THE NI~TVERSITY OF MICHIGAN RESEARCH INSTITUTE ANN ARBOR, MICH IGAN Quarterly Report No. 1 July 1 to September 30, 1958 THE CHEMISTRY OF BORON HYDRIDES AND RELATED HYDRIDES R. W. Parry R, C. Taylor C. E. Nordman For the Pro3ject Staff UMRI Project 2793 WRIGHT AIR DEVELOPMENT, CENTER CONTRACT NO. AF 33(616)-5874 PROJECT NCO 30o48, T ASK NO. 70321 WRIGHT-PATTERSON AIR FORCE BASE, OHIO December 1958

TABLE OF CONTENTS Page PROJECT STAFF iii ABSTRACT iv I. THE CHEMISTRY AND STRUCTURE OF THE DIAMMONIATE OF DIBORANE AND ITS DERIVATIVES 1 A. Background 1 B. Further Studies of the Compound Resulting from the Reaction Between B4Hlo- 2NH3 and HF 1 C. Attempts to Synthesize K2[BH2NHCH3] 2 II. STUDIES OF THE AMINE-BORANES 3 A. Further Vapor-Pressure Studies of the Amine-Boranes 3 B. Dipole-Moment Studies of the Amine-Boranes 5 C. Raman Spectral Studies of the Amine-Boranes 7 III. THE REACTION OF PHOSPHORUS TRIFLUORIDE-BORANE AND AMMONIA 11 A. Background 11 B. Experimental 12 IV. THE CHEMISTRY AND STRUCTURE OF THE DIAMMONIATE OF BORANE CARBONYL 13 A. Background 13 B. Experimental 13 V. TETRABORANE-THE CHEMISTRY OF ITS REACTIONS WITH VARIOUS BASES AND THE STRUCTURE OF THE PRODUCTS FORMED 14 A. Background 14 B. The Structure of Ammonia-Triborane, H3NB3H7 (Low-Temperature Form) 15 C. Physical Properties of H3NB3H7 and Me3NB3H7 15 D. The Structure of B4Hlo.2NH3 by X-Ray Crystallography 16 E. The Reactions of B4Hl0 with KBH4, KOH, and KOCH3The Preparation of KB3H8 16 F. The Raman and Infrared Spectra of Tetraborane 20 VIo STUDIES ON BRIDGE COMPOUNDS RELATED TO BORON HYDRIDES 21 A. Further Studies on F3PAlC13, Obtained from A12C16 and PF3. The Fluoride Shift in the PF3 Adducts 21 B. The Reaction of PF3 and Al2Me6 22 C. The Aluminum Hydride-Bis Trimethylamine Complex 22 REFERENCES 24 ii

PROJECT STAFF Ie Chemical Group R. W. Parry G. Kodama J. R. Weaver E. R. Alton J. C. Carter C. W. Heitsch M. Yamauchi IIo Spectroscopic Group R. C. Taylor C. Cluff R. Amster R. Brown IILo X-Ray Group C. E. Nordman C. Peters E. Chang iii

ABSTRACT Additional data obtained on the product from the reaction of B4Ho10.2NH3 and HF suggest that the new compound may be a fluoroborate of the cation H2B(NH3)2+ rather than a fluorides New data and more careful methods of data reduction give values of 18o8, 18.5, and 13.6 kcal/mole for the heats of sublimation of (CH3)NF2BH3, (CH3)2N1BHs3, and (CH3)3NBH3, respectively. Dipole-moment values for the amine-boranes in other solvents are being obtained. A Raman spectral study of the amine-boranes is described. PF3BH3 reacts with ammonia to give (NH2)3PBH3. Failure of H3NBH3 to react with PF3 and ammonia under comparable conditions argues against a reaction mechanism involving replacement of F3P by NH3 in the complex. Preliminary X-ray structural data on (NH2)3PBH3 have been obtained. Infrared and Raman spectra of H3NB3H7 and Me3NB3H7 have been obtained as part of a spectral study. Preliminary X-ray crystallographic data on B4H i0 ~2NM13 are reported. Tetraborane reacts with, KBH4 by a symmetrical cleavage process and with KOH and KOCH3 by nonsymmetrical cleavage processes to give KB3H8. PF3 reacts with Al2Me6 but the products formed have not yet been identified. H3Al(NME3)2 has a dipole moment of 1.3D in benzene solution, Preliminary X-ray crystallographic data are presented. iv

I. THE CHEMISTRY AND STRUCTURE OF THE DIAMMONIATE OF DIBORANE AND ITS DERIVATIVES A. BACKGROUND The type reaction: [H2B(NH3)21[B3H8] E+ HX _2_> [H2B(NH3)2]X + H2 + Et2OB3H7 has been used to prepare the chloride, bromide, and iodide salts of [H2B(NH3)2+]. An extension of the reaction to prepare the fluoride salt, [H2B(NH3)2]F, appeared to be obvious and routine, but unexpected difficulties have arisen in carrying out the reaction and characterizing the product obtained. In previous reports analytical problems were mentioned and it was reported that molecular weight values of the supposed [H2B(NH3)2]F were near 120 or roughly twice the expected value of 66. Attempts to obtain a more definitive characterization of the compound have been continued. B. FURTHER STUDIES OF THE COMPOUND RESULLTING FROM THE REACTION BETWEEN B4H10o2NH3 AND HF In an earlier report the reaction of L4,Hlo*2NH3 and HF in diethyl ether at -78~C was described. The compound formed was tentatively identified as [H2B(NH3)2]F on the basis of the reaction: [H2B(NH3)2]F + LiBH4 Me2O LiF + [H2B(No3)2]BH4 C i E]TH3 > LiFlJ+ [H2B(NH3)2][BH4](soln) extraction - The solid B2H6-2NH3 and Li were identified. by their X-ray powder patterns. The complete characterization of the compound was complicated by analytical problems. The molecular weight value as measured by vapor-pressure depression of an Me20 solution was around 120 as compared to a theoretical value of 66. The best new values for the percentage of ammonia are 25.2 and 26.2 as compared to 51.5 expected for [H2B(NH3)2]F. The observed value for hydrogen is 1.76% as compared to an expected value of 3.04%. For boron the -Value.s 17.3 as compared'u- 164- expectedo The data taken together indicate that the characterization as [H2B(NH3)2]F is probably incorrect. An attractive alternative possibility for the formula of the compound is [H2B(NH3)2]BF4. Theoretical values for this compound would be: Mol wt = 134, %NE3 = 25.3, H- = 1.49, B = 16e1. The results are still inconclu

