THE UNIVERSITY OF MICHIGAN COLLEGE OF LITERATURE, SCIENCE, AND THE ARTS Department of Chemistry Technical Report DIFLUOROPHOSPHINE LIGANDS, THEIR PREPARATION, PROPERTIES, AND CHEMISTRY Ralph William Rudolph ORA Project 06785 supported by: PETROLEUM RESEARCH FUND OF THE AMERICAN CHEMICAL SOCIETY GRANT NO. PRF 2089-A3 administrated through: OFFICE OF RESEARCH ADMINISTRATION ANN ARBOR May 1966

This report was also a dissertation submitted in partial fulfillment of the requirements for the degree of Doctor of Philosophy in the University of Michigan, 1966.

To My Wife, Margaret, and Our Parents ii

Acknowledgements The author wishes to express special thanks to the graduate and postdoctoral students in the research groups of Messrs. Parry and Taylor. Their advice and assistance helped in many aspects of this research. Worthy of particular thanks are C. F. Farran, for his help with portions of the vibrational studies, and Frank Parker, for the determination of several nmr spectra. I also wish to express my gratitude to Professor R. C. Taylor for many stimulating and constructive discussions. My deepest expression of gratitude is reserved for Professor R. W. Parry. One cannot imagine working for a more patient and understanding person. Without his apt suggestions and expert guidance this research would not have come to fruition. It is indeed a reward to work for such an inspired and dedicated teacher. I can only emulate the high professional standards which he exemplifies. The generous financial support of Union Carbide, Tecumseh Products Company, The Lotta B. Backus Fellowship Fund, The National Science Foundation, and The Horace H. Rackham School of Graduate Studies is gratefully acknowledged. iii

Table of Contents HISTORICAL BACKGROUND......................... 1 STATEMENT OF THE PROBLEM.....o................... 6 RESULTS AND DISCUSSION Io Synthetic Methods for the Preparation of Difluorophosphine Ligands................... 7 A. Fluorination of the Corresponding Halophosphine.............................. 8 B. Deamination of Dimethylaminodifluorophosphine................................ 9 Co Aminolysis............................... 10 D. Metathesis............................... 10 E. Coupling................................ 11 IIo Characterization of New Difluorophosphines A. Characterization of Difluoroiodophosphine, PF2I.............................. 12 Bo Characterization of Cyanodifluorophosphine, PF2CN.. o........................... 21 C. Characterization of — Oxo-bisdifluorophosphine, F2POPF2...................... 29 D. Characterization of Tetrafluorodiphosphine, P2F4......................... o o 39 Eo Characterization of Difluorophosphine, PHF2 o o o o o.............. o........ o o o...... 51 III. Some Chemistry of the New Difluorophosphines A o General.......................o... o.. 65 B. Solvolysis of Difluorophosphine 1) Hydrolysis.......................... 66 2) Aminolysis..o.ooo........... o.............. 68 iv

RESULTS AND DISCUSSION (cont) C. PHF2oLewis-Acid Adduct Formation 1) Reaction of HI and PHF2............... 70 2) Reaction of B2H6 and PHF2............ 74 Do Base Displacement by Difluorophosphine 1) Displacement of CO from Ni(CO)4...... 90 2) Displacement of PH3 and PF3........... 94 IV. Relative Base Strengths of PHs, PHF2, and PF3 Towards BH3 A o General.................................. 95 Bo Structures of the PHF2, PH3, and PF3 Borane Adducts............................ 96 1) Difluorophosphine Borane.....o........ 96 2) Phosphine Borane.................... 96 3) Trifluorophosphine Borane............ 108 C. Phosphine Borane and Trifluorophosphine Borane Systems 1) Phosphine and Trifluorophosphine Borane......oo.... o.................. 118 2) Observations Concerning PH3 BHs and PF3 BH3 o o o..... o o o...o..... o. 119 D. Conclusions........o..o. o............. 124 V. The Nature of the P-B Bond.................. 127 VI Summary......o o..................... 142 EXPERIMENTAL Io General Procedureso...........o............. 143 IIo Starting Materials.oo o...................... 143 III. Preparation and Reactions of PF2I A. Preparation.o............................ 144 v

EXPERIMENTAL (cont) B. Reaction of PF2I and Cu20 - Synthesis of F2POPF.............................. 145 C. Reaction of PF2I and CuCN - Synthesis of PF2CN................................ 146 D. Reaction of PF2I and Mercury - Synthesis of P2F4.................................. 147 E. Reaction of PF2I, HI, and Mercury Synthesis of PHF........................ 147 IV. Reactions of PHF2 A. Hydrolysis of PHF2a............o........ 149 B. Aminolysis of PHF2...0000....00......00000..... 149 C. Reaction of PHF2 and B2H6................ 150 D. Reaction of PHF2 and Ni(C0)4............. 151 V. Base Displacement Reactions Ao PHF2 and PF3sBH3....0..0................ 152 B. PHF2 and PH3sBH3o...o o.....oo........... 152 C. PH3 and PF3sBH3................ o..... 154 APPENDIX A - Spectroscopic Notation and Procedures 156 APPENDIX B - Band Shape Analysis for Difluorophosphine 0....o............................... 160 APPENDIX C - The C-N Stretching Vibration in PF2CN..................................... 163 APPENDIX D - Analysis of X2AA'X'2 Systems......... 166 REFERENCES oo.................................... 172 vi

