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QUARTERLY REVIEWS NATURE OF METAL SOLUTIONS By M. C . R. SYMONS PH.D. (THE UNIVERSITY SOUTHAMPTON) THIS Review is concerned with the physical nature of blue solutions which result when alkali or alkaline-earth metals dissolve with no net reaction in solvents such as ammonia amines and ethers. Concentrated solutions of these metals in ammonia are metallic in appearance and general prop- erties and will only be considered when appropriate to the main discussion. The reactivity of these solutions will be mentioned when it bears upon the problem but this topic has been discussed in detail by Birch and Herchel Smith,l and there have been several other r e v i e w ~ . ~ ~ ~ ~ ~ Kraus whose pioneering work is still the most extensive and reliable in many of the fields of study has stressed2 that these unique solutions have attracted relatively little experimental attention and their importance to fundamental theory warrants much more careful study.One object of this Review is to examine recent results and theories and hence to suggest avenues of investigation which might prove fruitful. In section 1 some relevant facts are recorded stress being laid on recent work. In section 2 mention is made of the various theories that have been forthcoming concerning the nature of these solutions and in section 3 these theories are compared and discussed in the light of the facts. Section 4 concludes the Review with brief reference to possible future developments and certain related phenomena. There are many properties of these solutions which resemble those of “solutions” of alkali metals in solid alkali halide crystals.Because the latter phenomenon is far better understood some reference will be made to such additively coloured crystals which will be considered in detail in a forthcoming re vie^.^ 1. Facts All alkali and alkaline-earth metals other than beryllium dissolve to some extent in ammonia. Other solvents include primary and secondary amines and diamines certain ethers in particular diethers such as ethylene Birch and Herchel Smith Quart. Rev. 1958 12 17. Kraus J. Chem. Educ. 1953,30 86. Becker Linquist and Alder J. Chem. Phys. 1956 25 971. Bingel Ann. Physik 1953 12 57. Doyle and Symons Quart. Rev. 1959 in the press. 1 99 100 QUARTERLY REVIEWS glycol dimethyl ether MeOCH2CH2*OMe,s alcohols and even oxygen- free water.’ Solutions in alcohols and water are extremely unstable and little can yet be said about them.Blue solutions also result when alkali metals are added to molten salts such as halides hydroxides and amides 8 p 9 these solutions present an interesting meeting point between the solutions under discussion and additively coloured crystals but unfortunately little is yet known of the relevant physical properties. When ammonium and tetra-alkylammonium salts in ammonia are electrolysed blue colours are detected but the solutions are stable only in the latter case.8 These solutions deserve attention since they might well provide a means for choosing between the two models for dilute solutions outlined in section (2). Electrical Properties.-Little recent work has been reported and this subject is fully reviewed elsewhere.2,8 Conductivities show a marked minimum at about 0 .0 5 ~ for ammonia and 0 . 1 ~ for methylamine,1° which can be viewed as the merging point between concentrated solutions mentioned earlier and dilute solutions which have electrical properties that are often described as “salt-like”. The reader is referred to these earlier reviews for details; suffice it to say here that in the region between 0 . 0 1 ~ and infinite dilution solutions of metals are truly “salt-like” in their electrical properties the negative ion whatever its nature being responsible for the conduction of about six-sevenths of the total current. It can be argued from the results of conductivity studies that ion-pairing is extensive :2 however since considerable ambiguity still clouds the physical significance of this statement when simple salts are considered it is probably unsafe to infer much from this conclusion.In solvents of low dielectric constant such as ethers direct contact between ions is to be expected and has been beautifully demonstrated by Adam and Weissman.ll Such a phenomenon is of great importance here and will be discussed in detail later. Ionic Volumes.-When alkali metals dissolve in ammonia there is a remarkable increase in volume amounting in the case of concentrated solutions to about 43 CM.~ per “mole” of electrons. Unfortunately nearly all the work which has been done relates only to concentrated solutions (down to about IM). Kraus et aZ.12 suggested that their results could be extrapolated to about 40 ~ m . ~ in the dilute range but this very long extra- polation is certainly unjustified as it stands.One recent experiment does appear to support this view however Stosick and Hunt have obtained a value between 20 and 70 ~ r n . ~ for a solution containing about 10-3~- * Down Lewis Moore and Wikinson Pruc. Chem. Soc. 1957 209. Jortner and Stein Nature 1955,175 893. Fernelius and Watt Chem. Rev. 1937,20 195. Johnson and Bredig J. Phys. Chem. 1958,62,604. lo Evers Young and Panson J. Amer. Chem. Soc. 1957 79 51 18. l1 Adam and Weissman ibid. 1958 80 1518. la Kraus Carney and Johnson ibid. 1927,49 2206. SYMONS METAL SOLUTIONS 101 ~0dium.l~ Considerable weight has to be put on this result to justify the extensive calculations which have been made;14*3 the datum is of very great significance however and it would be far more satisfactory if more results were available for a range of different dilute solutions.Concentrated solutions of potassium15 and lithium16 in ammonia behave similarly. The net volume change for potassium is somewhat smaller than that for sodium but if suitable adjustment is made for the volume changes which occur when atoms are converted into ions and for electrostriction effects,14 then the final volume changes are comparable for all three metals. The very high compre~sibilities~~ and low viscosities18 of concentrated solutions parallel the large volume expansions but again results for dilute solutions have not been obtained. Magnetic Properties.-Freed and Sugarman'sl9 very careful work on the static magnetic susceptibilities of potassium solutions down to about O-OO~M together with somewhat less extensive work by Huster20 on sodium solutions has recently been supplemented by a study of the electron-spin resonance spectra of these solutions down to about O - O O ~ M .~ ~ These alternative methods give broadly the same results except that the diamagnetism due to the electron centres is included in the former but not in the latter measurements. The small difference in susceptibility obtained by these two methods is therefore a measure of residual diamagnetism and has been used by Hutchinson and Pastor21 to give information concerning the solute species. The major result of these studies is that extensive electron-pairing occurs such that for solutions above about 0. IM pairing is nearly complete whilst below about 0.001~ the concentration of paired electrons is negligible.For a given concentration pairing decreases with increasing temperature and at least at room temperature sodium solutions show less tendency to form electron-pairs than potassium solutions of equal molarity. Once again concentrated solutions resemble liquid metals in their magnetic properties. One of the most striking features of the electron-spin resonance spectra is the extreme narrowness of the line. For dilute solutions of sodium or potassium in ammonia at -33" the width between points of maximum slope ( ~ H M s ) is about 0.05 gauss.21 This width decreases with an increase in temperature and increases when methylamine is used as solvent and also when ammonia is replaced by ND3. When ammonia solutions are frozen the metal is precipitated as very small clusters and the resonance is then a l3 Stosick and Hunt ibid.1948 70 2826. l4 Lipscomb J. Chem. Phys. 1953,21 52. l5 Johnson and Meyer J. Amer. Chem. SOC. 1932 54 3621. Johnson Mayer and Martens ibid. 1950,72 1842. l7 Maybury and Coulter J. Chern. Phys. 1951 19 1326. Kikuti J. SOC. Chem. Ind. Japan. 1944 47,488. l9 Freed and Sugarman J. Chem. Phys. 1943,11 354. *O Huster Ann. Physik 1935 33 477. 21 Hutchison and Pastor J. Chem. Phys. 1953 21 1959. 102 QUARTERLY REVIEWS function of the metal itself.22 If however the solutions are saturated with sodium iodide they solidify to clear blue glasses on cooling and the electron resonance signal changes to a single line having ~ H M S = 3.5 gauss.23 . Only certain metals in amine solvents give paramagnetic solutions. Thus lithium in methylamine gives an intense electron resonance absorption potassium a weak one whilst sodium in ethylenediamine is diamagnet ic.24 No paramagnetism can be detected for solutions in ethers.25 Nuclear Magnetic Resonance.-McConnell and Holm have studied the nuclear resonance spectra from 23Na 14N and lH in solutions of sodium in ammonia and found very large shifts to low fields for 23Na and 14N but the proton resonance was not detectably different from the standard.26 These measurements were made on a low-resolution instrument so that relatively small “chemical shifts” would not have been detected.These shifts were described as Knight shifts because they closely resemble the shifts found for metals however it is a general phenomenon for nuclear resonance spectra to be both broadened and shifted when there is the chance of close proximity between species containing unpaired electrons and the magnetic nuclei concerned.The shift may in a general sense be thought of as a modification of the magnetic field experienced by the nucleus under observation by the extra field due to the magnetic electrons. This shift is often much greater than the more familiar “chemical shifts” which occur in the absence of paramagnetic material. The shifts observed for 23Na and 14N decrease on dilution and could not be detected when the concentration of sodium was less than about 0 . 1 ~ . At this concentration the sodium valency electrons are almost en- tirely paired and the results unfortunately cannot be extrapolated to the dilute region. McConnell and Holm consider that the shifts are due to contact hyperfine interactions which give a finite density for the unpaired electrons at the magnetic nuclei and they express their results in terms of a parameter Pi the contact density of an electron on the atom in these solutions.This is compared with estimated values for Pio the contact density for an electron permanently in a given orbital on atom i. They conclude that the unpaired electrons are about 5 x times as effective in shifting the sodium line as 3s electrons on sodium and 0.1 times as effective as unpaired electrons in a 2s state on nitrogen. There is a significant decrease with dilution in the value of P N a but PN remains nearly constant in the concentration range 1 - 0 - 0 . 1 ~ ~ ~ ~ These very important results will be discussed in detail later. Electronic Spectra.-Generally there are only two inter-se absorption 22 Levy Phys.Rev. 1956 102 31. 23 Clark and Symons J. 1959 26 Wilkinson Cotton Fischer Down and Moore A h . 133rd A.C.S. Meeting 1958 26 McConnell and Holm J. Chem. Phys. 1957,26 1517. Fowles McGregor and Symons J. 1957 3329. 12L. SYMONS METAL SOLUTIONS 103 bands in dilute metal solutions one or both of which are invariably d e t e ~ t e d . ~ " ~ ~ These bands which are found in the 7000 and 15,000 cm.-l regions will be referred to as the infrared and visible bands respectively. These bands are so intense + 4 x lo4) that even when 0-l'mm. cells are used the solutions still have to be very dilute (ca. 10"~) for measure- ments near the band maxima. This point is of great importance since nearly all other physical measurements are made with more concentrated solutions.The solutions appear blue irrespective of which band is present. Very dilute solutions of all alkali and alkaline-earth metals in ammonia have identical spectra consisting only of the infrared band.27,28 The posi- tion of maximum absorption and the band width depend on temperature and concentration but not on the nature of the metal. Addition of a strong electrolyte results in a very small shift tc higher energy but does not alter the band width. High concentrations of added salts give rise to a new band of low intensity with a peak at about 12,500 The infrared band is very broad and is asymmetrically broadened on the high-energy side. When the high-energy edge is examined in more concentrated solutions (ca. 10-3~) it is found to extend into the near ultraviolet region and there is a pronounced shoulder in the 15,000 cm.-* region.23 Thin films of more concentrated solutions in ammonia show the visible and ultraviolet bands at 20"q though these are both at higher energies than usual because of the very low temperature.2B In methylamine and mixtures of ammonia and methylamine either or both bands appear depending upon the choice of metal temperature and solvent composition.Both bands in methylamine are shifted to somewhat higher energies compared with those in ammonia and the positions and widths are again found to be a function of temperature and concentration though they are not sensitive to the nature of the metal In ethyl- amine27 and certain diamines2* the visible band predominates and this is the only band found for solutions in ethers.6 It has been reported that the spectra of the very unstable blue solutions in water have a band maximum at 11,000 cm.-l.There is a close link between the spectra and magnetic properties of these solutions. Paramagnetism is only detected in those solutions which have a band in the infrared region and it has accordingly been postulated that this band is characteristic of the units containing unpaired electrons whilst the visible band is due to units containing e1ectron-pai1-s.~~ Thus very dilute solutions in ammonia in which all the solute electrons are un- paired show only the infrared band. More concentrated solutions in which electron-pairing is significant show both bands Again solutions of lithium in methylamine are paramagnetic and show both bands but other 27 Blades and Hodgins Canad.J. Chem. 1955 33 411. 28 Jolly U.S. Atomic Energy Comm. Nat. Sci. Foundation Washington D.C. 1952 29 Bosch 2. Physik 1954 137 89. U.C.R.L. 2008,3. 104 QUARTERLY REVIEWS metals in amines and ethers which show only the visible band are dia- magnetic. In all these solutions no other band appears before the onset of intense absorption by the solvent. The similarity to alkali halide crystals con- taining F or F’ centres is striking. (An Fcentre is a single electron and an F’ centre an electron-pair in a halide ion vacancy.) Both these centres give rise to intense absorption bands in the visible region but in contrast with metal solutions the F‘ or electron-pair band is always found on the low- energy side of the F band. Extinction coefficients and oscillator strengths are comparable and the long-wavelength sides of the crystal bands are again asymmetrically br~adened.~~ As with the solutions no other bands attributable to these centres are detected before the onset of the funda- mental absorption bands of the crystals.Photo1ysis.-There is no net decomposition when solutions of sodium in ammonia are photolysed either within the absorption bands or down to the onset of absorption by the solvent (about 40,000 cm.-l). However the solutions are readily bleached when irradiated with light of wavelength less than 250 mp to give sodamide and hydrogen although ammonia itself is not affected under these condition^.^^ Rigid solutions of sodium in ammonia containing sodium iodide are apparently unaffected by visible light,a3 but rigid solutions of lithium in a glass containing methylamine whose spectra are dominated by the visible band (600 mp) are readily bleached by light of wavelength less than 500 mp.When the glass is softened and recooled the 600 mp band reappears showing that no net decomposition has occurred.31 As the 600 mp band is bleached so a broad band with a maximum in the near infrared region appears which must be caused either by the ejected electron31 or possibly by the electron left behind.24 X-Ray Scattering.-The small-angle scattering of X-rays from con- centrated solutions of lithium sodium and potassium in ammonia has been investigated by Schmidt.32 Unfortunately reliable results could only be obtained in the 1-OM region but Schmidt draws some very interesting conclusions. He finds scatttering units with dimensions of the order of 15 A (but somewhat different for different metals) and states that only a few types of centre can be present.There does not seem to be any scattering which can be associated with large cavities (radius ca. 3 A) independently of the metal. The scattering shows a maximum intensity at a scattering angle well removed from zero. Schmidt concludes that the centre must therefore have regions of both excesses and deficiencies of charge. Energies.-Jolly has discussed some thermodynamic functions for metal-ammonia solutions. He estimates AH (4e3 = 40.5 kcal./mole 30 Ogg Leighton and-Bergstrom J. Amer. Chem. SOC. 1933,55 1754. 31 Linschitz Berry and Schweitzer ibid. 1954 76 5833. 32 Schmidt J. Chem. Phys. 1957 27 23. SYMONS METAL SOLUTIONS 105 A H f (el) = 43.5 kcal./mole AFf = 44.4 kcal./mole and AH = 6 kca1.l mole for the process where e and e symbolise the units containing the unpaired and paired valence electrons of the Hutchinson and Pastorz1 calculate A H = 2-3-3-9 kcal./mole for reaction (1).For the same process in mixed ammonia-methylamine solvents Blades and Hodgins2' estimate A H = 4-6 kcal./mole. Photoelectron emission from dilute solutions of sodium in ammonia gives a threshold of about 34 kcal./rn01e.~* It has been suggested that this figure gives a rough measure of the solvation energy for electrons in these solutions. 2. Theories e -f 2e . . . . . . . . (1 1 It is convenient to consider these solutions in three groups depending upon concentration. These will be described as concentrated solutions (above about OM) medium solutions (between 1 .0 ~ and 0.05~) and dilute solutions (less than 0.05~). The metallic properties of concentrated + Sol va t ed e Cavity sodium ion monomer Ex pa nded- meta I Expanded-metal monomer Expanded - meta I d imer solutions set them apart from the others and most of this discussion will be concerned with medium and dilute solutions. In many solvents these are the only solutions formed. There are two models for medium and dilute solutions which are fundamentally different but have an equal claim for serious consideration. These will be referred to as the "expanded-metal" theory and the cavity 33 Jolly Chem. Rev. 1952,50,351; U.C.R.L. 2201 1953. 34 Hasing Ann. Physik 1940 37 509, 106 QUARTERLY REVIEWS theory. The former presents a solution containing solvated positive ions with the valence electrons of the metal moving in “expanded orbitals” around these solvated ions.This theory is not new but has recently been given prominence by Becker Lindquist and Alder3 who explain the formation of electron-pairs by postulating an “expanded-metal” dimer in which two solvated cations are weakly bound together by two electrons of opposed spins. The cavity theory favoured by Kraus presents a solution containing solvated metal ions and solvated electrons. The solution is thought to resemble that of any strong electrolyte except that in place of negative ions in cavities of oriented solvent molecules there are electrons in similar cavities. Under certain conditions two electrons with opposed spins are thought to occupy the same cavity. These are overall pictures there is a lot to commend both models and they will be considered in turn in section 3.Many authors have attempted to make one or other of these theories more precise and some of these detailed descriptions will first be considered. An attempt will be made to keep these two aspects separated since there has been a tendency to assume that when experiment has shown that a particular description of one of the theories is in error then the whole model is proved false. Before presenting these mathematical descriptions some alternative models should also be mentioned. The Colloidal-metal Theory.-From time to time it has been suggested that the metal is present as a colloid probably with a net negative charge The case was put quite strongly by K~-iiger,~~ who pointed out that small colloidal metal particles should be coloured but his evidence was strongly criticised by Freed and T h ~ d e .~ ~ Hunt and his co-workers favour this but there seems to be little evidence for it at present. Carefully prepared solutions in liquid ammonia or amines do not show a Tyndall cone. Colloids showing a band maximum in the 15,000 cm.-l region prob- ably would be too small to scatter light but colloids absorbing at 7000 cm.-l should certainly show such a cone. The colloid theory offers no explanation for the appearance of two distinct bands in certain solutions but one could argue that only one of these is a colloid band. Since the infrared band has been associated with the unpaired-electron species it cannot be a colloid band since the number of unpaired electrons in a colloid particle would be small.Let us therefore consider the postulate that the visible band is due to colloidal metal. It is a general property of the colloid band due to excess alkali metal in alkali halide crystals that the position of maximum absorption is strongly dependent upon the nature of the metal and host crystal but both the peak and width are independent of temperature and the band is unaffected when the crystal is irradiated 35 Kruger ibid. 1938 33 265. 36 Freed and Thode J. Chem. Phys. 1939,7 85. 37 See Meranda Diss. Abs. 1957 17 249. SYMONS METAL SOLUTIONS 107 with visible light.38 The visible band in metal solutions has properties which are opposite to these in each case (section 2). Ammonium-metal Theory.-The term is misleading for dilute solutions.Certainly if a solution contained ammonium ions and sodium ions there might be a competition between these positive centres for the electrons provided that one postulated the “expanded-metal” model mentioned above. However if ammonium salts are added to metal-ammonia solutions hydrogen is rapidly evolved and the solutions are decolourised. In other words conditions which favour the formation of NH radicals also favour hydrogen evolution probably because these neutral radicals provide a low-energy pathway for decomposition which does not require the trans- ient formation of hydrogen atoms :23 e f NH,+ SNH,; 2NH4 -t 2NH,+ H . . . . . . . . (2) If NH formation were important in solutions containing no added ammonium salt26 there would have to be an equal concentration of amide ion in the solutions.The absence of any absorption in the 340 mp region in these solutions means that in fact the concentration of amide ion is negligibly small since the amide ion has an intense absorption band in this region.30 P i t ~ e r ~ ~ in an attempt to explain the results of nuclear resonance studies,26 has postulated the formation of the ion NH3- in which the un- paired-electron orbitals are “3s-like for nitrogen with the outer node at the N-H bond radius”. He goes on to show that this would explain both the large shift for nitrogen and the small shift for hydrogen. Since this is not a bonding orbital for ammonia molecules in the gas phase the only way in which it can become one in solution is by polarisation of the surrounding medium. If this is complete then the NH,- ion will presum- ably be solvated by an oriented shell of solvent molecules and the model becomes very similar to the cavity model with an extra ammonia molecule at the centre of the cavity.It is not clear from the text whether or not this is the intended The alternative would be a model in which the electron moves amongst solvent molecules so rapidly that orientational polarisation cannot occur. Such a state of affairs is probably similar to the ionized state envisaged by Becker et aL3 This “conduction electron” model is probably incompatible with spectroscopic results since the very well- defined infrared band must be associated with these electrons and it is hard to see how such a precise transition could be associated with a mobile electron of this sort. Also one might have expected high conductivity in frozen solutions but dilute solid solutions seem to be non-c~nducting.