Volatile Silicon Compounds

OTHER TITLES IN THE SERIES ON INORGANIC CHEMISTRY

Vol. 1. CAGLIOTI (Ed.)—Chemistry of the Co-ordination Compounds Vol. 2. VICKERY—TAé? Chemistry of Yttrium and Scandium Vol. 3. GRADDON—An Introduction to Co-ordination Chemistry

VOLATILE SILICON COMPOUNDS

by

E. A. V. EBSWORTH University Chemical Laboratory

Cambridge

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1963

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PREFACE PREVIOUS accounts of volatile silicon compounds have in the main been written from the organic or the inorganic point of view. Such a division is particularly unfortunate where silicon compounds are concerned, since it leaves compounds like chlorosilane, SiH3Cl, to the inorganic chemist, while trimethylchlorosilane, Me3SiCl, is considered an organic compound· I have tried to present a study of volatile silicon compounds irrespective of whether they contain carbon; anyone who wants a fuller treatment of the organic derivatives is referred to Professor Eaborn's excellent book Organosilicon Compounds.

This book is primarily intended for research students, though I hope it may be helpful to undergraduates in their final year. In a subject as extensive as this, there are bound to be omissions, many of them inadvertant; the balance of what I have included is a reflection of my own interests. I have combined discussion of molecular structure and chemical properties, concentrating on material which has not been reviewed before and summarizing points which are discussed at length elsewhere. The treatment of the compounds of germanium and tin is not intended to be in any way complete; it has been included so that the behaviour of silicon compounds may be considered in the context of the Periodic Table and of the compounds of neighbouring elements. Lead compounds have been perhaps rather arbitrarily excluded, save for some reference to organo-lead hydrides; this is because relatively few lead compounds are strictly analogous to the simple substituted silanes.

Some comment should be made about the way in which I have used free energy calculations based only on changes in bond energy in a number of reactions. I know well that such calculations give no absolute measure of the free energy change unless the entropy term is included; none the less, they may be useful in helping to decide why certain reactions do not occur. If, for instance, two compounds might react together in two different ways, each reaction giving rise to the same number of molecules of gaseous product, then the reaction in which the more bond energy is released should be the reaction with the more favourable free energy change. If in practice the other reaction is found to be preferred, the former reaction is likely to have an unfavourable activation energy.

1

2 VOLATILE SILICON COMPOUNDS

The use of force constants in studies of molecular structure presents a considerable problem. In principle, the force constant of a bond is as important a property as, say, the bond length; unfortunately, however, the calculation of force constants from the observed vibrational spectra of any but the simplest of molecules is a difficult and complex process, the result is liable to vary with the type of force-field assumed and the precise physical meaning of the parameter obtained is not absolutely clear. On the other hand, there is a clear correlation in simple molecules between bond order and force constant, at least where single, double and triple bonds are compared; moreover, any measurable property which is likely to be of use in discussions of molecular structure is not lightly to be ignored. I have therefore referred to work on force constants even for relatively complicated molecules, though I do not believe that much significance should be attached to small differences. I have only described Siebert's method of predicting the force constants of single bonds, though I know that there are other formulae of this sort; this is not meant to be a monograph on force constants, and I do not think that an elaborate discussion of the different formulae would be justified here.

I have tried to avoid using that unfortunate word "stable" without the essential qualification "to (some reagent or set of conditions)". Where the word is used without qualification, it is meant to describe the stability of the compound in question with respect to decomposition, polymerization or dissociation given—in other words, with respect to reaction with itself.

I have described π-bonds in a way which differs slightly from that most frequently used in published work. I have called an ethylenic π-bond a (p-p)Tc-bond; the π-bond between nitrogen and boron in N-dimethyl-aminoboron dichloride is called a (p-+p) π-bond, the π-bond between nickel and carbon in nickel carbonyl is a (d-*p) π-bond, while that in tri-silylamine is a (p-+d) π-bond. This notation conveys a little more about what the electron-distribution in the bond is believed to be than does the more commonly used method of description.

Finally, there is the vexed question of nomenclature. Several systems have been used for naming silicon compounds, and I do not propose to describe them all here; an account of some of them is to be found in Professor Eaborn's book. My main concern has been that the names I have used should leave the reader in no doubt as to the structural formulae of the compounds he is reading about. For the rest, I have preferred names for halogen- and organosubstituted mono- and disilanes based on the

PREFACE 3

word "silane"; I have not used this system where pseudohalogen substi-tuents are concerned, because I think the name "tetraisothiocyanatosilane" is much clumsier than "silicon tetraisothiocyanate". I have avoided the "silazane" system for amine-derivatives, preferring to call these "silyl-amines"; similarly, I have called the sulphur and selenium derivatives "sulphides" and "selenides" rather than "thianes" and "selenanes", because I believe that the amine-sulphide-selenide system is easier to understand for a reader who is not already familiar with organosilicon chemistry. I have capitulated over the oxygen derivatives, calling these^sil-oxanes rather than ethers.

The final chapter is in part a summary of material to be found in the rest of the book; in order to avoid repetition, only a few source-references have been included at the end of that chapter, but the source for any unreferenced statement should be easily found in the body of the book.

Thanks are due to many of my friends and colleagues, at Cambridge and elsewhere, for help in the writing of this monograph. Dr A. D. Buckingham, Dr L. E. Orgel, Dr A. G. Sharpe, Dr W Sheppard, Dr T. M. Sugden and Dr J. J. Turner have all given me the benefit of their expert opinions; I am most grateful to Mr J. S. Griffith, Dr K. MacKay, Dr D. C. McKean and Mr S. Frankiss for permission to quote observations or ideas as yet unpublished. Mr Frankiss, Dr A. Hass and Mr M. J. Mays read parts of the manuscript and the proofs, and made a number of helpful comments. Most of all, I am indebted to Dr A. G. Maddock, who has always been ready with advice about any difficulty, and has assisted and encouraged me at every stage.

E. A. V. EBSWORTH

CHAPTER 1

INTRODUCTION: ATOMIC PROPERTIES

BEFORE turning to a discussion of particular compounds, it is as well to consider the atomic properties of silicon and the other Group IV elements in relation to one another. Although these atomic properties, which are often determined from a study of the free atoms or of the elements themselves in their standard states, may be profoundly modified by compound formation, they often provide a surprisingly reliable basis for a discussion of the properties of molecules ; moreover, it is sometimes possible to find an explanation for the differences between apparently analogous compounds of two or more elements in terms of differences in atomic properties. This may be a dangerously speculative process, however, and must be regarded with caution.

All the elements in question—carbon, silicon, germanium, tin and lead—have the outer electronic configurations in their ground-states of (ns2np2). The group valency of four is reached by the formal promotion of an s-electron to an empty /^-orbital: (ns2np2->ns1npz). All these elements form compounds derived from the valency-state of four; a valency of two becomes relatively more stable as the atomic weight of the element increases. Carbon and silicon form no compounds derived from the divalent state that are stable at room temperature, if carbon monoxide is excluded; germanium (II) is rather ill-defined, and several of the compounds of formula GeX2 may contain germanium-germanium bonds*, but tin (II) is well-characterized. This increase in the relative stability of the lower oxidation state with atomic number has been attributed both to increasing stabilization of the ns- relative to the wp-electrons of the valence-shell with increasing atomic number, and also to the fact that the heavier elements form weaker covalent bonds(la). All of the compounds with which this monograph is directly concerned are at least formally derived from the valency state of four, but the stabilization of the lower state has an important effect upon the hydride chemistry of germanium (IV) and tin (IV). The MH bonds are strongly reducing, and compounds such as trichlorogermane

* Though Gel, has the Cdl2 structure/*)

4

INTRODUCTION — ATOMIC PROPERTIES 5

are liable to decompose to give germanium (II) chloride and hydrogen chloride(2):

GeHCl3 = GeCl2+HCl

The bonds from a saturated carbon atom are usually regarded as formed by ^-hybr id orbitals, with minor changes in hybridization in unsym-metrically-substituted compounds like chloromethane (though it has been suggested that there is some d- and even /-character in the hybrids)(3). Carbon, having no more orbitals in its valence-shell, has a co valency-maximum of four, but many compounds containing multiply-bonded carbon are of course well known. The position with silicon and the heavier elements is rather different. These all have «d-(and germanium and tin have nf-) orbitals in their valence-shells; these orbitals could in principle be used in forming bonds, but in the neutral atoms they are very much more diffuse and of higher energy than the ns- and np -orbitals(4). The σ-bonds from silicon (IV), germanium (IV) and tin (IV) are therefore normally regarded, like those from carbon, as formed from 5*p3-hybrids. It has, however, been shown that the production at the central atom of a positive charge, which may be quite small, has a strong contracting effect upon the d-orbitals of the valence-shell; such a positive charge can be induced by the presence of electronegative substituents like fluorine bound to the central atom in question*4»5·6). If this happens, the ^/-orbitals can be sufficiently contracted to become of energy and spatial extent appropriate for mixing with the bonding s- and ^-orbitals; hence an increase in the maximum number of or-bonds formed is possible, and this explains the well-known acceptor properties of tetrafluorosilane, which readily forms addition-compounds such as SiF4.2NH3 (in which the σ-bonds are almost certainly built from .sp^-hybrids). If the d-orbitals of silicon in tetrafluorosilane can be used to form additional σ-bonds, however, then they could also be used to form the σ-bonds of the parent compound; in other words, the silicon-fluorine σ-bonds of tetrafluorosilane itself are likely to have considerable ^-character. The extent of J-mixing will of course vary with the polarization of the ^/-orbitals involved; since the appropriate wave-functions are unknown, it is difficult even to guess at the extent of such mixing, but the comparison of σ-bonds from carbon and silicon is made even more difficult by the introduction of this unpredictable parameter.

Besides this effect on σ-bonds, the J-orbitals of the heavier elements are important in another way. Silicon, germanium and tin do not appear


to form (p-p)n-bonds9 analogous to the π-bonds of ethylene*. Symmetrical tetraphenyldichlorodisilane, for example, gives what is probably a dimer when treated with sodium(7).

2Ph2SiCl.ClSiPh2+4Na -> Ph2Si—SiPh2

I I Ph2Si—SiPh2

The reason for this may be bound up with inner shell repulsions, with the relatively poor overlap between 3p and 2p or 3p and 3/rcr-orbitals (though this could be improved by d-hybridization), and with the relatively greater energies of σ-bonds from silicon (and the heavier elements) to electronegative species when compared with bonds from the same species to carbon(9) ; thus, instead of forming a silicon-oxygen "double" bond, as in the formal silicon analogue of acetone, silicon prefers to form the two σ-bonds that lead to the formation of the polymeric silicones:

RN

R / S i = O

R/I

•R

Si/ Si \ R

On the other hand, the heavier elements have empty J-orbitals in their valence shells. Two of these are of π-symmetry relative to the tetrahedral σ-bonds of the saturated element, and so can combine with the π-orbitals of any attached atom or group; if the latter π-orbitals contain electron-pairs, (as in the halogen atoms, or the dimethylamino group), their energy will be lowered by this interaction, and a (p->d) π-bond will result(4):

other y-bond

Although this is rather different from (p-p)n-bonding9 it will have an important effect upon the structures and properties of silicon compounds. The delocalization of the π-electrons of the attached group will lead to a displacement of negative charge towards the silicon atom; since an attached group of this sort is always more electronegative than silicon, the polarity of the system will be reduced. The overall bond between the

* From the cracking of tetramethylsilane, a compound has been obtained which was believed to contain a carbon-silicon double bond; it has been given the structure. Me2Si=CHSiMes

(ea), but further studytêh> has shown that the molecule has the structure

MeaSK >SiMej. X C H /


two species will of course be strengthened by such π-interactions; more important, since overlap is likely to be greatest between a J-orbital and a pure -orbital, differences might be expected in the hybridization of an element of groups V or VI bound on the one hand to carbon and on the other hand to silicon, depending upon whether the lone pair or lone pairs were accommodated in hybrid orbitals (as in water or dimethyl ether) or in pure p-orbitals (where more efficient π-bonding would be possible). The extent of (/?->rf)7>bonding is likely to depend on the diffuseness of the J-orbitals concerned and on their principal ^quantum number, since the amount of overlap will depend on both these properties; some d-orbital contraction is probably necessary, so that π-bonding is likely to be most important when silicon is bound to some very electronegative group with π-orbitals containing electron-pairs(4). A detailed study of how this interaction would be expected to vary with such factors as the principal quantum number of the rf-orbitals has yet to be made.

The difference between the π-bonding properties of silicon and carbon can perhaps be made clearer by comparing the radicals triphenylmethyl, PhgC, and triphenylsilyl, Ph3Si.. The former is strongly stabilized by delo-calization of the unpaired electron over the rings; the latter has not been characterized but is almost certainly much less stable (see chapter 4), probably because the 3/?-orbital of silicon does not interact sufficiently with the π-orbitals of the ring to stabilize the system. On the other hand, the radical-ion PI14SÌ" might well be appreciably more stable to oxidation than its carbon analogue, because of (p-*d) delocalization of the unpaired electron.

Besides these properties, there are some others which will be of importance in the discussions which follow. The atomic radii of the elements are given in Table 1.1; the most surprising thing about the values is the relatively small increase in radius from silicon to germanium.

TABLE 1.1.—ATOMIC RADII(10)

Element

Carbon Silicon Germanium Tin

Radius (A)

0-77 117 1-22 140

This can be put down to the interpolation between the two elements of the first transition series, which gives rise to the "scandinide contraction".


The relatively small increase in atomic radius for an increase in atomic number of 18 leads to a greater electron-density in the germanium atom than migh t have been expected, and this may in its turn affect the relative electronegativities of germanium and silicon.

Electronegativity is one of those irritating concepts like covalent character which sound as if they mean very much more than they do. It may be defined in a number of ways(11-18), most of which have at least some connection with Pauling's definition111*0 : "The power of an atom in a molecule to attract electrons to itself." Electronegativity so defined is impossible to measure; it has therefore been supposed that certain physical or chemical properties of bonds can be related to the difference between the electronegativities of the bonded atoms, and it is as a result of attempts to use different properties in this way that most of the many different scales of electronegativity have arisen. As far as the elements of Group IV are concerned, all scales agree in making carbon the most electronegative, but there is a difference of opinion about the relative electronegativities of the other elements. In Pauling's original table(lla), silicon, germanium and tin were given roughly equal electronegativities, but silicon was by a little the most electronegative; in his modified table, the electronegativities of silicon, germanium and tin are all equal(11). His values are based on thermochemical measurements, but several other scales, using quite different properties to measure electronegativity, have given values which agree with his(12· 13>14). On the other hand, there are methods of calculating electronegativity which take into account the changes in relative electron-density in different parts of the Periodic Table caused by the various transition series, and according to these scales supported by NMR measurements) germanium is more electronegative than silicon or tjn(ie,i7,i8) p ^ 0f t h e trouble is that electronegativity is essentially a qualitative concept; it can be made quantitative by definition with respect to some specific property, but such definition at once makes it less general in application, for almost any measurable property is affected by other things besides electronegativity, and these other factors are in practice impossible to allow for. Where large differences in electronegativity are concerned, there is usually general agreement; since the electronegativity of an atom is unlikely to be constant, but will almost certainly depend upon the other groups bound to it, and upon such variables as orbital hybridization(18) (which for the heavier elements are almost impossible to predict or measure), it is only the large differences that have any general significance; for the rest, silicon, germanium, and tin should


probably be taken as of more or less equal electronegativity. The case of lead is rather different; some scales make it only little less electronegative than carbon(18), while others make it the least electronegative element of the group(12'13). It is not clear why this difference arises, nor which value represents the true state of affairs the more clearly.

These considerations can be used in a quasi-theoretical discussion of the structures and properties of the compounds of silicon and their analogues, but in any such discussion it is necessary to remember the limitations of the experimental methods at present in use in structural chemistry. From the point of view of chemical behaviour, the electron configuration of a molecule is its most important property; on the other hand, there are few experimental methods for determining this. X-ray diffraction, it is true, determines electron densities directly, but the bulk of the electron density of an atom is concentrated round the nucleus, and the changes in electron density that represent bonds are so small as to be almost impossible to observe. Electron spin resonance and electronic spectroscopy can be used to study electronic behaviour; unfortunately, very few compounds of silicon are paramagnetic, and there are not many with absorption bands in the readily-accessible region of the ultraviolet. It is therefore necessary to make inferences about the electron-distribution within such molecules by studying nuclear properties.

Interbond angles and bond lengths can now be determined quite accurately in many simple molecules(20); it is still difficult to obtain precise values for M H distances and HMH angles except in the hydrides MHn*, but for bonds between heavier elements some very precise information is available. A notable example of the sort of precision possible in favourable cases is the value for the silicon-chlorine distance in chlorosilane, which is given as 2-0479 ± 0O007Â(22)f · The precision of these measurements drops off sharply as the molecule concerned becomes more complex; electron diffraction, X-ray diffraction and microwave spectroscopy are all useful under different circumstances and can all give very accurate values, but X-ray diffraction has been relatively little used to study simple silicon compounds, since most of these are liquid at room temperature.

Vibrational spectra have been used to determine MH bond lengths and HMH angles, and afford a most valuable way of measuring these parameters; molecular symmetries have also been deduced from vibrational

* See, however, Ref. 21. t Some spectroscopists, however, doubt whether these measurements have physical

significance to less than 001 Â.


spectra, but here the results (while useful in giving corroborative evidence) can be misleading. In the case of disiloxane, for example, the vibrational spectrum led to the attribution of the wrong structure<23). In the infrared spectrum of this molecule the symmetrical skeletal stretching mode is missing, and so the heavy atom skeleton was taken to be linear; subsequent work has shown that the skeleton is in fact bent(24»25>2e).

From the interbond angles and bond lengths observed, it is possible to make some deductions about the electron configurations of the molecules studied, but these deductions rest on a number of assumptions. The interbond angles from heavier elements can often be explained in more than one way(27«28), while the interpretation of the lengths of bonds from the elements of Group IV other than carbon is at present in a thoroughly unsatisfactory state (see pp. 50,80). Some other nuclear properties are more directly affected by changes in electronic environment; among these are the properties measured by nuclear magnetic resonance and nuclear quadrupole resonance spectroscopy. Only nuclei with quadrupole moments give quadrupole resonance spectra; the only compounds of the elements of Group IV that have been systematically studied in this way are the halides, and the interpretation of the results for the compounds of silicon, germanium and tin involves too many uncertainties to be of much use in elucidating electronic structures*29'30»31*. The interpretation [of nuclear quadrupole coupling^constants is discussed further on pp. 52 to 53. Nuclear magnetic resonance is more generally useful, but here again the results are extremely difficult to interpret. None the less, it is possible that major advances in the understanding of the electronic structures of relatively simple molecules may come from interpretation of magnetic resonance measurements. Spin-spin couplings, for example, have been interpreted in terms of orbital hybridization(32).

It is therefore clear that a quantitative discussion of the differences in properties between the compounds of silicon and of the other elements of the group is too much to hope for; from the account that follows, which is of necessity more qualitative than might be desired, it is hoped that some general points may emerge that will prove to be of fundamental significance.

REFERENCES (1> H. M. POWELL and F. M. BREWER, / . Chem. Soc, 197 (1938).

(1*> R. S. DRAGO, J. Phys. Chem., 62, 353 (1958). <■> C. W. MOULTON and J. G. MILLER, / . Amer. Chem. Soc, 78, 2702 (1958).


(8> L. PAULING, The Nature of the Chemical Bond, Cornell Univ. Press, 3rd. ed., p. 126 (1959).

(4> D. P. CRAIG, A. MACCOLL, R. S. NYHOLM, L. E. ORGEL and L. E. SUTTON, / . Chem. Soc, 332 (1954).

<5> D. P. CRAIG and D. W. MAGNUSSON, / . Chem. Soc, 4895 (1956). (e> D. P. CRAIG, Chemical Society Symposia, Bristol, (Special publication, number 12),

343 (1958). (ea) G. FRITZ and J. GROBE, Z. anorg. all. Chem., 311, 325 (1961). <eb' G. FRITZ et al. to be published. (7> H. GILMAN, D. J. PETERSON, A. W. JARVIE and H. S. WINKLER, / . Amer. Chem.

Soc, 82, 2076 (1960). (8) K. S. PTTZER, / . Amer. Chem. Soc, 70, 2140 (1948). <e> R. S. MULLIKEN, / . Amer. Chem. Soc, 72, 4493 (1950.

<10> L. PAULING, op. cit., (réf. 3), p. 225. <") L. PAULING, op. cit., (réf. 3), Chapter 3. <lla> Idem, ibid, Chapter 2. (2nd. ed., 1939). <12> R. S. DRAGO, / . Inorg. Nucl. Chem., 15, 237 (1960). <13> W. GORDY, / . Chem. Phys., 14, 305 (1946). <14> W. GORDY, Phys. Rev., 69, 604 (1946). <15> J.K. WILMSHURST, / . Chem. Phys., 27, 1129 (1957). <16> R.T. SANDERSON, / . Chem. Phys., 23, 2467 (1955). (17> R. T. SANDERSON, / . Amer. Chem. Soc, 74, 4792 (1952). <18> A. L. ALLRED and E. G. ROCHOW, / . Inorg. Nucl. Chem., 5, 264; 269 (1958). (19> P. T. NARASIMHAN and M. T. ROGERS, J. Amer. Chem. Soc, 82, 5983 (1960). (,°) See, for example, "Interatomic distances and configuration in molecules" (Chem

ical Society Special Publication, number 11, 1958). <21> C. C. COSTAIN, / . Chem. Phys., 29, 864 (1958). <22) B. BAK, J. BRUHN and J. RASTRUP-ANDERSON, Acta Chem. Scand., 8, 367 (1954). (23> R. C. LORD, D. W. ROBINSON and W. SCHUMB, / . Amer. Chem. Soc, 78,1327 (1956). <24> R. F. CURL and K. S. PITZER, / . Amer. Chem. Soc, 80, 2371 (1958). (26) J. R. ARONSON, R. C. LORD and D . W. ROBINSON, / . Chem. Phys.t 33, , 1004 (1960). (26> D. C. MCKEAN, R. TAYLOR and L. A. WOODWARD, Proc Chem. Soc, 321 (1959). <27> R. S. MULLIKEN, / . Amer. Chem. Soc, 77, 884 (1955). (28) w # GORDY, Technique of Organic Chemistry, Vol. 9, Interscience, N.Y., p. 120 (1956). <2e> W. GORDY, Disc. Faraday Soc, 19, 14 (1955). <80> Β. P. DABLEY, ibid, 255. <31> J. K. WILMSHURST, / . Chem. Phys., 30, 561 (1959). (32> N. MÜLLER and D. E. PRITCHARD, / . Chem. Phys., 31, 1471 (1959).

CHAPTER 2

THE SiH BOND

1. FORMATION

THE methods of forming SiH bonds that are of general preparative importance may be divided into two main groups. Reactions of the first group are hydrolytic in nature, and involve the treatment of a compound containing negatively-polarized silicon with a protonic acid; the second group is made up of reactions between a metal hydride and some compound in which silicon is bound to a more electronegative element.

The hydrolytic method was used in the early studies of silanes. Magnesium suicide was found to react with a mineral acid in aqueous solution to form a mixture of silicon hydrides(1); the essential reaction may be described by the equation:

\ \ - S i - + H + = - S i H / /

but the process is almost certainly more complicated than this, and is discussed in more detail when the preparation of the higher silanes is considered in Chapter 3. Various solvent-systems have been used with success, including liquid ammonia(2) and anhydrous hydrazine(3). The silicon compounds that will react in this way are, however, few in number; the most important of them are the suicides of the metals of Groups I and II, and the alkali metal silyls. Most of these compounds are not available commercially and have to be specially prepared; moreover, the hydrolysis often gives rise to a complex mixture from which the required compound must be obtained by fractional distillation or some other method of separation. Finally, the yield of any one product is often low. Methods of the second type have therefore gained greatly in popularity, particularly since the metal hydrides have become commercially available at reasonable prices. Here the essential reaction is between a positively-polarized silicon atom and what amounts to a hydride ion:

\ * + «5- - \ - S i - X + H" = - S i H + X-/ /

12

THE SlH BOND 13

Alkali metal hydrides dissolved either in organic solvents*4·5* or in salt-melts(6) are effective, but the most commonly-used source of hydrogen is lithium aluminium hydride. This, dissolved in some ether (usually chosen with reference to the physical properties of the required product) has been found to react with a wide range of silicon compounds to form silicon-hydrogen bonds(5). Although the silicon-carbon(7) and silicon-silicon(8)

bonds are usually unaffected (but see Ref. 7), these metal hydrides have been found to react with a wide range of compounds containing silicon bound to fluorine(9), chlorine*10'11*, bromine(4), iodine(4»12), cyanide(4), oxygen(13>14), sulphur(4), and nitrogen(15); even quartz evolves some silane when treated with lithium aluminium hydride{16). The required product is usually obtained in good yield and reactions are relatively free from the formation of byproducts.

Deuterated silanes have been prepared using deuterated metal hydrides(17), and also by the hydrolysis of magnesium suicide with heavy water and DC1(18).

Besides these general methods, compounds containing hydrogen bound to silicon are formed in a number of more specific reactions. Trichlorosilane, for example, is made by passing a stream of hydrogen chloride gas over silicon heated to about 400°; some tetrachlorosilane is formed at the same time(19). Methyldichlorosilane is manufactured by a similar process; methyl chloride and hydrogen are passed over a heated mixture of copper and silicon(20). The pyrolysis of tetraalkylsilanes gives rise to some compounds which contain hydrogen bound to silicon(21); SiH bonds are also formed in the reactions between phenylchlorosilanes and certain Grignard reagents (22'23). Finally, the disproportionation of silicon hydride derivatives leads to redistribution of the silicon-hydrogen bonds already present; these disproportionation reactions are considered later (p. 31).

Attempts at partial hydrogénation of silicon compounds have usually failed. Little or no partly-reduced material was obtained from the reaction between lithium aluminium hydride and an excess of tetrachiorosilane(24)*, phenyltrichlorosilane(25), dimethyl-trichlorosilylamine(15), or hexachloro-disiloxane<26»27), while the last-named compound when passed with hydrogen over heated aluminium at 500° gave only traces of partly hydro-genated product(27). It has been reported that tetrachlorosilane and formaldehyde react together at 400-500° in the presence of γ-alumina to form chloro- and dichlorosilane(28), but attempts to repeat this experi-

* Dichlorosilane is said to be obtained when tetrachlorosilane is treated with lithium borohydride in tetrahydrofuran solution(27a).


ment have given only chloro- and dichloromethane, with no product containing hydrogen bound to silicon(29>30). Since chlorosilane(31) and chloromethane(32) have very similar vapour pressures over a range of temperature, it is possible that the material originally identified as chlorosi-lane was really chloromethane, and that dichloromethane was similarly taken for dichlorosilane(29). The production of large amounts of chlorosilane in the presence of hydrogen chloride, carbon monoxide and a catalyst at 400° would be surprising. In a subsequent report(28a), a copper catalyst is mentioned instead of the alumina catalyst originally described, but the effectiveness of this has not been reinvestigated.

When trichlorosilane vapour, mixed with hydrogen, was passed over granulated aluminium at 300-400°, mono- and dichlorosilanes were produced in considerable quantities, together with a trace of monosilane and some aluminium chloride(33). This is apparently a successful partial hydrogénation, but since aluminium chloride catalyses the disproportionation of halosilanes, some of the partly-hydrogenated material may have come from disproportionation of the trichlorosilane rather than from reaction between this compound and hydrogen.

The hydrides of germanium and tin have been prepared using reactions very like those used for making silanes. The acid hydrolysis of magnesium germanide(34) and stannide(35), and reduction of germanium and tin compounds with complex metal hydrides(5>36>37) are preparative methods that have been frequently employed. Although lithium aluminium hydride in ether solution is a good reagent for preparing monostannane(37), it does not give yields of monogermane greater than 40%(3e). Lithium tri (tert.-butoxy)-aluminium hydride is more efficient(38), and it has been found that germanium (IV) oxide<39) and tin (II) chloride<40) can be reduced to the hydrides by sodium borohydride in acid aqueous solution. The reaction with germanium (IV) oxide has been studied in some detail<3β·39) ; when heavy water was used as solvent for sodium borohydride containing no deuterium, the monogermane produced contained about 45 atom-% of deuterium among the hydrogen bound to germanium(39). It has been shown that sodium borohydride does not exchange hydrogen with water at pH 12 (41), but the behaviour in acid solution has not been studied.

Various other reactions have been used to produce hydrides of germanium and tin. Aqueous solutions of tin compounds can be reduced to hydrides electrolytically(42) or with a mineral acid and some electropositive element like zinc or magnesium(43); the latter reaction can also be used to prepare germanium hydrides(44). Trichlorogermane is prepared

THE S lH BOND 15

by passing hydrogen chloride over heated germanium (II) oxide(45); the action of hydrogen chloride upon a mixture of copper and germanium heated to about 400° also gives trichlorogermane(45a), a reaction analogous to that used to prepare trichlorosilane.

It is only very recently that hydrides of lead have been prepared and characterized. Trimethylplumbane and dimethylplumbane are made by treating trimethylchloroplumbane or dimethyldichloroplumbane with lithium aluminium hydride in ether at —90 110O(4e). Although evidence had been presented before that strongly suggests the formation of lead hydrides(47,48), this is the first time that such compounds have been isolated in a pure state.

2. PHYSICAL PROPERTIES

Strictly speaking, the properties of a particular bond should never be considered in isolation; they must always depend to some extent upon the nature of the rest of the molecule of which that bond forms a part. None the less, it is often possible to associate some measurable quantity with a particular bond—internuclear distance is an obvious example; the way in which such a property changes as the rest of the molecule is altered can then give useful clues about the nature and extent of interactions within the molecule as a whole. Bond properties that have been treated in this way include bond energy, bond length, the resistance of a bond to stretching, and the bond dipole moment; the field gradient at one or more nuclei (obtained from nuclear quadrupole resonance spectroscopy) and the influence of the bonding electrons and neighbouring magnetic nuclei upon the interaction of either or both nuclei concerned with a magnetic field (investigated by nuclear magnetic resonance spectroscopy) have also been considered in discussions of bond properties. Where two or more bonds are concerned, the bond angles and higher-order magnetic interactions may be studied. Of these various properties and methods, nuclear quadrupole resonance is of no use where the silicon-hydrogen bond is concerned, for neither hydrogen nor the common isotopes of silicon have nuclear quadrupole moments; hydrogen, however, has a nuclear spin of J, and so nuclear magnetic resonance is in principle a useful tool. Moreover, 29Si, which also has a nuclear spin of \, is present in natural silicon in about 5% abundance; although very little use has been made of silicon resonances, 29Si-H couplings can easily be obtained from the proton resonance spectra of silicon hydride derivatives(12). Bond dipole


moments are very difficult to determine unambiguously and have been little studied in silicon compounds*, and not many values for SiH bond energies are available. All the other properties mentioned have been investigated for SiH bonds at least to some extent as a function of the rest of the molecule, but little systematic information is available for compounds of germanium and tin.

TABLE 2.1

Property

Bond energy, kcala

Bond length, Â: (a) Observed (N ID,

MD3H) (b) Cale, uncorr.e

(e) Cale, corr/ Force constant, md/Â:

(a) Observed? (b) Predicted?

Proton resonance chemical shift, ppm relative to gaseous methane

CH4

99a

1 0936 ± 0005c 114 MO

54 5-4

0

SiH4

76a

1·480±·001<* 1-54 1-51

3 0 3-7

• -3·00Λ

GeH4

69ö

1·523±·001* 1-59 1-55

2-8 3-6

-2-85'

SnH4

60ö

1-701 i-001d

1-77 1-73

1-7* 2-9

aRef. 50. bThese values were obtained using an unusual high-temperature calorimetrie method(51). cRef. 52. dRef. 53. eRef.54. 'Pauling's electronegativities*55* were used in correcting according to Ref. 56. ?Ref. 57. ÄRef. 58. 'Ref. 59.

The simple hydrides MH4 are convenient reference-compounds, and their fundamental properties are therefore considered first. Such values as are available for the bond energies, bond lengths, force constants and proton resonance frequencies are given in Table 2.1, together with those for methane and also with certain calculated values. Most of the available evidence(5]'60) suggests that the energy of the MH bond drops as M becomes heavier; this trend is observed in other groups. The bond lengths increase in the same sense, but the values for silicon and germanium are very close together; this is probably a result of the "scandinide contrac-

* The SiH bond moment has been estimated at about 1·4-1·50(49> and 1-0D(103) by different methods; a very recent determination based on infrared intensity measurements gives a value of 1-6D in monosilane(49a).

THE SiH BOND 17

tion", and is also observed in the atomic radii of the two elements. The agreement between the measured lengths of the M-H bonds and those calculated from the atomic radii is not very good; rather better agreement is obtained if a correction is made of the form suggested by Schomaker and Stevenson(56), using Pauling's values for the electronegativities of the elements concerned, but the validity of such a correction has been effectively questioned(6l). It should be noted that any increase in the electronegativity of germanium(62) would make the corrected value for the GeH bond distance agree less closely with that observed.

The force constants are compared with values predicted using Siebert's formula(57). Since this formula was derived from the alkyls(63), it is in a sense the MH and the MC bonds that are being compared. While methane's force constant is very close to the predicted value, those for the other hydrides are all low. Siebert(57) has pointed out that low force constants are usually found in those hydrides in which hydrogen would be expected to form the negative end of the bond dipole; the force constants of silane, germane and stannane fit in with this empirical generalization. It must be remembered that it is the expected sign of the bond moment that is concerned; the sign itself is almost impossible to determine experimentally.

The relative values of the proton resonance chemical shifts are rather surprising. In saturated aliphatic compounds the chemical shift of a proton bound to carbon usually depends upon the electron-withdrawing power of the other substituents at the carbon atom, unless special effects such as multiple bonding or hydrogen bonding complicate the issue(64); the greater the electron density at the hydrogen atom, the higher the field at which resonance is observed. Thus in the series of halomethanes from iodomethane to fluoromethane the proton resonance moves to successively lower fields(65). By all standards silicon is less electron-withdrawing than carbon, at least where or-bonds are concerned; this means that in the SiH bond the electrons would be expected to concentrate round the hydrogen nucleus more closely than in the CH bond, and consequently that the protons would be more shielded in silane than in methane. Much the same reasoning applies when methane and germane are compared; it is therefore something of a surprise to find that the proton resonance in both silane and germane appear on the low field side of the resonance in methane.

At present there is no very satisfactory explanation for this. Attempts have been made to describe the chemical shift of a nucleus in a complex molecule by breaking down the circulation of electrons within the molecule


as a whole into circulations round the individual nuclei(ee). Using a treatment of this kind, four factors may contribute to the chemical shift of a particular nucleus: the diamagnetic shielding at the nucleus in question; electron-circulations round the next-door nuclei; anisotropies in near-by bonds; and a paramagnetic term, arising from quantum-mechanical mixing of the ground state with excited states involving unpaired electrons. The effects of environment and of permanent electric fields need not be considered here, for the measurement was made in the gas phase and neither molecule has a permanent dipole moment; the question of ring currents does not arise. Comparing the chemical shifts of the protons in methane and silane from this point of view, the first factor should lead to the SiH resonance appearing to high field of the CH resonance, since silicon is less electronegative than carbon. The second term averages to zero in each case, because of the tetrahedral symmetry of the molecules, while calculations show that the third term is much too small to account for the observed low-field shift(12). There remains the paramagnetic term. It could be argued that the J-orbitals of silicon will overlap the 2/>-π orbital of hydrogen, and so lower the energy-gap between the σ-bonding orbital and the π-orbital; this would increase the contribution of the excited state to the ground state, since this contribution depends inversely on the excitation-energy(67), and would lead to a larger paramagnetic term in silane than in methane (where any such [/?-► d] π-interactions are likely to be much less important). Since the paramagnetic term expresses a shift to low fields, this explanation could account for a low-field shift from methane to silane; unfortunately, however, calculations show that a change of this sort, taking reasonable values for the parameters concerned, can only account for about one-tenth of the observed shift(12). It therefore appears that a simplified approach of this kind is not adequate here, possibly because it implicitly neglects the effect of overlap upon the chemical shift(68).

So much, then, for the properties of the simple hydrides. The next thing to consider is how these properties vary when some of the hydrogen atoms are replaced by other groups. After the changes in the individual properties have been discussed, some attempt can be made to see how far these changes can be related to one another.

Bond length—The lengths of the silicon-hydrogen bonds in a number of simple compounds are given in Table 2.2, together with the CH bond leugths in analogous derivatives of methane. The most striking thing about the values for both silicon and carbon is that they are insensiivet

THE SlH BOND 19

TABLE 2.2.—SiH BOND LENGTHS IN VARIOUS COMPOUNDS (Â)

Compound

SiH,D, SiD3H SiH8F* SiH,F,t S1HF3 SiH.Cl

SiH,Br

SiH,I CH3SiH,Ft CH,SiHFtt CH,SiH, (CH^SiH, (CH^SiH

r(Si - H)

1·480±·001 1-474 1-471 ±007 1·455±·01 1-476 1·481±·001 1-481 1·483±'0Ο1 1·483±·001 1-473 1·474±·005 1·485±·005 1-483+005 1-489+001

Method

V V M M V M V M M M M M M M

Ref.

53 69 70 71 69 72 69 72 72 73 73a 74 75 76a

Compound

CH,D CH3F

CH2F2 CHF3 CH3CI

CH3Br

CH3I

CH3CH3

r(C - H)

1·0936±·0005 1097 1096 1098 1096

1095

1096

1-102

Method

V M M M M

M

M

V

Ref.

52 77 78 79 77 72 77

77

80

M =» microwave, V = vibrational spectroscopy. * The SiD distance in SiD8F is 1·473±·003 Â, as against 1·479±·001 for SiD.Cl

and 1-478 ±-001 for SiD8Br(74>. t A subsequent recalculation from the same measurements gives SiH distances

of 1-467 for SiH2F2, 1-477 for CH3SiH2F and 1471 for CH8SiHF2(80a).

to changes in substitution within the range of compounds considered here; almost no significant alterations can be detected. There is a suggestion that bonds may become shorter as fluorine atoms successively replace hydrogen, but the changes are extremely small and may not be real. The SiH bond lengths are effectively constant in the series of methyl silanes. The CH bond length does not change appreciably in the halo-methanes; there is a possibility that the SiH bond lengths in the mono-halosilanes may increase with the atomic number of the halogen atom, but once more it must be emphasized that the experimental values are being strained to and indeed beyond the limits of their precision. More accurate determinations must be made before trends of this sort can be established with any certainty.

In the monohalosilanes the HSiH angles are all within 15' of 111°(72); the HCH angles in the monohalomethanes, on the other hand, increase from the fluoride to the iodide by about 1°30,(77).

Vibrational properties—Sines it is difficult to determine force constants in complex molecules, the stretching-frequencies of SiH bonds are often

2·


used to give some measure of the resistance of the bonds to stretching. These frequencies usually lie between 2100 and 2300 cm"1, and so are relatively free from the effects of mechanical coupling with other molecular vibrations. For di- and trisubstituted silanes in dilute solution in carbon tetrachloride, it has been found possible to express the SiH stretching frequencies in terms of a set of parameters, called E-values, one of which is assigned to each substituent(81); the E-values themselves are determined from reference-compounds, and may then be used to predict the frequencies in other compounds containing different permutations of the same groups. The SiH stretching-frequency in methylethylchlorosilane, for example, is given by the sum of the E-values for methyl, ethyl and chloro-groups. Where chlorine, bromine and organic groups are concerned, the method is very successful; it has scarcely been tested for compounds containing silicon bound to fluorine, nitrogen or iodine, since much of the information available about these compounds has been obtained in the vapour phase, and the E-values only apply to solutions. The presence of oxygen bound to the silicon atom in question introduces complications; the oxygen atom transmits the effects of other groups bound to it, and allowance has to be made for this. It would be interesting to see if nitrogen behaves in the same way. The E-values are related to Gordy's electronegativities, but the relationship is not linear; they are also related to Taft's à-constants (82), which are themselves derived from a study of the rates of hydrolysis of esters(83). Attempts have also been made to relate the SiH stretching-frequencies to the electronegativity of any substituents at silicon*84»85*. The stretching-frequencies in SiH3-compounds have not been so thoroughly studied. Those for the halides in the vapour phase are given in Table 2.3 with the values for the halomethanes; it is clear that where degenerate and mean frequencies are concerned the CH-frequencies are very much the more sensitive to substitution. In more complex molecules it is often difficult to be sure which SiH mode gives rise to a particular band; this difficulty is made worse because symmetrical and unsymmetrical modes are usually rather close together, and so it is not possible to make useful comparisons.

As far as deformation frequencies are concerned, these are much more likely than the stretching-frequencies to be disturbed by mechanical coupling with the vibrations of other groups. In simple SiH2-compounds there is some correlation between the electronegativity of the substituents and the SiH2 "wagging" frequency, the frequency being higher in compounds containing electronegative substituents(89). For the SiH3-group, the symmet-

THE S lH BOND 21

rical and asymmetrical modes are usually close together, sometimes separated by only a few wave-numbers; in the halides the parallel and perpendicular bands can be distinguished by the rotational detail in the

TABLE 2.3. —MH3 STRETCHING-FREQUENCIES IN COMPOUNDS MH3X

Compound

CH3F CH3CI CH3Br CH3I SiH3F SiH3Cl SiH3Br S1H3I GeH3Cl

Frequency, cm-1

Ai

2964 2966 2972 2970 2206 2201 2200 2192 2121

E

2988 3036 3056 3060 2196 2195 2196 2206 2129

Ax-f- 2E 3

2980 3013 3028 3030 2199 2197 2197 2201 2127

86 86 86 86 69 69 69 87 88

latter, but in more complex molecules it is extremely difficult to assign a band to a particular mode with any confidence. It is therefore impossible to decide whether the symmetrical deformation modes vary with the electronegativity of the substituent, a relation that holds reasonably well for the methyl compounds of a number of elements(90). The SiH3-deformation frequencies seem to be roughly characteristic of the other atom present, but the range of variation is much smaller than in the methyl compounds.