sive, but all data are reasonably consistent with the fluoroborate interpretation. An authentic fluoroborate sample will be made for comparison. It is interesting to note that the reaction of B4Hjo02NH3 with HF is completely different from the reaction of the other hydrohalogen acids and B4Hlo.2NH3. C. ATTEMPTS TO SYNTHESIZE K2[BH2NHCH3] 1. Earlier arguments arising from the reaction between H2BNH2 and NH3 in the presence of Na led to the suggestion from this laboratory that a compound Na2[H2BNH2] might be formed as an intermediate. It was suggested that a related compound K2[H2BNHCH3] might even be prepared as a stable compound by the reaction between metallic potassium in liquid ammonia and [H2BNHCH3]3$ It had been reported in an earlier report from this laboratory that decolorization of a potassium solution in liquid ammonia took place without H2 evolution when a large excess of [CH3HNBH2]3 was added to the ammonia solution. Attempts to verify the earlier observation have indicated that all decolorization is accompanied by H2 evolution, and the previously described process has not been verified. Furthermore, the products from the B2H6-2NH3-K reaction do not decolorize a solution of K in liquid ammonia without H2 evolution. 2. Experimental Run 1. —A sample of (CH3NHBH2)3 (about 0.94 mM) was dissolved in about 3 ml of liquid NH3. The system was frozen down, and a sample of potassium metal, weighing 22.5 mg (equivalent to 0.565 mM) and sealed in a glass bulb by appropriate vacuum techniques, was added to the system and the bulb was broken. The system was warmed to -780C and held at this temperature for 2 days. At the end of that time the blue color had disappeared and 0.860 mM of H2 gas was measured. The H/K ratio is 3.04. Run 2. The Reaction of Diammoniate of Diborane with Excess Potassium.B2H6-2NE3 was prepared from 0.75 mM of B2H6 by conventional methods. To a liquid ammonia solution of the above reagent, 1.605 mM of K was introduced by conventional techniques [ratio B2H6.2NH3/K = 1/2.14]. The system was warmed to -78~C. Although 0.368 mM H2 was evolved (0.98 equiv. H/B2H6.2NH3), the blue color due to the measured excess of K remained. No direct evidence to support the existence of the compound K2[H2BNH2] has been obtained.

II. STUDIES OF THE AMINE-BORANES A. FURTHER VAPOR-PRESSURE STUDIES OF THE AMINE-BORANES Vapor-pressure measurements on the mono-, di-, and trimethylamine-boranes determined by the Knudsen effusion method have been reported previously.l However, since the heats of sublimation derived from the slope of the least-squares fit of the Clausius-Clapeyron equation appeared somewhat inconsistent, further measurements have been made to check the previous results and to extend the temperature range of the data. Experimental details have been given previously' and will not be repeated. Results were calculated from the relationship.2 P(mm) = 17.14 W where W is the weight of the substance in grams effusing in time t in seconds through an orifice of area A square cm. M is the molecular mass and T is the absolute temperature. The constant, K, is a correction factor necessary to allow for the back reflection of molecules from the cylindrical walls of an orifice having a finite thickness. Values of this constant obtained by Clausing are tabulated by Dushman;3 in the present work, the two orifices used had K values of 0.935 and 0.865, respectively. Since the data previously reported had not been calculated using this factor, the corrected values of the earlier data together with those obtained more recently have all been included in the accompanying Table I. The data have also been fitted to the integrated form of the ClausiLi -Clapeyron equation by least-squares methods to give the vapor-pressure equations below. In the case of trimethylamine-borane, vapor-pressure data determined by conventional manometric techniques are available in the literature.4 Although these data are for higher temperatures and do not overlap the present results, the two sets of data appear consistent on a log P vs. l/T plot. The equation for this compound consequently has been fitted to all data and covers the widest temperature range of the three substances. Methylamine-borane: Loglo P(mm) = - 4114(1 +0.05) + 11.411 (1 +0.065) T range = O0 to 45~C Dimethylamine-borane: Logo P(mm) = 4054(1+0.058) + 12.544 (1+ 0.042) range = O~ to 35~C

Trime thylamine- borane: Loglo P(nm) = - 962(1 o.o04) + 9.894 (1+ 0.014) T range = O~ to 90~C TABLE I Methylamine-Borane Dimethylamine-Borane Trimethylamine-Borane Temp, Pressure, Temp, Pressure, Temp, Pressure, OK mm K mm OK mm 273.6 2.39 x 10-4 273.8 7.78 x 10-3 273.4 9.84 x 10-2 29208 163 x 10-3 285.7 2032 x 10-2 274.2 1.04 x 10o1 300oo,o0 6.73 x 10-3 285.7 2.52 x 10-2 283.4 3.29 x 10o 30000 6.60 x 10-3 291.1 4.12 x 10-2 283.4 3.27 x 10-1 30_5.4 8.42 x 10-3 294.0 5.90 x 10-2 289.4 5.15 x 10-1 306.9 1.03 x 102 295.2 6.64 x 102 289.4 5.16 x 10'1 306.9 8.49 x 10-3 299.2 1.11 x 10-1 296.4 8.96 x 10lo 309.1 1.41 x 10-2 302.5 1.88 x 10-1 297.5 8.19 x 10-1 313.2 1.78 x 10-2 306.0 2.49 x 10-1 298.0 7.94 x 10o314.4 1.96 x 10-2 306.0 2.53 x 10-1 The uncertainties in the constants of the equations represent standard deviations calculated from the residuals between the points and the lines. Derived values for the heats of sublimation of the compounds are listed in Table II. The entropy of vaporization values are the standard entropy of vaporization at one atmosphere pressure and 25~C. TABLE II HEATS OF SUBLIMATION OF SOME AMINE-BORANES Substance AHsubl' AS2 (CH3)NH2~BH3 18.8 ~ 1.0 kcal 38.9 eou. (CH3)2NH:BH3 18.5 ~ 0.7 44.3 (CH3)3N:BH3 13.6 + 0.2 32.2..~~~~~~

