 1268; (6) Sutin J. Phys. Chem. 1960 65 Benson and Sutcliffe Trans. Furuday SOC. 1960 56 246; Benson Proll Sutcliffe j 6 Mathews Hefley and Amis J .Phys. Chem. 1959,63 1236. 64 7766. and Walkley Discuss. Faruduy SOC. 1960 29 60. 222 QUARTERLY REVIEWS Two-equivalent Redox Reactions The transition metals normally exhibit stable oxidation states differing by one electron and react with each other by 1-equivalent steps. On the other hand the stable oxidation states of the post-transition elements usually differ by two electrons (e.g. Snll-Snlv; TP-TlI*I; Hg22+ -2Hg9. The question thus arises whether reactions involving these redox couples [e.g. (2) and (3)] occur in a single step or by successive I-equivalent steps. In considering this it is convenient to discuss first non-complementary reactions (2-equivalent oxidant + 1 -equivalent reductant or vice versa) and then complementary ones (2-equivalent oxidant + 2-equivalent reductant).Non-complementary reactions.-The simple mechanisms which might be expected to operate in non-complementary redox reactions a typical one involving say the oxidation of A+ to A2+ and the reduction of B2+ to B are of four types 1. One-step termolecular mechanism. II. 2A+ + B2+ + 2A2+ + B A+ + B2+ + A2+ + Bf A+ + Bf + A2+ + B Bimolecular mechanism initial 2-equivalent step. A+ + B2+ + A3+ + B A+ f A3+ -f 2A2+ 2A+ + A + A2+ A + B2+ -+ A2+ + B Bimolecular mechanism initial I-equivalent step. 111. IV. Bimolecular mechanism initial d is proportionation. In each of the last three cases the earlier step which involves the formation of an unstable species would presumably be the slower. Shaffer's principle of equi-valency ~ h a n g e ~ ' * ~ ~ which notes that non- complementary reactions are often slower than similar complementary ones (e.g.the slow reduction of TlIII by Fe*I or of CeIv by TP compared with the rapid reduction of TP1 by SnI1 and of Ce Iv by Fe") finds rational- isation in terms of this picture. Thus the first of the above mechanisms is expected to be slow because it involves a termolecular step and the other three because they involve the formation of unstable intermediates (B+ A3+ and A respectively). Not unexpectedly however because of the other specific factors which influence the rates of redox reactions and because of the possibilities for stabilisation of intermediate oxidation states the equi-valency principle has only limited validity and many apparent excep- tions are to be 57 Shaffer J . Amer.Chem. SOC. 1933 55 2169; J. Phys. Cherzi. 1936 40 1021; 58 Halpern Cunad. J. Chem. 1959 37 148 Cold Spring Harbor Symp. Quant. Biol. 1939 1 50. HALPERN ELECTRON TRANSFER IN SOLUTION 223 There is evidence that at least the first three of the above mechanisms occur. The termolecular mechanism (type I) corresponding to the rate- law k[A+I2 [B2+] is readily distinguished kinetically from the others and several cases involving the oxidation or reduction of metal ions by oxygen and hydrogen respectively (e.g. 2 FeII + 02;59 2Pu111 + 02;60 2Ag+ + HZ6l) have been observed. However with the possible exception of one of the paths in the oxidation of CoII by PbIv in acetic there is no known instance of a reaction involving three metal ions which proceeds by this mechanism. In the limiting case both type I1 and type 111 mechanisms yield bimole- cular rate-laws k [A+] [B2+] and a kinetic distinction between them becomes possible only in the favourable case where reversal of the first step is fast enough to compete with the second one for then inhibition by one or other of the products is observed.An example where an initial 1-electron step has been demonstrated62 is the reaction 2Fe" + T P -+ 2FeIII + TI1 which is inhibited by FelI1 according to the rate-law k1k2 [FeI1l2 [TllI1]/ (k-l[Felll] + k2[Fe11]) which fits the mechanism kl k-1 Fell + TI111 + Fell1 + TI11 . . . . . . . (1 2) Similar mechanisms have been demonstrated for 2Co"I + TlI,63 and for 2V1v + TIIII,~~ and are considered likely for many other reactions on the basis of the relative plausibilities of the two alternative intermediates ; for example U" is considered more likely than FeI as an intermediate in the reaction U'v + 2Fe"I + Uvr + 2FeII.64 On the other hand inhibition by Pb" suggests that one of the paths in the oxidation of CoT1 by PbIV in acetic acid involves an initial 2-equivalent step with the formation of Co" as intermediate.55 Mechanisms of type IV are rare.They might be expected to operate in cases where the species A and A3+ are both known to have considerable stability e.g. in the reaction 2VIv + TPII+2VV + TP. The kinetics of this reaction,34 however point to a type I1 mechanism with at most minor contributions from the alternative path involving disproportionation of V I V . One instance of electron transfer which apparently proceeds by a disproportionation mechanism is the Agl-Agrl isotopic exchange.65 The observed rate-law in this case k [AgIII2 is consistent with the disproportion- ation 2Ag" + Agl + Ag"' as the principal exchange path and leads to the surprising conclusion that the rate constant for this step (1 .O x 1031.mole-' 5 9 George J. 1954 4349. 