SOME SALTS of VERY STRONG ACIDS. by Kenneth Charles Moss. a Thesis Presented in Partial Fulfilment of the Requirements for the D

Home , Fluoroboric acid, Silver perchlorate, Tetrafluoroborate

SOME SALTS OF VERY STRONG ACIDS.

By Kenneth Charles Moss.

A thesis presented in partial fulfilment of the requirements for the degree of Doctor of Philosophy of the University of London*

AUGUST 1962. -1- ABSTRACT

An investigation of monohydroxyfluoroborates was undertaken with reference to the sodium, potassium and tetraalkyl ammonium salts as the existence of these, compounds is in some doubt. Sodium and potassium monohydroxyfluoroborates were prepared and characterised, but the tetraalkylammonium salts could not be prepared. Several fluoroborate metal hydrates were prepared and their X-ray photographs indexed. They are found to be isomorphous with the corresponding perchlorates. It was found that silver(2) fluoride reacts with boron trifluoride to give silver(1) fluoroborate. A study oil the solubility of anhydrous first-row transition metal perchlorates, fluoroborates, trifluoroacetates, hexafluoro -phosphates, -vanadates, -silicates and -titanates in organic solvents such se benzene, toluene,ether and nitromethane was carried out. The d-d spectra of the solutions of these salts in ether were obtained. The preparation of solid anhydrous fluoroborates, perchlorates and trifluoroacetates from non-aqueous solvents was investigated and mum perchlorates and trifluoroacatates were obtained. Magnetic measurements were made on the solid trifluoroacetates and on solutions of the perchlorates in ether. -12- The stability of complexes of phosphorous pentafluoride with Group V triaryls was investigated, and.it was found that only triphenylphosphine forms a stable complex. Phosphine-metal-fluorides of Pt, Pd and Ir were prepared from phosphlne-metal-hydrides and are the first complexes to be prepared with both phosphine and fluorine bonded to the same metal atom. The existence ofid4PLIPd(0) has been disproved and io shown to be iAPLIPdHs. The compound APhiPt[C0i2F2 has also been prepared. No correspondin; first-row transition meta/ complexes could be prepared. .,-3- ACKNOWLEDGEMENTQ.., ? The work in this thesis was carried out in the Inorganic Chemistry Laboratories of the Imperial College of Science and Technology, London, and in the Royal College of Science and Technology, Glaegow. I wish to express my sincere thanks to my supervisor, Dr. E. R. Roberts,and to Dr. D.W.A. Sharp for his continued .help and encouragement. Thanks are due also to Dr. L. Pratt for help with n.m.r. measurements and to the Microanalytical Departments in both London and Glasgow. I am'indebted to Imperial Chemical Industries Ltd., Fisons Ltd., and the United States Navy (European Research Programme, Contract No. N 62258-3061) for financial support. -4- CONTENTS. Page. Abstract 1. Acknowledgements. 3. Introduction. 5. Chapter I. Some Fluoroborates and Monohydroxyfluoroborates, 11. Chapter 2. The Preparation of Some 1st Row Transition Metal Salts of Strong Acids. 67. Chapter 3. Some Physical and Chemical Properties of lst Row Transition Metal Fluoreborates, Perchiorates and Trifluoroacetates. 104. Chapter 4. Some Phosphine Fluoride Complexes. 148. Chapter 5. Molybdenum and Tungsten Phthalocyanines, 179. References. 184. -5- INTROD. UCTIO N

The study of, fluorine chemistry, intensified because of the nuclear programme developed during the last World Wary has received further stimulation from the preparation of fluorocarbons and from their interesting properties. These fluorOearbons have many uses such as inert plattice, aerosols, dielectrics, lubricants, solvents and pharmaceuticals. The study of fluorine derivatives of carbon opens up a vast new field of chemistry in the same way that the boron hydrides would seem to do. The importance of complex anions in. catalysis (F2iedel and Crafts type) is becoMing increasingly rare obvious0 Strong acids usually provide such anions. The strength of an acid depends - not only on the nature of the acid itself but also on the.medium in which it is dissolved. For example nitric acid acts as• an acid in water but as a base in hydrogen fluoride.

HNO3 + nH20 [11,ni + ion r + NO3—

HNO3 + nHF [H21105 if (FnIlmeir

The complex fluoro acids are all extremely strong acids due to the nature of the very electronegative fluorine. -6- Fluoroboric acid, for example, can only exist in solution when it is protonated;. no evidence for a 1:1 complex between hydrogen fluoride and boron trifluoride has been found (McCauley and Lien 1951). Isolation by itself would involve divalent fluorine or pentavalent boron or a free proton; similar conditions prevail for other fluoro-anions e.g. PPG -. A few examples of fluorine bridges are known however; the antimony pentafluoride- arsenic trifluoride complex has been given the structura

(Muetterties and Phillips 1957) and arsenic fluorosuiphonate the structure

FN N 0 0 .2A\ As F SO 0 0/ 0S---F 1 1 0 0 0/ \F (Muetterties and Coffman 195a). The value of complex halogen anions as catnlysi- e in organic chemistry and the preparation of organo-metallic compounds has been developed only comparatively recently. The discovery of the catalytic properties of anhydrous aluminium -7- chloride by Friedel and Crafts (1877) was the beginning of an intensive investigation into its properties and mechanism. Industrial use of hydrogen fluoride-boron trifluoride mixtures in the separation of aromatic isomers through formation of protonated systems is well established and has been thoroughly established by workers including Brown and Brady (/952). Complexes of the type ArH" BP4 have actually been isolated (Olah, Kuhn and Pavlath 1956) and have been extended to salts r- . such as F BFI (Olah, Noszaki and Pavlath

\._-NO 2 I H 1957). The solubility or miscibility of aluminium chloride and the fluoride systems is the key to their importance. The solubility of silver perchlorate (Hill 1921) and fluoroborate in organic solvents is well known and the latter has been shown to have catalytic properties in difficult alkylation reactions. Silver fluoroborate and ethyl bromide catalyse the formation of 1,3 dioxo- lenium salts from 2-alkyl-19 3 dioxolanes (Meerwein, Hederich and Wunderlich 1958). The discovery of Ziegler catalysts has stimulated the study of the solubility of inorganic salts in organic 8m solvents. Sharp and Sharpe (1956 a) have shown the solubility of a number of silver fluoro salts in aromatic systems. The bonding in such systems is undoubtedly due to Tr -bonding from the aromatic system and back bonding through the full 3d orbitals of the silver ion to the anti-bonding 1- orbitals of the aromatic system as in silver-olefin complexes. The main part of this thesis is devoted to a study of the solubility and properties of let row transition metal salts of strong acids in organic solvents of low dielectric constant. All the anions studied, except for C104- and CF3CO2are complex fluoro anions. These solutions, particularly solutions of trivalent titanium, are potentially important as catalysts. It is expected that these solutions may also be valuable starting points for the preparation of organo-metallic compounds. The solutions in benzene and toluene could be very important in the formation of carbonylst for example,which are known to proceed, in some cases, through an intermediate benzene complex. There are many Tr bonded complexes known but comparatively few have a fluorine atom bonded to the same transition metal as the Tr bonding ligand. This type of complex with the other halogens are well characterised -9- and in these complexes there is the possibility of back bonding to the halide ions because of the availability of unfilled d orbitals which are of not too high an energy. This back-donation is unlikely in fluorine complexes due to the non-availability of d orbitals to receive electrons. However, it is believed that, providing the correct methods for preparation can be found, Trbonded ligand-metal-fluorine complexes are as stable as the corresponding chlorides although experimental evidence suggests that the form of these complexes may be different from the complexes of the other halides where there may be additional back-bonding to the halogen atom. Platinum and rhodium carbonyl fluorides (Sharp 1960) have been prepared and their properties indicate that they are indeed different from other metal carbonyl halides of platinum and rhodium where the foimal oxidation number is always low. The platinum dicaibOnyloctafluoride is considered as having a co-ordination number of ten whereas rhodium dicarbonyltrifluoride is considered as being dimeric with possibly fluorine bridges. This thesis describes the preparation of some phosphine-metal- fluoride complexes and some phosphine-metal-carbonylfluoride complexes which are found to be very stable. -11- CHAPT_ERI

SWE FLUOROBORATES AND MONCHYDROXYFLUOROBOR:tT,Z.

At the beginning of the nineteenth century intensive investigation into the properties of the alkali and alka- line-earth metals was sparked off by Sir Humphrey Davy's discovery of the electrolytic processes of their simple molten salts. At that time neither fluorine nor boron had been isolated. However Davy (1808a) did isolate a dark combustible solid from the electrolysis of boric acid and a few months later the same product was obtained by heating boric oxide with potassium (Davy 1808b)0 At the same time Gay-Lussac and Thgnard (1811) prepared boron trifluoride by a reaction involving fluate 'of lime and glacial boracic acid' but they confused this gas with hydrogen fluoride. In 'Reserches Physicochimiques' (Gay- Lussac and Thgnard 1811) there is a detailed description of the preparation and properties of fluoroboric acid made from fluospar and boric acid. Many of the syntheses of boron trifluoride are just slight modifications of the methods just mentioned. The greatest practical difficulties experienced by the -12- investigators were in freeing the gas from impurities, usually hydrogen fluoride and silicon tetrafluoride. Other methods of preparation described involve the use of nitrogen triflum'ide and boi-On (Ruff 1931) and boron nitride and sodium fluoride (Moeser and Eidmann 1902). Although boron trifluoride was shown to have valuable catalytic properties e.g. in the polymerisation of propylene and isobutylene, as early as 1873 (Butlerow and Gorainow), not much interest was shown until 1927 when Hoffmann and Otto published the first of several papers on the polymerisation of olefins. These articles aroused considerable interest, but the use of boron trifluoride was restricted because it was not available commercially. The first large scale manufacture of the gas was started in 1936 in Cleveland' by the Ha'rshaw Chemical Company. Swinehart (1939, 1940) developed the efficient process which is the basis for its commercial production.

Nat B4 07 *H2 0 ÷ HP ( aqueous )----> Wag 0(13113 )4 + 1.2118 0 or H3 B03 + 6NH4 HF2 ----3 4NH3T + 11H20 + (NH4 )2 0.(BF2)4 Water is removed by heating and the solid mass is treated with the calculated amount of cold fuming sulphuric acid (20% sulphur trioxide) and then gradually warmed. -13- One of the first reactions tried with boron trifluoride was, naturally, its reaction with water. Much confusion has arisen about the species present in aqueous solution and the salts which can be prepared from it. In particular the existence of monohydroxyfluoroborates is disputed. The present work was undertaken with the hope that some of the confusion could be cleared up with special reference to potassium, sodium and tetraalkyl ammonium monohydroxyfluoroborates.

Sodium and Potassium Monohydroxvfluoroborates Gasselin (1894) and Gay-Ltissac and Thgnard (1809) formulated the product of the reaction of borontri- 0 fluoride withwater as B .3HF. Other workers \OH reported the same compound (Berzelius 1843, Basorow 1874) but it was not until 1933 that Meerweinand Pannitz rewrote the formula as BF3 .2H20. In two papers (1933, 1934) the preparation and properties of the dihydrate are described. It is a clear non-fuming liquid with a melting point of 6.2° (Greenwood and Martin 1951). Chemically the liquid dissolves metals, oxides and carbonates and forms adducts withdoxan, BP3 .21120.C4H0 02, and cineole (Meerwein and Pannitz 1934). More important -14- is the fact that the dihydrate can absorb another mole of boron trifluoride to form the monohydrate, BP3.H20, m.p. 5.4-6.0°. The first definite proof of an hydroxonium compound is given by Volmer (1924),; The crystal structures of ammonium perchlorate and perchloric acid monohydrate were shown to be isomorphous, and in analogy with this Klinkenberg and Ketelaar (1935) suggested that the boron trifluoride dihydrate should be considered as H3 e (BF3 OH)-: The X-ray photograph of BF3.2H20 taken at -60° was shown to be isomorphous with a'nmonium perchlorate and fluoroborate. It is orthorhombic with four molecules per unit cell. H304-(BFOH)- a = 8.74 + 0.061 NH4B1P4 a = 8.89 + 0.051 b = 5.64 + 0.03A b = 5.68 + 0.051 e = 7.30 + 0.101 c = 7.21 + 0003A (Klinkenberg and Ketelaar 1935) The structure [H2F]i. [BF2(0H2)2)- has been suggested as an' alternative representation (Sowa, Kroeger and + Nieuwland 1935) although no other example of an H2F ion is really known, but electrolytic data (Greenwood and Martin 1951) supports the formulation H30+ BF30H-. When the dihydrate is electrolysed there is no evidence for the

-15- evolution of fluorine or hydrogen fluoride at the cathode. In fact hydrogen is evolved at the cathode and oxygen at the anode in the ratio 2:1. The electrode processes can be written as:- cathode H20+ + H20 + 1112 anode 2BF20H- BF2H20 + BF2 + 102 The boron trifluoride generated is immediately absorbed by the dihydrate remaining to form the monohydrate. The relative mobilities of the ion could, in principle, be determined by analysis of the liquid in the cathode and anode compartments. The cathode liquid becomes progress- ively more dilute and the anode liquid more concentrated as electrolysis proceeds. Pord and Richards (1956), however, have found from nuclear magnetic resonance studies that in slowly cooled specimens of the dihydrate there is no ionization. The formulation of salts prepared from the dihydrate is contradictory in the literature and the existence of monohydroxyfluoroborates of the alkali metals is in some doubt; therefore a study of the sodium and potassium alts was carried out. Meerwein and Pannitz (1934) have claimed that suspensions of sodium hydroxide and methoxide react with boron trifluoride etherate to form sodium monohydroxy=

-16- fluoroborate and methoxyfluoroborate respectively. Klinkenberg (1937) using practically the same method supports his X-ray data for sodium monohydroxyfluoroborate with analytical data, powder photographs of mixtures which have the composition NaBF5 OH and vapour pressure measurements. The following cell dimensions are given:- WaBF4 a = 6.25 ± 0002A NaBF3 OH a = 6.24 + 0002A 0 0 b . 6.77 + MLA b = 6.82 ± Oc.01A 0 c 6.82 + 0001A0 c = 6.85 + 0.01 Further attempts by Wheeler and ',Tuttle (1954) to prepare the monohydroxy salt were unsuccessful sodium fluoroborate being obtained in each case. However another method of preparation is given by Ryss and Slutskaya (1952). In the present work sodium monohydroxyfluoroborate was prepared, according to these authors, from sodium bifluoride and boric acid. The X-ray photograph was taken and indexed on the orthorhombic system; the cell 0 0 0 dimensions were a = 6.31A; b = 6080A; c = 6083A. found to agree fairly well with those of Klinkenberg (1937); furthermore infrared spectra data added weight to the _1 argument. Bands occur at 3550, 3425 and 1645 cm. which must be due to hydroxyl; the principal band is also split probably because of hydrogen bonding. Spectra were also taken in the far infrared region. Sodium fluoroborate _1 shows a band at 549 cm. and a split band at 526 cm. 1 the shoulder being at 521 cm. However the spectrum of sodium monohydroxyfluoroborate has only one band s -1 at 520 cm. with ashoulder at 570 cm. The salt is also more soluble in water than the fluoroborate. Potassium monohydroxyfluoroborate has been prepared by the reaction of potassium bifluoride and boric acid (Ryss 1946, Wamser 1948). The salt B21'01(21.5E120 prepared by Travers and Malaprade (1930) was found to be identical to potassium monohydroxyfluoroborate by Wamser (1948) on comparison of X-ray photographs. Using the method of Ryss (1946) the potassium salt was prepared and its X-ray photograph taken. It was indexed on an orthorhombic 0 system with cell dimensions, a = 7085A; b = 7.35A; 0 c = 5.721. The infrared spectra of the hydrovfluoro- borate again shows hydroxyl bands, at 31209 2860 and 01 1630 cm. The principal band in this compound is also split, probably because of hydrogen bonding. In the far infrared region a band occurs at 520 cm. with a shoulder at 515 cm. Potassium fluoroborate has only -18- _1 _1 one band at 520 cm. with a shoulder at 532 cm. The monohydroxyEalt is much more soluble in water than the fluoroborate and gives no precipitate with nitron acetate An aqueous solution hydrolyses in the presence of alkali at room temperature:- KBF:5 011 + 3KP + KB02 + 2H20 whereas the fluoroborate only slowly hydrolises

SPECTRA IN PAR I911 KBP4 NABP4

526 S2Ga NaBFell IMF 3 OH

59 0 590 526 -19- The differences between the spectra of sodium and potassium fluoroborate in the far infrared have aroused a lot of interest. Cote and Thompson (1951) quote the following values for the two salts in this region as being:-

KBFd No SIT 521 cm. 518 cmo ) Cote and Thompson ) (1951) 53 522 ) ) 551 ) 520 cm, 521 cm° _m ) ) Present work 532 526 ) ) 549 )

The fluoroborate ion is regarded as being tetrahedral and has therefore nine normal modes which include a single vibration (V1 ), one twofold degenerate vibration (V2) and two, threefold degenerate, vibrations V3 and IT,4, V1 - the symmetrical breathing B-F

V3 - the degenerate B-F stretching oscillation

VV2)) - the deformational motions. The differences in the sodium and potassium fluoroborate spectra have been explained by Cote and Thompson (1951) on -20-7 the hypothesis that the degeneracy of the vibrations Vs V4 are removed by the effect of the crystal field. The crystal structures of ammonium and potassium fluoro- i borates have been shown to be D2h (Hoard and Blair 1935) whilst Klinkenberg (1937) suggests that the class of sodium 17 fluoroborate is Dza. It is also known that ammonium, potassium and rubidium perchlorates have the BaS06 structure whilst sodium perchlorate is like anhydrite (CaS0,1 ). By analogy we expect sodium monohydroxyfluore- borate to be different from the potassium salt and X-ray and infrared data would seem to confirm this. The differences which occur with the sodium salts seem to be a general effect connected with the smaller size of its ion. In effect the site symmetry of the BF4 7 BF3 0H- and C10' ions in the sodium salts may be different from those in the potassium and ammonium szlits. In conclusion we can say that potassium and sodium monohydroxyfluoroborates do exist and have been characterised by both infrared and X-ray powder photograph data. -21- Tetra Alkyl Ammonium Monohydroxyfluoroborates It has been shown that the maximum quantity of boron trifluoride absorbed per mole of alcohol is always very nearly 1 mole of the gas at atmospheric pressure (Gasselin 1894). O'Leary,Wenske (1933) prepared mercuric methoxytrifluoroborate from the alcoholate and mercuric oxide but were unable to prepare the sodium salt. From a study of the vapour pressure curves of the BF3.Me0H system it was concluded that a definite equilibrium exists:- BF3 Me0H -4===3 OH3 OH BF3 A maximum in the curve for a dilute solution was taken as an indication of the existence of a constant boiling mixture. The conductivity curves for concentrated solutions (Nieuwland, Vogt, Poohey 1930) had a daarp break in the vicinity of 50 mole % boron trifluoride again showing compound formation as above. Other attempts at forming salts of the [BF3OCH3]- ion by O'Leary and Wenske (1933) proved abortive and indeed the Hg2+ salt was found to be unstable readily decomposing, to mercuric fluoride. Measurements on the activation energies of ionic migration (Ep) and viscous flow (E1) of the mono- and dimethyl- alcoholates of boron trifluoride by Greenwood and Martin -22- (1953a) indicate that the values of Ell and El are similar in both cases. These results suggest that similar configurational changes occur in both the boron trifluoride complexes and hence that the sizes of the cation and anion are similar. Where Erb and Ep differ by large amounts conduction is unipolar and occurs by the rapid migration of small cations through the spaces in the larger anion network. Table 1 shows clearly the large differences between El and Ep for some halides (Greenwood and Martin 1953b).

TABLE 1 mop. Ep E9 Ery,x NaCl 800° 2070 9.4 305 KCI 768 3.26 7.8 2.4 NaBr 740 2058 10.6 41

Electrolysis studies (Greenwood and Martin 1953a) on the alcoholates provide results that establish that in the 1:2 complex the ions are of similar size as [CH3 OH2] and [BF3 0CHu r and that the 1:1 complex should be regarded as methoxytrifluoroboric acids le[BF3 00113 ]-, solvated by a further molecule. -23- 2[131'3.01130H] [BF3CH3OH2] [BP30CH3 r

Tetramethylammonium monohydroxyfluoroborate has been described (Wheeler, Beaulieu and Burns 1954) as being the result of the reaction of methoxyfluoroboric acid on the tetramethylammonium halides. The same product is obtained by titrating monohydroxyfluoroboric acid with a 10% aqueous solution of tetramethylammonium hydroxide. Identical X-ray patterns were obtained for the prodacts obtained by both methods, m.p. 414°. The corresponding tetrafluoroborate (Wheeler and Sandstedt 1955a) has been prepared from tetramethylammonium chloride and boroa trifluoride etherate using the general method of Schlesinger et al.(1953). 4(0113)20.BF3 3(CH)4 NCI-->3(CH3 )NBP4 30H3C1 B(0CH3)3 + (CH3)20. The same product, m.p. 418° 4- 1°, was obtained, but in a reduced yield, with the bromide. Similar reactions with the tetraethyl, tetra-n-propyl and tetra-n-butyl, ammonium halides yielded the appropriate fluoroborates. The n-butyl salt was also prepared by Witschonke and Kraus (1947) by the interaction of ammonium fluoroborate and tetra n= butylammonium hydroxide, mop. 161-168°. Wheeler and Sandstedt (1955b) report a value of 155°. The same series -24- of aalts have also been prepared by the reaction of the alkylammonium halides and hydroxides with monohydroxy® or methoxyfluoroboric acid (Wheeler and Sandetedt 1955b) In view of:the fact that the previous authors had only been dole to isolate one monohydroxyfluoroborate complex: namely tetramethylammonium monohydroxyfluoroborate: it was decided to repeat the work. Accordingly tetramethylammonium hydroxyfluoroborate was prepared following the methods of Wheeler: Beaulieu and Burns (1954) by the action of (a) methoxyfluoroboric acid on tetramethylammonium chloride (sample 1); (b) hydroxyfluoroboric acid on the chloride (sample 2). Tetramethylammonium fluoroborate was also prepared using 40% fluoroboric acid. Other substituted quaternary ammonium fluoroborates were prepared from Et4NX;(1.-Bu)4NX; (n-:%.)4 NX (where X is 01, Br or I) using both 40% fluoroboric acid and methoxyfluoroboric acid. It was found that the samples 1 and 2 of the monohydroxyfluoroborate gave identical X-ray powder patterns to that of the sample of tetramethylammonium fluoroborate. These in turn were identical to the powder pattern obtained by Wheeler Beaulieu and Burns (1954) for tetramethylammonium monohydroxy fluoroborate. A comparison of the 'd' values for the photographs are shown in table 2.