List of Tables'able Page 1 Characterization of Difluoroiodophosphine...o 14 2 Physical Constants of PF2I................... 15 3 Infrared & Raman Spectra of PF2I o..........o 21 4 Characterization of Cyanodifluorophosphine... 23 5 The Infrared Spectra of PF2CN, PF2C1 and P(CN)3 o o......oo..oo o.............o.oo. o o 28 6 Characterization of — Oxo-bisdifluorophosphineo..oo...oo.....o.o.o................ 31 7 Physical Constants of F2POPF2....o o o..oo 32 8 A Comparison of the Vibrational Spectra of F2POPF2 and F2P(0)-O-P(0)F2. o-. -...-...... 35 9 Characterization of Tetrafluorodiphosphine... 41 10 Vibrational Spectra of P2F400...000 -......o.. 48 11 Characterization of Difluorophosphine.o..-ooo 53 12 Comparisonof Coupling Constants and Geometry for PH39 PF3, F2PNMe2, and PHF2o.... 57 13 Comparison of Physical Properties for PH3, PF39 and PHF2 o...oo.o o o.. o o...... o. o. 58 14 Physical Constants of PHF2000 o....o......... 59 15 Infrared Spectra of PHF2..................... 63 16 Infrared Spectrum of PF2H-HIoOoo.....o..o... 73 17 Characterization of Difluorophosphine Boraneo 76 18 Physical Constants of PHF2~BH3.000., 0 ooooo 77 19 Coupling Constants of PHF2~BH30...0000..o.. 85 20 Fundamental Vibrations for PHF2oBH3.......0. 87 vii

Table Page 21 The Infrared Spectrum of PHF2aBH3............ 88 22 The Infrared Spectrum of (PHF2)Ni(C0)3...... 92 23 Symmetry and IR-active C-0 Stretching for LxNi(CO 4-x, x = 14.......... 93 24 Vibrational Spectra of PH3-BH............ 106 25 Vibrational Spectra of PF3~BH3o3...o...o.... 117 26 Equilibrium Data for PH3oBH3....o..... 120 27 NMR Data for PF3oBH3, PHF2oBH3, and PH3BH3.. 126 28 31 p and 19F NMR Data for PHXF3-x (x = 0,1,), and their Boranes, and 31 Data for PHxMe3-x (x = 0->3).......................... 136 viii

List of Figures Figure Page 1 31p NMR Spectrum of PF2 I()........ 16 2 19F NMR Spectrum of PF2I()................. 8 3 Infrared Spectrum of PF2I(g)................ 19 4 Raman Spectrum of PF2I().................. 20 5 19F NMR Spectrum of PF2CN()............... 24 6 31p NMR Spectrum of PF2CN(j)................ 25 7 Infrared Spectrum of PF2CN(g)........ 27 8 Infrared Spectrum of F2POPF2(g)............ 34 9 19F NMR Spectrum of F2POPF2()........... 37 10 31p NMR Spectrum of F2POPF2()...00. 358 11 19F NMR Spectrum of F2PPF2()....... 42 12 31P NMR Spectrum of F2PPF2()......... 43 13 Some Possible Configurations of F2PPF2......45 14 Infrared Spectrum of F2PPF2(g).......... 46 15 Raman Spectrum of F2PPF2(j)................. 47 16 EPR Spectra of F2PPF2..o. o.............. 50 17 1H NMR Spectrum of PHF2(,)............. 54 18 19F NMR Spectrum of PHF2()............... 55 19 31p NMR Spectrum of PHF2().. 56 20 Infrared Spectrum of PHF2(g) o...oo........ 61 21 Infrared Spectrum of PHF2(s)......... 64 22 Infrared Spectra of F2PNMe2, FHPNMe2, HP(NMe2)2, & HNMe2............ 69 ix