6o Baker and Newton J. Phys. Chem. 1957,61 381. 61 Webster and Halpern ibid. 1957 61 1239. 62 Ashurst and Higginson J. 1953,3044. 63 Ashurst and Higginson ibid. 1956 343 Betts Canad. J. Chem. 1955 33 1780. 65 Gordon and Wahl J. Amer. Chem. SOC. 1958,80,273. 224 QUARTERLY REVIEWS sec.-l at 0") is at least 100 times that for direct electron transfer between Agl and Ag". It is also appropriate to mention in this context the reaction Hgz2+ + T1"I + 2Hg" + TP whose kinetics k[Hg22+][Tl*11]/[HgII] imply the mechanism HgO + TI111 -+ Hgll + TI1 The rate constant for the latter step (10".mole-1 sec.-l at 25 ") is at least lo6 times that for direct oxidation of Hg,2+.66 The very rapid isotopic exchange between Hgz+ and Hg2+ is also believed6' to proceed through the dismuta- tion step (14). In some cases notably where substitution-inert species are involved chemical evidence may serve as a guide to mechanism. Thus the observa- tion68 that CrlI is oxidised to Craq3+ by 1-equivalent oxidants such as CulI and FeIII but to a binuclear species (probably Cr-0-Cr4+) by 2-equivalent oxidants such as H20 and T P suggests that the latter reaction proceeds through an initial 2-electron transfer Crll + TI111 -+ Crlv + TI1 Crlv + Crll -+ (Crlll) or alternatively through a modification of the termolecular mechanism involving a binuclear Cr" species .. . . . . . Hg,,+ + HgO + Hgll (1 4) (1 5) . . . . . . . . . . . . . . . (1 6 ) (1 7) . . . . . . . . . . . . . . . . . . 2Crll -+ (Crll) (1 8) (Crll) + TI111 -+ (Cr1ll) + TI' ' (19) . . . . . . In this connexion reference should also be made to the oxidations of N,H4 and S032-,69 which tend to exhibit different stoicheiometries with 1- and 2-equivalent oxidents i.e. - e N2H3 -f i N 4 H 6 -f iN + NH3 -2e NZH4 - Ij L N,H -+ N - 2e - e SO,- -+ &S,062- -[I SO,,- -+ SO (SO,2-) These systems have proved useful in assessing the relative 1- and 2-equi- valent oxidising tendencies of different oxidants. 66 Armstrong Halpern and Higginson J. Phys. Chem. 1956 60 1661; Armstrong and Halpern Canad. J . Chem. 1957,35 1020.67 Wolfgang and Dodson J. Phys. Chem. 1952 56 872. 68 Arden and Plane J. Amer. Chem. Soc. 1959 81 3197. 6 9 (a) Kirk and Brown ibid. 1928,50 337; Higginson Sutton and Wright J. 1953 1380 1402; J. 1955 1551 ; (6) Higginson and Marshall J. 1957 447. HALPERN ELECTRON TRANSFER IN SOLUTION 225 Complementary Reactions (2Equiv. Oxidant + %Equiv. Reductant).- The reactions in this class of which the Tlr-T1lrl e~change~~s~l and oxida- tion of UIV by T P (ref. 70) are typical commonly exhibit simple bi- molecular rate-laws e.g. k[U’V] [TlIII] which are consistent with either a single 2-equivalent step k Ulv + TI111 -+ Uvl + TI1 . . . . . . . (20) or with a sequence of 1 -electron steps ; k . . . . . . UlV 4- TI111 -+ UV + TI11 * (21) Uv + TI11 -f Uvl + TI1 . . . . . . .(22) Fast In the limiting case of the second mechanism where the two intermediates react with each other before they can diffuse out of the solvent cage in which they are formed a direct distinction between the two mechanisms becomes very difficult and of questionable rneaning.’l In more favourable cases where the intermediates do escape into the surrounding solution their detection (e.g. through some competing reaction) and thus confirma- tion of a mechanism of the second type becomes at least in principle possible. One of the implications of the comparison on which the principle of equi-valency change is based is that reactions between 2-equivalent oxidants and 2-equivalent reductants occur by a concerted 2-equivalent step for otherwise they would be expected to be on the whole even slower than non-complementary reactions.This and the absence of evidence for the existence of intermediate oxidation states provide some indirect support for the view that many reactions such as the TlLTl1II exchange do proceed by 2-equivalent mechanisms. The entropy correlations discussed by Higginson el aZ.34 also favour this conclusion ASS for the TI+-TP+ exchange (-J - 20 e.u.) being more negative than that expected for the rate-determining step T1+ + T13+ -+ 2T12+ for which AS” is probably quite positive. None of this evidence however is conclusive and there remains considerable uncertainty about the detailed mechanisms of these reac- tions. On the other hand numerous examples are now known42 of reactions in which a 2-equivalent redox change is accomplished apparently in a single step through the transfer of a hydride ion (e.g.R,CH.O- + Mn0,- + R2C0 + HMn042-) or of an oxygen atom (e.g. NO2- + OCl- + NO,- + C1-). For such reactions at least if not for reactions of the electron- transfer type it is now generally recognised that the Michaelis’ principle of “compulsory univalent oxidation” is without universal validity. 