Inspection of the heat of sublimation values in Table II shows that the values for the mono- and dimethylamine compounds are approximately the same, while that for the trimethylamine complex is significantly lower. This suggests that a single hydrogen bond involving a hydrogen attached to nitrogen is important in the crystal structure of the first two compounds. Apparently the error in the earlier values of AHsubl, which differed significantly from the above results, can be attributed to the scatter of the points in conjunction with the rather short temperature range over which the measurements were made. B. DIPOLE-MOMENT S7JDIES OF THE AMINE-BORANES 1. Theoretical Studies Because of the low vapor pressures of the amine-boranes, vapor-phase measurement of the dipole moments of these molecules appears to be impossible to effect with the equipment available here. Measurements in solution are always subject to the criticism that local solvent field effects may produce measured values which differ appreciably from the vapor-phase values which are of primary theoretical importance. One of the better approaches to this problem is the measurement of the dipole moment of each molecule in several solvents and a semiempirical, semi-theoretical reduction of the data. A number of investigators have considered the theoretical problem represented by solvent-solute interaction. One can say that, to a zero-order approximation, all the resulting equations are the same, the square of the permanent dipole moment being related approximately linearly to the change in dielectric constant with solution concentration. The precise form of the relation varies from one equation to another depending on the assumption, stated or implied, as to the nature of the interaction between a particular dipolar molecule and the rest of the solution. In the conventional treatment, for example, one assumes that the molar polarization is the sum of various additive components. The physical picture corresponding to this assumption is cloudy, and its experimental weakness is shown by the fact that the dipole moment measured in solution differs from that measured in the vapor phase by an amount which may be roughly related to the dielectric constant of the solvent. Several theoretical treatments of the problem, namely those of Onsager, Boettcher, and Scholte, have arisen from a set of clearly drawn physical assumptions regarding dipole-solvent interaction. It is well known that, although Onsager's equation gives reasonable values for the dipole moment of molecules measured in the pure liquid state, it does not show a similar improvement when applied to dilute solutions of polar molecules in nonpolar solvents. In fact, values for the dipole moment obtained from Onsager's equation usually deviate somewhat further from the vapor-state values than those obtained by the conventional method. It appears that Onsager's equation gives a fairly reliable picture of dipole-dipole interaction in pure liquids, but it offers little help in the interpretation of the interaction between dipoles and nonpolar solvent molecules.

Boettcher modified Onsager's treatment by introducing the assumption that the volume of the molecule should be regarded as smaller than its partial molar volume in the liquid or, in other words, that there is free space in the liquid. Such an assumption simply introduces two additional parameters into the equation and should result in a better fit of the experimental data regardless of the validity of the postulates made. The equation of Scholte introduces two further parameters. In addition to the average molecular radius and average polarizability of the molecule involved in Boettcher's equation, it recognizes the anisotropy of the molecule and includes constants for length and polarizability along the axis of the dipole. In order to evaluate the treatments of Boettcher and Scholte, the data of Pilpel on solutions of acetophenone and benzonitrile in benzene were used. After a rather exhaustive evaluation, it was concluded that the analysis of Boettcher was not particularly appropriate to the immediate data and ultimate problem under consideration. A method of plotting was devised by which Scholte's equation could be tested using the dielectric constant data in the dilute solution region. If data were restricted to solutions containing less than 10 mole percent solute, a modified form of Scholte's equation fit the data well and gave the correct value for the dipole moment of the solute molecule. For concentrations above 10 mole percent, both the acetophenone and benzonitrile showed rather sudden and large deviations. The validity of the modified Scholte equation in reducing data on the amine-boranes is to be investigated further and will be reported in a subsequent papero 2. Experimental MeasUrements (a) Preparation of Methylamine-Borane.-In the course of these investigations, methylamine-borane was prepared for the first time as a stable solid which melted sharply at 54~QC Although mentioned in an earlier report, the preparative procedure has not been previously described; hence it is included herein. Diborane was distilled twice at the temperature of methylcyclohexane slush to remove higher hydrides. About 50 millimoles of B2H6 were added slowly at -78~C to 15 ml of liquid tetrahydrofuran [previously dried over LiA1H4]. Commercial anhydrous methylamine (Matheson) was dried over sodium and about 2 ml of the liquid were distilled into the reaction tube containing the tetrahydrofuran solution of B2H6. A slush bath at -1120C was placed around the reaction tube and the system was allowed to stand. The temperature rose gradually. The solvent and excess amine were removed when the temperature reached 450C; the system was then allowed to warm to room temperature for a few minutes to aid in solvent removal. Finally the reaction tube was surrounded with ice and the last traces of amine and solvent were removed by pumping on the system with the highvacuum pump for 10 hours. Trimethylamine-borane and dimethylamine-borane should sublime under these conditions.

A cold finger was then inserted into the reaction tube, the sample was warmed to room temperature, and MelNE2BH3 was sublimed onto the cold finger as large rectangular crystals. After about half of the product had sublimed, the residue liquified. The sublimation was then discontinued and the residue discarded. The pure product melts sharply at 54~C. It is soluble in ethyl ether and dioxane, and dissolves with difficulty in benzene. Analysis of the sample showed: hydridic H = 6.64%, B = 23o8%; calculated values for CH3NH2BH3 are: hydridic H = 6.68%, B = 24,10%. (b) Calibration of Equipment. —In all previous dipole measurements, only a small portion of the total range of the measuring condenser was used; hence deviations from linearity were unimportant. However, in the extended measurements now in progress, a wider range of the condenser scale is involved and calibration was desirable. The condenser scale has been calibrated using conventional methods. The precision refractometer was also calibrated using the glass reference piece supplied with the instrument. (c) Dipole-Moment Measurements. —The moments of mono-, di-, and trimethylamine-boranes have been measured in benzene, ethyl ether, and dioxane. Monomethylamine was prepared as described above and transfers were made in the dry box. Dimethylamine-borane from Callery Chemical Company was sublimed several times and sublimed into the weighing cell before use. Trimethylamine-borane, prepared from prepurified B2H6 and Me3N, was sublimed into the reaction vessel. The three solvents, reagent grade benzene, "anhydrous" ether, and dioxane, were refluxed with LiAlH4 for from 6 to 12 hours, then put on the vacuum line and distilled through a frit into a glass storage vessel. Solvent was transferred to the weighing cells by means of a hypodermic syringe, Data from the runs have not been completely analyzed, but significant differences between measurements in different solvents have been observed anid are being considered in relation to the theoretical problem outlined in the earlier section. C. RAMAN SPECTRAL STUDIES OF THE AMINE-BORANES Reference has been made in previous reports to the conclusion, based on work with the amine-boranes, that the frequency assigned to the boron-nitrogen dative bond in the literature is too high.'The evidence supporting this conclusion is summarized as follows.5 The bond in question is one of the more stable of the dative bonds to boron and in complexes of boron trifluoride with ammonia6 and nitrogen heterocycles,7 frequencies lying in the range between 980 and 1120 cm-l have been assigned. In the complex of boron trimethyl with ammonia, a value of 1105 cm-l has been chosen8