~~ Detailed Descriptions of the Expanded-metal and the Cavity Model.- Becker et aL3 made two suggestions about the expanded-metal postulate.3* Compton Phys. Rev. 1957 107 1271. 38 Pitzer J. Chem. Phys. 1958 29 453. 108 QUARTERLY REVIEWS When discussing spectra they used a model in which the electron is con- strained to move on the surface of a sphere defined by the hydrogen atoms of the first layer of oriented ammonia molecules surrounding the cation. They suggested that the transition is one which involves a change in angular momentum and showed that the infrared band is predicted satis- factorily on such a model. However when discussing energies they made use of a different model in which a hydrogen-like 1s function is used such that it has a maximum at 3-5 8,.McConnell and Holm26 postulated a very diffuse orbital and compared the system with that of silicon containing traces of phosphorus. By loss of one electron to form P+ the phosphorus is able to replace silicon and the unpaired electron is then held in an orbital centred on the positive charge but extending a considerable distance from the phosphorus atom. The simplest description of the cavity model is that of an electron (or an electron-pair) in a spherical b o ~ . ~ * J ~ The lowest energy level in such a box is h2/8mrO2. In addition to this there are terms due to the polarisation of the dielectric medium by the electron including orientation of permanent dipoles around the cavity and for the surface-tension energy.Lipscomb14 developed this model in order to account for the volume expansion men- tioned in section 1. He estimated from the experimental data that cavities of radius 3-2 A were required and his detailed calculations gave a value of about 4.8 A for the radius of the spherical box. Stairs*l modified this theory by assuming a finite depth for the box he assumed a value of 3.2 8 for the radius and calculated values for the energy of the system. A somewhat different approach was made by P l a t ~ m a n ~ ~ who used Platzrnan and Frank’s which was developed to explain the absorption spectra of halide ions in solution. In this model the electron moves in a discrete centrosymmetric orbital defined by the potential field of the polarised solvent molecules around the electron.The medium is treated as a continuous dielectric and the wave function is a 1s function similar to that suggested by Becker et al. for the “expanded- metal” m0de1.~ He suggested42 that the different bands observed in various solutions are all part of a series of transitions of the electron in this poten- tial well but the considerations given in section 1 show that this theory is untenable. A very similar discussion has recently been presented by J ~ r t n e r ~ ~ who estimated a value of 3-6 8 for the cavity radius assuming four solvent molecules around the cavity. Hence he calculated hu = 1.1 ev for the 2p c 1s transition. This value should be compared with the experimental result of 0.8 ev. Lipscomb concluded that a further attraction potential was required and suggested that interactions between the electrons and the protons of the oriented polarised solvent molecules might be responsible for this dis- 40 Hill ibid.1948 16 394. I1 Stairs ibid. 1957 27 1431. 42 Platzman unpublished results quoted in footnote to p. 423 of ref. 43. Is Platzman and Frank 2. Physik 1954 138 41 1. Jortner J. Chem. Phys. 1957 27 823. SYMONS METAL SOLUTIONS 109 crepancy. Kaplan and Kitte145 have developed this idea and presented a model in which the unpaired electrons exist entirely in delocalised molecular orbitals on all of the protons which define the cavities. The N-H bonds are considered to be partially ionic thus permitting the electron to occupy the 1s level though higher levels are also considered. This model which was based on the highly successful model for F-centres proposed by Kip et aZ.,d6 was used to explain the results of electron spin resonance studies.21 One aspect of this work is that the authors assumed that the electrons occupy cavities which are already present in the liquid rather than cavities which are built up around the electrons in the same way as ions become solvated.They considered the line-width of the electron spin resonance spectra and assumed that only the protons contribute to this by hyperfine inter- actions. Their result is of the right order of magnitude only if they take into account the narrowing effect caused by rotation of the ammonia molecules the value for the width estimated when this motion is neglected being 13 f gauss where f is a measure of the effective s character of the molecular orbital on hydrogen.In view of the many approximations involved the numerical results can only be taken to mean that the model is reasonable for mobile solutions. The experimental result of 3.5 gauss for AHMs for rigid is appreciably less than the predicted value however unless is very much less than' unity which seems unlikely. 3. Discussion The various alternative models outlined in section 2 are based upon well- authenticated analogues in the solid state and i t is helpful to make certain comparisons. In the case of phosphorus-doped silicon mentioned above the electron is trapped at the cation centre because there is no alternative centre of any binding power in the crystal. However in an alkali-halide crystal the trap is an anion vacancy rather than a cation. In solution one can say that the electron can choose between solvated cations and the potential wells formed by oriented solvent molecules which can be formed in bulk solvent as the result of polarisation by the electron.Both alterna- tives seem reasonable. The mathematical descriptions outlined above have been used together with a variety of others to describe colour centres in solids with varying success. The results of electron spin resonance and double resonance studies on F-centres suggest that those models in which the orbital for the unpaired electron is considered to be diffusely spread over a large volume of solid are una~ceptable.~~,~~ The same argument may well be true for solutions. Concentrated Solutions.-The isolation of solids such as Ca(NH,), which have all the properties of a metal suggest that the expanded-metal theory must be substantially correct for concentrated solutions.That is 45 Kaplan and Kittel ibid. 1953 21 1429. 46 Kip Kittel Levy and Portis,.Phys. Rev. 1953 91 1066. 47 Wertz Auzins Weeks and Silskee ibid. 1957 107 1535. ** Feher ibid. 1957 105,51122. 110 QUARTERLY REVIEWS that the weakly held valency electrons bind together units consisting of metal ion and about six solvent molecules oriented around them. Pitzer has given an interesting treatment of these solutions and has shown how this model may be used to explain most of the known properties including the phenomenon of phase separation in concentrated solution^.^^ In this way the properties of concentrated solutions are at least qualita- tively explained. The marked volume changes are a natural outcome just as when sodium ions become sodium metal there is an apparent radius increase of about 1.2 A for each ion so when solvated sodium ions [say Na(NH3I6+] become “expanded metal” made up of Na(NH,) units an apparent increase in radius of 0.9 A would account for the experimentally observed expansion.Medium Solutions.-In this range (1 M-O-O~M) the solute valency electrons remain almost completely paired but the metallic properties of concentrated solutions give way to salt-like properties. The “expanded- metal” theory gives a picture in which dimer units predominate (see Figure). The conductivity falls rapidly as the average distance between dimers becomes too great for effective electron tunnelling but the change in paramagnetism is small since dissociation into monomer units or metal ions and electrons is still unimportant.but if Schmidt’s conclusions are correct the cavity theory does not. The latter theory tells us that on dilution the concentrated solutions break down to give simple solvated sodium ions together with electron-pairs in separate solvent cavities (e cavities). The electronic absorption band in the 15,000 cm.-l region must be ascribed to dimer units if the expanded-metal theory is correct. Becker et al. have estimated that this unit should absorb in the 7000 cm.-l region but the model used is very crude and the lack of agreement can hardly be used as an effective argument against the theory. If the e cavity theory is correct then the visible band must be caused by the excitation of one of the electrons to a higher level in the cavity. Whilst this explanation may be satisfactory for solutions in ammonia it is less so for solutions in solvents with very low dielectric constants.That is because in this concentration range the e cavity theory is modelled upon a solution of a 1:2 electrolyte. Such electrolytes are generally only sparingly soluble even in ammonia and when they dissolve ion-pair formation is extensive. In ammonia this does not necessarily mean that the ions are adjacent but in solvents of very low dielectric constant the ions are probably solvated as dipoles except in very dilute solutions.11~50 This phenomenon which will be described as dipolar contact has been This theory adequately explains the results of X-ray scattering 49 Pitzer J. Amer. Chem. Soc. 1958 80 5046. 5 0 Hughes Ingold Patai and Pocker J.1957 1206. SYMONS METAL SOLUTIONS 111 detected by spectrophotometry for salts such as tetra-n-butylammonium iodide even with concentrations as small as 10-4~.51 It is probable therefore that in ethers dipolar contact would be almost complete for 1:2 electrolytes. As Kraus has pointed out,2 one cannot picture such ion-pair formation in the usual way for metal solutions it is quite im- possible to construct a solvated-dipole model from sodium ions and electron-pairs. One is forced to conclude that under these conditions the electrons will move in an orbital centred on the sodium ion. If we go further and postulate that the equilibrium Na+(solv.) + Na-(soh.) + ZNa(solv.) . . . . . . . . . (3) lies well to the right then we have returned to the “expanded-metal” model for these solutions.In that case the band at 15,000 cm.-l found in metal-ether solutions must be ascribed to the dimer and it is a logical though not compelling extrapolation to suggest that the band in this region found for medium solutions in ammonia is likewise due to “expanded-metal” dimer.52 Dilute Solutions.-For solutions in ammonia this is the region in which electron-pair units break up to give single electron species. For solutions in ethers this does not occur and for solutions in amines it may or may not occur depending upon conditions. We will examine this process for ammonia solutions in detail. The picture we have to adopt for the “expanded-metal” theory is that on dilution dimers dissociate to give paramagnetic monomer units. While this process symbolised in the Figure accounts qualitatively for the magnetic phenomena it does not explain the great increase in electrical conductivity on dilution.Becker et al. overcame this difficulty by postulat- ing equilibrium (4) Monomer + Na+(solv.) + e . . . . . . . . (4) They compared this dissociation with that of a weak electrolyte and suggested that the electron ‘‘ becomes associated with other protons of bulk solvent ammonia molecules”. It is not clear in what state these “free” electrons are supposed to exist but the authors do not mean to imply that the electron is in a cavity of oriented solvent molecules. Indeed they con- sider that the cavity theory is fraught with many shortcomings and reject it entirel~.~ At one stage in their calculations they assume that the electron has direct access to the entire volume of the solution.Good evidence for the presence of this monomer unit in small con- centration in concentrated and medium solutions comes from the shift in the 23Na nuclear resonance spectrum indeed it would be very hard to explain this shift in any other way. However although this result shows 61 Grfiths and Symons unpublished results. ’’ Symons J. Chern. Phys. 1959,30. 112 QUARTERLY REVIEWS that the monomer unit can exist in metal solutions it does not prove that the unpaired electrons in dilute solutions are bound in this way rather than in cavities constructed of oriented solvent molecules (e cavities). More information can be obtained from a consideration of the spectra of dilute solutions in ammonia. In order to explain adequately the conduct- ivity of dilute sodium solutions Kraus has estimated a dissociation con- stant of 0.05 for the “ion pairs” at -33°.2 This is about the same as the value estimated from magnetic data by Becker et al.for the dissociation of rnon~mer.~ The latter authors suggest that the infrared band found in dilute solutions is due to excitation of the unpaired electrons to a higher level around the sodium ions. The important feature of this argument is that the monomer unit is responsible for the colour. A calculation using the equilibrium constants quoted3 shows that when the overall concentration of metal is less than 10-3~ the concentration of monomer is negligible compared with that of “free” electrons. For example for an overall con- centration of 10-4~ the concentration of monomer is about ~ O - ’ M . ~ ~ Nevertheless such a solution is still blue and since the molar extinction coefficient at 6700 cm-l is about 4 x lo4 the optical density of such a solution in a 1 cm.cell would be 4 at the peak. The oscillator strength of this band is23 about 0-7 if one assumes that all the unpaired electrons are responsible for the absorption the very small amount of monomer sup- posed to be present in these solutions could not possibly be the cause of this band and it would be necessary to postulate that practically all the electrons were still bound to sodium ions befme such a situation could be entertained. In other words one would have to reject the value of 0.05 for the dissociation constant and the negligibly small concentration of “free” electrons would have to have extremely high mobilities to account for the conductivity of these solutions.Against this conclusion is the definite decrease on dilution in the value of P N a deduced from the Knight shift for 23Na.26 Also the results from electron spin resonance studies on rigid glasses are not in accord with this possibility. If most of the paramagnetic centres are monomer units it is hard to understand why the spectrum is devoid of hyperfine The spectrum was recorded as the first derivative of the absorption curve and would be very sensitive to the presence of poorly resolved hyperfine lines. McConnell and Holm’s values26 of PNa and PNao being used it can be shown that splitting between outermost lines of the quartet expected for 23Na would be about 5 gauss. Interaction with neighbouring magnetic nuclei (lH and 14N) could broaden these four lines beyond the point of resolution but it is hard to see how a single symmetrical line of width 3 gauss could result.Indeed since this quartet is clearly resolved for fluid solutions containing Ph2CO-Na+ one might expect to have detected a quartet in fluid solutions of sodium in ammonia. There is another argument against the monomer theory. The addition of a sodium salt to a dilute solution of sodium in ammonia should have little effect upon monomer units and if the infrared band were due to electronic SYMONS METAL SOLUTIONS 113 transitions of the nionomer then the only change expected would be a slight shift or broadening. On the other hand if the band were due to excitation of electrons in solvent cavities then excess of sodium ion could shift the equilibrium depicted in Fig.2 to favour monomer and hence one might hope to observe a new band characteristic of the monomer. When sodium iodide is added a new band is observed in the 12,500 cm.-l region in addition to the 6700 crn,-l band.23 This band is therefore tentatively assigned to monomer. The band at 15,000 cm? has already been assigned to dimer units. To check the reasonableness of these suggestions one can compare these spectra with the well-established spectra of alkali-metal atoms and m01ecules.~~ The atoms have a doublet with very small separation in the 14,000 cm.-l region (P t S ) and the molecules have two bands one on either side of this (lCU+ t lCg+ and t- lC0+). There is a trend to lower energies on going from sodium to casium for all bands and the separation between the two bands for the molecules decreases.For example the bands for K and K2 are at 13,000 cm.-l (average) and 11,671 and 15,370 cm.-l respectively. The allocation for the solution bands therefore seems reasonable except that one might have expected a doublet for the dimer. However the 15,000 cm.-l band is broad and asymmetric and may well be both bands. It would not seem reasonable to assign the 6700 cm.-l band to monomer and the 15,000 cm.-l band to dimer if this analogy is good. These considerations lead to the conclusion that the blue paramagnetic unit in dilute solutions I s an electron in a solvent cavity and not the “expanded-metal” monomer.52 This being combined with conclusions drawn in the previous section it appears that of the four possible units (expanded-metal dimer and monomer; e2 and el cavities) only two ex- panded-metal dimers and el cavities need be considered as major com- p o n e n t ~ .~ ~ There is good evidence for the existence of monomer units but these do not seem to be major components under any conditions. No good evidence for e2 cavities has been found. It has been postulated that the absorption band in the visible region was due to excitation of electrons in e It was mentioned in section 2 that these units are comparable with F’ centres in alkali halide crystals. However the band characterising F’ centres is of lower energy than the Fcentre band and by analogy one might predict that the band associated with e cavities would be of lower energy than that due to single electrons in similar solvent cavities. It is hard to see why the energy of the e2 transition should be more than double that of the el transition.Consequences.-This hybrid model should be examined in the light of The large volume increment found for dilute solutions is accommodated those properties not so far considered. 114 QUARTERLY REVIEWS more satisfactorily than by the “expanded-metal” monomer theory Lipscomb’s argurnentsl4 are not invalidated. The Knight shift for 23Na nuclei are satisfactorily accommodated. Although the concentration of monomer units is always small it is quite sufficient to account for the shifts reported the only modification required is that the orbitals for the unpaired electron must be somewhat less diffuse than postulated by McCoiinell and Holm.26 The lack of shift for protons is a puzzling feature whatever model is adopted.One definite conclusion is that Kaplan and Kittel’s detailed picture cannot be correct that is the unpaired electron does not spend its time moving in s orbitals on the protons of the cavity. One can equally well reject the mathematical description given by Becker et aL3 when discussing spectra. Absence of a Knight shift for protons does not mean that protons and electrons do not interact. The shift expected for normal dipolar interaction is to high fields and therefore opposite to the shift caused by hyperfine contact interaction. If these effects are comparable then no shift would be Just this sort of cancellation was found by Phillips Looney and Ikeda for the effect of paramagnetic salts on the hydroxyl proton resonance of certain alcohols.63 The large Knight shift for nitrogen is also hard to understand on any of the models described.The explanation given by P i t ~ e r ~ ~ is satisfactory provided that one postulates normal solvation for NH3; but the major bonding forces must still be the same as for the cavity model and a possible alternative explanation would be that the electrons move in part around those oriented solvent molecules which define the cavities in orbitals of the type postulated by P i t ~ e r . ~ ~ The net effect would then be similar to that expected for NH3- and the model would resemble that for F centres in alkali-halide crystals.46 4. Future developments Most of the reasoning given in section 3 is based upon magnetic and spectrophotometric data and many of the postulates made could be tested by applying these techniques simultaneously to a variety of solu- tions.Electron spin double resonance48 might well solve many problems if applied to frozen glasses of various concentrations and magnetic studies of photolysed glasses in which the 15,000 cm.-l band has been bleached would be of great interest. The effect of added salts on the conductivity would be worth investigating if our tentative conclusions are correct there should be a decrease in conductivity with increase in concentration of added salt. In conclusion brief reference will be made to the postulate that the first excited state of solvated iodide ions may be compared with an electron in a cavity having an iodine atom at its centre.54 The radius of the 53 Phillips Looney and Ikeda {bid. 1957 27 1435. 64 Smith and Symons Trans. Faraday SOC.1958 54 339 346. SYMONS METAL SOLUTIONS 115 solvent shell will not be precisely that expected for an el cavity in the given solvent since it is predetermined by the size of the iodide ion. Nevertheless it should provide a potential well capable of binding the electron and the fact that an excited state exists for solvated iodide ions means that such preformed cavities can retain electrons. Since the formation of the cavity in the first instance is primarily a function of the negative charge this means at least that the el cavity model is reasonable. Of the two models proposed for this excited state the square-well seems to fit the data better than the hydrogen atom-like s-orbital However it is most probable that neither model does justice to the problem and the same is true of the simplified description of el ~avities.~l**~ These mathematical descriptions are however fairly simple and provided that the experimental data are reasonably well explained they are of considerable utility. Thanks are offered to Dr. H. C. Clark Ik. W. Doyle and Dr. G. W. A. Fowles for enjoyable discussion.