These empirical studies have two main uses. They can be helpful in predicting the frequencies of new compounds; this is much more satisfactory where SiH stretching vibrations are concerned than with deformation modes. In the second place, they may help to indicate whether it is the mass of a substituent that determines its effect upon the vibration concerned, or whether electronic influences are more important. It seems probable that for SiH, as for CH deformation frequencies it is the electronic properties of a substituent that are the more important(90), but there is not enough useful experimental information to decide how far this is true.

Magnetic properties—The chemical shifts for protons in SiH3-com-pounds vary with the nature of the remaining group in a way which is rather different from that observed in methyl compounds(12).In the mono-halomethanes, for example, the observed chemical shifts in infinitely dilute solution are proportional to the Pauling electronegativity of the substit-


uent(e6)*. The chemical shifts of the protons in the halomethanes and the halosilanes are plotted in Fig. 2.1 against the Pauling electronegativity of the halogen atom; the SiH resonance is much less sensitive than the CH resonance to changes in the substituents, and although the values

8-0

7-0

6-0

5-0

2 0 2-5 3-0 3· 5

Pauling electronegativity

4-0 4-5

FIG. 2.1.— Proton resonance chemical shifts in the halomethanes and halosilanes, plotted against the Pauling electronegativity of the halogen atom.

there for the heavier halosilanes fall on a single straight line, the value for fluorosilane deviates substantially from this. If Taft cr*-constants(83)

are taken instead of Pauling electronegativities, a very good straight line is obtained for the halosilanes, but the effect is spoilt if other silyl compounds such as methylsilane are included. Some factor must influence the proton resonances of the silyl compounds that does not affect their methyl analogues; this is probably connected with silicon's rf-orbitals.

As the hydrogen atoms in methane are successively replaced by halogen, the proton resonances shift to lower fields; this is illustrated in Table 2.4

* Magnetic anisotropy of the C-X bond has a small but perceptible effect(Ma).

THE S lH BOND 23

by the values for the series of chlorinated(91) and fluorinated(92) methanes. In the analogous derivatives of silane the same is qualitatively true for the chlorides(12), although the chemical shift is changed much less by each successive chlorine atom than in the analogous carbon compounds. In the fluorosilanes, however, the resonance shifts to high field with increasing fluorine substitution. This has not yet been satisfactorily explained.

TABLE 2.4.—CHEMICAL SHIFTS OF MH3X, MH2X2, and MHX3, IN τ UNITS UNLESS OTHERWISE STATED ( C e H 1 2 = 8*56)

C H 3 o CH2Cl2

a

CHCl3a

CH3F0 CH2F2c CHF3c

6-97 4-66 2-69 5-74 4-55 3-75

SÌH3CP SiH2a2ò SiHCV S1H3F0

SiH2F2ö S1HF30

5-41 4-60 3-93 5-24 5-29 5-47

aMeasured relative to cyclohexane as internal standard(91). ôMeasured and given as above(12). cMeasured and given as above<M>.

As far as the σ-bonds are concerned, the central atom would be expected to become more electron-withdrawing with each additional fluorine; the remaining hydrogen atoms would then become successively less shielded, and so the proton resonance should move to] successively lower fields. This is what is observed in the fluoromethanes. It has been pointed out(93>

that "double bond character", or (/j-*d)7c-bonding between fluorine and silicon might reduce the effect of fluorine substitution upon the polarity of the SiH bonds, and so might reduce the low-field shift expected; this cannot explain the observed high-field shift with increasing fluorine substitution unless the degree of π-bonding per fluorine atom increases as the number of fluorine substituents increases. On the other hand, any (/?-► -bonding will of necessity affect the J-orbitals of the silicon atom; this will in its turn affect the mixing of excited states with the ground state, and so will alter the paramagnetic term contributing to the chemical shift (see p. 18). If the energy of the i/-orbitals is effectively raised by (p-+d)n-bondmg then the paramagnetic term would be reduced in magnitude, and a high-field shift with increasing fluorine substitution might result; since the paramagnetic term is rather small for hydrogen chemical shifts<12>, this explanation is not very convincing. The problem of the hydrogen chemical shifts in unsymmetrical molecules such as these is much more complicated


than in silane, and factors such as the effect of fluorine substitution upon the overlap in the SiH bonds, upon electron-circulations at silicon, and upon permanent molecular dipole moments may also be involved, but it is not immediately obvious why chlorine substitution should lead to a low-field, and fluorine substitution to a high-field shift.

Couplings with magnetic nuclei—Values for J(29SiH) have been obtained for a large proportion of the known SiH3-derivatives of the elements of Groups IV to VII of the Periodic Table(12). The coupling is very roughly characteristic of the Periodic group of the other element present; it increases from Group IV to Group VII, and is rather less within any one group for the first element than for the second (and in two cases measured) the third. The same general pattern is shown by the coupling constants with 13C in the analogous methyl compounds(94), but there are some differences in detail. In the spectrum of disilyl sulphide the 29SiH satellites are not single peaks but quartets, with a multiplet separation of 0·70±0Ό4 c/sec; the satellites of disiloxane could not be resolved under the same conditions, and appeared as sharp, single peaks, so there may be some interaction between the silyl groups in the sulphide that does not occur in disiloxane(12). In the fluorosilanes, the F-H couplings increase with increasing fluorine substitution(12), just as in the fluoromethanes(92). Both the 29SiH and the HH couplings in the halosilanes increase in magnitude with increasing halogen substitution; there is a roughly linear relation between the two quantities for all the compounds for which both have been measured except the two fluorosilanes, disiloxane and trisilyl-amine (see Fig. 2.2)(12). It has been suggested that the HH couplings in monosilane and the other derivatives are negative* 12-95).

Correlation of trends—There is no correlation between the mean stretching frequencies and the lengths of the CH bonds in the monohalo-methanes ; fromfluoromethane to iodomethane the bond lengths remain more or less constant, while the stretching frequencies increase. In the mono-halosilanes, however, the SiH bond length may increase very slightly from the fluoro- to the iodo-compound, but the mean stretching frequencies are effectively the same. If there were any close correlation between the length and the stretching frequencies of the SiH bonds, the mean frequencies in bromo- and methylsilane ought to be roughly the same; in fact, they differ by 27 cm-1.*69·96) On the other hand, in trifluorosilane, where the SiH bond is appreciably shorter than in monofluorosilane, the stretching frequency of 2315 cm-1 is considerably higher(97). Perhaps more precise determination of the bond lengths may clarify the position.

THE SlH BOND 25

There is, however, a much closer correlation between MH stretching frequencies and MH3 angles in the monohalomethanes and monohalo-silanes(97a).

3 0 0

2 8 0

260

2 4 0

220

200

S i H j l K ^ S i H j S r

* " ( S i H 3 ) 2 0

SiH,Br ,

SiH2F2

20

C/S

3 0

FIG. 2.2.—H-H coupling constants plotted against 29Si-H coupling constants for a number of simple silicon compounds.

It is rather easier to compare the chemical shifts and the infrared stretching frequencies of silicon-hydrogen bonds, for there is more useful information available. In the alkylchlorosilanes there is a linear relationship between the two properties(98). For arylalkylsilanes the vibrational


frequency is much the less sensitive to substitution(98); in the series of compounds from dimethylphenylsilane to triphenylsilane the vibrational frequency changes by only 4 cm-1, as against a change in chemical shift of 1-0 ppm. The relationship, however, may still be roughly linear; in all of these series of compounds, an increase in infrared stretching frequency accompanies a shift in nuclear resonance to low field. In the f luorosilanes, on the other hand, the vapour-phase infrared stretching frequencies increase markedly with increasing fluorine substitution, while the hydrogen resonances shift a very little to higher fields(12> 3°.69.97>. It therefore appears that there is no general correlation between chemical shift and vibrational frequency for silicon-hydrogen bonds, although in certain series of compounds it may be possible to discover empirical relationships between the two properties.

3. CHEMICAL PROPERTIES The chemical behaviour of the SiH bond is determined by many factors.

When considered in relation to the CH bond, however, three points are of particular importance. The first, naturally enough, is energetic. The energy

TABLE 2.5.— ENERGIES OF BONDS FROM CARBON, SILICON, GERMANIUM AND TIN, MEASURED IN KCAL, WITH THE COMPOUND (GIVEN IN BRACKETS) USED IN THE DETERMINATION

Attached atom

H C Si F Cl Br I O S N

Carbon

99 83 76(SiQ

116(CF4) 78(CC14) 68(EtBr) 51 (Mel) 86 (general) 65 " 73 "

Silicon

76(SiH4) 76(SiQ 51 (silicon)

135(SiF4) 91(SiCl4) 74(SiBr4) 56(SiI4)

108(SiOt) ?70(SiS2)c

?77(MeeSi2NH)d

Germanium

69a

81(GeCl4) 66(GeBr4) 51(GeI4)

104(GeO,)&

Tin

60« 52(SnMe4)

76(SnCl4) öSiSnBrJ 65(SnMe,I)

The values are from Cottrell (Ref. 50), unless otherwise stated. aSee footnote (b) to Table 2.1; the value is from Ref. 51. feRef.99. There is uncertainty about the energy of the GeO bond, which may be much

less than that of the SiO bond (see Ref. 99a). ^Pauling gives 54 kcal (Ref. 100), and Kriegsmann (Ref. [101 ), [using vibrational

spectra, puts the SiS bond energy in (Me8Si)8S at 63 kcal. dThis value is estimated from vibrational spectra.

THE S iH BOND 27

of the SiH bond is less than that of the CH bond, while the energies of the silicon-carbon and silicon-silicon bonds are also less than those of the carbon-carbon and carbon-silicon bonds respectively (see Table 2.5). On the other hand, where the more electronegative elements are concerned, such as oxygen or the halogens, the energies of bonds to silicon are greater than those to carbon. It follows that SiH bonds are more likely to react to form bonds between these electronegative elements and silicon than are CH bonds to react in an analogous way; in other words, the SiH bond is likely to be a stronger reducing agent. At the same time, these thermo-dynamic factors do not by themselves determine the chemistry of molecules. The CH bond does not react at room temperature with a number of systems towards which it is thermodynamically unstable. Where the SiH bond is concerned, activation effects appear to be less important; this once more is particularly true for reactions with compounds containing electronegative elements.

Finally, there is the questionTof the mode of reaction to consider. Most of the reactions of the SiH bond probably involve the displacement of a hydride ion from silicon. The analogies in chemical behaviour between the SiH bond and carbon-halogen bonds have been noted before(102), and may be correlated with the probable polarity of the SiH bond (Si+-H-)(49·103'104). Reactions of this type are usually catalysed by Lewis acids, and in many ways resemble the hydride displacements of organic chemistry(105), but there is an important difference in mechanism. Although the organic reactions often involve the formation of carbonium ions, repeated attempts to detect the presence of siliconium ions have been unsuccessful·106»107)*, and polymolecular intermediates are usually postulated in these reactions of silicon compounds. A number of reactions of the SiH bond in a triorganosilane are stereospecific<108).

There are one or two reactions of SiH bonds, notably the formation of the alkali metal silyls(109»110) and the catalysed H/D exchange between silane-c/4 and hydrogen chloride, which probably involve the displacement of protons rather than hydride ions from silicon. A number of reactions are also known where radicals are almost certainly concerned; the reactions in which silicon hydrides are added across multiple carbon-carbon bonds are the most notable of these(m).

The particular reactions of the SiH bond are discussed in more detail below. Where possible, monosilane has been taken as a reference compound;

* Evidence has been put forward for what has been called a limiting siliconium ion mechanism in one case(107a).


the behaviour of substituted silanes towards the same reagents under the same conditions may then in principle be used to deduce the effect of substitution at silicon upon the chemical behaviour of the SiH bond. Unfortunately this is usually a pious hope rather than a profitable exercise. It is rare that the conditions under which the reactions have been studied are sufficiently similar for comparison, and the physical properties of monosilane, in particular its volatility and its low solubility in polar media, may influence its chemistry in a misleading way. A further word of caution is necessary. In places the reducing properties of such compounds as disilane or siloxene are referred to. These compounds contain other reducing groups, such as Si-Si bonds, besides SiH bonds; therefore their behaviour towards many reagents, though interesting, is not a reliable guide as to the reactivity of the SiH bond. Finally, it must be emphasized that to treat bonds as isolated from the rest of the molecule is in itself a drastic approximation; like so many approximations, it is useful in that it helps in explaining the reactions of complex compounds, but its limitations must always be remembered.

1. Stability towards heat and irradiation—Monosilane begins to decompose at about 400°, forming (ultimately) hydrogen gas and silicon(112). Appreciable amounts of disilane have been obtained by heating monosilane at low pressure to about 470-480°(113); this implies that silyl radicals, SiH3-, may be formed in the process of thermal decomposition.

Monosilane does not absorb ultraviolet radiation at wavelengths down to 1850Â, and is not decomposed by radiation of this wavelength. In the presence of mercury vapour, however, mercury resonance radiation will decompose monosilane, forming hydrogen gas and a solid containing silicon and some hydrogen(114). In the presence of mercury vapour, mercury resonance radiation will also induce the free-radical addition of silanes to multiple carbon-carbon bonds (see below). Silyl radicals have apparently been obtained in an argon matrix at 4-2°K by ultraviolet photolysis, and their electron resonance spectrum, consisting of the expected quartet, has been described(114a).

2. Reactions with electronegative elements—Silicon hydrides are not affected by nitrogen gas at 25° or lower temperatures, and are usually prepared in an atmosphere of nitrogen. Monosilane reacts with active nitrogen, however, giving hydrogen as the only volatile product, though with methylsilanes some ammonia and hydrogen cyanide are formed as well(115). Although specific studies with elementary phosphorus, arsenic, antimony, sulphur, selenium or tellurium do not appear to have

THE SiH BOND 29

been made, the behaviour of these elements towards other silyl derivatives shows that any reaction with SiH bonds must be extremely slow at room temperature in the absence of a solvent*27· 11β·117· 118). Monosilane reacts very violently with oxygen even at — 180O(119); the system mono-silane/oxygen has been studied kinetically in some detail(120). Pure monosilane may not catch fire spontaneously in air(114), but the gas as usually prepared burns with a yellow flame when released into the atmosphere, leaving a brown powdery residue that contains some SiH bonds. The products of the complete combustion in a plentiful supply of oxygen are of course silica and water(120).

The reaction between monosilane and fluorine does not seem to have been investigated, but with chlorine and bromine there is a very violent reaction at room temperature, and the tetrahalosilane is produced*121·122):

SiH4+4Br2 = SiBr4+4HBr

Controlled bromination at low temperatures has been used to prepare bromosilane(122·123):

SiH4+Br2 = SiH3Br+HBr

No reaction was observed between iodine and monosilane at a pressure of half an atmosphere either at room temperature or after two hours at 50°; in the presence of aluminium iodide, however, considerable amounts of mono- and diiodosilanes were formed in the same time at 50O(124)*. Hydrogen was also produced, showing that some of the substitution must have involved hydrogen iodide formed as a result of reaction with iodine (see below, reactions of silane with hydrogen halides) :

SiH4+I2 = S1H3I+HI

SiH4+HI = SiH3I+H2.

Trialkyl and triarylsilanes react with iodine at room temperature in polar media, and though the reaction is catalysed by Lewis acids, no catalyst is necessary; diethylsilane, for example, when treated with iodine at 0° and warmed to room temperature gave both mono- and diiododiethylsilane when the reaction-mixture was distilled in the presence of copper wire(125). The reluctance of monosilane to react with iodine is therefore surprising, and may merely reflect the difference in the phase in which the reactions take place. The order of reactivity of the halogens with SiH bonds is that which would be expected from a consideration of the energies of the sili-

* But see Ref. 124a.


con-halogen, halogen-halogen and hydrogen-halogen bonds; this suggests that the reaction with fluorine would be extremely violent. Thiocyanogen, a pseudohalogen though not an electronegative element, does not attack SiH bonds at room temperature(125a).

The reactivity of the SiH bond towards oxygen appears to be reduced if some of the hydrogen atoms bound to silicon are replaced by organic groups or halogen atoms. The explosion temperatures of the methylsilanes at a given pressure increase as the number of methyl groups increases(126); cyclohexylsilane may be distilled safely in the air(24), while trichlorosilane (although its flash-point is low) can be stored in stoppered bottles at room temperature and handled quite safely in dry air. On the other hand, rate studies of the reactions between triorganosilanes and iodine, together with the observation that trichlorosilane does not react with iodine at room temperature, suggest that in this system at least the different substit-uents have qualitatively different effects upon the reactivity of the SiH bond(127). It has been shown that electron-withdrawal from silicon hinders this reaction, implying that nucleophilic attack at silicon is not the primary step; Deans and Eaborn(127) have suggested that the reaction proceeds through electrophilic attack by iodine at the hydrogen atom bound to silicon, the function of the catalyst being simply to polarize the iodine molecule.

3. Reactions with hydrides—(i) Group IV. Monosilane does not react with hydrocarbons at room temperature, and the system CH-SiH is chemically stable; at high temperatures or in the presence of a catalyst such as a Lewis acid, however, silicon hydrides will react with a range of hydrocarbons, both aliphatic and aromatic*128"-131*. The reaction between methane and trichlorosilane is typical; it occurs under pressure and at about 500° in the presence of boron trichloride(129) :

475° CH4+HSiCl3 = CH3SiCl3+H2

Press., BC1,

Alkyl halides react in much the same way with silicon hydrides(132), though this reaction should properly be considered in Section 5:

290° SiHCls+MeCl -> MeSiCl3

A1.0,

The reaction between benzene and trichlorosilane(130) appears at first sight to be similar:

275° C6H6+HSiCl3 = CeH5SiCl3+H2

bomb

THE SlH BOND 31

The first step in the reaction with the aromatic ring may, however, consist of addition to two adjacent carbon atoms, followed by elimination of hydrogen; in that case the reaction would be an example of the addition of silicon hydrides to multiple carbon-carbon bonds, which is discussed in more detail in Section 7.

The sharpness of the proton resonance lines in the nuclear resonance spectra of the partly-deuterated monosilanes dissolved in cyclohexane(12)

shows that at room temperature and in the absence of a catalyst there is no appreciable hydrogen exchange, though monosilane and bromosilane exchange hydrogen with silylpotassium in an ether solvent(132a). Ethyldi-chlorosilane exchanges its hydrogen atom bound to silicon in the presence of a platinum catalyst at 150°, but there is no exchange at that temperature without the catalyst(133). In certain derivatives of monosilane, however, the hydrogen atoms and other groups bound to silicon may be moderately labile at room temperature. Almost all of the halides containing two or three SiH bonds show some tendency to disproportionate; chlorosilane, for example, forms monosilane and dichlorosilane(31) :

2SiH3Cl = SiH4+SiH2Cl2

For most of these compounds, disproportionation is slow at room temperature, though the process is apparently accelerated by almost any impurity; pure samples of iodosilane, for example, can be kept for many months in sealed tubes at room temperature without decomposition*27 »134). The derivatives of monosilane that are most liable to break down in this way are the fluorides<19), but even with these the tendency to disproportionate is less than was at first believed(135) ; the reaction is catalysed by hydrogen flouride(135). Derivatives of disilane disproportionate much more readi-ly(8'136). The disproportionation of the halosilanes is strongly catalysed by aluminium halides(31) and by base(137~139).

MacDiarmid(27'140) has suggested that the decomposition of the fluorides involves an intermediate in which two silicon atoms are linked by what amounts to a coordinate bond and a hydrogen bridge:

H—Si Si—H I \ H / I F H H

Similar intermediates have been proposed as taking part in the catalysed disproportionation of some alkylsilanes(141). Since fluorine substitution


increases the power of a silicon atom to accept electrons from an external donor (see p. 66), and since the fluorosilanes have been thought to be associated in the liquid phase (but see p. 54), this mechanism seems plausible, although it does not account for the catalytic effect of hydrogen fluoride. At the same time, care must be taken not to over-emphasize the differences in this respect between the fluorosilanes and the other halosilanes, for the latter compounds are also liable to disproportionate. In the catalysis by aluminium halides, it is possible that bridged intermediates are formed between the silicon atom and the aluminium halide(141), rather as in the aluminium halide dimers, while bases probably act by coordination to silicon. Why this should have a catalytic effect, however, is not clear. At all events, this catalysis by base and by Lewis acids is a very general feature of the chemistry of silicon compounds.

A somewhat different form of disproportionation has been observed in silyl derivatives of elements of Groups V and VI. Although trisilylamine, disiloxane, disilyl sulphide and disilyl selenide are relatively stable compounds*31· 118> U2), they all have some tendency to decompose when kept at room temperature for long periods, forming monosilane and a solid residue*30·124). The process is probably induced by the presence of traces of impurity, since thoroughly purified samples of these compounds have been kept at room temperature for several months without appreciable decomposition(27). It appears that a hydrogen atom migrates from one silicon atom to another:

x(SiH3)2S = xSiH4+(SiH2S)x

The reaction between disilyl sulphide and trimethylamine(27) is very similar; in this case, however, the amine combines with the solid residue to give what is probably a polymeric adduci:

xMe3N+x(SiH3)2S = xSiH4+(SiH2S,NMe3)x

The same kind of process was detected by Stock in the decomposition of samples believed to contain mono- and disilylamines(142):

x(SiH3)2NH = xSiH4+(SiH2NH)x

Here the reaction may have been catalysed by the ammonia present in the mixture. The decomposition of silyl trifluoromethyl sulphide(143) is also analogous, though it differs in that the migrating atom is fluorine and not hydrogen :

S1H3SCF3 = SiH3F+SCF2

THE SiH BOND 33

This too is very sensitive to the presence of impurity. No such decomposition has been reported for the N-methyl-silylamines (SiH3)2NMe(144) and SiH3NMe2

(123), though the latter compound is not very stable at room temperature(123). It is probable that the well-known ß-elimination reaction of tetraorganosilanes is also similar (see p. 83-5); an example of this is given below:

Et3SiCH2CH2Cl = Et3SiCl+CH2CH2

The mechanism of these processes in unknown, nor is it clear whether the migration is inter- or intramolecular. The proton nuclear resonance spectrum of disilyl sulphide, it is true, shows that there is some interaction between the hydrogen atoms bound to different silicon atoms(12), but this may well be quite unconnected with the reaction under discussion.

(ii) Group V. Monosilane does not react with pure, dry ammonia(1), but the reaction is fairly rapid in the presence of amide ion; hydrogen is evolved, and the ultimate product is probably some condensed derivative of tetraminosilane(110, :

SiH4+4NH3 = Si(NH2)4+4H2

The barrier to reaction with ammonia must therefore be kinetic. Triethyl-silane also reacts with liquid ammonia containing alkali metal amides, but not with liquid ammonia alone(14e):

2Et3SiH+KNH2 = (Et3Si)2NK+2H2

It is interesting that the system SiHNH is apparently unstable, for no stable compound containing this grouping has been prepared(142). The decomposition of unstable derivatives of this type usually involves labile SiN bonds rather than any reaction of SiH bonds themselves, although a second mode of decomposition in which SiH bonds migrate was discussed in the previous section; the whole matter is dealt with in more detail in Chapter 5.

Since tetrasilyl hydrazine has been prepared and is relatively stable at room temperature(147), it is clear that the NN bond is not affected by SiH bonds in the same molecule at temperatures up to at least 100°.

Phosphine and iodosilane do not react together at room temperature in several hours(lie). If monosilane and phosphine are heated together to about 400° at low pressures, however, silylphosphine is formed(148):

SiH^+PH, = SiH3PH2+H2

(iii) Group VI. Monosilane does not react with pure water in quartz vessels at room temperature(1). Here too the barrier to reaction must be

3


kinetic rather than thermodynamic, for hydrogen is evolved in the presence of even traces of alkali:

\ \ - S i H + H 2 0 = -SiOH+H2

Unless a fairly large amount of alkali is present the hydrolysis is not complete; a jelly-like solid is formed that contains hydrogen bound to silicon(1). With strong alkali, however, the reaction is usually quantitative, and has been frequently used for estimating the hydrogen bound to silicon in a variety of compounds*1»27»142). Monosilane reacts in a similar way with alcohols, and mono- and dimethoxysilanes have been prepared in this wayd49). triorganosilanes react with silanols in the presence of appropriate catalysts, giving disiloxanes(150).

The reaction between SiH bonds and the hydroxyl groups of water or alcohols is catalysed by a variety of substances; among the effective catalysts are base(1»151), acid(1·152), silver ion(153), and copper powder*149· 154>155). Base catalysis has been studied in aqueous alcohol(98); kinetic studies have shown that under these conditions the reaction between triorganosilanes and hydroxyl groups involves both nucleophilic attack at the silicon atom and electrophilic attack of a solvent molecule upon the hydrogen atom bound to silicon. In a series of triarylsilanes the rates of hydrolysis vary with the Hammett σ-constants of the substituents<156); this shows that electron-release to silicon hinders the reaction, and that as a consequence the silicon atom must be more negative in the transition-state than in the free molecule. Similarly, trialkylsilanes containing flu-orocarbon groups in the sidechains are much more rapidly hydrolysed than the analogous simple alkylsilanes(157). In the acid-catalysed hydrolysis of triorganosilanes in aqueous alcohols, kinetic evidence suggests that the intermediate is of the form(9)

g

[H20... SiR3—H...H—OH2]+

Silver ion has a marked catalytic effect upon the reaction between alcohols and triorganosilanes*153*, although it is reduced in the process; there is no direct evidence to show whether it is the silver ion or its reduced form that is the catalytic agent. The mechanism of the catalysis, and of the catalytic effect of copper powder upon the reaction between SiH bonds and alcohols, is unknown.

Since SiH bonds are not extensively broken when SiH3-derivatives are treated with hydrogen sulphide at room temperature(27), it may be

THE SlH BOND 35

concluded that under these conditions there is little reaction between the two species. The system SiH-OH, like the analogous SiH-NH, appears to be inherently unstable, but the mode of decomposition is not clear. Since 90% yields of disiloxane can be obtained from the hydrolysis of disilyl sulphide(27), if silanol is formed as an intermediate it must disproportionate according to the equation:

2S1H3OH = H20+(SiH3)20

On the other hand, yields of disiloxane obtained from the hydrolysis of other SiH3-derivatives such as bromosilane(158>159) are always much lower, and both monosilane and hydrogen are produced. This suggests that other modes of decomposition are possible; for example, hydrogen might be formed by some process like the one below:

XS1H3OH = (SiH20)x+xH2

Alternatively, the acid produced in the hydrolysis may catalyse the reaction between SiH bonds and water.

Silanethiol (silyl mercaptan) has been prepared, but seems to be rather unstable ; it decomposes to give disilyl sulphide and hydrogen sulphide (27·140). This is in keeping with the stabilities of the analogous derivatives of Group V elements—silyl phosphine being a stable compound at room temperature, but the primary and secondary silylamines being so unstable as to elude attempts to isolate them.

(iv) Group VII. In the presence of aluminium chloride, monosilane reacts with hydrogen chloride in the gas phase at 100° to form both mono- and dichlorosilanes(31) :

SiH4+HCl = SiH3Cl+H2

The dichlorosilane may be produced by further substitution, or by dis-proportionation of the monochlorosilane, for this disproportionation is catalysed by aluminium halides(31). Monosilane and hydrogen chloride may be heated together to 200° without reaction if no catalyst is present(31). Hydrogen bromide(122) and iodide(160) react similarly with monosilane in the presence of the appropriate aluminium halide, and halogen derivatives of disilane(8) and a number of alkylsilanes(10'14) have been prepared in much the same way. Triorganosilanes, on the other hand, react with hydrogen halides in solution without catalysts*155.161). The reason for this difference is not clear, but perhaps monosilane would react similarly in solution; as with the reaction between SiH bonds and iodine, catalysis may only be necessary in the vapour-phase reactions.

3*


The reaction probably involves the displacement of negatively-polarized hydrogen from silicon; the catalyst may serve to polarize the hydrogen halide molecule, or to act as a bridge in the removal of hydride from silicon:

AIX3+HX = H+AlXr

SÌH4+H+ADC = S1H3X+H0+AIX3

Aluminium halides also catalyse hydride displacements in organic chemis-try(105).In the monosilane-hydrogen halide system,however,another reaction takes place at the same tiniie. It has been observed that if attempts are made to prepare chlorosilane-rf3

(17) or iodosilane-c/3(30) from silane-</4

and the hydrogen halide, then the halosilane produced contains a considerable proportion of hydrogen*17·30*; the residual silane has not apparently been examined to determine the extent of exchange in it, but there is no reason to doubt that exchange takes place independently of halogen substitution. This exchange almost certainly means that a positively-polarized hydrogen atom is displaced from silicon:

AIX3+HX = H+ADÇ

SiD4+H+ = SiHD3+D+

The H/D exchange may not need a catalyst, but the uncatalysed system has not been studied.

4. Reactions with oxides and oxyacids—The information available about the reactions between SiH bonds and oxides or oxyacids is rather patchy. Monosilane is said not to react with concentrated sulphuric acid at room temperature(1), but a number of other compounds containing SiH bonds do so(125'ie2). Solutions of sulphurous acid are reduced by siloxenes to sulphoxylic acid, and the reaction differentiates between SiH bonds and SiSi bonds (which do not react in this way)(163). Solutions of hydrogen peroxide and of perdisulphuric acid react violently with siloxenes(162); disilane(1) will reduce solutions of potassium permanganate (to manganese dioxide), while triethylsilane reacts with chlorsulphuric acid, nitric acid, and potassium dichromate when warmed with them(125). Oddly enough, triethylsilane does not react with solid potassium permanganate(125). Carboxylic acids will react with a number of organosilanes in the presence of Lewis acids, forming hydrogen and a carboxy-derivative of the silane(le4) :

R8SiH+R'COOH = R3SiOCOR'+H2

THE SiH BOND 37

At about 700°, SiH bonds react quantitatively with carbon dioxide, forming carbon monoxide(3). The carbon-oxygen bonds in certain ethers are broken by SiH bonds(164a).

5. Reactions with salts and halides—Aqueous solutions of a number of transition metal compounds are reduced by silicon hydrides or their derivatives. Disilane reduces cupric ion (to copper hydride), and ferric ion (to ferrous ion)(1); siloxenes precipitate tin, lead, arsenic, antimony, bismuth, nickel, palladium, platinum, copper, silver, gold, and mercury from solutions of their salts(162), while the reaction between silicon-hydrogen bonds and aqueous mercuric chloride has been used in estimating the former(165) :

R3SiH+2HgCl2+H20 = R3SiOH+Hg2Cl2+2HCl

Triorganosilanes have also been found to react with a number of metallic and metalloid derivatives, giving in many cases the metal either free or in a lower oxidation state, together with a compound containing the anion bound to silicon(125):

2R3SiH+2HgBr2 = 2R3SiBr+Hg2Br2+H2

This reaction emphasizes the analogy between the SiH bond and the silicon-halogen bonds (see Chapter 4). Where more than one SiH bond is present, stepwise replacement may be possible(166).

C6H13SiH3+AgNCS = C6H13SiH2NCS+Ag+£H2

Monosilane itself does not react with chloroform or with carbon tetrachloride, but the higher silanes are liable to react violently with these compounds in the presence of a catalyst such as aluminium chloride forming substitution-products in which the silicon chain is preserved intact(167>ie8). This suggests that the SiH bonds may be more reactive in the higher silanes than in monosilane, and this suggestion is to some extent supported by the reaction between disilane and iodine{169). This gave iodinated disilanes, and there was no sign that silicon-silicon bonds were broken in the reaction; a rough thermodynamic calculation indicates that of the two possible reactions given below, the latter ought to be energetically preferred if the SiH bond has the same energy in disilane as in monosilane:

Si2He+I2 = Si2H5I+HI

Si2He+I2 = 2SiH3I


This point is discussed further in Chapter IV. It is possible that the differences between the behaviour of monosilane and the higher silanes are due to differences in activation energies or in reaction conditions, for triethylsilane reacts with alkyl halides under reflux, forming the appropriate hydrocarbon and triethylchlorosilane(170):

Et8SiH+C6Hl3Cl = Et3SiCl+CeH14

Alkyl halides may also react with hydrogen bound to silicon in a different way, forming silicon-carbon bonds(132).

Monosilane might be expected to react with a number of non-metallic halides that can be reduced fairly easily, and it does indeed react with the chlorides of tin (IV), phosphorus (V) and antimony, forming tetra-chlorosilane(171) : it also reacts with boron trichloride in the presence of méthylène radicals to give diborane(172). Triethylsilane reacts slowly with antimony (III) fluoride at room temperature, presumably to give triethyl-fluorosilane(125) : on the other hand, mono-, di- and trifluorosilanes have all been prepared from the reactions between the appropriate chlorosilanes and antimony (III) fluoride so the reaction with SiH bonds must be relatively slow at room temperature™.

6. Reactions with metals and with organometallic compounds—Silicon hydrides react under certain circumstances with Grignard reagents, forming the metal hydride and a silicon-carbon bond*13·173·174* :

Pl^SiHa+PhMgBr = Ph3SiH+HMgBr

Monosilane itself reacts in much the same way with organosodium and organolithium compounds to form tri- and tetrasubstituted derivatives, the product depending to some extent upon the solvent(175). With triorgano-silyl-metallic compounds there is a complex reaction; when triphenyl-silylpotassium is treated with triphenylsilane, for example, tetraphenylsilane is formed, together with what probably contains polymeric material(17e) :

PhgSiH+PheSiK->Ph4Si

The reaction-products depend upon the metal used and upon the solvent; with lithium derivatives, for instance, there is extensive silicon-silicon coupling(176) :

Ph3SiLi+Ph3SiH = Ph3Si.SiPh3+LiH.

Reactions between alkali metals and hydrogen bound to silicon have been most extensively studied for triorganosilyl compounds, and in

THE SlH BOND 39

particular for triarylsilanes. Several of these compounds react either with a sodium/potassium alloy in diethyl ether(177_8) or with lithium dissolved in tetrahydrofuran(179), giving hydrogen and a silyl-metallic compound:

R3SiH+Li = R3SiLi+JH2

Sodium derivatives of monogermane(180) and monostannane(37) were successfully prepared in liquid ammonia solution, but attempts to prepare silylsodium or silylpotassium in this way have not been successful; the amount of hydrogen gas evolved shows that the monosilane is completely ammonolysed, though the solvolysis is relatively slow(110). Silylpotassium, however, has been obtained as a crystalline solid from the reaction between monosilane and potassium; it is also formed when disilane reacts with either potassium or potassium hydride, though in the former case at least there are other reactions which occur as well(132a). Diphenylsilyllithium has also been prepared (181) by the splitting of the silicon-silicon bond in the appropriate disilane with lithium metal. Digermane reacts with alkali metals to give the metal germyls, but as with disilane there are side-reactions, possibly of the form:

Ge2H6+M = Ge2H5M + l/2H2

7. Reactions with multiple carbon-carbon bonds—Silicon-hydrogen bonds in a wide range of silicon hydrides and their derivatives will add across multiple carbon-carbon(182~186) bonds. Addition takes place when the systems are heated or irradiated, and in the presence of peroxide catalysts ; these are typical conditions for reactions involving free radicals, and the products formed usually fit in with the rules for free radical reactions. The addition, however, is also catalysed by platinum metals and by some of their compounds; the products obtained under these conditions are often different from those when the same reactants combine in the presence of the peroxide catalysts, implying that the mechanism of the metal-catalysed process may not involve radicals(186) (but see Ref. 133). The reaction is of considerable industrial importance; ion of SiH bonds has also been observed across —C = N<186a) and > C = 0(186b) systems.

8. Brief summary of reactions of germanium and tin hydrides — Monogermane is less stable than monosilane both thermally(6ö) and to irradiation(187); monostannane is thermally less stable still(137-188). On the other hand, monogermane is much more stable towards oxygen than is monosilane*189), while monostannane only reacts with oxygen when heated, and its decomposition is inhibited at room temperature by the presence of small amounts of oxygen(37«190). Germanium-(191) and tin-


hydrogen(190) bonds are readily broken by halogen; they are more stable than SiH bonds towards ammonia*37*180*. Trialkylgermanes are remarkably resistant to alkali-catalysed hydrolysis in alcohol/water mixtures, while trialkylstannanes hydrolyse rather more quickly than trialkylsilanes and by a different mechanism(192). Monogermane reacts with hydrogen halides in the presence of aluminium halides to give halogermanes(193); chlorostannane has been prepared by the uncatalysed reaction between hydrogen chloride and stannane at —70° (193a).To wards oxides, oxyacids and halides, both GeH and SnH bonds behave as reducing agents, the SnH bond being rather the more powerful·190·194·195*. With organometallic compounds, germanium(196) hydrides react initially to give the Group IV-alkali metal compound and a hydrocarbon:

R3GeH+RLi = R3GeLi+RH

This is different from the behaviour of the analogous silicon(174«175) and tin(197) compounds. As has already been mentioned, the germanium-hydrogen(180) and tin-hydrogen(181) bonds are broken by alkali metal in liquid ammonia; trichlorogermane is believed to be extensively ionized in water even in the presence of considerable concentrations of mineral acid, giving GeCl3~ ions(197«198). Germanium(45a) and tin(200·201* hydrides will add to carbon carbon multiple bonds rather more readily than the analogous silicon hydrides.

9. Comparison of SiH, GeH and SnH bonds—The general chemical behaviour of the three species is similar, and many of the differences observed—such as the differences in thermal stability—can be put down to the probable differences in MH bond energies. The comparative stability of the hydrides of the heavier elements to oxidation and to hydrolysis is surprising; lack of precise thermodynamic data makes it impossible to decide how far this is due to activation effects. The frequent appearance of solid hydrides of germanium in the reactions of germanium hydrides(2o2)

can probably be explained in terms of the relative energies of the Ge-H and Ge-Ge bonds, since the solid is likely to contain cross-linked germanium atoms rather than divalent germanium; unfortunately the relevant bond energies are not known with sufficient accuracy to make such an explanation quantitative.

On the whole, there is a definite impression left by a consideration of the chemical properties of the hydrides of silicon, germanium and tin, that the GeH- and SnH-bonds are rather more liable to ionize to give anionic germanium or tin compounds and protons than are SiH bonds

THE S lH BOND 41

to behave in a similar way. This impression is partly due to the greater stability of the GeH and GeCl bonds to hydrolysis; it is impossible, for instance, to study the ionization of the SiH bond of trichlorosilane in water because the SiCl bonds are rapidly solvolysed. A careful study of the comparative acidities of the hydrides of silicon and germanium in nonaqueous media would be interesting in this context. The study of the change in nuclear resonance chemical shift with solvent of the proton bound to silicon in trichlorosilane provides a beginning to this(203); the results are, however, very difficult to compare with analogous results for chloroform, because the possibilities of different bonding hydrids from silicon and from carbon complicate the matter. No corresponding measurements with trichlorogermane have yet been described. It is notable that while derivatives containing the SnCl" ion are known, and the etherate Et2O.HSnCl3 has been prepared,(204) the parent hydride HSnCl3is unknown; the ease of ionization of the hydrogen atom in the trihalides of the Group IV elements: MHC1 3 ;ÄMCLJ+H+ has been associated with the relative stabilities of the divalent and the tetravalent states, the ion MCI" being regarded as a derivative of M(II) rather than M(IV)(198).

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THE SlH BOND 43

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44 VOLÀTILE SILICON COMPOUNDS

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THE S l H BOND 45

(12e> R. L. SCHALLA and G. E. MCDONALD, Nat. Advis. Comm. Aeronaut. Tech. Note No. 3405, 1955; Chem. Abs., 49, 7248 (1955).

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THE SlH BOND 47

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CHAPTER 3

THE HALIDES OF SILICON

1. FORMATION OF SILICON-HALOGEN BONDS

SILICON-HALOGEN bonds are formed in reactions between the free halogens and elementary silicon(1), silicon hydrides<2), siloxanes(3,4), silicon sulphides^, silicon selenides(4) and compounds containing silicon-carbon**' and silicon-silicon(6) bonds; in the reactions between hydrogen halides and silicon itself(7) or compounds containing silicon bound to hydrogen(2>8), oxygen(3'9), sulphur(4»10), selenium(4), nitrogen(11), phosphorus(12), carbon(13); in the reactions between various inorganic and organic halides and silicon-hydrogen bonds (14'15); in the reactions between organic halides and elementary silicon(16), and in the reactions between inorganic halides such as boron trifluoride, and silylamines, siloxanes, and other compounds in which silicon is bound to a donor atom (see Chapters 5, 6). Silicon halides have been prepared by most of these processes, and, in addition to these, a number of reactions in which one halogen atom already bound to silicon is replaced by another are useful synthetically. The latter groups are discussed with the chemical properties of the halosilanes.

2. PREPARATION OF HALOSILANES

The tetrahalides that contain only one species of halogen are prepared by the action of the free halogen or the halogen hydride on elementary silicon(1) or some silicon compound such as ferrosilicon(19), but more specific methods are necessary for the preparation of the mixed tetra-halides(18'19'20). Halogen! replacement is often involved, and the most commonly used methods are mentioned later in this chapter.

Trichloro-(7) and tribromosilane(21) are made by the action of the appropriate hydrogen halide on silicon heated either by itself or with copper. The mechanism of the reaction is not certain, but it seems likely that the reaction is analogous to the preparation of organohalosilanes from alkyl halides, where an alkyl derivative of copper is believed to be formed as an intermediate(22). The tetrahalide is produced with the tri-halosilane, and must be separated from it.

48

THE HALIDES OF SILICON 49

Mono- and dichloro- (23), bromo-(2) and iodosilanes(8) have usually been prepared by the reaction between monosilane and a hydrogen halide (see p. 35), or (in the case of the bromides(2>23) and iodides(24)) the free halogen. A method of preparing iodosilane from phenylsilane by the action of hydrogen iodide has recently been described*25»26*:

C6H5SiH3 + HI = C6H6 + S1H3I

If iodine is used, then substitution at silicon leads to the formation of phenyliodosilane instead of the required product. Treatment of a phenyl-halosilane with the hydride of a different halogen gives a mixed dihalo-silane(26a).