and in the complex of borane (BH3) with nitrogen trimethyl, [N(CH3)3], a value of 1250 cm-i has been proposed.9 By way of comparison, frequencies assigned to the covalent carbon-carbon bond in halogenated ethane derivatives, where the masses involved are comparable, fall in the range between 800 and 1000 cm-1 with very few exceptions. It is, of course, not always a proper comparison merely to look at the magnitude of frequencies in complete disregard of other factors. Nevertheless, chemical data seem to indicate a lower range for this stretching frequency of the boron-nitrogen dative bond. In the case of ammonia-borane, it has been possible to identify the B-N frequency unambiguously as a well-polarized band of medium intensity at 785 cm1l in liquid ammonia solution. The assignment can be made with a high degree of confidence for this compound in contrast to those reported in the literature because of the relative simplicity of the spectrum below 1200 cm-l. Confirmation of the assignment is amply supplied by the results obtained from the several isotopic derivatives studied. To a first approximation, the NH3 and BH3 groups may be considered rigid and treated as point masses. The boron-nitrogen frequency of the isotopic species should then be calculable from the reduced masses of the psuedodiatomic molecules. In Table III the observed frequency ratios have been compared with the appropriate mass ratios and it is seen that the agreement is excellent, the largest deviation being about 2%. The simple boron-nitrogen stretching force constant calculated from the hydrogen compound datum to this approximation is 2.8 x 105 dynes/cm. TABLE III COMPARISON OF THE BORON-NITROGEN STRETCHING FREQUENCY IN ISOTOPIC DERIVATIVES OF AMMONIA-BORANE IN LIQUID AMMONIA SOLUTION Compound vB-), */v H3N:BH3 785 __ H3N:BD3 737 0.94 0.95 D3N:BH3 754 0. 96 0.97 D3N:BD3 708 090 0. 91 D3N:B 1~D3 713 0.91 0.93 In the case of the methylamine derivatives, the band in question appeared with approximately the same intensity and depolarization characteristics but with decreasing frequency as the number groups on the nitrogen increased (Table IV). mal compound shifted to 676 cm-1 in the isotopic. The band near 1250 cm' as8~~~~~~~~~~~~~~~t

TABLE IV THE BORON-NITROGEN STRETCHING FREQUENCY OF THE METHYLAMINE-BORANES (DIMETHYL ETHER SOLUTION) Amine Frequency, cmAmmonia 755 Methylamine 726 Dimethylamine 696 Trimethylamine 667 signed to the B-N frequency by Rice et al.9 however, was found to shift only 2 cm-1, which is approximately the precision of the measurements. The literature assignment thus appears doubtful, particularly since the 1250-cml band is weak in the Raman effect. The spectral range within which the boron-nitrogen frequency is assigned in the present investigation is in good agreement with the frequencies available for other dative bonds to boron in complexes of borane with ligands such as carbon monoxide, phosphorus trifluoride, and dimethyl ether. The boron-hydrogen stretching region in the spectra of the amine-boranes studied is markedly similar. In each case there are two strong bands appearing at approximately 2270 and 2370 cm-1 plus a weaker, rather diffuse band at about 2330 cm-1. The actual figures are listed in Table V. TABLE V FREQUENCIES IN THE HYDROGEN STRETCHING REGION IN THE RAMAN SPECTRA OF THE METHYLAMINE-BORANES H3NBH3 MeH2NBH3 Me2HNBH3 Me3NBH3 Me3NBH3 (Ref. 9) 2273 s,1v 2266 s,p 2260 s,p 2268 s,p 2266 (10), p 2326 w,d 2328 w 2332 w 2330 w 2330 (3) dif. 2373 s,p 2367 s 2368 s 2369 s,p 2372 (10), d Symbols: s = strong, w = weak, dif. = diffuse, p = polarized, d = depolarized In ammonia-borane, the polarization measurements have shown clearly that two intense bands are both strongly polarized while the intermediate band appears to be depolarizedo Accordingly, the 2273-cm-L band has been assigned as the symmetric B-H stretching mode, the 2326-cm'r band as the asymmetric mode, and the 2373-cm1 band as the overtone of the intense band at 1175 cm1l in Fermi reso

nance with the symmetric B-H mode. In the amine-boranes, the polarization characteristics of these bands have been more difficult to determine since the degree of polarization is less. Rice and his co-workers,9 in their work on trimethylamine-borane, report the 2372-cm-1 band to be depolarized and assign the 2266- and 2572-cm-1 bands to the symmetric and asymmetric B-H bands, respectively, leaving the intermediate band at 2330 cm.1 for the overtone. This assignment does not fit with the ammonia-borane data. The polarization properties of the trimethylamine-borane Raman bands in liquid ammonia solution have now been examined more carefully and the results indicate that the highest band at 2372 is indeed polarized and should therefore be assigned to a totally symmetric vibration, namely to the overtone. This assignment is not entirely unequivocal because of the large uncertainty associated with the measurements but it is in agreement with the ammonia borane results. Experimental difficulties have prevented carrying out the projected work on the infrared spectrum of these compounds at low temperatures. 10