ISSN:0009-2681    

DOI:10.1039/QR9591300099

出版商:RSC    

年代:1959

数据来源: RSC

摘要:

ORGANOSILYLMETALLIC COMPOUNDS THEIR FORMATION AND REACTIONS AND COMPARISON WITH RELATED TYPES By DIETMAR WITTENBERG (BADISCHE ANILIN- UND SODA-FABRIK A. G. HAUPTLABORATORIUM LUDWIGSHAFEN/MAIN GERMANY) and HENRY GILMAN (IOWA STATE COLLEGE AMES IOWA U.S.A.) SOME aspects of the chemistry of organosilicon compounds have recently been the subject of several monographs and reviews.l The field of silyl- metallic compounds however has been mentioned only briefly and in few of these articles.lg$h Although the formation of triphenylsilyl-lithium as a reaction inter- mediate was first reported2 in 1933 useful methods for the preparation of organosilylmetallic compounds were not published until 195 1. Their chemistry and reactions have been intensively and extensively studied in subsequent years.Compounds of the general formula RR’R”SiM where R R’ and R” are organic radicals or hydrogen and M is a metal fallinto the category of this Review. Emphasis will be placed on the silyl-alkali-metal compounds. These highly active reagents-triphenylsilyl-lithium for example-can be easily prepared in solution. Comparable in reactivity with correspond- ing organometallic compounds such as Grignard reagents organolithium compounds or triphenylmethyl-sodium silylmetallic compounds often show dissimilarities in their mode of reaction. Their reaction with a variety of functional groups is potentially a source of new types of organosilicon compounds. Reactions are not necessarily restricted to the formation of silicon-carbon bonds it is possible to form bonds between silicon and many other elements.1. Preparation of silylmetallic compounds (a) Early Unsuccessful Attempts.-“Even after a very short experience it was evident that corresponding derivatives of the two elements in ‘(a) H. W. Post “Silicones and Other Organic Silicon Compounds,” Rheinhold Publ. Corp. New York N.Y. 1949; (b) T 6. Rochow “An Introduction to the Chemistry of Silicones,” 2nd edn. John Wiley & Sons Tnc. New York N.Y. 1951 ; (c) A. Ya. Yakubovich and V. A. Ginsburg Uspekhi Khim. 1949 18 46; ( d ) K. A. Andrianov A. A. Zhdanov S. A. Golubstsov and M. V. Sobolevskii ibid. 1949 18 145; (e) K. A. Andrianov and A. A. Zhdanov ibid. 1952,21 207; (f) H. Gilman and G. E. Dunn Chem. Rev. 1953 52 77; (g) A. G. Brook Chemistry in Canada Sept. 1955,43; (h) A. G. MacDiarmid quart..&^^^. 1956,10,208; (i) P.D. George M. Prober and J. R. Elliot Chem. Rev. 1956 56 1065; ( j ) I. J. Wilk J. Chem. Educ. 1957 34 463; (k) A. D. Petrov and V. F. Mironov “Darstellung und Eigenschaften von Siliziumkohlenwasserstoffen,” Akademie-Verlag Berlin 1955 ; ( I ) M. Kumada YGki G6sei Kagaku Kyokai Shi 1958 16 379. C. A. Kraus and H. Eatough J. Amer. Chem. SOC. 1933,55 5008. 116 WITTENBERG AND GILMAN ORGANOSILYLMETALLIC COMPOUNDS 1 17 question showed very considerable differences in their chemical proper- ties.” This viewpoint of Kipping3 after forty years of research in silicon chemistry seemed to be confirmed by unsuccessful attempts to prepare silylmetallic compounds employing methods which are useful for organo- metallic compounds. The reaction of metals with organic halides is perhaps the most funda- mental one for the preparation of organometallic compounds.* Trialkyl- and triaryl-silyl halides have been allowed to react with alkali metals in non-polar solvents or without solvent.Although the process may involve the formation of a silylmetallic intermediate the end products have been disilanes :5 2R,SiCI + 2Na -+ R,Si.SiR + 2NaCI Numerous examples of metallations of the type RH + R’Li -+ RLi + R‘H are known in carbon chemistry.6 Corresponding reactions of Si-H types with organolithium compounds resulted in the displacement of a hydride ion and the formation of a new silicon-carbon bond:’ R,SiH + R’Li -+ R,SiR’ + LiH Whereas triphenylmethylpotassium is formed from triphenylmethane and potassium amidey8 related reactions of triphenylsilane give silyl- amines Ph,CH + KNH -+ Ph,CK + NH Ph,SiH + LiNR -+ Ph,Si*NR + LiH Halogen-metal interconversion is another method of choice for the Silyl halides generally yield coupling products when allowed to react preparation of many organoalkali-metal comp~unds.~J* with organoalkali-metal compounds :11 R X + R’Li -+ RLi + R’X R,SiX + R’Li -+ R,SiR’ + LiX In special cases where halogen-metal interconversion occurs in silicon a F.S. Kipping Proc. Roy. SOC. 1937 A 159 139. “(a) W. Schlenk J. Renning and G. Racky Ber. 191 1 44 1178 ; (b) F. S. Kipping Proc. Chem. SOC. 1911,27 143; (c) H. Gilman and G. E. Dunn J. Amer. Chem. SUC. 1951 73 5077; (d) M. G. Voronkov and Yu. I. Khubodin Zhur. obschchei Khim. 1956,26 586; (e) M. P. Brown and C. *. A. Fowles J. 1958 2811. See the reviews (a) H. Gilman and A. Morton “Organic Reactions,” John Wiley & Sons Inc.New York N.Y. 1954 Vol. VIII p. 258ff.; (b) R. A. Benkeser D. S. Foster D. M. Sauve and J. F. Nobis Chem. Rev. 1957,57 867. ’ (a) H. Gilman and S. P. Massie J. Amer. Chem. SOC. 1946,68,1128 ; (b) R. N. Meals ibid. p. 1880; (c) R. A. Benkeser and F. J. Riel ibid. 1951,73 3472; ( d ) H. Gilman and E. Zuech ibid. 1957,79,4560; and referer.?ees cited therein. 13 C. A. Kraus and R. Rosen J. Amer. Chem. SOC. 1925 47 2739. @ H. Gilman B. Hofferth H. W. Melvin and G. E. Dunn ibid. 1950; 72,5767. lo R. A. Jones and H. Gilman in “Organic Reactions,” John Wiley & Sons Inc. l1 See ref. l(b) sections 3 and 4. R. G. Jones and H. Gilman Chem. Rev. 1954,54 835. New York N.Y. 1951 Vol. VI p. 339ff. 118 QUARTERLY REVIEWS chemistry as in the reaction between silyl halides and the stilbene- dilithium adduct,12 the Si-Li intermediate couples immediately with the silyl halide to give a disilane as the final product R,SiCI + (CHPhLi) -f PhCH=CHPh + LiCl + R,SiLi -+ R,Si*SiRa R,SiCI (b) Metal Cleavage of Non-functional Silanes and Disilanes.-In 195 1 Benkeser and Severson13 reported the cleavage of (aa-dimethylbenzy1)- triphenylsilane with sodium-potassium alloy in ether to give a mixture of triphenylsilylpotassium and (aa-dimethylbenzy1)phenylpotassium Similarly? ap-bistriphenylsilylcumene was cleaved by the alloy.13 In the same year it was found14 that triphenylsilylpotassium could be prepared free from other products by the cleavage of hexaphenyldisilane with sodium-potassium alloy in ether.The yellow-brown silylpotassium re- agent is obtained as a suspension Ph,Si.SiPh + 2K -+ 2Ph,SiK It was found advantageous to remove excess of alloy from the mixture by ama1garnati0n.l~ Although cleavage occurred neither with sodium disper- sion in xylene or dioxan,l4 nor with sodium amalgam or lithium in ethyl ether hexaphenyldisilane was cleaved by sodium-potassium alloy in hot n-butyl ether14 or though very slowly in benzene or light petr01eum.l~~ Triphenylsilylrubidium and triphenylsilylcasium have been prepared by cleavage of the disilane with rubidium and casium respectively in ethyl ether.l5 The first solutions of silylmetallic compounds were obtained in ethylene glycol dimethyl ether. It was found that in this solvent hexaphenyldisilane is cleaved not only by sodium-potassium alloy but also by sodium and lithium.16 Unfortunately these solutions are not stable and have to be used immediately after their preparation.Triphenylsilyl-lithium in ethylene glycol dimethyl ether gave a negative Colour Test 117 after one hour's refluxing. The ether cleavage product was methyltriphenylsilane.18 Although only triaryl types were accessible before tetrahydrofuran came into use as solvent the latter made possible the preparation of silylalkali- metal reagents containing aliphatic as well as aromatic groups. Triphenyl- sil yl-lithium,l gmet hyldiphenylsilyl-lithium,~Qdimethylphenylsilyl-li t hium,lS l2 M. V. George D. Wittenberg and H. Gilman J. Amer. Chem. Soc. 1959,81,361. l3 R. A. Benkeser and R. G. Severson ibid. 1951,73 1424. l4 H. Gilman and T. C. Wu ibid. p. 4031. 15(ca) H. Gilman T. C.Wu H. A. Hartzfeld G. A. Guter A. G. Smith J. J. Goodman and S. PI. Eidt ibid. 1952,74 561; (6) H. Gilman and T. C. Wu J. Org. Chem. 1953 18 753. la A. C. Brook and H. Gilman J . Amer. Chem. SOC. 1954,76,278. l7 H. Gilman and F. Schulze ibid. 1925 47 2002; G. Wittig Angew. Chem. 1940 53 243. l8 D. Wittenberg D. Akoi and H. Gilman J. Amer. Chem. SOC. 1958,80,5933. l9 H. Gilman and G. D. Lichtenwalter ibid. p. 608. Ph,Si.CMe,Ph + 2K -f Ph,SiK + PhMe,CK WITTENBERG AND GILMAN ORGANOSILYLMETALLIC COMPOUNDS 1 19 and 5-lithio-5-methyldibenzosilole,20 prepared from the corresponding disilanes in tetrahydrofuran are examples of this type. These solutions are fairly stable at room temperature for long times providing a distinct advantage over the use of ethylene glycol dimethyl ether as solvent.Although the cleavage of hexaphenyldisilane by lithium in tetrahydro- pyran is rather slow at room temperature,1ap21 this solvent may be advan- tageous in some cases because of its fair stability towards cleavage by triphenylsilyl-lithium even at high temperatures.18 Other solvents in which hexaphenyldisilane has been cleaved include dioxanla and pyridine.22 Cleavage was also accomplished by sodium in liquid ammonia.23 In each of these cases secondary reactions occurred between the Si-Li reagents and the solvent. [For a detailed discussion see sections 2(e) and 2(f) below.] TABLE 1. Silylalkali-metal compounds prepared by metal cleavage of symmetrical disilanes. Compound Solvent' Derivative Yield ( %) Ref. made with Ph,SiK Et2O PhBr 40-96 13 14 1st Ph,SiK Et2O Me,SiCl 45-75 13 14 15 Ph,SiK Bu~O Me,SiCI 47 14 p-TolPh2SiK Et2O p-To12PhSiCl 52 24 p-TolsSiK EtZO Ph,SiCl 26 24 p-TolSiK Et2O co2 56 25 Ph,SiRb EtZO Me3SiC1 48 15b P h S i C s Et2O Me3SiC1 26 15b Ph3SiNa GDME Me,SiCl 68 16 Ph3SiLi GDME Me,SiCl 72 16 Ph,SiK GDME MesSiCl 77 16 Ph,SiLi THF Me,SiC1 79 19 Ph2MeSiLi THF Me,SiCI 74 19 PhMe,SiLi THF Me,SiCl 47 19 VHF Me2S04 54 20 Q-D 2 Me Li * GDME = Ethylene glycol dimethyl ether; THF = tetrahydrofuran.t Also H. Gilman G. D. Lichtenwalter and D. J. Peterson unpublished studies. The metal cleavage of a hexa-alkyldisilane has not been accomplished to date and no stable solutions of trialkylsilylalkali-metal compounds useful for synthetic purposes have been prepared. Although hexa-alkyl compounds of tin and lead can be split by alkali metals to give the alkali- metal ~ a l t s ~ ~ $ ~ ~ the corresponding compounds of germanium and silicon Po H.Gilman and R. D. Gorsich ibid. p 3243. z1 M. V. George unpublished studies. 22 D. Wittenberg and H. Gilman Chem and Ind. 1958 390. 23 T. C. Wu and H. Gilman J. Org. Chem. 1958,23,913. 24 H. Gilman and T. C. Wu J. Amer. Chem. Soc. 1953,75 3762. 2K A. G. Brook and R. J. Mauris ibid. 1957,79 971. ** C. A. Kraus and W. V. Sessions ibid. 1925 47 2361. 120 QUARTERLY REVIEWS seem to be truly resistant to cleavage by alkali metals. Thus hexaetbyl- disilane is not cleaved by lithium in eth~larnine~~ or by sodium-potassium alloy (either alone or in a number of ~ o l v e n t s ) ~ ~ ~ ~ ~ ~ ~ ~ although the disilane is readily cleaved by halogens.28 No reaction was observed between triethylphenylsilane and sodium in liquid ammonia.27 Benkeser and Severson13 reported no reaction between (aa-dimethylbenzy1)trimethyl- silane and sodium-potassium alloy in ether.Trimethylphenylsilane has been cleaved at -50" by sodium-potassium alloy in tetrahydrofuran. Although it was possible to characterise phenylpotassium in the reaction mixture by preparation of derivatives no evidence was found for the presence of trimethylsilylpotassium.29 Some of the reactions achieved by such methods are listed in Table 1. Another possible route for the preparation of trialkylsilyl-metal com- pounds involved cleavage of unsymmetrical disilanes. Trimethylphenyl- silane in low yield together with tetraphenylsilane were isolated from the cleavage of 1 1 1 -trimethyl-2,2,2-triphenyldisilane by sodium-potassium alloy after preparation of these derivatives with bromobenzene 28 Me,Si-SiPh + 2K 3 Ph,SiK + Me,SiK __lf + PhBr Ph,Si + Me,SiPh Cleavage of l,l,l-triethyl-2,2,2-triphenyldisilane by lithium in tetra- hydrofuran followed by acid-hydrolysis gave low yields of triethylsilane and hexaethyldisilane in addition to triphenyl~ilane,~~ indicating that triethylsilyl-lithium was one of the intermediates.(c) Metal Cleavage of Various Functional Si1anes.-While disilanes are usually prepared by reactions of the corresponding chlorosilanes with alkali rnetal~,~ and silylmetallic compounds by metal-cleavage of disilanes methods have been investigated to combine these two steps by a proper choice of solvent and metal. Chlorotriphenylsilane was converted into triphenylsilylpotassium by the use of sodium-potassium alloy in ether,l*~I~~ whereas in xylene the reaction stops at the disilane stage.15b Similarly tri-p-tolylsilylpotassium was prepared from chlorotri-p-tolylsilane.25 By using lithium in tetrahydrofuran it is possible to prepare several silyl- lithium compounds directly from the corresponding chlorosilanes.Triphenylsilyl-lithium,31 methyldiphenylsilyl-lithium,16 and 1 -phenyl- 1- silacyclohexyl-lithium32 were formed by this method. It is interesting that trialkylchlorosilanes are converted into the disilane by lithium in tetra- hydrofuran ;33 no further cleavage has been observed. Since it was found33 that silyl-lithium compounds couple readily with 87 C. A. Kraus and W. K. Nelson ibid. 1934 56 195. 28 H. Gilman R. K. Ingham and A. G. Smith J. Org. Chem.1953,18,1743. M. B. Hughes unpublished studies. a0 G. D. Lichtenwalter and D. Wittenberg unpublished studies. H. Gilman D. J. Peterson and D. Wittenberg Chem. mdlitd. 1958 1479. s2 H. Gilman G. D Lichtenwalter and D. J. Peterson unpublished studies. D. Wittenberg and G. D. Lichtenwalter unpublished studies. WITTENBERG AND GILMAN ORGANOSILYLMETALLIC COMPOUNDS 12 1 chlorosilanes even at -70° the reaction is thought to proceed in two steps; first the formation of a disilane and secondly its cleavage 2Ph,SiCI + 2Li -+ 2LiCI + Ph,Si*SiPh - + 2Ph,SiLi TABLE 2. SilyIa/kali-metal compounds prepared from various organosilicon compounds. + 2Li Compound Prepared from Solvent* Derivative Yield Ref. prepared with (%) Ph,SiK p-To1,SiK Ph,SiLi Ph,MeSiLi PhEt,SiLi Ph,SiK Ph,SiK Ph,SiK (p-Ph*C,H,),SiK Ph,SiLi Ph,MeSiLi (Ph,Ge),SiLi Et,SiLi PhsSiK Ph3SiK Ph,SiK Ph,SiK Ph3SiLi PhsSiK Ph,SiK Ph,SiCl p-To1,SiCl Ph,SiCl Ph2MeSiC1 Ph.Et2SiCl Ph,Si*OEt Ph3Si*OMe Ph,SiH (p-Ph-CeH4),SiH Ph,SiH Ph2MeSiCl (Ph,Ge),SiH Et,Si*GePh Ph,Si*GePh Ph,Si*CMe,Ph Ph4Si Ph,Si*SiMe Ph,SkSiEt PhaSiCPh Ph4Si Et2O Et20 THF THF THF Et2O Et2O Et2O Et20 THF THF NH2Et Et,O Et20 Et20 THF Et20 NH2Et Et2O EtZO ~~ Me,SiCl 67 15b co2 69 25 Me,SiCl 83 31 Furan 27 18 Ph,SiCl 36 31 Ph,SiCl 34 32 PhBr 24 15b PhBr 79 34 PhBr 67 34 p-PhCsH,Br 48 34 Me2S04 63 t Ph,SiCl 66 t EtBr 63 + co 86 § co2 70 7 PhBr 39 13 Ph,SiCl 70 1% Me3SiCl 79 t PhBr 50 15b 28 PhBr 39 28 EtBr - i 7 * THF = Tetrahydrofuran.t R. D. Gorsich unpublished studies. 3 J. G. Milligan Ph.D. Thesis Brown University Providence R.I. 1934; J. G.9 H. Gilman and C. W. Gerow ibid. 1956 78 5823. 7 A. G. Brooke H. Gilman and L. S. Miller ibid 1953,75,4579. The use of chlorosilanes may be advantageous in many cases (cf. Table 2); however the yields of the silylmetallic reagents often appear to be slightly lower than from the corresponding disilanes. Furthermore chlorosilanes though commercially available require special precautions owing to their ease of hydrolysis. Alkoxysilanes such as eth~xytriphenylsilane~~~ and methoxytriphenyl- ~ i l a n e ~ ~ have also been used for the preparation of triphenylsilylpotassium with sodium-potassium alloy in ethyl ether as the cleaving agent. This reaction also may involve the intermediate formation of hexaphenyldi- silane. The reaction of N-ethyl(triphenylsily1)amine with lithium in ethylamine had been reported to yield solutions of triphenylsilyl-lithium.2 Later 3* R.A. Benkeser H. Landesman and D. J. Foster,J. Amer. Chem. Soc. 1952,74,648. Milligan and C. A. Kraus J. Amer. Chem. Suc. 1950,72,5297. 122 QUARTERLY REVIEWS obtained the same “red solutions” but were able to show that reduction of phenyl groups occurs rather than formation of Si-Li bonds. Various silicon hydrides on treatment with sodium-potassium alloy in ethyl ether or with lithium in tetrahydrofuran yielded the corresponding silylalkali-metal compounds. Tris-p-biphenylylsilylpotassium,34 triphenyl- sil y lpo t a ~ s i u m ~ ~ trip hen y lsil y 1-lit hium 36 and met hy ldip heny lsil y 1-lit h i ~ m ~ ~ have been prepared by this method. The process actually appears to involve cleavage of aryl groups from the silicon atom; the arylalkali-metal compound then reacts with the silicon hydride to form a non-functional silane which in turn is again cleaved by the metal R3SiH - RaSiHLi + RLi - RSiHLi + 2RLi + 2Li + 2Li RLi + R3SiH + R4Si R,Si + 2Li -+ R,SiLi + RLi n R,Si H Li + nLiH + (R,Si) nRSiHLi + nLiH + (RSiLi) G o r s i ~ h ~ ~ isolated tetraphenylsilane after reaction of triphenylsilane with lithium in tetrahydrofuran and showed that tetraphenylsilane in turn can be cleaved to a mixture of triphenylsilyl-lithium and phenyl- lithium.It was also demonstrated that lithium in tetrahydrofuran splits off two biphenyl groups from tri-o-biphenylyl~ilane.~~ By a similar mechanism tetraphenylsilane appears to be formed from phenylsilane as well as from diphenylsilane and from triphenylsilane with sodium- potassium alloy in ethe~-.~~J’ Although in all of these reactions the silyl-metallic compounds contain- ing SiH groups were expected as by-products no evidence for their presence was found.Instead a polymer and a metal hydride were isolated nR,SiHLi -+ (R,Si) + nLiH The reaction of triphenylsilane with lithium in ether seems to take a different course the main product being hexaphenyldi~ilane.~~ Triethyl- silane apparently does not react under similar condition^.^^ In liquid ammonia bistriphenylsilylamine was formed from sodium and triphenyl- silane.38 Other reactions of alkali metals with various silicon hydrides and with chlorosilanes which may involve the formation of silyhnetallic compounds as intermediates have been described by Kipping and his collaborator^^^ and by Benkeser and Triphenylgermyltriphenylsilane on treatment with sodium-potassium 35 R.A. Benkeser R. E. Robinson and H. Landesman ibid. p. 5699. 36 R. D. Gorsich unpublished studies. R. A. Benkeser and D. J. Foster J. Amer. Chem. SOC. 1952,74,4200. 38 H. H. Reynolds L. A. Bigelow and C. H. Kraus ibid. 1929,51,3067. 3Q A. R. Stele and F. S. Kipping J. 1928,1431 ; F. S. Kipping J. 1923 123 2590; 1924,125,2291 ; F. S. Kipping and J . E. Sands J. 1921 119,830; F. S. Kipping and A. G. Murray J. 1929,360; F. S. fipping and J. F. Short,J. 1930,1029. 40 R. A. Benkeser and D. J. Foster J. Amer. Chem. Soc. 1952,74,5314. WITTENBERG AND GILMAN ORGANOSILYLMETALLIC COMPOUNDS 123 alloy afforded a mixture of triphenylgermylpotassium and triphenylsilyl- potassium *l Ph,Ge.SiPh + 2K -+ Ph3GeK + Ph,SiK An analogous cleavage of triphenyl(triphenylmethy1)silane has been de~cribed.~~ Lithium in ethylamine has been used to split triethyl(tripheny1- germyl)~ilane,~' and tris(triphenylgermy1)silyl-lithium was obtained from tristriphenylgermylsilane and lithium in ethylamine ; however the latter silylmetallic compound was described as unstable in this ~ o l v e n t .~ ~ ~ ~ (a) Silylalkali-metal Compounds containing more than One Silicon Atom.-An early attempt to obtain pentaphenyldisilanyl-lithium from chloropentaphenyldisilane and lithium in diethyl ether or di-n-butyl ether was unsuccessful~5 while with sodium in refluxing xylene decaphenyl- tetrasilane was formed.45 Cleavage of octaphenyltrisilane by lithium in tetrahydrofuran gave a mixture of pentaphenyldisilanyl-lithium and triphenylsilyl-lithium,*6 identified after acid-hydrolysis as pentaphenyldisilane and triphenylsilane.Cleavage of decaphenyltetrasilane by lithium in tetrahydrofuran gave a mixture of triphenylsilyl-lithium pentaphenyldisilanyl-lithium and hepta- phenyltrisilanyl-lithium.46 It is noteworthy that apparently no diphenyl- silylenedilithium was formed in these cleavages since diphenylsilane was not isolated after acid-hydrolysis. Ph,Si.SiPh,-SiPh + 2Li -+ Ph,Si*SiPh,Li + Ph3SiLi Ph,Si.SiPh,-SiPh,-SiPh + 2Li -+ 2Ph3Si.SiPh,Li 3 Ph,Si.SiPh,.SiPh,Li + Ph,SiLi (e) Interconversions.