70 Harkness and Halpern J. Amer Chem. Soc. 1959,81 1526. ‘l Westheimer in McElroy and Glass “Mechanism of Enzyme ‘Action,” Johns 78 Michaelis Trans. Electrochem. Soc. 1958 80 1073 ; Cold Spring Harbor Symp. Hopkins Press Baltimore 1954 p. 321. Quant. Biol. 1939 1 33. 226 QUARTERLY REVIEWS In addition to the simple mechanisms considered above an inherent possibility for reactions between 2-equivalent oxidants and reductants is a chain mechanism initiated by an initial 1 -electron transfer and propagated by the two intermediate oxidation states.An example of this is the oxida- tion of U I V by oxygen for which the mechanism (23-26) has been ad- vanced :73 (23) (24) (25) Termination UV + HO,+ UVl + H,O . . . . . . * (26) Initiation UlV + 0 -f UV + HO . . . . . . . . Propagation Uv+O,+Uvl+HO . . . . . . . UlV + HO -+ UV + H,O . . . . . . . Similar mechanisms apparently operate in the oxidation of other 2- equivalent reductants (e.g. by oxygen but not in any known instance of a reaction such as (20) between two metal ions. A factor which may favour this type of mechanism over a direct 2-equivalent reaction in the case of oxygen is the change in spin multiplicity (triplet+singlet) which accompanies the reduction of oxygen to hydrogen peroxide.Multi-equivalent reactions Redox reactions in which there is a net transfer of more than two electrons (e.g. those involving the MnI1-MnV1I or CrlI1-Crvl couples) almost certainly proceed by mu1 ti-step mechanisms. An example of such a reaction is Crlll + 3Celv -t Crvl + 3Celll . . . . . . . (27) for which Tong and King74 found the rate-law k[CrlI1] [Ce1v]2[Ce111]-l suggesting the mechanism Celv + Crlll + Cell1 + Crlv (rapid equil.) . . . . . (27a) Celv + Crlv -f Cell1 + Crv (rate-determining) . . . . (27b) Celv + Crv +- Cell1 + Crvl (rapid) . . . . . . (27c) Some Indirect Redox Mechanisms In some cases redox processes proceed through mechanisms which do not involve direct reaction (i.e. direct transfer of electrons or of oxidizing or reducing groups) between the oxidant and the reductant.Such indirect mechanisms are frequently responsible for catalytic effects in redox systems. 73 Halpern and Smith Cunud. J. Chem. 1956,34 1419. 7 4 Tong and King J. Amer. Chem. SOC. 1960 82 3805. HALPERN ELECTRON TRANSFER IN SOLUTION 227 Electron release to solvent. One possible alternative to direct electron transfer is the release of an electron by the reductant to the solvent and its subsequent capture by the oxidant e.g. . . . . . . . . Cr2+ + Cr3+ + e- (284 (28b) . . . . . . . . . . Fe3+ + e- -f Fez+ This might be expected to occur with the very powerful reducing agents particularly in solvents such as ammonia which give rise to stable solutions of “electrons”. In aqueous or alcoholic solution electron release would probably result in reduction of the solvent with the formation of a hydrogen atom that could subsequently reduce the oxidant.The reduction of metal ions such as Ag+ and Fe3+ in aqueous solution by hydrogen atoms (generated radiolytically or photochemically or introduced from the gas phase) is indeed well There is no experimental evidence however that either free electrons or hydrogen atoms are intermediates in redox reactions between metal ions even with ions as strongly reducing as Cr2+. Intermediate oxidation or reduction of ligands. Oxidisable or re- ducible ligands can act in effect as “electron carriers” between metal ions. An example of this is the Br-catalysed Tll-TPII isotopic exchange where kinetic contributions of the forms k[T1Br2+] and k[T1Br3] (in each case zero-order in TP) have been identified41 with the exchange paths TIBr,+ + TI+ + Br .. . . . . . . (29) TIBr +TI+ + Br + Br- (30) . . . . . Corresponding 1 -electron oxidation or reduction of ligands to yield atoms or free radicals (e.g. Br Cl or OH) is usually less favourable energetically and such intermediates are not commonly encountered in thermal redox reactions between metal ions. One possible such case is the oxidation of TP by CeIv where hydroxyl formed by the reaction CeOH3+ + Ce3+ + OH is suggested as an intermediate.76*58 Intermediate oxidation or reduction of the bridging ligand (in this case without release from the bridged complex) also constitutes a possible mechanism of electron transfer in inner-sphere bridged c~mplexes,~~ e.g. (Cr3+-X2-Co3+) L (Cr3+-X-Co2+) 7 I Catalysis by metal ions.Ions such as Cu2+ and Ag+ of metals which exhibit two or more stable oxidation states may also serve to “transport” 75 Collinson Dainton Smith Trudel and Tazuke Discuss. Faraday SOC. 1960 29 188 ; Halpern ibid. 1960 29 252. v6 Armstrong and Halpern unpublished work; Gryder and Dorfman J. Amer. Chem. SOC. 1961 83 1254. (C r2+-X-Co3+) 7 (C r2+-X-Co2+) 228 QUARTERLY REVIEWS electrons in redox reactions through a chain mechanism in which the catalytic metal ion is successively oxidised and reduced. Such mechanisms are fairly common some examples being (a) Catalysis by CulI of the reaction,34 VIIr + FeIrI -+ VIV + Fe". The rate-law for the catalysed path k[v"~] [CUII] suggests the mechanism k Vlll + Cull + VlV + Cul Fast Cul + Fell1 -+ Fell + Cull . . . . . . . . (31 ) (32) .. . . . . . . (b) Catalysis by Agl of the reaction,34 TlI + 2Ce1V -+ TP" + 2Ce"I. The rate-law klk2 [Cexv] [TP] [Agl]/(k- [CeIII]' + k2[Tl1]) is consistent with the mechanism (c) Catalysis by CuIr of the oxidation of UrV by oxygen (reaction 23),73 probably through catalysis of the initiation step (23) by the mechanism UlV + Cull -f UV + Cul . . . . . . . . (36) . . . . . Cul + 0 +CuIl + 02-(H0,) (37) Other examples where similar mechanisms probably operate are FelI + 0,; VlII + O2 (both catalysed by Cu9;59,77 Mn042-Mn0,- [catalysed by Fe(CN),3-].65 In all these cases the sequence of electron transfers from reductant to catalyst and from catalyst to oxidant is apparently more efficient than direct electron transfer. The reasons for this and for the particularly widespread effectiveness of the c u 1 - C ~ ~ ~ couple in catalytic mechanisms of this type are not altogether clear.Theoretical Considerations The theory of electron-transfer reactions in solution has been considered by various authors. Libby78 attempted to account for some of the observed rate differences in terms of the Franck-Condon principle ; a quantitative treatment incorporating some of these ideas was subsequently developed by R. A. R. J. Marcus Zwolinski and Eyring,80 Weiss,81 and 77 Ramsey Sugimoto and De Vorkin J. Amer. Chem. SOC. 1941,63,3480. 7a Libby J . Phys. Chem. 1952 56 863. 7s Marcus J. Chem. Phys. (a) 1956 24 966; (b) 1957 26 867 872; (c) Discuss. Furuduy SOC. 1960 29 21. Marcus Zwolinski and Eyring J. Phys. Chem. 1954 58 432. Weiss Proc. Roy. SOC. 1954 A 222 128. HALPERN ELECTRON TRANSFER IN SOLUTION 229 Laidle?2 have given quantitative treatments of electron transfer based on electron-tunnelling models.OrgeP3 has discussed some aspects of the problem from the standpoint of the ligand-field theory. George and G~-iffith,~~ and Halpern and Orge1,43 have examined the detailed mechanism of electron transfer between metal ions with particular reference to the role of bridging ligands. Also relevant to the subject are various theoretical considerations relating to the somewhat simpler and better understood electron-transfer processes which occur in the gas p h a ~ e . ~ ~ ~ ~ ~ The following discussion of the subject while not complete summarises many features of the above treatments. The general approach follows in many respects that of Marcus.79 The treatment of even the simplest type of electron-transfer process in solution is considerably more complicated than that of the corresponding process in the gas phase in that account must be taken in a rather detailed way not only of the interactions of the reactants with each other but also of those with the surrounding medium.The distinction between reactant and medium in this sense is sometimes rather arbitrary. For example in the reaction the “reactants” may be defined so as to include the inner hydration shells or alternatively the latter may be regarded as part of the surrounding medium. Whatever the terminology employed however a realistic model must include as part of the reacting system not only the inner co-ordina- tion shells of the ions but also that region of the surrounding solvent with which they interact i.e.that which they polarise or over which the trans- ferring electron is significantly delocalised. Thus even in the simplest cases we are dealing effectively with reactions of rather large and complex “molecules”. The fact that the configuration of the transition state including the compositions of the inner co-ordination shells which may differ from those of the reactants is frequently not known with certainty further complicates the problem. In the absence of electronic interaction between the reacting species the reactants and products in a reaction such as (38) may be regarded as two different electronic states of the system (represented by the wave functions $R and $p) each corresponding to a distinct potential energy surface in a many-dimensional atomic configuration space whose co-ordinates include those of all the atoms of the two reactants and of the surrounding medium.To simplify our discussion we may consider only the first co-ordination shell of each ion which we assume to consist of six ligand molecules in a Fe2+aq + *Fe3faq -f Fe3+aq + *Fe2+aq . . . . . . . (38) 82 Laidler Canad. J. Chem. 1959,37 138. 