III. THE REACTION OF PHOSPHORUS TRIFLUORIDE-BORANE AND AMMONIA A.o BACKGROUND In the original studies on the preparation of H3BPF3,10 it was found that trimethylamine replaced the weak base PF3 if H3BPF3 and NMe3 were mixed in a one-to-one mole ratio. Addition of excess trimethylamine then resulted in an unspecified reaction between PF3 and NMe3. The process can be represented as: NMe3( excess) F3PBH3 + NMe3 - Me3NBH3 + PF3 New products As noted in the last report from this laboratory, ammonia reacts with F3PBH3 in a different manner. The reaction is represented by the equation: 6NH3 + F3PBH3 11Et2O- 5 3NH4F + (H2N)3 PBH3 -111 to -350C The new ether soluble product, triamido-phosphorus borane was characterized by analysis and molecular weight values. In a formalistic sense the reaction of H3N and F3PBH3 was compared to the reaction of H3N and F3PO and it was of some academic interest to decide whether the reaction was a direct one involving attack of H3N on F3PBH3, or whether the reaction proceeded in a more conventional fashion by the process: H3N + F3PBH3 I —- H3NBH3 + PF3 PF3 + 6NH3 - 4 5113NH4F + P(NH2)3 H3NBH3 + P(NEH2)3 --— }!(,I2N).'BH3 + H3N The first equation in the above sequence would be directly analogous to the reaction between trimethylamine and F3PBH3. A direct test of the above postulate could be made by mixing H3NBH3, PF3, and NH3. If reaction proceeded as described above, (H2N)3PBH3 and NH4F should appear as products. Experimentally it was found that no (H2N)3PBH3 could be detected in the products and some of the original H3RBH3 remained unchanged. These data support the hypothesis that the reaction proceeds through direct attack of RTN3 on F3PBH3 without initial rupture of the P-B bond. In this sense the analogy with F3PO is supported. 11

B EXPERIMENTAL lo The Reaction Between H3NBH3, PF3, and NH3 Ammonia-borane (0.363 millimole) was dissolved in about 2 ml of diethyl ether. An approximately equimolar quantity of PF3 (0.368 mM) was condensed into the reactor and the system was allowed to warm up to -78~C. No external signs of reaction appeared. The solution remained clear and transparent and no gas evolution was detected. The system was frozen with liquid nitrogen and a large excess of ammonia (4.31 mM) was condensed into the reactor. The temperature was then allowed to rise slowly to -78~C and this temperature was maintained for 10 hours. Finally the temperature was raised to -350C and kept at this temperature for several hours. The reaction mixture was filtered through the vacuum line filtration apparatus and the volatile components were removed from the filtrate and fractionated. The recovered ammonia amounted to 3.07 mM; thus 1.24 mM appeared to have been used in the reaction. The ratio of NH3 consumed to PF3 used is 1.24/0.368 or 3.37. The value is below the ratio of the 6 to 1 predicted by the equation for the complete ammonolysis of PF3. (See second equation in above sequence.) The solid product remaining after removal of the volatile components from the filtrate was examined by X-ray powder methods. NH3BH3 was detected in the solid but no lines for (NH2)3PBH3 were found. 2. X-ray Crystallographic Data on (NH2)3PBH3 A structural study using single-crystal X-ray methods is in progress on this compound. Single crystals have been prepared, mounted, and photographed. The crystals are monoclinic with the approximate cell parameters: a = 9.38A, b = 9.47A, c = 6.21A, P = 100.2~. The unit cell contains four molecules. The space group P21/C is unambiguous. Three-dimensional data have been collected and visually evaluated. The locations of the four phosphorus atoms in the unit cell have been found from hko and okl Patterson projections. The determination of the complete structure is in progress. 12

IV. THE CHEMISTRY AND STRUCTURE OF THE DIAMMONIATE OF BORANE CARBO30NYL A. BACKGROUND One of the few relatively simple boron hydride addition compounds whose structure remains mysterious is H3BCO-2NH3 or the diammoniate of borane-carbonyl. Reasoning from analogy, members of this laboratory suggested a structure analogous to ammonium carbamateo + H2N\ /Z0 H2N\ NH4 C NH4+] I C Ammonium Carbamate Diammoniate of Borane Carbonyl Experiments designed to further the chemistry of the diammoniate of borane carbonyl and to test the above structural proposal are in progress. B. EXPERIMENTAL H3BCO2NHF13 was prepared by adding an excess of ammonia to 2.98 mM of H3BCO, warming to -126~C for 30 minutes, warming to -78~C, and allowing the system to stand at this temperature for 18 hours. The ammonia was then removed at -78~C and finally at 0~C until the white solid remaining evolved no more NH3. This solid gave off a noncondensable gas, presumably H2, fairly rapidly at room temperature (approximately 0.1 mM/hr). The ammonia was then returned to the reactor and a liquid of low vapor pressure remained in the U-trap where the ammonia had been stored. The nonvolatile liquid was allowed to remain in contact with excess ammonia at -780C for three days. At the end of this time the ammonia was removed again from the U-trap and a white solid remained. It is assumed that the solid is H3BC02NH3 and that the liquid of low vapor pressure represents a reaction intermediate. To the sample of solid NH3 and H3BCO.2NH3 in the original reactor at -1960C a large excess of sodium was added, As the temperature rose slowly, the solid ammonia melted and sodium went into solution. From this system 1.07 mM of hydrogen were evolved quite rapidly. On warming to -450C for 40 hours 1.04 mM of additional hydrogen were evolved. Since some of the original H3BCO was removed with the ammonia into the U-trap, a precise analysis of the stoichiometry of the sodium reaction cannot be made from the data available, but the qualitative observations are in agreement with the observations of Burg and Schlesinger.4 The system is being explored further. Additional simple and exploratory chemical studies appear mandatory. 13

V. TETRABORANE-THE CHEMISTRY OF ITS REACTIONS WITH VARIOUS BASES AND THE STRUCTURE OF THE PRODUCTS FORMED A. BACKGROUND The chemistry of diborane, as interpreted in this laboratory, usually involves the bridge bonds as centers of reactivity. So-called symmetrical and nonsymmetrical cleavage processes have been recognized on the basis of the products obtained. Thus diborane will undergo two reactions as indicated below: HH H H /BXXB\/B + 2NR3- 2H-BNR3 H 7H H H Symmetrical Cleavage H\ H H H\ /NH3 H /H B'.B\ + 2NH3 LH H H H H\NIH Nonsymmetrical Cleavage Tetraborane likewise has a double bridge bond and the same type of cleavage patterns has been recognized. (See earlier reports from this laboratory.) H H H\ /Ht-H /H H R3N-B-H\ /H /B/ | /B\ + 2NR3 -— )H-B NR3 + I B\ H / H-B-H H H H-B-H H H H Symmetrical Cleavage H\ H-B-H H H NH3\ H /H/B_ | /B + 2NH3 -/B\ /B H / H-B-H \H H NH3 -B-H H Nonsymmetrical Cleavage Only chemical evidence was used in establishing the nonsymmetrical cleavage of the double bridge bond in tetraborane by ammonia. The symmetrical ammonia cleavage product could be made by indirect methods, e.g.: 1 )