Mono-, di- and trifluorosilane have generally been prepared from one of the other halosilanes. There is no reason to believe that the reaction between, say, monosilane and hydrogen fluoride would not give f luorosilanes, but hydrogen fluoride is an unpleasant substance to handle. It has therefore been the common practice to prepare partly-fluorinated monosilanes by treating the appropriate chlorosilane with some mild fluorinating agent like antimony trifluoride(7) or zinc fluoride(27). With compounds containing Si-H bonds there is always some reduction of the antimony halide in this reaction, but the yield is reasonably high (at least 50%). If iodosilanes are used as starting compounds, lead (II) fluoride is a satisfactory fluorinating agent(28). Monofluorosilane could be prepared free of hydrogen fluoride by the action of an excess of trisilylamine on boron trifluoride(29)» but this method has not apparently been used synthetically. A mixture of all the f luorosilanes has been obtained from the action of atomic hydrogen on tetrafluorosilane at 2000o(30). The tetrahalides of germanium(31) and tin(32) are prepared by reactions analogous to those used in making the silicon compounds. Partly-chlorinated and brominated monogermanes(33)

and monoiodogermane(34) are also known. Trichloro- and tribromoger-mane are made by the action of the hydrogen halide upon some compound of germanium like germanium (II) sulphide*35·38* ; the mono- and dichlorides and bromides(33) are prepared by treating monogermane with the appropriate hydrogen halide and aluminium halide, while iodogermane was prepared by the action of iodine on monogermane at low pressure(34). Mono- and difluorogermanes have recently been obtained from the appropriate bromogermanes and silver (I) fluoride(34a).

Trichlorostannane has not been characterized, but the etherate Et20.HSnCl3

has been prepared and a number of derivatives of the chlorostannite ion,

4


SnCl^", are also known(37); monochlorostannane has been made by treating monostaruiane with hydrogen chloride at low temperatures(37a).

3. PHYSICAL PROPERTIES

Once again it is convenient to consider first the physical properties of the SiX bonds of the simple tetrahalides, which are set out (with those of their carbon, germanium and tin analogues) in Table 3.1. One of the most striking anomalies is the high boiling point of tetrafluorostannane, which is probably a reflection of the greater polarizability and coordinating power of tin (IV) than of the other elements. Although tetrafluorostannane is more salt-like than the other tetrafluorides, it is not a true salt(59). While the boiling-points of the tetrafluorides rise with the atomic weight of the central atom, tetrachlorosilane and tetrabromosilane boil at lower temperatures than their carbon analogues; this suggests that some intermolec-ular forces which are strong in tetrachloro- and tetrabromomethane are reduced or absent in tetrachlorosilane and tetrabromosilane, and perhaps also in the corresponding germanes. The interaction may be of the (p->d)r*-type, between halogen atoms of different molecules. This would be insignificant in the fluorides because of the inaccessibility of the 3i/-orbitals of fluorine, but might be important in the heavier halides of carbon; where silicon, germanium and tin are concerned, any intramolecular (p -> d)n -bonding between the central atom and the halogens bound to it would make the /^-electrons of the halogens less available for intermolecular π-bonding, and so lead to weaker intermolecular forces(60).

When the observed bond lengths are compared with those calculated from Pauling's atomic radii, the silicon-halogen, germanium-fluorine and carbon-fluorine bonds are all found to be shorter than predicted, though in some cases the discrepancy is small. This shortening has been attributed both to the polar character of the bonds concerned, and also to (/?->ii)7c-bonding (sometimes referred to as double-bond character). The Schomaker-Stevenson correction(61) attempts to allow for the former effect by introducing a bond shortening term into the expression for the internuclear distance, based on the difference in electronegativity between the bonded atoms. If the Pauling electronegativities are used, the agreement between observed and calculated bond lengths for the carbon-fluorine, and for the silicon-chlorine, -bromine, and -iodine bonds becomes rather good, but the correction does not always improve the correspondence

TABLE 3.1.—PHYSICAL PROPERTIES OF THE MX BONDS AND OF THE MOLECULES MX4

1 Bond energy (a> kcal

Bond length, A,a» Observed

(t>>Calc. uncorr(,>

(c>Calc. corrO^

Force const. md/A (n ■»>

«"Obs.

(«»Predicted

NMR chem. shift<r> (ppm)

rcl. to CF3COOH (e. std)

eQq, c/s«s-f)

m.p., °C("»

b.p., °C<«>

F

c 1 116

1-317 ±005«»

1-49

1-35

6-7

61

—11-9

—185

—128

SÌ 1

135

1-54 ± 0 2 ( 0

1-89

1-69

5-9

4-2

83-3

—90

—96»

Ge |

-

1-67 ±-03«*)

1-94

1-74

5-5

4-1

990

—15

—36*

Sn

-

-

-

ci c 1

78

1-77")

1-76

1-71

3-6

3-4

-81-85

—23

705» 76

si 1

91

2-01±-02</>

216

205

2-7

2-4

-40-8

—70

57

Ge |

81

208 ±02«»

2-21

2 1 0

2-7

2-3

-51-32

—49-5

86-5

Sn

76

2-31 ±-01<»>

2-39

2-28

ß

2-5

1-8

-48-2

—36

114

Br

C

65

(EtBr)

1·94±·02«>

l-vl

1-84

1-5

2-9

1 536-0

94

dec.

Si

74

2-15 ±-02(»

2-31

2-22

2 0

2 0

294

5

155

Ge |

66

2-21 ±-02<fc>

2-36

2-27

1-9

2 0

352

26

186-5

Sn

65

2-44±02<*>

2-54

2-45

1-9

1-6

328

33

203

C

51

(Mel)

2·15±·02<*>

210

210

--

-2130

171

dec.

I Si

56

2-43 ±-02(*>

2-50

2-44

\J(Q)

1-6

-1329

124

290

Ge

51

2-50±-03<*>

2-55

2-49

14«»

1-5

-1492

144

348

Sn

65

(Me,SnI)

2·64±·04<*>

2-73

2-67

\.yo)

1-2

-1389

144-5

346

»«»Values from CottreU, Ref. 38. <6)Ref. 39. (c>Ref. 40. ""Ref. 41. <*'Ref. 42. ">Ref. 43. (">Ref. 44. «»'Ref. 45. ">Ref. 46. <»Ref. 47. <*>Ref. 48. <»Ref. 49 <»»> Ref. 50. ,n»Ref. 51. <*>Ref. 52. WRef. 53, 53a. (r»Ref. 54. <»>Ref. 55. «>Ref. 56, 57. WRef. 58


between the observed and calculated bond-lengths; moreover, the silicon-fluorine bond length is still much too short. Although the Schomaker-Stevenson correction has been severely criticised(e2), there seems to be no reason for rejecting either mechanism, and it is probable that the bond-lengths are affected by interactions of both kinds. The difference between r(SiX) and r(GeX) is much greater for the fluorides than for the other halides, implying that the interaction between silicon and fluorine is responsible for the "bond shortening" is less important between germanium and fluorine. The measurements, however, are not very precise. The GeF and Si F bond lengths differ by much more than do the atomic radii of germanium and silicon, it could be argued that this shows (p -> d)7>bond-ing to be weaker between fluorine and germanium than between fluorine and silicon.

The nuclear resonance chemical shifts of fluorine nuclei are largely determined by changes in the paramagnetic term(62), itself arising through interaction between the ground state and excited states involving unpaired electrons (see p. 18). This leads to an expected correlation between chemical shift and "ionic character" of the bond to fluorine; the more like a fluoride ion the fluorine atom is, the higher the field should be at which resonance is observed*62·63*. There is a roughly linear correlation between increasing fluorine chemical shifts (to high field) in the tetrafluorides of carbon, silicon and germanium and the decreasing Pauling electronegativity of the Group IV atom(64), and this has been taken as in accordance with the prediction given above. It is therefore unfortunate that the resonance for the fluoride ion appears to low field of both germanium and silicon tetrafluorides. It is hardly possible that the fluorine nuclei are more shielded in the tetrafluorides than in the fluoride ion, while changes in the paramagnetic term are not likely to be responsible for the anomaly; the paramagnetic term induces a shift to low field, and a high-field shift can only be produced in this way by reducing the paramagnetic term, yet in the fluoride ion (where the resonance is to low field of the others) the paramagnetic term is zero for reasons of symmetry. One possible explanation is that the species studied as fluoride ion was in fact something else; the measurement was made using aqueous potassium fluoride solution, where there may have been some hydrolysis. Solvation could also be at least partly responsible*62'63* but the shifts involved (50-100 ppm) are larger than most solvent effects(65).

Attempts to take into account the effect of (j? -> </)7r-bonding from fluorine to silicon upon the fluorine chemical shift have so far been un-

4 ·


satisfactory. An argument based on the interpretation of fluorine chemical shifts in terms of changes in diamagnetic shielding has been used to deduce that (p -> ^ -bond ing in fluorosilanes will lead to a low-field shift(62-63) ; since, however, the paramagnetic term is much more important than diamagnetic effects for fluorine chemical shifts(62), this treatment is unlikely to be adequate. The paramagnetic term depends on the reciprocal of Δ£>3, where Δ£ is a mean excitation energy; silicon has more relatively low-lying excited states than carbon, and therefore fluorine bound to silicon might on this basis be expected to give a resonance to low field of fluorine bound to carbon in an analogous molecule. Any (p -* α)π-bonding is likely to lower the energy of the non-bonding /^-electrons of fluorine, and so, by increasing Δ£, at least partly to offset the low field shift. At the same time, there is the 1/r3 factor to consider; this means that excited states centred on the fluorine atom are likely to be more important than excited states centred on the attached atom; alternatively, for a fixed ΔΕ the fluorine resonance should shift to high field with increasing M-F distance*. Though the interpretation of fluorine chemical shifts is not a simple matter, and there are other factors which would have to be taken into account in a thorough treatment of the problem, it seems clear that π-bonding should have some effect, whether through changes in diamagnetic shielding, excited state mixing or magnetic anisotropy; this makes any correlation between chemical shift and electronegativity of the adjacent atom(64) all the more likely to be fortuitous.

The quadrupole coupling constants of the tetrachlorides, bromides and iodides of the elements of Group IV could in principle provide valuable information about the bonding in these compounds (68~~70). The quadrupole coupling constant of a nucleus with a quadrupole moment is zero in an electronic environment which has spherical symmetry. It has been argued that under normal conditions it is only changes in the symmetry of the p-electrons of the atom's valence shell that will alter this, the effect of ^-electrons being zero, and of ^-electrons being negligible(e9). In the halide ions, for example, the nuclei are surrounded by closed shells, which always have spherical symmetry, and the quadrupole coupling constants are therefore zero ; in the free halogen atoms the /^-electrons of the valence-shell are one electron short of this configuration, and so the quadrupole coupling constants of these species are taken to correspond to an unbalance of one

* I am indebted to Dr. A. D. Buckingham for much of this interpretation.


electron in the /7-shell. Once this has been established, the observed quad-rupole coupling constant for a halogen nucleus in any compound can in principle be used to give a measure of the unbalance of the /?-electrons surrounding that nucleus, and so to give some idea of the polarity of the σ-bond which holds the halogen atom to the rest of the molecule. The more closely the halogen atom in that bond corresponds to a halide ion, the lower will be the quadrupole coupling constant; if, on the other hand, the quadrupole coupling constant is higher for a halogen atom in some compound than for the free atom, then the halogen atom in that compound should form the positive end of the bond dipole. In chlorine compounds,

Tor example, the observed quadrupole coupling constants are almost always less than for the chlorine atom; in the chlorine molecule the value is the same as for atomic chlorine, while it is greater than this only if chlorine is bound to oxygen or fluorine(72). Unfortunately, however, the quadrupole coupling constants are affected not only by these changes in σ-bond polarity, but also by changes in the hybridization of the halogen's σ-orbit-als; moreover, any π-bonding between the halogen atom and the group to which it is attached will also reduce the quadrupole coupling constant. The interpretation of the measurements in terms of molecular structure is therefore a somewhat subjective process. Some authors have made assumptions about hybridization and calculated bond polarity(69~71), while others have made assumptions about bond polarity and calculated hybridization<73). Within the former of these categories quite widely different assumptions have been made(74«75), leading to different interpretations of the observed values; moreover, some of the more fundamental assumptions of the theory have been criticized*76*. The quadrupole coupling constants of the tetrachlorides, bromides and iodides of silicon, germanium and tin are all considerably less than for the analogous carbon compounds, which does strongly suggest that the bonds from the heavier elements of halogen are (as expected) more polar than the CX bonds; it is also generally agreed that there is some "double bond character" [(/>-xf)7r-bond-ing] in the halides of silicon, germanium and tin, but the detailed interpretation of the different values is still uncertain. There is a roughly linear relationship between Allred and Rochow's electronegativity values for the Group IV elements(77) and the quadrupole coupling constants of the tetrahalides, but this may (as with the fluorine nuclear resonances) be a chance correlation.

The force constants for some of the tetrahalides are compared in Table 3.1 with the values predicted by Siebert's formula. There is a very rough


correlation between a force constant which is considerably higher than predicted and a bond considerably shorter than the sum of the atomic radii, and the high force constants have been associated with (p^>d)n-bond-ing, since it is claimed that ionic character, while shortening the bond, would not give rise to a high force constant(62). There is little apparent reason why the predicted values for the force constants, which were derived from a study of alkyls, should necessarily be expected to apply exactly to the halides, and the force constants calculated from observed vibration-al frequencies vary quite widely with the method used to calculate them; differences between force constants calculated from observed spectra and the values estimated from atomic parameters are probably of qualitative rather than quantitative significance.

Physical Properties of the Hydride-halides

The melting- and boiling-points of the hydride-halides of carbon and silicon are given in Table 3.2. Of these, difluorosilane has the highest boiling-point of the f luorosilanes, but the same is true of dif luoromethane among the f iuoromethanes. Although it has been suggested that the f luorosilanes are associated in the liquid phase, the recalculated value for the entropy of vaporization of difluorosilane does not support this suggestion; Onyszchuk's vapour pressure measurements for fluorosilane itself only indicate very slight association(79), while the proton and fluorine nuclear resonance spectra of mono-, di- and trif luorosilanes are very little affected by dilution(93). Nuclear resonance chemical shifts are particularly sensitive to association, and the very mall changes observed over a tenfold change in concentration—from 90% to 10% solutions—implies that any association in mono-, di- and trifluorosilanes must be very alight*.

The lengths of the silicon-fluorine bonds increase significantly as hydrogen replaces fluorine (see Table 3.3); the same may happen, though to a lesser extent, to the silicon-halogen bond lengths in the chloro- and bromosilanes, but there are not enough precise measurements to be sure of this. In terms of the electrostatic explanation for the short SiX bonds, this lengthening could be explained by saying that the silicon atom becomes less electron attracting as hydrogen replaces halogen, because it is less positively-polarized. It could also be argued that replacement of halogen

* Dilution shifts for proton resonances of the fluorosilanes in tetramethylsilane are all less than 01 ppm, though in trichlorofluoromethane a dilution shift for trifluoro-silane of 0-25 ppm was observed; the largest fluorine dilution shift is 4 ± 3 ppm, for difluorosilane(98).


by hydrogen would lessen the contraction of the G?-orbitals of silicon, and so reduce (p->d)n-bonàing, and there is no direct evidence to favour the one explanation against the other; in the fluorinated methanes, however, where π-type interactions are likely to be less important because of the inaccessibility of carbon's 1p and 3d orbitals, the same increase in the length of the bonds to fluorine is even more marked. In the chlorinated and brominated methanes the carbon-halogen bond lengths do not alter so much or so consistently with changing substitution, partly because of the increasing size of the halogen atoms concerned.

TABLE 3.2.—PHYSICAL PROPERTIES OF THE HYDRIDE-HALIDES OF SILICON

Melting point, (degrees) Boiling point, (degrees) Entropy of vap.

at b.p. (kcal/moP)

Melting point, (degrees)

Boiling point, (degrees) Entropy of vap.

at b.p. (kcal/mol°)

SiH3F SiH2F2

99(d) 26(d>

SiH3Br

—94(0> 1·9<σ>

2H( e>

SiHF3

_122<a> —131<a> —76<e) _944<e>

20(e) 21( e )

SiH2Br2| SiHBr3

1

— 7 0 · 1<*> 66{o)

21·8(/>

—730» 112·2<β>

21·0(β>

SiH3Cl

— 118(Ö)

—30·4ω 19·8('»

SiH3I

—57·0(*> 45·6<*> 21·6(β>

SiH2Cl2

—122(Ö)

8·3<'> 21·4('>

SiH2I2

150<e> 20·8(β)

SiHCl3

—126-5< > 3ΐ·7(β)

20·8(*>

SiHI3

8«) 220dec.<*> 38(e)

(a>Ref. 7. «»Ref. 23. (*)Ref. 78. (d>Ref. 7. Onyszchuk(78> obtained a boiling point of 86° and an entropy of vaporization of 22-5 kcal/mol° for this compound. (e)Ref. 80 </>Ref. 81. (0)Ref. 2. <*>Ref. 21. «>Ref. 8. (^Ref. 82. (fc>Ref. 83.

The force constants for the silicon-halogen bonds in the halosilanes are fairly close to those observed in the tetrahalides; the calculation of precise force-constants is at present not sufficiently reliable for a satisfactory study of changes in force constant with substitution to be made. The stretching-frequencies of the silicon-halogen bonds appear in a region of the spectrum where mechanical coupling between these and other vibrational modes is to be expected, and so a comparison of these observed frequencies in a series of molecules is liable to be misleading.

The chemical shifts of the fluorine resonances in fluorosilanes change irregularly as fluorine is replaced by hydrogen ; a rather similar irregularity in fluorine chemical shifts has been observed in the methyl- and ethyl-


TABLE 3.3.—CHANGES IN MX BOND LENGTH WITH SUBSTITUTION IN THE HALOSILANES AND HALOMETHANES (VALUES IN Â)

CF4«*> 1·317±·005

CHF3(e>

1·332±·005 CH2F2«>

1·358±·001 CH,F(«>

1·385±·001

SiF4<» 1·54±·02

SiHF3('>

1·565±·005 SiH2F2

(/)

1·577±·001 SiH3F<«>

1·594±·005

CCl4(o

1-77 CH3C1(*> 1-781

CBr4<*) l-942±-002

CH3Br(ff) 1-939

SiCl4«*> 201 ±-02 SiH3Cl(Ä>

2·0479±·0007 SiBr4</>

2·15±·02 SiH3Br<*»

2·209±002

«*>Ref. 39. (0>Ref. 40. <ORef. 42. <*)Ref. 43. (e)Ref. 84. (/>Ref. 85. <*>Ref. 86· (*)Ref. 87. «>Ref. 88. (')Ref. 89. (*>Ref. 46. (')Ref. 47. (™>Ref. 90. (n>Ref. 91. <*>Ref. 92.

fluorosilanes(e6»67), and the observed chemical shifts in the three series of compounds are given in Table 3.4. In the alkylfluorosilanes the effect has been explained*66·67* in terms of a competition between changes in σ-bond polarity (which would be expected to lead to a high-field shift

TABLE 3.4. —FLUORINE CHEMICAL SHIFTS IN THE FLUOROSILANES, METHYLFLUOROSILANES AND ETHYLFLUOROSILANES, IN PPM RELATIVE TO CC1 3 F

R

H(a> CH8

(»)

R8SiF

217 155-3 174·7<ο

R2SiF8

151 1300 142-5

RSiF8

109-5 134-2 138-6

SiF4

161-8

(a>Ref. 93. (Measured in dilute solution in CC13F). (0>Ref. 66, 67. These were measured relative to CF8COOH as external standard, extrapolated to infinite dilution in CC14, and have been rather arbitrarily corrected to the Φ scale(Ma) using the value for CF3COOH as external standard of 78-5. (c)Ref. 93a.

as fluorine was replaced by hydrogen) and (/?->d)7r-bonding. If the total amount of π-bonding per silicon atom were more or less constant, then the π-bonding per fluorine atom would increase as the number of fluorine atoms was reduced. Since it was assumed that π-bonding would lead to a low-field shift by reducing the electron density at fluorine, the observed irregularity could be accounted for in terms of competition between the two effects. It has already been pointed out that π-bonding will not neces-


sarily lead to a low-field shift (see p. 55); the lack of correlation between the SiF bond lengths and the fluorine resonance chemical shifts, (for values, see Tables 3.3 and 3.4), suggests that some specifically magnetic interactions are more likely to be responsible for the irregular changes in chemical shift. The magnetic anisotropies will be different in the different molecules, for example, and this could well account for at least some of the anomaly. The HF couplings increase as hydrogen replaces fluorine, but the ^SiF couplings do not change regularly as the number of fluorine atoms bound to silicon is increased: the change in the latter from trifluorosilane to tetrafluorosilane is larger than the change from fluorosilane to trifluorosilane (see Table 3.5). The quadrupole coupling constants of

TABLE 3.5—COUPLINGS IN THE FLUORINE NMR SPECTRA OF THE FLUOROSILANES The values, which with one exception are from Ref. 93, are in c/s

Compound

SiH3F SiH2Fa

SiHF3

SiF4

J29SiF

229±0-6 297±3 274·£±0·3

178(a>

JHF

45·8±0·1 60·5±0·1 95·8±0·1

(a>Ref. 54.

the halogen nuclei in monochloro- and monobromo-derivatives of methane, silane and germane are given in Table 3.6. These values were obtained in the vapour phase from microwave spectra, and so are not strictly comparable with the values for the tetrahalides, but they also correlate roughly with the electronegativities of the Group IV elements as measured on the scale put forward by Allred and Rochow. The cupole moments of the monohalosilanes, like those of their methyl analogues, are all

TABLE 3.6.—eQq VALUES in MONOHALOMETHANES, -SILANES AND -GERMANES, IN C/S

CH3X SiH3X GeH3X

Cl

75-13«*) 40·0(α)

46<α>

Br

577·0(α> 336(α> 380«*>

I

1934(&> 1240(θ

(<*>Ref. 94. «»Ref. 94a. «^Ref. 94b.


TABLE 3 . 7 . — D I P O L E MOMENTS OF SOME HALOMETHANES A N D HALOSILANES, IN DEB YES

M O S T OF THE VALUES WERE MEASURED BY MICROWAVE SPECTROSCOPY

! F

CH 3 X SiH3X C H X 3

S iHX 3

1.79(a) l-27(c> 1·65(*> 1-26«*)

Cl

1·87<*> 1·30(α.β)

0·95<Λ)

0·85<»

ΒΓ

1 ·80 (Ο)

1·32('> 0·99(ί> 0-79(fc)

I

1-65«»

«*>Ref.95. <ö)Ref.96. <ORef. 97. «*>Ref. 98. (e>Ref. 99. WRef. 100. <*>Ref. 101. (*>Ref. 102. (*>Ref. 103. (»Ref. 104. <*>Rcf. 105.

roughly equal (see Table 3.7); the interpretation of dipole moments of molecules like these in terms of molecular structure is at present an uncertain and subjective process, but the low dipole moments of the halosilanes may arise from {p -► d)K-bonamg. Such halogermanes as are known do not differ in any unexpected ways from the halosilanes in their physical properties; the trialkylfluorostannes, on the other hand, are high-melting solids, and are believed to contain organotin cations such as Me3Sn+(106).

4. CHEMICAL PROPERTIES

Halosilanes are of great importance in the preparation of silyl and other silicon compounds, for they enter into a wide range of chemical reactions; from silicon halides it is possible to prepare compounds containing silicon bound to hydrogen(107), carbon*108· 109), silicon (110), germa-nium(111), nitrogen*11»112), phosphorus(113), oxygen(9), sulphur(4), selenium(4), and a number of groups such as the pseudohalogens(8). The chemical properties of the halosilanes may conveniently be considered under five headings: redistribution reactions, in which exchange occurs with other halogens (or pseudohalogen groups) bound to silicon or some other non-metal; reactions with elements; reactions with hydrides; reactions with metal compounds; and reactions with electron-donors to form adducts. In most cases in which a silicon-halogen bond is broken, the reactivity increases with the atomic weight of the halogen atom, as might have been anticipated from the energies of the silicon-halogen bonds ; where reactions with heavy metal compounds are concerned, in which the halide of the heavy metal is formed, the large free energies of formation of the heavy metal iodides make iodosilanes still more reactive than the other silicon


halides. Towards some reagents, such as zinc fluoride, chlorosilanes(27)

are more reactive than iodosilanes(93). Most, if not all, of the reactions mentioned above are probably initiated

by nucleophilic attack at silicon, and the intermediates may contain 5-coordinated silicon; it could indeed be argued that the adducts of the fifth category (such as SiF4.2NH3) are merely stabilized intermediates of this sort. There is little kinetic evidence in support of this idea, although no alternative mechanism is more definitely established*. The mechanisms of the solution reactions between halosilanes and hydroxides, NH-com-pounds or metal salts are all probably similar, and are unlikely to proceed through the formation of siliconium ions(114) ; nothing is definitely known about the mechanisms of the vapour-phase or of the solid-vapour reactions. The reactions are discussed in more detail below.

1. (a) Thermal stability—The tetrahalides of silicon that contain only one halogen species are relatively stable to heat, but (as might be expected) the thermal stability becomes less as the atomic weight of the halogen increases. Tetrafluorosilane, for instance, is not decomposed at very high temperatures (of the order of 2000°)(30); tetrachlorosilane, while stable in quartz apparatus at temperatures up to at least 800°, decomposes at higher temperatures in the presence of hydrogen; chlorosilanes such as Si10Cl22 have been isolated from the system(115). Very pure silicon has been obtained from the thermal decomposition of tetraiodosilane(116), which, however, is thermally much more stable than carbon tetraiodide(117). The hydridehalides and the mixed halides are less thermally stable. The disproportionation of the hydride-halides has already been discussed (see p. 31).

(b) Exchange reactions etc.—The mixed halides are also liable to disproportionate, but as a rule the compounds are stable at room temperature and disproportionation does not occur except in presence of a catalyst or when the system is heated(20*118). Trifluoroiodosilane is as might have been expected one of the least stable of these compounds(119). Mixtures of two different tetrahalides do not exchange halogen atoms except when heated*20,118'119,120) (or, perhaps, in the presence of a catalyst); under those conditions, an equilibrium is set up of the form:

2SiX4+2SiY4 = SiX3Y+2SiX2Y2+SiXY3

* At present there is not enough evidence to decide whether the reactions are bimo-lecular processes of the SN2 type, or whether the rate-determining step involves breakdown of an intermediate, formed in a fast pre-equilibrium, which contains a 5-coordinated silicon atom.


These have been described as entropy-driven reactions(120), since the total number of bonds from silicon to any of the other species present remains constant; this description implies that bond energies are independent of the rest of the molecule, and this may not be exactly true, particularly in view of the changes in such properties as bond length with substitution. Halosilanes also enter into exchange-reactions with halides or pseudo-halides of such elements as aluminium, germanium, phosphorus and antimony(120) (see also Chapter 7); the exchanges are not usually rapid at room temperature(120a).

Monofluorosilanes interact with boron trifluoride at temperatures below —120° to give what appear to be equimolar adducts; nothing is known about the structures of these compounds(120b).

2. Reactions with elements—Tetrafluorosilane is reduced by atomic hydrogen at temperatures over 2000°(30), the products including all the partly-fluorinated monosilanes, monosilane itself and a small proportion of higher silanes. Trichlorosilane reacts to some extent with hydrogen in the presence of aluminium at 250°, but here it is difficult to separate reduction from disproportionation induced by any aluminium chloride formed(121); again, monosilane and the chlorinated monosilanes are produced. Tetrachloro-(122) and tetraiodo-(123)silane may be reduced to silicon by the action of hydrogen at temperatures above 1000°. On thermo-dynamic grounds, some of the halosilanes would be expected to react with elementary halogens; these reactions do occur, and have been extensively used in the sythesis of mixed halides of silicon, but once again fairly vigorous conditions are usually necessary*19,124,125). Tetrachlorosilaae(12e)

does not exchange chlorine with liquid chlorine at 0°, nor with gaseous chlorine at 100°; tetraiodosilane does not exchange iodine with elementary iodine in solution in organic solvents at temperatures below 130°, but exchange is fairly rapid in molten tetraiodosilane at 140O(127). Trialkyliodo-silanes form addition-compounds with molecular iodine, just as do alkyl iodides(122); the alkyl compounds may, however, be rather more stable.

Silicon-halogen bonds are stable to oxygen at room temperature, but oxyhalosilanes have been prepared by heating halosilanes and oxygen to high temperatures*41·128·129). There appears to be some reaction between iodosilane and sulphur at room temperature, but the behaviour of the system is rather unpredictable(130); no reaction was detected between iodosilane and either selenium or tellurium(4·130). None of the silicon halides reacts with nitrogen under ordinary conditions, but tris(trichlorosilyl)amine was formed with other products when a glow discharge was passed through


a mixture of tetrachlorosilane and nitrogen(131). Iodosilane reacts slowly with white phosphorus, and with arsenic and antimony, at temperatures up to 100°; the products of the reactions are extremely complex, but compounds such as disilyliodophosphine have been obtained in very small yield, together with monosilane, disilane, hydrogen, phosphine and di-iodosilane(132»133). There was some evidence that trisilylphosphine was formed in small amounts in the reaction between iodosilane and phosphorus.

Halosilanes usually react with electropositive elements to form the metal halide and compounds containing silicon-silicon bonds. Hexaethyldisilane, for example, has been prepared from triethylchlorosilane and sodium(110) :

2Et3SiCl+2Na = Et3SiSiEt3+2NaCl

Some polymeric silicon compounds have been made in this way; tri-bromosilane reacts with magnesium to give the solid hydride (SiH)x

(134) :

2xHSiBr3+3xMg = 2(SiH)x+3xMgBr2

When monohalosilanes react with alkali metals, however, it is monosilane rather than disilane that is formed, together with some polymerized silicon hydride and hydrogen gas(135):

SiH8Cl+Na-SiH4+NaCl+H2+(SiH2)x

Iodosilane, however, reacts with sodium amalgam to give disilane(135a). The exact nature of the products depends upon the alkali metal used. Halosilanes do not form stable Grignard-type reagents with magnesium, and numerous attempts to isolate such compounds have been unsuccessful·8» 1 3 β · 1 3 7 ) ; bromosilane, for example, reacts with magnesium rather reluctantly, forming monosilane. Recently, however, some evidence has been obtained for the formation of unstable triorganosilylmagnesium hali-des(137). Some alkali metal derivatives of triorganosilanes have been prepared from triorganohalosilanes and alkali metals(138), but it is possible that a disilane is formed first, and that this then reacts with excess of the alkali metal(138):

2Ph3SiCl+2Li = 2LiCl+Ph3SiSiPhe

Ph3SiSiPh3+2Li = 2Ph3SiLi

3. Reaction ; with hydrides—All the silicon halides may be reduced by hydride ion or by complex hydrides to silicon hydrides*93·107.139>; this reaction has already been discussed in Chapter 2. Some most interesting


products of the reaction between lithium borohydride and tetrachlorosilane have recently been mentioned in a review<139a). Among the products, the compounds Li(H3BSiCl3) (isolated as an etherate), Li2[(BH4)2SiClJ, and H3BSi2H4 were obtained. Chlorosilanes react with certain hydrocarbons under very vigorous conditions to form SiC bonds,' although the change in bond energy is unfavourable(139b):

R3CH+R^SiCl = R3CSiR3+HCl

Halosilanes react readily with compounds containing hydrogen bound to amino-nitrogen(112·140»141); the reaction with chlorides, bromides and iodides is of the form:

-SiX + 2^>NH = -SiN<" + ^>NH2X -/ y / x ■ κ

Most of the silylamines have been prepared in this way (see Chapter 5). The reaction is reversible, and part of the driving force comes from the formation of the ammonium salt. Fluorosilanes do not as a rule react in this way except in the presence of lithium(142), which presumably removes hydrogen fluoride from the system more effectively than excess of the amine. The behaviour of tetrafluorosilane and tetrachlorosilane towards ammonia is interesting in this context; ammonia and tetrachlorosilane react together to form ammonium chloride and silicon-nitrogen poly-mers(143), while with tetrafluorosilane an adduct of formula SiF4.2NH3

is produced, which is thermally stable up to 300o<144). It seems possible that an adduct of this kind is formed initially in all of the reactions, but that the thermodynamics of the fluoride adduct are unfavourable to further decomposition; there is, however, no evidence that the reason for the adduct's failure to decompose further is not kinetic.

In contrast to this, even iodosilane does not react with phosphine, at temperatures up to 100O(133).This may reflect both the low lattice energy of phosphonium iodide and the poor electron-donor properties of phosphine towards silicon; silicon tetrafluoride is said(145) to form a very weak adduct with phosphine of formula 3SiF4.2PH3 but NMR studies give no evidence for the existence of any such adduct(146). No reaction was detected between iodosilane and arsine at temperatures up to 100O(133).

The ready hydrolysis of chlorides, bromides and iodides of silicon is well known. The essential reaction is represented by the equation:

\ \ - S i X + H 2 0 = - S i O H + H X


This is'reversible; methyliodosilane, for example, is formed by the action of hydrogen iodide on 1,1'dimethyl disiloxane and gives l,rdimethyl disiloxane on hydrolysis(9). Almost all hydroxy-compounds will react with halosilanes in this way; organofluorosilanes react in a similar way, but the equilibrium concentration of the silicon halide is much higher(147). The kinetics of the hydrolysis of triorganosilanes have been studied fairly extensively^4· 148> ; the silicon atom is more negative in the transition-state than in the free molecule, but there is still doubt as to whether the reaction is bimolecular, or whether a fast pre-equilibrium leads to the formation of a species containing 5-coordinated silicon, which then breaks down more slowly. Tetrachlorosilane reacts with hydrogen sulphide at high temperatures, forming compounds such as Cl3SiSH and Cl3SiSSiCl3

(150). Organohalosilanes react with hydrogen sulphide in the presence of base (which presumably removes hydrogen halide from the equilibrium mix-t u r e ) ( 1 4 9 ) .

2R3SiX+H2S^(R3Si)2S+2HX

This reaction has been used in preparing sulphur-derivatives of triorganosilanes, but compounds containing.hydrogen bound to silicon are usually unstable in the presence of base.

Iodosilane(130) reacts only slowly with hydrogen sulphide at room temperature, forming polymeric products. Hydrogen chloride exchanges chlorine slowly with tetrachlorosilane at 90° in the vapour phase(151), and iodochlorosilanes have been prepared by the action of hydrogen iodide on tetrachlorosilane at high temperatures(152). Once more, the mechanisms of these two reactions are not known. Organochlorosilanes react with anhydrous hydrogen fluoride to give fluorosilanes(153):

Bu2SiCl2+2HF=Bu2SiF2+2HCl

Hydrogen bromide, on the other hand, only reacts with phenylchloro-silanes in the presence of aluminium bromide(154). Chlorobromosilanes have been prepared by the action of hydrogen bromide on tetrachlorosilane at high temperatures(20), and bromoiodosilanes are formed when tetra-bromosilane reacts with hydrogen iodide at high temperatures(154).

4. Reactions with metal compounds

(a) Grignard reagents and organometallic compounds—Halosilanes will react with Grignard or other similar reagents to form the metal halide and ouncompds containing silicon-carbon bonds*108·109). Stock used this


method to prepare methylsilane from chlorosilane and zinc dimethyl(23):

2SiH3Cl+ZnMe2 = 2SiH3Me+ZnCl2

The reaction has found extensive application in the preparation of organosilanes, although some sterically hindered Grignard reagents react with halosilanes to form SiH bonds (155>156). Triorganosilylmetallic compounds will also react with silicon halides, and in most cases the reactions are similar to those involving carbon-metallic compounds; trimethyl-chlorosilane, for example, reacts with triphenylsilyl-potassium to form 1,1, l-trimethyl-2,2,2-triphenyldisilane(157) :

(CH3)3SiCl+(C6H5)3SiK = (CH3)3SiSi(C6H5)3+KCl

The reactions between polyhalosilanes and silyllithium compounds are more complicated(158).

(b) With halide ions—Tetrafluorosilane vapour exchanges fluorine with solid lithium fluoride at temperatures well below the decomposition-temperature of lithium fluorosilicate(159). Organochlorosilanes will exchange chlorine with tetraalkylammonium chlorides either dissolved in dioxane or as solids(160), and tetrachlorosilane will also exchange chlorine with solid tetraalkylammonium chlorides, though the exchange is not very rapid(161). The mechanism of the solution reaction* is apparently similar to that of the hydrolysis of halosilanes(le2).

Triorganohalosilanes will react with some alkali metal pseudohalides, forming the pseudohalosilane(163):

(C6H5)3SiCl+NH4NCS = (C6H5)3SiNCS+NH4Cl

This type of reaction, however, has found relatively little application in inorganic silicon chemistry. Iodosilane, the most convenient starting material, appears reluctant to react with solid sodium or potassium salts (unless, of course, the anion present will oxidize the SiH bonds) (130>133); from the inorganic point of view, reactions with heavy metal salts have been much more useful.

There is one more class of reactions with sodium and potassium derivatives that is likely to be of increasing importance in preparing new silyl compounds; this is the reaction between halosilanes and alkali metal derivatives of hydrides, particularly of Groups IV and V, dissolved in some inert solvent like diethyleneglycol-dimethyl ether. As yet, only one or two systems have been studied; as a result, however, two elusive and interesting compounds have been prepared for the first time. Trimethyl-


chlorosilane was found to react with the sodium derivative of bis(tri-methylsilyl)amine to give tris(trimethylsilyl)amine(le4):

[(CH3)3Si]2NK+(CH3)3SiCl = [(CH3)3Si]3N+KCl

and the same halosilane reacted with phosphylsodium in solution to form tris(trimethylsilyl)phosphine(113) :

(CH3)3SiCl+NaPH2->[(CH3)3Si]3P+NaCl

The sodium and potassium derivatives of phosphine and arsine react similarly with trimethylfluorosilane at — 20O(ie4a). This latter reaction must involve some disproportionation of the phosphine derivatives; it is discussed in more detail in Chapter 5.

(c) Reactions with salts of heavy metals. Salts of silver, mercury, lead and other heavy metals will generally react with halosilanes*48»93· 130>133) : the reactions are of the form

SiX+AgY -> SiY+AgX

but the direction of reaction is of course determined by thermodynamic properties. Various authors*130»1β5·1ββ»1β7) have formulated "conversion series" for the reactions between silicon compounds and silver salts; the silver salts are arranged in order, so that the silyl derivative of any anion in the series can be prepared by the reaction between its silver salt and the silyl derivative of any group appearing earlier in the series (that is, to the left as written); similarly, the silyl derivative of this member of the series can be converted into the silyl compound of a later member by reaction with the silver salt of the latter. The series for SiHs-compounds, covering a number of groups besides halides, is set out below(167) :

[(SiH3)aTe^]SiH3I^(SiH3)2Se^(SiH3)2S^SiH3Br^SiH3Cl->SiH3CN->

SiH3NCS^SiH3NCO->(SiH3)20->SiH3F

The series more or less follows the order of bond and lattice energies, as far as these are known, and the ease with which conversion may be brought about at room temperature shows that activation energies in these systems are usually low (though this may not always be true). The series that have been evolved for trimethylsilyl- and triethylsilyl-com-pounds(les»16e) are much the same as the one given above, except for the positions of chloride and cyanide; the experimental evidence for deciding

5


the relative positions of these two groups is not clear-cut, and the energies of the systems SiCN/AgCl and SiCl/AgCN must be very much the same. It is possible that the SiCN bond energy is higher in SiH3CN than in MegSiCN, in view of the possibly different compositions of the compounds at room temperature (see Chapter 7). Similar series could probably be worked out for mercury (Π) and lead (Π) salts, but at present not enough experimental evidence is available to do this; it looks as if the series for silver and mercury (II) salts are very similar*167»1β8).

The reactions are effective in the presence or absence of solvents, and the conditions used vary widely; in some cases, refluxing in the presence of the heavy metal salt for some hours is apparently necessary, while in others a single passage of the vapour of the halosilane over the solid salt is enough for virtually complete reaction. Yields are often of the order of 90-100%; the mechanism of the reaction is unknown. It may differ in detail from the mechanism of the analogous reaction between halo-alkanes and silver salts, since no isomerie cyanides are formed with silver cyanide; this is perhaps a reflection of the greater lability of groups bound to silicon.

Fluorosilanes have also been prepared by the reactions between chloro-silanes and fluorides of zinc(27) or antimony (III)(7); here again nothing is known about the mechanism of the reaction.

(5) Formation of adducts—It has already been pointed out that the silicon atom can under certain circumstances act as an electron-acceptor, and will form addition-compounds with a number of donor groups. The best-known of these adducts is the fluorosilicate ion, but many others have been prepared and characterized. All of the simple tetrahalosilanes combine with nitrogen bases to form adducts; silicon tetrafluoride, for example, forms the compound SiF4.2NH3 with ammonia(144), and reacts with trimethylamine to give the adducts SiF4.NMe3 and SiF4.2NMe3(171). Tetrachlorosilane reacts with ammonia to give polymers containing SiN bonds(143), but it forms addition-compounds with many tertiary nitrogen bases, including pyridine. The compounds do not usually contain more than two molecules of amine to one of halosilane(172); earlier results, which suggested that three or four molecules of amine could be coordinated, were probably obtained in the presence of water. Tetrabromosilane reacts in much the same way, but tetraiodosilane combines readily with amines in a 1:4 (iodide :amine) molar ratio(172). The 1:2 compounds probably contain 6-coordinated silicon, although there is little direct structural evidence on this point; the infrared spectrum of the compound SiCl4.2pyri-


dine is very like that of the analogous germanium compound, which has been shown by X-ray analysis to have a frms-octahedral structure(173a)-The adduct Me3N.SiF4 is a poor conductor of electricity when fused» but it is not clear whether the silicon atom is 5- or 6-coordinated (with halogen bridging in the latter case)(174). The structures of the 1:4 adducts are harder to understand. It seems possible that iodine, rather than silicon, is the electron-acceptor, since carbon tetraiodide forms some adducts with donors of this sort(175), but once again there is no direct evidence to show what the structures really are. Although tetrachlorosilane does not usually combine with more than two molecules of base, it forms a remarkable compound with urea which contains six molecules of urea to each one of tetrachlorosilane(172). This compound has been formulated as a salt, but yet again without any direct evidence on the matter.