83 Orgel Report X Conseil Chim. Solvay Brussels 1956 p. 289. 84 Griffith and George in Boyer Lardy and Myrback “The Enzymes,” Vol. 1 Academic Press Inc. New York 1959 p. 347. 85 Massey and Burhop “Electron and Ionic Impact Phenomena” Oxford University Press 1952. 86 Gurnee and Magee J. Chem. Phys. 1957,26,1237. 230 QUARTERLY REVIEWS w P Atomic configuration co-ordinate FIG. 2. Schematic potential energy diagram (after Marcus).regular octahedral arrangement. Our system then will have three atomic configuration co-ordinates the two metal-ligand separations and the separation between the centres of the two ions. Since the first two of these have different equilibrium values for the reactants and products [e.g. the equilibrium Fe-0 separations in Fe(OHJG3+ and Fe(OH,)62+ differ by about 0.15 A] the two potential energy surfaces will have minima in different regions of atomic configuration space. In general however there will also be a region of intersection of the surfaces corresponding to non-equilibrium atomic configurations in which the energies of the two electronic states are equal. The situation is depicted schematically in Fig. 2 where a single atomic configuration co-ordinate is used.A more complete treatment would take account not only of the first co-ordination shell but also of the polarisation of the surrounding solvent which will differ for the two states. Electronic interaction between the two states without which transition between them would not occur leads to the usual splitting of the surfaces as indicated in Fig. 2. Instead of two intersecting surfaces we now have an upper and a lower surface separated in the vicinity of the hypothetical intersection region by twice the interaction energy HRP (= <C$R IH lC$p>) of the two states. The qualitative behaviour which ensues depends to some extent on the magnitude of this interaction energy and it is convenient to consider three different cases. (1) Non-adiabatic electron transfer. This corresponds to the case of very weak interaction of the two states (e.g.when the reactants are very far apart or when the reactants and the products have different spin multiplicities) so that splitting of the surfaces is negligible. Because of the Franck-Condon restriction a radiationless transition between the two HALPERN ELECTRON TRANSFER IN SOLUTION 23 1 states can occur only in the vicinity of the intersection region where their energies are equal.* Thus there must occur before electron transfer a rearrangement of the atomic configuration of the system to some non- equilibrium (i.e. vibrationally excited) configuration usually intermediate between that of the reactants and the products The energy required for this has been denoted the Franck-Condon reorganisation energy.In such a case the apparent free energy of activation can be regarded to a first approximation as made up of the three contributions namely A F ~ = d F S r e p + AFfreorg - RTln K . . . . (39) where AFfrep is the free energy of repulsion between the reactants in the activated complex d Ff reorg is the Franck-Condon reorganisation free energy and K is a transmission coefficient representing the probability that the system will remain on the lower surface (i.e. that electron transfer will occur) when it passes through the intersection region. When this probability is small K is given6 by (40) -4n2HRp 4n2H2~p . . . [ h V l s R - SPI ] h V I S R - - P l K = 1 - exp where I& - &I is the difference of slopes of the crossing potential surfaces Y is the velocity of crossing and HRp is the interaction energy at the crossing point.The intersection surface will of course extend over many configurations of the system corresponding for example to different separations of the reacting ions. For ions of the same sign both d F $ r e p and K will in general decrease with increasing separation with opposing effects on dFf. In the activated complex of most favourable configuration i.e. that for which AFS is a minimum the separation between the exchanging ions will thus correspond to the best compromise between the electrostatic repulsion resisting close approach of the ions and the low probability of electron transfer at large separations. The quantities in equation (39) particularly K are not readily computed for systems as complex as those of actual interest. Some calculations have been attempted for the Fe2+-Fe3+ and related exchange reactions approxi- mating K by the probability of electron tunnelling though a potential barrier of some assumed shape.80$82 These treatments are necessarily very approximate and based in some cases on models of questionable validity but the resuIts serve to emphasise at least qualitatively some features of the behaviour described above.