Et2O B3H7 + NH3 - ) H3NB3H7 + Et20 Its structure has been worked out by X-ray methods and earlier reports have given an account of this work as it progressed. B. THE STRUCTURE OF AMMONIA-ITRIBORANE, H3NB3H7 (LOW-TEMPERATURE FORM) The refinement of this structure by least-squares methods has been further improved by the use of individual isotropic thermal parameters for all atoms, except the hydrogens of the NH3 group. The final refinement yielded a value of 0.107 for the residuals. The structure of the B3H7 group resembles a fragment of the B4H1o molecule and is in agreement with the symmetrical cleavage hypothesis outlined above- however, rather significant (and not unexpected) distortions of the parent tetraborane structure are observed due to replacement of BH3 by NH3. As a result of these distortions, an alternative structural interpretation can be suggested, namely, that B3H7'NH3 can be considered as a bridge substituted diborane. NH3 \ H\, I/H /B\ /B\ H H H The above model suggests the interesting possibility of synthesizing R3NB3H7 from the reaction: R3NBH3 + B2H6 - R3NB3H7 + H2. Such a possibility is under examination. A paper on the structure of H3NB3H7 has been completed and will be submitted for publication. C. PHYSICAL PROPERTIES OF H3NB3H7 AND Me3NB3H7 1. Infrared Spectra of H3NB3H7 and Me3NB3H7 Preliminary infrared spectra of H3NB3H7 and Me3NB3H7 have been obtained. Due to the low solubility of these compounds in the customary solvents, the spectra have not been entirely satisfactory for detailed analysis. The results available indicate that, in the ammonia complex the bridge-stretching frequencies have been shifted from their abnormally high value in tetraborane down into their customary region between 1500 to 2000 cm-l1 However, in the spectra of the trimethylamine complex, a much smaller effect on these bands is apparent and their position is still unusually high. It is tempting to postulate that this evidence implies a stronger interaction between ammonia and triioorane than between trimethylamine and triborane, a conclusion which is in agreement with prejudices arising from coordination theory.ll It is obvious that a more detailed and careful analysis must be made for reliable conclusions. 15

2. Vapor Pressure of H3NB3H7 Vapor-pressure data reported previously for ammonia-triborane were not corrected for the orifice thickness (cf. Section II-A of this report). The revised values are tabulated below together with the vapor-pressure equation obtained by a least-squares fit of the integrated Clausius-Clapeyron equation, Temperature, ~K Pressure, mm 303.6 7.76 x 10-4 309.6 1.31 x 10-3 314.2 1.99 x 10-3 319.4 3.15 x 10-3 327.0 5.81 x 10-3 Logic P(mm) = 3739 (1 +0.0075) + 9.200 (1 ~ 0.0096) T (range: 30 to 550C) The derived heat of sublimation is AHsubl' = 17.11 + 0.13 kcal/mole and the standard entropy of vaporization at 250C and one atmois AS~ = 28.9 eou. It is of interest to note that the heat of sublimation is approximately equal to that of the amine-boranes having at least one hydrogen on the nitrogen atom. D. THE STRUCTURE OF B4Hlo'2NH3 BY X-RAY CRYSTALLOGRAPHY The currently accepted model for B4Hj0o2NH3 is based entirely on chemical evidence. Although such evidence is considered to be very strong, independent physical evidence on the structure is desirable. A single crystal of B4Hlo-2NH3 has been prepared, mounted, and photographed for single-crystal X-ray studies. The compound is orthorhombic with the approximate cell parameters a = 9.26A, b = 9.42A, c = 8.22A. The unit cell contains four molecules giving a calculated density of 0.809 g/cm3. The space group is either C9 -Pbn2,m. Com22 o h-Pbnm. Complete three-dimensional data have been recorded using an integrating Weissenbergcamera. The more than 400 observed reflections are now being measured using a photoelectric densitometer. E THE REACTIONS OF B4HLo WITH KBH4, KOH, AND KOCH3 - THE PREPARATION OF KB3H8 1. Background Up to the present time, NH3 is the only reagent which is known to give nonsymmetrical cleavage of the double bridge bonds in tetraborane. In diborane, G. Schaeffer and his co-workers12 have obtained evidence for the reaction of LiNH2 with B2H6 to give LiBH4 and H2BNH2, a process which can be interpreted in terms of a nonsymmetrical cleavage of the bridge bonds of B2H6. 16

H \H H \/B\, + LiNH2 - H2BNH2 + LiBH4 Similarly, Schlesinger, Brown, and their collaboratorsl3 reported that B2H6 reacts with NaOCH3 in accordance with the equation: 3NaOCH3 + 2B2H6 - -, 3NaBH4 + B(OCH3)3 In the original paper the reaction was interpreted in terms of a symmetrical cleavage: 2NaOCH3 + B2H6 ) 2NaCH30BH3 followed by further reaction with diborane: 2NaCH30BH3 + B2H6 -- 2NaBH4 + 2BH20CH3 and a disproportionation: 6BH20CH3 - 2B2H6 + 2B(OCH3)3 It is of interest to note, however, that one can interpret the products obtained as evidence for a nonsymmetrical cleavage process comparable to that observed for LiNH2. \ X\ / BI B + NaOCH3 >- H2B(OCH3) + NaBH4 H /7H H B2H6 + B(OCH3)3 14 Similarly, a report by Winternitz indicated that the reaction of B2H6 with a cold concentrated solution of KOH gave significant yields of KBH4, a fact which can also be interpreted in terms of a nonsymmetrical cleavage process. H\ H\ /H /B /B + KOH - H2BOH + KBH4 H /H \ All the foregoing interpretations relative to the nonsymmetrical cleavage of diborane are equivocal in that one cannot rule out a symmetrical cleavage followed by a subsequent reaction which gives the observed products. On the other hand, the cleavage pattern of tetraborane is subject to more certain interpretation since "symmetrical cleavage" gives H3B and 33H7 fragments rather than two H3B fragments as in diborane. Accordingly, the reaction of tetraborane with KOCH3, KOH, and KBH4 has been studied to determine the cleavage pattern expected with these reagents. 17