Figure Page 23 Infrared Spectrum of PF2HHI(g).............72 24 1H NMR Spectrum of PHF2'BH3()).............. 79 25 1H NMR Spectrum of PHF2BH3(1) One Member of Borane Quartet........... 80 26 1H NMR Spectrum of PHF2.BH3(i) One Member of Phosphine Doublet............ 82 27 11B NMR Spectrum of PHF2'BH3(j)........... 83 28 19F NMR Spectrum of PHF2.BH3()......... 84 29 Structure of PHF2'BH30......................86 30 Infrared Spectrum of PHF2 BH3(g)............ 89 31 Infrared Spectrum of (PHF2)Ni(CO)3(g)...... 91 32 11B NMR Spectrum of PH3.BH3().............. 98 33 1H NMR Spectrum of PH3'BH3()............... 100 34 Infrared Spectrum of PH3-BH3(s)............. 103 35 Raman Spectrum of PH3'BH3(s)................ 104 36 Raman Spectrum of PH3~BH3()..)......105 37 1B NMR Spectrum of PF3'BH3()....... 109 38 1H NMR Spectrum of PF3'BH3()............. 111 39 19F NMR Spectrum of PF3-BH3(j)......... 114 40 Expected 19F NMR Pattern for PF3.BH3........ 115 41 Infrared Spectrum of PF3.BH3(g)............. 116 42 Plot of log Kdio vs 1/K for PH3 BH3..... 122 43 Hypothetical Sigma- and Pi-Bonding Contributions to Adduct Stability........... 130 44 A Possible Back-Donation Mechanism........... 131 x

Figure Page 45 Hypothetical Contributions of Fluorine and Hydrogen to Phosphorus Lone-Pair Polarizability............................ 133 46 31p NMR Spectra PHxF3-x-BH3 (x = 0,1,3)..... 14 47 Principal Axes of PHF2..... l000o 160 48 Band Shape and Type for PHF2.............. 161 49 C-N Stretching Vibration of PF2CN(g)....... 164 50 Calculated and Observed'9F NMR Spectra of F2POPF2...................... -oo o 167 xi

Abstract DIFLUOROPHOSPHINE LIGANDS, THEIR PREPARATION, PROPERTIES, AND CHEMISTRY by Ralph William Rudolph The nature of the P-B coordinate bond in various phosphine-borane systems is presently the subject of considerable controversy. This study was concerned with the preparation, properties, and chemistry of a series of ligands, the difluorophosphines, PF'2X, which should in theory be closely related to PF3 in bonding properties. The routes leading to known difluorophosphines were summarized. The preparation and characterization of the new mixed phosphorus halide, PF2I, led to two new synthetic methods for the preparation of difluorophosphine ligands. The new species PF2CN and F2POPF2 were prepared by a metathetic reaction between PF2I and the appropriate copper(I) salt according to the equation: nPF2I + CunX - (PF2)nX + nCuI X Xn(x = valence of anion X). A coupling reaction effected the preparation of F2PPF2 and PHF2 according to the following equation: PF2I + RI + 2Hg - Hg2I2 + PF2R (R = H, PF2). The new species, difluoroiodophosphine, cyanodifluorophosphine, i-oxo-bisdifluorophosphine, tetrafluorodiphosphine, and difluorophosphine were characterized and xii

compared with similar known compounds such as PF2C1, P(CN)3, F2P(O)-O-P(O)F2, N2F4, PF3, PH3, and NHF2. In most cases, infrared, Raman, and nmr spectroscopy were used to obtain information concerning structure and bonding. Preliminary investigations explored the hydrolysis and aminolysis of PHF2. The new ligand, PHF2, was also found to form adducts with BH3 and HI, and effect the displacement of CO from Ni(C0)4 to form difluorophosphine-nickel-carbonyl-derivatives. Base displacement reactions and other evidence demonstrated that the relative base strength towards borane decreases in the order PHF2 > PF3 > PH3. NMR spectroscopic studies showed that the unexpected stability of PHF2.BH3 cannot be attributed to an unusual structure. The nmr data for the series PHxF3-x'BH3 (x = 0,1,3) were discussed and compared to data for the series PHxMes-x'BH3 (x = 0-3) in relationship to structure and bonding. The unexpected stability of PHF2'BH3 demands at least a two-parameter model to rationalize the behavior. Some schools of thought will use a synergistic balancing of sigma- and pi-bonding to explain the trend. However, it is postulated that an F —H —F interaction may exist in PHF2. Such an interaction is "pictured" as resulting in a marked loosening of the phosphorus lone-pair electrons and a concommitant increase in the base strength of PHF2 towards BH3 when compared with PF3 and PH3. The latter model also implies a pyramidal geometry for PHF2 with short F —F and H —F non-bonded distances. xiii