-Some coupling .reactions of silyl-lithium com- pounds with chlorosilanes produced the desired unsymmetrical disilanes in very low yields. The symmetrical disilanes were formed as the main product^.^^,^^ Subsequent have indicated that cleavages of disilanes by silyl-lithium compounds were involved.Hexaphenyldisilane although very sparingly soluble was readily cleaved by dimethylphenylsilyl-lithium in tetrahydrofuran to give triphenylsilyl-lithium and 1,1,2,2-tetramethyl- 1,2-diphenyldisilane PhMe,SiLi + Ph,Si.SiPh -+ Ph,Si*SiPhMe + Ph,SiLi PhMe,SiLi + Ph,Si.SiPhMe -+ PhMe,Si.SiPhMe + Ph,SiLi Similarly dimethylphenylsilyl-lithium cleaved 1,2-dimethyl-l 1,2,2-tetra- phenyldisilane to yield methyldiphenylsilyl-lithium and 1,1,2,2-tetramethyl- 1,2-diphenyldisilane. Triphenylsilyl-lithium and 1,2-dimethyl- 1,1,2,2-tetra- 41 H. Gilman and C. W. Gerow ibid. 1956,78 5823. 42 A. G. Brook H. Gilman and L. S. Miller ibid. 1953,75,4579. 43 J. G. Milligan Ph.D. Thesis Brown Umversity Providence R.I.1934. 44 J. G. MiUigan and C. A. Kraus J. Amer. Chem. SOC. 1950,72 5279. 45 T. C. Wu unpublished studies. 46 D. Wittenberg M. V. George and H. Gilman J. Amer. Chem. Soc. 1959,91 in the 47 For a detailed discussion see section 2(h) below. press. 124 QUARTERLY REVIEWS phenyldisilane resulted from the reaction of methyldiphenylsilyl-lithium with hexaphenyldisilane ZPhMe,SiLi + Ph,MeSi-SiPh,Me + ZPh,MeSiLi + PhMe,Si*SiPhMe ZPh,MeSiLi + Ph,Si.SiPh -+ 2Ph,SiLi + Ph,MeSi-SiPh,Me When 1 l,l-triethyl-2,2,2-triphenyldisilane was allowed to react with lithium in tetrahydrofuran triethylsilane and triphenylsilane in addition to a small amount of hexaethyldisilane were isolated after acid-hydroly~is.~~ The isolation of the latter indicates the formation of triethylsilyl-lithium which by cleavage of 1 1 l-triethy1-2,2,2-triphenyldisilane gives triphenyl- silyl-lithium and hexaethyldisilane Ph,Si.SiEt + 2Li -+ Ph,SiLi + Et,SiLi Et,SiLi + Ph,Si.SiEt -+ Ph,SiLi + Et,Si.SiEt A similar cleavage of pentaphenyldisilane by triphenylsilyl-lithium appeared to be an attractive route to the unknown species diphenylsilyl- lithium.Although hexaphenyldisilane was actually found in this reaction,46 diphenylsilyl-lithium seemed to decompose into lithium hydride and a diphenylsilylene polymer under the experimental conditions Ph,SiLi + Ph,Si.SiPh,H nPh,SiH Li -+ nLiH + (Ph,Si) Ph,Si.SiPh + Ph,SiHLi On the assumption that the silicon-silicon bond strength remains nearly constant with varying numbers of alkyl and aryl substituents the cleavage experiments suggest that the reactivity of silyl-lithium compounds increases with an increasing number of alkyl groups on the silicon atom Ph,SiLi < Ph,MeSiLi < PhMe,SiLi < Et,SiLi 2.Reactions of silylalkali-metal compounds (a) With Hydrocarbons.-Shortly after the first successful preparation of triphenylsilylp~tassium,~~J~ it was observed that this reagent adds to the olefinic linkage of trans-~tilbene.~~ After hydrolysis 1,2-diphenyl- 1-triphenylsilylethane was isolated 0 Ph,SiK + PhCH=CHPh -+ Ph,SiCHPh.CH,Ph H,O However it was found that the reaction of triphenylsilyl-lithium with trans-stilbene in ethylene glycol dimethyl ether did not yield this product; instead a number of more complex compounds including 1 ,Zdiphenyl- 1,2-bistriphenylsilylethane and 1,2,3,4-tetraphenyl-1 -triphenylsilylbutane were formed.49 The results suggest that the addition product formed initially behaved as an active organometallic reagent which could react with other stilbene molecules as has been observed in other ca~es.~O~~l 4a H.Gilman and T. C. Wu J. Amer. Chem. SOC. 1953,75,234. 48 A. G. Brook K. M. Tai and H. Gilman ibid. 1955,77 6219. 50 K. Ziegler H. Grimm and R. Willer Annalen 1939 542 90. 51 K. Ziegler and H. Kleiner ibid. 1929 473 57. WITTENBERG AND GILMAN ORGANOSILYLMETALLIC COMPOUNDS 125 Treatment of tolane with triphenylsilyl-lithium yielded in addition to polymeric material some 1,2-diphenyl- 1,2-bi~triphenylsilylethane.~~ Formation of this product suggests that two moles of the Si-Li reagent added across the triple bond of tolane 2Ph3SiLi + PhC= CPh + Ph,Si.CHPh.CHPh.SiPh HZO Triphenylsilyl-lithium in tetrahydrofuran and triphenylsilylpotassium in ethyl ether were also found to add to 1,l-diphenylethylene and to triphenyl- ethylene.52 Attempts to add the same reagents to tetraphenylethylene and to a variety of aliphatic and cycloalkenes were however U ~ S U C C ~ S S ~ U Recently it was found that triphenylsilyl-lithium adds smoothly to anthracene to give 9,lO-dihydro-9-lithio- 1 0-triphenyl~ilylanthracene.~~ Triphenylsilyl-potassium -sodium and -lithium were found to metallate triarylrnethane~.~~ Whereas tetraphenylmethane was unaffected by tri- phenyl~ilyl-lithium,~~ fluorene was readily metallated to yield triphenyl- silane and 9-fl~orenyl-lithium.~~ If the silylmetallic reagent is present in excess in metallations it slowly reacts further with the triarylsilane to yield a tetra-aryl~ilane.~~ (b) With Organic Halides.-In reactions of silylalkali-metal compounds with organic halides three different modes of interaction have been observed direct coupling ; halogen-metal interconversion; and dehydro- halogenation.(i) Alkyl and aralkyl halides. Although no alkyl fluorides have yet been investigated reactions of triphenylsilyl-lithium with butyl and ally1 chloride give high yields of the coupling products butyltriphenylsilane and allyltriphenylsilane re~pectively.~~ With alkyl bromides fair yields of Ph,SiLi + BuCl -+ Ph,SiBu + LiCl coupling products have been obtained by using triphenylsilylpotassium in ether. On the other hand alkyl bromides react with triphenylsilyl- lithium mainly by halogen-metal interconversion to give hexaphenyl- di~ilane.~~ On reaction of triphenylsilyl-lithium with 1,4-dibromobutane only 3 % of tetramethylenebistriphenylsilane was isolated hexaphenyldi- silane being the main product.56 Methyl iodide and triphenylsilylpotassium afforded 33 % of methyltriphenylsilane and 52 % of he~aphenyldisilane.~~ Low yields of isopropyltriphenylsilane were obtained from triphenyl- silyl-lithium and isopropyl chloride ;55 none was isolated however when isopropyl bromide was used and no cyclohexyltriphenylsilane was formed from cyclohexyl bromide and triphenylsilylpotassium.5 The hexaphenyl- disilane is the result of a halogen-metal interconversion when t-butyf 63 T.C. Wu and D. Wittenberg unpublished studies. 6a 0. L. Mans unpublished studies. 64 A. G. Brook and H. Gilman J. Amer.Chem. Soc. 1954,76,2338. 65 H. Gilman and D. Aoki,. J. Org. Chem. 1959,24,426. ti7 A. G. Brook and S. Wolfe ibid. 1957,79 1431. D. Wittenberg and H. Gilman J. Amer. Chem. Soc. 1958 80 2677. I26 QUARTERLY REVIEWS bromide and triphenylsilyl-lithium are allowed to interact. In contrast t-butyl chloride gives mainly triphenylsilane under the same conditions ;55 this process probably involves dehydrohalogenation Ph,SiLi + Me,CCI -f Ph,SiH + Me,C=CH + LiCl The reaction of benzyl chloride with triphenylsilylpotassium gave fair yields of benzyltriphenylsilane in addition to bibenzyl and hexaphenyldi- silane 15y5 Ph,SiK + PhCH,CI 4 Ph,SiCI + Ph.CH,K Ph.CH2K + Ph,SiK + Ph-CH,CI 3 PhCH,-CH,Ph + Ph,SiCH,Ph + Ph,Si-SiPh Attempts to prepare triphenyltriphenylmethylsilane from triphenyl- methyl chloride and triphenylsil ylpotassium yielded liexaphenyldisilane and hexaphenylethane instead.42 This again demonstrates the occurrence of halogen-metal interconversion.The use of alkyl sulphates instead of alkyl halides in coupling reactions with silylalkali-metal compounds greatly diminishes or entirely eliminates any side halogen-metal interconversion-type reactions. For example when 5-lithio-5-methyldibenzosilsle20 was treated with methyl sulphate only 5,5-dimethyldibenzosilole was isolated (ii) Aryl halides. Bromobenzene has been widely used for preparing derivatives from triphenylsilylpotassium;13~14J5~34y55~57 however tetra- phenylsilane has rarely been obtained in yields exceeding 50 %. Slightly higher yields have been obtained by using chlorobenzene i n ~ t e a d .~ ~ . ~ ~ Fluorobenzene reacts extremely slowly,15 possibly because of metallation in the ortho-position. Extensive and excellent studies by Brook and W ~ l f e ~ ~ have demonstrated that with aryl bromides and chlorides halogen-metal interconversion can account for the side-products which were isolated. By adding triphenylsilylpotassium to a mixture of bromo- benzene and benzophenone they were able to capture phenylpotassium as trip h en y lme t hanol Ph,SiK f- PhBr 4 Ph,SiBr $- PhK -+ Ph,C.OH As a method of choice 2-triphenylsilyldibenzothiophen and 9-ethyl-3- triphenylsilylcarbazole have been obtained from the corresponding bromides and triphenylsilylpotassium.58 Similarly tetra-p-biphenylyl- silane was prepared from 4-bromobiphenyl and tri-p-biphenylylsilyl- potassium,34 while p-fluorophenyltriphenylsilane was obtained from p - bromofluorobenzene and triphenylsilyl-lithium.55 Brook and Wolfes7 were able to show that the yields and nature of the products from silylmetallic compounds and organic halides are greatly influenced by the mode of addition.Addition of the silylmetallic compound to the halide the reverse addition and slow simultaneous addition of the reagents have been employed. It was found that simultaneous addition gave the highest yields of coupling products. The yields of interconversion products seem to increase while the yields of coupling products decrease 58 R. Meen and H. Gilman J. Org. Chem. 1955,20,73. WITTENBERG AND GILMAN ORGANOSILYLMETALLIC COMPOUNDS 127 when the silylpotassium reagent in ether is replaced by the corresponding silyl-lithium reagent in tetrahydrofuran.(c) With Aldehydes and Ketones.-In general silylmetallic compounds react with aliphatic aldehydes and ketones in the same manner as do organolithiurn and Grignard reagents to give silyl-substituted alcohols. Triphenylsilylmethanol results froin the reaction of triphenylsilylpotassium with formaldc hyde ; 1 -trip henylsil ylet hano16 O and 1 -trip henylsilyl- propan- 1-01~~ were formed on reaction of the corresponding silyl-lithium reagent with acetaldehyde and propionaldehyde respectively. Similarly triphenylsilyl-lithium was shown to add normally to a ~ e t o n e ~ ~ $ ~ ~ cyclohe~anone,~~ nonadecan-2-one octadecan-3-one penta- decan-8-one and tricosan-l2-0ne,~* e.2. Ph,SiLi + COMe -+ Ph,SiCMe,-OH H 2 0 An important side-reaction is the abstraction of an acidic hydrogen atom from the aldehyde or ketone by the silyl-lithium compound to give the corresponding silicon hydride as b y - p r o d ~ c t .~ ~ j ~ ~ In one case the yield of triphenylsilane from the reaction of triphenyl- silyl-lithium with dibenzyl ketone was 78 %. None of the a-silyl-alcohol was Rather surprising results were obtained from reactions involving silylalkali-metal compounds with aromatic ketones. Triphenylsilyl- potassium was found to add to benzophenone to give after hydrolysis diphenylmethoxytriphenylsilane rather than the expected diphenyltri- phenyl~ilylmethanol.~~ Since silylmetallic compounds add normally to aliphatic ketones and since the closely related triphenylgermylmetallic compounds give normal addition products with ben~ophenone,~~?~~ it did not seem likely that inverse addition had taken place.It was postulated therefore that normal addition had occurred followed by rearrangement. Triphenylsilyl-lithium methyldiphenylsilyl-lithium and dimethylphenyl- silyl-lithium in tetrahydrofuran solution were also found to give the cor- responding alkoxysilanes instead of alcohols.68 The “rearrangement hypothesis” was supported by experiments by Brook,63 who succeeded in preparing diphenyltriphen ylsilylmethanol and showed that under mild basic conditions this compound rearranges to give diphenylmethoxytri- phenyl silane Ph,SiLi + Ph,&O -+ [Ph,Si.C(OLi)Ph,] -f Ph,Si.O.CPh,Li - Ph,Si.OCHPh H2O 58 H. Gilman and T. C. Wu J. Amer. Chem. SOC. 1954,76,2502. 6o D. Wittenberg and H. Gilman ibid. 1958 80 4529. ” A. G.Brook C. M. Warner and M. E. McGriskin ibid. 1959 81,981. 6z H. Gilman and D. J. Peterson J. Org. Chew. 1958 23 1895. 63 A. G. Rrook J. Amer. Chent. SOC. 1958 80 1886. 64 H. Gilman and G. D. Lichtenwalter ibid. p. 2680. 65 H. Gilman and T. C. Wu ibid. 1953 75 2935. 66 H. Gilrnan and C. W. Gerow ibid. 1955 77 5740. A. G. Brook and N. V. Schwartz unpublished studies. H. Gilman and G. D. Lichtenwalter J. Arner. Chem. Soc. 1958 80 607. 128 QUARTERLY REVIEWS It is noteworthy that in carbon chemistry the corresponding “Wittig rea~rangement”~~-~~ takes a c mplet y different course ; phenyl-substituted ethers rearrange upon metallation to the corresponding alcohols Ph.CH,-OR --+ PhCHLiaOR -+ PhCHR-OLi -+ PhCHR-OH PhLi H,O that the deeply coloured benzophenone ketyl is formed in the reaction between triphenylmethylsodium and benzophenone.The question whether the deep blue colour of reaction mixtures of silylmetallic compounds with benzophenone is also due to a metal ketyl or to the C-M type addition product is still unanswered. Brook and his c o - ~ o r k e r s ~ ~ recently showed that under certain conditions triphenylsilylalkali-metal reagents react with 2 mols. of benzophenone ; after hydrolysis tetraphenyl-2-triphenylsiloxyethanol is isolated. The reaction might possibly involve an equilibrium of the following type Schlenk and Ochs Ph,Si-OCLiPh + Ph,CO + Ph,Si.O-CPh,CPh,*OLi The formation of small amounts of benzpinacol from the reactions of benzophenone with methyldiphenylsilyl-lithium or dimethylphenylsilyl- lithium may involve a similar intermediate.68 Carboxylation of a reaction mixture from triphenylsilylpotassium and benzophenone gave benzilic acid and triphen ylsilanol.74 The reaction may possibly involve a cleavage similar to that observed in carboxylation of a-lithiobenzyltriphenylsilane 75 Ph,Si.O-CPh,K ___j Ph,Si.OH + Ph,C(OH).CO,H CO,,H,O From the reaction of triphenylsilyl-lithium with benzaldehyde similar “abnormal” addition products were found. 76 In addition to benzyloxytri- phenylsilane hexaphenyldisilane 1,2-diphenylethane- 1,2-diol and 1,2- diphenyl-2-triphenylsiloxyethanol were isolated. (d) With Derivatives of Carboxylic Acid~.-Brook~~ found that triphenyl- silylpotassium and benzoyl chloride give a low yield of benzoyltriphenyl- silane the first reported a-silyl-ketone in addition to hexaphenyldisilane. From the reaction of triphenylsilyl-lithium with acetyl chloride acetyl- triphenylsilane was isolated in a low yield together with 1 ,1-bis(tripheny1- silyl)ethanol triphenylsilane and triphenyl-( 1 -triphenylsiloxye thyl)silane the rearrangement product of 1,l -bis(triphenylsilyl)ethanol:60 Ph,SILi + CH,.COCI -f Ph,Si.COCH - -+ + Ph,SiLi (Ph,Si),C(OH)CH - -+ Ph,Si.CH(CH,).OSiPh Rearr.6 9 G. Wittig and L. Lohmann Annalen 1942 550 260. 7 0 G. Wittig and W. Happe ibid. 1947 557 205. 71 G. Wittig and R. Clausnizer ibid. 1954 588 145. 72 G. Wittig and E. Stahnecker ibid. 1957 605 69. 73 W. Schlenk and R. Ochs Bet-. 1916 49 612. 7aA. G. Brook N. V. Schwartz M. E. McGriskin N. Wolfish and M. Gold 76 H. Gilman and H. Hartzfeld ibid. 1951 73 5878. 76 D. Wittenberg and T. C. Wu unpublished studies.77 A. G. Brook J . Amer. Chem. Soc. 1957,79,4373. unpublished studies. WITTENBERG AND GILMAN ORGANOSILYLMETALLIC COMPOUNDS 129 In this connection it should be pointed out that triphenylmethylsodium does not add to aliphatic acid chlorides; enolisation occurs with the quantitative formation of triphenylrnethane. 78 Acetic anhydride and ethyl acetate on treatment with triphenylsilyl- lithium gave results62 analogous to those observed in the acetyl chloride reaction. Triphenyl-(1-triphenylsiloxypropy1)silane was isolated when propionyl chloride was allowed to react with the same silyl-lithium reagent while phenylacetyl chloride was predominantly eno1ised.62 Acrylonitrile was polymerised by triphenylsilyl-lithium.62 Lithium cyanide and tetraphenylsilane were among the products ob- tained on reaction of benzonitrile with triphenylsilyl-lithium.The nitrile group in this reaction behaves as a pseudo-halide:62 PhCN + Ph,SiLi -+ Ph,Si + LiCN The reaction of methyl triphenylgermanecarboxylate with triphenyl- silyl-lithium gave an excellent yield of triphenylgermyltriphenylsilane 79 Ph,SiLi + Ph,Ge-C0,Me -+ Ph,Si.GePh + LiOMe + CO (e) With Epoxides and Ethers.-Although the reactions of silyl-lithium reagents with aliphatic ketones aldehydes and acid chlorides [sections 2(c) and (d)] have proved useful methods for the preparation of certain a-silyl-alcohols a variety of p-silyl-alcohols have been synthesised from the reactions of silyl-lithium reagents with epoxides. ao When triphenylsilyl-lithium was allowed to react with ethylene oxide propylene oxide styrene oxide and cyclohexene oxide good yields of 2-triphenylsilylethanol l-triphenylsilylpropan-2-ol 1-phenyl-2-triphenyl- silylethanol and 2-triphenylsilylcyclohexanol respectively were obtained.8o Similarly methyldiphenylsilyl-lithium added to styrene oxide and cyclohexene oxide to give 2-(methyldiphenylsily1)- 1 -phenylethanol and 2-(methyldiphenylsilyl)cyclohexanol respectively a O SiMe Ph + Ph,MeSiLi - H** If these few experiments allow generalisation silyl-lithium reagents seem to add to epoxides in the selective manner characteristic of organolithium W.Schlenk and E. Bergmann Annulen 1928,464 1. 7 n H. Gilman and C. W. Gerow J. Amer. Chem. SOC. 1955 77 4675. *O((a) H. Gilman D. Aoki and D. Wittenberg ibid. 1959 81 1107; (b) D. Aoki unpublished studies. 130 QUARTERLY REMEWS compounds 817 82 Likewise triphenylgermyl-lithium has recently been reported to add to certain epoxides.83 Though the reaction of aryl-lithium compounds with epichlorohydrin has been used as a convenient synthesis of arylpropylene chlor~hydrins,~~ no corresponding compound has been isolated by utilising triphenylsilyl- lithium. Instead there were found allyltriphenylsilane,80 1-chloro-3- triphenylsilylpropan-2-01 1,3-bistriphenylsilylpropan-2-ol and triphenyl- silanol in addition to 2,5-bis(triphenylsilylmethyl)- 1,4-dioxan 2Ph3SiLi + 2CH -CHCH,CI - 2Ph3SiCHi$HCH,Cl OLi \2 1 0 h U - 2LiCt t Ph,Si-CH2-HF’ ‘$H2 H,C ,CH.CHiSi*Ph 0 2Ph3Si Li + CH -CHCH2CI - Ph,SICH&H-CHiSiPh \20’ OL i - Ph,Si.OH + Ph,Si-CHiCH=CH A high yield of hexaphenyldisilane together with some allyltriphenyl- silane and triphenylsilanol were obtained when triphenylsilyl-lithium was allowed to react with epibromohydrin ;*O the reaction apparently involved a halogen-metal interconversion.85 The reaction of Grignard reagents and organolithium compounds with trimethylene oxide has been shown to give good yields of 3-substituted P,h ,Ph PQ ,Ph Reagents 1 CrO, then SO,CI,. 2 AICi,. alcohols.86 Similarly with triphenylgermyl-lithium 3-triphenylgermyl- S. J. Cristol J. R. Douglas and J. S. Meek J. Amer. Chern. SOC. 195! 73 816. 82 For a detailed discussion of reactions of Grignard reagents with epoxides see the excellent reviews by N. G. Gaylord and E. I. Becker Chem. Rev. 1951 49 413; S. Winstein and R B. Henderson in Chapter I Vol. I of Elderfield’s “Heterocyclic Com- pounds,” John Wiley & Sons Tnc.New York W.Y. 1950; and M. s. Kharasch and 0. Reinmuth “Grignard Reactions of Nonmetallic Substances,” Prentice Hall Inc. New York N.Y. 1954. 84 H. Gilman B. Hofferth and J. Honeycutt J. Amer. Clzem. Suc. 1952,74 1594. H. Gilman C. W. Gerow and M. R. Hughes unpublished studies. See section 2(b). S . Searles J. Amer. Chem. SOC. 1951 73 124. WITTENBERG AND GILMAN ORGANOSILYLMETALLIC COMPOUNDS 13 1 propanol was obtained.s3 Triphenylsilyl-lit hium produced the correspond- ing 3-triphenylsilylpropanol in a 77 % yield.18 This alcohol on oxidation yielded p-triphenylsilyl propionic acid. The corresponding acid chloride on treatment with aluminium chloride in nitrobenzene underwent cyclisa- tion to give 1,2,3,4-tetrahydr0-4-oxo-l 1-diphenyl-1-silanaphthalene the first reported derivative in the silanaphthalene series.