_25_

TAhLt_a;..

(C16)411BP3OH (C1:15)4NBF3OH (CH3)417B1130H (CH5)4NBF4 Wheelers Sample 1 Sample 2 Beaulieu and Burns 1954.

5.80 (2) 5.92 (3) 5.82 (6) 5.87 (4)

rn 5.04 (1) 4.78 (4) 4.83 (2) 4.83 (10) 4.78 (6) 4.61 (1) 4.63 (1) 4.58 (11) 4.12 (10) 14.12 (10) 4.08 (100) 4.06 (10) 3.74 (1) 319 (1) 3.13 (1) 313 (20) 3.11 (7) 2.94 (1) 2090 (1) 2.77 (3) 2.79 (5) 2.76 (10) 2.77 (3) 2.64 (1) 2.61 (1) 2.49 (5) 2.50 (8) 2.48 (40) 2.41 (8) 2029 (2) 2.31 2.29 (20) 2.29 (5) 2.21 (6) 2.13 (1) 2.13 (4) d 2013 (3) 1.94 (1) 1.949 (5) 1.93 (20) 1.940 (4) 1865 (1) 1.85 (1) 1.852 (1) 1.655 (3) 1.64 (18) 1.646 (3) 1.563 (1) 1.55 (3) 1.558 (1) 1.467 (1)

1.36 (1) Ca. The figures in parenthesis are visual estimates of intensities. The strongest line being taken as 10; in the data of Wheelers Beaulieu and Burns the strongest line is taken as 1000 The fact that the X-ray patterns are identical does not offer definite proof that tetramethylammonium mono- hydrofluoroborate cannot be prepared by these methods since the sizes of the 0H- ion and F- ion are similar and therefore the two compounds ABF. and ABF3OH could have almost identical powder patterns assuming that there is no distortion due to hydrogen bonding. Analytical data is not decisive as the carbon and hydrogen are practically the same for both ABF4 and ABF3OH. Theoretical values fOr (C1-15) 6101130H;3 CD30019(A; ii98018A B.6.92%; F935-841 (C115)4NBF6 ; 0929.84%; 1197.4.5%; B96.83; F947.20%.

Wheelery Burns and Beaulieu (1950 support their work with the following analytical data for (0113)4NBP5OH O v29.63LA H97.99,4 B97.30vA; F935.36°A 29,93k 7980%; 6.901; 35963%

Apart from the fluorine figures the values could apply to (CHE)4NBF4. However fluo-rine analyses are always difficult to do with any accuracy (particularly in the presence of -27- boron)p and not too much reliance can be placed on their validity. The reported melting points are also very close 414° E(C113)4NBFOH3 and 1x.18° k(0115)4NBF4]. As further proof the infrared spectrum of all the alkylammonium salts were taken. The samples of g(C113)4NBPOH9 showed no evidence at all of hydroxyl bands in the normal regions whereas potassium and sodium monohydroxyfluoroborate (as already shown) do have such bands. In all the spectra of the alkylammonium fluoroborates studied there is no splitting of the main bands. From this study we conclude that monohydroxyfluoroborates of the alkylammonium salts have not been prepared so far. The evidence obtained shows that tetrafluoro- borates are formed in reactions involving fluoroboric9 monohydroxyfluoroboric and methoxyfluoroboric acids.

Some Although West (1935) has stated that the fluoro- 2+ a+ c 2 borates M(BF4)Q.6H20 (M .-73 Mgt+fl Mn Fe Co ' Ni ; 2+ a+ Zn Oci ) are isomorphous with the corresponding perchloratesp none of their lattice constants appear to have been recorded. Consequently the hydrates were -28- prepared and photographed in the prebent work. The lattice constants and the measured and calculated densities are recorded in Table 3, where they are compared with the values obtained by West (1935) for 2+ 51+ 24. 2+ u+ the perchlorates. The Mg Mn Fe v CO Ni and a+ Zn salts are hexagonal and are very similar in size to the corresponding perchlorates; the cadmium salts have a closely related trigonal structure with a one half of that shown in Table 3, the true value is doubled for comparison with the other salts. West (1935) has shown that the copper salts are not isomorphous with the other divalent salts and the present work has confirmed this. Because of the Jahn-Teller effect one would expect the crystal structure of the copper salt to be different. It is found that in almost all cases, for octahedral 2+ CU compounds, the octahedron has four equal hinds and two longer. Lithium fluoroborate exists in at least two forms, LiBF4.11209 exists above 23° and is tetragonal a = 5.74g ; o = 4.88g. LiBF6.31.120 crystallises from aqueous solution, beloW 23°, and is hexagonal isomorphous with the corresponding perchlorate according to West (1935). -29- The only phase that we could crystallise from solution is hexagonal9 a = 9.90A°9 c = 5.53e 9 but it is not isomorphous with the perchlorate trihydrate.

Table M21-(C104)a'.6Hm0 (West 1935) C measured calculated density density

Mg 15.52e 5.26e 1.98/ /099 Mn 15.70 5.30 2.102 2.10 Fe 15.58 5.28 2.147 2.17 Co 15.52 5.20 2.198 2.22 Ni 15.46 5.17 2.252 2025 Zn 15.52 5.20 2.252 2.26 Cd 15,92 5.30 2.368 2.38

M2+(BFj2..6H20 m21+ a C measured calculated density density Mg 15.36A° 5.38A? 1.849 1.85 Mn 15.46 5.44 1982 1.98 Fe 15.49 5.33 2.038 2002 Co 15033 5.22 2.081 2.11 Ni 1532 5.16 2.136 2.16 Zn 15.24 -5.30 2.120 2.16 Cd 15.96 + 5.58 2.292 2.12 Prepared and indexed by Russell (1959). aa2tia value doubled. -30- Silver Pluoroborate. Silver fluoroborate was first prepared in an anhydrous state by Sharpe (1952) from silver borate and bromine trifluoride. Sharp (1957) indexed the salt on an orthorhombic lattice (a = 7.75Ie s b = 6.542; c m 7.16/e). This salt has been prepared, in the present works from a solution of silver fluoride in anhydrous hydrogen fluoride under a pressure of boron trifluoride. Clifford and Kongpricha (1957) have prepared silver fluoroborate from the same reactants at atmospheric pressure. The X-ray powder photograph of the white solid, obtained by removing the solvent, was indexed on an orthorhombic system with lattice constants the same as those obtained by Sharp. Clifford, Beachell and Jack (1957) observed that a solution of fluoroboric acid gave an insoluble blue compound in the reaction with argentic fluoride. In the present work it was found that argentic fluoride in anhydrous hydrogen fluoride under a pressure of boron trifluoride gave silver(l)fluoroborate and not the silver(II) salt. -31- EXPERIMENTAL

The X-ray powder photographs described in this chapter were taken using a.9 cm, powder camera. The radiation used is indicated at the top of each set of results. The samples were packed in lithium borate, Lindeman tubes. As nearly all the samples were hygroscopic packing, of necessity, was carried out in the dry box. All calculations and indexing were done on a Ferranti 'Sirius' computer, using programmes compiled by Mr. D. R. Russell. Infrared spectra in the sodium chloride region were taken on a Grubb-Parsons, model S4, infrared spectrometer. Mulls in nujol and hexachlorobutadiene were used. For spectra in the far infrared a Grubb Parsons DM2 spectrometer was used with a caesium iodide prism. The samples were taken in nujol mulls on silver chloride plates. Sodium monohydroxyfluoroborate. Sodium monohydroxyfluoroborate was prepared by mixing sodium hydrogen fluoride and boric acid in 2:1 molar quantities at 0° with 50 ml. water per mole of -32- boric acid used (Ryss and Slutskaya 1952). After leaving for 2 hours at 0° (with occasional stirring) the solution was filtered and the sodium salt precipitated by the addition of a four-fold volume of alcohol. After thorough washing with alcohol the white solid was dried over phosphorus pentoxide in a vacuum. Polythene apparatus was used to prevent fluorosilicate impurities. Sodium hydrogen fluoride was freshly prepared from sodium peroxide and 40% aqueous fluoric acid in a platinum dish. Potassium monohydroxvfluoroborate and fluoroborate. Potassium monohydroxyfluoroborate was prepared by using the method of Ryss (1946) and Wamser (1948) again using polythene apparatus. Powdered boric acid was dissolved in a concentrated solution of potassium hydrogen fluoride at 0°. The salt was precipitated out with alcohol. Potassium fluoroborate was precipitated by neutralising 40% fluoric acid with aqueous potassium hydroxide and dried at 130°. Methoxyfluoroboric acid. The acid was prepared using the method of Greenwood and Martin (1953a). A cylinder of boron trifluoride was supplied by the Imperial Smelting Company. The gas was purified by passing it through a mixture of concentrated sulphuric acid and boric acid to remove hydrogen fluoride, and it was then bubbled through a layer of mercury into pure anhydrous methanol until it was saturated. The methoxyfluoroboric acid do formed was purified by fractional crystallisation (mop. Boron trifluoride dihydrate. The dihydrate was made by passing 1 mole of the purified gas into 2 moles of water cooled in an ice-bath (McGrath, Stack and McCusker 1944). At first absorption was slow but soon became rapid. When approximately 0.25 mole of the gas had been absorbed boric oxide precipitated out but this redissolved as the absorption of boron trifluoride continued. The clear9 non-fuming, hydrate was purified by fractional crystallisation (m0p0 'Tetramethylammonium hydroxytrifluoroborate'. Sample I - was prepared by the addition of methoxyfluoroboric acid to solid tetramethylammonium chloride. The compound was precipitated by the addition of ether followed by recrystallisation from water and drying at 110°.

34- (CH3 )4 NBF3 OH Cale. C9 30.19%; H98.18% Found: 30.82%; 7.90% (CHANBF4 Calc. 29.84%; 7.45%

Sample II - prepared from borontrifluoride dihydrate and tetramethylammonium chloride was purified as for Sample I. Analysis: 09 30.24%; H9 7.77%. Tetramethylammonium fluoroborate was prepared by the addition of 40% fluoroboric acid with tetramethylammonium chloride. It was recrystallised from water.

(CH3 )4 NBF4 Cale. 09 29.84%; H9 7.50% Found: 30.48%; 7.83%

Other alkylammonium salts. The preparationis of (C2110 )410F4; (n-C3117)4 NBF4 and (n-C4130)4 BF4 were effected in two ways:- (1)by addition of 40% fluoroboric acid to the solid halide9 (2)by addition of methoxyfluoroboric acid to the halide. The reactions proceeded quickly at 20° and on addition of ether the salts precipitated out. -35- Recrysta Compound from mop. Found Literature mop.

(CH3)4 NBP3 OH water IMO 414° 7heeler9 Beaulieu and Burns 1954 (CHs NB,ct water - 418+1° Wheeler and Sandstedt 1955a (C2B04NBF4 aq.methanol 364° 365367° Wheeler )and )Sandstedt )1955b ) (n-C3117)4 NBF4 methanol 248° 250© ) (n-C4110)4NBF4 50% methanol 161° . 155155°) 161.8° Witschonke and Kraus 1947.

Infrared Spectra.

KBF3 OH (nujol) 515(8)9 520(8)9 690(m )9 750(w)9 795(s) 830(m)9 900 970(b)9 1092 1155(b)9 1400(w)9 1630(m)9 1650(m)9 2300(w)9 3120(m).

KBF3OH (hexachlorobutadiene) 687( ), 755(w)9 900 970(b), 1090 1155(b)9 1400(s)9 1625(w)9 2300(09 2860(m)9 3130(s). NaBF3 OH(nujol) 520(w), 570(s), 771(8)9 803(8)9 900 967(b) 1080 1160(b), 1210(m), 1650(w), 3425(m), 3550(s). -36

NaBF3 OH (h.c. b. d) 770(m), 900 1250(b)9 290000 2 3125(m), 3550(s) . KBF4 521, 534, 771(s), 1032(m) , 1058(m), 1072(s), 1302(0'9 2148(w). EaBF4 518 522, 5519 770 779( 8) p 1036(m), 1058(m), 1076(s), 1305(m), 1332(s)

1( CH3 )4BF3 OH ' I and (CH3 )4NBF3 OH - II, ( nuj ol ) 521(s),533(09 665(w) 763(09 945( )9 1025 1050(b), 1285(m), 1400(09 1410(09 h.o.b.d. 1030 1075(b), 1285(m), 1400(m), 1425(09 1480(m). (CH3 )4NBF4 (nujol) 521(09 533(8)9 665(w), 763(w), 945(s), 1025 1050(b), 128300, 1400(m), 1420(0 h0c.b.d. 1030 1075(b) .1285(m), 1400(m), 1425(09 1480(0.

(C2H5 )4NBF4 . 521 ( s)9 533( a)9 790(s)9 1015 1050(b)9 1100(m), 1185(w) , 1282009 /307(09 1370(0 , 1400(09 1455(w), 1467(m), 1790(w). -37- 'Me41\113110 OH I Co radiation

Intensity sine 9 2 0.0238 5.80 4 0.0350 4.78 1 0.0376 4.61 10 0.0470 4.12 1 0.0768 3.19 2 0.0819 3.13 3 0.1040 2.77 5 0.1285 2.49 2 0.1523 2.29 1 0.1764 2.13 1 0.2132 1.94

91.1e4 NBIPz OH' - II Co radiation

Intensity sin2 9 d 3 0.0228 5.92 1 0.0315 50 04 8 0.0342 4.83 1 0.0373 4.63 10 0.0472 4.12 6 0.0810 - 3.74 1 0.0928 2.94 -38- Intensity sine 0 d Co radiation 5 0.1026 2.79 1 0.1142 2.64 8 0.1279 2.50 6 0.1496 2.31. 4 0.1755 2.13 5 0.2107 1.949 1 0.2301 1.865 3 0.2921 1.655 1 0.3275 1.563 MS4NBF4 Co radiation Intensity sin2e d 4 0.0232 5.87 6 0.0351 4.78 1 0.0382 4.58 10 0.0485 4.06 7 0.0829 3.11 3 0.1041 2.77 8 0.1375 2.41 5 0.1526 2.29 3 0.1720 2.13 4 0.2126 1.940 -39- Intensity 13in20 Co radiation 1 0.2332 1.852 3 0.2954 1.646 1 0.3294 1.558 0.3718 1.467 nBuzliTP,6 Co radiation

Intensity sin2 0 d 8 0.0098 9.04 6 0.0314 5.05 10 0.0525 3.90 3 0.0609 3.62 2 0.077? 3.21 1 0.1573 2025 1 0.1746 2014 2 0.2049 1.98 Et4NB194.. -40- Co radiation Intensity. ein28 d 9 0.0231 5,89 10 mina 3.15 8 0.0820 3.12 6 0.0964 2.88 5 0.1062 2.74 3 0.1175 2.61 4 0.1244 2.54 7 0.1516 2.30 9 0.1605 2.23 2 0.1695 2.17 8 0.1815 2,10 1 0.1933 2.03 2 0.2039 1.98 4 0.2795 1.69 3 0.2852 1.67 2 0.3159 1.59 1 0.3327 1.551 1 0.3492 1.514 1 0.3664 1.478 1 0.3970 1.420 3 0.4101 1.397 1 0.4536 1.328 3 0.5828 1.172

-41- WaBP3 OH. C4 radiation d Intensity sinaG(obs.) (hak,l) sin20(calc.) 3.18 8 0.0590 002 0.0596 2.86 7 0.0728 102 0.0724 2.81 7 0.0751 2.66 4 0.0841 2.24 10 0.1177 300 0.1152 2.222 9 0.1204 2.186 3 0.1243 212 0.1235 2.128 6 0.1312 301 0.1301 2.026 4 0.1447 311 0.1428 2.001 4 0.1484 103 0.1469 1.975 2 0.1524 - 1.801 2 0.1832. 023 0.1849 . , 1.736 1 0.1973 123 0.1977 1.685 1 0.2093 1.600 1 0.2323 1.555 1 0.2457 333. 0.2444 1.339 1 0.3312 150 0.3303 NaBF3 OH was indexed ao an orthorhombic lattice. A = 0.01287 + 0.00005 B = 0.01276 4- 0.00005 C = 0.01494 41- 0.00006 a = 6.80; b = 6.83; c = 6.31 -42- KME3 OH Oi radiation d Intensity sin20(obs..) (h,1E4) sin20(calc.) 4,53 1 0.0290 011 0.0290 4.21 1 0.C635 - 4.02 1 0,0368 200 0.0304 3.70 4 0.0436 020 0.0440 3.37 10 0.0525 120 0.0536 3.13 1 0.0606 021 0.0620 2.96 8 0.0676 211 0.0674 2085 10 0.0731 002 0.0720 2.66 5 0.0838 012 0.0830 2,39 1 0.1038 301 0.1044 2.25 3 0.1172 031 0.1170 2017 1 0.1264 131 0.1266 2.085 2 0.1367 230 0.1374 1.878 1 0.1684 312 0.1694 1.739 1 0.1965 420 0.1976 KM0P3 OH was indexed on an orthorhombic system. A = 0.00965 + 0.00002 B = 0.01101 ± 0.00014 C = 0.01815 + 0.00028 a = 7985A); b = 7.35N; c = 5.72AP -43- KBP4 Cu radiation d Intensity sin20(obs.) (h,k,l) sin2 O(ca1c.) 3.01 7 0.0655 112. 0.0677 2.85 4 0.0732 200 0.0736 2.64 8 0.0856 003 0.0864 2.43 5 001010 122 001004 2.35 1 04080 031. 0.107? 2.10 4 00.344 .0.32 0.1365 1.98 3 0.1510 - 1088 4 .0.1675 300 0.1656 1.797 10 0.1841 041 0.1840 1.772 8 001892 1.728 3 0.1990 024 0.1972 1.559 2 0.2445 1.516 1 0.2586 105 0.2584 1.368 4 0.3177 304 0.3192 1.331 1 0.3355 243 0.3344 1.300 3 0.3515 333 0.3501 1.196 2 0.4157 315 0.4165 1.175 3 0.4304 216 0.4301 1.114 1 0.4787 063 0.4788 1.065 2 0.5244 353 0.5245 1.031 2 0.5592 530 0.5581 0.951 1 0.6569 524 0.6572 0.940 1 0.6733 610 0.6733 0.929 1 0.6882 208 0.6880 0.922 1 0.6994 370 006997 KBP4 indexed on an orthorhombic system. A = 0.0184 B = 0.0109 C = 0.0096 a = 5.68; 13= 7.37; c = 7.85 -44- Hydrated fluoroborates

The hexahydrates were prepared by adding the appropriate carbonate to 40 fluoroboric acid in a platinum dish. When excess carbonate had been added the solution was warmed until the reaction was complete and then filtered. After concentrating on a steam-bath until crystallisation occurred the crystals were dried over concentrated sulphuric acid in a vacuum des04tor. The lithium trihydrate was obtained by pumping off the excess water from the solution under vacuum at 18°. When crystallisation occurred the trihydrate was dried c. over concentrated sulphuric acid in a vacuum des04tor at 19° for four weeks. The salts were subsequently handled in a dry-box in preparing the X.ray samples. -45- LiBF4 .H2 0 Cu radiation dh,ka intensity sin20(obs.) (11.4,1,) sin20(ca1c.) 4.860 6 0.0250 001 0.0250 3.375 9 0.0621 - 3.191 10 0..0581 - .-

2.799 1 0.0755 - 0=1,

2.689 1 0.0817 e=1 2.571 4 0.0894 210 0.0900 2.375 10 0.1049 211 0.1150 2.258 3 0.1159 - 2.014 9 0.1427 220 0.1440 1.875 2 0.1688 221 0.1690 1.811 5 0.1802 310 0.1800 1.777 4 0.1871 301 0.1870 1.733 4 0.1967

1.670 8 0.2118 OCD 1.588 6 0.2342 320 0.2340 1.453 3 0.2797 312 0.2800 1.431 3 0.2887 400 0.2800 1.409 5 0.2977 203 0.2970 10388 2 0.3068 410 0.3060

1.351 3 0.3236 L=. -46- dLatat t i Intensity sin2 e( alma.) (201,1) Edna e( c alc . ) 1.298 3 0.3490 331 0.3490 1.277 3 0.3622 420 0.3600 1.219 5 0.3976 004 0.4000 1.192 2 0.4159 104 0.4180 1.174 2 0.4290 1.124 2 0.4681 510 0.4680 1080 1 0.5068 1.064 1 0.5222 520 0.5220 1.057 1 0.5684 512 0.5680 1.020 1. 0.5750 440 0.5760 1.014 2 0.5963 - - 0.996 2 0.6027 441 0.6010 00990 2 0.6135 530 0.6120 00981 1 0.6509 0.9528 1 0.6576 o 0.9479 1 0.6670 610 0.6660 0.9412 20.68800.6862 404 0.9280 2 0. 7092 414 O. 7060 O. 9136 1 0.7245 334 0.7240 0, 9031 1 0.7450 621 O. 7450

0.8906 1 0. 7532 61511 -47- d Thtensit_y e( obs.) (11 4,1) sin a(c 11,k ,1 1 3 0.8857 1 0.7829 0.8679 1 0.7915 0.8640 1 0.8238 0.8464 2 008379 533 0.8370 0.8398 1 008654 - 0.8256 2 0.8886 613 0.8910

Most lines can be indexed on a tetragonal lattice; these are due to LiBF4 .112 0 which is the stable form above 23°C0

A = 0.01802 + 0.00004 C = 0.02496 + 0.00027 where sine 0 = A(h2 410) + 012

a = 5.74 + 0.03A° A= 4a2 o = 4.88 4- 0.03A° C j41.c2 -48 - LiBF4.3H20 Cu radiation d h,k,1 Intensity sin20(obs.) (h,k.11 sin20(calc.) 4.312 8 0.0319 200 0.0324 4.003 8 0.0391 3.690 4 0.0436 111 0.0438 3.383 10 0.0519 201 0.0519 2.479 2 0.0966 220 0.0971 2.32/ 5 0.1101 202 0.1104 2.259 6 0.1162 221 0.1166 1.993 5 0.1491 401 0.1489 1.840 3 0.1751 222 0.1751 1.691 2 0.2075 402 0.2074 1.644 1 0.2194 330 0.2184 1.616 1 0.2269 420 0.2265 1.600 1 0.2317 322 0.2317 1.550 1 0.2467 421 0.2460 1.534 3 0.2518 510 0.2508 (511 (0.2703 1.477 2 0.2715 (223 (0.2726 1.393 1 0.3052 422 0.3049 1.383 1 0.3100 - - LiBF4.31120 was indexed on the hexagonal system. A = 0.00809 + 2) where sin2 0 = A(h24-10) + C12 C = 0.01945 4.- 4) a= 9.90+ 0.04; c= 5053+ 0.04 Density = 1.05 -49- Mg(BF4)2.6H20 Cu radiation

h,k,1 Intensity sin2G(obs.) (hek.1) sin2G(calc.) 5.039 5 0.0232 210 0.0246 4.183 10 6.0338 300 0.0316 3.811 10 0.0407 220 0.0421 3.425 3 0.0504 - 3.034 1 0.0642 320 0.0667 2.820 10 0.0743 410 0.0737 2069 3 0.0812 002 0.0820 2.495 3 0.0949 330 0.0948 2.398 1 0.1028 510 0.1026 2.266 10 6.1151 2.190 2 0.1232 2.091 1. 0.1351 520 0.1369 2.016 1 0.1453 1.964 1 0.1532 610 0.1509

1.905 2 0.1628 !SO

1.831 7 0.1763 Goa 1.783 1 0.1859 003 0.1845

1.707 5 0.2027 cp 1.655 1 0.2158 540 0.2141 1.614 1 002258 800 0.2246 -5C- dh,k,1 alItnEkty sin26(obs.) (lak.1) sincplcalc.) 1.583 2 0.2359 720 0.2352 (810 (0.2562 1.516 2 0.2571 (413 (0.2582 1.440 1 0.2851 900 0.2843 1.411 1 0.2969 542 0.2961 1.393 1 0.3048 901 0c.3048 1.270 2 0.3663 902 0.3663

Mg(BF4)2.6H20 was indexed on a hexagonal system. A = 0.00351 + 0.000024 C = 0.02047 + 0.00015 a = 15.36 0.05A9g c = 5.38 + 0.02AP. Density = 1.85 -51_ Mn(BP4)2.6H20 Cr radiation dhok 1 Intensity sin20(obs.) (h,k_q1) sin20(calc.) 4.82 1 0.0285 200 0.0287 4.67 2 0060 4043 4 0.0668 - 4.19 5 0.0744 3.88 9 0.0371 220 0.0878 3.69 8 0.0960 211 0.0956 3.44 1 0.1105 301 0.1103 3.33 2 0.1182 400 0.1171 3.07 3 0.1388 320, 0.1391 2084 10 0.1617 401 0.1615 2.69 1 0.1810 500 0.1830 2.29 4 0.2502 421 0.2494 2.205 3 0.2689 430 0.2708 2.093 1 0.2985 2.040 1 0.3141 610 0.3148

1.937 1 0.3484 tem

1.860 2 0.3780 cora

1.842 2 0.3855 Cr.