Historical Background When Chatt and Williams prepared (PF3)2PtCl2 and [(PF3)PtCl2]2 they noted their close physical and chemical similarity to the corresponding carbonyl complexes.(1,2) (3) It had been suggested by Pauling( that bonding in the metal carbonyls involves the filled d-orbitals of the transition metal in double bonding. Thus, it is not surprising that the similar coordinating ability of PF3 and CO led Chatt to postulate that filled d-orbitals on the acceptor metal atom were essential for coordination with PF3. He felt that the electronegative nature of the fluorines attached to phosphorus would markedly reduce the ability of PF3 to coordinate via a classical dative bond. Thus, the existence of a second type of bonding was postulated to explain the nature of the chemical bond in trifluorophosphine complexes, i.e., a double bond which involved a weak a-bond as well as a w-bond utilizing the filled d-orbitals of the metal and the vacant 3d-orbitals in PF3. Since boron and aluminum do not have filled d-orbitals, Chatt did not expect them to form stable complexes with PF3. He was unable to prepare trifluorophosphine complexes of aluminum chloride or aluminum bromide and observed that Booth and Walkup could not effect the formation of PF3'BF3(. Subsequently, in direct conflict with the predictions - 1 -

- 2 - (5) of Chatt, Parry and Bissot prepared PF3'BH3( and found it to be similar in many respects to BH3'CO(6. To explain the existence of PF3'BH3, Graham and Stone(7) expanded the arguments which Bauer (8and Burg(9) had forwarded to explain the existence of BH3oCO and the stability of phosphinoborane polymers. They argued that PF3~BH3 and BH3~CO might owe their existence to supplementary w-bonding which involved the overlap of the vacant 3d-orbitals of phosphorus and a pseudo v-orbital formed by the hydrogen atoms of the borane group. In this light, Graham and Stone thought that it was significant that PF3 did not bond to other Group-III acceptor molecules, like BF3, yet does form stable complexes with (10) subgroup elements, e.g., Ni(PF3)4, where bonding is believed to be multiple in character. Surprised at the existence of PF3'BH3 and BH3~CO, the arguments of Chatt were subsequently modified(1) so as to be similar to the "hyperconjugation" effect(12) of organic chemistry and to be more in line with the ideas of Burg, Graham, and Stone. Although Chatt(1'2) originally reported that aluminum chloride does not form a complex with PF3, under different conditions Alton(l13) has been able to characterize PF3~AlCl3 Thus, the predictive value of the w-bonding arguments has been reduced and a modification of assumptions is demanded. In fact, about the time of Alton's work(13), Stone(l1) conceded that "by using a strong

- 3 enough acceptor atom (boron is rather weak in this respect compared with aluminum or gallium) it might well be possible to make a PF3 adduct of a Group III acceptor molecule in which the dative bonding could be described in terms of a classical a-bond." Alton 3)offered an alternative rationalization to correlate the difference in behavior of PF3 toward various acids. It was noted that bonding to phosphorus is strongly field-dependent.( 5) Since phosphorus is a relatively large nucleus, a considerable gain in energy may result if a phosphine ligand is able to closely approach the reference acid where the field strength would be greater. Phosphorus also should form stronger bonds to a given reference acid as the polarizability of the lone-pair on the phosphorus increases. Thus, the alkyl phosphines with polarizable lone-pair electrons are strong bases towards most reference acids, but in the case of PF3 even though the electron pair is still available, its bonding propensity is reduced by its lower polarizability and bonding will result only in the case of close approach by the reference acid. Alton estimated that the reorganization energy required to deform BH3 into a bonding configuration (sp to'sp ) was significantly smaller than for the corresponding change in BF3s Thus, a closer approach of borane is possible and bond formation is possible, whereas, the energy required to deform BF3 and allow close approach precludes bond formation.

-4 - Apparently the deformation energy of aluminum chloride is also small enough to be overcome by coordinate bond formation. The ability of aluminum chloride to dimerize to A12C16 (like borane to B2H6) is probably an indication of its low reorganization energy. Although some of Alton's ideas are presented in the more recent literature(l), most authors still appear to favor a supplementary d - pv back-bonding to explain the existence of PF3~BH3o Likewise, the ability of BH3 to reverse the "normal" order of coordinating ability (N > P; 0 > S) has been explained by Graham, Stone, Coyle and (7,17-19) Burg(779) in terms of similar w-bonding arguments. It should be noted however, that in these cases the "reversal" can be explained in terms of a change in size of the donor atom and a change in the nature of the donor electron pair so that the polarizability of the pair and the field strength of the reference acid are of basic importance. Shore and co-workers have recently shown that this apparent "reversal" of coordinating ability is a function of the strength of the reference acid(20) Thus, the nature of the chemical bond between phosphines and Main-Group-III acceptors, in particular borane, is still subject to considerable disagreement. Since the original work of Chatt and Wilkinsonl2l on the PFs complexes of platinum and nickel, interest in the similarity of PF3 and CO as ligands in transition metal chemistry has grown considerably. Subsequent work

- 5 - on such systems has been conducted by Clark, Kruck and others(2-23) Among the transition metal complexes of PF3, where there are filled d-orbitals on the metal to participate in a dw - pw interaction with the vacant 3d-orbitals on phosphorus, there appears to be no serious disagreement with the existence of double-bonding which supplements the dative P-metal bond in a synergistic manner. There is some debate, however, as to the influence which such double bonds have on compound properties. For example, the Ni-P stretching force constant, in a recent Raman-spectroscopic investigation (24)of Ni(PF3)4, was found quite clearly in the range for single bonds. Since difluorophosphine ligands are closely related to PF3, their synthesis and coordination chemistry have been of recent interest; Schmutzler(25) has recently published a comprehensive review on the fluorides of phosphorus including a summary of the chemistry of difluorophosphine ligands.