*' Wittig and Ruckertss reported that tetrahydrofuran is unaffected by triphenylmethylsodium but is readily cleaved by the same reagent in the presence of triphenylboron.Normants9 found that Grignard reagents cleave tetrahydrofuran at 200" to give primary alcohols of the type R* (CHJ 4*OH. Triphenylsilyl-lithium as well as methyldiphenylsilyl- lithium cleaved this solvent slowly at the boiling point,36 and more readily at 125" to give 4-triphenylsilylbutan- 1-01 and 4-(methyldiphenylsilyl)~ butan- 1-01,~~ respectively. Conversion of these alcohols into the bromide- and subsequent treatment with lithium offers a new route to silacyclopen- tanes via an intramolecular cleavage-cyclisation :58 None of the expected 5-triphenylsilylpentanol was isolated when tri- phenylsilyl-lithium was prepared in tetrahydropyran and allowed to react with the solvent at elevated temperatures instead a high-melting polymer was obtained.lB Ethylene glycol dimethyl ether was cleaved by triphenylsilyl-lithium to give methyltriphenylsilane in high yield.The corresponding cleavage of dioxan gave ethylenebistriphenylsilane in addition to ethyltriphenyl- silane tetraphenylsilane and unidentified products :I8 /' \ CHt-CH "0 4 Ph,SiCH,.CH,.SiPh ZPh,Siti + 0 / Only a small amount of benzyltriphenylsilane the ether cleavage product was obtained when triphenylsilyl-lithium was allowed to react with benzyl methyl ether. The main reaction involved a metallation of the ether . The reaction of phenyl-lithium with 1,4-epoxy- 1,4-dihydronaphthalene a highly strained cyclic ether has been reported to yield 2-phenylnaphtha- lene.91 The corresponding reaction with triphenylsilyl-lithium gave the expected silicon compound 2-naphthyltriphenylsilane in low yield. CH ?-c H * 87 P. B. Talukdar and D. Wittenberg unpublished studies. G. Wittig and A. Riickert Annalen 1950 566 104. 8B H. Normant Cornpi. rend. 1954 239 15 10. G. Wittig and L. Pohmer Chem. Ber. 1956 89 1334 G. Wittig Angew. Chem. 1957 69 245, 132 QUARTERLY REVIEWS Naphthalene and triphenylsilanol were obtained as the main products. These were probably formed as a result of the aromatisation of an alcohol formed as an intermediate. Cleavage of alkyl dialkylaminomethyl ethers with Grignard reagent^^^,^^ and organolithium compoundsg4 has been described as a convenient synthesis of certain tertiary amines.Similarly triphenylsilyl-lithium was found to cleave 1-(n-butoxymethy1)piperidine to give 1-(triphenylsilyl- met hy1)piperidin e :I C,H,.O-CH,.NC,H, + Ph,SiLi -f Ph,Si.CHa.NC5HI + C,H,.OLi (f) With Nitrogen Compounds.-During early attempts to prepare silylmetallic compounds Kraus and Eatough2 treated bromotriphenyl- silane with lithium in ethylamine to obtain N-ethyltriphenylsilylamine. Trimethyltriphenylsilyltin2 and he~aphenyldisilane~~ have been cleaved by sodium in liquid ammonia and triethyltriphenylgermylsilane by lithium in eth~lamine.~' However since all silylalkali-metal compounds seem to react readily with ammonia ammonia and ethylamine are probably not suitable solvents for their preparation. When silyl-lithium compounds prepared in tetrahydrofuran were allowed to react with several primary and secondary a r n i n e ~ ~ ~ an instan- taneous reaction occurred and a negative colour test" was obtained.Working up the reaction mixture gave lithium hydride and good yields of silylamines Ph,SiLi + R,NH + Ph,SiH + RzNLi + LiH + Ph,Si-NR N-n-Butyl- 1 1 1 -triphenylsilylamine 4-triphenylsilylmorpholine and 1,4- bistriphenylsilylpiperazine are a few representative compounds prepared by this method. The reaction apparently proceeds in two steps the first being a metallation of the amine and the second a hydride displacement by the amide anion. Analogous reactions of triphenylsilane with lithium dialkylamides have been reported. Diphenylamineg5 and N-(diphenyl- methyl)anilineS6 were also metallated by triphenylsilyl-lithium.I n these cases however no secondary coupling occurred. Whereas alkyl- and aryl-lithium compounds attack pyridine in the 2-position 97 and selective nucleophiles such as benzylmagnesium chlor- ideg88.99 and allylmagnesium bromideloo give very low yields of 4-sub- stituted products triphenylsilyl-lithium adds smoothly to pyridine to give good yields of 4-triphenylsilyl- 1,4-dihydropyridine a compound 92 G. M. Robinson and R. Robinson J. 1923 123 532. 93 S. P. Massie Iowa State Coll. J. Sci. 1946,21,41 (Chem. Abs. 1947,41 3044). 94 A. H. Haubein ibid. 1943,18 48 (Chem. Abs. 1944,38 716). 95 H. Gilman and G. Lichtenwalter unpublished studies. s6 D. Wittenberg M. V. George T. C. Wu D. H. Miles and H. Gilman ibid. p. 4532. 97 K. Ziegler and A. Zeiser Ber. 1930 63 1847; Annalen 1931 485 174.8* W. L. C. Veer and S. Goldsmidt Rec. Trav. chiin. 1946 65 793. 99 R. A. Benkeser and D. S. Holton J. Amer. Chem. Soc. 1951 73 5861. loo H. Gilman J. Eisch and T. S. Soddy ibid. 1957 79 1245. WlTTENBERG AND GILMAN ORGANOSILYLMETALLIC COMPOUNDS 133 which is easily oxidised to 4-triphenyl~ilylpyridine.~~ None of the 2- substituted isomer was detected. In contrast triphenylmethylsodium and the even more reactive diplienylmethylsodium do not react with pyri- Triphenylsilyl-lithium has also been found to add to acridine in the 9,10-positions to yield after hydrolysis 9-triphenyl~ilylacridan.~~~ One distinct chemical difference between Grignard reagents and organolithium compounds is demonstrated by their modes of addition to the azomethine linkage of benzophenone mil.It has been shown that phenylmagnesium bromide does not react in ether but under forced conditions a lateral-nuclear 1,4-addition occurs to give N-(a-2-biphenylyl- benzy1)aniline. 0 4 3 1 O5 Pheny 1-li t hium however with benzopheno ne anil gives a 1 ,Zaddition product N-(triphenylmethyl)aniline.106J07 The com- pound formed on the reaction of triphenylsilyl-lithium or -potassium with benzophenone anil was N-diphenylmethyl-N-phenyl- 1 1 I-triphenylsilyl- amine an “abnormal” addition product containing a Si-N bond :96 dine 101,102 Ph,SiK + Ph2C= NPh -+ (Ph,Si)Ph,C*NKPh -f Ph,CH*NPh.SiPh The “abnormal” addition may be visualised to occur analogously to the related reactions of aromatic ketones with silylmetallic compounds by assuming that the initial step involves a “normal” addition and is followed by a rearrangement.Although phenylmagnesium bromide phenyl-lithiuni and phenyl- sodium react with azobenzene at room temperature yielding hydr- azobenzene and biphenyl reactions at low temperatures have been found to give mediocre yields of triphenylhydrazine the addition product.108-11n Triphenylsilyl-potassium and -lithium add smoothly to the azo-linkage to give NN’-diphenyl-N-triphenylsilylhydrazine in high yields. 96 The com- pound was synthesised unambiguously from NN‘-dilithio-NN’-diphenyf- hydrazine and chlorotriphenylsilane in tetrahydrofuran H2O Ph,SiK + PhN=NPh -+ Ph,Si*PhN-NHPh t Ph,SiCI + PhNLi-NLiPh Ha0 H,O The same compound was obtained in fair yield on reaction of tri- phenylsilyl-lithium with azoxybenzene.lll In this reaction in which triphenylsilanol was formed as a by-product the mechanism possibly lol K.Ziegler and H. Wollschitt Annulen 1930 479 123. lo2 E. Bergmann and W. Rosenthal J. prakt. Chem. 1932 135 267. lo3 H. Gilman and G. D. Lichtenwalter J. Urg. Chem. 1958 23 1586. l o b H. Gilman J. E. Kirby and C. R. Kinney J . Amer. Clzem. SOC. 1929 51 2252. lo5 R. C. Fuson R. J. Lokken and R. L. Pedrotti ibid. 1956,78 6064. lo* H. Gilman and R. H. Kirby ibid. 1933 55 1265. J. Eisch and H. Gilman Chem. Rev. 1957,57 525. lo8 H. Gilman and J. C. Bailie J. Org. Chem. 1937 2 84. loQ F. M. Beringer J. A. Farr jun. and S. Sands J . Amer. Chern. SOC. 1953 75 110 P. F. Holt and B. P. Hughes J. 1954 764; 1955 1320. M. V. George P. B. Talukder C. W. Gerow and H. Gilman unpublished studies. 3984. 134 QUARTERLY REVIEWS involves reduction of the N-0 linkage and subsequent addition of a second molecule of the silyl-lithium reagent to the azobenzene Ph,SiLi -+ PhN=NPh -+ Ph,Si*OLi + PhN-NPh __I_ Ph,Si-PhN-NLiPh Ph,SiLi Whereas phenyl isocyanate with triphenylsilylpotassium gave large amounts of hexaphenyldisilane in addition to some sym-diphenylurea phenyl isothiocyanate reacted with two molecules of triphenylsilyl- potassium to give a compound which may be bis(triphenylsily1) ketone anil :112 2Ph,SiK + PhN=C-=S -+ K,S 4- (Ph,Si),C=NPh Reactions of silylmetallic compounds with nitriles have been discussed in section 2 (d).(g) With Organosulphur Compounds.-Only a few reactions between silylmetallic compounds and organosulphur compounds have been in- vestigated and rather unexpected results have been obtained.Butyl-lithium has been reported to metallate diphenyl sulphide in the ortho-position.l13 With phenylsodiurn,ll4 as well as with diphenyl-lithium- dibenzothiophen was obtained as a secondary product. The reaction of diphenyl sulphide with triphenylsilyl-lithium followed by carbo- xylation gave benzoic acid and thiophenol in addition to hexaphenyl- disilane Ph,S + Ph,SiLi -+ PhLi + Ph,Si*SPh Ph,Si*SPh + Ph,Sili -f PhSLi + Ph,Si.SiPh The formation of these products is best explained by a cleavage of the phenyl-sulphur bond to give phenyl-lithium and triphenyl(phenylthi0)- silane ; the latter then couples with triphenylsilyl-lithium to form hexa- phenyldisilane and lithium phenyl sulphide. The last step is analogous to the reported cleavage of triphenyl-(ptoly1thio)silane by triphenylsilyl- lithium to give hexaphenyldisilane and lithium p-tolyl sulphide.ll i~118 Hexaphenyldisilane and tetraphen ylsilane were isolated after reaction of diphenyl sulphone with triphenylsilylpotassium.With triphenylsilyl- lithium m-phenylenebistriphenylsilane triphenylsilanol hexaphenyldi- siloxane benzenesulphinic acid thiophenol and hydrogen sulphide in addition to hexaphenyldisilane and tetraphenylsilane were isolated from the reaction mixtures. Benzoic acid and p-triphenylsilylbenzoic acid were identified in carboxylated reaction mixtures.l16 The mode of formation of these compounds is still unknown. As the initial step a cleavage of the phenyl-sulphur bond has been postulated since only benzoic acid 112 T. C. Wu unpublished studies. 113 H. Gilman and R. L. Bebb J . Arner.Chem. SOC. 1939 61 109. lX4 V. A. Luttringhaus G . Wagner-v. Saaf E. Sucker and G. Borth Annalen 1945 115 G. Wittig and E. Benz Chem Ber. 1958 91 873. 116 D. Wittenberg T. C. Wu and H. Gilman J. Org. Chem. 1959,23 1898. 117 H. Gilman and D. Wittenberg J. Amer. Chem. SOC. 1957 79 6339. 118 D. Wittenberg H. A. McNinch and H. Gilman ibid. 1958 SO 5418. 557 54. WITENBERG AND GILMAN ORGANOSILYLMETALLIC COMPOUNDS 135 triplienylsilanol and benzenesulphinic acid are formed at low tempera- tures Ph.SO,.Ph 4- Ph,SiLi -f PhLi 4- Ph,Si.SO,*Ph -f Ph-CO,H 7- Ph.SO,H i Ph,Si.OH CO, H,O An attempt to synthesise triphenyl(phenylsu1phonyl)silane from tri- phenylsilyl-lithium and benzenesulphonyl chloride resulted in the forma- tion of hexaphenyldisilane and lithium benzenesulphinate 2Ph3SiLi + Ph-SO,CI -f Ph,Si -+ Ph-S0,Li -b LiCl (h) With Organosilicon Compounds.-(i) Silicon halides.The reaction of silylmetallic compounds with a variety of chlorosilanes has been used in syntheses of disilanes especially unsymmetrical ones which are inaccessible by other methods. 1,1,l-Trimethyl-2,2,2-triphenyldisilane has been pre- pared by reaction of triphenylsilyl-pota~sium,~~-~~J~~~~ - l i t h i ~ m ~ ~ ~ -rubidium15b and - c a ~ i u r n ~ ~ ~ with chlorotrimethylsilane. Similarly 1,l ,I- triethyl-2,2,2-triphenyldisilane has been prepared from triphenylsifyl- potassium and chlor~triethyl~ilane,~~~~~~~~~~~ l,l,l-trihexadecyl-2,2,2-tri- phenyldisilane from triphenylsilyl-lithium and chlorotrihexadecylsilane,~lg and 1,l-dimethyl-l,2,2,2-tetraphenyldisilane from triphenylsilylpotassium and methyldiphenylsilyl A number of hexa-aryldisilanes containing phenyl andp-to1 yl groups was obtained by coupling a triarylsilylpotassium with a triarylchlorosilane.24 Partially chlorinated disilanes have been synthesised by reactions involving triphenylsilylpotassium with one equivalent of silicon tetra- chloride trichlorophenylsilane and dichlorodiphenylsilane.The products were 1 1 l-trichloro-2,2,2-triphenyldisilane,24 1 I-dichloro- 1,2,2,2-tetra- phenyldi~ilane,~~ and chloropentaphenyldisilane,15 respectively. When two equivalents of triphenylsilylpotassium were allowed to react with dichlorodiphenylsilane octaphenyltrisilane was formed:15J20 Ph,SiK 4- Ph,SiCI -+ Ph3Si*SiPh,C1 - + Ph,Si.SiPh,.SiPh Ph,SiK Reaction of triphenylsilyl-lithium with trichlorosilane gave a small amount of tristriphenylsilylsilane the first reported branched-chain p~lysilane.~~ Although coupling reactions of triarylsilyl-lithium compounds with aryl- and alkyl-substituted chlorosilanes seem to proceed without com- plications reactions of silyl-lithium reagents containing alkyl as well as aryl groups with chlorotriphenylsilane yielded mixtures of several di- ~i1anes.l~ Further the nature of the products was dependent upon the mode of addition.Slow addition of dimethylphenylsilyl-lithium to chlorotri- phenylsilane gave 1,l -dimethyl- 1,2,2,2-tetraphenyldisilane. When the ll@ D. H. Miles unpublished studies. lZo H. Gilman and J. J. Goodman J . Arner. C11ent. Soc. 1953 75 1250. 136 QUARTERLY REVIEWS reverse addition was employed hexaphenyldisilane and 1,1,2,2-tetra- methyl-1 ,Zdiphenyldisilane were formed.lg The products appear to indicate that halogen-metal interconversion occurred between the Si-Li and Si-Cl compounds but subsequent studies have confirmed that cleavages of disilanes by silyl-lithium reagents are involved.In the presence of excess of the silyl-lithium reagent the initial coupling product is cleaved by the silylmetallic compound to give a symmetrical disilane and a less reactive silyl-lithium compound30 [see section l(e)] PhMe,SiLi + Ph,SiCl -+ PhMe,Si-SiPh + LiCl PhMe,SiLi + PhMe,Si.SiPh -+ PhMe,Si.SiPhMe + PhJiLi Ph,SiLi + Ph,SiCI -+ Ph,Si-SiPh + LiCl Large amounts of hexaphenyldisilane have been isolated after reactions of triphenylsilyl-lithium with silicon tetrachloride trichlorosilane ethyl silicate dichlorodiphenylsilane and chl~ropentaphenyldisilane.~~ This effect is much more pronounced with silyl-lithium reagents in tetra- hydrofuran than with silylpotassium reagents in ethyl ether.A mechanism involving the cleavage of a disilane by the silylmetallic compound has been proposed Ph,SiLi + SiCI4 -+ Ph,Si.SiCI ~ -+ Ph,Si.SiPh + LiSiCI + Ph,SiLi nLiSiCI -+ nLiCl + (SiCI,), Ph,Si*SiPh,CI + Ph,SiLi -+ Ph,Si.SiPh + Ph,SiCILi nPh,SiCILi -+ nLiCl + (Ph,Si) When (diphenylmethoxy)diphenylsilane obtained from diphenylsilane and benzophenone,121 was allowed to react with triphenylsilyl-lithium diphenylmethanol was formed in high yield in addition to some penta- phenyl disilane. Some hexaphenyldisilane was formed owing to further attack of pentaphenyldisilane by the silyl-lithium reagent [see section l(e)] .Related cleavages of disilanes by organometallic compounds have been known since 1869 when Friedel and Ladenburg122 investigated the reaction of diethylzinc with hexaiododisilane and obtained the expected hexaethyl- disilane in addition to some tetraethylsilane. Similarly tetraphenylsilane was the only product from hexachlorodisilane chlorobenzene and ~ 0 d i u r n . l ~ ~ When octachlorotrisilane was treated with phenylmagnesium bromide the products were hexaphenyldisilane and tetrapheny1~ilane.l~~ ( i i ) Silicon hydrides. Reactions of excess of triphenylsilylpotassium with phenylacetylene as well as with several phenyl-substituted alcohols gave high yields of tetrapheny1silane.l2j Similarly triarylmethanes when treated with an excess of a triarylsilylpotassium compound gave tetra-arylsilane~.~~ A subsequent investigation showed that in all of these experiments the lZ1 H.Gilman and D. Wittenberg J. Org. Chem. 1958 23 501. lZ2 C. Friedel and A. Ladenburg Compt. rend. 1869,68,923 ; Annalen 1880,203,251. lZ4 W. C. Schumb and C. M. Saffer J. Amer. Chem. Sac. 1939,61,363 lP6 H. Gilman and T. C. Wu ibid. 1953 75 2509. L. Gattermann and K. Weinlig Ber. 1894 27 1946. WITTENBERG AND GILMAN ORGANOSILYLMETALLIC COMPOUNDS 137 tetra-arylsilane is formed by reaction of a triarylsilylmetal with a triarylsilane :126 R,SiK + R,SiH -+ R,Si + “other products” Examination of a number of reactions proved that it is possible to cleave more than one aryl group from a triarylsilane.f26 The reaction is even more complex when the silylmetallic compound and the silane have different aryl groups Thus te traphen ylsilane te tra-p- to1 ylsilane and phenyl tri-p-tol ylsilane were isolated from the products of reaction of triphenylsilylpotassium with tri-p-tolylsilane in ethyl ether.Tetraphenylsilane and phenyltri-p-tolyl- silane were formed from tri-p-tolylsilylpotassium and triphenylsilane.126 The reaction of triphenylsilyl-lithium with triphenylsilane is slower and gives a high yield of tetraphenylsilane in tetrahydrofuran at elevated temperature^.^^ The same reaction in ethylene glycol dimethyl ether seems to take a different course hexaphenyldisilane as well as tetraphenylsilane were isolated.126 The coupling of triphenylsilyl-lithium with triphenylsilane to form hexaphenyldisilane is not unexpected since treatment of triphenyl- silane with organolithium compounds is a well-recognised method for preparing tetrasubstituted silanes Ph,SiLi + Ph,SiH -+ Ph,Si*SiPh 3- LiH The formation of tetra-arylsilanes in these reactions however is surprising and is still not completely understood.(i) With Inorganic Compounds.-(i) Hydrolysis. The products formed by hydrolysis of silylmetallic compounds depend on the conditions employed. Benkeser and Seversonf3 have shown that decomposition of triphenyl- silylpotassium with hydrochloric acid yields triphenylsilane. In the presence of sodium-potassium alloy treatment with dry hydrogen chloride diluted with oxygen-free nitrogen has been re~0mmended.l~ Only tri- phenylsilanol was obtained on reaction of triphenylsilylpotassium with water.15 It seems certain that under these conditions the silicon hydride formed initially is attacked by the strong base formed during hydrolysis to yield the silanol and gaseous hydrogen :12’ R,SiK + R,SiH R,Si + R,SiR’ + R’,SiR + R’Ji + “other products” Ph,SiK + H,O + Ph,SiH + K O H Ph,SiH + H,O - + Ph,Si-OH + H If water an alcohol or other acidic compound is employed in an amount insufficient for complete hydrolysis of the silylmetallic reagent interaction may take place between the silicon hydride formed and the silylmetallic compound [see section 2(h)] .126 A. G. Brook m d H. Gilman ihid. 1954 76,- 2333. 12’ For a theoretical discussion of the conversion of silicon hydrides into silanols under conditions of aqueous hydrolysis see H. Gilman and G . E. Dunn J . Amer. Chem. SOC.. 1951. 73. 3404. 138 QUARTERLY REVIEWS (ii) Curboxylation.Carboxylation of triphenylsilylpotassium was found to yield an unstable acid which decomposed at its melting point to give carbon monoxide and together with other products triphenyl- silan01.l~ Later studies described experimental conditions for the prepara- tion of pure triphenylsilanecarboxylic acid in good In a basic medium the acid is readily decomposed to carbon nionoxide and triphenylsilanol. It is possible to convert triphenylsilyl-lithium into triphenylsilanol in excellent yields by this method.128 Base-catalysed eliminations and thermal rearrangements of triphenylsilanecarboxylic acid its derivatives and related compounds have been published recently. 