The adducts are solid compounds, which dissociate on heating and differ in volatility; many are readily hydrolysed by water or hydroxy-compounds(172).

The hydride-halides of silicon also form 1:1 addition-compounds, but here the thermal stability of the adduct is usually very much less. The halomonosilanes and halomonomethylsilanes all combine with trimethyl-amine, giving solid products which dissociate completely in the vapour phase into the starting-materials(17e~177). These compounds can be regarded either as containing 5- or 6-coordinated silicon atoms (structure I below), or as analogous to the quaternary methyl ammonium salts (structure II). At present it is not clear which structure is correct. What evidence

Λ H 3 S i / [H3SiNMe3]+r-

NNMe3

(I) (Π) there is seems to favour the ionic formulation; the volatilities of the adducts of the monohalosilanes, for instance, increase as the atomic weight of the halogen atom increases (132>1β8»175), while for molecular addition-compounds the reverse might be expected to be true. Since, however, the heats of formation of the various adducts are not known, this argument does not have a great deal of force. The iodosilane: trimethylamine adduct gave a conducting solution when dissolved in acetonitrile(132), but the method of filling the conductivity-cell may have allowed scine hydrolysis, with the resulting formation of

5*


trimethylammonium iodide, a strong electrolyte(17e). The question is at present open, and its answer would be of considerable interest. It could be settled by conductivity-experiments in which the adduct was prepared under vacuum in the conductivity-cell. Monosilane does not form an adduct with trimethylamine(177).

The acceptor properties of silyl groups as far as external donors are concerned are sharply reduced if hydrogen bound to silicon is replaced by alkyl groups*174·176*; there was no sign from a study of the UY spectra of the ligands that either pyridine or /j-toluidine formed adducts with triethylfluorosilane, hexamethyldisiloxane or dimethyldiethylsilane at room temperature in cyclohexane solution(178).

In addition to the 1:1 compound mentioned, iodosilane forms an adduct with trimethylamine of approximate formula SiH3I.2NMes. This is stabler than the equimolar adduct, and may contain 6-coordinated silicon; its solution in acetonitrile did not conduct electricity appreciably(132).

The analogous derivatives of phosphine have been much less extensively studied. Tetrafluorosilane is said to form a compound with phosphine of formula 3SiF4.2PH3

(145), though this has not been conf irmed<14e> ; iodosilane combines with trimethylphosphine, triethylphosphine and trimethyl-arsine to give equimolar adducts(132'133), while similar compounds have been obtained from iodosilane and mono- and dimethylphosphines(179). Phosphine itself does not combine with iodosilane at temperatures up to 100O(132). Ethers do not combine readily with halosilanes at low temperatures(180), and under drastic conditions the carbon-oxygen bond is liable to be broken(m); sulphides are even weaker donors towards silicon. Tetrahydrofuran, for example, forms a weak adduct with iodosilane, dimethyl ether forms a weaker adduct still, while dimethyl sulphide appears to form no adduct at all(180). Diphenylsulphoxide forms 2:1 molar adducts with tetrabromo- and tetraiodosilane, which decompose irreversibly at room temperature; no adduct was obtained with tetrachlorosilane, substitution occurring instead, while no reaction was observed with tetra-fluorosilane(180a).

* 5. HALIDES OF GERMANIUM (IV) AND TIN (IV)

The halides of these elements are qualitatively like the halosilanes in their chemical properties, but there are some important differences in detail. To begin with, the dihalides are relatively more stable(182·183)·


The halogen atoms are much more labile; it has been shown that mixtures of two different tetrahalides of germanium or tin are liable to exchange halogen atoms at room temperature, though the rate of exchange depends upon the particular compounds concerned(118,184a). The susceptibility to hydrolysis and to ammonolysis decreases from silicon to tin(185~7), and halosilanes are often hydrolysed more rapidly than halogermanes(188); ammonolysis of the halides of tin may be complicated by the formation of ammonoanions(187). Germanium and tin halides react with Grignard reagents to form organogermanes and organostannanes(182'184), and with some heavy metal salts to form the halide of the heavy metal and the organogermanium or organo tin compound; conversion-series for trialkyl-germanium(189) and trialkyltin(190) compounds have been worked out empirically, and are very like the analogous silicon series which was described earlier in this chapter. There are one or two interesting differences, however; relative to the chlorides, bromides and iodides, the sulphides, cyanides, isothiocyanates and isocyanates become less stable from silicon to tin. This shows that as the atomic weight of M increases, the M-CN, M-NCS, M-NCO and M-S bond energies drop more rapidly than the M-halogen bond energies.

Many adducts of the tetrahalides of germanium and tin are known; tin in particular is a strong acceptor in the stannic halides, and forms numerous adducts with amines(187), alcohols, ethers and various oxygen-and sulphur-containing donors(191~194), in which one, two or more molecules of donor are present to each molecule of tin halide. Fluorogermane forms an equimolar adduct with ammonia at room temperature*34*), while tri-methylchlorostannane forms adducts with organic secondary and tertiary amines in which one molecule of base may be associated with one or two molecules of halide(195). The structures of these adducts have not been conclusively established, although the compounds of stannic fluoride containing two molecules of donor to each molecule of SnF4, are probably cw-octahedral in some cases at least(184); the ligands in GeCl42py are in a frans-octahedral arrangement(173a). The ligands of the SnF4 adducts exchange fairly rapidly in certain solvents(194). Complex chloroanions, such as SnClg, and SnCl-, are known(18e), in addition to the chloro-stannite ion already mentioned briefly in Chapter 1(195); the chloro-germanite ion is also known(193), and hydroxy-complexes such as [Ge(OH)xCl5_Je are formed when tetrachlorogermane is dissolved in hydrochloric acid(185), but the hexachlorogermanate ion is apparently not formed. Hexachlorosilicates are not known at room temperature.


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CHAPTER 4

COMPOUNDS C O N T A I N I N G SILICON BOUND TO ELEMENTS OF GROUP IV

1. THE SII ICON-CARBON BOND

So much information is available about the formation and properties of bonds from carbon to silicon and the other elements of Group IV that only a very brief summary of it can be given here. Many reviews are available, among the most important being Eaborn's recent book about organosilicon compounds(1), an article on organogermanium compounds(2)

written nine years ago and a much more up-to-date one dealing with organic derivatives of tin(3). A survey of material dealing with organogermanium compounds published between 1950 and 1960 has recently appeared(3a).

Silicon may be bound to carbon by a variety of methods. Starting from elementary silicon, organosilicon compounds are produced when alkyl or aryl halides are passed over heated silicon*1·4* :

2CH3C1 + Si = (CH3)2SiCl2 The reaction is catalysed by a number of substances, particularly transition metals, and appears to be analogous to the preparation of trichloro-silane from hydrogen chloride and heated silicon(5). It is probable that in the copper-catalysed reaction a copper alkyl is formed as intermediate, and that this or a copper halide reacts with the heated silicon to initiate the formation of the organohalosilanes(6). The details of this reaction are of considerable industrial importance.

Compounds containing SiH bonds react with many organolithium, organosodium and organopotassium compounds and with some Grignard reagents to form SiC bonds:

SiH4 + 2EtLi = Et2SiH2 + 2LiH (Ref. 7) SiH< + 4C6H5Na = (C6H5)4Si + 4NaH (Ref. 7)

C6H5SiH3 + CeH5MgBr = (C6H5)2SiH2 + HMgBr (Ref. 8) Silicon-halogen bonds react in this way much more readily, and this is one of the most widely-used of the methods of making SiC bonds:

2SiH3Br + BrMg(C=C)MgBr = SiH3C=CSiH3 + 2MgBr2 (Ref. 9)

76

SILICON BOUND TO ELEMENTS OF GROUP IV 77

The mechanism of this reaction has been studied in some detail(10); a cyclic intermediate is believed to be involved. In compounds of the type ArMe2SiX, electron-release to silicon hinders reaction, implying that the silicon atom is more negative in the intermediate than in the free molecule. Silicon alkoxides(11) and some siloxanes(12) also react in this way with Grignard reagents and with organolithium compounds.

The most important of the remaining methods for forming silicon-carbon bonds is the reaction between an unsaturated carbon compound and a silicon hydride. Typical of this type of reaction is the formation of ethyltrichlorosilane from ethylene and trichlorosilane(13):

HSiCl3 + C2H4 = CH3CH2SiCl3

It is discussed in more detail on p. 39. Among other preparative methods, silicon-carbon bonds are formed in the reaction between iodosilane and silver cyanide(13a):

S1H3I + AgNC = SiH3CN + Agi but the cyanide group is usually treated as a pseudohalogen, and the reaction will be considered in that section of Chapter VII.

Germanium- and tin-carbon bonds may be formed by the action of organic halides on the heated elements in the presence of an appropriate catalyst(14«15), by the action of hydrides or halides of germanium or tin on orgaaometallic compounds*16·17*, and by the addition of GeH or SnH bonds to unsaturated organic compounds*18·19*. All of these methods are analogous to the reactions for forming silicon-carbon bonds described above. The addition of the hydrides to multiple carbon-carbon bonds is rather easier where the heavier elements are concerned. Physical Properties

Some of the physical properties of the tetramethyls of carbon, silicon, germanium and tin are given in Table 4.1. In one or two cases the appropriate property has been measured very much more precisely for some other similar compound, when both values are given. The boiling-points are not remarkable. The observed force-constants agree well with the predicted values, which is no surprise since the formula used in calculation was based on the observed force-constants of a large number of alkyls, including these. Although the precise values for the M-C bond lengths obtained from the monomethyls do not agree very well with the sums of the appropriate covalent radii, the agreement is greatly improved by the Schomaker and Stevenson electronegativity correction(28).


TABLE 4.1.—PHYSICAL PROPERTIES OF THE TETRAMETHYLS

Bond energy, ία> kcal

MC bond length (a) Observed, Â (b) Cale. uncorr.(e) (e) Cale. corr.W Force constant^ md/Â (a) Observed (b) Predicted Proton resonance Chemical shift(A) (ppm). 18C coupling, c/s 18C chem. shift('> ppm rei to C6H6

m.p. °C b.p. °C

C

83

1·54±·02<Ο>

1-54 1-54

4-2 41

910 124«)

97-1 —16·5<*>

9.5<fc)

Si

72

1·89±·02<Ο 1-94 1-88

2-9 2-8

1000 119(0

129-0 —102</>

ZVn

Ge

51 (GeEt4)

1·98±·03Μ> 1-99 1-93

2-7 2-7

9-67 126<»

130-4 —88(m>

43 (m)

Sn

54 (Sn Et4)

2·18±·03(Λ> 2-17 212

2-2 2-2

9-95 128<»

137-6 —54<*>

78(n)

(a> Values from Cottrell (Ref. 20). (*> See Ref. 21. <«> See Ref. 22; in methyl-silane(28> the value is 1·8668±·0005Α. (d) Ref. 24. The values for methylgermane(25>, and methylstannane(26) are respectively 1·9453±.0005, 2-143±002Â. (e) Ref. 27· </> Ref. 28. <*> Ref. 29. (Λ> Ref. 30. «> Ref. 31. (» Ref. 32. <*> Ref. 33. «> Ref. 34. <*> Ref. 2. <*> Ref. 3.

This is the only one of the two standard mechanisms for "bond-shortening" (that is, for explaining the difference between observed and calculated bond lengths) which is likely to be relevant here; (/?-+</) π-bonding is less likely to be important than in compounds of silicon, germanium or tin with elements of Groups V, VI or VIT, because there are no lone pairs on the methyl groups to take part in such interaction. On the other hand, a type of hyperconjugation is certainly possible between the electrons of theCH-bonds and the d~ (or perhaps /-) orbitals of the central atom; since the extent of this will depend on the diffuseness of those rf-orbitals, it is likely to differ in the different elements, because orbitals of different principle quantum number will be involved. This type of interaction is likely to affect the 13C and proton resonance chemical shifts appreciably, and the 13C-H couplings; the theory of ^C chemical shifts and of 13C-H couplings is not yet thoroughly understood, and the observed values cannot therefore be interpreted satisfactorily in terms of molecular structure. The chemical shifts of the protons bound to the methyl groups of the tetramethyls of


carbon, silicon, germanium, tin and lead have been used as measures of the electronegativity of the central atom(30); this rests on the assumption that any π-type interaction between the protons and the central atom is likely to be either negligible or constant, an assumption which for the reasons given above is open to question.

A systematic study of the changes in any of these properties with sub stitution is at present possible only for the SiC bond length in methyl-silanes and methylfluorosilanes (see Table 4.2).

TABLE 4.2.—SILICON-CARBON BOND DISTANCES IN A NUMBER OF COMPOUNDS, MEASURED BY MICROWAVE SPECTROSCOPY UNLESS OTHERWISE STATED

Compound j Bond length, Â

<*>(CH8)4Si <*)(CH3)8SiH <c>(CH3),SiH, <d>CH,SiH8

1-89 ±-02 1·868±·002 1·867±·004 1·8668±·0005

Compound Bond length, À

(e>CH3SiH2F ('>CH3SiHF2* WCH^CHSiH,, <i>SiH3CN

1·848±·005 1·833±·002 1-853 1-848

* A re-calculationuio) from the same measurements gives a value of 1-84±01 Â <*> Measured by electron diffraction (Ref. 22). <» See Ref. 35. (c> See Ref. 35,

36. <<*> See Ref. 23. <β> See Ref. 37. c/> See Ref. 38. <*> See Ref. 39. (»> See Ref. 40, 41.

The values obtained from microwave spectra indicate that the carbon-silicon qond length is almost unaffected by the number of methyl groups present. The replacement in monomethylsilane of one or two of the hydrogen atoms bound to silicon by fluorine leads to appreciable shortening of the silicon-carbon bond, but this could be interpreted as a consequence of the greater induced polarity of the bond.

Changes in the MC Bond with Unsaturation Unsaturation at carbon might be expected to have a marked effect

on the length and force-constant of the MC bond, through the change in hybridization at carbon; the force constant for the SiC bond even in so simple a molecule as silyl cyanide is, however, rather difficult to calculate precisely. The lengths of the SiC bonds in this molecule(40,41) and in vinyl-silane have been determined, and are given in Table 4.2; there is a slight shortening of the SiC bond from methylsilane to silyl cyanide, but the shortening is much less than the change in carbon-carbon bond length from ethane to acetonitrile (1-543 Â(42) to 1-458 Â(43)).


If carbon can be assigned covalent radii that depend upon the hybrid character of the a-orbitals(44), then either the carbon-silicon bond in methylsilane is abnormally short, or that in silyl cyanide is abnormally long. The 29SiH and 13CH couplings in methylsilane(4δ) do not support the suggestion that the carbon or the silicon atoms in this molecule use abnormally large proportions of ^-character in the C-Si bond(41)*; the stretching frequency for the C-Si bond in silyl cyanide(4e) is lower than in methylsilane, but the C-C stretching frequency in acetonitrile is also lower than in ethane(4ea).

Any back-coordination has so far escaped detection by physical means, although there is strong evidence (based upon studies of base-strengths(47)

and acid dissociation constants(48)) that it occurs in silicon-aromatic systems. This evidence will be discussed in more detail in the next section; it suggests that the π-orbitals of an aromatic system can become conjugated with the d-orbitals of a silicon atom directly bound to the ring. If this is so, then any such process might be expected to have an appreciable effect upon the ultraviolet absorption spectrum of the aromatic system, and a comparison between the spectra of analogous compounds in which carbon and silicon are bound to an aromatic group ought to reveal appreciable differences. This is not found to be so in practice. When a trimethyl-silyl group replaces hydrogen bound to a benzene ring, there is a shift to lower frequencies of about the same magnitude as that produced by substitution of a i-butyl group(49). The spectra of triphenylmethane and triphenylsilane are very similar(49); moreover, there is no evidence from the UV spectrum of the compound that the rings in tetraphenylsilane interact with one another through the silicon atom<60»61). Although the number of strictly analogous carbon and silicon compounds that has been studied is not large, there is at present little or no evidence from ultraviolet spectroscopy for back-conjugation in silicon-aromatic compounds. There appears to be a similar lack of any evidence from the UV spectra of aromatic compounds of divalent sulphur for the use of J-orbitals in conjugation of the lone pairs of electrons at the sulphur atom with the aromatic systems(52) (but see Ref. 51).

Silicon-containing ketones of the formula R3SiCOR absorb in the

* The 2eSiH coupling constant in methylsilane(45> is 194 c/s, as against 202 c/s in mono-silaoe; the "CH coupling constant is 124 c/s, as against 126 c/sinethane^'.The 29SiH coupling constant in silyl cyanide<"> is 240 c/s; in acetonitrile(4ea) the "CH coupling constant is 136 c/s, so the change in MH coupling from H8MCH, to H,MCN is in the same direction whether M is carbon or silicon.


ultraviolet at much lower frequencies than their carbon analogues(53)

and are strongly coloured. The band affected corresponds to the (/ζ-π*) transition of the carbonyl group, which in a-silyl ketones such as Ph3SiCOPh occurs at a frequency about 5000 cm-1 lower than in the analogous a-carbon ketones (such as Ph3CCOPh). This large shift is in striking contrast to the small effect of silicon-substitution upon the frequencies of the ultraviolet bands of phenyl groups.

In the Oî-π*) transitions of ketones, one electron of the carbonyl oxygen atom is transferred to a π-orbital, so that the cr-acceptor and π-donor properties of the group are greatly increased. Now there is a good deal of evidence, summarized in the next subsection, to show that a triorgano-silyl group can act as a σ-electron donor and as a π-electron acceptor; that is, it can take part in both of the interactions which could stabilize the (η-π*) excited state of a keto-group relative to its ground-state and so lower the energy associated with the (w-π*) transition. Since alkyl groups are weaker σ-electron donors than silyl groups and very weak π-electron-acceptors, the difference in the spectia can be explained in terms of the simultaneous operation of cr-donor and π-acceptor effects of the tri-organosilyl group.

The transition associated with the band of phenyl groups at about 4000 cm-1 is of the (π-π*) type; the small effect of the presence of a silicon atom bound to the ring upon the frequency of this band could be explained if the silicon atom interacted more or less equally with the groimd state and the excited state. Although a-silyl carboxylic acids and silyl cyanides are colourless, the (η-π*) bands in their ultraviolet spectra might be expected to show marked shifts to low frequencies when compared with those of analogous carbon compounds*.

Dipole moment studies have sometimes been cited as showing that back-coordination occurs in silicon compounds, and particularly in silicon aromatic compounds(55'56). The interpretation of the dipole moments of even very simple molecules in terms of molecular structure, however, is hedged about with uncertainty; although the results may be significant when a comparison is made in a series of very similar molecules, it is hard to predict what would be expected to happen to the dipole moments of aromatic compounds when silicon (a much more polarizable atom) replaces carbon. The "group moment" of the trimethylsilyl group varies in a series of ^-substituted benzene derivatives*55*, and becomes less as the electron-releasing power of the other substituent increases; this can certain-

* I am extremely grateful to Dr. L. E. Orgel for the ideas in these two paragraphs.

6


ly be interpreted in terms of a variable π-type interaction between both substituents and the benzene ring, but the interpretation assumes that there is no change in the distribution of electrons in the silicon-carbon σ-bond, and also that the "group moments" of the other substituent are the same in these compounds as in their carbon analogues. Either or both of these assumptions may be wrong, and the evidence from dipole moments can be taken at best only as corroborative.

The Electronic Effect of the MezSi-group

The trimethylsilyl group is apparently electron-releasing as far as σ-bonds are concerned. When insulated from any possible π-donor such as nitrogen or a benzene ring by at least one saturated carbon atom, replacement of tert.-butyl groups by trimethylsilyl groups increases the base strengths of amines(67), weakens carboxylic acids(68), and generally activates aromatic rings to electrophilic substitution<59)*. At the same time, a rigorous comparison with the tert.-butyl group is made difficult by the different steric requirements of the two systems. When, however, a silicon atom is directly bound to an aromatic nucleus, the silyl group behaves either as an electron source or sink; />-trimethylsilylbenzoic acid, for example, is a stronger acid than its tert.-butyl analogue(48), and the /^-substituted trimethylsilyl phenols and anilines are respectively more acidic and less basic than the unsubstituted compounds(60). The Hammett cr-constants for the trimethylsilyl group vary with the nature of the other substituent and the type of reaction both for the meta- and for the para-positions but are positive only for the para-position in certain cases(eo). These properties indicate that the /^-trimethylsilyl group may be electron-withdrawing when the other substituent has π-electrons that can become part of the conjugated system, and that the electron-withdrawing power depends to some extent upon the other substituent present. There appears to be some electron-withdrawal even in trimethylsilylbenzene itself, since the ortho-and para-positions are only weakly activated towards electrophilic substi-tution(61), while if there were no π-type electron-withdrawal the activation of the ring would be expected to be much greater.

Chemical Properties The silicon-carbon bond is surprisingly inert. Its bond energy is not

very high, and is very much less than the energies of the SiO and SiF bo nds

* A trimethylsilylmethyl substituent causes a marked shift to lower frequencies of the ultraviolet bands of a benzene ring(59a).


(see Table 2.5); none the less, bonds from silicon to saturated aliphatic groups persist through a variety of chemical reactions, and are moderately stable to such strong reagents as alkali. The reason for this reluctance to react must involve activation energies, since many of the compounds are thermodynamically unstable to substances with which they do not react at room temperature; it is possible that the apparently large activation energies are connected with the absence of lone pairs of electrons and of low-lying excited states on the carbon atom. If the carbon atom bound to silicon is unsaturated, the silicon-carbon bond is appreciably more reactive.

Eaborn(62) has pointed out that the silicon-carbon bond is chemically rather like the hydrogen-carbon bond. Silicon-carbon bonds may be broken by electrophilic or by nucleophilic attack; the former is more effective in removing alkyl groups from silicon, and the latter in breaking bonds between silicon and aromatic nuclei. In general, substitution at silicon or at carbon makes the SiC bond more easily broken, and the behaviour of ß-substituted alkylsilanes is particularly interesting in this context(63). Many of these react very easily with base,forming an olefine and halide ion:

(CH3)3SiCH2CH(Cl)CH8+OH- = (CH3)3SiOH+CH2CH(CH3)+Cl-

This reaction, which appears to have analogies in the chemical behaviour of a number of other silicon compounds, is discussed in more detail below.

1. Thermal stability—Tetraalkylsilanes are stable up to about 500°, and indeed methyldibutylsilane is said to begin to decompose only at 1000°(64) ; the products of thermal decomposition include such compounds as mono-silane(e6), indicating that CH as well as SiC bonds are broken in the process. Substituted alkylsilanes are much less thermally stable, and in particular those containing ß-halogen atoms are liable to decompose even when distilled at atmospheric pressure(63):

Me3SiCH2CH2Cl == Me3SiCl+CH2CH2

This reaction, which is related to the hydrolytic reactions of ß-halotetra-alkylsilanes, seems to be analogous to the decomposition of silyl trifluoro-methyl sulphide(ee) :

SiH3SCF3 = SiH3F+SCF2

and to the hydrogen-atom migration discussed on p. 32. The mechanism is unknown, although bimolecular intermediates have been postulated, and there is evidence from Raman spectra that there is some intermolecular

6·


interaction in the pure liquids that is absent from dilute solutions(e7)*· A rather similar a-elimination has recently been reported(65):

CFCl2-CF2SiCl3 2^° CFCl:CFCl+FSiCl3

An intramolecular mechanism was proposed for this process. 2. Reactions with elements—Halogens often break the bonds from

silicon to aromatic carbon systems(69), particularly in the presence of aluminium halides; since aluminium halides are also active in this way(70)' there is liable to be some confusion about the species involved. Bonds from silicon to aliphatic carbon atoms are much less readily broken by halogen(71)<

and it is possible to chlorinate alkyl groups bound to silicon with chlorine gas and a suitable catalyst(72). This is remarkable, and illustrates the importance of activation effects, for according to a rough thermodynamic calculation, based upon standard values(20) for the relevant bond energies and ignoring entropy changes, the SiC fission ought to be thermodynamically favoured by at least 10 kcal, (see also p. 91). Silicon-carbon bonds are rather stable to oxygen(73), but they can be broken by alkali metals(74).

3. Reactions with hydrides—Hydrogen halides remove silicon atoms from aromatic rings, particularly in the presence of aluminium halides, and the reaction has been used in the preparation of iodosilane from phenyl-silane(76). Acids in general break silicon-aromatic carbon bonds, and the process, called desilylation, has been extensively studied by Eaborn(76). Lewis acids, such as aluminium halides, are also effective(70), but base reacts much less readily(77).

Alkylsilanes are more stable to hydrogen halides, even in the presence of an aluminium halide catalyst, and monomethylhalosilanes have been prepared by the action of a hydrogen halide, in the presence of the appropriate aluminium halide, upon monomethylsilane at 100O(78). Tetraalkyl-silanes are stable even to 40% alkali at room temperature, but the silicon-carbon bond in hexachloromethylsilane is broken by cold water(79):

Cl3CSiCl3+4H20 = Cl3CH+Si(OH)4+3HCl

The silicon-carbon bond in this molecule is also broken by lithium aluminium hydride at 0°, although dichloromethyltrichlorosilane can be reduced by lithium aluminium hydride to dichloromethylsilane under the same conditions(80).

Substitution in an alkyl group bound to silicon generally makes the Si-C

* See also ref. 67a.


bond more easily broken, but when there is a substitution like oxygen or chlorine ß to the silicon atom, the SiC bond is much less stable than in other monohaloalkylsilanes, and is readily broken by base(e3-81):

R3SiCH2CH2X+OH- = R3SiOH+CH2CH2+X-Although some carbon compounds react similarly, the ß-elimination takes place much more readily in silicon compounds; this has not yet been satisfactorily explained, although there is some evidence (see above) for unusually strong intermolecular interactions in ß-halotetraalkylsilanes(67)*.

4. Reactions with halides—The derivatives of a number of elements such as aluminium(70) or mercury(82) may break silicon-carbon bonds in aliphatic or (more readily) aromatic compounds. As has already been pointed out, Lewis acids such as aluminium halides catalyse the dispro-portionation of a number of silicon compounds(83), and it has been suggested that bridged intermediates with structures similar to those of the dimeric aluminium alky Is may be formed in these reactions(83). Silicon-carbon bonds are much more labile in the presence of aluminium halides(83a).

The bonds from carbon to germanium and tin, while relatively inert, are generally less stable than silicon-carbon bonds, particularly to attack by halogens(2-3). They are also more readily oxidized. Organo-cations of tin are well-characterized; the heavy atoms in the ion Me3Sn+ are probably coplanar, while those in the ion Me n"*-1" are probably in a linear arrangement(84)f. No analogous species of silicon are known; alkylgermyl cations may possibly be formed in solution, but if so they are extensively hydrolysed(85). It thus appears that germanium and tin both form cationic and anionic species more easily than silicon (see also Chapter 2). There is evidence from a study of the dissociation-constants of para-substituted aromatic acids that there is π-bonding between trimethylgermyl or tri-methylstannyl groups and benzene rings(48), while α-germyl ketones show a shift in ultraviolet absorption similar to that observed in the analogous silicon compounds*53*.

2. THE SILICON-SILICON BOND There are only two important methods for the controlled synthesis

of silicon-silicon bonds. Hexaorganodisilanes have been prepared by an

* A "limiting siliconium ion" mechanism has been suggested for the hydrolytic reaction, based on kinetic studies(67a).

t Though some of the evidence relating to the structure of trimethyl tin acetate has been questioned(84a), there seems little doubt that the fluoride contains planar Medications*841».


extension of the Würtz-Fittig reaction, in which triorganohalosilanes are treated with sodium, sodium amalgam or sodium-potassium alloy(8e):

2R3SiX+2Na = R3SiSiR3+2NaX

This method is most useful for making symmetrical hexaorganodisilanes, though compounds such as MegSi-SiMea'SiMe^SiMeg have also been made in this way, and the trisilane MegSi-SiMe^SiMeg has been obtained by the action of sodium-potassium alloy on a mixture of trimethyl-chlorosilane and pentamethylchlorodisilane(8ea) :

2Na+MegSiCl+Me3Si-SiMe2Cl = MegSi-SiMe2-SiMeg+2NaCl

When monohalosilanes are treated with alkali metals or amalgams, it is usually monosilane rather than disilane that is formed(87) together with hydrogen and a solid that is believed to be a polymeric hydride (SiH2)x:

2SiH3Cl+2Na = SiH4+(SiH2)+2NaCl

odosilane, however, reacts with sodium amalgam to give disilane<87a>. The silicon-silicon bond in disilane itself is apparently unstable to sodium or sodium amalgam(87), which may explain why silicon-silicon bonds are not usually formed in this reaction. When monohalosilanes are treated with magnesium, any reaction is like that just described(88). A number of solid polymeric silicon compounds of formulae (SiXg) and (SiX) have, however, been made by the action of electropositive metals on polyhalo-monosilanes(89'90) :

xSiBr4+xMg = (SiBr^+xMgBra

The reaction between sodium and diphenyldichlorosilane gives what appear to be ring compounds, Ph8Si4 and Ph12Si6

(90a). Hexaorganodisilanes have also been prepared by the action of tri

organohalosilanes on triorganosilylmetallic compounds(91):

Me3SiCl+KSiPh3 = Me3SiSiPh3+KCl

The course of this reaction depends on the alkali metal used; with silyllithium compounds a complex mixture of products may be obtained(92), suggesting that there may be some disproportionation of the silicon-silicon and silicon-carbon bonds in the presence of the lithium reagent(93). So far this method has chiefly been used to make hexaorganodisilanes, and with one exception has been restricted to compounds containing silicon bound to aromatic rings. Trisilanes could in principle be prepared in this way:

2Ph3SiK+SiH2Cl2 = Ph3SiSiH2SiPh3+2KCl

SILICON BOUND TO ELEMENTS OF GROUP I V 87

Side-reactions, however, reduce the yields of the higher silanes in such systems(94).

The reactions between silylpotassium and monobromosilane or tetra-bromosilane do not appear to give higher silanes of the expected configurations(94a).

Silicon-silicon bonds are also formed in the pyrolysis of monosilane(95)

and of tetraalkylsilanes(65); in the pyrolysis of disilanes (at rather lower temperatures)(9e); in the addition of monosilanes to unsaturated systems under the influence of irradiation(97); and in the polymerization of the unstable dihalides SiX2

(98). Compounds of formulae such as SieCl14 and Si5Cl12 have been obtained by the action of trimethylamine on hexachloro-disilane(98a), but the higher hydrides have so far been prepared almost exclusively by the hydrolysis of an alkaline earth metal silicide (99~102)

in the presence of a protonic acid. This is an uncontrolled method of preparation, giving monosilane, hydrogen, disilane and varying proportions of higher silanes; the relative amounts of the different products depend very much upon the conditions of hydrolysis. When magnesium suicide is solvolysed with ammonium bromide in liquid ammonia at about —40°, a yield of monosilane (based upon the amount of silicide taken) of well over 50% is obtained(100), together with some disilane but almost no higher silanes. At lower temperatures the yield of monosilane drops considerably, and a good deal of hydrogen is produced. With hydrogen halides in anhydrous hydrazine at temperatures over 50°, yields of monosilane of 50% or over are also obtained, and the addition of water to the solvent reduces the yield; again, little or none of the higher silane fraction is obtained<101). Hydrolysis with an aqueous mineral acid, on the other hand, gives a greater proportion of higher silanes, but the overall yield of silicon hydrides is rather less<99) (though it has recently been improved*102*). The mechanism of the reaction is uncertain. It has been suggested that the diradicai SiH2

is formed first, and that this then polymerizes or reacts with the acid or solvent*; in support of this idea, the compound SiH2(MgOH)2

(103) has been isolated from the products of the reaction between magnesium silicide and alcoholic hydrogen chloride. Moreover, some discrete higher silanes have been obtained when the polymeric (SiH2)x is heated(104). On the other hand, it is difficult*100* to account for yields of monosilane of greater than 50% obtained from the reactions in ammonia and in hydrazine unless the reaction between this intermediate and hydrogen bound to nitrogen pro-

* It is not clear whether the SiHa-species is a true diradicai, a radical-ion (SiH^), a dianion (SiHH or a polymer.


cedes more readily than the polymerization. If this were so, it would also account for the very small amounts of products containing silicon-silicon bonds which were formed in the same solvents. It would have to be supposed that in water the reaction with the hydrogen atoms of the solvent proceded less readily, and so there was more scope for polymerization. Since, however, "magnesium suicide" is not a very definite compound, and may vary in composition from one method of preparing it to another, it is perhaps unwise to draw conclusions about the course of the reaction from the results of different workers using different solvent systems without taking into account the possibly different samples of suicide used.

A few compounds containing silicon-germanium and silicon-tin bonds have been prepared by the reaction between halosilanes and organogermyl-or organostannyl metallic compounds, or from silyl metallic compounds and halo-germanes or stannanes(10510e). Several organodigermanes(107)

and -distannanes(108) have been prepared by similar methods. The hydrolysis of magnesium germanide by acids in solution in water(109), ammonia(110), or hydrazine(111)has been used to make inorganic germanes,and monomers have been characterized containing up to five germanium atoms(109); here also the proportion of higher germanes is much higher when aqueous acids are used. Higher germanes have also been obtained by passing monogermane through an ozonizer(mo). Distannane, Sn2He, has recently been prepared from the action of a mixture of stannite and boronate on aqueous acid at 0°; a much larger amount of monostannane was formed at the same time(u2).

Physical Properties

Some of the physical properties of the hydrides M2H6 of carbon, silicon, germanium and tin are given in Table 4.3, with some of the properties of the M-M bonds as measured in these compounds. There is little that calls for comment. The observed MM bond distances agree with the values calculated from the atomic radii, which is as well, since neither π-bonding nor electronegativity corrections would be appropriate here. The MM force constants also agree reasonably well with those predicted using Siebert's rules(29).

The proton resonance spectrum of trisilane has recently been analysed in detail(120). It is very like that of propane, but there are two important differences. In the first place, the resonance for the silyl (SiH^) protons papears to low field of the silylene (SiH2-) resonance; in propane the methyl protons give resonance to high field of the méthylène protons(121«122).


TABLE 4.3.—PHYSICAL PROPERTIES OF THE HYDRIDES M2H6

C

MM bond energy«*), 1 83 Kcal

MM bond length, Â (a) Observed (b) Calculated«/)

MM force constant (a) Obs., md/Â (b) Predicted«)

m.p. °C b.p. °C

1.543(0) 1-536«) 1-54

4-6W 4-1

—183«) —89

Si

51

2·32±0·03λ«*> 2-34

1·7<α) 1-9

—132-5(fe) - 1 5

Ge Sn

45 1 39 in MeeSna

2·41±·02(β> 2-44

1·3(Μβ,Θβ2)(Λ) 1-8

—109<'> 29

2-80

10(MeeSna)(Ä) 1-2

dec. at room tempera-

ture(m>

(«) Values from Cottrell (Ref. 20). <*» Ref. 42. <c> Ref. 113. «*> Ref. 114; the value in SitCVlw> is 2 .24±06Â. (e> Ref. 116. </> Ref. 27. <*> Ref. 117; values of 1-3 md/À in MeeSi„ from Ref. 118. <*> Ref. 118.* «> Ref. 29. <'> Ref. 33. <*>Ref. 99. <f> Ref. 119. <»> Ref. 112.

* Value of 1-3 md/A in Ge,He(ref. 118a).

Furthermore, there is a chemical shift of about 0-03 ppm between the 29SiH3

protons and the 28SiH8 protons, the latter being the less shielded; the 29SiH2 and 28SiH2 resonances are at almost exactly the same field values. Similar isotopie shifts between protons bound to 12C and 13C have been detected(123), but they are almost always an order of magnitude smaller; the same is true for the simple SiH3X and SiH2X2 compounds.

In the higher members of the series of hydrides, isomerism is possible, and of course has been well-known for years with hydrocarbons. Stock and other early workers suspected that samples of "hexasilane" were in fact mixtures of isomers(124), and fractions corresponding to isomerie tetra-, penta- and héxasilanes have recently been isolated from crude mixtures of higher silanes by gas chromatography(125). The identity of the different isomers cannot be regarded as established beyond doubt, however; the nuclear magnetic resonance spectra, on which identification was largely based, would be extremely complex except in the case of the silicon analogue of neopentane, and analysis would require a thorough study of the 9SiH satellites. A particularly interesting feature of the reaction


between hexachlorodisilane and amines is that the chloropenta- and hexasilanes formed appear to be isomerically homogeneous(98a).

Chemical Properties

The silicon-silicon bond can be broken by electrophilic or by nucleo-philic attack. Several organic disilanes have been shown to react with halogens, forming monohalosilanes(12e):

MegSiSiMea+Ig = 2Me3SiI

On the other hand, disilane itself is unstable in the presence of sodium-potassium amalgams or alloys<87), and (though the silicon-silicon bond in hexamethyldisilane is not attacked by alkali metals) hexaphenyldisilane reacts with sodium-potassium alloy in ether to give triphenylsilylpotas-sium(91):

Ph3SiSiPh3+2K = 2Ph3SiK

In addition to these processes, which may be essentially heterolytic, it seems very probable that silyl (SiH3·) or silylene (-SiH2·) radicals are formed in the pyrolysis of disilanes(96·128). Some of the specific reactions are considered in more detail below.

1. Thermal stability—Both disilane<128) and hexachlorodisilane(96

decompose at temperatures above 450°. When heated in the presence of copper to 350°, hexachlorodisilane is decomposed to silicon and tetra-chlorosilane, but in glass tubes and without the metal polymerization occurs, and chlorosilanes containing three, four, or more silicon atoms are formed. Similarly, higher silanes are formed in the pyrolysis of disilane itself(128). The mechanisms of these processes have not been conclusively established, but it seems very likely that silyl or silylene radicals are involved as intermediates(128). The higher silanes (containing three or more silicon atoms) decompose at lower temperatures*99·124·128) and even at room temperature over long periods, forming (ultimately) hydrogen, monosilane and a solid that consists of silicon and the polymeric (SiH2). The decomposition is probably very susceptible to catalysis by traces of impurity, since divergent accounts of the stability of tetrasilane at room temperature have been given*99·124·129). As the molecular weight increases, the higher silanes also become susceptible to photodecomposition. The polymeric hydride gives some monomeric hydrides when heated(104), while "dibromosilicon", SiBr2, gives hexabromodisilane when heated to about 220O(98).


The organic disilanes are relatively stable to heat. Hexaphenyldisilane, for example, is stable at about 350o(130), in marked contrast to hexaphenylethane, which dissociates at much lower temperatures into triphenyl-methyl radicals(131). The enhanced stability of the silicon compound is to be associated partly with the decrease in steric strain introduced because the silicon-silicon and silicon-aromatic bonds are longer than the corresponding bonds in hexaphenylethane, and partly because of the reluctance of silicon to enter into (p-p)n-bonding (a point which is discussed further in Chapter 1).

2. Reactions with elements—Hexaalkyldisilanes react with bromine or iodine to give trialkylhalosilanes(12e'132):

Me3SiSiMe3+I2 = 2Me3SiI

The reaction occurs at about 50° in the case of the iodide; a molecular mechanism has been suggested. Hexaaryldisilanes react similarly with bromine(133) but much less readily. Disilane itself reacts with iodine at room temperature, but the silicon-silicon bond appears to be unaffected; poly-iododisilanes are the only reaction-products(134). This is strange, since from the thermodynamic data available it appears that the reaction

R3SiSiR3+I2 = 2R3SiI

has a greater driving force than the SiH substitution:

R3SiSiR2H+I2 = R3SiSiR2I+HI

by about 10 kcal (ignoring entropy changes); it therefore seems possible that the SiH bond is weaker in disilane than in monosilane, although a decrease in the bond energy of as much as 10 kcal is unlikely. If substitution by iodine, in preference to breaking the silicon-silicon bond, can be explained as the initial reaction, then the stability of the Si-Si bond in the substituted product to further attack by iodine can be put down either to steric factors, or to some electron-release from iodine to silicon. The Si-Si bond is rather less reactive towards chlorine than might have been anticipated from a consideration of bond energies. The various methylchlorodisilanes, for example, can be chlorinated at carbon without fission of the silicon-silicon bond(132·13δ), although the latter process should be energetically the more favourable by about 40-50 kcal (once more ignoring entropy changes). The stability of the silicon-silicon bond under these conditions increases with the number of chlorine atoms bound to silicon; this can be accounted for in terms of steric hindrance, though elee-


tronic effects (such as (p-*d) π-bonding) may also be involved. Hexa-halodisilanes react with halogens at high temperatures, and this method has been used to prepare mixed halosilanes(135a).

With oxygen, the higher silicon hydrides react very violently, and are completely oxidized to silica and water(99«124) ; hexaorganodisilanes, on the other hand, are relatively inert towards oxygen(130). This must be an activation effect, since oxidation is thermodynamically possible. It is not easy to separate the activities of the SiSi and the SiH bonds in the higher hydrides towards oxygen; the first step in the oxidation may well be attack at one of the SiH bonds. The reactions of silicon-silicon bonds with other elements of Group VI do not appear to have been studied.

Hexaalkyldisilanes are not affected by alkali metals(127), but the silicon-silicon bonds in hexaaryldisilanes(91«93) are broken by lithium, sodium or potassium in the appropriate solvents (such as tetrahydrofuran), and diphenylsilyllithium has been made in this way(13e). Stock, who studied the reaction between disilane and sodium amalgam or sodium-potassium all0y(87) w a s oniy abie to detect decomposition-products such as hydrogen and silicon. There was little reaction with sodium amalgam and almost complete decomposition with sodium-potassium alloy after a month at room temperature. Tri- and tetrasilanes are also decomposed by sodium amalgam at room temperature(129). Silylpotassium however has recently been obtained from disilane and potassium in an ether solvent(94a).