Statement of the Problem Since the nature of the coordinate bond in various phosphine-borane systems has been the subject of considerable controversy, it was desired to study a series of phosphines in which the w-bonding propensity, theoretically, should be relatively large. The difluorophosphines should prove ideal for such a test of theory because the electronegative fluorines lower the energy of the 3d-orbitals on phosphorus and thus make them more readily available as acceptor orbitals. Therefore, a study involving the synthesis, characterization, and chemistry of several difluorophosphines, PF2X, was undertaken. As a result, the interesting ligand difluorophosphine, PHF2, was discovered and subjected to further study. In particular, the stable species PHF2~BH3 was compared with PH3~BH3 and PF3~BH3 in an attempt to gain a greater understanding of the nature of the P-B bond. - 6 -

Results and Discussion I. Synthetic Methods for the Preparation of Difluorophosphine Ligands A portion of the present study was spent in the development of new preparations for difluorophosphine ligands. Section I of the discussion is an attempt to classify the known and newly developed preparations into various "type-reactions"; however, two novel syntheses of dimethylaminodifluorophosphine which could not be adequately classified have also been reported(26'27) One involves the reaction of boron trifluoride with trisdimethylaminophosphine (26) 2BF3 + P(NMe2)3 -> F2PNMe2+ 2Me2NBF2. The other is the aminolysis of trichloromethyldifluoro(27) phosphine(27) C13CPF2 + Me2NH - CC13H + F2PNMe2. The remaining routes to difluorophosphines have been classified as A) Fluorination of the Corresponding Halophosphine, B) Deamination of Dimethylaminodifluorophosphine, C) Aminolysis, D) Metathesis, and E) Coupling. Several new difluorophosphines have been prepared during the course of this research - difluoroiodophosphine (PF2I)l 4-oxo-bisdifluorophosphine (F2POPF2), cyanodifluorophosphine (PF2CN), tetrafluorodiphosphine (F2PPF2), and difluorophosphine (PHF2). The characterization of these new - 7 -

-8 - species is discussed in Section II. A. Fluorination of the Corresponding Halophosphine. fluorinating*. RPX2 flurinating RPF2; X= C1, Br, I agent (Eqn. 8-1) The early patent literature(28) (1936-39) describes this method for the preparation of difluorophosphines, however, the description of F2PN(C2H5)2 differs considerably from that mentioned in subsequent research(29). More recently, Martin(30'31) and co-workers used the method of equation 8-1 for the preparation of difluorophosphites, ROPF2o Other studies indicate that the fluorination of alkyldihalophosphines(32'33) does not proceed smoothly but is accompanied by oxidation-reduction and the formation of fluorophosphoranes, RPF40 When more electronegative substituents are placed on the phosphorus, or in the case of dichlorophosphites (ROPC12) and dialkylaminodichlorophosphines (R2NPCl2), the fluorination occurs more smoothly without appreciable oxidationreduction(26 29-31 354-37) Booth(38 39) used a similar type-reaction fluorinat ing* PX3 fluorinating- PF2X; X = C1, Br agent for the preparation of mixed phosphorus halides. Among those prepared were PF2C1 and PF2Br from PC13 and PBr3, respectively. More recently, Holmes and Gallagher(4) * Typical fluorinating agents used are SbF3, AsF39 and NaF.

-9 - simplified Booth's procedure, but an even more convenient preparation of halodifluorophosphines will be discussed in Section I-Bo The fluorination of P(NCO)3 and P(NCS)3 with SbF3 to give PF2(NCO) and PF2(NCS)1(41) respectively, might also be classified with the above reactions since isocyanates and isothiocyanates are considered to be halogenoids. For the present work the general synthetic method of equation 8-1 was used for the preparation of dimethylaminodifluorophosphine, a starting material; however, the procedure was not used for other syntheses since, as the above discussion indicates, the fluorination of halophos(25) phines has been rather exhaustively studied(2). B. Deamination of Dimethylaminodifluorophosphine. F2PNMe2 + 2HX - PF2X + Me2NH~HX; X = Cl, Br, I This route to halodifluorophosphines was developed by Fleming(42) and Moye(43) in our laboratories. Subsequently, it was reported independently(4445). Halodifluorophosphines have been prepared previously by the partial fluorination of the appropriate phosphorus trihalide(38-0) However, these methods yield mixtures which are difficult to separate. The action of a hydrogen halide on dimethylaminodifluorophosphine provides a unique route for the preparation of pure PF2X ligands. Of special interest in this study was difluoroiodophosphine which had been implied by Cavell(44) but never