253 129 Ph,SiK + CO -f Ph,Si.CO,K -+ Ph,Si-CO,H -+ Ph,Si.OH + CO (iii) Oxygen. Reactions of triphenylsilyl-lithium with oxygen have been studied21 for tetrahydrofuran and tetrahydropyran solutions at tempera- tures of 25" 0" and -25".The products were triphenylsilanol triphenyl- silane and hexaphenyldisilane. Higher temperatures resulted in an in- creased amount of hexaphenyldisilane and a decreased amount of tri- phenylsilanol. The mechanism of the reaction is not clear but may involve a free-radical path Ph,SiLi + 0 -+ Ph,Si.O-OLi - 2Ph,Si*OLi Ph,SiLi Ph3Si-0.0Li 4 Ph,Si.O- + WOLi PhsSi.O. + Ph,SiLi -f Ph,Si-OLi -t Ph,Si. Ph,Si. + Solvent -f Ph,SiH 2Ph,Si. -+ Ph,Si-SiPh (iv) Sulphur. Triphenylsilyl-lithium in tetrahydrofuran reacts smoothly with sulphur to give the lithium salt of triphenylsilanethiol. Subsequent reactions of this intermediate with methyl iodide benzyl chloride and benzoyl chloride have afforded methylthio- benzylthio- and benzoylthio- triphenylsilane respectively in good yields,130 e.g.Ph,SiLi 4- S -A Ph,Si*SLi -+ Ph,Si.SMe -4- Lil Me1 3. Compounds containing silicon bonded to non-alkali metals (a) Magnesium.-Cusa and Kipping131 in 1933 obtained dicyclohexyl- phenylsilane by reaction of trichlorophenylsilane with cyclohexylmagne- sium bromide. This unexpected result was explained by postulating the intermediate formation of a silyl-Grignard reagent from chlorodicyclo- hexylphenylsilane and cyclohexylmagnesium bromide by means of a halogen-metal interconversion. A corresponding reaction between 128 A. G. Brook and H. Gilman J. Amer. Chem. Soc. 1955 77 2322. 130 G. D. Lichtenwalter unpublished studies. I 3 l N. W. Cusa and F. S. Kipping J . 1933 1040. A.G. Brook ibid. p. 4827. WITTENBERG AND GILMAN ORGANOSILYLMETALLIC COMPOUNDS 1 39 chlorotriphenylmethane and phenylmagnesium bromide has been reported by Gomberg and Cone:132 Ph,CCI+ PhMgX -+ Ph,C.MgX -+ Ph,CH H,O Ph(C,H,,),SiCI + C,H,,MgBr + Ph(C,H,,),Si.MgX A recent re-inve~tigationl~~ of Kipping’s results showed that several sterically hindered Grignard reagents such as those containing the cyclo- hexyl cyclopentyl isopropyl or isobutyl group react with trichlorophenyl- silane to give compounds of the PhR,SiH type and further that the silicon hydride is formed before hydrolysis of the reaction mixture. The cor- responding unsaturated hydrocarbons cyclohexene cyclopentene pro- pene and isobutene were also identified as reaction products PhR,SiCI + C,H,,MgBr 4 PhR,SiH + C,Hl0 + MgBrCl Similarly triphenylsilane and cyclohexene were isolated after reaction of chlorotriphenylsilane with cyclohexylmagnesium bromide.57 No silyl- magnesium compound could be detected in these reactions.However when tetrahydrofuran was employed as the solvent a different course of reaction was observed. Hexaphenyldisilane was isolated in high yield when chlorotriphenylsilane was refluxed in this solvent with cyclohexyl- 2-methylcyclohexyl- or phenyl-magnesium bromide.133b Chlorotrimethyl- silane under corresponding conditions remained unchanged. The formation of triphenylsilylmagnesium bromide as an intermediate was confirmed by employing a mixture of chlorotriphenylsilane and chlorotrimethylsilane which yielded l,l,l-trimethyl-2,2,2-triphenyldisilane the unsymmetrical coupling product.Ph,SiCI + 2C,H,,MgBr -+ Ph,SiMgBr + C,& + C&O + MgClBr Ph,SiMgBr + Ph,SiCI + Ph,Si*SiPh + MgClBr Ph,SiMgBr + Me,SiCI -+ Ph,Si.SiMe + MgClBr It has been observed that magnesium dissolves in a solution of silyl iodide in isopentyl ether. The formation of silane hydrogen and silicon as the hydrolysis products was interpreted as meaning that an unstable silylmagnesium iodide was formed.13* Silyl bromide however was not to react with magnesium in butyl ether. Neither triethyliodo- ~ i l a n e l ~ ~ nor chlor~triphenylsilane~~~ reacts with magnesium in ether. Preliminary experiments however indicate that formation of silyl- magnesium reagents is also possible from the corresponding silyl-lithium compound and anhydrous magnesium bromide in tetrahydr0f~ran.l~~ (b) Mercury.-An unstable volatile compound possibly SiH,*HgI has been reported from the reaction products of silyl iodide with mercury M.C. Harvey W. H. Nebergall and J. S. Peake J. Amer. Chem. SOC. 1957 134 H. J. ErnelCus A. G. Maddock and C. Reid Nature 1939,144,328; J. 1941 353. 136 E. R. Van Artsdalen and J. Gavis J. Amer. Chem. SOC. 1952 74 3196. lS8 C. Eaborn. J. 1949. 2755. M. Gomberg and L. H. Cone Ber. 1906,39 1461. 79 2762; (b) T. G. Selin and R. West Tetrahedron 1959,5 97. 4 140 QUARTERLY REVIEWS but no details of its properties were give11.l~~ Results obtained from reac- tions of silylmetallic compounds with organomercury compounds and mercury halides13' also indicate the existence of unstable silicon-mercury compounds. The reaction of triphenylsilyl-lithium with diphenylmercury in a 1 1 ratio yielded after carboxylation benzoic acid tetraphenylsilane and mercury.In an analogous reaction with di-p-tolylmercury the products isolated were p-toluic acid triphenyl-p-tolysilane and mercury Ph,SiLi + RzHg -+ RLi + Ph,Si-HgPh -f R.CO,H + Ph3SiR + Hg CO, H 2 0 Similarly tetraphenylsilane and metallic mercury were obtained from triphenylsilyl-lithium and phenylmercuric chloride. The formation of chlorotriphenylsilane and hexaphenyldisilane from mercury(I1) chloride and triphenylsilyl-lithium was explained by a similar mechanism Ph,SiLi + HgCI -+ LiCl + Ph,Si*HgCI +- Ph,SiCI + Hg (c) Germanium and Tin.-Two different methods have been employed for the preparation of silicon-germanium and silicon-tin bonds the reaction of a silylmetallic compound with a germanium or tin halide and the reaction of a silicon halide with a germyl- or tin-metallic com- pound.Triphenyltriphenylsilylgermane was obtained from the reaction of triphenylsilylpotassium with bromo- or chloro-triphenylgermane.*l This compound was isolated also after reaction of triphenylsilyl-lithium with methyl triphenylgermanecarboxylate. 79 Reactions of triphenylgermyl- potassium -sodium and -lithium with chloro- or bromo-triethylsilane gave triphenyltriethylsilylgermane.2 7~138 A quantitative yield of tristriphenylgermylsilane was obtained from triphenylgermylsodium and trichloro~ilane.~~~~~ Lithium in ethylamine con- verted the compound into the corresponding silyl-lithium derivative which on subsequent coupling with ethyl bromide gave ethyltristriphenyl- germylsilane 3Ph,GeNa + SiHCI -+ (Ph,Ge),SiH -+ (Ph,Ge),SiEt The same silyl-lithium reagent when allowed to react with chlorotri- p hen y It in yielded trip henyltr is(trip henylgermy1)sily ltin the first reported compound containing a Ge-Si-Sn linkage :139 (Ph,Ge),SiLi + CISnPh -+ (Ph,Ge),Si*SnPh Trimethyltriphenylsilyltin was formed from triphenylsilyl-lithium and chlorotrimethyltin.2 Triphenyltriphenylsilyltin has been obtained both 13' M.V. George G. D. Lichtenwalter and H. Gilman J. Amer. Chem. SOC. 1959 81 978. lg8 H. Gilman and C. W. Gerow ibid. 1955,77 5509. 13Q J. J. Goodman unpublished studies. WITTENBERG AND GILMAN ORGANOSILYLMETALLIC COMPOUNDS 14 1 from triphenylsilylpotassium with ~hlorotriphenyltin,~~~ and from chloro- triphenylsilane with triphenylstaiinyl-lithium.140 Various stable silylmetallic compounds are listed in Table 3.(a) Other Metals.~-Cyclopentadienyldicarbonyl(trimethylsilyl)iron a compound which is thought to contain a true silicon-iron bond has been prepared from the corresponding sodium derivative and chlorotrimethyl- ~i1ane.l~~ The compound decoinposes at 200" and reacts readily with oxygen (C,H,)(CO),FeNa + Me,SiCI + NaCl + (C,H,)(CO),Fe.SiMe TABLE 3. Stable silylmetallic compounds Compound Prepared from M.p. Yield Ref. ( %I Ph,Si.GePh Et ,Si-GePh3 (Ph,Ge),SiH (Ph,Ge),SiEt ( Ph,Ge),SiBr (Ph,@e),SiCl (Ph,Ge),SiOH (Ph,Ge),S!NH (Ph,Ge),Si*SnPh Ph,Si.SnMe Ph,Si*SnPh (CaH,)(CO)2Fe-SiMe3 Ph,SiK + Ph,GeCl 354-355" 43-63 Ph,SiLi + Ph3GeCOzMe 357-359 84 Ph,GeNa + Et,SiBr 93.5 - Ph3GeK(Li) + Et,SiCl 97-98 4&63 Ph,GeNa + SiHCl (Ph,Ge),SiLi + EtBr 283 63 (Ph,Ge),SiH + Brz 242 - (Ph,Ge),SiNH + HCl 230-231 - (Ph,Ge),SiBr + NH,OH 197 - (Ph,Ge),SiBr + NH 206 - (Ph,Ge),SiLi + Ph,SnCl 340-342 15.5 Ph,SiLi + Me,SnCl Liquid - Ph,SiK + Ph3SnCl 296-298 76 (C5H J (CO) ,FeNa 70 42 170-171) 1 87- 1 8 8 loo Ph,SiCI + Ph,SnLi 289-291 71 + Me3SiCl 41 79 27 138 44 44 44 44 44 44 139 2 15b 140 141 When silyl iodide was allowed to react with metallic zinc for a year in a sealed tube a non-volatile liquid was obtained which showed the properties expected for silylzinc iodide.134 Preliminary e x p e r i m e n t ~ ~ ~ ~ l ~ ~ indicate that several compounds containing silicon bonded to non-alkali metals are formed in reactions of silyl- lithium compounds with anhydrous metal halides in tetrahydrofuran.Many of these new silylmetallic compounds however are unstable and only their decomposition products have been isolated so far.Research in this field is still in progress. 4. Analytical procedures (a) Colour Tests.-The most widely applied colour test for reactive organometallic compounds involves their reaction with Michler's ketone1 which gives after hydrolysis and oxidation with iodine in acetic acid a 140 H. Gilman and S. D. Rosenberg J. Amer. Chem. SOC. 1952 74 531. 141 T. S. Piper D. Lemal and G. Wilkinson Naturwiss. 1956,43 129. 142 QUARTERLY REVIEWS blue or blue-green colour. When triphenylsilylpotassium in ether was tested,15 the supernatant liquid gave a Negative Colour Test I whereas the precipitate gave an intense blue-violet colour in both the organic and the aqueous layer. Similarly solutions of silyl-potassium -sodium and -lithium reagents in ethylene glycol dimethyl etheP and tetrahydr~furanl~ have been found to give a positive colour test I.Colour Test HA the reaction of an organometallic compound with p-brornodimethylaniline followed by treatment with benzophenone and concentrated hydrochloric acid with production of a red colour has been used to distinguish between highly reactive alkyl-lithium and less reactive aryl-lithium compounds. Silyl-lithium gave a slightly positive test only under special conditions. This phenomenon however is not due to a lower reactivity on the part of the Si-Li reagents but on the contrary is a reflection of their higher reactivity which im- mediately gives rise to coupling products. (b) Quantitative Methods.-In some cases yields of silylalkali-metal compounds have been determined by titration of aliquot parts of a solution of reagent with In order to determine the alkoxide which may be formed by interaction of the silylmetallic compound and the solvent the double-titration was e m p l ~ y e d .~ ~ . ~ ~ In this method the total alkali is determined by the usual titration and a second aliquot portion is added to benzyl chloride and subsequently titrated with acid to determine the alkoxide content. As a method which is thought to give the most reliable values the procedure of Ziegler and his c o - w ~ r k e r s ~ ~ ~ is recommended. An aliquot part of the solution is allowed to react with n-butyl bromide after hydro- lysis the bromide content of the aqueous layer is determined by Volhard’s method. 5. Relative reactivity of silylmetallic compounds and other comparisons with related types (a) Cleavage Reactions.-Owing to their dissociation into free radicals hexa-arylethanes are cleaved even under mild conditions for instance by sodium amalgam.Symmetrical tetraphenylethane required sodium- potassium alloy for its conversion into diphenylmethylpotassium while 1,2-diphenylethane and related diarylethanes are resistant toward cleavage by metals.145 Similarly it was found that the rate of cleavage of symmetri- cal disilanes by lithium increases as the number of phenyl groups attached to the silicon atoms increases.lg Despite the extremely low solubility of hexaphenyldisilane in organic solvents only a few hours are required for its 14* H. Gilman and J. Swiss J. Amer Chem. SOC. 1940 62 1847. 148 H.Gilman and A. H. Haubein ibid. 1944 66 1515. 14‘ K. Ziegler F. Crossmann H. Kleiner and 0. Schafer Annalen 1929,473 21. 145 F. Runge “Organornetallverbindungen” Wissenschaftl. Verlags-GmbH. Stutt- gart 1944 p. 45. WITTENBERG AND GILMAN ORGANOSILYLMETALLIC COMPOUNDS 143 complete cleavage. Lithium in tetrahydrofuran splits sym-dimethyltetra- phenyldisilane in less than one hour while the corresponding cleavage of sym-tetramethyldiphenyldisilane requires several hours. Hexa-alkyldi- silanes are not cleaved by metals. Hexa-alkyldigermanes also seem to resist cleavage by metals whereas alkali metals readily split hexamethyl- ditin. 5e Increasing atomic weight and atomic radius should be factors leading to weakening of the M-M bond in R,M compounds. Moreover replace- ment of alkyl by aryl groups should aid in weakening the M-M bond because of the greater electron-withdrawing power of the aryl groups.The aforementioned cleavage studies verify these theoretical aspects. Since delocalisation of the negative charge on the aryl groups will appreciably stabilise the anion the compounds may be arranged in the following order of stability ; Decreasing stability is equivalent to increasing reactivity. A correspond- ing reactivity series of silylmetallic compounds has been deduced from cleavage of disilanes by silyl-lithium compounds [see section l(e)] . (b) Metallation.-At present it seems impossible to draw general conclusions concerning the reactivity of silylmetallic in comparison with organometallic compounds since no kinetic studies have yet been per- formed.Moreover comparison of corresponding reactions indicates that a silylmetallic can be more reactive than a given organometallic compound toward one substrate and less reactive toward another. The anions of organoalkali-metal compounds if compared under similar conditions can be arranged in a series of increasing proton Ar3M- > Ar,AlkM- > ArAlk,M- > AIk,M- affinity :146,14 7 Bu- > Ph- > Ph*CH,- > Ph,CH- > Ph3C- > PhmNH- > RO- Thus alkylsodium compounds have been found to metallate benzene ; phenylpotassium metallates toluene to give benzylpotassium; and benzyl- potassium will metallate diphenylmethane. Future studies may determine the exact position which triphenyl-silyl -germyl and -tin anions would hold in such an expanded series at present little information is available.Phenyl-lithium and butyl-lithium have been found to metallate triphenyl- germane,148 while triphenylgermyl-lithium has been found to metallate compounds such as fl~0rene.l~~ Related metallations involving triphenyl- tin and triphenylstannyl-lithium have been described.150 A metallation of J. B. Conant and G. W. Wheland J. Amer. Chem. Sac. 1932 54 1212. lP7 W. K. McEwen ibid. 1936 58 1 124. 148 H. Gilman and C. W. Gerow ibid. 1956,78 5435. 14@ H. Gilman and C. W. Gerow J. Org. Chem. 1958,23 1582. 150 J. d'Ans H. Zimmer E. Endrulat and K. Lubte Naturwiss. 1952 39 450; G. Wittig F. J. Meyer and C. Lange Annalen 1951 571 167; H. Gilman and S. D. Rosnberg J. Amer. Chem. Soc. 1953 75 3592; H. Gilman and L. A. Gist jun. J. Org. Chem. 1957,22 689. 144 QUARTERLY REVIEWS triphenylsilane has not yet been accomplished since silicon hydrides on treatment with various anions displace a hydride anion rather than abstract a proton.7 9 9 Although alkylpotassium compounds such as phenyl- potassium and even benzylpotassium are known to metallate ethyl ether readily triphenylsilylpotassium is stable in this solvent indicating its lower reactivity. On the other hand triarylmethanes are readily metallated by silylalkali-metal e.g. fluorene by triphenylsilyl-lithium. 33 One might estimate that the proton affinity of the triphenylsilyl anion is between that of the benzyl and the triphenylmethyl anion and perhaps comparable with the corresponding dip henylme t hyl anion. (c) Addition.-Various organometallic compounds have been found to add to olefins. For example aa-dimethylbenzylpotassium adds to anthra- cene stilbene and 1,l-diphenylethylene but not to triphenylethylene.144J51 Triphenylsilyl-lithium adds to anthra~ene,~~ ~ t i l b e n e ~ ~ and 1 l-diphenyl- ethylene,52 as well as to tri~henylethylene.~~ Triphenylgermyl-lithium adds to octadec-1-ene and 1,l-diphenylethylene but not to ~ti1bene.l~~ Tri- phenyltinlithium is unreactive in all similar addition rea~ti0ns.l~~ Compounds containing C-Li Si-Li and Ge-Li linkages add to the carbonyl group of aldehydes and ketones add to the azo- and the azo- methine linkage and cleave organic oxides such as ethylene oxide and trimethylene oxide.Corresponding Sn-Li reagents are rather unreactive and only their addition to ethylene oxide has been ac~omplished.15~ In the series of the triphenyl derivatives Ph,C-M Ph,Si-M Ph,Ge-M Ph,Sn-M the silyl- and germyl-metallic compounds seem to be the most reactive in addition.Their reactivity seems to be comparable with that of aa-dimethylbenzylpotassium. Steric effects as well as differences in electron-distribution in the carbon-metal and metal-metal bonds in these compounds seem to be responsible for the specific differences observed in some of their additive reactions. (d) Coupling and interconversion Reactions.-Reactions of halides with organometallic compounds are known to proceed mainly by two routes coupling and halogen-metal interconversion. Coupling of C-M Si-M Ge-M Sn-M and Pb-M types with alkyl and aryl halides have been widely applied for synthetic purposes. Interconversion has been used for the preparation of certain organolithium compounds4Jo which are inaccessible by other methods.Alkyl-lithium reagents convert aryl bromides and iodides into aryl- lithium compounds. Aryl-lithium reagents show interconversion with benzyl-type halides as well as with aryl iodides. Germylmetallic compounds have not been fully investigated only coupling and no interconversion products have been reported from their reactions with halides. Some stannyl-lithium compounds have been found to undergo halogen-metal 151 K. Ziegler and K. Bahr Ber. 1928 61 253. 15a H. Gilman and C. W. Gerow J. Amer. Chem. Sac. 1957,79 342. 158 H. Gilman and S D. Rosenberg J. Org. Chem. 1953 18 1554. WITTENBERG AND GILMAN ORGANOSILYLMETALLIC COMPOUNDS 145 interconversion with certain aryl iodides.154 Silylmetallic compounds as described in section 2(b) exhibit an extremely high order of reactivity in interconversion.Silylpo tassium and especially silyl-lithium reagents give interconversion products not only when allowed to react with aryl bro- mides and iodides but also when they react with aryl chlorides alkyl bromides and alkyl iodides. The phenylthio- and the phenylsulphonyl radical seem to act as pseudo-halogens since phenyl-lithium is formed in the reactions of triphenylsilyl-lithium with diphenyl sulphide and with diphenyl sulph one. 116 Silylmetallic compounds also show a high order of reactivity in coupling reactions. When triphenylsilyl-lithium and phenyl-lithium were allowed to compete for chlorotriphenylsilane at -50° hexaphenyldisilane was isolated in 85 % yield and tetraphenylsilane in only 1.5 % yield.Reasons for the increased reactivity of silylmetallic compounds become clearer on comparison of bond energies of corresponding carbon and silicon compounds. Despite the fact that Si-H and Si-C bonds are slightly weaker than C-H and C-C bonds silicon forms stronger bonds than carbon with electronegative elements such as 0 N F C1 Br and 1.v Recent studies have shown that the triphenylsilyl radical readily abstracts a chlorine radical from chlorobenzene a reaction not observed with hydrocarbon radi~a1s.l~~ The same tendency of silicon to combine with electronegative elements gives rise to interconversion when silylmetallic compounds are allowed to react with organic halides. (e) Conclusion.-Organosilylmetallic compounds are a relatively new development in chemistry. Fortunately they can now be prepared directly and with ease.The uncommon diversity of their chemical behaviour is striking. They promise to open broad avenues for the synthesis of organo- silicon types which are either novel or have been inaccessible hitherto. The authors are deeply grateful to Bernard J. Gaj William J. Trepka and 154 H. Gilman and S. D. Rosenberg ibid. p. 680. 155 J. Curtice H. Gilman and G. S. Hammond J. Arner. Chem. Soc. 1957 79 4754. Justin Diehl for assistance.