(2) Adiabatic transfer weak interaction. This corresponds to the limiting case where electronic interaction between the reactants in the * The Franck-Condon restriction applies also to electron transfer between polyatomic species in the gas phase. It has been estimated,s6 for example that the cross-section of the H2-H2+ and N2-N2+ electron transfers are reduced by factors of 0.31 and 0.94 respectively owing to the different equilibrium internuclear separations of the molecules and corresponding ions, 232 QUARTERLY REVIEWS activated complex is sufficient for the system to remain on the lower surface when it passes through the intersection region (i.e.K x l) yet weak enough for the contribution of the interaction energy to the lowering of the activation energy to be neglected. The products are now formed from the reactants adiabatically yet the splitting of the surfaces is small so that to a good degree of approximation the activated complex can be identified with the lowest crossing point of the hypothetical non-interacting surfaces and AF$ can be computed without explicitly evaluating the interaction energy. suggested that this approximation (which requires that the interaction energy in the activated complex lie within a rather narrow range roughly between 0-01 and 1 kcal./mole) is valid for many actual electron-transfer reactions in solution and has developed a quantitative theory based on it.In its original form this makes the further simplifying approximations. (a) Each reactant (i.e. the metal ion plus its co-ordination shell) is treated as a rigid sphere inside which no changes in interatomic dis- tances occur during the reaction. Thus only reorganisation of the surrounding medium which is treated as a continuous unsaturated dielectric contributes to A Ffreorg. (b) The separation r12 between the centres of the two reactants in the activated complex is taken as equal to the sum of their radii (rl + r2). This leads to the following expression for AF where the first term corresponds to AF$rep and the second to AFSreorg +m2X .. e1e2 AF$= - Dsh2 (41) ) / A . . . ' (42) e1e2 - e1'e2' Dsr12 where 2rn + 1 = - . . (43) el e2 are the reactant charges and el' e2' the product charges Ae (= el' - el) is the number of electrons transferred AFo is the standard free energy of the electron transfer step q is the refractive index of the solvent and D is the static dielectric constant. For an isotopic exchange reaction where AF" = 0 el = e2' e2 = el' rl = r 2 = r equation (41) reduces to (44) Values of AFT for several reactions computed from this expression are listed in Table 4. Reasonable agreement with experiment is found for HALPERN ELECTRON TRANSFER IN SOLUTION 233 reactions of covalent complexes such as Mn0,2=Mn04- Fe(CN) 64- Fe(CN),3- but for the aquo-ions dFScalc is considerably too low.This probably reflects the larger error introduced in the latter cases by failure to take account of the reorganisation of the inner co-ordination shells.* The assumption of adiabatic electron transfer (i.e. K = 1) represents another possible source of error in these calculations which is more likely to be serious for aquo-ions than for complexes such as cyanides in which the d-electrons are extensively delocalised. TABLE 4. Comparison of calculated and experimental AFI AFS (exp) Reaction (A) Temp. (kca1.l mole) Mn042-Mn04 - 2.9 1" 12.8 F~(CN):I(F~CN),~- 4.5 4 12.7 Fe2+-Fe3+ 3.4 0 16.3 Co2+-Co3+ 3.4 0 16.4 Fe(C,H;),-Fe(C,H,),+ 4.10 - 75 6-0 7 9 9 9 3.54 -75 6.0 AFS (@:rep/ ( c W dF (kca1.l mole) 9.2 0-17 10.1 1.0 9.8 0-5 9.9 0-5 5.3 - 6.1 - values.Ref. 79b 796 796 79b 14 14 More recently has extended the theory to take account also of reorganisation of the first co-ordination shell and of electrolyte effects but these refinements have not as yet been quantitatively applied. (3) Adiabatic transfer strong interaction. This corresponds to the situation which prevails in ordinary chemical reactions where covalent bonds are broken and new ones simultaneously formed. Redox reactions which proceed through the transfer of covalently bonded atoms (e.g. oxygen atoms or hydride ions) belong in this category as do electron- transfer reactions between metal ions which are strongly coupled by inner- sphere bridging ligands. In such cases lowering of the activation energy by interaction between the reactants in the transition state is likely to be appreciable and cannot be neglected.Absolute rate calculations as with ordinary chemical reactions thus become prohibitively difficult. A quantitative treatment of electron transfer based on a similar model has also been formulated by Hush.s8 Factors Mecting Rates of Electron Transfer.