- 10 - characterized. During this work it was found to be stable enough for a complete characterization and was used as a reactive intermediate for the introduction of the PF2moiety into new molecules. A discussion of the synthetic methods which utilize the lability of the P-I bond can be found in Sections I-DE. C Aminolysis. PF2X + 2HNR2 - F2PNR2 + R2NH'HX; X = F, C1 The interaction of an amine and a phosphorus halide is a general method for the preparation of aminohalophos(46547) phines (4647) which has not been applied extensively to (42) phosphorus trifluorides. Fleming(42) developed this f(37) route to F2PNMe2, but Van Doorne7 found that PF3 and N-methylaniline do not undergo this type-reaction. This (44) method has been independently reported by Cavell 44 who found that, in addition to PF3, PF2C1 also yields F2PNMe2 on aminolysis with dimethylamine. These studies indicate that this may be a convenient route to selected aminodifluorophosphines, and that aminolysis will probably proceed more smoothly as the halogen substituent is varied from Cl through I, the order of increasing P-X bond lability; however, this postulate was not tested during the present work. D. Metathesis. PF2I + CuX - PF2X + CuI

- 11 - This route to difluorophosphines was discovered during the course of this study when small amounts of F2POPF2 were isolated from the reaction mixture after PF2I had been treated with powdered copper metal in an attempt to prepare P2F4o As subsequent investigation showed the F2POPF2 must have resulted from small amounts of Cu20 in the copper metal since treatment of PF2I with Cu20 gave F2POPF2 in good yields. Apparently this metathetic reaction with cuprous salts may be extended to some degree since when PF2I was treated with cuprous cyanide, PF2CN was isolated and characterized. E. Coupling. PF2I + RI + 2Hg - PF2R + Hg2I2 The preparation of the parent member of the difluorophosphine homologous series, PHF2, was effected by this route with HI as the source of hydrogen. The reaction is similar to that used by Burg and Mahler(48) for the preparation of (CF3)2PH from (CF3)2PI and a proton source in the presence of mercury. (49) Previously, Bennet et al.9 had isolated tetrakistrifluoromethyldiphosphine, (CF3)2PP(CF3)2, from a similar coupling-reaction by treating (CF3)2PI with mercury. It is not surprising then that tetrafluorodiphosphine, F2PPF2, has also been isolated in this study from an analogous reaction between PF2I and mercury. It

- 12 - is felt that a wide variety of compounds containing the PF2-moiety might be made by analogous coupling-reactionso In fact, preliminary results indicate that F2P-CH2-CH=CH2 is produced when PF2I and allyl iodide are shaken with (50) mercury o IIo Characterization of New Difluorophosphines. The characterization of the new difluorophosphine ligands prepared in this study is described below. Techniques such as nmr, infrared, and mass spectroscopy, as well as more classical methods were used to obtain a complete characterization. A typical experiment describing the preparation of each new species is described in the Experimental Section. A. Characterization of Difluoroiodophosphine, PF2I. Difluoroiodophosphine was obtained in 94% yields from the reaction of HI and F2PNMe2 according to the equation F2PNMe2 + 2HI - PF2I + Me2NHOHI. (44) Although Cavell 44 indicated that PF2I was too unstable for characterization, it was found that at 25~ in the gas phase, in a pyrex container, after 1 day, and with p(initial) = 37 mm, PF2I was ca. 1% decomposed; after 9 days it was 19% decomposed. However, the rate of disproportionation according to the equation 3PF2I - 2PF3 + PI3 was pressure dependent since with p(initial) = 280 mm,

- 13 - after 1 day the PF2I vapor was 6% decomposed, and after 8 days 87% decomposition had occurred. An unequivocal characterization of PF2I is obtained fron the elemental analysis, vapor density molecular weight, and mass spectrum, tabulated in Table 1 A, B, and C, respectively. The mass spectrum confirms the vapor density molecular weight for the monomer PF2I and displays the expected fragmentation pattern. In the mass spectrum observed experimentally there were a number of additional minor peakso These extraneous peaks were ascribed to small amounts of impurities which resulted from the reaction of PF2I with traces of water in the vacuum system or to traces of unreacted F2PNMe2 in the PF2I sample. The vapor pressure data for difluoroiodophosphine are compared in Table 2 with those calculated from the vapor pressure equation. Also included in Table 2 are the freezing point, boiling point, and Trouton constant for PF2I. The phosphorus and fluorine nmr spectra of PFFI (Figure 1) display the expected spin-spin splitting patterns with P-F coupling somewhat smaller than in PF3 51) The 1:2:1 triplet (JpF = 1337 cps) centered 242.2 ppm downfield from OPA in the S31 nmr spectrum of PF2I (Figure 1) shows that two equivalent fluorines (21 + 1 5) are directly bonded to phosphorus; the doublet (JppF 1340 ops) displayed 31.9 ppm downfield from TFA