1.716 3 0.4440 CPO OP Mn(BF4)2.6H20 can be indexed on the hexagonal system. A = 0.00732 + 0.00001) sine 6 = A (h2+kh+k2) + C12 C = 0.0444 + 0.0006 ) a = 15.46 +0.01A° a2 >2,,r3A 2 c = 5.44 + 0.09A° Cr = Calculated Density = 1.98. -52- Fe(BP6)2.6H20. Co radiation 4. Intensity sin20(obs.) (h„lc,1) sin'O(calc.) 5.17 8 0.0300 210 0.0312 4.18 9 0.0460 201 0.0460 3.83 10 0.0539 220 000534 3.68 2 0.0590 - 3.47 7 0.0666 301 0.0683 3.07 6 0.0847 320 0.0846 2.82 10 0.1008 401 0.0994 2.70 2 0.1100 500 0.1113 2.28 8 0.1548 302 0.1529 2.20 6 0.1650 430 0.1647 2.095 3 0.1826 402 0.1840 2.031 2 0.1943 431 0.1929 (521 (0.2018 1.982 1 0.2040 (412 (0.2063 1 0.2162 0 0.218/ 1.924 53 1.839 8 02371 422 02 374 (701 1 .798 1 0.2480 (531 0.2463 1.712 5 0.2733 602 0.2730 1.526 2 0.3443 622 0.3342 1.177 2 0.5782 813 0.5787 Po(BF4)2.61120 was indexed on a hexagonal lattice. A = M0445 C = 0.0282 4. 0.0006 a= 15.49g c = 5.33g Calculated density = 2.02. Ni(BF4 )2 . 6H2 0 Cu 'radiation d Intensity alEIL011,1 ( h Oc j ) sin?. 0( calc . ) 2.61 3 O. 0872 002 0.0896 2025 8 0.1171 - 2014 4 0.1274 520 0.1318 2.03 4 0.1438 610 0.1453 530 1.90 2 0.1642 700 0.1656 1.83 5 0.1766 620 0.1788 1.80 5 0.1841 - - 1.73 1 0.2000 003 0.2016 1.68 10 0.2114 630 0.2129 1.598 2 0.2327 -

1.571 1 002407 OM. 1.519 1 0. 2574 640 0.2569 1.442 2 0.2860 820 0.2839 1.390. 5 0.3077 650 0.3076 1.269 7 0.3688 750 0.3684 1.223 1 0.3970 930 0.3955 1.175 2 0.4302 760 0.4293 1.157 1 0.4442 1.078 1 0.5110 950 0.5104

1.057 2 0.5320 1=1.1. -h4- d Intensity sine @I (11.401) sin20(calc.) 1.037 1 0.5532 11,390 0.5509 1.012 1 005797 996,0 0.5780 10000 1 0.5944 0.963 1 0.6409 129 390 0.6388

Ni(BF4)2.6H20 was indexed on a hexagonal system.

A = 0.003388 0.00002 C = 0.0224 4,- 0.0002

a = 15.32 0003A° c = 5.16 mate Calculated density = 2016 -55- Zn( BP,/ )2 .6H2 0 Cu radiation

d Intensity sin2 e( obs.) (h,,k ,1) sine G(21.1c 0 )

6.69 8 0.0133 200 0.0136 4094 7 0.0244 210 0.0239 4.41 5 0.0306 300 0.0307 4.16 7 0.0344 201 0.0348 3.85 10 0.0401 220 0.0409 3.65 7 O. 0446 310 0.0433 3.41 3 0.0510 301 0.0519 3.04 4 0.0642 320 0.0648 2.82 10 0.0746 401 0.0758 2065 3 0.0848 002 0.0848 2.546 2 0.0917 330 0.0921 2.460 1 0.0982 202 0.0984 2.375 1 0.1054 510 0.1057 2.267 10 0.1156 302 0.1155 2.179 9 0.1252 222 0.1257 2.069 3 0.1388 402 0.1394 2.021 1 0.1456 610 0.1466 1.970 2 0.1531 521 0.1542 (530 1.917 4 0.1671 ( 700 0.1671 1.835 7 0.1764 332 0.1769 10762 1 001914 003 0.1908 -56- d Intensity 121.12_2LL2.1121 (119kPl) sin26(calc.) 1.699 5 0.2060 203 0.2044 1665 1 0.2143 213 0.2147 1.590 1 0.2351 313 0.2351 1.553 2 0.2464 403 0.2454 1.525 3 0.2556 232 0.2556 1.449 3 0.2830 333 0.2829 1.406 2 0.3006 632 0.2996 1.351 1 0.3255 523 0.3238 1.328 1 0.3372 613 0.3374 1.274 3 0.3660 623 0.3681 1.251 1 0.3797 224 0.3803 1.190 1 0.4194 723 0.4193 1.165 1 0.4380 813 0.4397 1.135 1 0.4614 851 0.4611 1089 1 0.5015 770 0.5013 1.072 2 0.5167 624 0.5765 /.009 1 0.5832 960 0.5831 1.000 1 0.5944 325 0.5948 0.988 1 0.6087 734 0.6086 Zn(BF4)26H20 was indexed in a hexagonal lattice. A = 0.00341 4- 0.00002 C = 0.02115 4. - 0.00012 a = 15.24 4- 0.04AP5 c = 5.30 + 0.02A° Calculated density = 2.16. -57- Cd(BF4 )2 .6H2 0 Cu radiation

d Intensity sin2 0( obs. ) (h ,lc 91) sin2 e( calc o )

5.66 10 0.0184 001 0.0191 3.23 6 0.0565 111 0.0566 2.62 9 0.0864 210 0.0875 2.38 6 0.1042 211 0.1066 1. 98 a 0.1515 220 0.1500 1. 90 8 0./629 301 0.1625 1.80 10 0.1820 311 0.1816 1.64 2 0.2210 203 0.2219 1.54 3 0.2495 1.44 3 0.2839 303 0.2844 1.355 4 0.3219 223 0.3219 1.314 4 0.3425 114 0.3431

1.240 5 0.3844 CCM. t 1.200 3 0.4096 323 0.4094 1.154 8 0.4437 -

1.118 1 0.4729 .1= 1.103 2 0.4860 520 0.4875 1.083 6 0.5037 404 0.5056 1.032 1 0.5546 611 0.5566 -58- d Intensity, sin2 Giobs. ) (11A21) s:in.lo . 0.993 1 0.5990 440 0.6000 0.957 3 0.6453 334 0.6431 0.943 2 006639 0.879 3 00 7640 540 0.7625

CdO3P. )2 0 61 2 0 was indexed on a hexagonal lattice.

A = 0.01246 ÷. 0.00009 C = 000191 0.0003

a = 7098 ± 0.032 c = 5.58 + 0.042 Calculated density = 2.12. -59- Co(BF4 )2 .6H2 0 (Russell 1960)

d Intensitz sin2 O(obso) (holt,l) ein2 G(calc0)

9.34 1 0.0091 425. 6.67 7 0.0181 002 0.0182 4.95 8 0.0327 101 0.0340 4.38 6 0.0417 003 0.0410 4.163 7 0.0462 102 0.0476 2.843 10 0.0532 3.651 - 7 0.0600 3.425 5 0.0683 103 0.0704 3.092 4 0.0836 3.047 3 0.0862 2.802 10 0.109 104 0.1022 2„)649 4 001139 005 0.1138 2.532 3 0.1247 - 2.471 1 0.1311 113 0.1292 2.377 1 0.1416 105 0.1432 2.263 9 0.1562 203 0.1586

2.185 8 0.1675 CV. 2.073 3 0.1862 2.017 3 0.1966 1.969 2 0.2064 210 0.2058 60- (h,Jr,l) sin20(calc.) 1913 3 0.2186 1.885 1 0.2251 212 0.2240 1.834 5 002376 - 1.820 5 (L2413 - 1.782 1 0'02518 107 0.2524 1.734 1 0;2658 300 0.2646 1.614 1 003068 303 0.3056 1.473 1 0.3691 216 0.3696 1.443 1 0.3839 310 0.3822 1.397 1 0.4106 208 0.4088

Co(BF4)206H20 indexed on a hexagonal system.

A = 0.00455 ± 1 C = 0.0294 A- 1 a = 15.53 ± 0.02; c = 5.22 4. 0.01 Calculated density = 2011 -61- Cu(BF4)2.6H20 Cu radiation d Intensity. sinee(obs.) 4.52 10 0.0291 4.16 8 0.0343 3.78 9 0.0416 3.20 10 0.0580 3.03 3 0.0646 2.93 1 0.0690 2.80 9 0.0756 2.71 1 0.0812 2.585 4 0.0890 2.471 2 0.0973 2.267 7 0.1396 2.195 .4 0.1277 2.123 3 0.1234 2.053 2 0.1319 1.974 2 0.1525 1.902 2 0.1624 1843 5 0.1750 1.768 1 0.1901 1.717 6 0.2015 _620 rjt Intensity taince(obso)

10634 1 0.2225 1.487 1 0.2689 1.377 1 0.3135 1.296 1 0.3690 -63- AMS2121111LSMS1.42=1MILJnallt02 Argentous fluoride was prepared using a modified method of Anderson9 Beck and HilIebert (1953). Silver carbonate was added in small quantities to 40A hydro- fluoric acid contained in a platinum dish. When excess solid had been added the solution was warmed on a steam bath to complete reaction. The solution was filtered into twice its volume of anhydrous methanol and was refiltered into a large volume of anhydrous ether. The bright yellow precipitate was quickly filtered and washed

with anhydrous ether and then pumped under vacuum at 600 for half an hour. The solid was stored in a dark bottle in the dry box. Argentic fluoride was prepared by fluorinating silver powder in a stream of fluorine at 3000 for one hour (Ruff and Giese 1934). A dark brown solid was obtained. Silvgr Fluoroborate 1 gm. of argentou6 or argentic fluoride was put in a. 25 m10 Bassett and Lindsey high pressure bomb (in the dry box). Approximately 10 ml. of anhydrous hydrogen fluoride was distilled into the bomb using the apparatus (A) in -64- diagram 1. Boron trifluoride was condensed in a U-trap cooled in liquid air and was then transferred to a vacuum line ard the gas was then redistilled into the bomb using apparatus B shod in diagram 1.

DTApRAM 1.

'Instantor' valve hexagonal nut

1/4 bore copper tubing To HF To bomb cylinder

To vacuum line

standard B14 socket To bomb (brass) -65- The bomb was rocked for 24 hours and then vented at 600 to remove most of the hydrogen fluoride. It was then evacuated at 60'4 for 1 hour on a normal vacuum line. In each case a white solid was obtained which was found to be AgBFG by analysis and X-ray. Analysis AgBP4 Found Ag = 55.1% (from AgF) Calculated Age 55.4% from AgtF2 Found Ag = 54.2% Agl3F4 (from AgF) d Intensity:cii0, . e h,k,i Sin.a cale. 4.2 1 0.0336 -- 3074 1 0.0425 - - 3025 8 0.0564 021 0.0563 315 6 0.0599 2094 8 0.0688 0 - 2079 2 0.0763 121 0.0757 2.70 8 0.0813 2.357 2 0.1069 103 0,1085 2.253 3 0.1170 202 001172 2.185 5 0.1245 220 0.124.0 2.126 7 0.1315 131 0.1337 2.032 10 0.1439 032 0.11440 1.971 10 0,1529 123 0.154.9 1.915 7 0.1620 132 0.163h. .742 3 0.1959 Oa 0.1955 1.608 1 0.2297 321 0.2309 1.550 1 0.2474 005 0.2475 1.488 1 0.2684. 105 0.2669 . 1.462 1 0.2782 330. 0.2790 1.419 1 0.2952 14.3 0.2941 -66- AgBloc (from AO's). Intensity Sin20 11 9k91 sine' (oale). 4.18 3 060341 3.74 2 0.0425 - - 3.48 2 000490 012 0.0512 3.38 2 0.0519 012 0.0512 3.22 5 (L0572 021 0.0563 3.12 5 0.0607 /02 0.0590 2.92 6 em0697 2.807 4 000754 121 0,0757 2.678 2 0.0828 2.342 1 001083 103 0.1085 2.241 1 0.1184 202 0.1172 2.175 2 0.1257 220 0.1240 2.121 10 0.1320 131 0.1337 2.080 1 0.1373 023 0.1355 2.022 10 0a454 032 0.1440 10971 8 0..1530- 123 0.1549 1.951 3 0.1562 123 0.1549 1.902 5 0.1643 222 0,1636 1.808 1 001.818 230 0.1820 1.738 1 0.1968 311 0.196/ 1.4.54 4 02810 134 0.2822 1.406 1 Q.3005 0 51 0.2999 10374 1 0.3150. 125 0.3133 1.356 1 0.3234 410 0.3220 0 - 7.75 .1 b = 6.54 = 7.16 A -67- CHAPTER

THE PREPARATION OF SOME lst ROW TRANSITION METAL SALTS OF STRONG ACIDS: Until fairly recently water has been considered unique in its solvent properties because of its ability to dissolve materials of such a wide variety of types9 structures andcompositiono. Without a doubt the availability, the physical properties and the ability to form solvates has encouraged much attention to be focussed on water as a ?uni- versal° solvent. The fact that Arrhenius? views on electrolytic dissociation were limited to aqueous solutions did, no douly%foster this idea. Ostwald went so far as to say that water was in a class by itself in so far as its ability to form electrolytic solutions or to bring about ionization. It raust9 however, be admitted that so far no other solvent has been found which approaches water in its solvent behav- iour. A realization that water is not unique in its characteristics during the last 60 years has helped enorm- ously in the development and understanding of the chemistry of solutions. The first intensive investigations into non- aqueous solvents were started just before the beginning of this century. Attention was first directed to alcohols9

-68- esters and ketones as .:solvents for inorganic compounds; this was followed by the studies on liquid ammonia aid sulphur dioxide. Solubility of salts in a solvent ie decided by how far the sum of the solvation energies of the ions compensate for the work done in breaking up the ionic lattice. If entropies are neglected this breakup can be represented thust- lia(solid) U (gas) gas) +L -H

M(solv.) + X-(solvo) +L = U o Ho where L is the heat of solution at infinite dilution, U is the lattice energy and H is the solvation energy of the gaseous ions. For mazy yearso a high di-electric constant was considered necessary if a solvent was to be a good solvent;since the electrostatic interactions of the ions will be greatly reduced by such a solvent. On this hypothesis the higher the dielectric constant the more likely the solvent to dissolve salts. This has now been realised as being only one aspect of the problem and the importance of the chemical aspects of solution are now beginning to be appreciated. The best solvents for compounds -69- ionised both in the solid aid solution states are nearly. always those which have unshared pairs of electrons, examples being,, water, ammonia and sulphur dioxide. Organic solvents such as ethers g ketones and amines show remarkable tenden- cies to dissolve salts although they have low dielectric constants; saturated hydrocarbons shOw-no comparable solvent properties and it is reasonable to infer.that the presence of lone paira of electrons is an important. factor. Por a salt to be soluble in etber, acetone or an amine the ideal .requirements are:- a very small cation (of high polarising power). or a. low.lattice,ene rgy which is found mainly with large anions. . Thus lithium perchlorate and :fluoroborate, sodium iodide and tetramethylammonium perchlorate are readily soluble in such solvents. Recently it has been shown that many silver fluoro- salts are soluble in unsaturated hydrocarbons, that is in solvents. with.< bonding systems e.g. silver fluoroborate is. soluble in benzene and toluene. However the sodium and pot-. assium salts are insoluble although they have similar ionic ..., radii (Agt1,0R; Na 0.9a; e 1.33/) ..from which one infers that the reason for the solubility of the silver sat must be more, complex. than a mere polarisation effect. There are a considerable number of compounds known in which the the 11- electrons of unsaturated hydrocarbon systems are used in bonding. Familiar examples are the cyclopentadienyl complexes ferrocene and cyclopentadienylmanganesetricarbonyl; ead dibenzene chromium. The silver salt and these examples all have one thing in common that the sodium and potassium salts do not possessand that is filled d orbitals. Bonding in these complexes must involve these d electrons. THE SOLUBILITY OF INORGAXIC SALTS IN AROMATIC HYDROCARBONS . The forces that hold the lattice of a covalent salt together are merely weak Van der Waals forces and solubility in aromatic hydrocarbons is clearly possible since the small solvation energy will compensate for the work done in destroying the lattice. Mercury salts of trifluoroacetie acid (Swarts 1939) and trichloroacetic acid (Davidson and Sutton 142) are both soluble in benzene. An examination of _ . the dipole moment favours a simple covalent structure in which the C13C.00 groups rotate around the 0-Hg-Hg-0 axis. Many silver salts of strong acids have been found to be soluble in.a fairly wide range of unsaturated hydrocarbons. Hill (1921) studied the solubility of silver perchlorate in benzene, toluene and chlorobenzene. The degree of electrolytic dissociation was found to be very small, of the same -71- order as that of water. The salt was found to be insoluble in chloroform and carbon tetrachloride. A stable complex AgC104.00H0 was isolated and a toluene complex AgC104.C?Ha is formed in the AgC104 C?Ha system (Spurgeon 19111). More recently a whole series of silver perchlorate-hydrocarbon complexes have been prepared by Peyronel et alia (1956) involving all the xylenes and some polynuclear hydrocarbons, all of these complexes are extremely hygroscopic. Olefin complexes have been characterised by Comyns and Lucas (1957). Silver trifluoroacetate and heptafluorobutyrate are soluble in aromatic hydrocarbons and complexes with benzene, toluene and the xylenes have been prepared (Woolf 1954; Swarts 1939; Tildesley and Sharpe 1954). Other silver salts soluble in benzene include AgSNCF3 (Haszeldine and Kidd 1954), silver trifluoromethylantimonate„ silver bistrifluoromethylphosphin- ate (Moss 1955) and AgO2C.COls which decomposes rapidly in solution with precipitation of silver chloride (Davidson end Sutton1942). Andrews, Keefer and co-workers have found evidence for the formation of complexes between silver salts and aromatic hydrocarbons. Particular attention has been focussed on the solubility of aromatic hydrocarbons and their alkyl derivatives in aqueous silver nitrate (Andrews and Keefer -72- 19k9c9 1950a9 1950b9 1952) evidence for the formation of ionic species AgAr7/. aid AgaAr+2 has been found and the order of stability of these complexes followS the order of the hydrocarbons when considered as eleotron donors. Lien Load McCaulay(1950a) have used this compleX formation as the basi6 of a commercial separation of the xylenes and ethylbenzene. ti vapour-liquid system is used; the liquid phase being an

ITIVA0 mixture. ,?epRration is simple as the figiAgF uomplex with ethylbenzene is much less stable than the others.