-Perhaps of greater significance than the rather restricted quantitative applications of the theoretical treatments cited above is the qualitative insight which they have provided into the factors which affect the rates and mechanisms of electron-transfer reactions and which account for some of the observed trends. Among these are the following. * Corrected values of dFtcalc. == 18 kcal.jmole which include the reorganisation free energies of the first co-ordination shells have been reporteds7 for the Fez+-Fe3+ and the Co2+-Co3+ reaction.Although details of the calculations were not given the method used is probably that described in Ref. 79c. Hush Discuss. Furaday Soc. 1960,29 113 116. 87 Marcus Trans. New York Acad. Sci. 1957 19,423. 234 QUARTERLY REVIEWS (1) Electrostatic repulsion. Both theoretical and experimental con- siderations suggest that dFSrep (given by e,e2/D6r) is usually small at least for reactions of the outer-sphere type in aqueous solution. Thus there appears to be no widespread correlation between rates of electron-transfer reactions and the charges of the reacting ions and many reactions between highly charged ions of the same sign are very fast e.g. Fe(CN),4-- Fe(CN),3- Fe ~hen,~+-Fe hen,^+.For the reactions listed in Table 3 the contributions of dFzrep to the total dFzcalc range from 15 % to 50 %. In inner-sphere reactions however because of the closer approach of the metal ions Coulomb repulsion between them may become more important particularly when the bridging ligand is uncharged. The very large difference (-106-fold) between the rates of oxidation of Cr2+ by CO(NH,),OH,~+ and Co(NH3),0H2+ is undoubtedly due in part to this. Also because of the dependence of d F S r e p on the dielectric constant this factor is likely to become more important in passing from water to less polar media; this may account at least in part for the generally slower rates of electron transfer in non-aqueous solutions. (2) Reorganisation energy. The theoretical predictions concerning the importance of this factor are well borne out by experimental evidence.In general slow electron transfer is observed when the reactants and products differ appreciably in atomic configuration i.e. in the geometries or dimensions of their co-ordination shells. In recent years the ligand-field theorys3 has contributed greatly to our understanding of this factor as it affects the reactions of transition-metals ions and complexes. This theory leads to the expectation of larger changes in metal-ligand bond lengths and hence larger reorganisation energies for the transfer of e than of t2 d-electrons between octahedral complexes. This factor undoubtedly accounts for the very slow electron transfer between Cr2+ and C?+ compared with say V2+-V3+ or Fe2+-Fe3+. ( 3 ) Electron “conductivity” of ligands.Unlike simple electron-transfer reactions between atomic species in the gas phase which depend on direct overlap of the donor and acceptor atomic orbitals reactions between metal ions generally involve electron transfer through intervening ligand and/or solvent molecules. The electronic interactions which determine the probability of electron transfer in such systems are thus rather complex and at best only qualitatively ~nderstood.~~ In a general way we expect the ease of electron transfer to depend on the extent of delocalisation of the donor and acceptor metal orbitals in the activated complex through mixing with orbitals of the intervening ligands.” From this standpoint ligands such as water and ammonia which contain saturated single bonds * Recent investigations of the nuclear magnetic and paramagnetic resonance spectra of transition-metal compounds provide direct information about this.Thus it is found that even in highly ionic substances such as MnF, the unpaired d-electrons of the metal are partly delocalised (to the extent of 5-10%) over the surrounding fluoride ions.89 Owen Discuss. Faraday Sm. 1958,26 53. HALPERN ELECTRON TRANSFER IN SOLUTION 23 5 are expected to be much less effective in "conducting" electrons between metal ions than unsaturated ligands such as CN- and phenanthroline whose complexes are characterised by a high degree of covalency and electron delocalisation. While the experimental evidence cited (Table 1) is on the whole in accord with this the comparisons involved (e.g. Feaq2+-Feaq3+ and Fe ~hen,~+-Fe hen,^+) are complicated by the