3. Reactions with hydrides—Disilane reacts with hydrogen chloride, bromide and iodide in the presence of the approdate aluminium halide, forming halodisilanes ; monochloro-, bromo- and iododisilanes have been prepared in this way(137-138). There are possibilities of isomerism in the disilanes which contain more than one halogen atom, but the separation of isomers would be extremely difficult since the halides concerned are very liable to disproportionate; the disproportionate does not usually involve the silicon-silicon bonds .The stability of the silicon-silicon bond in the presence of hydrogen halides is a little surprising, since (ignoring, as usual, any entropy changes) rough thermodynamic calculations suggest that the SiSi bond might be unstable to hydrogen chloride, bromide or iodide. It seems probable that the observed stability of the silicon-silicon bond under these conditions is due to an unfavorable activation energy. The SiH substitution reaction occurs more readily between hydrogen iodide and disilane than between hydrogen iodide and monosilane, but it is not clear whether this difference is purely kinetic or whether changes in the bond energies are involved. Alkylchlorodisilanes react with hydrogen chloride


at high temperatures to break the SiSi bond(139); the compound Si5Cl12

reacts with hydrogen chloride at room temperature to give tri- and tetra-chlorosilane(98a).

The silicon-silicon bond in disilane is broken only slowly by water in quartz apparatus(99), but aqueous alkali reacts quantitatively with SiSi bonds to give hydrogen(99):

R3S1SÌR3+2OH- = 2R3SiO-+H2

With solid sodium hydroxide, however, the silicon-silicon bond may persist after all of the silicon-hydrogen bonds have been broken, giving rise to curious hyposilicates which are strong reducing agents and are liable to give hydrogen when treated with water(99). Disilane itself is said not to react with ammonia at room temperature(99), but hexachlorodisilane disproportionates when heated with ammonium salts, or with tertiary amines or their hydrochlorides, forming polysilanes, tetrachlorosilane and trichloro-silane(98a 140·141). In the presence of a trace of trimethylamine, for example, tetrachlorosilane and Si6Cl14 were formed quantitatively at room temper-ature(141). It has been suggested(140) that the reaction proceeds by an ionic mechanism:

ClaSiSiCla+RsN^RaNSiCla+ClaSi

Cl?Si-+Cl3SiSiCl3->Cl3SiSiCl2SiCl3+Cl

Cl- +R3NSiCl3->SiCl4+R3N

Any (p -* (Γ)π bonding from chlorine to silicon might be expected to stabilize the electron-deficient silicon atom produced in the heterolytic first step. This reaction is one more of the large group of amine-catalysed disproportionation reactions of silicon hydrides and their derivatives.

There is little direct evidence for reactions between silicon-silicon bonds and aliphatic hydrocarbons, but when hexamethyldisilane is heated to 600° it rearranges to give a product with two silicon atoms linked by a méthylène bridge(142):

MegSiSiMeg = Me3SiCH2SiMe2H

This reaction is thermodynamically reasonable, and the drastic conditions required suggest that the activation energy must be high.


Since disilane can be prepared from hexachlorodisilane and lithium aluminium hydride(138), it appears that the silicon-silicon bond is relatively stable to attack by this reagent. The yields, however, are not high (about 30-40%); the other products of the reaction are not described, so it is not possible to say whether the low yield is caused by reaction with the ether solvent or by decomposition of the disilane under the influence of the complex hydride. In the latter case, the expected products of decomposition would be monosilane and polymeric (SiH^.

4. Reactions with other compounds—Other reactions of the silicon-silicon bond have been relatively little studied, particularly in inorganic compounds, but one or two points of interest are apparent. In hexaorganodisilanes, the silicon-silicon bonds are surprisingly inert to concentrated sulphuric acid(99), and Stock reports that there is no apparent reaction between disilane itself and the same reagent. It is hard to believe that the system is thermodynamically stable; in fact, although silicon hydrides are sometimes reluctant to react with concentrated sulphuric acid, the reactions once they can be persuaded to begin are usually violent<148). There must be some activation barrier which prevents the reactions between disilanes and this oxidizing agent under fairly vigorous conditions. There is probably some rearrangement of the silicon-silicon bonds in chlorodisilanes in the presence of aluminium halides at high temperatures; chloromethyl-pentamethyldisilane disproportionates in the presence of aluminium chloride at room temperature, forming a chlorosilane in which two silicon atoms are linked by a méthylène group(132):

MeaSiSiMegCHgCl = Me3SiCH2SiMe2Cl

This is yet another of the many disproportionation and rearrangement reactions of silicon compounds that are catalysed by aluminium halid-es(83,144)

Tri- and tetrasilane have been shown to react with chloroform(129)

in the presence of aluminium chloride at room temperature; the reaction-products consist largely of chlorinated silanes, the silicon-silicon bonds being unaffected. Tetrasilane also reacts with bromo- and iodoform and phosphorus tri-iodide at room temperature(129), an aluminium bromide catalyst being necessary in the first case. In none of these reactions is the main step the fission of the silicon-silicon bonds, for the bulk of the products apparently consists of a mixture of halogenated higher silanes; these could not be separated from one another or identified. Since monosilane does not react with chloroform or with phosphorus tri-iodide at


room temperature*99·129*, the difference in reactivity again suggests that the SiH bonds may be weaker in the higher silanes. The difference in the phase of the reactions could also be responsible, though, and a much more thorough investigation is necessary before any definite conclusions on this point can be drawn; from the relevant bond energies, monosilane itself would be expected to react with chloromethanes. Silicon-silicon bonds may also be broken by organometallic(146) and organosilylmetallic(93)

reagents. In the ring compounds Ph8Si4 and Ph12Sie, the ring is broken by mer

cury (II) and tin (IV) chlorides giving the straight-chain phenylpoly-silanes C^SiPh^Cl where n is 4 or 6(90a).

A number of compounds are known that contain silicon-germanium and silicon-tin bonds(105,10e). These are all at least partly organic; one, (Ph3Ge)3SiH, contains hydrogen bound to silicon(147), but most are of the type R3S1.MR3. Relatively little is known about the chemical behaviour of the silicon-metal bonds in these compounds, but they seem to be of the same general order of stability as silicon-silicon bonds in analogous di-silanes.

As mentioned earlier, the higher hydrides of germanium have been characterized as far as Ge5

(109), and evidence has been obtained for the formation in the hydrolytic preparation of compounds containing at least six germanium atoms. In addition to these discrete molecular compounds, a solid hydride (GeH2)x is known(148), analogous to the polymeric (SiH2)x. Hexaorganodigermanes and hexaorganodistannanes react more readily with halogens and with alkali metals than do the analogous silicon compounds, but the reactions are usually of the same general form; digermane reacts with sodium in liquid ammonia to give germylsodium, GeH3Na(149), while hexamethyldistannane reacts with trifluoromethyl iodide, forming trimethyliodostannane and trimethyl-trifluoromethylstannane(15o). The impression left by this discussion of the properties of the M-M bonds is that they are less reactive than might have been expected from their bond energies; once more, the reason for this must be sought in a study of the mechanisms and activation processes of the various reactions.

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(1957); Chem. Abs., 51, 16283 (1957). ( 1 4 5> A. STOCK and P. STIEBELER, Ber., 56, 1087 (1923). ( 1 4 6> H. GILMAN and G . D . LICHTENWALTER, / . Org. Chem., 24, 1588 (1959). (14?) J. G. MILLIGAN and C. A. KRAUS, / . Amer. Chem. Soc, 72, 5297 (1950). ( 1 4 8 > P. ROYEN and R. SCHWARZ, Z. anorg. all. Chem., 215, 295 (1935). (14β) K . M . MACKAY, Thesis, Cambridge, 1960. < 1 5 0 > H. C. CLARK and C. J. WILLIS, / . Amer. Chem. Soc, 82, 1888 (1960).

CHAPTER 5

COMPOUNDS C O N T A I N I N G SILICON BOUND TO ELEMENTS OF GROUP V

1. SILYLAMINES THE compounds considered in this section contain silicon bound to tertiary nitrogen. Adducts of halosilanes and organic amines, in which the nitrogen atoms probably form four single bonds, are dealt with in the last section of Chapter 4; isothiocyanates and other similar compounds in which the nitrogen atom bound to silicon is multiply bound to some other group are discussed in Chapter 7. Silyl hydrazines and silylamides are therefore included in this section although they are not strictly silylamines. A review of silicon-nitrogen compounds has recently been published(1).

Preparation Silylamines are most commonly prepared by the reaction between a

chloro-(la),bromo-(2) or iodosilane(3) and ammonia or some amine that contains hydrogen bound to nitrogen. N-dimethylsilylamine, for example, is made from bromosilane and dimethylamine:

SiH3Br+2HNMe2 = SiH3NMe2+Me2NH2Br

Trisilylamine is prepared from chlorosilane and ammonia:

3SiH3Cl+4NH3 = (SiH3)3N+3NH4Cl

and tetrasilylhydrazine from iodosilane and hydrazine(3) :

4SiH3I+5N2H4 = (SiH3)4N2+4N2H5I

The reaction may at least formally be broken down into steps in which hydrogen halide is first formed and then removed by excess of the starting -amine:

SiH3Cl+NH3 = SiH3NH2+HCl 1 HC1+NH3 = NH^Cl 2 SiH3NH2+SiH3Cl = (SiH3)2NH+HCl 3

101


There is, however, little or ho direct experimental evidence for the formation of monosilylamine (see p. 110). The steps in which hydrogen halide is formed are probably reversible, for the SiN bonds in most silylamines are broken by hydrogen halide (s:e p. 111). Fluorosilanes react in a similar way in the presence of lithium metal(4), which probably removes hydrogen fluoride from the system more efficiently than an excess of amine.

There are important differences between this and the analogous reaction oft haloalkanes with ammonia or amines. Haloalkanes react initially to give quaternary ammonium salts(5); these may be further substituted in the presence of an excess of the haloalkane, but substitution is by no means complete, for the alkylamines are strong bases:

MeI+NH3 = MeNH3I

MeNH8I+MeI = Me2NH2I+HI etc.

The reaction-product therefore consists of a mixture of hydrohalides of the various alkylamines. With one or two exceptions, however, the silylamines are weak bases; they are decomposed by hydrogen halides and do not form quaternary compounds with local excesses of halo-silane(1"3). In the presence of an excess of halosilane, therefore, the NH-groups of the original amine may be completely replaced by silyl groups; trisilylamine, for example, is obtained in 60% yield from the reaction between chlorosilane and ammonia(la). [Tris(triorganosilyl)amines are not formed in the reactions between triorganohalosilanes(e) and ammonia, probably for steric reasons (see p. 103).] N-dimethylsilylamine is the strongest base of the silylamines that contain SiH3- groups(2); it is interesting that this amine could not be prepared from dimethylamine and chlorosilane· This may be because the amine is sufficiently basic to catalyse the dispro-portionation of local excesses of chlorosilane, without forming a sufficiently stable quaternary derivative to remove any chloro- and dichlorosilanes from the system*7·8*.

The reaction is smooth and rapid in the vapour phase or in solvents such as diethyl ether, although in preparing silylamines containing hydrogen bound to silicon it is usually necessary to avoid large local excesses of amine; the reason for this is discussed on p. 109. The mechanism of the process is uncertain. Halosilanes form adducts with tertiary amines, and it seems reasonable to suppose that similar adducts are formed in the course of this reaction; the adducts might then lose hydrogen halide to give the silylamines(2):

SILICON BOUND TO ELEMENTS OF GROUP V 103

HNMe2 MeNH SiH3Br+Me2NH -> H3Si<T -> SiH3NMe2+Me2NH2Br

Br It is not clear whether the second molecule of amine is included in the adduct or not, nor is there any evidence to suggest what the rate-determining step is. A mechanism of this sort implies that any amine which reacts with a halosilane must be an electron-donor, for the intermediate is formed by electron-donation and the hydrogen halide is removed by a molecule of base. Since, as will be made clear, disilylamines are weak bases and do not form hydrochlorides, the formation of trisilylamine from chlorosilane and ammonia is rather surprising; the disilylamine formed in the (formal) second step (equation 2 above) would not be expected to react readily with chlorosilane unless an additional molecule of ammonia were involved. Perhaps the trisilylamine is formed by dis-proportionation of the disilylamine rather than by its reaction with chlorosilane (see p. 109).

Although tetrasilylhydrazine is formed in the reaction between iodo-silane and hydrazine(3), the analogous reaction between hydrazine and triorganohalosilanes gives products which do not contain more than one triorganosilyl group bound to either nitrogen atom(9). A number of compounds of this sort have been prepared, as well as hydrazine derivatives of diphenyldichlorosilane(10) ; one of them is believed to have the structure

Ph Ph

HN X X NH I I

Ph Ph

Silyl hydrazines which contain two triorganosilyl groups bound to a single nitrogen atom have been prepared by the reaction between a triorgano-halosilane and the alkali metal derivative of an incompletely N-substituted silyl hydrazine(11):

R3SiNHNR^R3SiNLiNR^^(R3Si)2NNR2

Tris(trimethylsilyl)amine has been prepared from trimethylchloro-silane and the sodium(12) or lithium(12a) derivative of bis(trimethylsilyl)-


amine, and several other tris(triorganosilyl)amines have been made by these methods:

Me3SiCl+(Me3Si)2NNa = (Me3Si)3N+NaCl

A compound which has been described as disilylcyanamide is obtained when iodosilane reacts with silver (I) cyanamide(13):

2SiHJ+Ag2NCN = (SiH3)2CN2+2AgI

Structural studies, however(13a), show that the molecule is of the carbo-diimide type, SiH3NCNSiH3; the reaction between trimethylchlorosilane and silver (I) cyanamide also gives a disilylcarbodiimide, which has been obtained from the reaction between bis(trimethylsilyl)amine and carbonyl chloride(13b>:

(Me3Si)2NH +COCl2-* [(Me8Si)2N]2CO-> (Me3Si)20 +Me3SiNCNSiMe3

Triorganosilylmetallic compounds react with primary and secondary amines to form the metal hydride and silylamines as ultimate products(14):

Ph3SiLi+R2NH = LiH+Ph3SiNR2

They also react with such substances as azobenzene and benzophenone anil, giving a silyl hydrazine or a silylamine on hydrolysis(15)*:

Ph3SiK+PhN=NPh-5ph3Si-PhNNHPh

Ph3SiK+Ph2C=NPh^Ph2CH-NPh-SiPh3

Where germanium and tin compounds are concerned, the position is rather different. No volatile germylamines have been isolated from the reaction between chlorogermane and ammonia(16), but alkyl and aryl halogermanes react with amines to give triorganogermylamines(17), and tris(triphenylgermyl)amine has been made from ammonia and triphenyl-bromogermane(18). The electron-acceptor properties of tin (IV) lead to the formation of adducts as ultimate products of the reaction between triorganohalostannanes and amines, and discrete stannylamines have not been isolated, though compounds like SnCl3.NH2 are believed to be formed in the ammonolysis of tin (TV) chloride(19).

* This illustrates the stability of some compounds containing Si-N bonds to water unless acid is present.


TABLE 5.1.—PROPERTIES OF THE C-N AND Si-N BONDS

Bond energy, kcal Bond length, Â; (a) Obs. in (MH3)3N (b) Calc.(e> uncorr. (e) Cale, corr.</> Bond angle at N in (MH3)3N Force constant, mD/Â (a) Observed (b) predicted

C-N

73(a)

1·47±·02(θ 1-51 1-47

108±4o ( c )

4-7(s) 4-7

Si-N

?77(Ö>

1·74±·02«*> 1-91 1-81

119-6ìl°(d> ±2 4-1<Ä)

3-3

«*> Ref. 20. <*> Ref. 21. <o Ref. 22. «*> Ref. 23. (e> Ref. 24. ('> Ref. 25. (o) Ref. 26. <*> Ref. 27.

Physical Properties Some of the physical properties of the Si-N bond and of the simple

silylamines are given in Table 5.1. The value for the Si-N bond energy, which was calculated from the vibrational spectrum of bis(trimethylsilyl)-amine, is about what would be expected, but the value cannot be considered reliable; no estimates of even this degree of reliability are available for the Ge-N and Sn-N bond energies. The length of the Si-N bond is appreciably less than the sum of the covalent radii, even when an electronegativity correction has been made, and the force constant is higher than that calculated from Siebert's formula for a single Si-N bond; these two observations have been taken as showing that there is (p -» d) π-bonding between nitrogen and silicon(27). Additional evidence for this comes from the geometry of the bonds from nitrogen(23>28), from the change in proton resonance chemical shifts in the N-methylated silylamines(29), and from a study of the chemical properties of the silylamines in general*.

The angles of the bonds from nitrogen in trisilylamine suggest that the σ-bonds are formed from what are essentially sp2~hybria orbitals. Now nitrogen normally uses ^-hybrids when forming three σ-bonds, the lone pair of electrons occupying the fourth of the roughly-equivalent orbitals. In trimethylamine, for example, the CNC angle is 108°, as against the tetrahedral angle of 109° 24'. If, however, the groups bound to nitrogen have empty orbitals of π-symmetry relative to the σ-bonds, then the

* The Si-N bond moment appears to be very small, affording further evidence for (/?-></)7r-bonding between nitrogen and silicon<28a>.


energy of the lone pair in a pure /^-orbital of nitrogen can be lowered by donor π-bonding, and the nitrogen atom may adopt a planar arrangement of σ-bonds to make this possible (see Chapter 1). This probably happens in N-dimethylaminoboron dichloride, Me2NBCl2(30); the planar arrangement of heavy atoms in trisilylamine can be explained in this way by invoking (p-+d) π-bonding between silicon and nitrogen:

Some slight widening of the SiNSi angle as against the CNC angle might be expected for steric reasons; moreover, since the replacement of methyl groups by the more electronegative fluorine leads(31) to a narrowing of the bond angles at nitrogen from 108° in trimethylamine to 102° 30' in nitrogen trif luoride, it could be argued that some widening of the angles in trisilylamine is to be expected because silicon is more electropositive than carbon. Neither of these mechanisms seems likely to account for the substantial change observed, and it has not been shown that the replacement of carbon by electropositive groups would necessarily increase the bond angles at nitrogen beyond the tetrahedral value. Despite the categorical statement that the skeleton of tris (trimethylsilyl) amine is non-planar(1), there is no reason to believe that the structure of this molecule differs as far as the Si3N system is concerned from that of trisilylamine. In the electron-dififraction study of the latter molecule, Hed-burg(23) concluded that the skeleton might deviate from planarity by about 1°; it has been suggested, largely on the basis of chemical evidence(12), that the skeleton of tris (trimethylsilyl) amine is not quite planar. In the vibrational spectra of both molecules, the symmetrical SiN stretching mode, which should be infrared inactive, appears as a weak band(12-31a) ; this could mean that the skeletons are nonplanar, though other expia-


nations are possible. The weight of evidence leads to the conclusion that the SiNSi angles in both molecules are much the same, and at least within a degree or so of 120°.

TABLE 5.2—PHYSICAL PROPERTIES OF THE N-METHYL-N-SILYLAMINES

m.p.,° b.p.,°

m.p.,° b-p..·

m.p.,° b.p.,°

(SiH3)3N —106(a)

49(d)

(SiH3)2NMe —124<e>

32</>

SiH3NMe2 3(e)

(MeSiH^N —107«» 109<ô)

(MeSiH2)2NMe —■115(0)

80<ö>

MeSiH2NMe2

150«» 4 5( Ô )

(Me3Si)3N 70-71(o

76/12 mm(o

(Me3Si)2NMe

148/740 ram'«'

Me3SiNMe2

—107(*> 86(Ö)

Me3N —117-1^

2·8("

<a> Ref. la. <w Ref. 33. «> Ref. 12. «*> Ref. 34. <·* Ref. 2. ω Ref. 8. <ο> Ref. 35 (Ä) Ref. 36. ci) Ref. 37

The configuration of the heavy atoms in tetrasilylhydrazine has not yet been determined, but the vibrational spectrum of this compound strongly suggests that the molecule has higher symmetry than tetramethyl-hydrazine(28), an observation which is also consistent with (p -> d) π-bonding between silicon and nitrogen. The structures of the N-methylsilyl-amines have not been determined either, and these are of considerable interest. The evidence from nuclear magnetic resonance and chemical properties indicates that there is a progressive increase in the ^-character of the lone pair orbital at the nitrogen atom as two of the carbon atoms in trimethylamine are successively replaced by silicon, but it has not yet been shown that there is any difference between N-methyldisilylamine and trisilylamine in this respect*. If there is such a continuous change, then the bond angles at nitrogen should change in the same sense from about the tetrahedral value in trimethylamine to 120° in trisilylamine.

In the three compounds Me3N, Me2NSiH3 and MeN(SiH3)2, the nuclear resonance of the methyl protons shifts successively to lower fields(29), suggesting that the nitrogen atom becomes more electronegative; this is exactly what would be expected if an increase in the ^-character of the lone pair orbital led to a corresponding increase in the ^-character of the σ-bonding orbitals, for the electronegativity of an orbital increases with its s-character. The change in chemical shift (0-5 ppm) is in fact comparable

* See, however, ref. 44a.


with the change in the chemical shift of the protons of methyl groups bound to saturated (sp*-) and ethylenic (sp2-) carbon atoms(32). In both cases the observations must be interpreted with caution, for the overall symmetry of the molecules changes with the change in hybridization, and the electron-circulations within the molecules are likely to be appreciably affected by any π-bonding. The resonance of the protons bound to silicon in trisilylamine and in the two N-methylsilylamines are much less affected; in N-dimethylsilylamine the resonance is 0-1 ppm to high field of the others, which were both at τ = 5-55 ± 0-02.

There is, therefore, physical evidence which strongly suggests that the lone pair of electrons in silylamines is involved in 0-W)7r-bonding. Additional evidence could probably be obtained from the 14N nuclear quadrupole resonance in these compounds, (although the interpretation of this might be difficult) and from a study of the hydrogen atoms bound to nitrogen in the primary and secondary triorganosilylamines. These should appear to be more acidic than in their carbon analogues, since the σ-orbitals binding them to nitrogen would be more electronegative. Such a study of the OH-groups of silanols has been made, and will be described in the next chapter. Since, however, the lone pair of electrons at nitrogen is made less available to external electron-acceptors because of (p -> d) π-bonding with silicon, hydrogen bonding should be weaker in primary and secondary silylamines than in their carbon analogues.

Among the other physical properties, the most remarkable is the high melting point of N-dimethylsilylamine (see Table 5.2). This compound may also be associated to a very small extent in the vapour phase, but there is no value available for the entropy of vaporization of the liquid. It has been suggested(2) that the apparently large intermolecular interactions may arise through the electron-donor properties of the nitrogen atom; donor-acceptor bonds could be formed from one nitrogen atom to the silicon atom of another molecule:

Mev /SiH3

M e / ^SiHaNMea leading to dimerization or polymerization. If such a process occurs in N-dimethylsilylamine, then the amine should form a loose adduct with trimethylamine ; until this has been experimentally tested, or until the crystal structure has been determined, the mechanism of the interaction will probably remain in doubt. The melting-point of N-dimethyl-methyl-silylamine is about — 150O(33), showing that there is little such interaction


in this compound, and the entropy of vaporization at the boiling point is about that expected for a normal liquid; on the other hand, it is known that silicon-methylation reduces the cr-acceptor properties of silicon towards external donors (see p. 68).

Chemical Properties The silicon-nitrogen bonds in most silylamines are readily broken

by halogen-containing electron acceptors such as hydrogen halides or boron trifluoride; trisilylamines and N-methyldisilylamines behave as weak bases towards such acceptors as diborane or trimethylboron, but some addition-compounds of N-dimethylsilylamines have been prepared. These adducts appear to be less strongly bound together than the analogous compounds of trimethylamine, but there is very little quantitative thermo-dynamic information available about them. Any (p^>d) π-bonding between nitrogen and silicon would be expected to make the lone pair of electrons at nitrogen less available to external acceptors; it appears that the extent of delocalization of the lone pair increases with the number of silicon atoms bound to nitrogen. Some of the reactions of silylamines are discussed in more detail below.

Thermal Stability (a) Tertiary silylamines and tetrasilylhydrazines—Most of these com

pounds are stable to decomposition at room temperature in sealed apparatus for considerable lengths of time(1>3>8), and at temperatures up to 100° for several hours; certain SiN polymers are stable at very much higher temperatures. N-dimethylsilylamine, however, decomposes slowly at room temperature(2). The decomposition-products have not yet been identified, but it seems possible that the main process is one of dispro-portionation:

2Me2NSiH3 = MeN(SiH3)2+Me3N

Even trisilylamine is liable to decompose at room temperature over a period of months, giving monosilane and a solid polymer (see p. 32.)

(b) Primary and secondary silylamines—No compound containing the system SiH-NH has been described which is stable at room temperature. There is evidence that disilylamine, (SiH3)2NH, is formed in the reaction between chlorosilane and an excess of ammonia(1), but it breaks down at room temperature to form monosilane, trisilylamine, a solid polymer and some hydrogen. Stock suggested that two modes of decomposition


were of importance(1) ; the one, a hydrogen migration, has already been discussed (p. 32):

. x(SiH3)2NH = xSiH4+(SiH2NH)x

The other involves SiN disproportionate: 3(SiH3)2NH = 2(SiH3)3N + NH3

The former process is irreversible, and this is why local excesses of primary amine or ammonia should be avoided in the preparation of di- or tri-silylamines containing hydrogen bound to silicon.

There is little direct evidence for the formation of monosilylamine in the reaction between chlorosilane and an excess of ammonia. Methyl-chlorosilane reacts with ammonia in very much the same way, but in this case a little di(methylsilyl)amine has been obtained even from the reaction between ammonia and a small excess of methylchlorosilane(33'36). This secondary amine decomposed at room temperature to give tri-(methylsilyl)amine and ammonia or mono(methylsilyl)amine, but there was no evidence of a decomposition like the former öf those described above.

The reason why these primary and secondary silylamines are unstable is by no means clear. It is possible that the impression of instability is wrong, and that the decomposition-reactions are catalysed by the ammonia or primary amine with which all samples of these compounds so far prepared have been contaminated; many disproportionation-reactions of silyl compounds are base-catalysed. The mechanism of the SiN disproportionate is unknown, although a donor-acceptor complex of the type shown below may be involved:

H H\ ySiH3 i X y N \ / SiH3

H3Si xSiH8 N-methyldisilylamines, however, are weak bases towards such acceptors as boron trimethyl·33·38*, and the ready formation of such a complex by disilylamine would therefore be a little strange. Perhaps the hydrogen atom bound to nitrogen plays an essential part in the formation of the complex, through some sort of hydrogen bonding.

Mono(triethylsilyl)amine is a stable compound at room temperature(39), and bis(triethylsilyl)amine is also known; mono(trimethylsilyl)amine, however, is unknown, and no direct evidence of its formation ÌA the reaction


between ammonia and trimethylchlorosilane has been recorded*e). This variation in relative stability of primary and secondary amines is probably associated with some effect of the size of substituents at silicon upon the formation of an intermediate, for although bis(trimethylsilyl)amine neither disproportionates nor reacts with trimethylchlorosilane even under very vigorous conditions (500°), tris(trimethylsilyl)amine is a stable compound(12).

The reactions between primary or secondary triorganosilylamines and primary or secondary aliphatic amines are at least superficially like the disproportionation reactions of the primary or secondary silylamines themselves. In general, mono(triethylsilyl)amine or bis(trimethylsilyl)-amine reacts with a primary amine RNH2 to give a mixed secondary silylamine as product; ammonia is distilled away(40»41):

Et3SiNH2+RNH2 = Et3SiNHR+NH3

This reaction, which also bears some resemblance to the alcoholysis of Si-N bonds(41a), is strongly catalysed by ammonium salts, and it has been suggested that the catalytic agent is really a hydrogen ion formed by dissociation of the ammonium salt; this could attach itself to the lone pair of electrons at the nitrogen atom from which silicon was to be displaced:

H Et H I I 1+

Et3SiNH2+H++RNH2 = RN->Si-NH = RNHSiEt3+NH3+H+ H Et2 H

The intermediate is very like the one proposed above for the disproportionation of disilylamine. Bis(trimethylsilyl)amine reacts similary with primary and secondary amines, and the reaction is strongly affected by steric factors(4la); it also reacts with tetrachlorosilane, though slowly, giving MesSiNHSiCl8(41b). In view of the catalytic action of the ammonium salts, it seems that mono(trimethylsilyl)amine is most likely to be isolated if it can be removed from the presence of any ammonium salt as soon as it is formed.

Reactions with Hydrides Anhydrous hydrogen halides react with silylamines even at —80° to

break the silicon-nitrogen bond*1·2·8'33*, forming an ammonium salt and a halosilane. This is effectively the reverse of the preparative reaction:

(SiH3)3N+4HCl = 3SiH3Cl+NH4Cl


Presumably part of the driving force is provided by the removal from the system of the amino-reaction product. Many silylamines have been shown to react in this way, and in only two cases has the formation of addition-compounds been reported. N-dimethyl(trichlorosilyl)amine and bis(N-dimethyl)dichlorosilyldiamine both react with excess of hydrogen chloride at low temperatures to form solid addition-products which contain one molecule of hydrogen chloride for each atom of nitrogen present in the silylamine molecule(42):

HCl+Me2NSiCl3= Me2NSiCl3.HCl

The products have been formulated as hydrochlorides, although their structures have not been determined. It is difficult to see why these silylamines should form hydrochlorides while others do not, particularly since N-dimethyl(trimethoxysilyl)amine, which is in some ways electronically similar, is decomposed by hydrogen chloride at room temperature(43). Possibly the so-called hydrochlorides should be regarded rather as addition-compounds of dimethylamine and tetrachlorosilane.

Triorganosilylamines react with water and with alcohols in an analogous way; the reaction between water and bis(trimethylsilyl)amine, for example, has been used to prepare trimethylsilanol(6):

(Me3Si)2NH+2H20 = NH3+2Me3SiOH

The reaction is catalysed by acids, and inhibited by the presence of small concentrations of hydroxyl ion. Trisilylamine, on the other hand, reacts violently with water, giving ammonia, hydrogen and silica(la); this is probably because the ammonia formed in the initial hydrolysis makes the water alkaline, and this leads to hydrolysis of the SiH bonds. Hydrogen sulphide reacts with silylamines to give thiols(44):

Et3SiNH2 +2H2S = Et3SiSH +NH4HS

The analogous reaction with hydrogen selenide might provide a method of preparing "selenols". Mono(triethylsilyl)amine reacts very much more readily than bis(trimethylsilyl)amine with hydrogen sulphide; the latter amine does not react with butylmercaptan, but gives trimethylsilyl butyl sulphide when treated with sodium butylmercaptide(40):

(Me3Si)2NH+2NaSR = 2Me3SiSR+Na2NH

This last reaction, taken with the failure of butyl mercaptan to react,


suggests that nucleophilic attack by sulphur at silicon is a critical step in the reaction-mechanism.

Reactions with Electron Acceptors

(a) Diboraneiu&\—Trisilylamine does not react with diborane at —80°. N-methyl disilylamine reacts at —80° to give an equimolar adduct, which dissociates and decomposes when warmed, while N-dimethylsilylamine gives a similar adduct which decomposes at room temperature into mono-silane and Me2NBH2.

(b) Trimethylboron.—Trisilylamine(38), tri(methylsilyl)amine(33), N-me-thyldisilylamine(33) and N-methyl-di(methylsilyl)amine(33) do not react with trimethylboron at temperatures between —180° and 25°. N-dimethyl-silylamine(2), -methylsilylamine(33) and -trimethylsilylamine(33), however, all form equimolar adducts with trimethylboron at —80°. These adducts are apparently completely dissociated in the vapour phase, and have dissociation pressures approaching an atmosphere at room temperature; from the slope of the dissociation-pressure curve, the heat of dissociation of the N-dimethylsilylamine : trimethylboron adduct was estimated at 8-5 kcal, as against 17-6 kcal for the corresponding compound of trimethylamine( 2).

(c) Trimethylaluminium^^. — Trisilylamine gives an equimolar adduct with trimethylaluminium at —46°; this decomposes at 0°, giving mono-silane. N-methyldisilylamine and N-dimethylsilylamine give similar adducts, the last-named being the most thermally stable; none of these dissociates reversibly over a sufficient range of temperature for a reliable heat of dissociation to be calculated.

(d) TrimethylgalliumU4*\— Trisilylamine does not react with trimethyl -gallium at 0° or —80°; N-methyldisilylamine gives a solid equimolar adduct at 0° which decomposes slowly in the solid state and much faster on melting (at 12°) while N-dimethylsilylamine gives a similar adduct which melts at about 50° and also decomposes slowly as a sob'd but much more quickly in the liquid state.

(e) Boron halides.—Trisilylamine reacts with an equimolar amount of boron trifluoricje at —80° to form a compound which on warming gives

a mixture of boron trifluoride, fluorosilane, and a less volatile material(45). In the presence of excess of boron trifluoride, however, the decomposition into fluorosilane and N-disilylaminoboron difluoride is almost quantitative:

(SiH3)3N+BF3 = (SiH3)2NBF2+SiH3F

8


The mechanism of decomposition of the adduct which is probably formed as an intermediate is uncertain. Tri(methylsilyl)amine reacts in the same way with boron trifluoride<33) ; in this case the formation of an adduct when equimolar amounts of the reactants were mixed at —80° was clearly indicated by the appearance of solid at this temperature, since the melting-points of the products and the reactants are all less than —96°. No solid was observed when an initial molar excess (2:1) of boron trifluoride was taken, which implies that the decomposition of the adduct may involve an additional molecule of boron trifluoride(36).

N-methylated silylamines react similarly to form what are probably equimolar adducts; these then decompose, the N-dimethyl derivatives being appreciably the more stable:

MeN(SiH3)2+BF3 = MeN(SiH3)2BF3 = MeN(SiH3)BF2+SiH3F

Me2NSiH3+BF3 = Me2NSiH3BF3 = Me2NBF2+SiH3F

Of the substituted boron fluorides formed, N-methyl-N-silylaminoboron difluonde loses an additional molecule of fluorosilane to form a polymeric amino-boron fluoride(46).

xMeN(SiH3)BF2 = (MeNBF)x+xSiH3F

N-disilylaminoboron difluoride(45) and N-di methylsilyl)aminoboron di-fluoride(33»36) are relatively stable compounds. The latter reacts with trimethylamine to form a roughly equimolar adduct, indicating that the boron atom is still an electron-acceptor, while the stretching frequencies of the BF bonds are close to those observed in boron trifluoride itself(36)

There is therefore little evidence of interaction between the boron atom and the lone pair of electrons at nitrogen. The reactions between boron trifluoride and a number of other silylamines are generally similar(46)·

The other boron halides react in much the same way with silylamines(38). In a number of cases there is some evidence that an equimolar adduct is formed at low temperatures, but no adduct has been obtained which dissociates reversibly over a range of temperature.

( / ) Halosilanes—Tetrasilylhydrazine does not react with iodosilane at temperatures between —180° and +100°, unlike tetramethylhydrazine (which forms an equimolar adduct)(3). Tri(methylsilyl)amine and N-methyl-di(methylsilyl)amine do not react with methyliodosilane over the same temperature range(33), but N-dimethyl(methylsilyl)amine and N-dimethyl-


(trimethylsilyl)amine both react readily with methyliodosilane at temperatures below 0°, forming solid equimolar adducts(33). The vapour pre sure of the N-dimethyl(methylsilyl)amine adduct at 0° was about 1 cm, and the compound dissociated in the vapour phase into the original reactants:

MeSiH2NMe2+MeSiH2I^(MeSiH2I)-MeSiH2NMe2

The vapour of the adduct formed between methyliodosilane and N-di-methyl(trimethylsilyl)amine, on the other hand, was not homogeneous and its infrared spectrum suggested that it may have contained methyliodosilane, N-dimethyl-methylsilyl-amine and N-dimethyl (trimethylsilyl)-amine. This could be readily explained if the adduct were ionic, for then there would be two possible modes of dissociation, unlikely to differ appreciably in energy:

+ Me2NSiH2Me+Me3SiI I - *

S Me2NSiMe3+MeSiH2I

The presence of trimethyliodosilane in the vapour would be difficult to detect by infrared spectroscopy. If the adduct were not ionic, but contained a five-coordinated silicon atom, then the complex mixture obtained on dissociation is more difficult to explain. It is true that the presence of base usually catalyses disproportionation of bonds from silicon, but methyl-chlorosilane and methyliodosilane do not disproportionate rapidly in the presence of trimethylamine in the vapour phase at room temperature(3e). Radical-exchange reactions like the one given below are well known in silicon chemistry*47»48), and the formation of the two silylamines could be explained on the basis of a process of this sort:

4Me3SiNHPh+SiBr4 = 4MeaSiBr+Si(NHPh)4

These reactions differ from the process in question in that they are not usually rapid at room temperature.

(g) Iodomethane—Iodomethane does not react with tri(methylsilyl)-amine at temperatures between —180° and 100°, but it forms equimolar addition-compounds with N-dimethyl-methylsilylamine and with N-di-methyl-trimethylsilylamine(33'3e). The reaction is slow at room temperature, unlike the reaction between halosilanes and amines (which is usually rapid at temperatures well below 0°); the latter amine reacts appreciably the faster. The structures of the products are not known, but unless their

8*

Me2N' / ' SiH2Me

^SiMe«


formation is accompanied by splitting of a silicon-nitrogen bond, or unless they contain nitrogen-iodine donor-acceptor bonds, they seem likely to be true salts.

Thus, despite the absence of measured thermodynamic quantities, there is qualitative evidence supporting the proposed order of base-strengths for the silylamines:

Si3N < Si2NMe < SiNMe2 < NMe3

It is remarkable how often the stabilities of adducts to irreversible decomposition fit in with this order, implying that the irreversible decomposition is connected with reversible dissociation.

Other reactions: Recent dielectric work suggests that there is rapid exchange between chloro- and aminogroups bound to silicon in dilute solution in benzene(48a>. N-dialkylsilylamines react reversibly with carbon disulphide to give thioamide derivatives*48** :

R2NSiR3+CS2 = R2NC(S)SSiR3

The reaction is reversed by heating the product to 100°. The compounds so far discussed have mostly been monomeric, and

contained only one bond to nitrogen from each silicon atom. Many silyl-diamines(48), silyltriamines(49), and silyl tetramines(48) have been prepared, usually by methods like those discussed above; their chemical properties are also usually very like those of the silylamines themselves. The reaction between N-triphenyl-silyltriamine and hydrogen iodide, for example, has been used to prepare triiodosilane(49):

HSi(NHPh)3+6HI = HSiI3+3PhNH2I

Polymeric silylamines, analogous to the silicones, are well known and are of some commercial importance, being used in lamp-basing cement(50), in the waterproofing of leather(51) and as hypnotics(52); the chemical behaviour of the Si-N bonds in these compounds appears to be little affected by the different degree of molecular complexity. Although the reactions between monomeric silylamines and transition metal ions have been as yet little studied, some interesting reactions involving coordination from the nitrogen atoms of poly silylamines to cobalt (III) atoms have recently been described(53).

Properties ofNH-groups Bound to Silicon

The NH-stretching vibration of liquid bis(trimethylsilyl)amine is at 3380 cm_1,(54) and is at much the same frequency in the vapour phase(55),


so that there is no strong hydrogen bonding in the liquid material. The same compound is inert to sodium in boiling xylene, or in liquid ammonia(6), but it reacts with methyl magnesium bromide to give methane and what is presumably the Grignard-type derivative of the silylamine(6) :

(Me3Si)2NH+MeMgBr = MeH+(Me3Si)2NMgBr

Bis(trimethylsilyl)amine forms a silylaminolithium derivative when treated with lithium phenyl(11), and gives a sodium derivative when treated with sodium and styrene in dioxane as solvent(12) ;NH-groups of partly-substituted silyl hydrazines react in the same way with lithium phenyl(11). The stability of the Si-N bond to organometallic reagents is in contrast to the reactions of silicon-oxygen and silicon-sulphur bonds with such compounds as lithium phenyl(5e).

Little is known about the chemical properties of the germylamines. Some of them are readily hydrolysed by water to give germoxanes, and most of them are relatively unstable at temperatures above that of the room, losing ammonia or an amine and leaving what is presumably an organogermanium-nitrogenpolymer(1718). Bis(triphenylgermyl)amine reacts with hydrogen peroxide to form bis(triphenylgermyl) peroxide(67).

2. SILYL PHOSPHINES AND ARSINES

Very few compounds have been prepared that contain phosphorus or arsenic bound to silicon. The direct reaction between white phosphorus and iodosilane gives a very complex mixture of products, from which silyl-diiodophosphine, SiH3PI2, has been obtained; disilyliodophosphine and (possibly) trisilylphosphine were also formed, but could not be isolated(5859). Silylphosphine has been prepared by heating a mixture of silane and phosphine at low pressure to about 400°(60).Trimethylchlorosilane reacts with solutions of alkali metal derivatives of phosphine dissolved in diethy-leneglycol dimethyl ether to form mono-, bis- and tris(trimethylsilyl)-phosphine(el-4), and analogous reactions have been used to prepare a number of esters containing silicon-phosphorus bonds(65)*. Attempts to prepare triphenylsilylphosphines from triphenylsilyl-metallic compounds and phosphorus halides have not so far succeeded(ee) (though this type of reac-

* Trimethylfluorosilane reacts similarly with phosphylpotassium at — 20o(Ma); trisilylphosphine has been made from chlorosilane and phosphylsodiumie5b>.


tion has been used to make dibenzyl triphenylsilyl phosphonate(66)), but triphenylsilylphosphines have been prepared from triphenylchlorosilane and diphenylphosphyllithium(67):

Ph3SiCl+LiPPh2 = Ph3SiPPh2+LiCl

P-dialkyl(trimethylsilyl)phosphines have been made similarly from tri-methylchlorosilane and dialkylphosphyllithium salts(68).