Silver fluoroborate monohydrate has been shown to be soluble in benzene and toluene by Warf (1952)9 who was also able to prepare a solution of silver fluoroborate by passing boron trifluoride through a suspension of silver fluokide in toluene. A similar preparation using benzene (Heyns and Paulsen 1960) resulted in the isolation of AgBF4008116. Isolation of the pure fluoroborate is possible by decomposing the complex at 50Q/1m.m. Olah and Winn (1960), using nitromethane have obtained the salt in precise- ly the same manner. Sharpe (l952) prepared the salt from silver borate and bromine triflueride and demonstrated its solubility in benzene; toluene and ether. Russell and Sharp (1961) have shown that silver fluoroborate and fluorophos- phate can be prepared by passing the appropriate gas through -73- a suspension of silver fluoride in liquid sulphur dioxide. In the present work it was found that silver fluoroborate will not dissolve in sulphur dioxide unless boron trifluoride is present. A fairly unstable complex seems to be formed in solution with S02/BF3„ which decomposes as sulphur dioxide evaporates. Sidgwick (1950) has pointed out that the solubil- ities and general properties of the hexafluorophosphates are very similar to those of the perchlorates. Sharp aild Sharpe (1956a) have shown this to extend to the solubility of Ag1T0 in aromatic hydrocarbons. The same authors have shown that AgS00;AgAsPe; AgSbFe, AgNbFEI and AgTaPs are also soluble in benzene, toluene and xylene; complexes with these hydrocarbons were isolated. Until the present work was started no silver salts containing two silver atoms per molecule had been shown to be soluble in aromatic hydrocarbons. How- ever, we have prepared silver hexafluorosilicate and hexafluorotitanate and have found that they are soluble, to a small extent, in benzene, toluene and ether. .As the lattice energy of salts with a divalent anion should be much greater than the lattice energies of salts with.a univalent anion one would expect the solubility to be less. If the solubility of silver salts in aromatic 74

hydrocarbons is due to the readily available frelectrons it is to be expected that silver ions would also form olefin derivatives; these are well known although mainly in aqueous solution. Ebers et alia (1937) have shown their existence .by partition studies as the complexes formed are very unstable. More recently Comyns and Lucas (1957) have prepared cyclo- hexene derivatives of silver perchlorate and nitrate; the former salt also reacts with the pinenes. Stable silver olefin complexes prepared at low pressure and room temperature from siver fluoroborate have been described by Quinn and Glew (1962). Perchlorate and hexafluoroantimonate complexes with propylene are also described, the order of stability being C10:4',BP4-.4 SbF6-. Raay and Schwenk (1959) have patented a process for the separation of gaseous olefin hydrocarbon mixtures using siver fluoroborate and fluorosilicate solutions. Cuprous olefin complexes are, on the whole, more stable than the silver ones. Addition compounds of cuprous chloride are formed with ethylene, propylene and iso-butylene under pressure (Tropsch aid Mattox 1935). Andrews and Keefer (1948, 1949a, 1949b) have prepared water soluble cuprous complexes with unsaturated acids and suggest that thestability of the copper bond appears to be dependent on the steric -75 influence of. groups attached to the double bond whereas steric factors do not appear significant with aromatics. The copper7acetylene complex of. Schreiber aid Reckleben(1911) is a true acetylide. Whaley and Starkey (1943) prepared what they thought was phenyl copper by refluxing aryldiazonium fluorobor4tes in benzene and in toluene with coppery but Warf (1957) on repeating this work obtained different results. When the reaction is done under oxygen-free conditions a pale yellow compound can be isolated on the addition of pyridine, this was identified as cuprous pyridine fluoroborate9 on aeration a blUe compound Cu(py)4,0Fa2was formed. The product is in fact cuprous fluoroborate. The same worker obtained solutions of silver and cuprous fluoroborate by passing boron trifluoride through suspensions of silver fluoride and mixtures of copper and cupric fluoride in toluene respectively. Solutions of cuprous nitrate have been prepared by shaking solutions of silver nitrate in acetonitrile with copper (Morgan 1923). Sharp and Sharpe (1956b) using a similar me2thod have obtained solutions of cuprous BP4-; SO3F-; PFG-; AsP8-; NbF3 arid TaPci- in aromatic hydrocarbons. Although these solutions are stFill_l.e in the absence of moisture no solid complxes co.Ald be isolated. In the present work solutions of most of the 1st -76- row transition metal perchlorates, fluoroborates and trifluoroacetates have been prepared in benzene and toluene. These solutions have been obtained by shaking a solution of the appropriate silver salt with an excess of the metal or metal halide. The results obtained were so promising that the experiments were extendbd to AgPF69 AgVF0; Ag2SiF6 and Ag2TiF6 although the range of metals investigated was restricted; There are no previous mentions of either silver hexafluorosilict4e or hexafluorotitanate, the former has been characterised by X-ray powder photography and is found to have a cubic lattice (a=6.12A) isomorphous with the potassium, ammonium and rubidium salts. Preparations of the fluorogermanate were unsucces- ful as it was found very difficult to free the white product, prepared from QeF4 AgF, from the last traces of solvent (anhydrous hydrogen fluoride). The following table summarises the results obtained by shaking the silver salts with the 1st row transition metal halides in benzene and toluene. Colours indicate the co?our of the solution; in some cases the solutions obtdined were so dilute that no colour was apparent, in these cases a 4. indicates a successful reaction. Reactions which did not work are marked -.

Ti3 q' C3 4. Cr3 `!' 11112 + FO2 P03 CO2 Ali` ."-• Cu"

Form of TiC13 VC13 CrC13 Mn /70C12. PeC13 CoC12 NiC13 CuCla ;Cu Metal Used.

B pink green. + colour. green brotn pink-blue yellow yellow C104 less T pink green • colour- green brown pink-blue yellow yellow less

B e - + colour- - - + BP?, less T - - + colour= _ — + + + less B rod green green colour- Pe" brown blue yellow blue-green Cu2 + CF3 CO3 - less T red green green colour- Pe3 4. brown purple yellow dark blue Cu2'1' less

PPG

ITF0 - B

a- B SiF3 T Brcaday cpeaking the cupric, cobaltous and nickel salts were more 'soluble in the hydrocarbons than the others, and it was possible to obtain stronger solutions of trifluor- acetates than perchlorates or fluoroborates. In all cases solids were precipitated but due to the fact that they were mixed with excess of the metal or metal chloride and silver chloride they could not be isolated except in the case of cupric trifluoroacetate.- Cuprous perchlorate and trifluoroacetate could not be prepared either by shaking the silver chloride salt with copper or cuprouwr by shaking the cupric solution with copper. However Swarts (1939 ) has reported the prep aration of cuprous trifluoroacetate but makes no reference to its solubility in aromatic hydrocarbons. The failure to Tj3+ obtain solutions of V3+q- fluoroborate maybe due to the instability of the fluoroborate with respect to the fluoride. In the toluene solutions the precipitates were, in many cases, in the form of a sludge& There are very few other reports of other inorganic salts which are soluble in aromatic hydrocarbons; most of the published work is concerned with the halides of metals which can be thought of as Lewis acids, i.e. as electron acceptors. The aluminium and ferric halides are typical examples. There is a report that a' solution of aluminium bromide in benzene will dissolve potassium bromide to give a conducting solution ( Plotnikov and Jakubson 1928 ). Similar results are obtained by the same authors (1930) for solutions in toluene and o- and p xylene. Cuprous bromoaluminate is also soluble in benzene and toluene„ a 1:1 complex having been isolated. Although it is reasonable to expect the cuprous salt to be soluble there being co-ordination between the Cu+ ion and the toluene molecule ( thus stabilising the Cu+ ion ) it is unlikely that a similar species could be present in the potassium salts. Aluminium perchlorate has been prepared by reacting aluminium chloride and silver perchlorate in benzene (Acerete and Lacal 1954 ) the anhydrous salt being obtained on removing the solvent. The only other salts studied in aromatic hydrocarborth are those of gallium. During the last few years much work haa been done in trying to establish the valency states of gallium. Particular attention was paid to gallium dichloride, as thallium and indium have well established I+ and 3+ states the apparent divalency of gallium was an anomaly. As a result of Raman,nuclear magnetic resonance studies the identity of the dichloride was shown to be [ Ga I ] + [ GaIII u14 . This and other gallcu.s -80- salts are soluble in benzene e.g. Ga AlO10,Ga Al Br4. Rundle and Corbett (197 ) have shown that benzene solutions of the dichloride have very similar properties to that of silver perchlorate. Of the thallous salts investigated Warf (1952) has shown that the perdhlorate and fluoroborate are insoluble in toluene and Sharp (1957 ) has shown the insolubility of the Plc; and SoF6 - salts. Thallous trifluoroacetate however is soluble in benzene (Swarts 1939) and in toluene, m-xylene an mesitylene Sharp(1957). Amongst the other halides which do dissolve in aromatic hydrocarbons are those of mercury. A number of other halides form complexes With aromatic hydrocarbons and include AsC16 ; InBro SnC14 ; PC16 ; WF6 ; TaC16 and NbC16. Separation procedures for the xylenes and for the methyl benzenes using liquid hydrogen fluoride - boron trifluoride mixtures (McCaulay, Shoemaker and Lien 1950, McCaulay and Lien 19519 1952) and isotope exchange reactions between para substituted phenols and benzenesulphonic acid ( Gold and Satchell 1955) have been explained by postulation of a protonated aromatic system. There are two possible ways of protonating an aromatic nucleus a) by loose attachment of a proton to the -81- benzene ¶ electron system and b) by having a bond between the proton and the ring system.

a) b) • + , r H : 1

H 114- Olah, Kuhn and Paviath ( 1956 ) have actually isolated complexes of the type Are BF:.

SOLUBILITY OF lst ROW TRANSITION METAL FLUOROSORATES v TRIFLUOROACETATES AND PERCHLORATES IN ETHER.

As stated at the beginning of this chapter one of the factors which were thought to be necessary for a solvent to be a good one was a high dielectric constant. Subsequent work has shown that this is only one aspect of the problem and that a pair of unshared electrons is an important factor in eeciding the suitability of a solvent. To demonstrate this, the present work was extended to prepare solutions, in a manner already described, in a solvent of low dielectric constant. Anhydrous diethylether was chosen other solutions were prepared in nitrome'hane. The following table gives the value of the dielectric constant for some solvents ( Weissberger 1955 ) -82- Diethyl ether benzene toluene acetone nitromethane water 403 203 2.4 1901 3509 80.4 The fact that many inorganic salts are soluble in ether and the formation of etherates is well known. Covalent mercuric chloride and bromide are soluble in ether and the corresponding magnesium salts form etherates Mg X2 2 Et20. Swarts (1939) has shown that mercuric thallo'asand cupric triflucroacetate are soluble in ether and Sidgwick (1950) lists many more acid halides and oxyhalides that dissolve in ether and form etherates. From all this evidence the donor properties of ethers must be considered well established. Jander and Kraffcyk (1955 ) have proposed auto ion- , ization of ether to account for the observed conductivity of solutions of a series of ionic and covalent compounds in ether. Reactions of lithium aluminium hydride in ether ( Jander and Kisaffczyk 1956, Vliberg and Schmidt 1951) have been explained on this hypothesis. Grignard reagents have been shown to conduct electricity in etheralmlutions demonstrating the existence of ions ( Evans and Lee 1933- 1934 ) but the solubility of salt in ether must be considered as solvationof the ions by co-ordination with -83- the lone pair of electrons on the oxygen atom of the ether molecule. Few silver salts form complexes containing a metal-, oxygen bond apart from chelate compounds. Among the exceptions are the complexes of fluoroacids (Sharp 1957) and the perchlorate. Bathe (1931) has found evidence of the hydrated ion A6 (0H2)21- in solution and the lack of hydrated silver salts may be due to their covalent character. V(illstater and Pummerer (1904 •) have prepared both pyrone and dioxane complexes with silver perchlorate and an X-ray determination of the latter ( Prosen and. Trueblood 1956 ) shows that each silver atom is surrounded by a regular octahedron of dioxane oxygen atoms. Conduct- ivity of the silver perchlorateiacetone system indicates the possible existence of the molecular compound Age104 Me2C0(klochko and Uchurkhanov 1953 ). Sharpe (1952 ) showed that silver fluoroborate is,. soluble in ether and the solubility of silver salts of fluoroacids has been shown to extend to AgPF6 ; AgAsFo ; AgTaF6 ; AgSbF6, and AOlibF6 by Sharpf and Sharpe (1956a9 b) who also showed that cuprous salts of these fluoroacids can be prepared by displacement of silver with copper. -84- Clifford 9 Beachell and Jack (1957 ) have prepared rare earth fluoroborates by reacting the fluorides with boron trifluoride etherate but they do not say whether the salts were soluble or not. In the present work metathetical reactions were carried out on the same metals and acid radicals in ether as previously described for the solutions in benzene and toluene. Monnier (1957) has done similar experiments using the bromides of a few metals and silver perch3orate and claims to have isolated the anhydrous salts from solution but gives no analytical data to support this. We were not able to isolate the anhydrous salts by removal of the solvent. Cuprous trifluoroacetate could not be prepared; immediate disproportionation taking place with- the formation of cupric trifluoroacetate and an unidenti- fied black substance. Similarily cuprous perchlorate could not be prepared ; no copper would go into solution as long as conditions were strictly anhydrous. Vanadiur (III) fluoroborate sOlution could not be prepared even after shaking for 14 days. Although Ti III trifluoroacetate could be prepared in solution fairly easily the -85— perehlorate and fluoroborate were very unstable. Reaction was swiftest with the cobaltous salts: the reactions :— :Ag C104 Co(C10,4 CoC12 + Ag BF4 AgC1 + Co(BF4)2 Ag02 QCF2 Co( 02=113 )2 were complete in one hour using 0.75 gm. of the, silver salt and an excess of the cobaltous salt. Reactions with chromic chloride, however, were only complete after 5 or 6 days shaking. The rate of reaction seems to be solely dependent on the physical nature of the halide used. The reactions in ether were, in all cases, complete in a shorter time than when aromatic hydrocarbons were used, and the solutions obtained were much more concentrated. The reactions proceeded just as quickly with nitromethane as solvent. .Chromous chloride would not react to form chromous salts in nitromethane; as it reacts with ether reactions in this solvent were not attempted. The products of the reactions of chrpmous chloride with ether are not fully understood; Hein, Farl and Bar (1930) suggest that CrC120C2B6 is a product but butane which is also expected to be formed has not been detected. McAvoy (1962) has -86- detected acetaldehyde as a product and suggests that there are at least two others one containing a carbon- chromium bond and the other being an ethoxy compound° The following table summarises the results.

Ti3 Mn2 C r3 Cr. Fes Fes 4. CO2+ Nit •• cua+ Cuz Form of TiC13 VC13 CrC13 1.1n FeCl3 FeC12 CoC12 NiC12 CuCla Cu metal used

C104- bluish green green colour- yellow green purple yellow green less

BF4 green colour= -- purple yellow green colourless less

Oa C.CF3 bluish green green colour= reddish brown purple yellow green disproportion less brown !a3+ soln. ates Cu

CO SiF0 2 sky- yellow yellow blue green + p.p. green

PPG purple soln.

0 bluish pink yellow solno green

sky- yellow TiF0 2- blue 4. pop.

Colours indicate colour of soin, obtained. - indicates unsuccessful reactions, -88-- THE NATURE OF THE BONDING IN AROMATIC HYDROCARBONS

The silver salts of fluoroacids have been shown to be ionic by consideration of X-ray and infra-red studies ( SharP:1957 ). Similar studies on silver perchlorate ( Braeksn and Harcng1930 ) plus conductivity and cryoscopic measurements ( Hill 1921 ) have shown that silver perdnlorate is almost certainly ionic, whilst silver trifluoroacetate is probably covalent.. In the cases of the fluoroaclds and perchlorate salts the bonding in aromatic hydrocarbon solutions must involve the silver ion and the aromatic system. Mulliken (1952 ) has discussed the formation of such a complex in considerable detail. The possible positions, that are considered, for interaction betWeen the Ag+ ion and the aromatic system are; above the centre of the ring ; above one bond in the ring ; and abOve one corner of the ring. Ai:sandwich structure betWeSh the silver ion and two benzene rings was not consicfeied although this type of structure is now well estabiished. Mulliken, in considering the symmetry conditions, excludes the contro-symmetric model on the ground that such an interaction would involve -89- considerable electronic exitation ( c* n4 4 E.V.). Similar considerations plus geometrical distribution of the M.O. of the benzene ring rule out the last possibility. The model that Mulliken prefers is that in which the silver ion is above one of the C-C linkages* Evidence to support this ,theory has been given in two papers on the X-ray crystal structure determinations of the AgC104. 06116 complex ( Rundle and Goring 1950, Smith and Rundle 1958 ). The conclusions reached are that the Ag+ ion is above one of the C-C bands and slightly to one side and that the benzene ring is distorted. A shortening of the C-C tend involved with the Ag+ ion is observed (1.355i as opposed to 1.43R)* In the molecule so described each benzene ring is bonded to two silver ions and each silver ion to two benzene rings* l'ulliken does not consider that the geometrical requirements of such a riacro-molecule will be altered much from his model* Taufer, rurray and Clevland (1941) have studied the Raman spectra of benzene in aqueous silver perchlorate and find that there is a smaller frequency laift than that observed for ordinary olefins indicating that the influence of the Ag+ ion is directed towards more than one double blend. The lowering (by 13cm-1) of the totally -90- symmetrical vibration is too small to account for the attachment of appreciable mass anywhere else than on the sixfold symmetry axis . Aqueous silver nitrate complexes with olefins show shifts of 60-70 em.-1 and with actylenes of 100 cm.-1 in- the Raman spectra. Rundle and Corbett (1957) use nulliken's model to explain the solubility of Ga[GaC14 ] and Ga[A1C12in benzene. However in the model propoSed the 0.-a ion is above the centre of the ring. In the case of Aeand H the lowest orbital available for donation into is an s orbitall movement of the cation to a position of lower symmetry is necessary for bonding. The lowest acceptor orbitals which can accept TT electrons from the benzene ring are the p orbitals in the case of gallium. The Ga [us C14 ]C61 complex is not isomorphous with AgC1041 .C6H6 although their condueivities are very similar in solution ( Mcrullan and Corbett 1958 ) This picture of the bonding is too simple because it does not take in4 o account the actual bond energy, of the complex. This would not be expected to be ,very large9 in fact9 no largei than any other charge transfer complex. Smith and Rundle (1958 ) estimateit as 15..7 kcal/mole whereas Tildesley and Sharpe (1953) give a value of about -91- 50 k. cal mole.-1. The iodine -benzene charge transfer complex has a bond energy of about 2 k cal ( Cromwell and Scott 1950 ) which is the normal value for the bond formed in simple charge transfer complexes. Andrews and Keefer (1952 ) were unable to find the expected charge transfer band for silver-aromatic complexes in the ultraviolet but Orgel suggests that it should occur at 23003. where there is 100% absorption for a normal aromatic hydrocarbon.. Cteric hindrance does not seem to be important9 although Mulliken (1952) has tried to correlate this with the stability of 1:1 2sifver'complexes in aqueous silver nitrate. However, Tildesley and Sharpe (1954) have:shown that silver heptafluorobutyrate is soluble in durene and hexamethyl- benzene and Sharp (195?) has shown that AgPP6 disv;olves in mesitylene. One would expect maximum steric hindrance in these molecules so some other factor must be: considered in the bonding. To add to this9 Perguson (1956) has shown by infra-red studies that in the halogen-benzene complexes the most likely configuration is that in which the haloren is coincidental with the benzene Ce axis in agreement with the partially polarised molecule "g4'-I8- -92- suggested by Dewar (1946a v 1946b). The energy of the.Ag+ - benzene interaction indicates as already explained, something more than a simple charge- transfer complex. The most reasonable explanation for this high bond strength is back donation of electrons from -• the filled d- orbitals of the silver ion to the anti- bending etrbiiala of the benzene. A symmetrical model would now be possible because excited orbitals can now be involved in the ir bond donation. This sort of back donation was first proposed by Chatt (1949) and Chatt and Duncanson (1953) to explain why platinum-olefine complexes can be - formed whereas trimethyborane-olefine complexes cannot - The fact that platinum can form these complexes and the borane cannot suggests that the d orbitals must be involved. In describing the ethylene- platinum complex-'it is suggested that the overlap of the 5d6s6p2 orbitals withthe Of-orbital of the olefin form the ci-bond, and that a i bond is formed by overlap of a filled 5d orbital-and a if lwlecular orbital of the olefin. -93-

Exactly the same type of bonding can occur in the silver -aromatic complexes and would account for their stability. The 5s orbital on the sil- r ion can accept electrons from the -Tr orbitals of the benzene, and candonate 4 d electrons into the eorbitals of the benzene ring. It also explains why sodium and potassium perchlorate and lithium fluoroborate are not soluble in aromatic hydrocarbons. The same reasoning can be used to explain the solubility ( small though it may be) of the first row transition metal perchlorate9 fluoroborate9 Ilexa fluorophosphates9 vanadates9 silicates and titanates prepared in this work. The instability of Ti3 solutions in aromatic hydrocarbons is explained on the general instability of the Ti3+ salts to oxidation and also to the fact that the ion has only one 3d electron. The studies on the solutions in benzene and toluene have been only a preliminary survey and much more intensive invest- iga4 ions into their properties are necessary before any definite conclusions can be drawn on the mechanism of -94-- their dissolution. The bonding in the ether solutions must be considered as donation of the lone pair of electrons on the oxygen atom to the metal ion. Such a co-ordinate link would be expected to have about the same energy as a covalent band and the complexes so formed to be reasonably stable,' -95- EXPERMENTAL All the solutions mention in this section were made up in a dry-box using anhydrous solvents and salts. The solutions were contained in 100 ml. flasks and had well greased ground glass stoppers wired on. In the case of the titanium solutions special precautions were taken and are described below. Shaking was done on a mechanical shaker with times of shaking varying from a few hours to fourteen days. The solutions were filtered in the dry box using a sintered glass crucible ( porosity 4 ) and a 'hand bulb connected to a Buchner flask. Often it was necessary to filter two or three times because of 'the finely divided condition of the precipitated silver chloride. Each solution was tested for silver and chloride and by boiling off the solvent and then testing in the conventional way. All starting materials were stored in the dry box. Prvmaration of starting materials Solvents Ether, benzene and toluene were all dried with sodium wire and the dried solvents stored still over sodium, in the dry box. Vitromethane was dried over calcium chloride and then distilled into molecular sieves. The atmosphere in the dry box was kept free from moisture by open trays of phosphorus pentoxide. -96- Silver Perchlorate. The anhydrous salt was purchased from B.D.H. and was further dehydrated by continuouspumping over phosphorous pentoxide for 24hours. The salt was then left under vacuum over P205 for several days before being transferred to the dry box. Silver Fluoroborate. Approximately one quarter of the silver fluoroborate used was prepared from silver fluoride and boron trifluoride in anhydrous hydrogen fluoride as described in chapter 1. However due to the very corrosive nature the valve used to control the flow of gas was not very reliable, and consequently the preparation of Sharpe (1952) was used. Dry silver borate (1.5g) was paced in a silica bottle fitted with a standard B14 cone, Bromine(5m1) was added to act as a moderator in the reaction, and bromine trifluoride was added, drop by drop, through a copper funnel. When excess had been added the bromine and the remaining bromine trifluoride were removed under vacuum on a normal vacuum linew. The last traces were removed by heating to 130° for half an hour. Silver Trifluoroacetate. Anhydoue trifluoroacetic acid was added to an excess of silver carbonate plus a little water in a platinum dish. The solution was -97- warmed on a steam bath to ensure complete reaction and then filtered. The white solid obtained from evaporating the solution on a. steam bath was dried in vacuo over phosphorus pentoxide. The silver compound was then extracted with anhydrous ether and recrystallised twice. Silver hexafluorqphosphate This was prepared by mixing silver powder, phosphorus pentoxide ( a slight excess ) and bromine and then adding bromine trifluoride. After removing bromine and excess trifluoride by heating under vacuum at 150° a purewhite product was obtained. Silver hexafluoro adate A mixture of silver powder and vanadium trichloride in equivalent quantities were reacted with bromine trifluoride ( Emeleus and Gutmann 1949 ). A yellow product was obtained by heating at 130° in vacuum for an hour. Silver hexafluorosilicate Silicon tetrafluoride was prepared from silica and fluorosuiphonic acid (Edwards 1954 ) by heating to 200° with a little water. The tetrafluoride after passing through a reflux condenser was collected in a U-trap cooled in liquidair. It was then distilled on a vacuum line into a bomb containing a solution of silver fluoride in anhydrous hydrogen fluoride. After 12 hours the solution was removed and -98- white silver hexafluorosilicate obtained. Aga,Pi.P6 Pound Ag = 60.9% Calc Ag = 60.4% An X-ray powder photograph showed that it has a cubic lattice a = 8.12 A d intensity sin"°6 (obs) hkl sireq,calcl 3.27 10 0.0558 211 0.0540 2.86 10 0.0728 220 0.0720 2.69 6 0.0819 221 0.0810 2.46 5 0.0989 311 0.0990 2.36 4 0.1070 222 0.1080 2.28 3 0.1141 320 0.1170 2.18 2 0.1255 321 0.1260 2.04 1 0.1429 400 0.1440 1.944 3 0.1572 410 0.1530 1.893 1 0.1663 331 0,1710 1.676 4 0.2115 422 0.2160 1.620 2 0.2265 430 0.2250 1.583 3 0.2371 510 0.2310 1.527 1 0.2548 --- 0.2520 Clifford, Beachell (1957) prepared a white solid using the same reactdnts at atmospheric pressure but did not analyse it. They suggested that it was AgSiF6. -9/3= This salt has not been otherwise characterised. It is isomorphous with.the,corresponding potassium9 rubidium and ammonium hexafluorosilicates. K2SiF6 (11H4)2 SiF6 Rb2 S iF6 Ag2 a=8.13A01 8.34A02 8.45A°2 8.12A° 8017A°2 8.38A°3 1 Cox 1953 2 Ketelaar 1935 3 Bozarth 1922 Silver hexafluorotitanate Titanium tetrafluoride was prepared from titanium dioxide and bromine trifluoride (Emeleus and Woolf 1950). The product had to be heated at 180° to remove all the excess bromine trifluoride. The stoichinometric quantities of titanium tetrafluoride (0.50 gm) and silver fluoride (1.02 gm) were weighed into a bomb and anhydrous hydrogen fluoride (10m1) was distilled on. The bomb was rocked for 18 hours;when the solvent was removed white silver hexafluorotitanate was obtained. Ag2 TiF6 calc Ag = 57.20 found Ag = 57,0% Titanium trichloride - was donated by Laporte Titanium Ltd. and had been prepared by the reduction of the tetrachloride at 900° in hydrogen. All manipulations -99= This salt has not been otherwise characterisdo It is isomorphous with. the ,corresponding potassium, rubidium and ammonium hexafluorosilicates.