ISSN:0009-2681    

DOI:10.1039/QR9591300116

出版商:RSC    

年代:1959

数据来源: RSC

摘要:

QUANTITATIVE STUDIES OF HYDROLYTIC EQUILIBRIA By LARS GUNNAR SILL~N (DEPARTMENT OF INORGANIC CHEMISTRY ROYAL INSTITUTE OF TECHNOLOGY STOCKHOLM 70.) Introduction IF an iron(Ir1) salt say the nitrate or perchlorate is dissolved in water the solution has an acidic reaction. Thus reaction with the water a “hydro- lysis” has taken place:” xFe3+ + yH20 + Fex(OH),(3z-Y)+ + yH+ (1) Iron(II1) is by no means unique in this respect almost all cations undergo this reaction perceptibly in aqueous solution e.g. Be2+ AP+ UOZ2+ Cu2+. Many explanations have been offered but not always supported by suffi- cient experimental evidence. Some of the older textbooks for instance state that uncharged hydr- oxide is formed Fe3+ + 3H20 $ Fe(OH) + 3H+ This requires either microcrystals or a true solution of the hydroxide and both these explanations can be ruled out by applying the law of mass action even to rather crude measurements.It was suggested by Werner1 and by Pfeiffer2 that protons are split off from the water molecules bound to the cation (“aquo-acidity”) e.g. or more briefly Fe(H20),3+ + H 2 0 + Fe(H20),0H2+ + H,O+ Fe3+ + H20 + FeOH2+ + H+ (2) The equilibrium constant of this reaction would be the acidity constant K, of the Fe3+ ion. Niels Bjerr~m,~.~ who was the pioneer in this field as in many others determined K for Cr3+ as early as 1906. Bronsted and V~lqvartz,~ in 1928 from kinetic measurements of [H+] and solubility * Editor’s note As written in equation (1) this reaction appears to be fission (lysis) of water rather than fission by water the latter being the sense in which the term “hydrolysis” is generally used in other fields.The reaction however assumes the form of hydrolysis by water if it is written in the old (inadequate) non-ionic form such as FeCl + H20 3 FeC12.0H + HCl The word “hydrolysis” is used in this Review for any such reactions in which water takes part as is customary in other publications in this field. A. Werner Ber. 1907 40 272; “Neuere Anschaungen auf dem Gebiete der anorganischen Chemie” Vieweg and Son Braunschweig 2nd edn. 1909 p. 238. P Pfeiffer Ber. 1907 40 4036. N. Bjerrum Kgl. danske Videnskab. Selskub. Skrifter Nat.-mat. Afd. 1906 4 1; 2. phys. Chem. 1907,59 336; 1910,73 724. J. N. Bronsted and K. Volqvartz Z . phys. Chem. 1928 134,97 N. Bjerrum Thesis 1908 Copenhagen pp. 110-117. 146 SILLEN QUANTITATIVE STUDIES OF HYDROLYTIC EQUILIBRIA 147 data calculated K for AP+ Fe3+ and six other cations.As a result and perhaps also under the influence of the success of the Bronsted-Lowry acid-base concept in other fields it was thereafter often assumed that reactions such as (2) are the general explanation of cation hydrolysis. If this were so it would provide an extremely simple picture of the be- haviour of elemental ions in aqueous solution. Consider for instance the series of ions with 2s2p6 structure Na+ Mg2+ AP+ Si4+ P5+ S6+ C17+. The electrostatic field close to the ions increases strongly from Na+ to C17+ since the radius decreases and the charge increases; thus it would be increasingly hard for protons to remain near the central ion. The field of Na+ barely suffices to direct the surrounding water dipoles whereas Mg2+ and APf hold 6 water molecules strongly enough to bring them into the lattice when a salt crystallises.The increased repulsion of protons is illustrated by the species that may reasonably be assumed to exist in acidic and alkaline solution at pH = 0 Mg(H20)62+ Al(H20)63+ PO(OH), SO,-OH- c104- at pH = 14 Mg(H,O),OH+ A1(H2O),(OH),- Si02(OH)22- P043- Obviously the simple electrostatic picture goes a long way towards explaining the facts. However it is certainly an oversimplification. Already in 1908 Niels B j e r r ~ m ~ ~ had found that the hydrolysis of Cr3+ also gives polynuclear complexes i.e. complexes with several metal atoms Cr2(OH)24+ etc. This work was however for a long time not the centre of attention. Jander,5 by diffusion measurements later gave qualitative evidence for polynuclear complexes of many cations.Nevertheless and in spite of the fact that polynuclear complexes have long been recognised among the hydrolysis products from anions such as molybdate and silicate the idea of aquo-acidity with mononuclear products predominated for many years Previous work on hydrolytic equilibria of cations and anions is listed in the recent Tables of Stability Constants for inorganic ligands,g Tables 1-4 (ligand OH-); anionic hydrolyses will be found in Tables 5,6,7 8 and 50. The present Review deals mainly with work carried out in Stockholm in the last 10 years. Symbols and Equations.-The basis of all the work to be described is the law of mass action. Consider two reagents A and B (omitting the charges) which can form one or several complexes A,B, each with a formation constant Bpg.Let a be the concentration of free A and b the concentra- tion of free B. If the activity factors are kept constant for instance by using a concentrated ionic medium we may choose the standard states so For a review see G. Jander and K. F. Jahr Kolloid-Beih. 1936 43 295. J. Bjerrum G . Schwarzenbach and L. G. SillCn “Stability Constants Part 11 Inorganic Ligands” Chem. SOC. Special Publ. No. 7,1958 published under the auspices of I.U.P.A.C. Sod2- C104- 148 QUARTERLY REVIEWS that they are equal to unity and then use concentrations in place of activities. The law of mass action then gives [A,Bq] = JS,,[A]P[B]Q = ppqapbq . . . . (3) If B is the total concentration of B and 2 the average number of A atoms bound per B atom we have B = [B] + Cq[A,B,] = b + Cq/Ipqa”bq .. . . (4) BZ= Cp[A,B,] = ZpPpqapbq . . . . . . . . ( 5 ) We shall moreover define a convenient variable q 9 = log(B/b) = log(1 + Cq/3pqa~bq-1) . . . . . . .(6) For the special case of hydrolysis let A be OH- and B the metal ion; here however we shall still often leave out the charges of complexes. The general formula of a complex may conventionally be written as B,(OH),; it must be stressed at once that equilibrium measurements cannot distinguish between species which differ only in the number of solvent molecules or ions of the ionic medium. For instance it is not possible to distinguish from equilibrium measurements in a perchlorate medium alone between Fe2(OH)24+ Fe,04+ Fe2(OH)2C10,3+ or a mixture of these and similar species such as Fe2(OH),(H20).(C104-),4-~.To conform with the published Tables6 we shall denote by *Paq the equilibrium constant for the reaction written with (H20 - H+) as a reagent rather than OH-:? where h is the concentration of H+ Equations (4) (5) and (6) then are changed to the forms The problem is to find the sets of numerals p q that correspond to complexes present in appreciable amounts and the equilibrium constants *ISDq for their formation. 7 In this Review as in ref. 6 /3, is the equilibrium constant for the formation of A,B from the reagents A (the ligand) and B (usually a metal ion cf. eqn. 3). For mononuclear complexes A,B q = 1 and the second subscript is usually left out p2 being written instead of Pz1 etc. The squilibrium constant for step-wise formation of a mononuclear complex Ap-IB + A f A,B is denoted by K,.An asterisk on a p or a K denotes that the reactions are written with HA-H* as ligand instead of A (for instance in the F-Fe3+ system * p 2 is the equilibrium constant for Fe3+ + 2HF + FeF,+ + 2Hf). For OH- complexes HA is of course water *P, is the equilibrium constant defined by eqn. (7) and *K1 *K2 etc. are the acidity constants as in equations (8) and (9). SILLEN QUANTITATIVE STUDIES OF HYDROLYTIC EQUILIBRIA Experimental methods To provide numerical values for p q and */I in equations (4a-6a) we have to measure B 2 b and h over as wide a range as possible. B (= the total concentration of B) is always known from the amount of component B added to the solution. h (= [H+]) may often be measured with a hydrogen or quinhydrone electrode and almost always with a glass electrode.2 (= the average number of OH groups bound per atom of B) is obtained from h and the simple argument that the number of protons in the solution should be the same as that originally added. In many cases b and thus 7 [= log (B/b)] may be obtained by using cells with a metal (Agf Hg22+) amalgam (Cd2+ Pb2+ Bi3+ In3+) or redox (Hg2+ TI3+ Fe3+) electrode. The experiments are conveniently carried out as “titrations” to say 100 ml. of a solution S are made successive additions of a solution T (for practical reasons sometimes of two different solutions TI and T2). The compositions of the solutions are so chosen that B the concentration of B in the mixture remains constant whereas 2 is varied by the addition of acid or base. Moreover the ionic medium is kept as constant as possible usually by the aid of sodium perchlorate.After each addition the equili- brium values for h and if possible b are measured by suitable electrodes. After a small correction for the liquid junction potential Ei which is proportional to h the e.m.f.s give the concentration directly and may be calibrated with solutions of known h and b. It seems for instance that with 3~-(sodium) perchlorate as medium one may replace as much as 0.6 mole per 1. of Na+ by H+ before deviations in the activity factors of metal cations correspond to more than h0.2 mv in the e.m.f.s.’ All the data used for the calculations refer to clear solutions (no colloid or precipitate may be present) where it has been proved that the same values are obtained from whichever direction the equilibrium between the dissolved species is approached e.g.from higher or from lower 2. To increase 2 a base is added often OH-. Sometimes however e.g. for Fe3+ a local excess of OH- gives rise to a local precipitate which dissolves very slowly. It has then been proved practicable to add HC03- instead it is proved that complex-formation by carbonate is negligible by bubbling carbon dioxide and nitrogen alternately through the solution and observing the effect on the e.m.f.s8 By such titrations at constant total concentration of B and in a constant ionic medium but with varying 2 values (first used by I. Leden9 in a study of a number of cadmium complexes) it is possible to obtain in a limited time a much larger number of experimental results than could be obtained in the same time by the older “point-wise” method.Moreover from the course of a titration curve it is possible to correct for small errors in the 149 7 G. Biedermann and L. G. Sillen Arkiv Kemi 1953,5 425. * B. 0. A. Hedstrom Arkiv Kemi 1953,6,- 1_._ 150 QUARTERLY REVIEWS analysis for instance to determine accurately the excess of acid concen- tration in a metal-salt solution.l* This incidentally gives the titration method an advantage over the time-honoured method of measuring the pH of a solution of a “pure salt” (see the case of Fe2+ belowll). Treatment of results For each system studied it is essential to obtain a series of measurements as accurate as possible and over as wide a range of concentrations (B h) as possible. The results are conveniently displayed as in many cases in this Review as graphs of 2 against log h (or as graphs of 7 against log h) for specified values of total concentration of reagent B.These families of curves will be represented in the text below by Z(1og h)B and y(log h)B. If a certain set of *PDq is to be acceptable as the final solution of the problem the curves Z(1og h)B calculated by use of these constants must within the experimental error agree with the experimental results over the whole range studied. It is important that the whole range rather than a limited part of it shall be covered. Even when the numerical equations derived give acceptable agreement with the experimental results over the whole range the question still remains whether the explanation in terms of chemical reactions is unique or not. Therefore it has been necessary to devise mathematical and graph- ical methods for treating the results that are free from preconceived opinions as to what the complexes should be and to apply as many such independent methods as possible.For details of the mathematical and graphical methods the original papers12 should be consulted. Sometimes the set of curves gives important information at a first glance. For instance if the data inZ(1og h) coincide for different values of B and thus are independent of B then only one value of q is represented. In that case all the complexes present in appreciable amounts are homonuclear ; then as a rule q = 1 and the complexes are mononuclear. [If 7(log h) is independent of B the complexes must be mononuclear.] This is exceptional with OH- but common with other ligands; for OH- the 2 and 9 curves generally change with B so that polynuclear complexes must also be present.If the curves are parallel with a constant spacing ( A log B)/ (A log h)Z = t then Z and 7 are functions of the single variable x = log B - t log h and it can be shown that all complexes present in appreciable amounts can be written in the form B((OH),B), the “core + links” formula.12a However it requires a more accurate analysis of the curves to find whether lo C. Berecki-Biedermann Arkiv Kemi 1956,9 175. l1 B. 0. A. Hedstrom Arkiv Kemi 1953,5,457. l2 L. G. Sillen Acta Gem. Scund 1954 8 (a) 299 (6) 318; S. Hietanen and L. G. Sillkn ibid. p. 1607; B. 0. A. Hedsiiorn ibid. 1955,9 613; G. Biedermann and L. G. Sillen ibid. 1956 10 1011; F. J. C. Rossotti and H. S. Rossotti ibid. 1955 9 1166; L.G. Sillhn ibid. 1956 10 186; F. J. C. Rossottl H. S. Rossotti and L. G. Siilen ibid. p. 203; L. G. Sillen ibid. p. 803. SILLEN QUANTITATIVE STUDIES OF HYDROLYTIC EQUILIBRIA 15 1 only one or two values of n are represented or whether there are several such perhaps a series of complexes.12* In the following pages our results for several systems will be discussed. The chemical picture will be stressed and for some systems the degree of agreement between experiment and theory will be indicated. Since the same species and equilibrium constants have been deduced by several independent mathematical approaches-often also several independent experimental methods-it seems that the main reactions are generally well established; it is however possible that species existing in minor amounts escaped detection.Cationic systems Mononuclear Species.-Mercury(i1) is one of the few ions which on hydrolysis fulfil the criterion for mononuclearity 2 and 7 are functions of log h only and independent of B. The two independent sets of data Z(1og h) obtained by use of a glass electrode and r(log h) obtained by use of a redox electrode may be explained by an acid dissociation in two steps; the same values are obtained for the acidity constants :13 HzOHgOHg2+ + HOHgOHz+ + H+; *Kl = */I1 (= */Ill) . (8) HOHgOHz+ + HOHgOH + H+; *Kz = *Pz/*/I1 (= */3zl/*/311) (9) The symbols given for the equilibrium constants are those used in the Tables of Stability Constants6 (cf. footnotet on p. 148). For other polyprotic acids such as H,PO and HzS04 the ratio between successive acidity constants is often around For PO(OH), for instance the logarithms of the constants are -2.1 -7.2 and -12.3.-This is easily understood it must become harder to pull protons from molecules of increasing negative charge PO(OH)3 P02(OH)z- P030H2-. By analogy one might perhaps expect that for mercury(r1) a proton would be more easily split from an ion of charge +2 than from one of charge + l . The first acidity constant *Kl (at 25"; in 0.5~-NaClO,) is lO3*' so the second might be expected to be of the order lo-*. In fact it is 10-2*6 and this is a case where the second dissociation constant of an acid is greater than the first. This means that the following disproportionation is favoured 2HgOH+ + Hg(OH)2 + Hg2+; K = *Kz/*K,=lO1*l . (10) Fig. 1 shows the distribution of HglI and for comparison Pv over various complexes for variations of log h.It is seen that the fraction of mercury(r1) present as the intermediate complex HgOH+ is at most about 14%. If we assume that there are N equivalent possible sites for OH groups on a HgU atom and that the probability that one of these is oc- cupied by an OH is independent of whether the other sites are occupied lS S. Hietanen and L. G. Sill& Actu Chem. Scund. 1952 6 747. 152 QUARTERLY REVIEWS or not then one niay derive the “statistical” value for the equilibrium constant K in equation (10). We have 5 0 - 0 - x - If we assume N = 2 we find that log Kstat. = log 1/4 = - 0.6; for higher values of N log Kstat. would have values between - 0.6 and - 0.3. 3 ‘ I I I 1 1 I I I 1 I 1 1 0 2 4 6 8 10 /2 14 PH FIG. 1. Distribution of Pv and HgII over diferent species in solutions of various pH.For each pH one can draw a vertical line; the section of this line falling within theJield of each species is proportional to the amount of that species present at equilibrium e.g. at pH = 3-0 we have 12% of H3POa and 88% of H2P042-; at the same pH we have 59 % of Hg2+ 12 % of Hg(OH)+ and 29 % of Hg(OH) 2. 1 Hg2+; 2 Hg(OH)+; 3 Hg(OH),; 4 H3P04; 5 H2P04-; 6 HPO,,-; 7 P043-. [Reproduced with permission from J. Inorg. Nuclear Chem. 1958 8 193.1 In fact although the electrostatic effects would have tended to decrease log K it is greater than the statistical value. (Our conclusion was in- cidentally later confirmed by work in Schwarzenbach’s school at Ziirich.l4) It appears at first sight possible to relate the high value for K with the sp-bonds in Hg(r1) complexes to consider that the two symmetrical species are favoured by mesomerism but this view cannot be maintained in face of the accumulated data.The Table (p. 167) is a survey of equilibrium constants for mono- and di-nuclear hydroxo-complexes studied at Stockholm. Some of the products are only secondary while in the available concentration range polynuclear complexes predominate. In other cases however the data have been accurate enough to afford *K and *K2 independently and so to give K in equation (10). It is true that TP+ and In3+ with electronic structures similar to Hg2+ also have log K greater than the statistical value; but so have Fe3+ and Sc3+ with very different l4 G. Anderegg G. Schwarzenbach M. Padmoyo and 0. F. Borg Helv.Chim. Acta 1958 41 988. SILLEN QUANTITATIVE STUDIES OF HYDROLYTIC EQUILIBRIA 153 electronic structures. It seems therefore that this behaviour is typical of OH2 acidity as distinguished from the OH acidity of say PO(OH)3 and S02(0H),. Perhaps the second complex is not M(OH)2 but rather MO FeO+ HgO InO+ etc. (cf V02+). One specific case that of the iron(rr) ion has general interest. An early inve~tigatorl~ very carefully applied what was then the standard method for cation hydrolysis. He prepared a very pure specimen of iron@) perchlorate recrystallising it many times in the absence of air and measured the pH of solutions of this salt in air-free water. He then calculated *K (the equilibrium constant for Fe2+ + H,O + FeOH+ + H+) as [Other workers,15" using pure solutions of iron (11) chloride found a value of 10-7.9.] The same reaction was studied by Bengt Hedstrom,ll by what has become our standard method.He prepared a solution containing iron(I1) and perchlorate ions and a small amount of H+ in excess. From the changes in log h on addition of increasing amounts of base he calculated accurately the amount of H+ in the solution; the part of the titration curve where Fe2+ and FeOH+ were both present gave *Kl namely 10-9*5 iron(rr) is thus a much weaker acid than was previously thought. It is not hard to explain the discrepancy it is difficult to avoid presence of small amounts of H+ and Fe3+ which may be adsorbed on the crystals of Fe(ClO,),; although the resulting solutions do not change their pH after a number of recrystallisations the acidity is then much too high.Unfortunately most older work and some more recent on cation hydrolysis is founded on measurements of pH in "pure salt" solutions. Dinuclear and polynuclear complexes Iron(IrI).-The hydrolysis of the iron(rrr) ion has been studied by many workers,6 who however have until recently usually tried to explain their results in terms of reaction (2) only with an equilibrium constant *Kl. Now by various methods it is possible to obtain approximately constant values for *Kr over limited ranges of h and total iron concentration. However accurate measurements over a wide concentration range show that the calculated *K varies by more than can be accounted for by activity factors or by a second mononuclear complex Fe(OH),+. From deviations in both spectrophotometric16 and magnetic rnea~urements~~ it was concluded that there must exist some other species presumably a polynuclear complex Fe2(OH)y(33C-Y)+.Bengt Hedstrom* in Stockholm studied the hydrolysis of iron(1rr) using glass and redox electrodes and deduced very straightforwardly from both sets of results that the main product in his experimental conditions [ 3 ~ (Na)ClO,; 25"] had the formula Fe,(OH),4+ and that two mononuclear l6 F. Lindstrand Diss. Lund 1939. 150 K. H. Gayer and L. Woontner J. Amer. Chem. SOC. 1956,78 3944. l6 T. H. Siddall tert. and W. C. Vosburgh J . Arner. Chem. SOC. 1951 73,4270. l7 P. W. Selwood personal communication (1952). 154 QUARTERLY REVIEWS complexes occurred as by-products becoming important at low concen- trations Fe3+ + H20 + FeOH2+ + H+; log “KI = -3.05 FeOH2+ + H20 + Fe(OH),+ + H+; log *K2 = -3.26 log *Bz2 = -2.91 2Fe3+ + 2H20 + Fe,(OH),4+ + 2H+; Once the formulae of the complexes and the approximate equilibrium constants were known it was possible to explain the magneticla and spectrophotornetric results,lg though it seems that these methods alone did not permit determination of x and y in Fe,(OH) with certainty.0.8 0- 6 ry 0.4 0.2 0 -2 -3 - 4 -5 -6 jog h Average number Z of OH bound per Be as a function of log h. All experimental points are given. Completely filled or open symbols represent points from diferent titrations half-filled symbols represent points from back-titrations (decreasing Z ) . A hydrogen electrode was used for and a quinhydrone for other symbols. Full curves are calculated by assuming Bes(OH)sa+ as the only complex with log * p33 = - 8-66.Broken curves are calculated with the constants given in the text by assuming Be20H3f Be(0H) 2 and Be3(OH)s3+ to be present. Total Be concentration; B; 1 48 m ~ ; 2 19 mM; 3 10mM; 4 5 m; 5 2.5 m ~ ; 6 1 mM. [Reproduced with permission from Acta Chem. Scand. 1956 10 990.1 The Table lists also a number of other ions where a complex B,(OH)2 or B 2 0 is formed either as main product or together with mono- or poly-nuclear species. The column “ t KZ2” gives the “dimerisation con- stant” of BOH to B,(OH)z. This often has a fairly high value in spite of L. N. Mulay and P. W. Selwood J. Amer. Clzem. SOC. 1954,76,6207; 1955,77,2693. FIG. 2. l@ R. M. Milburn and W. C. Vosburgh J. Amer. Chern. Soc. 1955,77 1352. SILLEN QUANTITATIVE STUDIES OF HYDROLYTIC EQUILIBRIA 155 the fact that the reaction is an association of two positive ions.The fact that strong forces are involved is demonstrated by Mulay and Selwood's discovery18 that the dinuclear Fe 2(OH) 24+ is diamagnetic whereas the mononuclear FeOH2+ is strongly paramagnetic like Fe3+. Beryllium.-It has been known since Berzelius's time that solutions of beryllium salts remain clear on the addition of up to 1 OH- per Be2+. Previous work on this reaction can be divided into two groups that20 in which the main product is held to be Be,(OH)22+ (or its equivalent Be202+) and that21 in which it is claimed as Be4(OH)44+ (or Be4024+). The hydrolysis of beryllium (in ~ M - N ~ C ~ O at 25") was studied by modern methods by Kakihana and whose results [in the form of Z(1og h)B curves] are reproduced in Fig.2. This Figure shows that Z increases (ie. more OH is bound per Be) with decreasing log h (Bconstant) and with in- creasing B(1og h constant). The first result must be so; the second indicates that there are polynuclear products. The detailed measurements are incom- patible with the assumption of either Be2(OH)2z+ or Be4(OH)44+ as the main product whereas mathematical analysis of them strongly indicates Be3(0H)33+. Assuming this as the single product gives fair agreement over a fairly wide range (full curves in Fig. 2) but deviations at the two ends of the curves lead to the formulae and formation constants of two further complexes present in minor amounts (but still not those assumed by pre- vious workers). The broken curves in Fig. 2 have been calculated by assuming the following equilibria :22 3Be2+ + 3H,O + Be,(OH),3+ + 3H+; log *p33 = -8.66 log *p12 = -3.24 log *p2 = -10.9 2Be2+ + H 2 0 + Be20H3+ + H+; Be2+ + 2H20 + Be(OH) + 2H+; Fig.3 is a distribution diagram like Fig. 1. For each species in the solution a field is given. If a vertical line is drawn at a specific value of log h the section of this line falling within each field denotes the fraction of total beryllium present as the corresponding complex at the log h in question. Since polynuclear complexes are formed the distribution depends also on the total concentration B diagrams are given for B = 1 10 and 100 milliniolar. The predominance of Be3(OH)33+ suggests that it contains a ring of 3Be and 30H with 6H20 to fill out the co-ordination tetrahedra of beryllium. The other two complexes have narrower fields of existence; it seems likely that other complexes might also appear presumably with still narrower ranges if more sensitive experimental methods were available.ao M. Prytz 2. anorg. Chem. 1929,180 355; 1931,197 103; R. A. Gilbert and A. B. Garrett J. Amer. Chem. Soc. 1956 78 5501. 