The only silylarsine containing the SiHs-group that has so far been characterized is silyldiiodoarsine(68), which was obtained from the reaction between arsenic metal and iodosilane. Disilyliodoarsine and trisilylarsine may also have been formed, but they were not isolated(68,M). Evidence was obtained for the formation of trisilylarsine in the reactions between odosilane and potassium or mercuric arsenides, but again the compound was not isolated(59). Tris(trimethylsilyl)arsine has been obtained from the reaction between trimethylfluoroarsine and arsenylpotassium(65a), and (very recently) trisilylarsine from chlorosilane and arsenylsodium(65c).

Physical Properties The boiling points of silylphosphine (12-7°) and tris(trimethylsilyl)-

phospine (243°) are not remarkable; the stretching frequency of the silicon-phosphorus bond in silylphosphine is 454 cm~1.(69) There is little other useful structural or physical information about these compounds.

Chemical Properties

These have been relatively little investigated. The fully-alkylated silyl phosphines are thermally stable at temperatures up to about 250° in the absence of air(e3), though it has been suggested that tris(trimethylsilyl)phos-phine is rather less thermally stable(64); silylphosphine decomposes into a complex mixture of products when heated(60). Silylphosphine reacts with caustic soda to give silane and phosphine, and the silicon-phosphorus bond is broken by hydrogen bromide, no addition-compounds being isolated(60). Tris(trimethylsilyl)phosphine, however, reacts with diborane and pentaborane to give addition-compounds*63»64^ This suggests that the phosphorus atom is an electron-donor, and is in keeping with the earlier observation that a substance which was believed to be trisilylphosphine

* It also forms an adduct with sulphur*wa).


formed a solid addition-compound with iodosilane, from which it could not be separated by fractional distillation(59). Similar difficulty was encountered in the attempts to prepare trisilylarsine from iodosilane; it therefore appears that the lone pair of electrons on the phosphorus and arsenic atoms are available to external acceptors, and that there is less(/?->d)7r-bonding in the silyl-arsenic and -phosphorus compounds than in the silylamines. An alternative possibility is that the addition-compounds contain 5-valent phosphorus or arsenic; this is unlikely, since the systems involved are powerfully reducing. The addition-compounds formed by pho phines and arsines with halosilanes are described in Chapter 4. Tris(trimethyl-stannyl)phosphine was obtained from the reaction between trimethyl-bromostannane and phosphylsodium in diethyl ether at —20°, but its chemical properties do not seem to have been studied(65a).

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Chemistry, Pergamon Press, pp. 52, 58 (1959). <88> E. A. V. EBSWORTH and H. J. EMELÉUS, / . Chem. Soc, 2150 (1958). <84> K. KELLY, U . S. Bureau of Mines Bulletin, N o . 383 (1935). (35> J. R. N O L L , J. L. SPEIER and B. F. D A U B E R T , / . Amer. Chem. Soc, 73, 3867 (1951). (36> E. A. V. EBSWORTH, Thesis, Cambridge, 1957. (a?) J. G. A S T O N , M. L. SAGENKAHN, G. J. S Z A S Z , G. W. MOESSEN and H. F. Z U H R ,

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CHAPTER 6

COMPOUNDS CONTAINING SILICON BOUND TO ELEMENTS OF GROUP VI

1. COMPOUNDS CONTAINING SiO BONDS THE number of compounds known which contain silicon bound to oxygen is vast. The silicones, of considerable commercial importance, are based upon skeletons of alternating silicon and oxygen atoms, and there is almost no end to the number of ways in which units of different sizes and with various groups bound to silicon may be linked together. Besides these polymeric species, numerous monomeric siloxanes and silanols have been prepared; within the class of siloxanes of general formula R3S1OX, the atom X may be hydrogen(1), oxygen(2), sulphur<3), chlorine(4), nitrogen(5), phosphorus(6), arsenic(7), boron(8), aluminium(9), carbon*10·11*, silicon*10·11*, titanium(12), vanadium(5), chromium(13) or almost any element which forms directed bonds to oxygen. There is yet another possible variable in the number of oxy-groups bound to each silicon atom. In a monograph of this nature it is impossible to deal adequately with all of these different groups of compounds; the account that follows will deal chiefly with the properties of disiloxanes, silanols, alkoxysilanes and phenoxysilanes, but compounds of other types will be referred to from time to time. Anyone wishing for more information about the other species should consult one of the several more elaborate reviews of silicon-oxygen compounds*10·11·14*. Formation

Silicon-oxygen bonds are the ultimate products of the oxidation or hydrolysis of almost all silicon compounds. For the controlled formation of SiO bonds, hydrolytic methods are usually employed; some compound containing silicon bound to halogen(15), sulphur(le), nitrogen(17) or hydro-gen(18) is treated with a hydroxy-compound, the general reaction being of the form:

SiX+ROH ^ SiOR+HX

The reaction is reversible, and the presence of acid or base may be helpful both catalytically and in removing HX from the system. Tri-

122

SILICON BOUND TO ELEMENTS OF GROUP VI 123

methylsilanol, for instance, is conveniently prepared by the acid hydrolysis of bis(trimethylsilyl)amine(17):

(Me3Si)2NH+HCl+2H20 = 2Me3SiOH+NH4Cl

while base was used in the preparation of ethoxysilane from the chlorosilane: trimethylamine adduct(19):

SiH3Cl.NMe3+EtOH = SiH3OEt+Me3NHCl

The equimolar iodosilane:trimethylamine adduct reacts similarly with methanol, giving methoxysilane{19a). The various hydrolytic reactions are discussed in the chapters that deal with the properties of the SiX bonds.

When R is hydrogen, condensation may occur, and the product may be a disiloxane rather than a silanol. Whether this happens or not will depend partly upon the conditions of hydrolysis and partly upon the other atoms or groups bound to silicon. The hydrolysis of monohalo-monosilanes, for instance, gives disiloxane and not silanol(20), and silanol itself has not been isolated, though it is presumably formed as an intermediate :

SiH3Cl+H20 = SiH3OH+HCl 2SiH3OH = (SiH3)20+H20

In the hydrolysis of disilyl sulphide, it is possible (though not very likely) that silanol is not formed at all, and that there is a simple exchange reaction :

(SiH3)2S +H20 = (SiH3)20+H2S

The condensation reaction is discussed on p. 127. Among other methods for preparing siloxy-compounds that have

been used with success, silver oxide has been found satisfactory as a reagent for converting chloro-, bromo- and iodosilanes to siloxanes(21). Bis(tri-organosilyl)-peroxides have been made from triorganohalosilanes and sodi umperoxide(2), while some trimethylsilyl compounds react with concentrated sulphuric acid to give bis (trimethylsilyl) sulphate(3). Compound scontaining more than one alkoxy- or aroxy group bound to each silicon atom have usually been prepared from the appropriate hydroxy-compounds and di-, tri- or tetrahalosilanes or compounds containing SiH2— or SiH3= groups, while the polymeric siloxanes are usually made by the hydrolysis of polyhalosilanes(10'11·14,. The conditions of hydrolysis are varied so as to give polymers of differing composition and structure.


Oxy-derivatives of germanium (IV) and tin (IV) are prepared by the hydrolysis of halides, sulphides or amines, just as are the analogous silicon compounds*22"25*. Compounds of the general type R3MOR', (R3M)20, R2M(OR')2 and various polymeric germoxanes and stannoxanes have been described, though less is known about them than about their silicon analogues. Digermoxane itself, (GeH3)20, is perhaps formed when iodo-germane is hydrolysed under controlled conditions(26), but it has not been isolated*26**. Mixed derivatives, such as hexamethyl- stannosiloxane, -germo-siloxane and -plumbosiloxane have been prepared from trimethylsiloxy-lithium and trimethylchloro-germane, -stannane or -plumbane<2eb) :

Me3SiOLi +Me3MCl = Me3SiOMMe3 +LÌC1 (M is Ge, Sn or Pb).

Physical Properties Some of the physical properties of the SiO bond and of disiloxane

are given in Table 6.1, with the corresponding properties of the CO single bond and of dimethyl ether. The SiO bond energy is greater than that of the CO bond, as is the estimated value for the GeO bond energy(36)

in Ge02* ; the observed force constants for the SiO bond in a number of

TABLE 6.1—PHYSICAL PROPERTIES OF THE CO AND SiO SINGLE BONDS MEASURED IN THE COMPOUNDS (MH3)20 UNLESS OTHERWISE STATED, WITH SOME OF THE PHYSICAL

PROPERTIES OF THE COMPOUNDS (MH3)20

Bond energy, kcal(a)

Bond length, Â: (i) Observed (ii) Cale, uncorr.(d> (iii) Cale. corr.(e> Bond angle at O, °

Force constant, md/Â (i) Observed (ii) Predicted

m.p. °C b.p. °C

C

86 ("general")

1-417 1-52 1 43 111·5±1·5(«

5·2</> 5-4 —142·5(Λ) —24·8<ΐ)

O

108 in silica

1·63±·(Η(*>ΐη(Μβ88ί)20 1 91 1-76 130±10(c> in (Me3Si)aOt

4-7 in (Me3Si)20(*> 3-7 —144») —15(*>

(e> Ref. 27. <*» Ref. 28. <*> Ref. 29. (d> Ref. 30. (e> Ref. 31. {f) Ref. 32. (°) Ref. 33. <*> Ref. 34. <*> Ref. 20, 35.

t The angle in (SiH3)20 is probably about 1450(32ft).

* See, however, page 26.


molecules are all higher than the value predicted using Siebert's formula and the SiO bond is shorter than the sum of the relevant atomic radii, even when corrected for the difference in electronegativity between oxygen and silicon. These observations are all consistent with a significant contribution by (p -> ύθπ-interactions to the silicon-oxygen bond* ; there is not enough experimental evidence to show how far this is true of the GeO bond (see also p. 133). Additional evidence to suggest that (p->d)n-bonding between silicon and oxygen is of importance comes from a study of the bond angles at oxygen in disiloxanes, and from the physical properties of the hydrogen atoms of hydroxyl groups bound to silicon.

In dimethyl ether the COC angle is close to the tetrahedral value, and this lends support to the description of the cr-bonds from oxygen in terms of 5/>3-hybridization, the remaining σ-orbitals being occupied by lone pairs. In disiloxane, on the other hand, the SiOSi angle is almost certainly <32a>37)

about 140-150° (though there has been some doubt about this)(3839), an angle which suggests that the cr-bonding orbitals from oxygen in this compound are intermediate between sp and sp2 in character; the lone pairs would then occupy almost pure ^-orbitals. Since the overlap between a π-type J-orbital and a /?-orbital is likely to be greater than between a J-orbital and an ^-hybrid, the change in hybridization at oxygen from sp* in dimethyl ether to sp-sp2 in disiloxane can be explained in terms of (/?->d)7T-bonding between oxygen and silicon; the widening of the angle beyond 120° suggests that both lone pairs must be involved to some extent in such interactions.

If, as has recently been suggested(39a), only one lone pair were involved in π-bonding, then the bond angle should be about 120°, and the oxygen atom should still be a relatively strong electron-donor. The observed angles in most cases of 140-150° imply that both lone pairs are to some extent delocalized; the extent of delocalization depends on several factors, such as delocalization energy, π-bond overlap, changes in σ-bond overlap with changes in .y-character, lone pair-lone pair repulsions, lone pairbond repulsions, and electrostatic effects arising from electron-delocaliza-tion, and the balancing of these will not necessarily lead to complete delocalization of one or both lone pairs. Only in cases where there is a single lone pair (such as trisilylamine or silyl isothiocyanate) does the lone pair appear to become sterically inert.

* A recent discussion of (>-></)Tc-bonding between oxygen and silicon is based very arge ly on arguments from bond lengths<36a).


The nuclear resonance of the hydroxyl proton in trimethylsilanol is to low field of that in tert-butanol(40), which is consistent with a greater degree of ^-character in the σ-bonds from oxygen in the silanol than in the carbinol*. At the same time, there is spectroscopic evidence that hydrogen bonding is stronger in silanols than in alcohols*41·42). This can be explained if the oxygen atom in a silanol forms its two σ-bonds from roughly .sp2-orbitals, the third of these hybrids being occupied by a lone pair of electrons, while the other lone pair (in a -orbital) is involved in π-bonding with the silicon atom. This means that one lone pair is available to electron-acceptors, while the hydrogen atom is more acidic than in analogous alcohols; this state of affairs could well lead to strong hydrogen bonding, and should be contrasted with the bonding-situation in secondary silylamines. In those compounds, there is only one lone pair of electrons, which interacts with two silicon atoms, and so is not available to external acceptors; hence hydrogen bonding would be expected to be weak.

From the commercial point of view, the most important physical properties of the alkylsiloxanes are associated with an unusual freedom of motion about the silicon-oxygen bonds. The reason for this is not clear, but it may be connected with the large atomic polarization in hexa-methyldisiloxane(43), and with the unusual intensity of some SiO bond stretching vibrations in the infrared(44); the skeletal bending frequencies of disiloxanes are low and difficult to observe spectroscopically(32a'37~39· ^). There are rotational and translational motions of the siloxane chains insilicones at low temperatures(45).

Chemical Properties

The silicon-oxygen bond is chemically rather unreactive. Some siloxanes react with halogens and with the organic derivatives of electropositive metals (such as lithium alkyls or Grignard reagents) to form silicon-carbon bonds and metal oxides; many react with non-metallic halides to give halosilanes. It seems likely that most of these last reactions involve donor-acceptor intermediates, but the oxygen atom of a siloxane is a very much weaker electron-donor than the oxygen atom of an aliphatic ether. This is to be expected if both of the lone pairs of electrons at oxygen are involved to some extent in Q?->d)7r-bonding. On this basis, both silanols and alk-

* Unfortunately, it is not clear whether the measurements were corrected to infinite dilution. If not, their significance is doubtful in view of the marked influence of hydrogen bonding upon chemical shift values.


oxysilanes should be better donors than siloxanes, and there is some evidence that this is so(39a).

Some of the reactions of siloxanes are considered in more detail in the subsections that follow, while the properties of the hydroxyl group bound to silicon are discussed in the final subsection.

Stability (i) Disiloxanes and alkoxy- and phenoxysilanes—The SiO bond may be

thermally stable up to temperatures of the order of 1000°, but the thermal stability of siloxanes varies widely with the nature of the rest of the molecule. Many of these compounds, and in particular the polymeric siloxanes, are liable to disproportionate when heated, giving products in which the total number of silicon-oxygen and silicon-carbon bonds is unchanged but the molecular species present have been altered; this disproportionation is strongly catalysed by acids and bases(4e·47). Disiloxane itself is relatively stable at room temperature, but is liable to some SiH disproportionation if kept for long periods of time; this process too is probably acid-catalysed(48).

(ιϊ) Silanols—The stability of the silanol group depends very much upon what other atoms are bound to silicon. Silanol itself is unknown, and its formation has never been conclusively established; no silanol, indeed, has been characterized that contains hydrogen bound to the silicon atom. As with primary and secondary silylamines, any silanol formed in the hydrolysis of a silyl compound decomposes either to form disiloxane and water, or to give monosilane and a solid polymer:

2S1H3OH = (SiH3)20+H20

XS1H3OH = xSiH4+(SiH20)x

In keeping with this, the yield of disiloxane obtained from the hydrolysis of a monohalosilane is usually not greater than 30-40% and some monosilane is usually formed at the same time(20). Yields of over 90%, however, are obtained from the hydrolysis of disilyl sulphide(1β); the difference may arise either because the hydrolysis is not a two-stage process or because the hydrogen sulphide formed as the other product of hydrolysis does not catalyse the second of the decomposition-processes of silanol. Triorganosilanols, however, are well known(1> n ) ; they usually show some tendency to condense to form disiloxanes, but (again as with the primary and secondary amines) this tendency becomes less as the organic groups bound to silicon increase in size. Trimethylsilanol condenses readily;


triethylsilanol, rather less readily(49), while triphenylsilanol condenses with some reluctance(50). The condensation reaction has been studied in some detail; it is catalysed by both base and acid, and it seems that protonic attack at oxygen and nucleophilic attack at silicon are both important in the overall reaction(w). The mechanism may be related to that of the reactions between disiloxanes and hydrogen halides, discussed below.

The disproportionation-reactions of siloxanes, which were mentioned above as strongly catalysed by both acid and base, may proceed through intermediate hydrolysis and subsequent condensation of some of the silicon-oxygen bonds.

Reactions The silicon-oxygen bond is liable to be broken by halogens at room

temperature*20·52), though disiloxane has been chlorinated at —125° to give some hexachlorodisiloxane(20), and hexaalkyldisiloxanes can be chlorinated(52'53) or brominated(54) at carbon with the free halogen. Triphenylsilylpotassium has been prepared by the action^of sodium-potassium alloy on methoxytriphenylsilane(55), showing that the SiO bond in the parent compound is unstable to potassium.

Most siloxanes react with hydrogen halides to give halosilane and a hydroxy-compound, though the equilibrium concentration of all of the halosilanes except fluorosilanes is small·56·57):

\ \ —SiOR+HX ^ —SiX+HOR / /

Where R is an alkyl radical containing an a-phenyl group, however, the products obtained are the haloalkane and a siloxycompound(58):

(MeCHPhO)4Si+4HCl = 4MeCH(Ph)Cl+Si(OH)4

It is possible that the halosilane and alcohol are formed initially, but that the alcohol reacts with excess of hydrogen halide, generating water which hydrolyses the halosilane; if, on the other hand, the initial step in the reaction is attack by hydrogen chloride at the oxygen atom, then the alkyl halide may be formed directly.

Alkoxysilanes(59·60) and phenoxysilanes(61) react with water to give the appropriate silanol, in a process like the reverse of the silanol condensation discussed in the previous subsection. This, too, is catalysed by both acid and base, and it seems probable that the mechanism is similar to that


of the silanol condensation(61). Siloxanes react with alkali to form silanols and the alkali metal silanolate(e2), which is effectively a similar process, as is the reaction between hexamethyldisiloxane and amide ion in ether or liquid ammonia(62):

2(Me3Si)20+2KNH2 = (Me3Si)2NH+2KOSiMe3+NH3

and the reaction between disiloxanes and phenols*62a). (R3Si)20 4- 2ROH = 2R3SiOR' + H20 .

The SiO bond is reduced by the hydrides and complex hydrides of alkali metals in organic solvents(63~5); disiloxanes(66) and alkoxysilanes(67) react with Grignard reagents and other organometallic compounds, the reactions usually being of the form:

\ \ —SiOR+R'M =—SiR'-fROM / /

Unusual reactions, however, leading to the formation of SiH bonds, have been observed(63).

Though Ι,Γ-dimethyldisiloxane does not react with iodomethane at temperatures up to 100O(57), most alkoxysilanes react with acid halides to give the alkyl ester and halosilane:

\ \ —SiOR +AcCl = —SiCl+AcOR / /

Among the effective acid halides are acetyl chloride(68),benzoyl chloride(69), thionyl chloride(70), phosphorus oxychloride(71), and phosphorus oxy-bromide(71). While the reaction between «-butoxytrimethylsilane and phosphorus oxychloride follows the equation given above, a-phenylethoxy-trimethylsilane reacts with phosphorus oxychloride to form the halo-alkane(71):

Me3SiOCH(Ph)Me+POCl3 = Me3SiOPOCl2+MeCH(Ph)Cl This, like the reaction between alkoxysilanes and hydrogen halides, suggests that the C-0 and C-X bond energies in substituted ethanes are effected differently by a-phenyl groups. Disiloxanes react very slowly with tetra-chlorosilane unless a catalyst is present, and the reactions between tetra-alkoxy- or tetraphenoxysilanes and tetrachlorosilane are slow even under vigorous conditions, but trimethylphenoxysilane reacts with tetrachlorosilane at room temperature, forming trimethylchlorosilane(72):

Me3SiOPh+SiCl4 = Me3SiCl+PhOSiCl3

9


Phosphorus tri- and pentahalides react similarly(71). Disiloxane reacts readily with the halides of boron(73) and aluminium(74).

In some cases solids appear when equimolar amounts of the reactants are mixed at low temperatures, indicating that adducts are formed, but these always decompose on warming to room temperature :

(SiH3)20+BF3 = [(SiH3)20->BF3] = SiH3F+SiH3OBF2

Ι,Γ-Dimethyldisiloxane and symmetrical tetramethyldisiloxane react similarly with boron halides(57); in the reaction between boron trifluoride and Ι,Γ-dimethyldisiloxane, methylfluorosilane is] produced after only a few minutes at —130°, showing that any adduct formed must have a very transient existence (57-73). Hexaalkyldisiloxanes react readily with boron halides in much the same way(57-76), and (on heating) with aluminium halides; hexachlorodisiloxane does not react with aluminium halides(78). The siloxyboron halides are monomeric species and are not always very stable; siloxyboron difluoride loses monosilane and fluorosilane at room temperature(73). Dimethylsiloxyaluminium, the material obtained from the reaction between disiloxane and dimethylaluminium bromide, is also rather unstable, and is dimeric in the vapour phase at room temperature(74) :

2(SiH3)20+Al2Me4Br2 = 2SiH3Br+[Me2'Al(OSiH3)]2

The siloxyaluminium halides, (SiH30)AlX2, are solid at room temperature and are also probably polymeric (74), but it is not clear whether the bridging involves the siloxy groups. Hexachlorodisiloxane does not react with boron trifluoride.

Disiloxane does not react with diborane at temperatures up to 25°, but there is evidence for the formation of a weak equimolar adduct with trimethylgaUium at —80°; towards these two acceptors, disiloxane is less basic than dimethyl ether, though perhaps a rather better donor than trisilylamine(26a).

The reactions between alkoxysilanes and boron halides usually give halosilane and alkoxyboron halide(79):

Me3SiOBun+BCl3 = Me3SiCl+BunOBCl2

While sec.butoxytrimethylsilane reacts in this way, secbutoxytrichloro-silane and ^ec.butoxydimethylchlorosilane react to give a siloxyboron dichloride and an alkyl halide(80):

Cl3SiOBusec+BCl3 = Cl3SiOBCl2+BusecCl


This means either that the presence of chlorine bound to silicon in this system affects the activation energies of the two steps in different ways, or that the SiO bond energy is increased when chlorine replaces a methyl group bound to silicon; moreover, this increase must be greater than any possible increase in the SiCl bond energy in dimethyldichlorosilane as against trimethylchlorosilane*. Alkoxysilanes also react with aluminium halides, and the compound (AlCl20)^Si has been obtained from the reaction between aluminium chloride and compounds of formulae (RCOCH=CHO)4Si(81).

Methoxysilane, MeOSiH3, forms no adduct with diborane, and reacts at low temperatures with boron trifluoride to give fluorosilane(19a).

Hexamethylgermanosiloxane, Me3GeOSiMe3, reacts with phosphorus oxychloride at —20° to give 80% of the trimethylchlorogermane and 20% of the trimethylchlorosilane predicted by the two equations below(81a).

-> Me3GeCl + Me3SiOPOCl2

Me3GeOSiMe3 +POCl3

->Me3SiCl + Me3GeOPOCl2

With aluminium chloride, however, only trimethylchlorogermane is produced(81b):

Me3GeOSiMe3+AlCl3 = Me3GeCl+Me3SiOAlCl2

Hexamethyldisiloxane reacts with oxides of elements such as chromium (VI) or selenium(VI), forming the appropriate trimethylsilyl ester; the reaction is in some cases very violent(81c):

(Me3Si)20+Se03 = (Me3SiO)2Se02

Hexamethylgermanosiloxane reacts similarly, but the mixed esters formed are extremely unstable(81d).

Silanols

Triorganosilanols have been prepared by the hydrolysis of halosilanes(1), silylamines(17) and disiloxanes(6e); the preparative method used to prepare a particular compound depends on how liable the silanol is to condense. Since the condensation is catalysed by acid and by base (see p. 127), those

* The facts would also be explained if the SiCl bond energy dropped with increasing chlorine substitution, or if the mechanism were different in the two cases. The former is unlikely, in view of the shortening of the SiCl bond with increasing chlorine substitution (see Chapter III).

9 ·


silanols most liable to condensation (that is, diols and silanols in which the organic groups bound to silicon are relatively small) must be prepared under very mild conditions. Hydrogen bonding and association in silanols has already been mentioned on p. 126. From the chemical point of view, silanols are more acidic than alcohols, though (according to titrations with tetraalkylammonium hydroxides) rather less so than most phenols(82); they react with aqueous caustic soda, for example, to give the alkali metal silanolate<49). Silanols react with acid halides, forming either halosilanes or oxysilanes:

R3SiOH+R'COCl = R3SiCl+R'COOH

RgSiOH+R'COCl = R3SiOCOR'+HCl The relative importance of these two processes depends to some extent on the conditions of reaction(83). Triethylsilanol exchanges O with H2

180 in five hours at room temperature(84). Adducts: A silicon atom to which three or four oxygen atoms are bound

si an electron-acceptor; adduct-salts have been prepared(84a) of certain spiro-siloxanes and alkali metal alkoxides, such as:

C H 2 - 0 O—CH9

Si< y

CH

CH2—O I O—CH2 Li+

and similar adducts have been postulated as intermediates in the exchange reactions between alkoxy-anions and alkoxide groups bound to silicon.

In the cyclic compounds of formulae ZSi(OCHRCHa)3N, where Z can be hydrogen, alkoxy, alkyl or aryl, there is some evidence for intramolecular coordination from nitrogen to silicon(84b).

Oxygen Compounds of Germane and Stannane

Hexaorgano-digermoxanes and -distannoxanes are well known, and behave chemically in much the same way as their silicon analogues. The M-O bonds are, however, rather more easily broken by halogen hydrides, by hydroxy-reagents such as acetic acid, and by other weak acids(8&-88).

(Et3Ge)20+2CH3COOH = 2Et3GeOCOCH3+H20

(R3Sn)20+H2S = 2R3SnSH+H20


Trimethylstannyl hydroxide is basic(89), in contrast to trimethylsilanol; dimethylgermanediol is more acidic than dimethylstannanediol in aqueous solution(90). In the compounds Ph3MOH (where M is silicon, germanium or tin), hydrogen bonding appears to be much less strong for the germanium and tin compounds than when M is silicon; this has been interpreted as showing that (/?-></)7r-bonding is less important between oxygen and germanium or tin than between oxygen and silicon(91).

Hexamethyldigermoxane reacts with boron trifluoride even at —80° to give trimethylfluorogermane, but trimethylmethoxygermane gives a solid equimolar adduct with boron trifluoride that dissociates reversibly over a range of temperature. This implies that the oxygen atom in the methoxygermane is a reasonably strong electron-donor, and so suggests that (/?-></)7r-bonding may be less important in the germanium compound than in alkoxysilanes(184a), but comparison is not easy in view of the absence of values for dissociation energies.

2. COMPOUNDS CONTAINING SILICON-SULPHUR BONDS

Compounds containing silicon bound to sulphur can be prepared by the reactions between halosilanes and metal sulphides(16), or from hydrogen sulphide and halosilanes(92), or silylamines(93); the first of these methods has proved the most useful for making inorganic silyl sulphides. Disilyl sulphide, for example, is prepared by allowing iodosilane vapour to pass over either mercury (II) or silver sulphide at room temperature(16):

2SiH3I+HgS = (SiH3)2S+HgI2

The yield obtained is high. Some organosilyl sulphides have been prepared by the analogous reaction between thio-derivatives of the alkali or alkaline earth metals and chlorosilanes in the presence of some organic solvent(94) :

Me3SiCl+Me3CSNa = Me3SiSCMe3+NaCl

Trimethylchlorosilane reacts with lead(II) mercaptides to give alkyl trimethylsilyl sulphides(94a) :

2Me3SiCl+Pb(SR)2 = 2Me3SiSR+PbCl2

Methods involving hydrogen sulphide have as yet been little used in inorganic silicon chemistry, though chlorosilyl sulphides, such as SiCl3SH, have been prepared by the reaction between hydrogen sulphide and tetra-chlorosilane at temperatures greater than 500O(92). Iodosilane was found not to react appreciably with hydrogen sulphide at room temperature(16);


organohalosilanes react with hydrogen sulphide in the presence of base, presumably because an equilibrium is set up which is displaced by the removal of hydrogen halide(96):

R3SiCl+H2S ^ R3S1SH+HCI

Since iodosilane reacts with most bases, this method has not been used for the preparation of disilyl sulphide*. Another reaction, which is more promising from the inorganic point of view, involves the treatment of a silylamine with hydrogen sulphide*93·94·97*:

Me3SiNHPh + H2S = Me3SiSH + PhNH2

This might well prove to be useful in the preparation of silyl mercaptan.

Physical Properties

Relatively little is known about the physical properties of the silicon-sulphur bond; some of the information that is available is summarized in Table 6.2, together with some of the physical properties of disilyl sulphide and of the analogous carbon systems.

The silicon-sulphur bond is somewhat shorter than the sums of the CO-

TABLE 6.2.—PHYSICAL PROPERTIES OF THE SiS BOND AND OF DISILYL SULPHIDE

Bond energy, kcal(a)

Bond length, À: (a) Observed

(b) Cale, uncorr(e>. (e) Cale, corr('>.

Force constant, md/Â: (a) Observed (b) Predicted(0)

m.p. b.p.

C

65 in Me2S

1·802±·002<θ 1-81 1-81

3·3<*> 3-2 —98-3«) 37·3(ί>

Si

70<ί»

2·14±·02 in Cl3SiSH«*> 2-21 215

2-2 in (Me3Si)2S(>*> 2-2 —70(') 59<'>

(a> Ref. 27. {ι» Pauling(98> gives 54 kcal, and Kriegsmann(") estimates the SiS bond energy from the vibrational spectrum of (Me3Si)2S to be 63 kcal. (c) Ref. 100. «*> Ref. 101. <*> Ref. 30. </> Ref. 31. <*> Ref. 32. <>*> Refs. 33, 99. <*> Ref. 102. <» Ref. 16.

* Silyl methyl sulphide has been prepared from methyl mercaptan and the equimolar iodosilane:trimethylamine adduct(eea).


valent radii of silicon and sulphur, but the shortening can be almost entirely accounted for on the basis of the usual electronegativity correction; this might lead to the conclusion that (p -> d) π-bonding was relatively insignificant in the SiS bond, and this conclusion is to some extent supported both by the correspondence between the observed and the predicted force-constants for the SiS bond, and by what evidence is available as to the bond angles at sulphur in disilyl sulphides. The angle in disilyl sulphide itself, though it has not been measured precisely, is almost certainly less than the corresponding angle in disiloxane(103'104,, while polymeric organo-silyl sulphides have SiSSi angles of the order of 100O(105). There is, on the other hand, quite strong chemical evidence to suggest that the lone pairs of electrons at sulphur are considerably involved in interaction with the silicon atoms(ie'57). This does not necessarily conflict with the narrow bond angles at sulphur, because sulphur can use rf-orbitals to form π-bonds when its σ-bonds are sharply-angled (as in thiophen), and so the bond angles are a less reliable criterion for the extent of (p -> d) π-bonding here than in the compounds of silicon with first-row elements such as oxygen. The argument from bond-lengths carries little weight, since the validity of the correction even in qualitative terms has been questioned(106). The low force-constant has qualitative significance, but the predicted force constants in general cannot be taken as more than rough guides as to what the force constants of single bonds would be. None the less, the weight of the physical evidence taken as a whole suggests that (p -> d) π-bonding is less important in the SiS bond than it is in silylamines and disiloxanes.

There is some additional evidence for interaction between the two silyl groups bound to sulphur in disilyl sulphide(107). The proton resonance spectrum of (28SiH3)2S should consist of a single line, for the two silyl groups are chemically equivalent. If one 28Si atom is replaced by 29Si, the chemical equivalence is preserved but the groups become magnetically non-equivalent; 29Si is present in natural silicon in about 5% abundance, and has a nuclear spin of £, so the nuclear resonance spectrum of disilyl sulphide (like those of dimethyl sulphide, disiloxane and dimethyl ether) consists of a single strong resonance line, with two weak satellites of equal intensity, almost exactly equally spaced on either side of the main peak. The satellites in the carbon compounds are due to protons bound to 13C atoms, and are weaker, since this isotope is present in a natural abundance of only about 1%. In the molecules which contain a single 29Si or 13C atom, coupling is possible between the protons of the two MH3-groups, which are not magnetically equivalent; since the concentrations of the magnetic


isotopes in question are so small, it is most unlikely that a significant proportion of molecules will contain two 29Si (or 13C) atoms, and so the resonance of the protons bound to 29Si in any silyl compound of formula (SiH3)2X might in principle be expected to appear as a quartet. The protons of the two SiH3-groups, however, are separated by four chemical bonds, a distance usually considered to be too great for observable coupling unless special interaction-mechanisms such as π-bonding are involved; none the less, the 29SiH satellites of disilyl sulphide appear as quartets when studied under conditions of very high resolution, with a multiplet separation of 0-70 ± *04c/s. This could be explained very neatly in terms of (p-+ d) π-bonding between silicon and sulphur, were it not for the fact that the 29SiH satellites of disiloxane (where such π-bonding is probably rather stronger than in disilyl sulpkide) appear as sharp, single peaks under conditions which allowed clear resolution of the satellites of the sulphide. It is therefore more reasonable to associate the splitting with d-orbitals of sulphur rather than of silicon, perhaps through a hypercon-jugative interaction with the SiH bonds. It is not clear whether there is any such splitting in the 13CH-satellite resonances of dimethyl ether and dimethyl sulphide; attempts to study these have not so far given conclusive results, partly because of the small natural abundance of 13C.

The physical properties of the SH-group bound to silicon have not yet been very thoroughly investigated.

Chemical Properties

The Si-S-Si system is thermally stable at temperatures up to as much as 300° in bis(triethylsilyl) sulphide(21); disilyl sulphide is stable in sealed apparatus at 70° for several hours, but samples may decompose at room temperature if kept for a few weeks, probably because of the presence of traces of impurity(107). The products of decomposition are monosilane and a colourless solid polymer. Silyl trifluoromethyl sulphide is very liable to decompose in the presence of traces of impurity at room temperature, giving fluorosilane and thiocarbonyl fluoride(108).

The condensation-type reaction of triorganosilyl mercaptans has not been extensively studied, but it appears to take place rather less readily than the analogous silanol condensation(109). Silyl mercaptan, SiH3SH, has been obtained in small amounts from the reaction between disilyl sulphide and hydrogen sulphide; it decomposes at temperatures below 0°, giving disilyl sulphide by what is presumably a condensation reaction(1β»110).

2SiH3SH = (SiH3)2S+HsS


Reactions—Hexaorganodisilyl sulphides are stable in air at temperatures as high as 279°(21), while disilyl sulphide does not react with dry air at room temperature. Disilyl sulphide forms traces of what may be disilyl disulphide, (SiH3)2S2, when treated with sulphur at room temperature, and it reacts with iodine to give iodosilane and sulphur(16110). The silicon-sulphur bond is readily hydrolysed by water; the reaction between hexaorganodisilyl sulphides and water has been used to prepare triorganosilanols, since the hydrogen sulphide produced does not effectively catalyse the condensation of silanols to siloxanes(21):

(R3Si)2S+2H20 = 2R3SiOH+H2S

Disilyl sulphide is almost quantitatively hydrolysed by water to disiloxane. and this reaction affords the most efficient way of preparing the latter compound(16) :

(SiH3)2S+H20 = (SiH3)20+H2S

Hexaorganodisilyl sulphides react with alcohols to give alkoxysilanes*94**· Disilyl sulphide reacts with hydrogen sulphide to form small amounts of silyl mercaptan, in what is presumably an equilibrium*16110) (see above). Trialkylsilyl alkyl sulphides exchange S-alkyl groups with mercaptans under reflux conditions*94*) :

R3SiSR'+R"SH = R3SiSR"+R'SH.

The product depends on the relative volatilities of the components. Disilyl sulphides react with hydrogen halides to give halosilanes and hydrogen sulphide; the reaction between disilyl sulphide itself and hydrogen iodide is quantitative at room temperature*16110).

(SiH3)2S+2HI = 2SiH3I+H2S

The silicon-sulphur bond is broken by lithium hydride in ether solution, but disilyl sulphide reacts only slowly with lithium aluminium hydride in w-amyl ether solvent to give monosilane(16'21). The silicon-sulphur bonds in hexaorganodisilyl sulphides are broken by oxides of nitrogen(m); contrary to an earlier report(112), disilyl sulphoxides are not formed in this reaction.

Disilyl and di(methylsilyl) sulphides are both relatively weak electron donors. Neither compound reacts with iodomethane* or iodosilane at

* Trimethyliodosilane has been obtained from the reaction between trimethylsilyl butyl sulphide and iodobutane under reflux<e4a):

BuSSiMe,+BuI = Bu,S+Me,SiI


temperatures up to 100°, and di(methylsilyl) sulphide does not react with trimethylboron over the same temperature-range(ie«57). No addition-compounds are formed with boron trifluoride at temperatures up to 20°, but di(methylsilyl) sulphide decomposes when heated with boron trifluoride or trichloride(57). In neither case has any evidence been obtained to suggest that an adduct is formed. Both disilyl sulphide and silyl tri-fluoromethyl sulphide react with trimethylamine(16·108110), forming silane (or fluorosilane) and solids that have been regarded as polymeric adducts:

x(SiH3)2S+xNMe3 = xSiH4+(SiH2SNMe3)x

xSiH3SCF3+2xNMe3 = xSiH3F.NMe3+(Me3NSCF2)x

Both of these reactions are examples of base-induced rearrangements. Disilyl sulphides react with some compounds of silver (I) and mercury (II), and the reactions can usually be predicted by reference to the appro* priate conversion series(21) (seep. 65); the reactions are reversible and are of the form:

(SiH3)2S+HgX2 ^ 2SiH3X+HgS

The conversion series do not always apply to substituted silyl sulphides such as silyl trifluoromethyl sulphide, nor do the reactions of silver and mercuric salts always follow the same pattern. It must moreover be borne in mind that the conversion series have usually been worked out from systems in which one component is removed from an equilibrium mixture by distillation or by treatment with further quantities of the second react-ant; thus iodosilane when distilled over solid mercuric sulphide is almost quantitatively converted into disilyl sulphide, but disilyl sulphide reacts with mercuric iodide in a closed system to give about 8% of the amount of iodosilane calculated assuming complete reaction:

(SiH3)2S+HgI2 *=? 2SiH3I+HgS

The conversion series cannot therefore be used to predict the equilibrium concentrations when two reactants are mixed in roughly equimolar quantities.

The properties of silyl mercaptans have been relatively little studied. The isolation and instability of silyl mercaptan, SiH3SH, is particularly interesting; since silylamine, disilylamine and silanol are all so unstable as to have eluded attempts to isolate them, while silyl phosphine is a stable compound at room temperature, and since (p -> d) π-bonding is strong in silylamines and siloxanes, probably weak in silyl phosphines (see p. 118)


and of moderate strength in silyl sulphides, it is tempting to associate the stability of the system SiHMH with lack of π-bonding between silicon and M. The difficulty of isolating mono- and disilylamines may, however, be connected with the relative volatilities of these compounds and their parent amines (see p. 109), and not be a true reflection of their stabilities with respect to disproportionation.

Hexaorganodigermyl(114) and hexaorganodistannyl(115) sulphides have been prepared, but their chemical properties have not been extensively investigated. Digermyl sulphide has been made from iodogermane and mercuric sulphide(26).

3. SILYL SELENIDES

Disilyl selenide is prepared by the reaction between silver selenide and iodosilane(16):

2SiH3I+Ag2Se = 2AgI+(SiH3)2Se

The compound boils at 85°, and decomposes slowly in the presence of traces of impurity at room temperature, forming monosilane and a polymeric solid.

Bis(trimethylsilyl) selenide, and similar derivatives of germanium and tin(l le), have been prepared by the reactions between the appropriate triorganochlorosilane, -germane or -stannane and sodium selenide in dry benzene:

2Me3SiCl+Na2Se = (Me3Si)2Se+2NaCl

Almost none of the fundamental properties of the silicon-selenium bond have been determined, but the vibrational spectrum of disilyl selenide shows that the skeleton is bent(103). The mean Si-Se stretching frequency in this molecule is about 390 cm- 1 . The compound is hydrolysed by water to disiloxane and hydrogen selenide, and reacts with hydrogen iodide to give iodosilane and hydrogen selenide(16). The silicon-selenium bond is broken by iodine(le):

2(SiH3)2Se+I2 = 2SiH3I+Se

Disilyl selenide does not react with iodomethane at temperatures up to 100°, showing that the selenium atom in this compound is not a strong electron-donor(16).

No volatile compounds containing silicon bound to tellurium have been prepared. Iodosilane reacts neither with tellurium metal nor with silver telluride at temperatures up to 100Ο(16·110).