K2S1P6 (Nli)2 SiFG Rb2 SiFe Ag2 SiP6 a=8.13A°1 8034A°2 8.45A°2 8012A° 8017A°2 8,38A°3 1 Cox 1953 2 Ketelaar 1935 3 Bozarth 1922 Silver hexafluorotitanate Titanium tetrafluoride was prepared from titanium dioxide and bromine trifluoride (Emeleus and Vloolf 1950). The product had to be heated at 180° to remove all the excess bromine trifluoride0 The stoichiometric quantities of titanium tetrafluoride <0.50 gm) and silver fluoride (1.02 gm) were weighed into a bomb and anhydrous hydrogen fluoride (10m1) was distilled on. The bonb was rocked for 18 hours*2 when the solvent was removed white silver hexafluorotitanate was obtained. Ag2 TiP6 ears; Ag = 57.2% found Ag m 57.0% Titanium trichloride was donated by Laporte Titanium Ltd. and had been prepared by the reduction of the tetrachloride at 900° in hydrogen. All manipulations -100- involving this compound were carried out in a nitrogen bag and all solvents had previously been freed from oxygen. The stoppers of all bottles containing solution experiments of this chloride were sealed with paraffin wax as an added safeguard in preventing oxygen and water entering the reaction flask. Vanadium Trichloride Vanadium trichloride is usually prepared by heating vanadium pentoxide and sulphur in chlorine and refluxing the product with sulphur mono. chloride ( Ruff and Lickfett 1911 ). However this method always involves removing sulphur at 300° in vacuum; a simpler method of preparation was developed during the present work starting from vanadium. Vanadium metal (5gm.) was placed in a silica tube tilted at an angle; the bottom end of which was connected by means of a ground glass joint to a flask containing sulphur monochloride (see diagram ). furnace

to dveyschel bottle containing conc. 112SO4

:i2 012 -101- A dry stream of carbon tetrachloride and oxygen free nitrogen was passed over the metal at 660°. VC14 formed rapidly and was carried as a dark red vapour into the flask ( cooled in ice water). At the end of the reaction any tetrachloride left on the walls of the reaction tube was gently warmed and the dark red liquid easily ran into the flask which was then fitted with a reflux condenser protected by a guard tube. The mixture of S2C12/VC14 was refluxed under dry (02 free) nitrogen for 18 hours and the excess S2C1,I, plus SC12 was distilled off under reduced pressure. The purple solid was washed five or six times with anhydrous carbon disulphith ( in the dry box ) and it was then heated under vacuum at 85° for two hours. The product was pure VC13 Analysis VC12 %C1 Found 67.4% Calculated 67.6% Chromic chloride - was prepared by heating chromium metal in a stream of dry chlorine. it was found that to get an 'active' product, i.e. one which would react to give chromic solutions in organic solvents4 it was necessary to carry out the chlorination at as low a temperature as possible, -102- Chromous chloride - was prepared by passing dry hydrogen chloride over chromium at 19000°. A white product, very sensitive to moisture, was obtained.

Mang.anese metal powder was used as it was found that manganous chloride prepared by dehydration of the hydrate in vacuo would not react. Ferric chloride used in these experiments was the normal commercially available anhydrous compound. Ferrous chloride - was prepared by reduction of ferric chloride with hydrogen followed by sublimation to purify the product. Cobaltous chloride ® was prepared by heating the finely ground hexahydrate in vacuum at 100° for 36 hours. A pale blue product was obtained. Nickel chloride Thionyl chloride was slowly added to the finely divided hexahydrate until in large excess. The mixture, adequately protected from moisture, was left for 7 days and then after refluxing for half an hour the thionyl chloride was distilled off. The anhydrous salt was washed well with anhydrous carbon disulphide and then heated in vacuum at 60° for 18 hourse Cupric chloride - the commercial anhydrous salt was used. In experiments to get cuprous solutions either -103- finely divided copper or freshly prepared anhydrous cuprous chloride was used.

-104- 0 ILA R T_B

SOME PHYSICAL AND CHEMICAL PROPERTIES O1 1ST ROW TRANSITION METAL FLUOROBORATES0 PERCHLO/ATES AND AND TRIPLUOROACETATES.

Spectrg pj Ethereal Solutione During the present work it was thought that an examination of the ultravioletp visible and Lf r infrared spectra would be of interest. No mathematical attempts have been made to explain the observed bands and any assignments are made purely on comparison with published spectra and on the relative strengths of the bands. The literature on the spectra of d-cI transitions is vast and is covered extremely well by Dunnos review (1960) and in the book by Jorgensen (1962). During the last few years there has been a considerable :advance in the understanding of the magnetic and spectroscopic properties of inorganic complexes. At the present there are three main theories which attempt to explain these properties. (a) Valence Bond Theory. This theory, put forward by Pauling (1935) is not very helpful in explaining epectga. -105- The reason for this is that the different valence b$nd resonance structures of molecules rarely approximate to the observed energy bvels although theoretically it should lead to the same reeults,. This theory pays no attention, in its usual form, to the existence of excited states. (b)Crystal Field Theory. This method was first used by Penney and Schlepp (1932) to investigate the apparent anomalies in the magnetism of crystals. Stri king successe have been obtained in applying this method to the absorptian bands which lie in the visible and ultraviolet regions in the case of octahedral complexes of the transition metals, and their associated tetragonal and rhombic distorted forms. (c) Ligand Field Theory. This approximation has best results when it is applied to a case where there is a strong interaction between the ligand and ion orbitals. The transitions involve either charge transfer (Milliken 1950) or involve orbitals having a large degree of localisation over the ligands and central ion. When discussing the atomic energy levels in a polyelectron atom or ion, the various forces acting within the species are placed in order of magnitude and calculations -106— ars carried out, starting with the largest force, by a series of approximations. For a polyelectron atom or ion there are, in general, three types of forces:- (a) central field forces (electrostatic), (b) interelectronic repulsion (electrostatic), (c) spin-orbit coupling forces (magnetic). The central field forces are coulombic in nature involving the attraction of electrons by a central positively charged nucleus and these forces are, of necessity, larger than the others (to ensure a stable entity). Forces b and o must also be taken into account, and if b>>c the levels obtained are said to belong to the Russell-Saunders or LS coupling scheme. If c>> b, as is true for heavy atoms (7....>30), then the final levels are said to belong to the jj scheme. The Russell-Saunders coupling scheme applies very well for those atoms with Z <30 and therefore includes the first row transition metals. One electron atom schemes for an atom are drawn up i.e. the configuration, and then the effect of interelectronic repulsion is calculated. This has the effect of splitting the otherwise highly degenerate configuration into groups of levels, having lower degeneracies, known as terms. These are specified by the total atomic orbit angular momentum (L) and the spin angular momentum (S). -107- i:.0 L " ii 1 2 3 4 5 6 S PDF

The orbital degeneracy is 2L + 1 as there are 2L + 1 directions in. which L may be orientated in a magnetic field.

S =1E and 8 has multiplicity `C = 28 1 .3i The term is denoted thus e.g. S = 19 L = 1 term is 3P (C atom ground term) S = 1a9 L = 3 term is F (CrIII ground term). 0 0 0 . The states Do P9 S9 VD9 VP9 QS are possible for the carbon atom but application of the Pauli exclusion principle shows that not all of these states are allowed. Those a allowed are the P9 VD and states. The possible 11L values which can be formed by the combination of ml and me are shown below. 11 = 1 the M values are ml = / 0 -1 I =1 2 1 0 1 2 1 0 -1 -2 1 0 -1 0 1 0 -1 0 -1 -2 0 m2 -108-

For S = 1 both electrons have the same spin quantum numbers, so that according to Pauli they must differ in their values of ML. Therefore none of the ML values from the diagonal can combine with S = 1. When S = 1 one is limited to the ML values 19 09 -19 giving rise to a P term. Thus there are two sets of ML values 2, 1, 0, -19 2 and 0 to combine with S = 0 giving and 9S terms. For the carbon atom the allowed terms are P, 9S, °D. The terms arising from the configurations d" are as follows (Dunn 1960) d'9 d6 2D d29 d8 3F 31' 9 G 9 D 9S

22 d3, (IT 6 F 4P 2H 2G 2P 2D(2) vp d4, d6 6D 3H 3G 3P(2) 3D 31)(2) 9 I oG(2) 9D(2) 9 5(2) d6 35 6G 6P 4D 4P 21 2H ?.G(2) 2P(2) 2D(3) 3P 25

Magnetic fields are generated if L and S are non zero and orientate their moments with respect to each other in the direction where their interaction energy is a minimum. This coupling gives rise to the inner quantum number J which is a measure of the total angular momentum (spin + orbit) of the atom. 109- J= L S9 L S 1 L S So for the carbon atom 3p s. J = 2i 1, 0

i.e. 3P29 3Pa9 5P0 levels CrIII J = /29 V2, V29 3/2 41$/29 i.e. 4-9/P /2 9 4/%5/2 *6Fa/2 levels.

The collection of all valuesof J having the same L and 59 i.e. arising from the same term is referred to as a mult iplet and each value of J associated with a given L value is called a component. Spectral transitions in atoms occur between the components of two different multiplets according to the selection rules.

AS = 0, AL = 09 19 AJ = 09 1 J = 0(IW = 0 When the coupling between the li is not too strong only one electron changes its quantum numbers in the transition so that a l'urther selection rule is Ali = ± 1 i.e. S d :„, f but S , p f . By applying Hundtue rules it is possible to determiniwhich of the components has the lowest energy. (1) Of the Russell Saunders states arising from a given electron configuration and allowed by the Pauli exclusion principle the most stable state will be the one with the greatest multiplicity. -110- (2) Of a group of terms with a given S value, the one with the largest L value has the lowest energy. (3) Of the states with given L and S values in a configuration consisting of less than half the electrons in a subgroup, the state with the smallest value of J is usually the most stable v and conversely for a configuration with more than half the electrons in a subgroup. Therefore the levels for the carbon atom are aPo, aPi,

51)20 QD2, ISO. Application of an external magnetic field can cause a further splitting of the energy levels; J is split into 2J + 1 equally spaced levels corresponding to the number of values that can be assumed by the magnetic quantum number ML. Diagram 1 (adapted from Eyringp Walter and Kimball 1944)- shows the various energy levels arising from the IP configuration

-111- DIAGRAM 1_ _ _ _ S - ii1;120

2S -I- 1 singlet; level L1 = 2 1 - 41)2 _ 0

M 1-

••••••••••••• . . M = 2S +/1...... -4:._/------triplet / . . 9.

level/ , . / / 3 Zp 0 ------6 I

np2 1.1 . S.1l S.Sii magnetic no electronic ontalYng coupling field interaction coupling present

CRYSTAL FIELD THEORY The first calculations on the effect of electric fields of various symmetries on field free atomic terms and degenerate orbitals were done by Bethe (1929). -112- If Ruosoll-Saunders coupling is assumed three different . — approximation orders arc posbiblo;- (1)Crystal field energy spin-orbit coupling, energy. The simplest case is the (1 9 configurations Titanium III has this configurations D3 Du the term being eD and the components 5 41 9 e A. When it is surrounded by six octahedrally placed ligande (such as water) the previously fivefold degenerate D term splits into two new levels denoted by symmetry symbols t2 for the threefold degenerate level and ce.g for the twofold level. The effect is shown in diagram 2: -113- DIAGRAM 2

6Dg. .Dq

===ta g in field-free ion in octahedral 5 fold degenerate. environment. The Mulliken symbols are used to designate the states of an ion in the crystal field. Their origin is in group theory but they can be regarded simply as labels.

State of free ion States in the Crystal field

S Al

P TI

D E + T2

P Ao + T1 + To

G E Ta E 2T1 Ta

-114- The following table dhows some of the states arising x y in an octahedral field from the configuration teg og (Dunn 1960)

Free ion Strong octahedral Levels arising in octahedral Configuration field configura- field. tion.

2B d Q 9 e2 t g nTag

d'3/, de 2

D °Tao. 3 Tag 9 TQg t2gag g (g, T2g +

sag uTT.Ig QEg 9T0g

d39 e58 2 tAg eg 22T5113 22T2g

egg 580' 4T2g 2A L8 -1-°A211 erNg g a 2qTe ta A2 B 2ITIg + 2T28 g g g

The weak field limit implies that the tag and eg hvels are not very far apart aad that therefore interactions between configurations such as tag and tageg must be considered. In -115- the liMit of very strongfields the energy gap between the one electron t2g and eg levels maybe so large that interactions between states arising from configurations involving both till2g and eg orbitals may be ignored. All the features of a spectrum in the strongly perturbing field limit cannot be explained by ignoring t2g o eg interactions and have to be considered by examining the intermediate field case where these interactions are taken into account. In a field free atom or ion or in a weak field case the maximum multiplicity is Bil but in a strong field case it is four. CoIII complexes are a good example of this effeoty in a weak field they have four unpaired spins but become diamagnetic in the strong field. The cases of strong fields and weak fields in tetrahedral fields have, been investigated in a similar manner. Ligand inequivalencies9 the Jahn-Teller effect spin-orbit coupling and crystal lattice interactions can all cause distortion from the regular symmetry as a result of further perturbation of the crystal field. The spectra of d-d transitions of ethereal solutions are summarised below. d 9 and (19 In both cases, if the perturbing field is perfectly 2 2 cubic, there is only onetransition Tlitg --,E for dc and -116- and 2Es-PT2g for d°. However there are perturbations which will affect the simple concept of one transition (19 and d° will be distorted dace to the Jahn-Teller effect and also for d° the spin-orbit coupling is about 800 cm, which will impose a further perturbation. 3 The d electron configuration of Cue is t2g es ) which in an octahedral environment will giVe a doubly degenerate ground state as there are two assignments of the eg electrons=.- (d z2) (dx2 m y2)9 and (d z2.). (dLe . 2) . If one considers that the d 2y2 orbital has only .one electron thil a loss of spherical symmetry will occur in such a way that the nucleus is less screened along the x and y directions than along the z direction. As a result the ligands in tha- x and.y planes are attracted with a force larger than the ligands along the z axisp producing four short coplanar bond's with the metal atom and two longer ones. Similarly there is only one electron in the d2° orbital an opposite distortion will occur. Experiientally only a few complexes are known with two short bonds and four long ones9 e.g. K200169 ICCuF3. Other configurationa Which can show this soxi of degeneraby and distortion are listed below

-117- (tag)' (eg)v Cr2÷ Mn

dv (tag)° (eg)9 Com Ni" d0 (tag) ° (eg)3 Cu Ag Due to the deviation from octahedral symmetry the energy levels are further split as shown in diagram 3. If the splitting is large enough there are three transitions which can occur in the visible and near infrared.

DIAGRAM 2B2

2E g

2 g

2B2g Free ion cubic field rhombic field

, 4+ The spectrum of Ou(Ra0)0 has a single broad maximum at 129600 cm. but it is clear that the envelope contains more than one transition. Bjerrum0 Ballhausen and Jorgensen (1954) suggest that absorption bands (extinction coefficient -118- as a function of wave number) follow a simple Gaussian error curve, and consequently are *metrical about their maximum value. Using this principle on the spectrum of Cu(H20)02+` they resolve the band into at least two symmetrical absorption curves with maxima at 12650 cm. and 9440 cm. The spectrum of bis(acetylacetonato)-coPper has been shown to have more than one band (Graddon 1960) and the related compound bis-(3-phenylacetylacetonatO)'=copper has been shown to have four d -d transitions ( Basu r Belford and Dickenson 1962).

The spectra of copper perchlorate r fluoroborate and trifluoroacetate in ether have broad bands in the following positionfi. Band in Solvent Reference -1 cm. ow—rmakm...* Cu(01002 12820 ether present work 13330 acetonitrile Underhill and 12200 ethyl acetate I Hathaway 1962 Cu(BF,D 2 12500 ether present work 13300 acetanitrile Underhill and 11620 ethyl acetate Hathaway 1962 Cu (020 CF3)2 12760 ether 13270 benzene present work 15d60 toluene -119- The only reasonably stable Ti° solution obtained in the present work was the trifluoroacetate. However it was not possible to obtain a complete spectrum before decomposition occurred. A weak band was observed at -1 , 3+ 17850 Cm. The Ti(H20)6 ion has a maximum at

209 300 cm. (Hartmann and Schlafer 1951b). d2 and (18 Vanadium III. This oxidation state was one of the first to be treated by Hartmann and Schlafer (1951a). Both the hexahydrate and the trisoxalato complexes have been studied in detail and have well defined transitions at 17000 cm. and 24,000 cm. These transitions are a assigned NM -4T2g(F) and 2T1g(F)---iPTILg(P)p (Orgel 1955)9 at first the second transition was thought to be 3Tg--;>4.A2g(P)9 but it is unlikely as it represents a 0 0 configuration with a vtwoq electron transition i.e. tzg

Og $/e3--9 eg . Two low lying singletstates tEg and qTag are expected from an .energy diagram of a d2 configuration but no trace of these have been found yet. -120-

3T2g(P) 5Tig(P) V(01044 14500 cm. 2.980 cme present work V(02COPOts 16150 23d00 v1I16H20 17700 25600 Furman and Garner 1950 III V 30C 16500 23500 Hartmann and Schlaf- er 1954 VIII6SCN- 16700 Furman and Garner 1951 Vanadium III flucoroborate solution could not be prepared. In all attempts a yellow solution containing no vanadium was obtained. Nickel II. Because of the large spin orbit coupling in nickel II complexes they have been very thoroughly investigated spectroscopically. The spectra of the octahedral complexes usually consist of three main absorption bands with two or three weaker transitions. The hydrated ion Ni(H20)321- has three main bands9 at _1 89500; 1)4.9500 and 269 000 cm. (Dreis ch and Trommer 19370 Kiss and Szabo 1943). Close examination of published spectra suggest that there are weak bands at 1990009 _1 229500 and 309000 cm. . The following assignments have been made. -121- 6Tms (F) 8,600 Oth.0-1 (F) 13,500 cm.-1 ',E (D) 15x400 cm,; gAig (G) 18 9500 ern. (P) 255500 omo

craguE a Other T2/3 g banc4s N1(C104)21 75560 12,580 11,080 23,800 6040 17"' in ether 5660 5140 Ni(02C0F5)2 8,820 13,150 13,780 - 24,400 8620 3z in ether 8300 7166 Ni(BF4 )2 8,400 12,480 14,500 - 24,100 7610 z in ether 5610 5 08D Ni(1120)62+ 85500 13,500 15,400 22,000 25,300 Ni(en)32+ 11,200 18,350 12,400 24,000 29,000

NiII 9M,KSOU 15,300 13,200 25,200 *AO NiIZ H2SO4 12,200 14,700 23,350

Present work Jorgensen 1954 a JOrgensen 1955 b -122- The yellow precipitate which is obtained when silver perchlorate and nickel chloride are shaken in toluene was washed with ether. The resulting solution gave a spectrum identical to that of nickel perchlorate prepared in ether, d3 and d7 Chromium III is the best known example of d3 and its spectrum with many ligands is known. The ground state is 4 A2g and is the only state of maximum multiplicity 3 arising from the configuration t2g The lowest lying 4 4 terms of the free ion are F, P and G in that orders so far six excited levels have been identified and they are:- -/ 4 A2 g 42 Eg 15,000 cm. ruby spectrum

-42 T1 g 159500 cm. t% 6 ®i T2g (F) 17,400 cmo hydrate _1 --)2 T2 8 22,000 cm. ruby

4 Ti (F)? 24,700 cm. hydrate

372000 cmo Large differences have been noted in the ruby spectrum and the spectra of chromium salts dissolved in glasses -123- Orgel (1957) explains the difference by assuming that there are large compressive forces in the ruby which create an anomalously large crystal field on the octahedrally co-ordinated chromium species. This idea is supported by the fact that a red colour is obtained in 'hard' corundum and spinel lattices and a green one in the 'softer' glasses. The spectra of chromium() perchlorate and trifluoroacetate in ether were obtained but preparation of the fluoroborate was very difficult and only a weak solution could be obtained.