- 14 - Table 1 Characterization of Difluoroiodophosphine A. Elemental analysis. P F I found~ 15.37 18.09 64.34 calcd~ 15.81 19.39 64.78 B. Vapor density molecular weight. obsvd: 192.9 g./mole calcd: 195.9 g./mole C. Mass spectrum. relative m/e peak height assignment 196 100 PF2I~ 177 15 PFI~ 158 6 PIT 127 71 I+ 88 1 PF3' 69 60 PF2a 63.5 4 2+ 50 29 PF+ 31 11 P 19 F

- 15 - Table 2 Physical Constants of PF2I A. Vapor pressure data. ~C mm (obsvd) mm (calcd) -63.5 5.0 5.0 -45.6 19.8 18.9 -30.6 48.2 48.7 -22.9 74.5 75.8 0.0 240.1 244.4 B. Equation. log p(mm) = -1514 + 7.929 T C. Trouton constant b.p. (extrapolated) 23.09 cal./ deg. mole 26.7~C m.p. -93.8 to -93.3~C

Figure 1 H _40.4 Me 31P NMR Spectrum of PF2I(j) @ 370 JpF = 1337cps OPA 6 = -242.2 ppm b = -242.2 ppm

- 17 - in the 19F nmr spectrum (Figure 2) further indicates that the fluorine nuclei are bonded to a single atom (I =) o It has been noted that P-F spin-spin coupling constants correlate qualitatively with the s-character of the abonding hybrid of phosphorus(51) When compared with JpF(PF3) = 1441 cps(51) the JpF(PF2I) = 1340 cps indicated that the P-F bonds in difluoroiodophosphine have less's-character" than those in PF3 which is usually considered to be close to an sp' hybrid. Couched in more fundamental terms, the F-P-F bond angle in PF2I is probably less than that in PF3 (1040)(52) The infrared and Raman spectra of difluoroiodophosphine are displayed in Figures 3 and 4, respectively. The observed frequencies and tentative assignments are tabulated in Table 3. Since both planar (C2v) and pyramidal (Cs) geometries for PF2I would be expected to display six fundamental vibrations, the geometry of the molecule cannot be confirmed from the number of observed vibrational frequencies. Five of the fundamentals were assigned from the infrared and the sixth was observed in the Raman effect at 202 cm', apparently too close to the cut off of the infrared spectrometer for observation. Polarization studies were precluded by the rapid decomposition of PF2I under the influence of the mercury arc light even at -40~o The resolution in the Raman spectra was apparently insufficient to resolve the two P-F vibrations expected near 830 cm-1 and the two other motions near

Figure 2 94.1 Mc 19F NMR Spectrum of PF2I( ) @ 370 v pTFA JPF = 1340 cps 6 = -31.9 ppm

Figure 3 The Infrared Spectrum of PF2I(g) frequency (cmI ) ______1000 800 600 400 200

- 20 - Figure 4 The Raman Spectrum of PF2I ( ) anti-stokes --— 1 ines 0 200 400 600 800 1000 1200 frequency (cm- )' I I I I' ---' * H 1 i nes

- 21 - 368cm,1 thus only four bands were observed. The P-F motions in the infrared were assigned on the basis of their similarity, both in vibrational-frequency and bandshape, to those of molecules like PF2C1, PF2Br, and PF3. Table 3 Infrared & Raman Spectra* of PF2I Infrared (g) Raman ()(-400~) freq. (cm ) freq. (cm-) & intensity note assgn. & intensity 850.4 s v PF s 828 w 845.3 vs vas PF 417.5 m R 411.8 m Q 5 FPF 404 s s 407.0 m P 378.5 m v PI 368 vs 372 w sh as PF2 5 PF2I 202 s * For an explanation of notation see Appendix A. B. Characterization of Cyanodifluorophosphine, PF2CN. Cyanodifluorophosphine was prepared by treating CuCN with difluoroiodophosphine as represented in the equation PF2I + CuCN - PF2CN + CuI. In contrast to P(CN)(53), a colorless, solghtly volatile solid, PF2CN is a colorless volatile liquid which is characteristically similar to the halogenoid-halides of