R E F E R E N C E S

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84, 959 (1952); Chem. Abs., 47, 3228 (1953). (77> A . H . C O W L E Y , F . FATRBROTHER and N . S C O T T , J. Chem. Soc, 717 (1959). (78> N . F . O R L O V , Dokl. Akad. Nauk SSSR. 114, 1033 (1957). Chem. Abs., 52 , 2742

(1958). ( 7 8 a ) H . G R O S S E - R U Y K E N , Angew. Chem., 66, 754 (1954). (7S>) W . G E R R A R D a n d J. A . STRICKSON, Chem. Ind., 860 (1958). (80> M . J. F R A Z E R , W . G E R R A R D and J. A . STRICKSON. / . Chem. Soc, 4701 (1960). (81> Y u . K . Y U R ' E V and G . B. E L Y A K O V , Zh. Obsch. Khim., 27, 176 (1957); Chem. Abs.,

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H . SCHMIDBAUR, I . RUTDISCH and P . B O R N M A N , Angew. Chem., 7 3 , 408 (1961). ( 8 1 d ) M . S C H M I D T and H . SCHMIDBAUR, Ber., 94 , 2137 (1961). <82> R . W E S T and R . H . B A N E Y , J. Inorg. Nucl. Chem., 7, 297 (1958). (83> N . S. N A M E T K I N , A. V. TOPCHTEV a n d F . F . M A C H U S , Dokl. Akad. Nauk SSSR, 87 ,

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SILICON BOUND TO ELEMENTS OF GROUP V I 1 4 3

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CHAPTER 7

OTHER SILICON COMPOUNDS

1. PSEUDOHALOGEN DERIVATIVES

MANY compounds have been prepared in which silicon is bound to pseudo-halogen groups such as cyanide(1>2), isocyanate<3) and isothiocyanate(4); analogous compounds of carbon(5), germanium(e) and tin<7) are also known. The silicon compounds are usually prepared by the action of a halosilane on the pseudohalide of some heavy metal like mercury, silver or lead; silyl isothiocyanate, for instance, is made from iodosilane and silver thiocyanate(8):

SiHgl+AgNCS = SiH3NCS+AgI

The halosilane must be chosen with reference to the conversion series that are described on p. 65 since the above reaction is an equilibrium and the extent of reaction will depend on the relevant bond- and lattice energies. Another reaction that has been used to prepare organosilyl pseudohalogen compounds is the radical-exchange between a halosilane and the pseudohalide of some other non-metal·9»10» η ) ; this is like the reactions between the halosilanes and non-metallic halides, and its preparative usefulness depends on the fact that an equilibrium is set up from which the most volatile component can be removed by distillation, thus driving the reaction to completion if the most volatile component is one of the desired products :

«C12H25SiCl+PO(NCS)3 ^ POCl3+wC12H25Si(NCS)3

Silyl pseudohalides can also be prepared by the action of pseudohalogen hydrides on silylamines; the reaction between bis(trimethylsilyl)amine and hydrogen cyanide, for example, gives trimethylsilyl cyanide(11):

(Me3Si)2NH+3HCN = 2Me3SiCN+NH4CN

This method has so far been relatively little used. Perhaps the most remarkable method by which a pseudohalide of silicon has been prepared is the reaction between urea nd trimethylchlorosilane at 300O(12):

Me3SiCl +CO 2 = MesSiNCO +NH4C1

144

OTHER SILICON COMPOUNDS 145

Pseudohalides of germanium and tin may be prepared by most of these methods*6·7· 10> n ) .

Some of the physical and structural properties of the simplest pseudohalides of silicon are given in Table 7.1, with those of their methyl analogues. The question of atomic arrangement in the silicon compounds will be considered first. "Normal" and "iso" cyanides and thio-cyanates of aliphatic groups are well known, and have distinct physical and chemical properties; so far no such isomerism has been shown to occur in silicon compounds, though the preparation of two isomerie silicon tetracyanates has been reported(3)*. The reaction between trimethyl-iodosilane and silver cyanide(20) gives the same product as the reaction between bis(trimethylsilyl)amine and hydrogen cyanide(11), though methyl

TABLE 7.1.—STRUCTURAL AND PHYSICAL PROPERTIES OF SOME SIMPLE PSEUDOHALOGEN DERIVATIVES OF SILANE AND METHANE

m.p°., b.p°.,

MC distance, Â CN distance, Â

m.p°., b.p°.,

MN distance, Â NC distance, À MNC angle, (degrees)

MeCN

_4!<a) 81-8(a> 1458(c) M57 ( o

MeNCS

35<*> 119 (e>

VArVn 1-22W 142«/)

SiH3CN

3 2(ò)

49.6(0) m

1·848«*> 1156(ii)

SiH3NCS

—51-8<« 84«»

1·73±·02(*> 1·20±·02<*> 180(*>

<«> Ref. 13. (ö) Ref. 8. (<» Ref. 14. «*> Refs. 15, 16. (e> Ref. 17. ('> Ref. 18 (0) Ref. 19.

isocyanide is obtained from the reaction between iodomethane and silver cyanide; it is therefore of interest to determine whether the silicon compounds are "normal" or "iso" in form. The microwave spectrum of silyl isothiocyanate(19) shows that the material studied consists] at least mainly

* The so-called "normal" cyanate was only obtained in ca. 5% yield; its infrared spectrum suggests that it is most probably a hydrolysis product, a disiloxyisocyanate(19a).

10


of a compound of structure SiH3NCS; there is no physical evidence to suggest that the material is not homogeneous, or that it contains any of the "normal" compound, SiH3SCN, and so it has been described here and elsewhere as the isothiocyanate. The structure of silyl cyanide has also been determined by microwave spectroscopy(15'16), and the results indicate a "normal" rather than an "iso" arrangement in this case; again, nothing was observed to suggest the presence of any of the other isomer and silyl cyanide behaves as a single, homogeneous compound. It has been suggested(11), however, that triorganosilyl cyanides consist of equilibrium mixtures of "normal" and "iso" cyanides; trimethylsilyl cyanide, for example, is said to consist of 90% of the normal and 10% of the isoderivative at room temperature, the proportion of the latter increasing as the temperature is raised. It is proposed that the two forms are in labile equilibrium(11), the intermediate being of the same sort as that suggested for the disproportionation of the fluorosilanes(21):

/ C N x 2R3SiCN ^ R3Si<; >SiR3 ^ 2R3SiNC

X N C / The halogen atoms in halosilanes are not labile at room temperature (see Chapter III), but in view of the exchange reactions that have been studied at rather higher temperatures(9_11) the proposed cyanide-isocyanide exchange does not seem unreasonable; the physical evidence on which the suggestion is based is not, however, very strong. There are bands in the infrared spectra of trimethylsilyl and triethylsilyl cyanides at 2100 cm-1, in addition to the much stronger band at 2190 cm-1; the latter bands have been assigned to the CN stretching modes of the "normal" cyanides, and the former (by analogy with some organic isocyanides) to the analogous modes of the isocyanides. The band at the lower frequency increases in intensity relative to the higher-frequency band in the spectrum of triethylsilyl cyanide as the temperature rises, and this has been interpreted as showing that the proportion of isocyanide in the mixture increases with increasing temperature; the lower-frequency band is apparently missing from the spectra of silyl cyanide(21a), triphenylsilyl cyanide(11) and methyl-silyl cyanide(22). This is not conclusive evidence for the presence in trialkylsilyl cyanides of molecular species. The additional bands at the lower frequency could be due to some overtone or combination mode (possibly intensified by Fermi resonance with the CN stretching fundamental), though this would not be expected to increase in intensity relative to the fundamental with increasing temperature; the absence of the band


from the spectrum of triphenylsilyl cyanide would be consistent with its assignment to some combination-mode involving the CH3- or CH2- groups. There is no other physical evidence that supports the proposed isomer-equilibrium process. The nuclear resonance of the protons bound to silicon in methylsilyl cyanide consists of a sharp quartet which is neither broadened nor shifted by more than 0-02 ppm by dilution from 95% to about 10% solution in cyclohexane(23); if there were significant exchange of the kind suggested, the shape and position of the nuclear resonance would probably be affected by dilution. On the other hand, the "isocyanide" band is missing from the infrared spectrum of methylsilyl cyanide, so it is not possible to draw any conclusions from the NMR results about the structures of trialkylsilyl cyanides. The molar refractivity^of triphenylsilyl cyanide suggests that the compound has the iso-structure(24), in contrast to the infrared work described above.

Infrared evidence indicates that silyl derivatives of the cyanate group contain SiN rather than SiO bonds(19a>25); this had previously been deduced from measurements of boiling points and refractive indices(3-2e).

Turning to other structural properties of the silyl pseudohalides, the SiC bond length in silyl cyanide is only a little less than in methylsilane; the point is discussed on p. 79. The linear arrangement of the heavy atoms in silyl isothiocyanate contrasts with the bond angle of 140° at nitrogen in the analogous methyl derivative, and can be explained in terms of (p->d) π-bond-ing between nitrogen and silicon*. The vibrational spectrum of silicon tetraisocyanate suggests that the molecule has tetrahedral symmetry, which implies that the Si-N-C-O groups are linear(19»25) ; this, besides being consistent with the structure of the isothiocyanate, affords indirect evidence for the "iso" -structure of the cyanate, since while several compounds are known in which nitrogen forms bonds at 180° to one another, there are very few linear R-O-R systems, and in none of those so far characterized is silicon bound to oxygenf. Germanium tetraisocyanate, it appears, has a different structure.

The entropy of vaporization of silyl isothiocyanate(8»27) implies that the compound is associated in the liquid phase; there is no evidence from

* The vibrational spectrum of silicon tetraisothiocyanate suggests that the NCS-groups are not quite linear; since the material was studied in the solid phase, however» the results are not strictly comparable with those for the SiH3-compound. The vibrational spectrum of trichlorosilyl isothiocyanate suggests that in the liquid phase the SiNCS skeleton is bent(2ea>.

t Some silicates have very recently been described which contain a linear Si-O-Si system(34a>.

10·


vapour density measurements to indicate association of the vapour, but only measurements at low pressures (~5cm) have been made. The value for the entropy of vaporization is a little uncertain, since there was some decomposition during the determination of the vapour pressures of the compound(8). The proton resonance shifts from τ = 5·43 ±·05 ppm in 95% solution in cyclohexane to 5·54 ±·01 ppm in 5% solution in the same solvent, implying that the intermolecular forces in silyl isothiocyanate are rather stronger than in most other silyl compounds(28).

The structures of trimethylgermyl and trimethylstannyl cyanides have not yet been determined. On the strength of infrared and chemical evidence it has been suggested that the germanium compound consists of an equilibrium mixture of cyanide and isocyanide<7); trimethylstannyl cyanide melts at 183° (as against a melting-point of 38° for Me3GeCN), and is probably rather more polar than its silicon and germanium analogues. Germanium derivatives of thiocyanogen have been described as isothio-cyanates, but the evidence on which this is based is not strong.

Chemical Properties

Stability to heat—Cyanides, isocyanates and isothiocyanates of silicon resemble the corresponding halides in thermal stability (see Chapter 3), unless the molecule in question contains hydrogen bound to silicon. Silyl cyanide is stable in sealed apparatus at room temperature for several weeks,but over longer periods it is liable to decompose, forming solid ma-terial(8»27). The decomposition is catalysed by traces of mercury; methylsilyl cyanide may be rather less stable(22). Silyl isothiocyanate decomposes slowly at room temperature in sealed apparatus, being about 50% decomposed after ten days; the decomposition-products include hydrogen, monosilane, disilane and a barely volatile liquid that has not been identified. Once more, methylsilyl isothiocyanate may be rather less stable(22). Attempts to prepare silyl isocyanate from iodosilane and silver cyanate were unsuccessful, the main products obtained being silicon tetraisocyanate and hydrogen(8»27); phenylsilyl isocyanate, PhSiH2NCO, has been prepared(28), however, so the system -SiH.NCO is not intrinsically unstable. It is interesting that a phenyl group bound to silicon apparently stabilizes the compound, in contrast to the effect of silicon-methylation (see above). The difficulty in obtaining silyl isocyanate may have been caused by its ready disproportionation to monosilane and silicon tetraisocyanate, but this is not very likely, since in that case about three times as much monosilane as tetraisocyanate should have been formed, and this was not ob-


served. On the other hand, the products of the reaction strongly suggest that silyl isocyanate must have at least a transitory existence; perhaps its decomposition is catalysed by the presence of excess of the silver salt. The reaction between iodosilane and silver selenocyanate gave only mono-silane, ammonia, and traces of unidentified material that contained no hydrogen bound to silicon(29) ; it seems probable that silyl selenocyanate was formed, but that this decomposed, perhaps again because of the presence of excess of the silver salt*.

The chemical properties of the silyl pseudohalides are in many ways like those of the halosilanes. Triorganosilyl cyanides react with lithium hydride or lithium aluminium hydride to form silicon hydrides(30) :

4RsSiCN+LiAlH4 = 4R3SiH+LiCN+Al(CN)3

They, and the other pseudohalides, are usually hydrolysed by water, forming a disiloxane or a silanol and the pseudohalogen hydride(30):

2R3SiCN+H20 = (R3Si)20+2HCN

They react with the salts of heavy metals in ways which can usually be predicted by reference to the appropriate conversion series; triethylsilyl cyanide, for example, reacts with mercuric oxide to form mercuric cyanide and hexaethyldisiloxane(20'30) :

2Et3SiCN+HgO = (Et3Si)20+Hg(CN)2

They exchange halogen or pseudohalogen groups with the halides of other non-metals (9»10'11), as described in the section dealing with preparative methods; triorganosilyl cyanides react with Grignard reagents to form tetraorganosilanes(30), and are likely to react in much the same way with other organometallic reagents. Silyl isocyanates react with alcohols to give alkoxysilanes(12).

R3SiNCO+MeOH = R3SiOMe+HNCO

This is in contrast to the reaction between silicon tetraisocyanate and secondary amines(31):

Si(NCO)4+4Ph2NH = Si(NHCONPh2)4

Silyl cyanides react with diborane(32), and with boron halides(33), at low temperatures to form equimolar adducts:

2SiH3CN+B2H6 = 2SiH3CNBH3

* Silyl isocyanate and isoselenocyanate have recently been prepared*83).


Me3SiCN+BF3 = Me3SiCN-BF3

These adducts decompose when wanned, giving a silane (or halosilane) and a substituted boron cyanide, which is often polymeric and which may itself decompose:

SiH3CN-BH3 = SiH4+[BH2CN]

Me3SiCN-BF3 = MesSiF+tBFsCN]

There is at present no evidence to show whether the adducts contain cyanide or isocyanide groups coordinated to boron.

Some of the reactions of silyl cyanides are relevant to the question of structure that was discussed'above from the point of view of the physical properties of the compounds concerned.Whentrialkyl-(30), and triarylsilyl(11)

cyanides are heated with sulphur, the corresponding triorganosilyl iso-thiocyanates are formed:

MegSiCN+S = Me3SiNCS

Since the product contains silicon bound to nitrogen, the simplest explanation of the reaction is that it occurs through attack of sulphur on an isocyanide:

R3SiNC+S = R3SiNCS

Such a reaction with a "normal" cyanide would have to involve some process like that given below:

R3SiCN+S -> R3Si S -> R3SiNCS

II N+

This is by no means impossible, but is much less probable than the other mechanism. If the system consisted of an equilibrium mixture of the "normal" and "iso" cyanides, and if (as seems likely) the former reacted slowly with sulphur (if at all), the complete conversion of the whole system to isothiocyanate would not require the presence of a large equilibrium concentration of the iso-compound. On the other hand, the systems were usually heated to at least 100° before the reaction-products were examined. Silyl cyanide does not react with sulphur at room temperature and a complex reaction occurs when the two substances are heated together(21a).

The two possible cyanide isomers would be expected to differ in their


reduction-products, the normal cyanide giving a primary aliphatic amine and the iso-compound giving a secondary silylamine:

R3SiCN->R3SiCH2NH2; R3SiNC->R3SiNHCH3

Trimethylsilyl cyanide is reduced by hydrogen and raney nickel at 140° under high pressure to bis(trimethylsilyl)-amine, a reaction which is more in keeping with the isocyanide structure(31) ; the conditions, however, are so vigorous that a normal cyanide could have reacted in this way, rearranging in the course of reaction. The only other chemical evidence that isocyanides are present comes from the reactions between trialkylsilyl cyanides and metal carbonyls(34). The carbonyls of metals such as iron are known to react with isocyanides to form compounds in which some of the carbonyl groups are replaced by alkyl isocyanides; alkyl cyanides do not react so readily. Since trialkylsilyl cyanides form compounds such as Me3SiNOFe(CO)4 with iron pentacarbonyl, this has been taken as showing that some at least of the trialkylsilyl isocyanide must be present in the material referred to as "trialkylsilyl cyanide" under the conditions of the reaction; the results can most simply be explained by postulating a labile equilibrium of the kind already discussed. While the evidence has some force, it must be borne in mind that the structures of the compounds formed by iron carbonyl with silyl cyanides have yet to be determined; it is possible that silyl cyanides would react differently from alkyl cyanides with metal carbonyls, and so these observations seem to be suggestive rather than conclusive.

As far as silyl cyanides are concerned, almost all of the chemical evidence as to structure comes from studies of the trialkylsilyl compounds, while the most definite physical evidence refers to silyl cyanide itself. The position regarding the structures of silyl cyanides is thus thoroughly confusing. The physical evidence shows that a compound of the structure SiH3CN certainly exists; the only physical evidence for the existence of silyl isocyanides comes from some features of the infrared spectra of trialkylsilyl cyanides which can almost certainly be explained in other ways. There is, on the other hand, strong chemical evidence for the presence of at least some isocyanide at room temperature in trialkylsilyl cyanides. It seems to be generally agreed that this amount is small; since the cyanide groups are well known to be labile when bound to silicon, it is likely to be very difficult to distinguish between chemical reactions of a labile "normal" cyanide and those of an "isocyanide" in labile equilibrium with the "normal" compound. In this context it is relevant to remember that the SiC


and the SiN bond energies in saturated compounds have been given roughly the same value(35»3e); unless the energies of the SiN= and SiC= bonds differ appreciably from these values (and the bond lengths indicate that this is unlikely), there will not be much thermodynamic difference between the stabilities of the normal and the isocyanides.

There is no chemical evidence to modify the assignment of the isostructure to silyl thiocyanates, and since the SiN bond is probably appreciably stronger than the SiS bond this is not surprising. Silyl cyanates have so far been given the iso-structure; this fits in with their reactions with amines. Thermodynamically the structure containing a silicon-oxygen bond seems likely to be more stable. The structures of trimethylgermyl and trimethylstannyl cyanides have not been determined; the germanium compound reacts readily with iron carbonyl in the same way as its silicon analogue, but the tin compound does not react readily in this way; the difference has been put down to the salt-like character of MesSnCN.

2. COMPOUNDS CONTAINING SILICON BOUND TO TRANSITION ELEMENTS

Trimethylchlorosilane reacts with the sodium derivative of π-cyclo-pentadienyldicarbonyl iron to give 7>cyclopentadienyl(trimethylsilyl) iron<37):

(C6H6)(CO)2FeNa+Me3SiCl = NaCl+(C5H5)(CO)2FeSiMe3

The compound decomposes at 200°. Tetramethylstannane reacts with iron pentacarbonyl to give the compound (CO)4Fe-Me2Sn-Fe(CO)4

<38).

3. COMPOUNDS CONTAINING SILICON-BORON BONDS

Triphenylsilylpotassium reacts with(B-trichloro, N-triorgano)borazenes to give products which are believed to contain silicon-boron bonds(39):

3Ph3SiK+B3N3Cl3Me3 = 3KC1+ Ph3Si

A MeN NMe

I I Ph3SiB BSiPh3

Me


The compounds react with moist air, and the silicon-boron bonds are apparently broken by bromine.

An anion containing a silicon-boron bond has been obtained by treating the triphenylsilyl anion with triphenylboron(39a) :

PI13SÌ-+BPI13 = [Ph3SiBPh3]-The tetramethylammonium salt was isolated, as was its germanium analogue. The compounds such as Li(SiH3BH3), obtained by the low-temperature. reaction between tetrachlorosilane and lithium borohydride, may contain silicon-boron bonds, but nothing is known about their structures(39b).

4. METAL DERIVATIVES OF SILANES

No compounds of silicon analogous to Grignard reagents have been characterized, although there is evidence that triphenylchlorosilane forms an unstable magnesium derivative(40) (see Chapter 3); unstable trior-ganosilylmercury compounds may be formed in the reactions between triphenylsilylmetallic compounds and mercuric halides(41) (see below), while iodosilane reacts with zinc metal to give an unstable compound, barely volatile at room temperature, which contains zinc.

Alkali metal derivatives of triorganosilanes(43), and of diphenylsilane(44), have been characterized, and monosilane itself gives a potassium derivative (45>45a). The triorganosilylmetallic compounds have been prepared by the action of alkali metals on symmetrical disilanes, trisilanes, alkoxy-silanes or halosilanes; silylpotassium is made from monosilane and potassium, or from disilane and either potassium or potassium hydride, in an appropriate ether solvent*45·45a). Silylpotassium is a white, crystalline solid which decomposes in vacuum at about 240°, and reacts with hydrogen chloride and with chloromethane according to the equation:

SÌH3K+RCI = S1H3R+KCI

Its reactions with bromosilane and with tetrabromosilane are complex, and there is extensive hydrogen exchange in the reaction with bromosilane; silylpotassium also exchanges hydrogen with monosilane. Some potassium borohydride is formed in the reaction between silylpotassium and di-borane(45a).

Triphenylsilylmetallic compounds react with halides of carbon(43)

(see p. 76), silicon(44) (seep. 86),mercury (41), and Group Velements*47·48*; with hydrogen bound to nitrogen, giving silylamines(43); and withhydroxy-


compounds, giving silanes or siloxanes<49). They add across multiple carbon-carbon<60), carbon-nitrogen(51), carbon-oxygen(61), and nitrogen-nitrogen(61) bonds; their chemical properties have been extensively review-ed(43). Alkali metal germyls and stannyls such as NaGeH3

(63), and NaSnH3

(54), as well as their alkyl and aryl derivatives(65»66), are also known; they have been prepared by the reactions between stannane, germane or digermane and the alkali metal concerned in liquid ammonia. Germylene-disodium, GeH2Na2, is' formed in the reaction between germane and sodium(45), while stannylenedisodium has been obtained as a solid which decomposes in vacuum(64-57) at 0°; both of these compounds react with the simple hydride concerned to give the monosodium derivative:

SnH2Na2+SnH4 = 2SnH3Na

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194 (1959). (2°) C. EABORN, / . Chem. Soc, 3077 (1950). (21> A. G. MACDIARMID, Quart, Rev., 10, 208 (1956). caia) H . R. LINTON and E. R. NDCON, Spectrochim. Acta, 10, 299 (1958). ( « ) w . KUCHEN, Z. anorg. all. Chem., 288, 101 (1956). <**) E. A. V. EBSWORTH, M. J. MAYS and S. G. FRANKISS unpublished work. <"> J. J. MCBRIDE, J. Org. Chem., 24, 2029 (1959).

OTHER SILICON COMPOUNDS 155 (24»> D . W. J. CRUICKSHANK, / . Chem. Soc, 5486 (1961). ( 2 5 ) F . A . MILLER and G. L. CARLSON, Spectrochim. Acta, 17, 977 (1961). <*% G. S. FORBES and H . H . A N D E R S O N , / . Amer. Chem. Soc, 70, 1043 (1948). ( 2 e a ) J. G O U B E A U and H. R E Y H I N G , Z . anorg. all. Chem., 294 , 96 (1958). (,7> A . G. M A C D I A R M I D , Thesis, Cambridge, 1954. <M> H . H . A N D E R S O N , / . Amer. Chem. Soc, 81 , 4785 (1959). (2·> E . A . V. EBSWORTH and M . J. M A Y S , unpublished work. (so> J. J. M C B R I D E and H . C. BEACHELL, / . Amer. Chem. Soc, 74, 5247 (1952). (S1> J. GOUBEAU and E. HEUBACH, Ber., 93 , 1117 (1960).

<"i E . C. EVERS, W . O . FREITAG, J. N . Κ Ε Π Ή , W. A . K R I N E R , A . G. M A C D I A R M I D

and[S . SUJISHL, / . Amer. Chem. Soc, 81 , 4493 (1959). (M> E . C . EVERS, W . O . FREITAG, W. A . K R I N E R , A G . M A C D I A R M I D and S. SUJISHI,

/ . Inorg. Nucl. Chem., 13, 239 (1960). (84> D . SEYFERTH and N . K A H L E N , / . Amer. Chem. Soc, 82 , 1080 (1960). (85> T. L . COTTRELL, The Strengths of Chemical Bonds, Butterworths 2nd edn. (1958). <M> H . KRIEGSMANN, Z. Elekt., 61 , 1088 (1957). (27> T. S. PIPER, D . LEMAL, and G. WILKINSON, Naturwiss., 43 , 129 (1956).

<"> R. B. K I N G and F . G. A . STONE, / . Amer. Chem. Soc, 82 , 3833 (1960). <w> A . H . COWLEY, H . H . SISLER and G. E . RIYSCHKEWITSCH, / . Amer. Chem. Soc,

82 , 501 (1960): Μ . V. G E O R G E , G. D . LICHTENWALTER and H . G I L M A N , / . Amer.

Chem. Soc. 81 , 978 (1959): D . M. SEYFERTH and H. P. K Ö G L E R , / . Inorg. Nucl. Chem., 15, 99 (1960).

<«*> D . M . SEYFERTH, G. R A A B and S. O. G R I M , / . Org. Chem., 26, 3034 (1961). (i»b> H . N o r a , Angew. Chem., 7 3 , 371 (1961). Uo> T. G. S E U N and R. W E S T , Tetrahedron, 5, 97 (1959). (41> M. V. GEORGE, G. D . LICHTENWALTER and H. G I L M A N , / . Amer. Chem. Soc,

81,^978 (1959). <**> H . J. EMELEUS, A . G. M A D D O C K and C. REED, / . Chem. Soc, 353 (1941). ' (4e> D . WITTENBURG and H . G I L M A N , Quart. Rev., 13, 116 (1959). ("> H . G I L M A N and W. STEUDAL, Chem. Ind., 1094 (1959). <4δ> K. M A C K A Y , Thesis, Cambridge, 1960. <«·) Μ . A . R I N G and D . RITTER, / . Amer. Chem. Soc, 83 , 802 (1961). <*·> D . WITTENBERG, M . V. GEORGE and H . GELMAN, / . Amer. Chem Soc, 8 1 , 4 8 1 2 (1959). <«') M. V. GEORGE, B . J. G A J and H . GELMAN, / . Org. Chem., 2 4 , 624 (1959). <") Μ . J. N E W L A N D S , Proc. Chem. Soc, 123 (1960). (°> R. A . BENKESER and R. G. SEVERSON, / . Amer. Chem. Soc, 73 , 1424 (1951). (60> H . GELMAN and T. C. W u , / . Amer. Chem. Soc. 75, 234 (1953). <81> D . WITTENBERG, M. V. GEORGE, T. C. W U , D . H . MÊLES and H. GELMAN, / . Amer.

Chem. Soc 80 , 4532 (1958). ("> H . GELMAN and T. C. W u , / . Amer. Chem. Soc, 76, 2502 (1954). <M> C. A . K R A U S and G. K. T E A L , / . Amer. Chem. Soc, 72 , 4706 (1950). (M> H. J. EMELÉUS and S.F.A. KETTLE, / . Chem. Soc, 2444 (1958). (M> C. A . K R A U S and L. S. FOSTER, / . Amer. Chem. Coc, 49, 457 (1927). (M> C. A . K R A U S and W. H. K A H L E R , / . Amer. Chem. Soc, 55, 3537 (1933). <*7> S. F. A. KETTLE, Thesis, Cambridge, 1958.

CHAPTER 8

CONCLUSIONS

FROM this survey of the structures and chemical properties of volatile silicon compounds, it is clear that there are many gaps in our knowledge of these systems which must be filled before a properly coherent picture can be presented. In terms of molecular structure, it is the theoretical rather than the experimental side of the problem that is deficient; a number of structural parameters of silyl compounds, particularly in the halides and simple alkyl derivatives, have been measured, but the interpretation of these is at present very nearly impossible, and the best that can be achieved is an attempt to see how far common influences are involved in different properties. There is a notable absence of any precise information about the structures of disiloxane, disilyl sulphide and disilyl selenide, and the structural examination of germyl derivatives has scarcely begun. The wide bond angles to silicon from oxygen and mtrogen are among the most important pieces of evidence (summarized in the second part of this chapter) to show the occurrence of (p-*d) π-bonding; for the rest, the bonds from silicon to electronegative elements are almost always shorter than the sums of the relevant atomic radii, but it is not at all clear how this should be interpreted. The short bonds generally go with bond energies which are greater than the energies of bonds from the same elements to carbon, and also with force constants which are greater than the values predicted by Siebert's formula(1) ; an exception to this is the SiH bond, which, although shorter than the sum of the atomic radii of silicon and hydrogen, has a lower force constant than the predicted value, and a bond energy substantially less than the energy of the CH bond. This lends some support to Kriegsmann's contention<2), that the high force constants indicate (j>->d) π-bonding, which is not possible with hydrogen (though the high H-0 force constant in water conflicts with this interpretation); the theoretical basis for this and for Siebert's rule is, however, a little obscure, and the position is not made easier by the diversity of methods for calculating force constants from observed spectra, which lead to widely differing results(3). Bonds from silicon seem to be rather less sensitive than bonds from carbon to changes in the rest of the molecule; this is particularly noticeable where

156

CONCLUSIONS 157

the SiH -stretching modes of monohalosilanes are concerned. The relative insensitivity of the chemical shift of hydrogen bound to silicon to changes in the rest of the molecule may be due to changes with substitution in two competing effects.

The information at present available about the energies of bonds from silicon is not very extensive, and where germanium and tin are concerned the position is still less satisfactory. The value for the SiS bond energy, for example, is uncertain, and there appear to be no precise measurements of the energies of the SiN or GeO bonds. This seriously hampers any discussion of chemical reactions in which these species are concerned.

From the point of view of chemical studies, there is an immense amount of data about certain systems, and little or none about others. If trade follows the flag, it has certainly been true of research in silicon chemistry that the banner of scientific advancement has been carried forward on a surge of commercial expansion, and in directions which have been to some extent determined by this. The immense amount of information now available about SiO and SiC systems is in marked contrast to the relative neglect of the compounds of silicon with nitrogen or sulphur. It would be interesting, for instance, to have a detailed study of the condensation-type reactions of SiNH and SiSH groups to compare with the relatively extensive information that has been obtained about the silanol conden-sation(4). There are major aspects of the preparative chemistry of silyl compounds that are still more or less untouched, and the recent preparation of iodogermane and chlorostannane shows that it will be possible to take the comparison of the hydride-derivatives of silicon, germanium and tin further than at one time seemed likely. None the less, the study of equilibria in the chemistry of silyl compounds seems likely to be less important in the future than the study of mechanisms.

In any comparison of the behaviour of silicon compounds with analogous derivatives of carbon and the other elements of the group, it is necessary to have some idea not only about the products of the reactions involved but also of the paths by which these products were formed. If only the products are known, then it is very difficult to know how far differences in chemical behaviour should be interpreted in terms of changes in thermo-dynamic properties and how far they should be put down to differences in activation energies; moreover, it is all too easy to assume that the mechanisms of two apparently related processes are the same, when there is no necessary reason why this should be so. This in its turn can lead to a false int rpretation of the differences observed. Bonds from silicon to


hydrogen, halogens, and elements of Groups V and VI are usually more reactive and more labile than the analogous CX bonds, while it seems from what few facts are available that GeX bonds are more labile than SiX bonds, but react rather less readily with oxidizing agents. In terms of particular systems, silicon-halogen bonds are hydrolysed by what seems to be an SN

2* mechanism, and the reaction proceeds very much more readily than in carbon chemistry; this is probably to be associated with the presence of unfilled d-orbitals in the valence-shell of the silicon atom(6). Mechanisms of this sort have been suggested for almost all reactions of silicon compounds with nucleophilic species, though there is frequently no evidence to show what the mechanism of a particular reaction really is; it is dangerously easy to ascribe all differences between the reactivities of silicon and of carbon compounds to the J-orbitals of silicon, and it is possible that other factors are more important than is sometimes realized(6,.The condensation reactions of silanols are catalysed by both acid and base, as is the hydrolysis of the SiH bond; it is possible that attack at the atom bound to silicon may be important in a number of reactions of silicon compounds*4"7*. If, for example, there is extensive (p->d) π-bonding between silicon and nitrogen, then addition of hydrogen halide may reduce the π-bonding by attack of H+ at nitrogen, which would in its turn leave the silicon atom more ready to accept an electron-pair from an external donor (in this case, a halogen atom). This type of action may also explain why aluminium halides are such effective catalysts in the displacement of X" from silicon; the X atom, be it hydrogen, halogen or any other group, becomes partly bound to the aluminium halide, leaving the silicon atom more ready to accept electrons from any other species in the vicinity. This sort of catalytic action has not yet been sufficiently studied; it would be interesting to see if boron halides catalyse the reaction between monosilane and hydrogen halides.

The effectiveness of base catalysis in these systems is less easy to understand. The most obvious way for base to attack silicon compounds is by coordination to the silicon atom, but there is no apparent reason why this should lead to disproportionation. Perhaps it helps a bonded group to become detached from the silicon atom in question by increasing the electron density at silicon. Among the reactions catalysed by bases and Lewis acids are the redistribution of bonds from silicon to hydrogen, halogen, oxygen, carbon and silicon.

* Or through the formation of an intermediate containing 5-coordinated silicon— see p. 63.

CONCLUSIONS 159

This leads to another important feature of silicon chemistry—the relative inertness of SiC and SiSi bonds. The bond energies are not large (see Table 2.5), and though the SiSi bond is non-polar in a symmetrical molecule, the SiC bond should be at least as polar as the SiH bond, and probably as polar as the SiBr bond, particularly if any(p->d) π-bonding between silicon and bromine reduces the bond moment in the latter. Speculations of this sort make it clear at once, however, that it is impossible to compare bond polarities in compounds like these, with no knowledge of the extent of π-type interactions or of their effect on the polarity of a given bond. All that can be said is that the SiC bond does not seem likely to be of unusually low polarity. None the less, if an aliphatic group is bound to silicon, the SiC bond is remarkably inert, particularly to attack by base. This may mean that the CH-bonds release electrons to the rf-orbitals of silicon (though this suggestion shows that the instinctive interpretation of reactions of this sort is in terms of nucleophilic attack at silicon); the presence of methyl groups bound to silicon reduced the tendency of the silicon atom in question to form additional σ-bonds with external electron-pair donors(8), so thus far at least the theory and facts are in agreement. There is, however, another possible factor to be considered—the absence of any electron-pairs or low-lying excited states in the attached group. If, as suggested in the previous paragraph, attack upon the group bound to silicon plays some part in the reactions of silicon compounds with nucleophilic systems, it is easy to see why the SiC bond should be relatively inert. It is interesting to note that if the carbon atom is unsaturated or has other substituents attached to it, the chemical activity of the SiC bond is increased.

The SiSi bond, too, is remarkably inert. Though in disilane it is fairly readily broken by base, it is much less active in the fully-alkylated disilanes. In particular, hexamethyldisilane reacts with concentrated sulphuric acid and ammonium chloride to give pentamethylchlorodisilane(9), the SiC bond breaking in preference to the SiSi bond. Considering the relevant bond energies, the most probable reaction appears to be an oxidative splitting of the SiSi bond. Moreover, although hexamethyldisilane reacts with chlorine to give trimethylchlorosilane, it is possible to chlorinate penta-methylchlorodisilane and other methylchlorodisilanes in the methyl groups, keeping the SiSi and the SiC bonds intact(10); the reaction conditions are typical of free radical reactions, and [this suggests that the reactions of the SiSi bond with halogens do not involve free radicals—a suggestion that is borne out by the results of kinetic studies. The stabilizing influence

160 VOLÀTILE SILICON COMPOUNDS

of chlorine substituents upon the SiSi bond is hard to understand, unless (which does not seem very likely) it is simply a question of steric hindrance.

Reactions which have been shown to involve free radicals are not common in silicon chemistry. The addition of SiH bonds to unsaturated systems is of this sort, and the cracking and polymerization of the simple hydrides and alkyls of silicon are probably also free radical processes, but these do not occur readily except under fairly vigorous conditions. The formation of siliconium ions has not been established in any of the reactions of silicon compounds*, and several attempts to make compounds which contain siliconium ions such as Pl^Si* have given either decomposition products (as in the case of thefluoroborate)(11) or compounds which are probably not salt-like [as with the perchlorate(1213) or with the silicon analogue of crystal violet, tris(/?-dimethylamino-phenyl)chlorosilane(14)]; moreover, no sign has been detected of the ionization of silicon compounds such as triphenylchlorosilane in solution in dimethylformamide<14). It must be concluded that "electron-deficient silicon" (that is, a silicon atom with less than eight electrons in its valence-shell) is an extremely reactive species. In this it contrasts with tin and (to a lesser extent) with germanium; organo-tin cations such as Me3Sn+ are relatively well-characterized. One feature of the attempts to make siliconium ions has been that the systems studied have almost always contained silicon bound to some electronegative element such as fluorine, oxygen or chlorine, where the bond to be broken is exceptionally strong, and where any alternative structure in which the covalent bonds were preserved is thermodynamically favoured. It would be interesting to study solutions of triphenylbromosilane or trichlorobromo-silane in some acceptor solvent such as fused gallium tribromide; unfortunately, however, a solvent of this sort is all too likely to catalyse disproportionate of the silicon compound in question.

It is therefore clear that silicon is more susceptible than carbon to attack by nucleophilic reagents, but that the mechanisms of a number of the reactions of silicon compounds are more complicated than appears at first sight. In the interpretation of the differences between silicon and carbon compounds, the behaviour of silicon as a σ-electron acceptor from external donors is probably of primary importance, but its function both as a σ-donor and as a π-acceptor when forming four formal single bonds may be of greater importance than has sometimes been allowed.

* Kinetic evidence has been obtained to support the proposal of a "limiting siliconium ion" mechanism in one case(10a>.

CONCLUSIONS 161

Future work may well include exchange studies as an important feature. There have been some recent investigations into the exchange reactions of silicon-hydrogen and silicon-halogen bonds(le~18), and this type of work may well prove to be of increasing importance; one problem towards whose solution it might provide most useful evidence is the question of the structure of the trialkylsilyl cyanides. With a knowledge of the reactions of silicon compounds and of the mechanisms of the processes involved, it may well prove possible to make a thoroughgoing comparison between silicon and carbon, but at present it is possible only to indicate the lines along which such a comparison may develop. Much more experimental information is needed about the chemical behaviour of the simple compounds of germanium and tin before these could be included in a detailed comparative discussion of that kind.

In the preceding section, the interesting question of (p-+d) π-bonding has been more or less omitted; in my opinion, this is one of the most significant of the causes of difference between analogous compounds of silicon and carbon. It is not easy, however, to show conclusively that this sort of interaction plays an important part in silicon chemistry, and so the rest of this chapter is given up to a brief and critical summary of the evidence given in more detail in the rest of this monograph. It must be emphasized that the opinions given as to the relative importance of the different pieces of evidence are entirely my own. The evidence falls naturally into two groups—physical and chemical in nature.

(a) Physical evidence

Short bonds (that is, bonds significantly shorter than the sum of the relevant atomic radii) may indicate (/?-></) π-bonding, ^ u t this is not the only explanation, (see p. 50). High force const ants{1-2), may also indicate (ρ->ά)π bonding, but there is little theoretical justification for this interpretation (see pp. 54, 156). Nuclear quadrupole coupling constants give information about a combination of effects, from which the extent of (p->d)n bonding can only be obtained if assumptions are made about relative electronegativities and hybridization (see p. 52-3). Bond angles. The wide bond angles at nitrogen in trisilylamine and silyl isothiocyanate, and at oxygen in disiloxanes, afford perhaps the strongest physical evidence for (p-+d)n bonding between silicon and these atoms. The narrower bond angles at sulphur in silicon-sulphur compounds, how-

11


ever, do not necessarily mean that there is little or no π-bonding between silicon and sulphur (see pp. 105, 125, 135 and 147). N.M.R. chemical shifts in complex molecules cannot at present be given a satisfactory interpretation, but the changes in proton resonance chemical shift with substitution in the fluorosilanes suggest that there is some unusual interaction between silicon and fluorine (see pp. 17 and 23). Dipole moments are extremely difficult to interpret, but some of the changes in group moments with substitution in aromatic silicon compounds suggest that there may be appreciable (/?-><f)7c-interaction m these systems, as do the small dipole moments of the halosilanes (see pp. 57 and 81). Ultraviolet spectra. The marked shift in the frequency of the (n—π) transition in oc-silyl ketones can be interpreted in terms of (p->d) π-bonding between silicon and the carbonyl group (see p. 81). Bond energies. The large energies of bonds from silicon to electronegative elements with lone pairs could be partly due to (/?->*/)7r-bonding, but once again there are other possible explanations. (b) Chemical evidence Donor properties. The electron donor properties of nitrogen, oxygen and sulphur are reduced by an attached silicon atom; this is strong evidence for π-interactions between silicon atoms and adjacent lone pairs, but is of no help in deciding the importance of such bonding between silicon and halogen atoms (see pp. 109, 130 and 137). Properties of p-Me3SiCtH4X. The dissociation-constants of p-trimethyl-silylbenzoic acid can most satisfactorily be interpreted in terms of π-bond-ing between the silicon atom and the ring, as can the variable Hammett σ-constant of the trimethylsilyl group (see p. 82). Hydrogen bonding. The strong hydrogen bonding in silanols, and the apparently weak hydrogen bonding in primary and secondary silylamines, can also be explained in this way, as can the acidity of the hydroxyl group when bound to silicon (see pp. 117 and 125).

Although it is difficult to say that any one of these points is conclusive, the weight of evidence strongly suggests that π-bonding is important between silicon and nitrogen, oxygen, fluorine, and (probably) chlorine, and between silicon and a benzene ring. There is a conflict of evidence about sulphur, but in view of the uncertainties involved in the interpretation of the physical evidence, I accept the conclusion from the chemical properties that the lone pairs of electrons at sulphur are probably involved in π-bonding with silicon. On the same basis, there seems to be little or no

CONCLUSIONS 163

such π-bonding between silicon and phosphorus, while there is not enough evidence to decide if there is significant π-bonding between silicon and bromine or iodine. A study on these lines of analogous compounds of germanium and tin would be most interesting. What evidence there is suggests that (/?->c/)rc-bonding from oxygen to germanium or tin is weaker than from oxygen to silicon*19*, but there is room for a great deal of experimental work on the sulphur and nitrogen compounds of germanium and tin.

REFERENCES (1> H. SIEBERT, Z. anorg. all. Chem., 273, 170 (1953). <2> H. KRIEGSMANN, Z. anorg. all. Chem., 299, 138 (1959). (3) see T. L. COTTRELL, The Strengths of Chemical Bonds, Butterworths, 2nd. ed.

Chapter 11 (1958). (4> W. T. GRUBB, / . Amer. Chem. Soc, 76, 3408 (1954). (5> A. D. ALLEN and G. MODENA, / . Chem. Soc, 3671 (1957). (e> J. E. BAINES and C. EABORN, / . Chem. Soc, 1436 (1956). (?) S. H. LANGER, S. CONNELL and I. WENDER, / . Org. Chem., 23, 50 (1958). (8> Ε. A. V. EBSWORTH and H. J. EMELÉUS, / . Chem. Soc, 2150 (1958). ( 9 ) M. KUMADA, M. YAMAGUCHI, Y. YAMAMOTO, J. NAKAJIMA and K. SHIINA, / . Org.