Potassium hydroxide, potassium methoxide, and potassium borohydride were each allowed to react with tetraborane in diethyl ether slurry at O0C. Edwards and his collaborators at Callery Chemical Company have established that NaBH4 reacts with tetraborane by a symmetrical cleavage process to give NaB3H8. The equation iso NaBH4 + B B % B2H6, + NaB3H8 H,' H-i-H H H It was reported by Edwards and Houghl5 that sodium hydride reacts in a similar fashion: H I H /HB-H\ /H 1 NaH + B /B/ NaB3H8 / \ H //H-B-H \H The KBH4-B4Ho10 reaction was employed in this study as a means of preparing an authentic sample of KB3H8. The reaction gave a solid product with a hydrogenboron ratio of 3 to 1 as determined by acid hydrolysis. The theoretical hydrolysis equation would be: KB3H8 + H+ + 9H20 -- 3 )B(OH)3 + 9H2 + K+ Thus the ratio expected for KB3H8 would be H2/B = 3/1 as observed. Complete analytical data are not yet available, and the absolute analytical values for B and H are about 20% low for KB3H8 suggesting the possibility of solvent in the sample. The product is slightly soluble in ethyl ether, but very soluble in tetrahydrofuran. The solid gives a definite and characteristic X-ray powder pattern. Although characterization is as yet incomplete, the evidence supporting the formula [KB3H8 + some solvent] is good. On the basis of still incomplete evidence, the reactions of KOH and KOCH3 and with tetraborane can be represented by the nonsymmetrical cleavage equations: H H H\iH —H\/H H\ /OCH3 s -B-H\ / KOCH + B \ -B _+ K /\ H/ H-B-H H H H, H H H\,~H-B-H\ /H H OC3 -B-\ 18

The processes are always accompanied by side or subsequent reactions. Appreciable quantities of H2 gas were evolved, probably as a result of reactions involving H2BOH and H2BOCH3. Neither H2BOH nor H2BOCH3 were isolated from the process. The latter compound has been described as a solid which decomposes slowly to B2H6 and B(OCH3)3. The solid identified as KB3H8 gave a powder pattern which was identical to the solid tentatively identified as KB3H8 resulting from the KBH4-B4H1o reaction. 2. Experimental (a) KBH4-B4H1O.-Commercial KBH4 (Metal Iydride) was allowed to react at O0C for 3 hours with a slight excess of B4Hlo in diethyl ether. The KBH4 was slurried in the ether; it is not soluble. The system was filtered in the vacuum line filtration system and the diethyl ether filtrate was separated. It left virtually no solid when the ether was distilled out. The solid on the filtering disc was then leached with tetrahydrofuran. From this filtrate a white solid, presumably KB3H8, was crystallized out as the THF was removed. Its characterization is as yet incomplete. (b) Reaction KOH-B4Hlo — Several pellets of commercial reagent grade KOH were placed in a reaction tube; the tube was fastened to the vacuum system, evacuated, and about 5 ml of diethyl ether were condensed into the system. A sample of B4H1o ( 5 mM) was condensed into the system and it was stirred constantly as the temperature gradually rose to O0C. The solid phase was separated by filtration; then it was leached with tetrahydrofuran. Removal of the THF by distillation left a solid, the X-ray powder pattern of which is identical to that of the product, presumably KB3H8, of part (a). (c) Reaction KOCH3-B4H1o.-Potassium metal sealed in a bulb by standard techniques (3157 mM) was broken into a reaction tube under a stream of dry nitrogen; then the tube was attached to the vacuum line. A sample of CH30H (slightly more than 3.57 mM) was condensed into the tube and allow to react with the potassium by raising the temperature slowly. The H2 evolved and the excess of CH30H was pumped from the system; about 5 ml of ethyl ether and 5~42 mM of B4H1o were condensed in the tube and the system was allowed to warm to -78~C and maintained at this temperature for 30 minutes. The KOCH3 slowly dissolved to give a clear solution; then a white precipitate began to form. The system was then allowed to warm up to -35~C; it was held at this temperature for 30 minutes; then filtered in the vacuum line filtration assembly. As the filtrate was warmed to room temperature a white precipitate formed, indicating incomplete reaction prior to filtration. The solid product was leached from the filter with tetrahydrofuran. When the solvent was removed, the solid gave an X-ray powder pattern which is identical to that of the earlier product identified tentatively as KB3H8. 19

F. THE RAMAN AND INFRARED SPECTRA OF TETRABORANE A study of the vibrational spectrum of tetraborane has been started. The infrared spectrum of this substance has been published but no interpretation has been given.6 The Raman spectrum has not been obtained. The spectrum of this boron hydride is unique since the stretching frequencies associated with the bridge hydrogen atoms do not occur in the region in which they are found in the other boron hydrides. The conventional bridge frequencies are found in the region 1500-2000 cm1l, yet this region is clear in the spectrum of B4H1o. The frequencies tentatively attributed to the bridge hydrogens in the preliminary spectrum appear in the region around 2150 cm1. The preliminary spectra gave evidence of B-2 and B-5 impurities in the tetraborane and considerable time was devoted to obtaining spectroscopically pure material. 20