fact that the geometrical configurations of the two oxidation states usually differ less in complexes of the latter type than in those of the former so that differences in reorganisation energy also contribute to the observed rate differences.Since the two effects arise to a large extent from related causes they are not readily separated experimentally and their relative importance is difficult to assess. (4) AF'. Equation (42) implies that the free energy dF" of reaction of the electron-transfer process enters into determination of the rate. The following physical interpretation may be placed on this where there is an overall decrease in free energy the requirement for atomic rearrangement before electron transfer (i.e. AFfreorg) is reduced since the product may now be formed in vibrationally excited states without the need for corresponding excitation of the reactants and the excess of vibrational free energy is dissipated as part of the overall free energy of reaction.The experimental evidence is on the whole in accord with this predic- tion. Thus electron-transfer reactions between dissimilar ions in which there is a-net free-energy decrease are in general faster (often too fast to measure) than the isotopic exchange reactions of either ion e.g. the oxida- tion of FeT1 by CeIV is faster than either the FeII-FeI" or the Celll-Celv isotopic exchange. It has been suggestedg0 that some of the variations (Table 1) in the rate of oxidation of Cr2+ by different Co"' complexes [e.g. CO(NH,),~+ and CO(NH,),OH,~+] are also due to this factor.Good correlations between AFf and AFo are exhibited by several series of redox reactions including electron transfer between Fez+ and different Fe(II1) c o m p l e x e ~ ~ * ~ ~ ~ as well as the oxidations of leucoindophenols by oxygen and of quinols by Fe3+ which are believed to involve electron- transfer mechanisms79b In other cases already cited correlations between ASS and AS" and between AH and AH" have been observed. (5) Multiple-electron transfers. As noted earlier there is some un- certainty as to whether 2-electron transfer reactions such as the TF-Tl1I1 exchange occur in a single step or through successive l-electron steps. The choice of mechanism in such cases will be influenced by the following factors. (a) Differences in AF". In reactions between 2-equivalent oxidants and 2-equivalent reductants these will normally be in the direction of favouring concerted 2-electron transfer.go Shimi and Higginson Discuss. Faraday SOC. 1960 29 122. g1 Ford-Smith and Sutin J. Amer. Chem. SOC. 1961 83 1830. 236 QUARTERLY REVIEWS (b) The lower probability of a 2-electron transition. While this has often been cited as an argument against 2-electron transfers there are grounds for questioning its general validity. Thus 2-electron-transfer processes such as A + A2+ + A2+ + A and Ne + Ne2+ + Ne2+ + Ne are known to occur in the gas phase with cross-sections only slightly lower to 4 in the cases cited) than those of corresponding resonant 1-electron transfers.86 A greater reduction in the relative probability of 2-electron transfer is however expected where interaction between‘ the reactants in the activated complex is very weak i.e.for reactions proceeding by a non- adiabatic mechanism. (c) Differences in A F f r e o r g . A greater difference in configuration and hence a larger reorganisation energy are expected when the two oxidation states differ by two electrons than when they differ by one. This is reflected in the ( Ae)2 term in equation (43) which implies four times as large a AF:,,,rg for a 2-electron transfer as for a corresponding 1-electron transfer. Because tlie other factors become relatively less important in such cases (a) is likely to dominate in reactions which proceed through strong-inter- action activated complexes e.g. by oxygen atom or hydride transfer or by inner-sphere bridged mechanisms. Indeed such mechanisms are commonly encountered in 2-equivalent redox reactions.On the other hand with “weak-interaction” activated complexes AFZreorg often makes the dominant contribution to the overall free energy of activation and in such cases [especially with reactions proceeding by non-adiabatic mechanisms where (b) also becomes important] I-electron transfers are likely to be preferred. For the Tl+-T13+ and related reactions particularly in the absence of a knowledge of the thermodynamic and other relevant proper- ties of T12+ the relative importance of the various factors is difficult to assess and the mechanisms in these cases remain uncertain. Valuable discussions with Professor H. C. Longuet-Higgins and Dr. L. E. Orgel are gratefully acknowledged.

ISSN:0009-2681    

DOI:10.1039/QR9611500207

出版商:RSC    

年代:1961

数据来源: RSC