4 Ti2g 4 4 Tig 2Eg 2Ti.g 2T, Tig "'g Cl(C104)3 15,950cm. 24,400cm. 31 780cm. Cr(02C.CF5)2 16,650 22,300 14,920

1r Cr(H20)634. 17,400 24,600 37,800 15,000 21 , 000 Cr(OX)33- 17,500 23,900 14,350 Cr ens5+ 21,850 28,500 14000 14,900 15,500 15,500

Present work it" Jorgensen 1954a Gates and King 1958. O tTorgensen 1954b Woldbye 1958 -124- Cobalt II. The hydrated ion, Co(H20)62+ has an absorption band in the infrared with a maximum at 8,100 cm._ (Dreish and Trommer 1937) and a broad complex band at about 19,000 cm. (Kiss and Csokan 1940). Orgel (1955) ®1 taking the value of Dq as 970 cm. predicts transitions _It at 8,100, 17,500 and 22 000cm. andeuggests that the 19,600 cm. band of the cobaltous ion includes both of the shorter wavelength quartet-quartet transitions and in addition some quartet-doublet transitions. The Co(11115 )02+ WP ion has an absorption maximum at 20,200 cma and the ethylenediamine complex, Co.en32+, has a maximum at 20,800 cm.®$ (Roberts and Field 1950) and because of the similarity of other ammonia and ethylenediamine complexes the two bands near 20,000 cmo are assumed to have the same origin ie. Ti g-0A2g transitions which give values for Dq at about 11,300 cm. -1 The known transitions are 4 T1g --14 T2g(F) /NJ 8,000 - 9,000 cm.-1 --Y)Eg N 11,000 cm. --/4 A2g tv 16,000 L. 189000 cm. • --0T1g(P)2 20,000 -'21,000 cm. The spectra obtained for ether solutions of cobaltous perchlorate, fluoroborate and trifluoroacetate each had a broad band in a region spreading from 14,000 to 20,000

-125- -1 cm. These bands are split with three or four maxima, -1 but no band was found at about 11,000 cm. 4T2g(p) 2A2g110 4Tigiai--- 14,490 cm.- Co(C104)2 7490 cm. 1 17,500 31,400 18,100 Co(02C.0F3)2 7760 14,350 17,380 16,180 199 200 27,000 co(BF4)2 7280 17,300 17,600 17,900

The precipitate obtained by shaking silver perchlarate with cobalt chloride in benzene was washed with ether and the solution thus obtained gave an identical spectrum to that obtained by carrying out the reaction in Ether. d4 and d6 No Orli solutions were attempted in ether as the ion reacts with the solvent (Hein, Farl and Bar 1930). Preparation in nitromethane and benzene were also unsuccessful; silver was produced in each case. A solution of ferrous trifluoroacetate could not be prepared; all attempts ended in the production of ferric trifluoroacetate in a manner analogous to the reaction of silver trifluoroacetate with copper. No reaction occurred when silver fluoroborate was shaken with either iron or ferrous chloride. A dark green solution of ferrous perchiorate could be obtained in ether but it was unstable and a spectrum could not be obtained before

-126- decomposition occurred to give a light green precipitate and a pale green solution. do Manganese II, in its complexes, has been exhaustively a studied. The d configuration is exceptional in that no spin allowed transitions are to be expected since the S ground state is not split significantly by the field and all excited states are of lower multiplicity. The fact that spin-allowed transitions do not occur is obvious from the colour of manganous compounds which are pale pink- much less intensely coloured than the other divalent transition metal ions. The assignments for the Mn(H20)621+ spectrum are:- 6 4 -1 ----) Ti g (G) at 18,800 cm. Ai g --34 T2g (G) 23,000 --41 Eg (G) 24,900 (G) 25,150 (D) 28,000 ---)01 Eg (D) 29,700 --y4 Tig (P) 32.,4004 (F) 35,400 --- 4 Tig (F) 36,900 ---,oT2g (F) 40,600 Spectra of manganous perchlorate, fluoroborate and trifluoroacetate were obtained in ether. Due to the -127- very intense charge transfer bands it was difficult to pick out the low intensity spin-forbidden bands. _1 Mn(C104)2 21,850 cm. ...1 -a Mn(BF4 )2 20,220 cm. 21,800 24,400 cm 31,500= Mn(02C.CF2)2 20,200 25,640

Iron III. The spectral data on this ion is sparse and contradictory mainly because of the uncertainty of the species actually present, especially in water. Jorgensen (1958) gives the first transition at 12,600 cm.-1 for the hydrate whilst Rubinowich and Stockmayer (1942) 1 give a value of 14,300 cm._ . This ion is one of the few which gives a charge transfer spectrum in water (at about 2380 I). It was found that a solution of ferric fluoroborate could not be prepared. On titration of ether solutions of silver fluoroborate withibrric chloride no colour appeared until excess of chloride had been added. Ferric fluoride is probably formed, the fluoroborate being unstable. A concentrated solution of ferric trifluoro© acetate in ether gave a strong absorption band at 10,750cmo 1 and inkenzene a band at 28,600 cm. - . Ferric perchlorate in ether gave a band at 29,500 cm. _a -128- Charge Transfer Spectra

All of the spectra discussed in the previous 0 section show intense, narrow bands at about 2,300 Ae Perchloric acid solutions in water do not absorb in this region (Wolfsberg and Helmholtz 1952) and studies on aqueous fluoroboric acid have shown that this too has no special characteristics from 2v1C7f.) 1. upwards (Scheffer and Hammaker 1950). We have found that trifluoroacetic acid has a charge transfer band at 48,700 cm. in water but in ether this moves up to 459 000 cm. . Perchloric acid in ether has bands at

459400 cm. and 389500 cm, $ in water these bands occur _1 at 449500 and 40,000 cm. In the case of the perchlorate ion Wolfsberg and Helmholtz (1952) suggest that it is an. oxygen rr electron which is transferred to a central ion, orbital, in a completely intraionic charge transfer.. The interpretation of charge transfer spectra is not at all clear yet but some success has been .had in explaining the charge transfer spectra of species such as

ReIO 2-, OsI6 2- and PtI62- (Jorgensen 1959) on molecular- orbital theory. The ibllowing table shows the observed charge-transfer bands of the ethereal solutions. All. values are in cm and the values in brackets indicate shoulders. V34 14112 + 1+ Cr3 e3 CO2 4' Ni2 Cue + Cu C104.— 43000 43,500 468500 45,500 45,400 459500 45,000 35,700 38,500 38,400 (37,000)

BF4 45,500 45,700 47,600 44,400 (41,200) 43,500 (389500) 39,30o (36,30o) 379 400

C CF3 45,500 45,500 45,400 46,500 45,50o 44,500 459 000 38,40o (43,500) 380300 (38 9400 (37,100) (37,000)

Tires a 44,400 44,400 37,10o Preparation of anhydrous Perchlorates, fluoroborates and trifluoroacetates. Perchlorates. Perchlorates of most metals are known)many in the form of hydrates. Until 'very recently silver perchiorate was the only known anhydrous lst row transition metal perchiorate. The salt is very hygroscopic and rapidly absorbs water to form the hydrate which can be dehydrated at 43°. Addison and Hathaway (1957) have prepared anhydrous cupric perchiorate by the repeated reaction of nitrosyl perchiorate oncopper perchiorate dihydrate, purification was effected by sublimation at 200° under vacuum between each reaction. In the present work anhydrous silver perchiorate and cupric chloride (excess) were shaken together in liquid sulphur dioxide. A green solution was obtained after 2 days, on removing the solvent and subliming the residue at 200° a pale green solid was obtained. It was very hygroscopic and dissolved readily in ether and benzene to give green solutions. These properties are the same as those described by Hathaway and Under hill (1961) for anhydrous cupric perchiorate except that Hathaway (1960) reports that it is insoluble in benzene. Longer periods of agitation and increase in temperature (80°) did not make the yield rise appreciably. r131-

Similar experiments with anhydrous nickel and cobalt chlorides gave positive results. The sulphur dioxide took on a faint yellow and pink colour respectively but the yields were so small that further attempts to isolate these pure were abandoned. However when ether was added to the resultant mixture after removing the solvent (SO2) the characteristic colour of solutions of the salts in ether were obtained. Results with ferric chloride were more promising, stoichiometric quantities of anhydrous ferric chloride and silver perchlorate react to give a light, orange precipitate in a brown solution. The colour of this precipitate changed to red-brown on removing the solvent. This was sublimed from silver chloride in vacuum and the brown sublimate gave an infrared spectrum with peaks similar to those obtained as by Hathaway and Underhill (1961).

Cu(C104)2 Fe(0104)2.6(NO3)0.4.0.2NO2C104 Fe(C104)3 1350 v.w. 1605 v.w. 1610 m 1270-1245 b 1285 b 1260 s 1130 s 1160 s 1130 s 1030 w 1018 w 1040 m 948 s 922 s 930 w 920 s 893 sh 850 m 665 m 662 m -132- However it was found that if the time of reaction was extended from 12 hours to 48 hours no orange precipitate was obtained; instead the solution was pale greenvitha light green precipitate plus a heavy brown-grey solid. No product could be sublimed out although the colour of the sulphur dioxide had suggested that reduction to krrous had taken place. An interesting phenomenon was noted whilst carrying out the preparations with the copper solutions. On allowing the green solution of Sulphur dioxide to stand for half an :hour or so the colour changed to reddish Irown. On shaking, the colour immediately reverted to green. A similar phenomenon has been noticed with potassium and caesium fluoroborates. A 10% aqueous solution of potassium fluoroborate is blue at 100° turning green at 90° and yellow at 60°. Precipitates of caesium fluoroborate show colour effects when the solution is warmed or diluted (De Boer and van Liempt 1927). This is explained as being due to the fact that the mean refractive index of CsBP4 is almost equal to that of the solution while the dispersion is different. Perhaps the explanation for the colour effects Observed for copper perchlorate in liquid sulphur dioxide are due to similar conditions existing. As the yielde obtained using silver perchlorate were so small reactions were attempted with the magnesium salt but no reaction would occur. At the present time there are no satisfactory methods of preparing anhydrous transition metal perchloratee. Hathaway and Underhill (1961) although able to prepare the pure cupric salt were unable to isolate the pure ferric salt. Reactions of anhydrous metal chlorides with anhydrous perchloiic acid result in the formation of hydrates (Hathaway and Underhill 1960) if there is reaction at all. The same authors found that anhydrous cupric nitrate would react with anhydrous perchloric acid but again the pure anhydrous salt could not be obtained. Removal of the solvent from a solution o an anhydrous perchlorate in an organic solvent (e.g. nitromethane ) is a possibility, but is fraught with danger.. Perchlorates are notorious for exploding when in contact with organic material. Even the preparation of a saturated methanolic solution of lead perchlorate has resulted in a fierce detonation (Willard and Kasner 1930). At the present time the best method of Preparation for anhydrous transition metal perchlorates- seemsto be by metathesis. Use of other, more suitable solvents, would speed reaction and longer periods of agitation would -134- ensure a complete reaOtion. Hydrogen fluoride may prove to be a good solvent for this type of metathesis. Starting from the fluoride and silver perchlorate (the former slightly soluble, the latter very soluble) reaction would be expected to proceed fairly easily. Of the products, the perchlorate would almost certainly be soluble as is silver'fluoride. The latter could rosily be removed from the product by washing with anhydrous methanol, and sublimation should give the pure perchlorate. Addison and Hathaway (1957) were the first to show the volatility of anhydrous cupric nitrate. Later it was shown that mercuric and ferric nitrate could also be sublimed (Addison, Hathaway and Logan 1958) although uranyl, zinc, cadmium, nickel and cobalt nitrates could not. The volatility of cupric and ferric perchlorates suggest that they can exist in the vapour state, as molecules. This would involve some degree of covalent bonding between the perchlorate ant metal ions. Normally the perchlorate ion is thought of as unlikely to form a complex with a cation and the bond, if formed, is usually considered as being purely ionic. As the perchlorate ion is tetrahedral it is reasonable to suppose that the bonding -135 involves one or more oxygen atoms for each perchlorate ion. Hathaway and Underhill (1961), from a study of the infrared spectra of anhydrous and hydrated perchlorates conclude that the perchlorate ions act as bidentate ligands and are strongly co-ordinated through two of their oxygen atoms in the anhydrous salts. It is possible that the solutions of anhydrous metal perchlorates in organic solvents contain bonding from the C104- groups as well as from the solvent molecules. Trifluoroacetates. Metathetical reactions, similar to those described in the previous section, were tried in an attempt to prepare anhydrous trifluoroacetates. Swarts (1939) claims to have prepared anhydrous copper and nickel trifluoroacetates by evaporation of aqueous solutions at 100° and states that they are volatile. The reaction of silver trifluoroacetate and cupric chloride in liquid sulphur dioxide proceeds rapidly and by filtration and removal of the solvent the pure green, anhydrous salt was obtained. Attempts at sublimation _5 (180;0/10 mm.) resulted in decomposition although Swarts (1939) was able to sublime the salt at 180°/1 mm. As expected the solid was soluble in benzene and ether. -136- When cobaltous Ohloride was used there was a reaction but the mauve trifluoroacetate was not very soluble and precipitated out. Fourteen days shaking was necessary for complete reaction. A mauve solid was obtained on , sublimation at 185'7 10 mm. but with considerable decomposition. The anhydrous salt could be extracted from the reaction mixture by continuous extraction with benzene and subsequent removal of the boiler liquid. Unfortunately there was considerable contamination by very finely divided silver chloride which had been washed through the extraction thimble. This method of purification would be suitable for perchiorates but again there is the difficulty of removal of silver chloride. Similar results were obtained for Ni", Cr$+ and Fe" reactions in sulphur dioxide. With trifluoroacetates removal of an organic solvent is not hazardous as it is with perchlorates. In the present work it was found possible to carry out the metathetical reaction in nitromethane and to remove the solvent completely by heating under a vacuum. Cupric, cobaltous„ nickelous, ferric and chromic trifluoroacetates were obtained in this manner. The infrared spectra of the trifluoroacetates were studied in the region 1F -050p sad in the following tables the peaks observed are compared with silver trifluoroacetate and trifluoroacetic acid. The spectrum of the latter was taken as a liquid film between silver chloride plates. Infrared in the region 2p to 15 )p

AgOgC.CF3 1677sh 1629s 14542 1207s 1197s 1145s 858s 795s 7348 cu(0,4c.c113). 4350(m) 1715s 1663sh 1622sh 1197-1157(b) 1098w 1053m 1015w 956m 858s 792s 732s 0o(0200F8)2 4350(w) 3545(m) 1668(8) 1220-1150(b) 1088w 972(q) 853e 792s 725s N1(000113)a 4280(m) 3460m 1705sh 1658s 1443w 1196e 1156s 1061m 853s 794s 728o Pe(0aCCF3)5 4290(w) 3330m 1666b 1563w 1196s 1156s 1014w

0r(020.0P3)3 4330(m) 3330m 1666b 1563w 11960 1156s 1020m 996w 862s 787e 730s -138. Infrared in the region 15i to 50)1 CF3CO2H 6908 597m 510m AgO2CF3 690s 597m 516s 4514 314s 2818 264sh Cu(0200F02 7.09s 606w 516s 503s 319s 250s Ni(02C0F3)2 694s 614s 526s 467s 327m 303s 256m 235b Co(0200F3)2 7048 610s . "524s 465s 395w 307w 286m Fe(02C.0P02 704s 616w 553sh 519M 4938 371w 342m 301s 233m 0r(02C.0F3)3 696m 621m 572m 526s 3710312b

The bands at about 1677 and 1623 cm. are due to asymmetrical CO2 stretches and the band at about 1450 cm. to symmetrical CO2 stretches. C-F bands definitely occur at about 12009 11909 1145, 597 and 470 cm. C-0 stretch _11 appears at about 985 cm. . The preparation of anhydrous transition metal trifluoro acetates by metathesis in nitromethane has been established as a good method of preparation. Metathetical reactions in liquid sulphur dioxide proceed readily only with Cu2+ ion This latter method is' better as the removal of the solvent is much more ready and contamination less likely. -139- Fluoroborates. No anhydrous first row transition metal fluoroborates are known other than the silver salt. Metathetical reactions involving silver fluoroborate and the anhydrous metal halides do not occur in liquid sulphur dioxide. Attempts using lithium fluoroborate (Shapiro and Weiss 1953) instead of the silver salt met with no success even when heated at 90° for 24. hr. An impure sample of cupric fluoroborate could be prepared by metathesis in nitromethane and then removing the solvent in vacuum. Some decomposition always occurred on attempting to remove the last traces of solvent. The salt could not be sublimed.

MagnQtie stUdieS. Measurements of the magnetic susceptibilities of the ethereal solutions of the perchlorates were made using the method of Evans (1959). The position of a line in the proton resonance spectrum of a molecule depends on the bulk susceptibility of the medium in which it is situated. For an inert substance in solution the shifts caused by paramagnetic ions are given by the expression AHM = 27,-/3. AK (Dickinson 1957) where AK is the change in volume susceptibility. _140- If a dilute solution of a paramagnetic substance containing about 2% of an inert reference (e.g. cyclohexane) is put in a nuclear magnetic resonance tube together with a capilllryp containing the solvent and the same concentration of the references on spinning the tube two proton resonance lines will be obtained from cyclohexane. These lines are obtained because of the difference in the volume susceptibilities of the solution and of the solvent,with the line from the more paramagnetic solution lying at higher frequencies. The mass susceptibilitytX 9 of the dissolved substance is then given by the expression ) rffm 4. X ° ° • (d ° Mciu where of is the frequency separation between the two lines in cycles sec. 9 f is the frequency at which the proton eg resonances are being studied9 in cycles sec. 9 m is the mass of substance contained in 1 mlo of solution9 0 is the mass susceptibility of the.solvents, do is the density of the solvent and dB that of the solution. Using cyclohexane or tetramethylsilane as the inert reference the -following results were obtained for ether solutions of the perchlorates of approximately 0.1M strength. Ion No. of Spin only Found for Range usually unpaired moment perchlor- found expert-' electrons ates mentally for ion. 5+ V 2 2.83BM 2076 BM 2075 - 2085 0+ Cr 3 3.88 4.43 3.70 - 3.90 a+ Mn 5.92 5.45 5.65 - 6.10 a+ Fe 5 5092 5.97 5.70 - 6.0 2+. CO 3 3.88 3.88 4.30 - 5.20 2+ Ni 2 2083 2.81 2080 - 3.50 a+ Cu 1073 1.81 1.70 - 2020

The magnetic moments of the solid trifluoroacetates described in the previous section were measured using the Gouy method and the moments are summarised Cu(020CF02 1.75 B0310 co(02cop02 4.77 Ni(o2ccp02 3.21 cr(oac.cp3).a 4.65 All the measurements indicate that the compounds are spin-free, although the high moment of chromic trifluoroacetate is unusual. -142-

EXPZEIMENTAL. The spectra of the solutions were taken on an Optika spectrometer in 1 cm. stoppered cells. Infrared spectra were taken on a Grubb-Parson S4 spectrometer in the NaCl region9 and on a Grubb Parsons DM2 spectrometer in the 15y -->50? region using caesium iodide optics.

The metathetical reactions involving liquid sulphur dioxide as solvent were carried out in Carius tubes. The tubes were baked out in vacuum at 450° and then filled with the solid reactants in a dry box, In all cases, except for the experiments involving ferric chloride9 excess of the metal halide ( 2 gm.) and about 1 gm. of the silver (or magnesium or lithium) perchlorate9 fluoroborate or trifluoroacetate were used. For ferric chloride the stoichiometric quantities were used. Sulphur dioxide (from a siphon) was purified by passing it through concentrated sulphuric acid and then through a 2 foot column of phosphorus pentoxide. The gas was then condensed on phosphorus pentoxide in a moisture protected u-trap cooled in liquid air; it was then redistilled on the vacuum line into the Carius tube. After sealing off the tube it -143- was placed in a metal sheath and shaken horizontdlly in a mechanical shaker. If heating was necessary a heating tape was wound round the sheath. At the end of the reaction the solvent was distilled off in vacuum. If filtering was necessary (as in the case of cupric trifluoroacetate) the Carius tube was frozen down and the tube fitted with drying tube was then opened. When all the solvent had liquefied it was filtered rapidly in the apparatus shown in diagram 4. Diagram 4 The evaporating solvent provided the force necessary to effect filtration; the lower half of the apparatus was then connected to a vacuum line and the remaining solvent distilled off.

orous glass disc, porosity no. 4. <—B14 cone -144-

Cupric Perchlorate. Cupric chloride (2 gm.) and silver perchlorate (1 gm.) were shaken in liquid SO2 (20 ml.) for 2 days9 after which time a pale green solution had formed. The solid obtained after removing the solvent was sublimed at 2000/10-5 mm. A light-green solid was obtained which fumed in moist air and was soluble in ether and benzene. Ferric Perohlorate. Ferric chloride (0.30 gm.) and silver parchlorate (1.14 gm.) were shaken in liquid SO2 for 18 hours. A bright orange precipitate formed in the light brown solution. After removal of the solvent the red-brown residue was sublimed at 140° for 5 hours under vacuum. The white sublimate was removed and the residue reheated at 180° , -0 for 12 hr. Finally the residue was sublimed at 230°/10 mmo and dark red-brown crystals were obtained. Cupric Trifluoroacetate. Cupric chloride (2 gm.) and silver trifluoroacetate (1 gm.) were shaken for 15 hours in liquid 302. The bright green solution was filtered and the solvent removed to give pale blue-green crystals of cupric trifluoroacetate. Analysis: Theoretical Cu = 22.1% Found Cu = 21.7% Magnetic measurement gave a magnetic moment of 1.75 B.M. Silver trifluoroacetate (1 gm.) and excess anhydrous cobaltous chloride were shaken for 2 days in anhydrous nitromethane. Afterascertaining that there were no chloride or silver ions in the solution it was filtered in the dry box. The solvent was then removed at room temperature under vacuum and the resulting sticky mass washed thoroughly withd ry 40-60 pet. ether. After heating at 70° under vacuum for 15 hr. the temperature was raised to 150° for five further hours. Beautiful mauve crystals were obtained. Analysis: Theoretical Co = 20.71; Found Co = 20.5% Magnetic moment 4.77 B.M. ickel Trifluoroacetate. This was prepared in exactly the same way. Apple-green crystals were obtained. Analysis: Theoretical Ni = 20.7% Found Ni = 21.3; Magnetic moment 3.21 B.k.

per.iToroaoetate. This compound Vale prepared by adding stoichiometric quantities of ferric chloride and silver trifluoroacetate dissolved in anhydrous n:.tromethane

to each other. A beautiful red-brown solution was obtained and by removing the solvent as already described brown crystals of ferric trifluoroacetate were obtained.