VI. STUDIES ON BRIDGE COMPOUNDS RELATED TO BORON HYDRIDES A. FURTHER STUDIES ON F3PAlC13, OBTAINED FROM A12C16 AND PF3. THE FLUORIDE SHIFT IN THE PF3 ADDUCTS 1o Background The compound Cl3AlPF3, formerly assumed to be too unstable for isolation, was prepared in this laboratory and has been described in earlier reports. A characteristic reaction of this compound is an internal shift of F- and C1- to give A1F3 and PC13 as ultimate products. This process looked particularly interesting since it represents another example of an internal shift of ligands in these Lewis acid-base complexes. The process can be formalistically compared to the shift of OCH3- and H- groups found in the previously described compound H3BN(CH3)20CH3, etc.17 It has been assumed that F3PAlCl3 is formed initially by interaction of the lone pair electrons of PF3 with the aluminum of AlC13. It is then assumed that a fluorine of PF3 is transferred to the Al of AlC13 (temporary coordination number of 5 for Al) and that this is followed by a shift of C1- to the P. It appears that it should be possible to isolate intermediates of the form PF2C1, PFC12, etc., by careful control of the conditions of the rearrangement. Such a study has been conducted and is described in the following subsection. 2, Experimental. The System PF3-A12C16 Attempts to isolate mixed phosphorus halides from the decomposition of F3PAlC13 have been unsuccessful. A sample consisting of 19.95 mM PF3 was condensed onto 0.58 mM of A12C16. After standing at room temperature under 8 atm pressure for three hours, there was no evidence of a liquid phase and thus of PC13 formation, but the complex C13AlPF3 had formed. From the reaction vessel 18.84 mM of PF3 was recovered, leaving 1.11 mM of PF3 consumed in the reaction or a ratio of 0.96 mole PF3 per or; alaa weight AlCl3,; The complex was never warmed above -112~C after removal of excess PF3 except as specifically noted below. About 35 mM PF3 was condensed on the sample and left at -1120C for three hours. During this period a cloudiness (presumably AlF3) began to develop in the PF3 solvent. Another 30 mM PF3 were added and after two more hours at -112~C there was a large amount of precipitate. PF3 was removed at -150~C and the molecular weight of the last portion of vapor to come off was measured. A value of 86 (theoretical for PF3 = 88) indicated the absence of PF2C1 or PFC12 in that fraction. No volatile fraction was obtained at -112~C; however, at -63~C a volatile fraction identified as 21

0.75 mM PC13 was separated. (V.P. PC13 = 1o mm at 0PC, PFC12 = 450 mm at O;~C sample = 41 mm at O~C.) The tube was then warmed to O0C and an additional 0.34 mM of PC13 was recovered representing about 98% of the original phosphorus consumed in compound formation. The fact that no mixed halides were ever isolated indicates that halide shift occurs more easily than dissociation of the complex. The foregoing observations indicate that in excess liquid PF3, the halogen shift occurs even at -1120C, but the reaction is sufficiently slow so that the molecular weight of the complex as determined in liquid PF3 is valid if only fresh solutions are considered. (See earlier reports.) The reaction time for preparation of the complex seems to be about 3 hours if very pure PF3 is used; however, PF3 with traces of HC1 (colors solution red in liquid state) seems to slow down halogen exchange in the complex and the system may be held for 4 hours at room temperature without appreciable exchange occurring. 35 The Decomposition of Pt(PF3)2C12 Earlier observations in this laboratory had suggested that the product obtained from the thermal decomposition of Pt(PF3)2C12 might be an authentic mixed halide PF2C1 or PFC12. Preliminary observations on this system have been inconclusive due to the presence of NH4C1 in the reaction system. B. THE REACTION OF PF3 AND Al2Me6 Preliminary studies on the system PF3-Al2Me6 have been carried out. Experimental techniques have been developed using grease-free systems. The reaction is more difficult experimentally than the PF3-A12C16 reaction. It is rather clearly established that a reaction occurs between PF3 and Al2Me6 but the stoichiometry has been poor and nonreproducible; products have not yet been identified. Some evidence for the formation of PMe3 has been obtained. C. THE ALUMINUM HYDRIDE-BIS TRIMETHYLAMINE COMPLEX 1. Dipole-Moment Measurements The dipole moment of the compound H3Al(NMe3)2 has been determined in benzene solution using the standard heterodyne beat apparatus described in earlier reports. A value of 1.3 Debyes has been obtained for the compound in benzene solution. 2. Single-Crystal X-Ray Studies A single crystal of the compound has been prepared, mounted, and photographed. 22

The data are now being reduced. Preliminary information as noted below is now available. (a) Lattice type is orthorhombic; a = 8.85A, b = 10.10A, 0 c = 12.93A. (b) Number of molecules/unit cell = 4. (c) Space group Ccmb or Cc2a. 23

REFERENCES 1. Contract No. AF 33(616)-3343, Project No. 6-(2-3055), Task No. 70321; quarterly reports for periods January 1 to March 31, 1958, and April 1 to June 30, 1958. 2, Dushman, S., Scientific Foundations of Vacuum Technique (John Wiley and Sons, Inc., New York, 1949), po 22. 3. Ibid., p. 99. 4. Burg, Ao B., and Schlesinger, H. I., J. Am. Chem. Soc., 59, 780 (1937). 5. Taylor, R. C., and Cluff, C. L., Nature, 182, 390 (1958). 6. Goubeau, J., and Mitschelen, H., Z. phys. Chem., 14, 61 (1958). 7. Luther, H., Mootz, D., and Radwitz, R., J. Pract. Chem., 277, 242 (1958). 8. Goubeau, Jo, and Becher, H. J., Z. anorg. allgem. Chem., 268, 1 (1952). 9. Rice, B,, Galliano, R. J., and Lehmann, W. J., J. Phys. Chem., 61, 1222 (1957). 10. Parry, R. W., and Bissot, T. C,, J. Am. Chem. Soc., 78, 1524 (1956). 11. Bailar, J. C., Chemistry of the Coordination Compounds (Reinhold Publishing Corp., New York, 1956), p. 128. 12. Schaeffer, G. W., and Basile, L. J., J. Am. Chem. Soc., 77, 331 (1955). 13. Schlesinger, H. I., Brown, H. C., et al., J. Am. Chem. Soc., 75, 188, 203 (1953). 14. See Schechter, W. H., Jackson, C. B., and Adams, R. M. (eds.), Boron Hydrides and Related Compounds, Callery Chemical Co. Report, second edition (May, 1954), p. 51. 15. Edwards, L. J., Hough, W. V., and Ford, M. D., in Proceedings of the Sixteenth International Congress for Pure and Applied Chemistry, Paris, 1957 (International Union, Paris, 1958), p. 474. 16. McCarty, L. V., Smith, G. C., and McDonald, R. S., Anal. Chem., 26, 1027 (1954). 17. See WADC Tech. Report No. 56-318. 24