21 R. Schaal and J . Faucherre Bull. SOC. chim. France 1947 927; J. Faucherre ibid. 1953 1117; 1954 253; P. Souchay ihid. 1948 143; M. Teyssidre and P. Souchay ibid. 195 1,545. 2 2 Ref. u of the Table p. 167. 5 156 QUARTERLY REVIEWS Tin(II).-In early work on the hydrolysis of Sn2+ it was assumed that either23Q the mononuclear SnOH+ the dinuclear Sn2(OH)2+ (or Sn202+) was formed. This system was studied by Tobias24 who measured h with glass electrodes and b (= [Sn2+]) with tin amalgam electrodes.The / 2 0.5 F 0 -2 -3 -4 -5 log h -6 -2 -3 -4 -5 -6 log h 100 2 F 0 -2 -3 - 4 -5 -6 log h Distribution of beryllium over complexes for total beryllium concentra- FIG. 3. tion B = 1 10 and 100 millimolar as a function of log h. 1 Be2+; 2 Be,0H3+; 3 Be3(OH),3+; 4 Be(OH),. [Reproduced with permission from Acta G e m . Scad. 1956,10 1002.1 two sets of data 2 (log h)B and 7 (log h)B (Figs. 4 and 5) independently gave the same set of species and equilibrium constants namely Older values 3Sn2+ + 4H20 + SII~(OH),~+ + 4H+; log *Pa3 = -6-77 - 2Sn2+ + 2H20 $ Sn,(OH),2+ + 2H+; log *pz2 = -4-45 *30 M. Gorman and P. A. Leighton,J. Amer. Cliem. SOC. 1939,61,3342; 1942,64,719; A. B. Garrett and R. E. Heiks ibid. 1941,63 562; C . E. Vanderzee and D. E. Rhodes ibid. 1952,74,3552,4806; (b) M.Prytz Z. arzorg. Chem. 1928,174,355. 84 Ref. fof the Table p. 167. --3 Sn2+ + H 2 0 + SnOH+ + H+; log *Kl = -3.9 -1.7 SILLEN QUANTITATIVE STUDIES OF HYDROLYTIC EQUILIBRIA 157 The solid curves in Figs. 4 and 5 were calculated by assuming only Sn3(OH)42+ which suffices to account for a large part of the data fairly well. The broken curves were calculated by.assuming also the two secondary products. FIG. 4. Hydrolysis of tin@); glass electrode used for measurements. Average number Z of OH bound per Sn as a function of log h (molar scale); B as a para- meter. Full curves are calculated by assuming Sn,(OW)42+ as the only complex with ldg *j143 = - 6.69. Broken curves are calculated by also assuming the presence of Sn,(OH),2+ and SnOH+ and using equilibrium constants given in the text.Sna+ concn. 1,4040 mM; 2 20.00 mM; 3 10.00 mM; 4 5.00 mM; 5 2-50 mM. [Reproduced with permission from Acta Chem. Scand. 1958 12 205.1 The two minor products happened to be the same as those proposed earlier as main products. Knowing the species and the approximate equilibrium constants Tobias recalculated the old data. In most of the old work the tin@) solutions must have been much too acidic because of a rather high content of tin(rv) which is harder to avoid than might be expected. The purest tin(r1) solutions seem to have been those of Milda P r y t ~ ~ ~ ~ who concluded correctly that a polynuclear complex was formed; but because she used too narrow a concentration range she missed the correct formula. The fact that Tobias obtained the same reactions and the same equili- brium constants by two independent experimental methods24 justifies some confidence in his results.Bismuth(m).-The hydrolysis of bismuth(I1r) was studied by Olin25 who measured b(= [Bi3+]) and thus 7 using bismuth amalgam electrodes. as Ref. j of the Table p. 167. 158 QUARTERLY REVIEWS His results (Fig. 6) indicate that the main product is a hexanuclear complex of charge 6+ Bi6(0H)1z+ (or e.g. Bi6066+). The deviations at the lowest 1 2 3 4 5 0 - l o g h FIG. 5. Hydrolysis of tin@); amalgam electrode used for measurements. q = log B/b as a function of log h B as a parameter. Curves are calculated for the same sets of constants and same Sn2+ concns. as in Fig. 4. values of 7 give evidence for another species BiOH2+. The curves in Fig. 6 were calculated by assuming the equilibrium constants [25"; 3~-(Na)C104 [Reproduced with permission from Acta Chem.Scand. 1958 12 205.1 6Bi3+ + 12H20 + Bi6(OH)1:+ + 12H+; log *p12,6 = 0.33 Bi3+ + H20 + BiOH2+ + H+; log "K = -1.58 The agreement is as good as could be desired in a range of B 0.1-50 millimolar i.e. a ratio of 1500. Since the solutions are rather acidic Z is a difference between large numbers and is not very accurate; the agreement however is good within experimental error. For Z> 2.0 slightly larger complexes are formed probably26 with 9Bi atoms and charges 5+ 6+ and 7+. About twelve years ago Grankr and the Re~iewer,~' studying the hydro- lysis of Bi3+ concluded that a series of species Bi{ (OH)2Bi}n3+n was formed. This also seemed to give fair agreement with the experimental results though not as good as that in Fig.6. Two factors contributed to the shortcoming. Tn this early work the concentration B was varied only five- fold (from 10 to 50 millimolar). Secondly there was a systematic un- expected experimental error the quinhydrone electrode which was used 2B Unpublished work at Stockholm. 2i F. Graner and L. G. Sillkn Acta Chem. Scand. 1947 1 631. SILLEN QUANTITATIVE STUDIES OF HYDROLYTIC EQUILIBRIA 159 for certain standardisations deviates from the ideal formula at higher acidities because p-benzoquinone is a base and attracts a proton.28 The resulting error (maximum 1-2 mv) shifted the experirncntal points so that the distance between the 7 curves corresponded to the general formula Bi{ (OH)2.0Bi},3+n rather than Bi( (OH),.qBi)n3+0‘6n. log n FIG.6. Hydrolysis of bismuth(@. 77 = log B/b as a function of log h. The curves are calculated by assuming the presence of the complexes Bi,(OH),,6+ and BiOH2+ with equilibrium constants given in the text. For clarity not all of the experimental points are shown for low values of v. 1,O.l mM; 2,0.5 mM; 3 l mM; 4,25 mM; 5 5 mM; 6,lOmM; 7 5 0 m ~ . [Reproduced with permission from Acta Chem. Scand. 1957,11,1452.] This experience stresses the importance of using a broad concentration range. When Olin later studied Bi3+ over a wider range of B a deviation was found which was traced back to the basicity of quinone. It should be mentioned however that the mechanism which was first proposed for bismuth and was erroneous in that case sufficed to explain the results very well for a number bf other ions especially indium thorium uranyl and scandium.Series of complexes For the cations discussed above one or at most two polynuclear species have been found; if two one of them has had a much wider “range of existence” in distribution diagrams such as Fig. 3 { Be3(0H)32+ Sn,(OH),2+}. For a number of other systems the predominating com- plexes seem to have the general formula B((OH),B}, with t constant and there is no preferred value of y1. Good agreement with experiment is obtained 28 G. Biedermann Act0 Chem. Scand. 1956 10 1340. 160 QUARTERLY REVIEWS by assuming that complexes with all values for n = 1,2,3 . . . are formed and that the equilibrium constants for their formation vary with n in some regular manner usually linearly. For In3+ BiedermanP' deduced the following equilibrium constants (as well as those for the mononuclear complexes see Table) from two independent sets of measurements Z(1og h)B and $log h)B (glass or quin- hydrone and In-Hg electrodes) In3+ + 2H20 + In3+ $ In(OH),In4+ + 2H+; log = -5.21 In{(OH)21n),n+3 + 2H20 + In3+ + In{(OH)$n);:f + 2H+; log K = -4.69 for r~>0 or in general (n + l)In3+ + 2nH20 +- 1r1{(0H),In),~+~ + 2nH+; -0.52 - 4 .6 9 ~ ~ I 1% *P2n,n+1 - The following reactions have been deduced exclusively from Z(1og h)B curves (n + l)U022+ + 2nH20 $ U02((OH)2U02),2+ + 2nH+; (n + l)Sc3+ + 2nH20 (n + l)Th4+ + 3nH,O log*p2n,n+l = 0-30-635n ref. 30 log *p2?a,n+l = 0.70-6.87n ref 31 log *&@+I - -7050n ref. 32 + SC((OH),SC},@+~ + 2nH+; + Th((OH)3Th},4+n + 3nH+; - It should be mentioned that there is evidence for other complexes as well namely (U02),0H3+ Th2(OH) 26+ Th20H7+ScOH2+ Sc(0H) ,+(Table refs.g and k) and that the equations for the equilibrium constants /3tnn,+l must be understood as approximations only. This simple approximation in spite of its having only two arbitrary constants gives fairly good agreement with experiment; it might perhaps 'have been improved by introducing more than two adjustable constants. Structure of the complexes When one complex predominates it seems reasonable to assume a closed structure (cyclic tetrahedral or octahedral etc.) and attempts are being made to determine their crystal structures. For aluminium earlier equilibrium studies by Cyril1 B r o s ~ e t ~ ~ indicated rather larger complexes contrary to previous opinion. Alkali can be added to a solution of alumin- ium chloride up to a ratio of 2.50H- per A13+ without causing permanent Ref.h of the Table p. 167. The mononuclear complexes were also studied by a distribution method by F. J. C. Rossotti and H. S. Rossotti Acta Chem. Scand. 1956 * 10 779. 30 Ref. c of Table p. 167. 31 Ref. g of Table p. 167. 3a S. Hietanen Acta Chem. Scand. 1954 8 1626. s5 C. Brosset G. Biedermann and L. G. SillCn Acta Chem. Scand. 1954 8 1917. SILLEN QUANTITATIVE STUDIES OF HYDROLYTIC EQUILIBRIA 161 precipitation. Such solutions have important cosmetical uses. The alu- minium complex they contain has been conventionally written as Al,(OH),+; it is formed rather slowly but once it is formed one may precipitate well-crystallised salts with e.g. Se042- and C2042-. In the hope of finding the structure of the complex “A12(OH),+” some of these salts were studied by X-ray methods.They were found26 to contain also alkali metal; a typical formula is NaA1,3(0H)32(S04)4(H20)z. The crystal structure of a corresponding selenate has been studied thoroughly by Georg Johansson.26 Aluminium was found to form separate groups with 13 A1 and charge +7; the formula conventionally Al13(OH)327+ is really A1130POH487+ and the structure is remini~cent~~ of that of dodecahetero- tungstates such as PW120403-. The equilibria in AP+ solutions are now under study by Biedermann and it seems probable that these large All ions will form a part of the final picture. In crystal structures units U604(OH)412+ and Ce6O4(0H),l2+ have been found35 as separate building-stones they contained an octahedron of U (or Ce) and a cube of 0 atoms.Recent X-ray of solutions seem to indicate that the “Bi6(OH)126+” complex may have a similar structure Bi604(0H)4“+. However no direct evidence has yet been found for existence of the Ce and U complexes in solution. n = l n =2 n-3 n - 4 0 0. 0.0 0 0 0 0 0 0 0 0 0 0 . 0 0 . 0 0. 0.0 0 0 0 n-5 0.0 0 0 0 . 00.00 0 0 0 . n - 6 a O . 0.0. O . O . O 0 0 0 0 0 0 0 FIG. 7. Possible structures of complexes U02[(OH)2U02],2+ with n = 1 to 6 shown as increasing fragments of a sheet from a- UO (OH),. @ uranyl groups with U in the plane of the paper and the two uranyl0 (not shown here) above and below perpendicular to the paper 0 oxygen atoms above (thick) and below (thin) the plane of the paper. [Reproduced with permission from Acta Chem. Scand. 1954,8 1914.1 In those cases where a series of complexes is formed it seems reasonable to assume open structures chains or sheets so that addition of each new “link” has approximately the same equilibrium constant.It is tempting to make comparisons with crystal structures of compounds which are known to be precipitated from such solutions. For instance Fig. 7 shows the geometry of the sheets in the crystal structure26 of cc-U02(OH)2. 34 J. F. Keggin Proc. Roy. SOC. 1934 A 144 75. 36 G. Lundgren Arkiv Kemi 1953,5 349; 1956,10 183. 36 H. Levy and K. A. Kraus 1958 personal communication. 162 QUARTERLY REVIEWS The black dots are the U atoms of UOZ2+ groups; the two oxygen atoms are one above and one below the plane of the paper. The open circles are OH groups in two layers one above and one below the level of the U atoms.If well-rounded fragments are cut from such a sheet their formulae will be (UO,),+,(OH),, as seen for n = 1-6 in‘Fig. 7. From the e.m.f. data the same formula was deduced for the complexes in solution and it is attractive to regard them as fragments of the UO,(OH) crystal structure pre-formed in the solution. Similarly the series of indium and thorium complexes may be related to the chains that have been found in the crystal structures of “basic” thorium3’ and indium Their composition would then be In(OIn),n+3 and Th(ogTh),n+4; by equilibrium measurements in aqueous medium one cannot distinguish between species with 20H- and lo2-. The assumption that Th4+ on hydrolysis forms chain-like complexes is supported by recent work on colloidal thorium hydr~xide,~ especially electron-microscopy which shows the particles to be threadlike.Anionic hydrolysis General.-Borates. It has long been known that polynuclear species are formed in certain anionic hydrolytic systems examples are borates germanates tellurates chromates molybdates tungstates and vanad- ates. There is no fundamental difference between anionic and cationic systems. The same graphical and mathematical approaches can be used for both and the experimental methods are similar except that in anionic systems as a rule it is necessary to work at higher pH and there are some extra experimental dficulties (e.g. exclusion of carbonate). In some systems with polyions both cations and anions are formed in the ordinary pH range for example with aluminium vanadium(v) and molybdenum(m).One of the first systems studied by our methods was the b o r a t e ~ . ~ ~ Besides the mononuclear species B(OH) and B(OH)g polyborate ions are known to exist especially at high total concentrations; they have previously been described in general as B40,2-. From the Z(1og h)B curves it was however deduced that in the more acidic range the predominant polynuclear species is a triborate of charge - 1 presumably B303(OH)4- a 6-membered ring compound. For higher 2 the accuracy of the equili- brium measurements in ref. 40 did not permit distinction with certainty between B303(OH)52- and Bp05(OH)42-; it is possible that both exist in solution and both have been claimed to exist in crystal lattices.41 Recent work by a “self-medium” method26 indicates that the tetraborate ion of ’ 37 G.Lundgren and L. G. Sillen Arkiv Kemi 1949,1,277; G. Lundgren ibid. 1950 2 535. 38 H. E. Forsberg Actu Chem. Scund. 1956 10 1287; 1957 11 676. 3B A. Dobry S. Guinand and A. Mathieu-Sicaud J. Chim. phys. 1953 50 501. 40 N. Ingri G. Lagerstrom M. Frydman and L. G. SillCn Acru Chem. Scund. 4 1 C. L. Christ and J. R. Clark Actu Cryst. 1956,9 830; N. Morimoto Mineralog 1957,11 1034. J.,J956 2 1. SILLEN QUANTITATIVE STUDIES OF HYDROLYTIC EQUILIBRIA 163 charge -2 and the triborate ion of charge - 1 are the two predominating polyions . measured by Sa~aki,~ [25"; 3hl-NaC1O4]. In this case it was practicable to let 2 be the average number of H+ bound per [A = H+ B = in equations (3-6)]. A related quantity is z = 2 - 2 the average charge per Mo atom. Molybdates. Fig. 8 gives Z(1og h)B curves for hydrolysis of -7 -6 -5 -4 -3 /oY[H'] Hydrolysis of molybdate ions.Average charge per Mo atom (z) plotted against log h. Points are for diferent total concentrations B of Mo. Circles with crosses in them (curve 2) refer to back-tritations. The curves are calculated by assuming the species and equilibrium constants given in the text. 7 2.5 mM; 8 1.2 m ~ ; 9 0.6 mM. [Mo] total 1 160 mM; 2 80 mM; 3 40 mM; 4 20 mM; 5 10 mM; 6 5 mM; [Reproduced with permission from J. Inorg. Nuclear Chem. 1959 9 94.1 At low log h 250 (z = - 2) at all values of B since is the pre- dominant species. The shape and spacing of the curves show conclusively that the first polynuclear complex formed is one with 7Mo and a charge of 6-. It is reasonable to identify this species with Mo@246- the para- molybdate ion which exists42a as a separate unit in the crystal structure of and Mo ,O 246-.Deviations at low concentrations of molybdenum indicate the presence of HMoO,- and at higher Z that of HMO,O,,~- The curves in 42 Y. Sasaki I. Lindqvist and L. G. SillCn J. Inorg. Nuclear Chem. 1959,9,93. d2a I. Lindqvist Arkiv Kemi 1950 2 325; Nova Acta Regiae SOC. Sci. Upsaliensis (NH4)6Mo,02dH20)4* Most of the results in Fig. 8 require only the two species 1950 C1. IV 15,l. 1 64 QUARTERLY REVIEWS Fig. 8 were calculated by assuming the following equilibrium constants (there is as yet no standard symbolism for such constants) + H+ + HMo04-; log K = 4.08 7M0042- + 8H+ + + 4H20; log K = 57.7 MO7Oz4'- + H+ $ HM0702g5-; log K = 4.33 HM0702g5- + H+ $ H2M070244-; l o g K = -3.7 Agreement is quite good.The species HMo702a- is well supported; some doubt may be expressed about H2M07024P in spite of the very good agreement since at the end of the range there must anyhow be some un- known species.42a 2 3 4 6 - / O g b FIG. 9. Hydrolysis of vanadium(v). Average number of hydroxyl groups 2 bound to each VO,+ ion as a function of log h for different total concentrations B and vanadium. The curves are calculated by assuming the presence of the species and equilibrium constants given in the text. [V] total 1 0.0200 M; 2 0.0100 M (+ reverse titration); 3 0.0050 M ; 4 0.0025 M. [Reproduced with permission from Acta Chem. Scand. 1956 10 964.1 Vanadium(v). Fig. 9 gives curves for vanadium(v) hydroly~is;~~ Z is the average number of OH- bound per V02+ ion. In highly acidic solutions VO,+ seems to be the only species (2 = 0); increasing the acidity from 0.05 to ~M-H+ does not change the vanadium absorption spectrum.With decreasing h 2 increases up to -1.4 and the solution becomes orange-coloured. An uncharged species corresponding to Z = 1 (V020H or HV03 or V205 etc.) might have been expected but there has been no evidence of its presence. The data indicate that the main reaction is the formation of a single polynuclear species namely an anion with 10 vanadium atoms. Assuming a complex with 11 vanadium atoms gives 43 F. J. C. Rossotti and H. Rossotti J. Inorg. Nuclear Chem. 1956 2 201; Acta Chem. Scand. 1956,10,957. STLLEN QUANTITATIVE STUDIES OF HYDROLYTIC EQUTLIBRIA 165 acceptable agreement for 2 = 0--1-1 but deviations at higher 2 values. With V or‘Vlz complexes neither the spacing between the curves nor their shape gives acceptable agreement with experiment.Moreover X-ray studies4* of orange vanadates crystallised from similar solutions indicate a multiple of 10 V per unit cell. The full curves of Fig. 9 were calculated by assuming the following equilibrium constants [ lw(Na)Cl O4 ; 25’1 lOVO2+ + 8H20 f H2Vlo02,4- + 14H+; log */114,10 = -6.75 H2V100284- f HV100285- + H+; HV100285- + V,,0286- + H+; log K = -3.6 log K = -5.8 The agreement is satisfactory. The formulae are written in the conventional way with as little water as possible. Structural studies may show them to contain more water. From a study26 of the alkali side of vanadate equilibria the following equilibrium constants have been deduced [0*5~-NaC1; 25’1 V04” + H+ + HVO,”; HV042- + H+ + H2V07-; log K = 12.6“ log K = 7*7b 2HV042- + H+ + HV2073- + H20; log K = 1 0 ~ 6 ~ 3HV0,2- + 3H+ S V30g3- + 3H2O; log K = 3007~ a By spectrophotometry.By e.m.f. measurements. Fig. 10 gives a schematic survey of all the vanadate equilibria the average charge z per V is plotted as a function of log h for different total concentrations of vanadium. We may neglect the difference in ionic medium between different parts of the diagram. In each of the ranges z = + 1 to - 0.6 and - 1 to - 3 there are rapid reversible equilibria. As long as the solution is kept within one of these regions equilibrium is obtained practically instantaneously on change of composition. If however a solution or a part of it has once been brought into the region z between - 0.6 and - 1.0 it requires a long time to attain equilibrium.This “instability range” may be one reason for difficulties that some pre- vious authors have experienced with vanadate equilibria. Similar instability ranges have been observed also for silicate equilibria.26 It is pertinent to ask why vo43- HV042- HV2073- and V30g3- pre- dominate of all the conceivable species. It may be these formulae seem arbitrary because they are written on the pattern of ortho- pyro- and meta-phosphates. It is true that some vanadates are isomorphous with phosphates indicating a tetrahedral V043- ion. However the more common co-ordination number of vanadium(v) seems to be 5 two oxygen atoms being bound in a linear group OVO and three more lying in the equatorial 44 I. Lindqvist personal communication; H. T.Evans jun. M E. Mrose and R. Marvin Arner. Min. 1955 40 314. 166 QUARTERLY REVIEWS v,o,3- FIG. 10. Schematic survey of the hydrolysis of vanadium(v). Average charge per V atom z plottedagainst log h. The lower leff part corresponds to Fig. 9 the upper right part to unpublished measurements by Ingri and Brito.26 Between them is an “instability range” with slow equilibria. plane. One may then imagine the following structures for the three species that predominate between z = - 1 and -2 (B = V02+ A = OH-) A A A A B A B V02(OH)32- = HV042- AB A (V02)2(OW,3- = HV,O,~- A A A B A A A B A B The equilibrium data as usual give only the average charge of a species and the number of vanadium atoms but not the amount of water. So there is no way of distinguishing between say HV042- and V02(OH)32- by equilibrium data alone.There is however no proof for these formulae. SILLEN QUANTITATIVE STUDIES OF HYDROLYTIC EQUILIBRIA 167 Conclusion In the systems discussed above it seems that the e.m.f. methods when applied over a wide concentration range and with a constant ionic medium have been able to give independently of other data sufficient information Logarithms of equilibrium constants for cations with mono- and di-nuclear Ion "K "K2 *P22 "K2I"K tK22 Other p q Ref. hydrolysis products. Be2+ vo2+ uo,2+ Fe2+ CU2f Cd2+ Hg22+ Hg2+ Sn2+ sc3+ Fe3+ 1n3+ TP+ Bi3+ Th4+ u4+ ("P2 - 6.0 -9.5 -9.0 - 5.0 -3.70 - 3.9 -5.1 -3.05 -4.4 -1.14 -1.58 -2.0 - (-8) I - 10.9) - - - - - - -2.60 I -5.1 -3.26 -3.9 -1.49 - - I - - 6.9 - 6.05 - - 10.6 - - - -4.45 - 6.2 -2.91 - 5.2 - - -4.7 - 3,3; 1,2 2n,n+ 1 - - - - - - 473 2n,n+ 1 2n,n+ 1 12,6 3n,n + 1 3n,n + 1 - - a b 11 10 d e 13 f g 8 h i j k 1 C The medium is usually 3~-(Na)clO, and the temperature 25" ; 1M-(Na)C1O4 for UOZ2+ Fez+ and Sc3+; 0.5~-(Na)C10 for Hg,2+ and Hg2+.For UOz2+ the temperature was 20". (a) H . Kakihana and L. G. Sillen Acta Chem. Scand. 1956 10 985. (b) F. J. C. Rossotti and H. S. Rossotti ibid. 1955 9 1177. (c) S. Ahrland S. Hietanen and L. G. SillCn ibid. 1954 8 1907. ( d ) Y . Marcus ibid. 1957 11 690. (e) W. Fording S. Hietanen and L. G. SillCn ibid. 1952 6 901. ( f ) R. S. Tobias ibid. 1958 12 198. (g) G. Biedermann M. Kilpatrick L. Pokras and L. G. Sillen ibid. 1956 10 1327. (h) G. Biedermann Arkiv Kemi 1956,9,277; Rec. Trav. chim. 1956 75 716. (i) Idem Arkzv Kemi 1953 5 441; Rec. Trav.chim. 1956 75 716. ( j ) A. Olin Acta Chem. Scand. 1957,11 1445; F. GranCr A. O h and L. G. SillCn ibid. 1956 10 476. (k) S. Hietanen and L. G. Sillen unpublished work. (I) S. Hietanen Acta Chem. Scand. 1956 10 1531; Rec. Trav. chim. 1956,75,711. to allow reliable conclusions about the species present and the equilibrium constants. Even if the e.m.f. method is perhaps the most versatile single method there is every reason to supplement it wherever possible by other equilibrium methods such as solubility and distribution studies and by spectral "finger-print" methods. With the last one may check the results obtained with e.m.f. data and also extend the measurements in the ranges where the accuracy of the e.m.f. data is unsuflicient. K z z = *pzz*Kl-z 2BOH+B,(OH)z. 168 QUARTERLY REVIEWS The structure of the polynuclear complexes formed on hydrolysis is of considerable interest Much remains to be done by diffraction work both on crystals and solutions.The equilibrium work gives only the dG of the reaction whereas it would be desirable to know the AH and AS components. It seems that the accuracy that may be obtained from the variation of equilibrium con- stants with temperature is rather low so an attack with calorimetric methods seems more promising. Acknowledgements. The research work described in the present Review has been supported in various ways by the Swedish Natural Science Research Council the Swedish Council of Technical Research and the Swedish Atomic Energy Commission. Recently very valuable support has been given by the Air Research and Development Command United States Air Force through its European Office.

ISSN:0009-2681    

DOI:10.1039/QR9591300146

出版商:RSC    

年代:1959

数据来源: RSC