Analysis: Theoretical Fe = 14.2% Found Fe = 14.6%

Chromic Trifluoroacetate. Using the same method as for cobalt trifluoroacetate dark green chromic trifluoroacetate was obtained from chromic chloride and silver trifluoroacetatee Analysis: Theoretical Cr = 13.3% Found Cr = 1309

Magnetic moment = 4.65 B.M.

Maanetic Measurements. Measurements on the ethereal solutions were made on a varian V A40 spectrometer at 56.4 Mc.sec. The spinning sample tubes had an external diameter of 5 mm. (i.d. 3.5 mm.)i the capillaries were constructed so that the volume of liquid in the capillary was approximately the same as the volume of the annulus i.e. 2.5 mm. external diameter. The capillary was filled with a 2% solution of cyclohexane (or 7-147—

tetramethylsillne) in ether and the annulus with the ethereal solution of the perchlorate containing 2A of the reference. Equal concentrations of the reference in the two solutions ensure that the proton lines are of equal intensity. The spinning sample tubes were adequately protected from moisture. Line separations were measured by the conventional side band technique using a Muirhead decade oscillator. The solutions were analysed by the normal analytical methods. A measured volume of the ether solution was poured onto 50 ml. water and the ether was then gently evaporated. After filtration analyses were carried out in the normal manner. Magnetic measurements on the solid trifluoroacetates were carried out using the Gouy method. An Oertling magnetic balance capable of weighing to 0.01 mg. was used. The specimen was suspended in the usual way by means of an aluminium collar hangIng on a nylon thread. The specimens were packed in a pyrex tube (i.d. 3.5 mm.) fitted with a tightly fitting ground glass stopper. All packing opera- tions were carried out in the dry box. The permanent magnet was put into position by means of an ordinary hydraulic jack. The strength of the magnetic field was found by calibration with Hg(Co(CNS)4]. -148—

CHAPTER - 4 Phosphene Fluoride Complexes_ ThiS work Was undertaken in an effort to prepare some new phosphine fluorides. The first part of this chapter is devoted to a-study of the complexes that phosphorous pentafluoride forms with the triar.yls of the Group V elements.- Also included in this work was . triphenylphosphineoxide and -iris (diethylaminomethyl ) phosphine and its oxide. Phosphorus pentafluoride has the general acceptor properties of a. Lewis,acid and Muetterties et alia (1960) have prepared a. number of complexes with bases such as ethers, sulphoxides, amines amides and esters, in all cases 1:1 complexes were formed. However the stabilities of these complexes are much less than the corresponding boron trifluoride complexes e.g. all.the BFu o amine complexes can be recrystallised from water but all of the PP5 complexes are readily decomposed. by water and alcohols. The lower stability of the PA'S complexes hes been explained as being partly due to steric factors. P5. 4- is only slightly larger than Ba + and .yet must.accomodate two more fluorine atoms around itself. All the experiments were carried out under pressure 149- with anhydrous sulphur dioxide as solvent. Of all the complexes formed only that with triphenylphosphine (1:1) was stable and could be recrystallised from methylene chloride without decOmposition. The complex with triphen ylphosphine oxide was found to decompose fairly slowly but an accurate analysis could not be obtained. The triphenylarsine complex was unstable and very sensitive to moisture and decomposed rapidly under vacuum. Triphenyl- bismuth reacted with phosphorus pentafluoride to give a whitish precipitate but if the reaction was allowed to proceed for 24 hours a purple colour appeared which on freezing down changed to green. Only pastes could be obtained from the reactions with tris(diethtlaminomethyl) phosphine and phosphineoxide. The reaction with triphenylamine gave a dark blue solution in liquid sulphur dioxide which after a few minutes turned to dark green. The dark blue colour is characteristic of the triarylaminium ion. Green crystals were subsequently formed which were fairly stabled however attempts to recrystallise the solid from chloroform resulted in decomposition. The very strongcolours observed in solution suggest that a charge-transfer complex was formeu,? Nuttall and Sharp (1962) have -150- found that /1731F4 and s35N.AgC104 are both pale green solids and Sharp. (1958) that the triphenylammonium salts of complex chloro - and bromo - acids are generally pale green solids. This suggests that triphenylamine was oxidised. Of the series ofhP„ i4As9 yJ533i9 triphenylphosphine has been shown, by dipole measurements, to be most tetrahedral in shape ( Klages and Langpape 1959 ). Jaffe (1952) has given spectroscopic evidence that shows that, if the central atom (lip,As„Sb,Bi) has an unshared pair of electrons, extensive conjugation occurs between the phenyl radical;leading to spectra in sharp contrast to those of the simple benzene derivatives. Tripheny- lamine shows a greater degree of conjugation than the other derivatives, and triphenylphosphineoxide has been shown to have practically no conjugation in its system at all. From this evidence it is expected that triphenylphosphineoxide should be the best donor. The complex AP:PF5 is found to be stable but the complex AP=0 PF5 decomposes slowly. The instability of the AAs, 'Sb,,,04/31 complexes should increase from As-PBI as the donor molecule becomes less tetrahedral, and this order was found to apply -151—

experimentally; The instability of the triphenylamine complex is not surprising as Sasaki, Kimiwa and Kubo (1959) have shown that the triphenylamine molecule has an almost flat trigonal pyrimidal structure with a short C—N bend which emphasises the conjugation of the molecule. Muetterties et alia (1960) have shown from 19F n.m.r. studies on the sulphoxide„ amine and thioamine complexes with PFts that the observed spectra are in agreement with a structure involving four coplanar equivalent fluorine atoms and one apical fluorine atom. it is therefore, reasonable to assume that the triphenylphosphine adduct has a similar octahedral arrangement. -152-

The infra-red spectrum of AP:PF5 showed a strong absorption band between 870 and 816 on-1 with peaks at 854 and 828 cm-1. The spectrum of 51 3P=0.PF3 showed a a, • broad band with a m;ximum at 840 cm-1 . The spectra of all the pastes obtained from the other aryls showed a broad band at approximately the same position. Daasch and Smith (1951) have tentatively assigned the P:P link in pentavalent PF compounds to 900-850 cm-1. Phosphine-Metal fluorides Although phosphine-metal halide compounds are well known when the halogen is bromine chlorine or iodine there are no previous reports of any fluorides. The only compounds knowivhich both phosphorus and fluorine are bonded to the same metal atom are the phosphorus trifluoride derivatives PF3 OsF2 y PF3IrF20 PF3PdF2 y PF3CoF ( Peacock 1960) and even here there is no structural confirmation to substantiate this formulation. In an attempt to prepare such complexes it was decided to study the reactions of some transition metal fluorides with triphenylphosphine in non-aqueous solvents. Sulphur dioxide was the solvent chosen as it is a good solvent for some of the higher valent metal fluorides. Accordingly tungsten•hexafluoride and triphenyiphosphine were dissolved in liquid sulphur dioxide. The colour of the solution -153- formed was bright red which then gradually faded to an orange colour precipitating pale yellow crystals over a period of two days. The red colour observed has been noted by several other workers when studying the solubility of the hexafluoride in organic solvents (Priest and Schumb 1948; Fischer 1930). After recrystallisation the solid was shown by analysis to be an almost 1:1 complex. Unfortunately at this time Muetterties (1961) published a paper on precisley the same subject in which is described a 1:1 complex WF6,AP. From n.m.r. studies a structure :-

- is proposed. Support for this formulation comes

m••••••.4 from the fact that tungsten hexafluoride compleXes of this ligand and of amines do behave as electrolytes in liquid sulphur dioxide. The infra-red spectrum is fairly complicated but there are three main bands not attributable to triphenylphosphine at 1015; 970 and 845cm-1. As a result of ,154-

Muetterties' publication further work on complexes of this type was abandoned. During the last two or three years many molecular hydrides have been prepared showing that metal-hydrogen linkages are much more common than was previously thought (Green 1960). Hydride species have been established for the metals of transition Groups V to IX. Nuclear magnetic resonance spectroscopy has been extensively used to observe the metal-hydrogen linkages, because there is a characteristic high field shift for these protons. Using this technique it has been shown that the metal atoms in many metallocenes, carbonyls and phosphinemetalcarbon- yls can be protonated in strong acid solutions such as concentrated sulphuric acid, trifluoroacetic acid and boron trifluoride monohydrate ( Davidson and Wilkinson 19609 Curphey et alia 1960). Reaction of phosphine® metal halides and cyclopentadienylmetal halides in solution with borohydride also result in hydride formation (Green, Streetand Wilkinson 1959; Fischer and Hristidu 1960). A more interesting method of producing hydrides is that using a mixture of a transition metal halide, a substituted phosphine or arsine and an alcohol (usually -155-

in the presence of an alkali). The products vary according to the conditions used and a wide variety of hydrides and carbonylhydrides have been prepared (Chatt and Shaw 1960; Vaska 1960; Vaska and Sloane 1960). The use of hypophosphorus acid has also resulted in the formation of hydrides such as [ Rh C12.H (AsPh2Me)2]. ( Lewis, Nyhoim and. Reddy 1960). More recent develop- ments with ditertiary phosphines and arsines have led to the preparation of Nickel (o), palladium (00 vanadium (o), chromium (0, molybdenum (o) and Tungsten (o) complexes in which there is not a metal-hydride bond (Chatt, Hart and Watson 1962, Chatt and Watson 1962). Greene,Strestand Wilkinson (1959) have shown by nuclear magnetic resonance measurements that a borohydride reduced solution of bis (tri-n-propylphosphine) nickel- dichloride in te'trahydrofuran gives a high field peak attributed to a hydride species. In the present work it was found that when anhydrous hydrogen fluoride gas was passed through such a solution a precipitate of nickel fluoride was obtained. Kemmitt (1961) had observed that when a methanolic solution of silver fluoride was added to a solution of (n-Pr2P)2NiC12 in the same solvent a Colourless solution was obtained. This was repeated and the precipitate was found to be a mixture of silver chloride and nickel fluoride. Because of the failure of these two methods the action of anhydrous fluoride on the phosphine nickel chloride was investigated. Using a variety of solvents e.g. anhydrous ethers tetrahydrofuran and chloroform it was found that nickel fluoride was formed in each case. Bis(triphenylphosphine) nickeldicarbonyl dissolved in tetrahydrofuran and in chloroform also gave nickel fluoride on treatment with gaseous hydrogen fluoride. If liquid anhydrous fluoride was used on the dicarbonyI a black tar was formed. Nickel fluoride was identified in each case by X-ray photography. From these results it was concluded that the nickelphosphinefluoride, if formed, must be unstable readily decomposing to the very stable nickel fluoride. Attention was next directed to the 2nd and 3rd row transition metal complexes. Malatesta and Cariello (1958) have reported the preparation of some platinum (o) complexes with triaryl phosphines and arsines. The reaction of triphenylphosphine with either the dihalogenobis( triphenylphosphine) platinum (II) compound or potassium tetrachloroplatinate in ethanolic potassium hydroxide gives alleged tetrakis ( triphenylphosphine) platinum (o). Further investigation into this complex by Chopoorian, Lewis and Nyholm (1961) led them to - . formulate it as a hydride of divalent platinum. It was also shown that in solution the complex lost first one and then two molecules of triphenylphosphine. This compound and (A5P)2PtH2 were prepared using the method of Chopoorian, Lewis and Nyholm (1961)-/0 the infra-red spectrum of bis(triphenylphosphine)-platinum (II) hydride showed bands at 1672 and 815 cm-1 in agreement with Chopoorian et.alia (1961). The [AP]I PtH2 dissolved in benzene to give a yellow solution; when anhydrous fluoride was passed through this solution for 15 minutes it underwent a series of rapid colour changes from yellow-igreen-oblue-wellow with the precipitation of a white solid. Triphenylphosphine was recovered from the benzene solution. The precipitate was found to be insoluble in all non-polar solvents but to be extremely soluble in chloroform, methylene chloride (lgm in lml. ) and nitrobenzene. After recrystallisation the white compound analysed to the formula bilaPj2 PtF2. The same product could be obtained by dissolving the hydride in liquid anhydrous fluoride. After leaving for -158-

15 minutes the acid was removed and the resulting solid recrystallised. Preliminary measurements on the 191 and31P n.m.ro spectra of this compound in methylene chloride support the formulation [AP]2 PtF2. However as the time taken to run a complete spectrum was about 5 hours and the compound exchanges with the solvent ( at high concentrations) over a period of time further measurements will be necessary in another solvent. The infra—red spectrum shows two sharp bands at 3682 and 2890 cm-1 and a strQn6 band at 1094and 1045cm-1. We formulate the structure of this compound as square planar (dsp2) with the two triphenylphosphine ligands cis to each other, although we cannot substantiate this with dipole measurements due to the insolubility of the compound in non—polar solvents. The reasoning for this formulation comes from analogy with the corresponding dichloride and also from the fact that this arrangement is that which the 'trans° effect would predict. The a bond forming part of the double bond between a donor and a transition metal is visualised as donation of a lone pair from the donor to the metal. The dative rrbond is formed by overlap of. a. filled d or dp hybrid of -159- the metal with a vacant p„d or dp hybrid of the donor

-3s3e

When the phosphine fluoride is heated gently in a solvent the colour deepens to red and a pink solid can be obtained. On recrystallisation several times the original white compound is recovered° This suggests that there is sOme isomerisation and the trans form is prepared although it is unstable.° '31) F heat /0P Pt recryst.s4

The compound (AID)2 PtH2 is also expected to be square planar with the hydride ions cis to each other° The substitution reaction takes place readily due, no doubt, t4i the fact that the triphenylphosphine group is strongly :trans directing and also to the fact that the -160- 0 F-,ion ( radius 1.34 A ) is much smaller than the H- . ion (2.12A) and thus more polarising. This greater polarising power of the F- ion should increase the strength of the Pt -P bond and the compound (4P)2. PtF2 has been found to resist attack by nitric acid and alkali. As a consequence of the strong bonding in the plane of the moleculel platinum metal should be electron deficient above and below the plane, and perhaps therefore able to accomodate other ligands. When bis ( triphenylphosphine) platinumdifluoride dissolved in anhydrous hydrogen fluoride is left under a pressure of carbon monoxide for 12 hours a pale yellow compound is formed. Analysis shows it to be (AP)2PtF2 (00)2. The 6 co-ordinate platinum II compound shows the same bands in the infra-red as does the difluoride and in addition has a strong carbonyl band at 2152 cm-1 With shoulders at 2105 and 2083 cm®g. Booth and Chatt (1962) have prepared similar compounds with on and cobalt phOsphine halide complexes. Their results show that, in general, the complexes of the more aliphatic phosphines react more readily with carbon monoxide to give products of greater stability than those with -1SI-

aromatic phosphines. In fact ( AP)2 Fe (CO)X2 ( where X = Cl,BrpI ) could not be prepared Cl PR3 The configuration is given on the basis of 1;R

Cl U spectra(2 sharp bands in the PR3 region of 2000 cm—') and on dipole moment measurements. Cobalt complexes could be prepared but with only one carbonyl added ono -In the configuration of the iron complex the most strongiy trans directing ligands are placed opposite the weakest ( see later section for I.R. data). Valatcstaand Cariello ( 1958 ) prepared a compound designated Pt (C0)2 (Pp3)2 from Pt(,4P)4. However as

the latter is now known to be PtH2 (,eP)4 it is very likely

that the carbonyl is really Pt(C0)2 /12 (JP)2. This solid slowly loses one molecule of carbon monoxide; the loss is more rapid in solution. A similar reaction takes place with PtH2 (4P)3 under 300 atn carbon monoxide; Pt(,KP)2(e0)2112.AP is presumably the product. On dissolution in ether a mole of carbon monoxide was evolved .1(1 on addition of alcohol orange—red crystals of (R(3P)3. PtCO ( according to Malatesta and Cariello ) were obtained; -162—

This could be (AP)3 PtCO A thorough investigation of the magnetic and spectroscopic properties of these compounds would be well worthwhile. The easy removal of one molecule of carbon monoxide suggests that they are attached perpendicular to the plane containing the hydrides and phosphines. CO P Pt P

By analogy with the carbonyl complexes prepared by Booth and Chatt (1962) the configuration of the platinum phosphinecarbonylfluoride is expected to be

The bonding in this compound can be described as ad 6s 6p2 for the planar part of the molecule containing the phosphines and carbonyls and the other two bonds as being a hybrid of the 6d and the other 6p orbital. -3.63-

These two bonds are expected to be weaker due to the fact that the other filled 5d orbital must point in the same direction. This explanation has been used by Harris and Nyholm (1957) to explain the bonding in [Au(Arsine)2 12 ]+. There are eight non-bonding electrons in Pt2+9 tug eg, six are located in the akelevel and two in the tb,i level The 5d x24.orbital can be assumed to combine with the 6s96px, 6py orbitals to give four planar hybrid bonds 5d6s6p2. The filled 5d22 orbital points towards the two remaining positions in the octahedron. Any ligands along the H axis are therefore expected to experience a replusion. X-ray results on the gold diar#sine complex by Harris and Nyholm (1957) show that two of the bonds are longer than usual, supporting this idea. In ligand field theory a 5p6d hybrid is predicted as being longer than a normal covalent bond using 5d orbitals. This bonding has been used to explain the configuration and diamagnetism of [ Pt(NH2)4(CH20102]012 (Harris and Stephenson 1957 )0' X-ray structural determinations have shown that the molecule has a configuration where the platinum nitrile distance is about 0• CHAgisi 3A as compared with 21 H NH3 usually found for Ptil distances. N- C On molecular orbital theory d2sp3 grouping can be used but here the two extra electrons from the platinum e9 orbitals must be placed in the antibonding d22 and d - 1,2 orbitals; rrbonding involving the non—bonding t29 orbitals increases the magnitude of Dq and is responsible for causing electron pairing in complexes. As bis (triphenyiphosphine) platinumdifluoride can addon carbon monoxide it was decided to try and prepare an analagous dichloride. When [051]2 PtC12 was heated at 120° for 15 hours under a pressure of carbon monoxide a bright green solution was obtained. A green crystalline solid was isolated with a very strong carbonyl peak at 2145 cm—i. However it was not possible to purify this .ompound. Attempts at purification by chromotography caused decomposition but the presence of at least three components was shown. One was identified as the starting material. Malatesta amd Angoletta (1957) have reported the preparation of Pd(0). (p3P)4. It has been observed that this yellow compound loses first one and then two molecules of the phosphine when dissolved in benzene. (fa;P)4 Pd (4P)5Pd ,T3P

(AP)2 Pd (O;P)2Pd These properties are exactly the same as those described for the platinum hydride complex. Fischer and Werner (1962) have prepared . the same compound from 05115Pd.C6H9 and triphenYlphosphine and have found a dipole moment of 2.30 ±0.01 D which is surprisingly large for a molecule allegedly symmetrically substituted. lieasure- ments of the molecular weight by cryoscopic methods gave a value of from 491 to 523. In the present work attempts to prepare the complex according to Malatesta and Angoletta (1957) were not very successful. Reduction of a suspension of bis(triphenylphosphine) pa]ladium dichloride in alcohol with an excess of phosphine was effected by the action of sodium borohydride. The yellow compound obtained was sensitive in air turning orange fairly quickly ( as was also reported by Talatesta and Angoletta.). The infra-red spectrum showed bands at 1962, 1886 and 1808cm-1.. Although the complex did not seem to react with carbon tetrachloride a dilute solution of bromine in the same solvent was immediately decolourised by it. On the evidence of the infra-red spectrum, its decomposition in solution, its high dipole moment and reac'ion with bromine we suggest that the compound should in fact be formulated -166- as [ p3P ]4 Pd(1)112. The compound reacted vigorously with anhydrous hydrogen fluoride to give a dark red solution. A tan coloured compound was isolated on recrystallisation from methylene chloride/ ether . The infra-red spectrum showed a strong double band at 1092 and 1053 cm-1 and sharp bends at 3080 and 3450 cm-1° Analysis showed that the compound is (02P)2 PdP2. Vaska (1961) has described a series of iridium hydrido complexes with triphenylphosphine and arsine. The preparations are very simpler involving heating (NH4)2IrC16 in an alcohol'or aqueous alcohol with the ligand. Depending on the temperature used for reaction the mono or dihydro complexes are formed.. Hayter (1961) extending this work has shown that the trihydride IrH5 (p2r)2 is formed by reduction of the monohydride IrHC12 (p P)3 with lithium aluminium hydride. The preparation of the trihydride has also been described by Malatesta9 Angoletta9 Araaneo and Canziani.(1961) by reducing a mixture of iridum tribromide and triphenylphosphine in alcohol with borohydride. These authors separate the two isomers I and II by crystallisation. 41)

I. I.R. bands at 2130 II. I.R. band at 2075 c.m.-1. and 1750 c.m.-1'

In the present work, using the method of Malatesta et alia (1961) the trihydro complex was prepared and dissolved in anhydrous hydrogen fluoride. A pale yellow compound was isolated from this reaction. This complex had a peak at 2125 cm-1 in the infra-red which must be due to the hydride and it nlso had a strong double bind at 1086 and 1047 cm-1. Analysis indicates that the compound is (AI% ITHF27 which is reasonable because other workers have found that two of the hydrides are easily replaced but the third is not. Similar rheniumphosphinehydrides have been prepared (Malatestav Freni and Valenti 1961 ). Reduction of bis(ttiphenylphosphine )- rheniumdiiodide wish sodium borohydride in ethanol gives red crystalline ReH3 (P03 )2. . - It was found in the present work that this reduction always caused considerable decomposition and the yield of hydride was always small. The red hydride reacted with hydrogen fluoride to give a small quantity of dark green crystals. The infra-red spectrum showed strong absorptions at 1093 and 1025 cm-1. Unfortunately on repeating this experiment several times only a black mass could be obtained. Reaction in a solvent such as benzene was equally unsuccessful. On bubbling the hydrogen fluoride gas through the solution the colour changed from red to green but again on attempting isolation of the complex a black mass was obtained. These experiments will have to be repeated as isolation of the fluoride should be possible. Infra-red spectra :The platinumphosphinefluoride and carbonylfluoride show a strong double band at 1045 and 1094 cm-1 and two sharp bands at 2850 and 3675 cm-1. Triphenyl phosphine has bands at 1025, 1069 and 1087cm-1 ; these bands are very sharp and of medium intensity. In the spectrum of the dihydride ([ 541:14 PtH2 ) these three bands occur but at lower intensity. A similar situation exists in the case of the palladium complexes. The dihydride shows three sharp bands of medium intensity at 1080; 1064 and 1023 cm-1 whereas the difluoride has a strong double band at 1092 and 1053 cm-1 plus a sharp band at 4310cm-1 . Ir11304P)3 -169- has sharp peaks at 1024, 1074 and 1087 c.m.-1 whereas the difluoride has a strong double band at 1090 and 1050 0.m.-1 plus sharp peaks of medium intensity at 3450 and 3080c.m.-1 Pinally the rhenium phosphine fluoride has a strong double band at 1092 and 1038 c.m.-1. The platinum phosphinecarbonylfluoride showed a strong carbonyl peak at 2152 c.m.-1 with much weaker shoulders at 2105 and 2083 c.m.-1. The following table compares the carbonyl frequencies of some phosphine metalcarbonylhalides. PtF2(C0)#03 )24- 2152 2105 2083 c.m.-1 PeC12 (C0)2(YEt3 )2 2014 1963 PeBr2 (C0)2 (PEt3 )2 2009 1958

FeI2 (C0)2 (PEt3 )2 2003 1953

FeC12 (C0)2 (PEt2y5) 2 2025 1972 FeC12 (C0)2( 2)2 2033 1980 (AP)2PtC12 + CO 2145 + present work: as Booth and Chatt (1962). Spectra of the platinum and palladium phosphine- fluorides in the region 164p--.)251 show the following bands. -170-

(10)2PtF2 .1P (05P)2PdF2 553 c.m.-1(m) 543 c.m. 1(m) 539 c.m.-1(m) 517 (s.) 508 (sh) 519 (s) 49? 510 (e) 481 481 (2) 482 (sh) 448 430 (w) 44? (m) 422 422 (w) 422 (m)

A method of preparing phosphine-metal-fluorides of the 2nd and 3rd row transition metals has been established in the present work using liquid hydrogen fluoride. It is doubtful if (41))4Pd exists; our results show that the formulation (541))4PdH2 is probably the correct one. 1 rI*1,

EXPERIMENTAL.