Chem., 21, 1264 (1956). (10> M. KUMADA, J. NAKAJIMA, M. ISHIKAWA and Y. YAMAMOTO, / . Org. Chem., 23,

292 (1958). (10a> L . H . SOMMER and G. A. BAUGHMAN / . Amer. Chem. Soc, 83, 3346 (1961). (11) B. STERNBACH and A. G. MACDIARMID, / . Amer. Chem. Soc, 81, 5109 (1959). (li> U. WANNAGAT and W. LIEHR, Angew. Chem., 69, 783 (1957). (18> U. WANNAGAT, F. BRANDMAIR, W. LIEHR, and H. NIEDERPRUM, Z. anorg. all. Chem.,

302, 185 (1959). (14> F. BRANDMAIR and U. WANNAGAT, Z. anorg. all: Chem., 280, 223 (1955). (15> A. B. THOMAS and E. G. ROCHOW, J. Amer. Chem. Soc, 79, 1843 (1957). (le> R. H. HERBER and A. W. CORDES, / . Chem. Phys., 28, 361 (1958). <17> A. F. REID and R. MILLS, / . Chem. Soc, 708 (1960). (18> R. H. HERBER and SHIH-CHEN CHANG, / . Inorg. Nucl. Chem., 17, 385 (1961). <w) R. OKAWARA, Proc. Chem. Soc, 383 (1961).

INDEX (Entries in roman type describe major references; entries in italics

refer to tables)

Acetonitrile bond lengths in, 79-80 physical properties, 145 as solvent for adducts, 67-68

Adducts of germoxanes (as donors), 133 of halogermanes (as acceptors), 69 of halosilanes

(as acceptors) with ammonia and organic amines,

5, 59, 62, 66-68; in preparing siloxanes, 12; in preparing silyiamines, 102; in preparing silyl sulphides, 134

with oxygen-containing donors, 68 with phosphines, 62, 68 with silyiamines, 115

(as donors), 60 of halostannanes (as acceptors), 69, 104 of iodoalkanes

(as acceptors), 67 (as donors), 60

of pseudohalogermanes (as donors), 152 of pseudohalosilanes (as donors), 149-52 of silicon-oxygen compounds

(as acceptors), 68, 132 (as donors), 130

of silyiamines (as donors), 102, 112-15 of silyl phosphines (as donors), 118-19

Alkali metal complex hydrides

in preparation of hydrides, 13-15 reaction with

chloromethylsilanes, 84 halosilanes, 13, 61-62 hexachlorodisilane, 94

silyiamines, 13 SiO systems, 13, 129 Si pseudohalides, 13, 149 SiS systems, 13, 137

germyls, formation, 39, 40, 95, 154 hydrides

and preparation of hydrides, 12-14 alkali metal silyls, 39, 153

reaction with halosilanes, 61 pseudohalosilanes, 149 SiO systems, 129 SiS systems, 137

silyls formation, 27, 38-39, 61, 90, 92, 153 hydrolysis, 12 reactions

(general), 153-4 with boron compounds, 152-3 with halosilanes, 64, 86-87 with multiple bonds to nitrogen, 104 with phosphorus halides, 117-18 with SiH bonds, 38

stannyls, formation, 40, 154 a-elimination, 84 Alumina, as catalyst, 13-14, 30 Aluminium

as catalyst, 13-14 bromide, as catalyst, 63, 94 chloride

as catalyst, 14,35,37,60,94 reaction with SiO compounds, 131

halides as catalysts, 31-32, 35-36, 40, 84-85,

92, 158

165

166 INDEX

reaction with SiC systems, 84-85; with SiO systems, 130-1

iodide, as catalyst, 29 trimethyl, reaction with silylamines, 113

Atomic radii (table), 7

Base catalysis of disproportionate reactions

general, 158 SiH rearrangement, 31-32, 34, 102,

115, 138 Si-halogen rearrangement, 31-32, 102,

115, 138 SiN rearrangement, 110-11 SiSi rearrangement, 87, 93 silanol condensation, 127-8

of SiH hydrolysis, 34 ß-elimination, 33, 83-85 Bond angles

in amines, 105-7 and molecular structure (general), 7,

9-10, 156; and (P-+d) π-bonding, 161

in monohalides, 19; and vibrational frequencies, 25

in oxygen compounds, 7, 9-10,124, 125 in pseudohalides, 145, 147 in selenides, 139 in sulphides, 135

Bond dipole moments of SiH bond, 15-17 of Si-halogen bonds, and nuclear quad

ruple resonance, 53 of SiN bonds, 105

Bond energies and activation energies

(SiC reactions), 82-84, 159 (SiO reactions), 131 (SiSi reactions), 91-92, 159

conversion series, 65, 69, 144 isomerism in pseudohalides, 152 (p-+d) π-bonding, 162 SiC reactions, 82-83 SiH reactions, 27, 37, 40 table of, 26

Bond length to carbon, 78; changes with substitution,

79; changes with unsaturation, 79-80, 147

determination of, 9-10 and electronegativity, see Schomaker-

Stevenson correction Ge to Ge, 89 to halogen, in hydride-halides, 54-55,

56; in tetrahalides, 50-51 to hydrogen, changes with substitution,

18-19; correlation with other changes, 24; in simple hydrides, 15-16

to nitrogen, 105 to oxygen, 124, 125 and (P-*d) π-bonding, 161 to pseudohalogen groups, 145, 147 to sulphur, 134-5

Boron halides as catalysts, 158 reactions with

SiO compounds, 130 silylamines, 113-14 silyl cyanides, 149-50

Boron hydrides, see Diborane, Penta-borane

Boron trichloride as catalyst, 30 reaction with

monosilane, 38 SiO compounds, 130 silyl sulphides, 138

Boron trifluoride, reaction with fluorosilane, 60 SiO compounds, 130-1 silylamines, 49, 113-14 silyl cyanides, 150 silyl sulphides, 138

Boron trimethyl, reaction with silylamines, 110, 113 silyl sulphides, 138

Carbon dioxide, 37 Carbon divalent, 4

INDEX 167

Carbon halides adducts of, 60, 67 hydride-halides

bond angles in, 19; and CH vibratio-nal frequencies, 25

CH bond lengths in, 19; and vibra-tional frequencies, 24

CH vibrational frequencies in, 21, 24-25

C-halogen bond lengths in, 54-55, 56 dipole moments of, 58 proton resonance in, 22-23 quadrupole coupling constants in, 57

reactions with silicon hydrides, 37, 94-95 tetrahalides

bond lengths in, 50-51 physical properties, 50

Carbon monoxide, 4, 14, 37 Conversion series

for Si compounds, 65-66, 138, 144, 149 for Ge, Sn compounds, 69

d-orbitals contraction of, 5, 7, 55 and hyperconjugation, 78, 136, 159 and π-bonding, see (p-d) π-bonding and σ-bonds, 5, 22, 158-9 and UV spectra, 80-81, 162

Diborane reaction with

siloxanes, 130-1 silylamines, 113 silyl cyanides, 149 silylpotassium, 153

Dimethyl ether bond angles in, 7, 124, 125 "CH satellites in H resonance spectrum

of, 135-6 electron donor properties of, 130 physical properties of, 124

Dimethyl sulphide "CH satellites in H resonance spectrum

of, 135-6 physical properties of, 134-6

Dipole moments of hydride-halides, 57-8

and (/>-►*/) π-bonding, 162 and π-bonding with aromatic systems,

81-82 Disproportionation (redistribution) reac

tions SiH migrating, 13-14, 31-33 Si-halogen migrating, 31-33, 58-59,

92-94, 102, 115, 138 SiN migrating, 33, 110-11 SiSi migrating, 87, 90, 93 silanol condensation, 123, 127-8

Double bonds to silicon, 6

Electron diffraction, 106 Electron donor properties

of germoxanes, 133 and (p-+d) π-bonding, 162 of siloxanes, 125-6, 130, 162 of silylamines, 17, 102-3, 108-9, 111,

113-16, 162 of silyl phosphines, 118-19 of silyl selenides, 139 of silyl sulphides, 137, 162

Electron resonance, 9, 28 Electronegativity, 8-9

and bond length, see Schomaker-Stevenson correction

andTF resonance chemical shifts, 51-52 andAH resonance chemical shifts, 21-23,

78, 106-7 and quadrupole coupling constants, 53,

57 and "vibrational frequencies of SiH

bonds, 20

Force constants of MC bonds, 77-78; and unsaturation,

79 of MH bonds, 15-17 of M-halogen bonds, 50, 53-55 of MM bonds, 88, 89 of MN bonds, 105 of MO bonds, 124, 125 of MS bonds, 134-5 predicted, and bond polarity, 17, 54;

π-bonding, 54, 156, 161

168 INDEX

Formaldehyde, 13

Gallium trimethyl, 113, 130 Gennanium, divalent, 4, 5, 40-41, 68

GeC compounds chemical properties, 85 formation, 40, 69, 77 physical properties, 78y 85

GeGe compounds chemical properties, 39, 95, 154 formation, 88 physical properties, 88-89

GeH compounds chemical properties, 39-41; reactions

with C=C bonds, 77; with halogen hydrides, 49; with organo-metallic compounds, 77

formation, 14-15 physical properties, 16, 17, 21

Ge-halogen compounds chemical properties, 67-69; reactions

with amines, 104; with mercury (II) sulphide, 139; with silylme-tallic compounds, 88; with siloxy-lithium compounds, 124; with sodium selenide, 139; with water, 124

formation, 40, 49-50, 85, 95 physical properties, 50-54, 57-58

Ge-metal compounds chemical properties, 88 formation, 39, 40, 95, 154

GeN compounds chemical properties, 117, 124 formation, 104

Ge pseudohalides chemical properties, 152 formation, 145 physical properties, 148 structures, 147-8, 152

GeO compounds formation, 117, 124 properties, 131-3

GeS compounds, 139 GeSe compounds, 139 GeSi compounds, 88, 95

Germanium compounds: halides

GeBr4, bond energy in, 26; physical properties, 50, 53

GeHBr8, formation, 49 GeH2Br2, formation, 49 GeH3Br, formation, 49; quadrupole

coupling constant, 57 GeCl2, formation, 5 GeCl4, adducts of, 69; bond energy

in, 26; hydrolysis of, 69; physical properties, 50, 53

GeHCla, decomposition of, 4-5; formation, 1Φ-15, 49; ionization of, 4(M1

GeH2Cl2, formation, 49 GeHsCl, formation, 49; GeH stret

ching frequencies in, 21 ^quadrupole coupling constant in, 57; reaction with ammonia, 104

GeF4, F resonance chemical shift in, 51; physical properties, 50

GeH2F2, formation, 49 GeH3F, adduct with ammonia, 69;

formation, 49 Gel«, crystal structure of, 4 Gel4, bond energy in, 26; physical

properties, 50; quadrupole coupling constant in, 53

GeH3I, formation, 49,157; hydrolysis, 124; reaction with mercury (II) sulphide, 139

hydrides GeH4, H resonance in, 17fjphysical

properties, 16; preparation, 14; reaction with hydrogen halides, 40, 49; with oxygen, 39; with sodium, 39, 154; thermal] stability, 39

Ge2He, physical properties, 89; reaction with sodium, 39, 95, 154

Higher germanes, formation, 88, 95 metal compounds

GeH3Na, formation, 39, 95, 154 GeH2Na2, formation, 154

INDEX 169

nitrogen compounds (Ph3Ge)2NH, reaction with hydrogen

peroxide, 117 (Ph3Ge)3N, formation, 104

organogermanes MeGeH3, GeC bond length in, 78 Me4Ge, physical properties of, 77-78 organogermanium cations, 85, 160

organohalides Me3GeCl, formation, 131 Me3GeF, formation, 133 Ph3GeBr, reaction with ammonia, 104

oxygen compounds (GeH3)20, possible formation of, 124 Me3GeOMe, reaction with boron tri-

fluoride, 133 Me3GeOSiMe3, formation, 124; reac

tion with halides, 131 Me3GeOGeMe3, reaction with boron

trifluoride, 133 Et3GeOGeEt3, reaction with acetic

acid, 132 (Ph3Ge)202, formation, 117 GeO, reaction with HC1, 15 Ge02, estimated bond energy in, 26;

and preparation of GeH4, 14 pseudohalides

Me3GeCN, physical properties, 148; reaction with iron pentacarbonyl, 152; structure of, 148, 152

Ge(NCO)4, structure of, 147 sulphur compounds

(GeH3)3S, formation, 139 GeS, react io with hydrogen halides,

49 selenium compounds

(Me3Ge)2Se, formation, 139 Grignard reagents

reactions with Ge-halogen bonds, 69 SiH bonds, 38, 76 Si-halogen bonds, 13, 63-64, 77 SiN compounds, 117 SiO compounds, 126, 129 Si pseudohalides, 149 Sn-halogen bonds, 69

silicon compounds analogous to, 61,153

Hammett σ-constants, 34, 82, 162 Hybridization, 4-8

and bond angles, 105-8, 125-6; length, 79-80

and NMR parameters, 10, 41 and quadrupole coupling constants, 53

Hydrazine, as solvent, 12, 87-88 Hydrogen bonding

in OH-compounds, 108, 126, 132-3 and (p-+d) π-bonding, 162 in silylamines, 108, 117

Hydrolysis of GeH bonds, 40-41 of Ge-halogen bonds, 41, 69, 124 of GeN bonds, 117, 124 of GeS bonds, 124 of germanides, 14, 88, 95 of SiC bonds, 84-85 of SiH bonds, 34, 40, 122, 158 of Si-halogen bonds, 35, 62-63, 122-3,

127-8, 131, 158 of SiN bonds, 104, 112, 122, 131 of SiO bonds, 129, 131 of Si pseudohalides, 149 of SiS bonds, 35, 122-3, 127, 137 of SiSe bonds, 139 of SiSi bonds, 93 of suicides, 12-13, 87-88 of SnH bonds, 40 of Sn-halogen bonds, 69, 124 of stannides, 14, 88, 95

Hyperconjugation, 78, 136, 159

Infrared spectra and NMR spectra, 25-6 and structures of

adducts, 66, 115 siloxanes, 10, 126 silylamines, 106-7 silyl pseudohalides, 146-7, 151

Isomerism in disilane derivatives, 92 in higher silanes, 89 in pseudohalogen compounds, 145-7,

150-2

170 INDEX

Lead hydrides, 15 Lewis acids

as catalysts general, 158 in SiC reactions, 85 in SiH reactions, 27, 29, 30, 36

see also Aluminium halides Liquid ammonia, 12, 39, 40, 87-88, 95,129 Lithium aluminium hydride, 13-15, 84, 94,

149 Lithium borohydride, reaction with tetra-

chlorosilane, 13, 62, 153 Lone pairs

and activation energies, 83, 111 and (P->d) π-bonding

(general), 7, 78, 162 in siloxanes, 125-6 in silylamines, 105-9 in silyl phosphines and arsines, 119 in silyl sulphides, 135

and (p->p) π-bonding, 114

Methane physical properties, 16, 17-18 reaction with trichlorosilane, 30

Methyl cyanide, see Acetonitrile Methyl isothiocyanate, physical properties

of, 145 Microwave spectra

and bond lengths, 9, 79, 145-6 and quadrupole coupling constants, 57 and structures of silyl pseudohalides,

145-6

Nuclear magnetic resonance chemical shift

F, in fluorosilanes, 54-57; in tetra-fluorides, 51-52

H, changes with substitution, 21-24; and electronegativity, 8; in Me4M, 78; in MH4, 16-18; and (P-*d) π-bonding in fluorosilanes, 23, 162; and (P-*d) π-bonding in silanols, 126; and (p-+d) bonding in silylamines, 105, 107-8;

solvent shifts and association, in fluorosilanes, 54, in silyl isothiocyanate, 148, in trichlorosilane, 41

coupling constants, 10, 15 in fluorosilanes, 24, 25', 57 and hybridization radii, 80 in Me4M, 78 in SiH-derivatives, 24, 25 in silyl sulphides, 33, 135-6

and electronic structure, 10 and isomerism in higher silanes, 89 and SiF4-PH8 adduct, 62 and structure of silyl cyanides, 147 spectrum of trisilane, 88-89

Nuclear quadrupole resonance general, 10, 15, 52-53 in hydride-halides, 57 and (P-+d) π-bonding, 161 in silylamines, 108 in tetrahalides, 52-53

Organotin cations, 58, 85, 160

p-orbitals and σ-bonds, 5, 7, 53 and (p-*d) π-bonds

general, 6-7 and bond angles in siloxanes, 126; in

silylamines, 105-8 ; in silyl pseudohalides, 147; in silyl sulphides, 135

and bond dipole moments, 105, 159 and bond lengths, in halosilanes,

50-51; in siloxane 126; in silylamines, 105; in silyl sulphides, 135; in tetrafluorogermane, 51

and dipole moments, 58, 81-2, 105 and donor properties of siloxanes,

125-6; of silylamines, 105, 108, 109; of silyl phosphines and arsines, 119

and force constants, 54, 156; in halosilanes, 54; in siloxanes, 126; in silylamines, 105; in silyl sulphides, 135

and H bonding, in OH compounds, 108, 126, 133; in silylamines, 108

INDEX 171

and mechanisms, 158, 160 and NMR chemical shift, F, 51-52,

56-57; H, 18, 23, 107-8 and nuclear quadrupole coupling con

stants, 53 and reactions of chlorinated disilanes,

92-93 stability of SiHXH derivatives, 138-9

summary of evidence for, 161-3 UV spectra, 80-81, 85

and (P-P) TT-bonds, 6, 7, 91 and (P-+P) π-bonds, 2, 106, 114

Pentaborane, reaction with tris(trimethyl-silyl)phosphine, 118

Pyrolysis (thermal stability) of halosilanes, 59 of higher silanes and derivatives, 90, 93 of monosilane, 28, 87 of organosilanes, 13, 83 of siloxanes, 127 of silylamines, 109 of silyl phosphines, 118 of silyl pseudohalides, 148 of silyl sulphides, 136

Quadrupole resonance, see Nuclear quadrupole resonance

Radical reactions of SiH systems, 27-28, 39, 78, 87 in silicon chemistry, 159-60 of SiSi systems, 90-91

Raman spectra, 83 see also Infrared spectra, Vibrational

spectra

s-character and bond length, 80 and orbital electronegativity, 107, 125

Scandinide contraction, 7, 16 Schomaker-Stevenson correction

general, 17, 50-51 and MC bonds, 77-78

and MH bonds, 16, 17 and MM bonds, 88, 89 and SiN bonds, 105 and SiO bonds, 124, 125 and SiS bonds, 134, 135

Silicon, divalent, 4, 87 SiAs compounds, 61, 65, 118-19 SiB compounds, 62, 152-3 SiC compounds

chemical properties, general, 82-83, 159; a-elimination, 84; ß-elimi-nation, 33, 83-85; reactions with elements, 48, 84; with halides, 85; with hydrides, 13, 48-49, 84-85; thermal stability, 83-84

formation, 27-28,31,38,62-64, 76-77, 126, 129, 153, 160

physical properties, 77-82 SiH compounds

chemical properties, general, 26-28; disproportionation, 13-14, 31-33, 35,127;exchange, 27, 31, 36,161 ; in higher silanes, 37, 92-95; ioni-zation, 40-41 ; reactions with elements, 27-30, 37, 48, 153; with hydrides, 30-36, 49, 92-93; with multiple bonds to carbon, 27-28, 31, 39, 77, 160; with organome-tallic compounds, 38-39, 76; with oxides and oxyacids, 36-37; with salts and halides, 37-38

formation, 12-14, 84, 94, 149 physical properties, 15-26

Si-halogen compounds chemical properties, general, 58-59;

disproportionation, 14, 35, 59-60, 92-93,102,115;exchange, 58-60; formation of adducts, 5, 59, 60, 62, 66-68, 102; reactions with elements, 6,13-14,60-61,86,117-18, 153; with halides, and pseudohalides, 38,49,58-60,64-65,144; with heavy metal salts, 65-66,104, 133,139,144; with hydride, 13,35, 61-63,101-4,122-3,127-8,133^; with organometallic compounds, 13, 63-64, 76-77, 86-87; with

172 INDEX

salts, 64-65, 88, 103-4, 118, 133, 139; thermal stability, 59

formation, 29-30,35-38,48-50,83-85, 91-92, 94, 111-16, 118, 129-32, 137-9, 150

physical properties, of hydride-hal-ides, 54-58; of tetrahalides, 50-54

Si-metal compounds, see Alkali metal silyls, Grignard reagents

SiN compounds chemical properties, general, 109; dis-

proportionation, 109-10; reactions with electron acceptors, 48, 102-3, 113-16; with elements, 48; with hydrides, 13, 102, 111-13, 123, 135, 144; with other compounds, 116-17

formation, 33,60,62.65,101-4,129,151 physical properties, 105-9

SiNH compounds chemical properties, 117 disproportionation, 103, 109-10 physical properties, 108, 116, 126

SiO compounds chemical properties, general, 128;

disproportionation, 127; formation of adducts, 130, 132; reactions with elements, 48, 128, 153; with halides and electron acceptors ,^, 129-31; with hydrides, 13, 48, '4128-9; with organome-tallic compounds, 77, 117, 129; thermal stability, 127

SiOH compounds chemical properties, 132 condensation, 35, 123, 127-9 formation, 123,*131 physical properties, 126

SiP compounds formation, 33, 61, 65, 117-18 properties, 35, 48, 118-19, 138

Si pseudohalides chemical properties, 13, 149-51 formation, 37, 65-66, 77, 144-5 physical properties, 145-8 structure, 145-8, 150-2

SiS compounds chemical properties, general, 136;

reactions with elements, 48, 137; with halides and electron acceptors, 48, 137-8; with heavy metal salts, 65, 138; with hydrides, 13, 35, 48, 123, 127, 137; with tri-methylamine, 32, 138

formation, 63, 65-66, 133-4 physical properties, 134-6

SiSe compounds, 48, 65, 139 SiSi compounds

chemical properties, general, 90, 159-60; reactions with base, 87, 89, 90, 93; with hydrides, 13, 92-94; with oxides and oxyacids, 36, 94; with other compounds, 94-95; thermal stability, 90

formation, 6, 28, 38, 59, 61, 85-88 physical properties, 88-90

SiSn compounds, 88, 95 Silicon compounds

halides SiBr4, adducts of, with amines, 66,

with oxides, 68; bond energy in, 26; bond length in, 56; eQq in, 53; physical properties of, 50; reaction with magnesium, 86; with hydrogen iodide, 63; with silylpotassium, 87, 153

SiHBr8, dipole moment of, 58; formation of, 48 ; melting point, boiling point of, 55; reaction with magnesium, 61

SiH2Br2, formation of, 49; melting point, boiling point, 55

SiH8Br, in conversion series, 65; dipole moment of, 58; eQq in, 57; formation, 29, 35, 49; H resonance chemical shift in, 22; H resonance couplings, 25; melting point, boiling point, 55; reaction with dimethylamine, 101, 103; with magnesium, 61 ; with organo-metallic compounds, 76; with silylpotassium, 31, 87, 153; SiBr

INDEX 173

bond length in, 56; SiH bond length in, 19; SiH stretching frequencies in, 21

SiCl4, adducts of, 66, 112; adduct with pyridine, structure of, 66-67; bond energy in, 26; bond length in, 56; eQq in, 53; exchange with Cl2, 60; with Cl, 64; formation, 13, 38, 90, 93; physical properties, 50; reaction with ammonia, 62, 66; with complex hydrides, 13, 62, 153; with diphenyl sulphoxide, 68; with formaldehyde, 13-14; with hydrogen bromide, 63 ; with hydrogen sulphide, 63; with nitrogen, 60-61; with siloxanes, 129; with silylamines, 111; thermal stability, 59

SiHCls, dipole moment of, 58; flash point of, 30; formation of, 13,48, 76, 93; H resonance chemical shift in, 23; melting point, boiling point, 55; reaction with benzene, 30; with chloromethane, 30; with ethylene, 77; with hydrogen, 14, 60; with iodine, 30; with methane, 30; with oxygen, 30; solvent shifts in H resonance, 41

SiH2Cla, formation of, 31, 35, 49, 102; possible formation from tetra-chlorosilane and formaldehyde, 13-14; H resonance chemical shift in, 23; melting point, boiling point, 55; possible reaction with potassium silyls, 86

SiH3Cl, in conversion series, 65; dipole moment, 58; disproportion-ation in, 31; eQq in, 57; formation, 35, 49; possible formation from tetrachlorosilane and formaldehyde, 13-14; H resonance chemical shift in, 22, 23; H resonance couplings in, 25; hydrolysis, 123; reaction with NH-groups, 101-3,109-11 ; with organometal-lic compounds, 63-64; with phos-phylsodium, 118; with sodium,

61, 86; SiCl bond length in, 9, 56; SiH bond length in, 19; SiH stretching frequencies in, 21

SiF4, adducts of, 5, 62, 66, 68; bond energy in, 26; bond length in, 56; exchange with F, 64; F resonance chemical shift, 51-52, 56; F resonance coupling, 57; physical properties, 50; reaction with atomic hydrogen, 60; thermal stability, 59

S1HF3, association in, 54; dipole moment of, 58; F resonance chemical shift in, 56; F resonance couplings in, 57; formation of, 38, 49; H resonance chemical shift in, 23-24; melting point, boiling point, 55; SiF bond length in, 56; SiH bond length in, 19, 24; SiH stretching frequency in, 24

SiH2Fa, association in, 54; F resonance chemical shift in, 56; F resonance couplings in, 57; formation of, 38, 49; H resonance chemical shift in, 23; H resonance couplings, 25; melting point, boiling point, 55; SiF bond length in, 56; SiH bond length in, 19

S1H3F, adduct with trimethylamine, 138; association in, 54; in conversion series, 65 ; dipole moment of, 58; disproportionate of, 31-32; F resonance chemical shift in, 56; F resonance couplings in, 57; formation, 38, 49, 83, 113-14, 130, 138; H resonance chemical shift in, 22-3 ; H resonance couplings, 24, 25; melting point, boiling point, 55; SiF bond length in, 56; SiH bond length in, 19; SiH stretching frequencies in, 21

S1H3I, adducts with trimethylamine, 67-68,5134; and conversion series, 65; disproportionation of, 31; eQq in, 57; formation of, 29, 49, 137; H resonance chemical shift,

174 INDEX

22; H resonance coupling, 25; melting point, boiling point, 55; reactions with antimony, arsenic, 61,118; with hydrazine, 101 ; with hydrogen sulphide, 63; with mercury (II) sulphide, 133, 138; with phosphorus, 61, 117; with silver cyanamide, 104; with silver cya-nate, 148; with silver cyanide, 77; with silver selenide, 139; with silver selenocyanate, 149; with silver thiocyanate, 144; with sodium amalgam, 61, 86; with tri-methylamine, 67-8; with trisilyl arsine, trisilylphosphine, 119; with zinc, 153; SiH bond length in, 19; SiH stretching frequencies in, 21

SiH2I2, formation of, 29, 49, 61; melting point, boiling point, 55

S1HI3, formation of, 116; melting point, boiling point, 55

Sil4, adducts of, 68; bond energy in, 26; eQq in, 53; physical properties of, 50; reaction with hydrogen, 60; thermal stability, 59

Halopolysilanes halodisilanes, dispro-portionation in, 31, 92; formation of, 92

SiaCle, reaction with base, 87, 93; with lithium aluminium hydride, 94; SiSi bond length in, 89; thermal stability, 90

Si2Bre, formation, 90 Si2Cl2Ph4, reaction with sodium, 6 Si2ClMeß, formation, 159; reaction

with chlorine, 159; with sodium, 86

Si5Cl12, formation, 87; isomerism in, 90; reaction with hydrogen chloride, 93

SieCl14, formation, 87, 93; isomerism in, 90;

SiioCl», formation, 59 hydrides

SiH4, absence of adduct with tri-methylamine, 68; bond energy in,

26; bond length in, 16-17; bond moment in, 16; exchange in, 27, 31,36; force constant, 17; formation, 14, 31-2, 60-61, 86, 109-10, 118, 127, 130, 136-9, 149-50; H resonance chemical shift in, 17-18; H resonance couplings, 24, 25; no reaction with chloroform or with phosphorus triiodide, 37, 94; physical properties, 16; reaction with ammonia, 33; with antimony pentachloride, trichloride, 38; with boron trichloride, 38; with hydrogen halides, 35, 49, 92, 158; with organometallic compounds, 38; with phosphine, 33, 117; with phosphorus pentachloride, 38; with potassium, 39, 153; with sulphuric acid, 36; with silylpotassium, 153; with tin (IV) chloride, 38; with water, 33-34; stability to irradiation, 28; thermal stability, 28

Si2He, formation, 61, 86-87; physical properties, 89; SiH reactions (general), 28; with alkali, 93; with ammonia, 93; with hydrogen halides, 92; with iodine, 37, 91; with lithium aluminium hydride, 94; with potassium, 39, 92, 153; with potassium hydride, 39, 153; with oxides and oxyacids, 36; with water, 93

Si8He, NMR spectrum of, 88-89; reaction with chloroform, 94; with sodium amalgam, 92

Si4H10, reaction with bromoform, with chloroform, with iodoform, 94; with phosphorus triiodide, 94; with sodium amalgam, 92

Higher silanes, formation, 12, 87-88; isomerism in, 89

Metal compounds silylpotassium, formation, 39, 92,

153; reaction with bromosilane, 31, 87, 153; with diborane, 153;

INDEX 175

with monosilane, 153; with tetra-bromosilane, 87, 153

triphenylsilylpotassium, formation, 38-39, 61, 90, 92, 153; reactions (general), 153-4; with boron compounds, 152-3; with halo-silanes, 64, 86-87; with multiple bonds to nitrogen, 104; with phosphorus halides, 118; with SiH bonds, 38

nitrogen compounds SiHsNHa, disproportionation of, 138;

possible formation of, 101, 110 SiH3NMea, disproportionation of, 33,

109; donor properties of, 102; formation of, 101, 103; H resonance chemical shift, 107-8; reaction with aluminium trimethyl, 113; with boron trifluoride, 114; with boron trimethyl, with diborane, with gallium trimethyl, 113; physical properties, 107-8

(SiH8)2NH, disproportionation of, 32, 109-10, 138; donor properties of, 103; possible formation of, 101, 103

(Me3Si)aNH, bond energy in, 26, 105; formation of, 111, 129; H bonding in, 117; lithium derivative, 103, 117; magnesium derivative, 117; no disproportionation, 111; reactions with amines 111; with Grignard reagents, 117; with hydrogen cyanide, 144-5 ; with phosgene, 104; with sodium mercap-tides, 112; with tetrachlorosilane, 111 ; with water, 112,123 ; sodium derivative, 65, 103-4

(SiH3)aCN2, formation of, 104 (SiH3)aNMe, disproportionation of,

32; H resonance chemical shift in, 107-8; physical properties of, 107; reaction with aluminium trimethyl, 113; with boron trifluoride, 114; with boron trimethyl, with diborane, with gallium trimethyl, 113

(SiH3)3N, formation of, 101-3; H disproportionation in, 32, 109; H resonance chemical shift in, 108; H resonance coupling constants, 24, 25; physical properties, 107; reaction with aluminium trimethyl, 113; with boron trifluoride, 113; with boron trimethyl, 113; with diborane, 113; with gallium trimethyl, 113; with hydrogen chloride, 111; with water, 112; SiN bond angles in, 105-7, 161; SiN bond length in, 105; SiN force constant in, 105

(Me3Si)3N, formation of, 65, 103-^; physical properties, 107; SiN bond angles and vibrational spectrum, 106; stability, 111

(SiH3)4Na, formation, 33, 101; reaction with iodosilane, 114; SiN bond angles and vibrational spectrum of, 107

organohalides Me3SiBr, formation of, 115 Me3SiCl, formation of, 83, 131, 159;

reaction with alkali metal phos-phyls, 65, 117; with alkali metal silyls, 64, 86; with ammonia, 111 ; with lead mercaptides, 133; with sodium salts, 65, 103-4, 133, 139, 152; with urea, 144; Würtz-Fittig, 86

Me3SiF, F resonance chemical shift in, 56; formation, 150; reaction with alkali metal arsenyls, phos-phyls, 65, 117-18

Me3SiI, formation, 90-91, 115; reaction with silver cyanide, 145

Cl3CSiCl3, reaction with lithium aluminium hydride, 84; with water, 84

organomonosilanes MeSiH3, H resonance couplings and

SiC bond length in, 80; reaction with hydrogen halides, 84; SiC bond length in, 78, 79-80; SiH

176 INDEX

bond length in, 19; SiH stretching frequency in, 24

CH2=CHSiH3, SiC bond length in, 79

CeH5SiHa, and preparation of mono-halosilanes, 49, 84

Me4Si, pyrolysis of, 6; physical properties, 77-79

Me3SiPh, electronic effect of Me3Si-group in, 82; UV spectrum of, 80

organopolysilanes MeeSi2, reaction with iodine, 90-91;

with sulphuric acid and ammonium chloride, 159; rearrangement in, 93

PheSi2, formation, 38; reaction with i ithium, 61; with potassium, 90; thermal stability, 91

Me polysilanes, formation of, 86 cyclo-Ph8Si4, formation of, 6, 86; reac

tion with halides, 95 cyclo-Ph12Sie, formation of, 86; reac

tion with halides, 95 oygen compounds

(i) alkoxysilanes SiH3OMe, formation of, 123; reac

tion with boron trifluoride, with diborane, 131

SiH3OEt, formation of, 123 (ii) silanols

"SiH3OH", possible formation as intermediate, 35, 123, 127, 138

Me3SiOH, acidity, 133; condensation, 127; formation, 83,112,123; H bonding in, 126-7; H resonance in, 126

Et3SiOH, condensation, 128; O exchange with water, 132

(iii) disiloxanes (SiH3)20, and conversion series, 65;

formation, 35, 123, 127, 137; H disproportionate in, 32; H resonance spectrum, couplings in, 25; H resonance spectrum, physical properties in, 124; 2*SiH satellites in, 135-6; reactions

with boron trifluoride, 130; with chlorine, 128; with diborane, with dimethylaluminium bromide, with gallium trimethyl, 130; SiOSi bond angle in, 10, 125, 156

(MeSiH^O, reaction with boron halides, 130; with hydrogen iodide, 63 ; with methyl iodide, 129

(Me3Si)20, no acceptor adducts with N bases, 68; formation, 104; large atomic polarization in, 126; reaction with oxidizing oxides, 131; with potassamide in liquid ammonia, 129; SiO bond angle in, 124; SiO bond length in, 124; SiO force constant in, 124-5

(Cl8Si)20, formation, 128; no partial reduction of, 13; no reaction with boron trifluoride, 130

(iv) solids silica, bond energy in, 26 siloxene, reactions of, 28, 36-37

phosphorus compounds SiH8PH2, disproportionation of, 138;

formation of, 33, 117; physical properties, 118; reactions with hydrogen bromide, with sodium hydroxide, 118

(SiH3)3P, possible adduct with iodo-silane, 118-19; formation of, 117

(Me3Si)3P, formation of, 65, 117; physical properties, 118; reactions with diborane, with penta-borane, with sulphur, 118; thermal stability 118

pseudohalides (i) cyanides

SiH3CN, and conversion series, 65-66; formation of, 77; isom-erism in, and infrared spectrum, 146-7; and microwave spectrum, 145; and reactions, 150; physical properties, 145; reactions with boron trifluoride, 150; with di-borane, 149; with sulphur, 150;

INDEX 177

SiC bond length in, 79-80, 145, 147; thermal stability, 148

MeSiH2CN, isomerism in, and infrared spectrum, 146; isomerism in, and NMR spectrum, 147

Me3SiCN, and conversion series, 66; formation of, 144-5; isomerism in, and formation, 145; isomerism in, and infrared spectrum, 146,151 ; isomerism in, and reactions, 150-1; reactions with boron trifluoride, 150; with hydrogen, 151; with iron penta-carbonyl, 151; with sulphur, 150

Et3SiCN, isomerism in, and infrared spectrum, 146; reaction with mercury (II) oxide, 149

PhsSiCN, isomerism in, and infrared spectrum, 146; structure of, and molar refractivity, 147

(ii) isocyanates S1H3NCO, and conversion series,

65; formation of, 148-9 PhSiHaNCO, formation of, 148 Me8SiNCO, formation of, 144 Si(NCO)4, isomerism in, 145; reac

tion with secondary amines, 149; structure of, 147

(iii) isothiocyanates SiHsNCS, association in, 147-8;

and conversion series, 65; decomposition of, 148.; formation of, 144; isomerism in, 146; physical properties, 145; structure of, and microwave spectrum, 146; structure of, and π-interactions, 147, 161

MeSiHaNCS, thermal stability of, 148

Me3SiNCS, formation of, 150 Si(NCS)4, structure of, and vibra-

tional spectrum, 147 (ii) selenocyanates

SiH»NCSe, formation, 149 sulphur compounds

(i) alkyl silyl sulphides

SÌH3SCH3, formation of, 134 S1H3SCF3, disproportionate of,

32, 83, 136; reaction with tri-methylamine, 138

Me3SiSCMe3, formation, 133 (ii) thiols

S1H3SH, disproportionate of, 35, 136, 138; formation of, 136, 137

Me3SiSH, formation of, 134 CI3S1SH, formation of, 63, 133;

SiS bond length in, 134 (iii) disilyl sulphides

(SiH3)aS, and conversion series, 65, 138; formation of, 133; H disproportionate in, 32; H resonance spectrum, 2*Si satellites in, 24, 135-6; physical properties of, 134; reaction with boron halides, with boron trimethyl, 138; with hydrogen sulphide, with hydrogen iodide, with iodine, 137; with lithium aluminium hydride, 137; with methyl iodide, 137; with sulphur, 136; with trimethylamine, 138; with water, 35,123,127, 137; SiSSi bond angle in, 135; thermal stability, 136

(SiH3)2S2, possible formation of, 137 (Me3Si)2S, SiS bond energy in, 26,

134; SiS force constant, 134, 135 selenides

(SiH3)2Se, and conversion series, 65; decomposition of, 32; formation, physical properties and reactions of, 139

(Me„Si)2Se, formation of, 139 transition metal compounds

Me3SiFe(CO)2(C6H5), formation of, 152

5-coordinated silicon, 59, 63, 67, 115, 158 6-coordinated silicon, 67-68 Siliconium ions, possible formation of, 27,

59, 85, 160

Taft o*-constants and H resonance chemical shifts, 22

178 INDEX

and vibration frequencies, 20 Thermal stability, see Pyrolysis Tin, divalent, 4, 41

SnC compounds chemical properties, 152 formation, 40, 69, 77, 95 physical properties, 77-78, 85 π-bonding in aromatic, 85

SnH compounds chemical properties, 39-40; reactions

with C=C bonds, 40, 77; with organometallic compounds, 40; 77; with sodium, 39-40, 154

formation, 14-15, 88 physical properties, 16-17

Sn-halogen compounds, chemical properties, 69; reactions

with amines, 69, 104; with heavy metal salts, 139, 145; with organometallic compounds, 77; with salts, 119, 124

formation, 40, 49, 77 physical properties, 50-54

Sn-metal compounds, 39-40, 154 SnN compounds, formation, 104 SnO compounds

chemical properties, 132-3 formation, 39-40, 69, 124 H bonding in, 133 physical properties, 132-3

SnP compounds, formation, 119 Sn pseudohalides

chemical properties, 152 formation, 145 physical properties, 148

SnS compounds, formation, 139 SnSi compounds, formation, 88, 95 SnSn compounds, formation and prop

erties, 88, 95 Sn-transition metal compounds, forma

tion, 152 Tin compounds:

halides SnBr4, SnCli, Snl4, bond lengths in,

50-51; eQq's in, 52-53; force

constants in, 53-54; physical properties of, 50

SnCl4, reaction with ammonia, 104 SnF4, adducts of, 69; bond length in,

50-51; physical properties, 50; volatility, 50

SnH8Cl, formation, 40, 50, 157 hydrides

SnH*, formation, 14, 88; physical properties, 16; reactions with alkali metals, 154; with hydrogen chloride, 40, 50; with oxygen, 39; thermal stability, 39

Sn2He, formation, 88, 95 metal derivatives SnH3Na, formation,

154 oxygen compounds

MesSnOH, basicity of, 133 Ph8SnOH, hydrogen bonding in, 133 Me8SnOSiMe8, formation, 124

organic derivatives MeSnHs, SnC bond length in, 78 Me4Sn, physical properties, 78; reac

tion with iron pentacarbonyl, 152 MeeSnt, force constant in, 89; reaction

with CFJ, 95 Me,SnBr, reaction with phosphyl-

sodium, 119 Me,SnCl, adducts of, 69; reaction

with trimethylsiloxylithium, 124 Me8SnF, structure, 85 Me3SnI, formation, 95

pseudohalides Me,SnCN, melting point, 148; no adduct with iron pentacarbonyl, 152

Trimethylaluminium, see Aluminium trimethyl

Trimethylamine adducts of, 66-68, 108, 123, 134 bond angles in, 105-6 and disproportionation reactions, 32,138 H resonance chemical shift in, 107-8 melting point, boiling point, 107 reaction with hexachlorodisilane, 87, 93

Trimethylboron, see Boron trimethyl Trimethylgallium, see Gallium trimethyl

INDEX 179

Ultraviolet irradiation, 28, 39 Ultraviolet spectra, 9, 68

and π-bonding, 162 in aromatic derivatives, 80-81 ina-substitutedketones, 80-81,85,162

and bond length, 9, 24 frequency variation with substituent,

19-21 and H resonance chemical shift, 25-26 see also Infrared, Raman

Vibrational spectra and bond angles, to N, 105-7; toO, 9-10;

to Se, 139; in Si(NCO)4, 147; correlation, 25

and bond energy, 26, 105, 134

Wurtz-Fittig reaction, 86

X-ray diffraction, 9, 67

Volatile Silicon Compounds

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