The following procedure was used for all preparations involving phosphorus pentafluoride and the phosphines phosphine oxides, arsines, bismuthines and stibines. Phosphorus pentafluoride was prepared by the thermal decomposition of 'phosfluorogen A' under vacuum. The gas Passed through two long traps(filled with glass wool and cooled in cardice acetone baths) and was then condensed in a u-trap cooled with liquid air. The fluoride was redistilled into a dry Carius tube containing the other reactant. Dry sulphur dioxide was then condensed on top of the mixture and the tube sealed off. At the completion of the reaction the excess fluoride and the solvent were removed under vacuum. Triphenylphosphine-phosphorous pentafluoride. An orange solution was obtained as the sulphur dioxide warmed up: after a few hours pale yellow crystals were deposited. The solvent was removed after 1 day, and the product washed with ether and then recrystallised from methylene chloride. Pale yellow crystals were obtained. Analysis. CieH/5P2P15 : Calc. C=55.6%, H=4.5% Found =55.0% 3.9%

Triphenylphosphineoxide-phosphorous ppm:Anna-y-1dg.. A rtn1 pink solid was obtained from the yellow solution formed after 24 hours. The product slowly decompCses and accurate analyses were ,repossible. Analysis. C1811151'204. Cale. C=53.5%, H=3.2% Found =51.3% =407% Tungsten hexafluoride-triphenylphosphine. '..-ungsten hexafluoride was prepared by heating tungsten in a stream of fluorine at 300° on a conven tional fluorine line. The volatile fluoride was collected in a trap and at the completion of the fluorination it was redistilled into a storage bulb which was then sealed off. Subsequently the fluoride was distilled into a Carius tube containing rigidly dried sulphur dioxide and triphenyiphosphine. As the solvent warmed up a deep claret coloured solution formed and after an hour yellowish crystals started to form. After 2 days the pale orange solution was frozen down and the solvent carefully removed under vacuum. A white product was obtained which could be crystallised from methylene chloride. Analysis. 'C18H151iF6. Cale. C=38.6%, H=2.7% Found =38.0% =2.1% Reactions of (n-nr/P),NiC1?. Bis(tri-n-propylphosphine) nickel dichloride was prepared by the addition of phosphorous trichloride in ether to a Grignard solution made from n-propyl bromide, magrieGitliA and ether according to the directions of -11/3-

Davies, Pierce and Jones (1929). The mixture was then cooled to 00 and cautiously treated with a solution of ammonium chloride in water. The ethereal solution was dried over anhydrous sodium sulphate and the ether distilled off inan atmosphere of nitrogen (Davies and Jones 1929). The phosphine was distilled under reduced pressure; the distillate was allowed to run into an alcoholic solution of nickel chloride. Bright red crystals of bis(tri-n-propylphosphine)nickel dichloride precipitated out. In all the reactions involving anhydrous hydrogen fluoride polythene apparatus was used. Flat-bottomed bottles with screw caps, fitted with welded inlet and outlet tubes, were used. For work involving solutions the gas was just bubbled through; if liquid anhydrous hydrogen fluoride was used, the gas was condensed by surrounding the bottle with cardice. After reaction the liquid HF was xemoved by a rapid stream of dry nitrogen. In the reactions where nickel fluoride was a product, X-ray powder photographs were taken to identify it. Infra red spectra were also taken. Bis tri hen 1phosphine)platinum difinnriaa Tetr-kis(triphenylphosphine)platinum dihydride was prepared as directed • by Malatesta and Cariello (1958). Potassium tetrachloro- -174- platinate (2.2g) in water(5m1) was added with stirring to a warm saturated ethanolic solution of triphenylphosphine (6.5g) containing potassium hydroxide (0.6g). The yellow precipitate was washed with warm alcohol, water and alcohol again and then dried in vacuo. lg of the hydride was dissolved in benzene (10M1) and hydrogen fluoride was bubbled through the solution for 15 minutes. Dry nitrogen was then passed through he solution to remove all excess HF. The solution was filtered and the precipitate was recrystallised from methylene chloride/ ether. A cleaner reaction is obtained if no solvent is used; the hydride dissolved readily in hydrogen fluoride to give a dark red solution. After a few minutes the HP war blown off by a stream of dry air and the white solid recrystallised as before. Analysis. (C1015 )2PtF2. Calc. C=57.1%, H=3.96%, Pt=25.7% Found =57.3% =3.85% =25.9% =56.4% 3.87% =25.1% Bis(triphenylphosphine)platinumdicarbonyldifluoride. Approximately 5m1 of anhydrous HF was condensed in a bomb containing (i3P)2PtF2 (0.5g) and e pressure of carbon monoxide applied. 'After rocking overnight HF and 0 were removed and the residue was taken up in methylene chloride and recrystall- -175= ised to give a pale yellow solid. Analysis. (C1efi15P)2Pt(00)2P2 Cale. 0=56.0%, H=3.69%, Pt=24.0% =56.3% =4.69% =23.6% =h7.0% =4.00% =24.1% Bis(triphenvlphosphine)platinumdichloride9 was prepared by the action of aqueous potassium tetrachloroplatinate on alcoholic triphenylphosphine with heating, in a molar ratio 1:2 (Grinberg and Razumova 1954). The mixture was sha ken for twenty hours to ensure complete reaction. The white precipitate was washed with hot water, ethanol and finally with ether (decomp. 308°). On heating the compound in methylene chloride at 120° for 15 hours9 under a pressure of carbon monoxide, a bright green solution was obtained. Green crystals were isolated, but not in a pure condition. Attempts at chromatography in methylene chloride showed the presence of at least three components. The green compound decomposed on the column and the eluate contained the starting material. Bis(triphenylphosphine)palladiumdifluoride. Attempts to prepare(f4P)4Pd as described by Malatesta and Angoletta (1957) from palladium nitrate and triphenylphosphine in benzene -176

proved unsuccessful. No reaction occured when freshly prepared palladous oxide was added to a saturated ethanolic solution of the phosphine; and reaction of(i3P)2PdC12 with excess phosphine and hydrazine occured only very slowly. It was found that the best method of preparation was to suspend the phosphine dichloride and triphenylphosphine (molar ratio 1:2) in alcohol and to reduce it with borohydride, and to recrystallise from ethanol. A pale yellow compound was obtained, decomposing at TOO to I051? It was found that the compound was unstable in air, turning orange quite quickly. The yellow solid decolourised a dilute bromine solution in carbon tetrachloride and had peaks in the infra red at 1962, 1886 and 1808 c m This compound,'which we formulate as (i3P)4PdH2„ dissolved in liquid anhydrous HF with considerable vigour, to give a. dark red solution. After 15 minutes the HF was removed by a rapid stream of dry air. The brownish residue was recrystallised from methylene chloride/ether to give a tan coloured solid. Analysis.(C181110P)2PdF2. Cale. C=6405%, H=4.48%, Pd=1509% =61.5%, =4.43%, =15.6%

TrisLtyiphen/lphosphine) irid iumhydrided ifluoride The trihydride, IrH3(,31,)3 , was prepared by adding a saturated ethanolic solution of iridium tribromide to warm saturated ethanolic solution of triphenylphosphine (in a molar ratio of 1:4). Sodium borohydride was slowly added to effect reduction (Malatesta, Angoletta, Araneo and Canziani 1961). The trihydride was dissolved in liquid anhydrous HP to give a red solution; the HP was removed by a stream of dry air after 15 minutes and the residue recrystallised from methylene chloride/ether to give a pale yellow solid. Analysis ir ( ji3p ).3Hp2 Cale. C=60.9%, H=4.86%, 1r=20.5% =58.3% =4.74%, =21.0% =20.0% Bis(triphenylphos_phine)rhenium Bis(triphenylphosphine)rhenium diiodide (Freni and Valenti 1961a) was prepared by dissolving perrhenic acid in alcohol and adding hydriodic acid. After warming for five minutes a saturated ethanolic solution of ig3P was added. On cooling, silver-green crystals were obtained; these were recrystallised from benzene/ethanol. The diiodide was suspended in warm alcohol and sodium borohydride was added slowly with vigorous stirring (Preni and Valenti 1061b). Graant,ily the suspension turned red. The red crystals were purified by recrystallisation from benzeneitalcohol. -178—

Oh treating the solid hydride with liquid anhydr— ous HF a dark green solution was obtained. However on wor1ing up the dark residue, left after removing the HP, only a small quantity of green crystals could be obtained. They were found to be insoluble in benzene and ether but soluble in methylene chloride. On repeating this experiment no more green product could be obtained. If HF was bubbled through a benzene solution of the hydride, the colour changed from red to green but no product could be isolated. -17 9-

CHAPTER--5.

MOLYBDENUM AND TUNGSTEN PHTHALOCYANINES. In a series of papers Linstead et alia (1934) described the first laboratory preparations of phthalocyanines from phthalamide, o-cyanobenzamide and o-phthalonitrile. Since then phthalocyanines of almost every metal have been prepared. These compounds have a great variety of uses in the dyeing and printing industry as dyes and pigments9 because of their intense colour and great resistance to fading. Phthalocyanines containing the metals molybdenum and tungsten have not been studied in any detail so far. It was thought that a study of these would be rewarding; because these two metals exist in a number of oxidation states a variety of phthalocyanines should be possible. This work was undertaken in the hope of prepaiing phthalocyanines of molybdenum in oxidation states 3, ii., 5 and 6 and tungsten in oxidation states 5 and 6. Reactions involving the metal halides aid molten o-phthalonitrile proceeded easily although the time necessary for reaction with molYbdenum dichloride was considerably long er than with any of the others. Reactions proceed slowly with MoOs and Mo02 and from the latter a compound analysing to the formula MoOPc was obtained. Purification of the -180- reaction products usually involved long extractions with benzene and alcohol followed by aniline or pyridine (in which phthalocyanines are slightly soluble). The crystals obtained were then sublimed at about 400D/10-e m.m. The results of reactions involving o-phthslonitrile are summarised in the following table. Time taken for reaction Product. with molten o-phthalonitrile.

MoC13 ___ Solidified after 1 hour. .10.1 01..0 dark bluel could not be purified.

Mo014 ___ Under N2; 20 minutes •••• ••• purple, could not be purified.

MoC15 --_ Under N2; very vigorous .111k .••••• dark blue, could reaction, . 5 minutes. not be purified. Mo012 ___ Very slow reaction; dark blue,could 6 hours. not be purified.

VirBr5 ___ Under V2; vigorous react- ••• Oft •••• dark green, could ion, Br2 evolved. 10 mins. not be purified; 17C16 ___ Under N2; vigorous react- dark green, could ion, C12 evolved. 5 mins. not be purified. Mo Very slow reaction. 2days.___ dark blue, could not be purified. Mo02 Under N2; slow reaction. dark blue, anal- 12 hours. ysing to MoOPc. -181-

Mo03 Slow reaction. 2.5 dark blue, could hours. not be purified.

Refluxing MoC12 and MoC13 with o-phthalonitrile in a variety of high boiling solvents, such as dimethyl- formamide, 1:2 propane diol, glycerol and dimethyl suiph- oxide did not result in the formation of phthalocyanines. However when TToC13 was refluxed in diethanolamine with o-phthalonitrile a dark blue precipitate was formed: on purification a compound analysing to LffoOHPc was obtained. Attempts to prepare molybdenum and tungsten phthalocyanines by refluxing the chlorides and bromides w.;th lithium phthalocyanine in dioxane and quinoline met with no success. LiPe was recovered intact in moetcases. As the analyses for Mo0HPc and MoOPc are identical,within the limits of analytical error, attempts to identify the oxidation state by magnetic measurements were attempted but no worthwhile results were obtained. The infra-red spectrum of Mo0HPe did have a weak band at 1612 cm-1 , this may be due to the OH group. We believe that the phthalocyanines which we have attempted to prepare polymerise, through oxygen bridges. This would account for the difficulty in purification. For this reason further -182- work on these compounds was discontinued. EXPERIMEDTAL.

0-phthalonitrile (supplied by I.C.I. Ltd.) was purified by distillation into methanol. MoC12, MoC13 , WC16 and W131.5 were supplied by the Climax Molybdenum Co. MoC15 was prepared by the chlorination of molybdenum.. metal. MoC14 was pfepared from molybdenum dioxide and chlorine in refluxing hexachlorobutadiene (Austin and Tyree 1960). Mo02 was obtained by reducing Mo05 with hydrogen at 470° for 36 hours. The reactions with molten phthalonitrile were carried out in boiling tubes (if the metal halide was stable) with mechanical stirring at temperatures between 220 and 2700. If the metal halide was unstable reaction was carried out in a three necked flask fitted with an air condenser and guard tube whilst a stream of dry nitrogen was past through. The crude product was finely ground and extracted first with benzene and then with alcohol. The phthalocyanine was then extracted with either aniline or pyridine, and the product om the boiler flask was then sublimed. Sublimatioa at 2000 for 24 hours with the probe at 20° was followed by slow sublimation at 4000 with the probe at.a temperature of 260°. Analysis. C321116N8CMo. calco C=61.6%, H=2.57%, N=18.0%, Mo=15.4%. Found C=61.5%, H=2.86%, N=1803%, Mo=15.3%.

Analysis. C321117N80Mo. Calc. C=61.6%, H=2.57%, N=18.0%, Mo=15.4%.- Found C=61.9%, H=2.78%, N=17.9%, Mo=1505% =61.1%, =2.54%, =18.2%, =15.5%

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Acta Cry8t. (1961). 14, 330 The lattice constants of some metal-fluoroborate hexahydrates. By K. C. Moss, D. R. RUSSELL, and D. W. A. SHARP, Inorganic Chemistry Research. Laboratories, Imperial College, London, S. T. 7, England

(Received 15 November 1960)

Although it has been stated (West, 1935) that the tion below 23 °C. and is hexagonal, isomorphous with Cuoroborates, .111(BF4)2. 6 H2O (M =Mg2+, Mn2+, Fe2+, the corresponding perchlorate (West, 1935). The only Co2+, Zn2+, Cd2+), are isomorphous with the cor- phase that we could crystallize from such solutions is responding perchlorates, none of their lattice constants hexagonal, a =9-90, c =5-53 A, but is not isomorphous appear to have been recorded. The lattice constants and with the perchlorate trihydrate. measured and calculated densities are recorded in Table 1, The hydrates were prepared from solutions of the where they are compared with the values obtained by appropriate carbonates in fluoroboric acid. X-ray powder West for the perchlorates. The Mg, Mn, Fe, Co, Ni, and photographs were taken with a 9-cm. camera using Zn salts are hexagonal and are very similar in size to Cu KN, Co or Cr radiation. the corresponding perchlorates: the cadmium salts have a closely related trigonal structure with a one half of Thanks are due to the Department of Scientific and that shown in Table 1; the true value is doubled for Industrial Research (D.R.R.) and the Hertfordshire comparison with the other salts. West has shown that the County Council (K.C.M.) for maintenance grants. copper salts are not isomorphous with other divalent salts. Theory (Orgel & Dunitz, 1957) would predict a distortion of the octahedra of oxygen atoms about the References Cu2+ ions. Lithium fluoroborate exists in at least two forms. ORGEL, L. E. & DUNJTz, J. D. (1957). Nature, Lond. 179, H2O, stable above 23°, is tetragonal, a = 5.74, 462. c =4.88 A. LiBF4 . 3 H2O crystallizes from aqueous solu- WEST, C. D. (1935). Z. Kristallogr. 91, 480.

Table 1. Lattice constants and densities

111"(C104)2. 6 H2O 17"(13F4)2. 6 H2O

Measured Calculated Measured Calculated At" density density a c density density Mg 15.52 A 5.26 A 1.981 1.99 15.36 A 5.38 A 1.849 1-85 Mn 15.70 5.30 2.102 2.10 15.46 5.44 1.982 1.98 Fe 15.58 5.24 2.147 2.17 15-49 5.33 2.038 2.02 Co 15.52 5.20 2.198 2.22 15.33 5.22 2.081 2.11 Ni 15.46 5.17 2-252 2.25 15.32 5.16 2.136 2.16 Zn 15.52 5.20 2-252 2.26 15.24 5.30 2.120 216 Cd 15.92* 5.30 2.368 2.38 15.96* 5.58 2.292 2.12 * See text.

Printed in Denmark at Fr. Raves kg?. Hofbagtrykkeri, Copenhagen 196z_,

J. Inorg. Nucl. Chem., 1960, Vol. 13, pp. 328 to 329. Pergamon Press Ltd. Printed in Northern Ireland

Tetramethylammonium tertafluoroborate (Received 10 November 1959)

ALTHOUGH it has been shown(1) that substituted ammonium tctrafluoroborates, RQNBF, (R = Et, Pr°, Bun), result from the reaction between tetra-alkylammonium halides and hydroxytrifluoroboric or methoxytrifluoroboric acids, it appears from the literature'2) that the corresponding reactions with tetramethylammonium halides yield tetramethylammonium hydroxytrifluoroborate. We have repeated these latter preparations and believe that the reaction of tetramethylammonium halides with hydroxy- or methoxy-trifluoroboric acids yields tetramethylammonium tetrafluoroborate. The salts were formerly characterised by analysis but the expected analytical figures for the two possible products are so similar (except for fluorine which is very difficult to determine accurately) that we do not consider that they constitute adequate proof of identity. The reported melting points of tetramethylammonium tetrafluoroborate, 4180,(1) and of tetramethylammonium hydroxytrifluoroborate, 4145,12' are almost identical and again do not give a great deal of information. We have found that X-ray powder photographs of tetramethylammonium tetrafluoroborate and the alleged tetramethylammonium hydroxytrifluoroborates are identical and moreover have d spacings the same as those reported by previous workers for Me4NBF3OH (Table 1). This is not definite proof of identity since the OH- ion being of similar size to the F- ion, two compounds, ABF4 and Al3F3OH, should have almost identical powder patterns assuming there is no distortion due to hydrogen bonding. As further proof, however, the infra-red spectra of all of the samples show no

TABLE 1.-d SPACINGS FOR TETRAMETHYLAMMONIUM FLUOROBORATES Values obtained by Me,NBF, "Me4NBF3OH"-1 "Me4NBFaOH"-2 WHEELER et a/.(21

6.31 (5) 5-89 (1) 6.28 (2) 5.82 (6) 4.85 (6) 4.835 (2) 4.78 (2) 4.83 (10) 4.09 (10) 4.07 (10) 4.12 (10) 4.08 (100) 3.12 (7) 3.145 (3) 3.13 (4) 3.13 (20) - - - 2.90 (1) 2-78 (5) 2.750 (2) 2.77 (3) 2.76 (10) - - - 2.61 (1) 2.497 (8) 2.497 (4) 2.493 (4) 2.48 (40) 2.295 (5) 2.310 (3) 2.305 (4) 2.29 (20) 2.136 (4) 2.120 (1) - 2.21 (6) F943 (3) 1.947 (3) - 1.93 (20) 1.885 (1) - - 1.85 (1) 1.650 (3) 1.654 (2) - 1.64 (18) 1.557 (1) - - 1.55 (3) 1.369 (2) - - 1.36 (1) 1.339 (1) - - - 1.301 (1) - - -

Figures in parentheses are estimated intensities based on the strongest line = 10 except for last column where strongest line was put = 100. hydroxyl bands. A sample of potassium hydroxytrifluoroborate clearly showed these bands at 3120, 2860, and 1630 cm-1. The principal band in this compound is split, probably because of hydrogen bonding. (1)C. M. WHEELER, JR. and R. A. SANUSTEDT, J. Amer. Chem. Soc. 77, 2025 (1955). (2)C. M. WHEELER, JR., R. D. BEAULIEU and H. W. BURNS, J. Amer. Chem. Soc. 76, 6323 (1954). Notes 329

Potassium hydroxytrifluoroborate was prepared from potassium bifluoride and boric acid!" The alleged tetramethylammonium hydroxyfluoroborates were prepared as previously describedo) from hydroxytrifluoroboric acid (Sample 1) and from methoxytrilluoroboric acid (Sample 2). (Found for 1: C, 30.8; H, 7-90 %. Found for 2: C, 30.2; H, 7.77 %. Cale. for Me4NBF3OH: C, 30.2; H, 8-34%. Calc. for Me4NBF4: C, 30.4; H, 7.51 %.) Tetramethylammonium tetrafluoroborate was prepared from 40% tetrafluoroboric acid and tetramethylammonium chloride (Found: C, 30.5; H, 7-83%). We thank the Imperial Smelting Corporation Ltd. for a gift of a cylinder of boron trifluoride and Hertfordshire County Council for a maintenance grant (to K. C. M.). Inorganic Chemistry Research Laboratories K. C. Moss Imperial College London S. W.7 D. W. A. SHARP

(3) C. A. WAMSER, J. Amer. Chem. Soc. 70, 1209 (1948).