- 22 - phosphorus(54) Although, at this time the characterization of PF2CN is not extensive, the vapor density molecular weight, and mass spectrum shown in Table 4 offer good evidence for the formulation PF2CN. Experimentally, a separation of PF2CN from small amounts of impurities like HCN, PF2I, F2PNMe2 in the PF2I, and P0-containing species, was difficult because of the similarity in volatility among the compounds; minor peaks attributable to these species, although not all noted in Table 4, did appear in the observed mass spectrum. The intensity of the HCN peak, as noted in Table 4, is appreciable; however, the spectrometer is very sensitive to HCN, and the slight leak in the inlet system of the instrument may have led to partial hydrolysis of PF2CN and production of HCNo Nevertheless, the salient feature of the mass spectrum is a fragmentation pattern consistent with the formulation PF2CN. In fact the presence of a peak at m/e = 43(PC+) is good evidence that PF2CN is a true cyanide as opposed to an isocyanide which would be expected to give a peak at m/e = 45(PN+) The L9F nmr spectrum of liquid cyanodifluorophosphine at -20~ (Figure 5) consists of a simple doublet centered 11.7 ppm upfield from TFA (external std)o The doublet splitting of 1267 cps is attributed to spin-spin coupling of the fluorine nuclei with the single phosphorus nucleus. While the fluorine spectrum shows that one phosphorus is present in the molecule, the 31p spectrum (Figure 6)

- 23 - Table 4 Characterization of Cyanodifluorophosphine A. Vapor density molecular weight. obsvd: 97.5 g./mole calcd: 95.0 g./mole B. Mass spectrum. relative m/e peak height assignment 95 27.1 PF2CN+ 88 2.0 PF3+ 76 14.2 PFCN 69 100.0 PF2 57 1.9 PCN+ 50 17.6 PF+ 43 1.8 PC+ 38 1.4 PFCN2+ 31 13.3 P+ 27 12.6 HCN 26 9.6 CN+ 19 1.6 F 14 2.5 N

Figure 5 94.1 Me 1 F NMR Spectrum of of ~~~~~~~H —>PF2CN( 2) @ -200 -F=JpF = 1267 cps TFA 6 = 11.7 ppm I I_______________ I ----— V —— I

Figure 6 40.4 Mc 31 NMR Spectrum of PF2CN(I) @ -200 \J1 JF = 1273 cps PF"~~ ~OPA 6 = -140.8 ppm

- 26 - clearly establishes the presence of two equivalent fluorines bonded to the phosphorus. The 1:2:1 triplet (JPF = 1273 cps) in the 31p nmr is centered 140.8 ppm downfield from OPA. Even at -20~ there was significant disproportionation of PF2CN ( ) as evidenced by the formation of white solids [probably P(CN)3] in the nmr tube and the appearance of PF3 in the nmr spectrum. The equation probably is 3PF2CN - 2PF3 + P(CN)3. The infrared spectrum of PF2CN gas is shown in Figure 7; in Table 5, the vibrational frequencies derived from the spectrum are compared with those for P(CN)3 and PF2C1 in order to aid in making tentative assignments. If the cyanide moiety is considered as a point mass, i.e., if the motions due to cyanide [v sN, oasPCN, 6sPCN] are neglected, the spectra of PF2Cl and PF2CN would be expected to be very similar; under these limiting conditions they are similar as can be seen in Table 5. In fact, the band shapes in the P-F and P-X stretching regions are also very similar. Notably, the P-C1 bending and wagging motions could not be detected under the conditions used to determine the gas phase spectrum of PF2Cl. Likewise, under similar experimental conditions, for gaseous PF2CN the corresponding P-CN wagging and bending motions did not appear in the expected region (250 - 320 cm'1) (55 56a) However, according to the

Figure 7 I - The Infrared Spectrum of PF2CN(S) frequency (cml ) I A / I I I / 2400 2000 1000 800 600 400

Table 5 The Infrared Spectra* of PF2CN, PF2Cl, and P(CN)3 PF2CN(g) PF2Cl(g) P(CN) 3 56a) freq. (cm-1) freq. (cm-1) freq. (cm-1) & intensity note assgn. & intensity note assgn. & intensity assgn. 2193.9 vs sCN 2204 vs vCN 869.4 s v PF 864.5 s v PF 866.6 vvs Vs PF 855.5 vvs VasPF as as 633 s 2' bPC3 S 636.4 R 550.3 R 605 vs v PC3 629.0 vs Q v PC 543.7 vs Q v PC1 622.1 P 538 P 585 s vsPC3 549 w 6 PCN ~~~~~~~~s ~~464 w 6PCN 460.5 w t asPCN 451 w 421 R 344.6 w t bFPF 412.2 m Q b FPF 404 s 302 6b PC1t 312 m b PC3 s S 276 w b PC3 259t casPC1t 267 m-w * For an explanation of notation see Appendix A. t Indication of P and R branches. t Not observed in spectra determined on PF2C1 vapor in this study. Taken from Delwaulle and Francois.(55)

29 - assignments of Miller(56b) for P(CN)3 the latter motions would fall below 200 cm-1, the limit of our instrument. Unless the two "missing" frequencies are accidentally degenerate with those noted in Table 5, only seven of the expected nine fundamental vibrations for PF2CN have been observed. The expected pyramidal geometry (Cssymmetry) for PF2CN cannot be ascertained from the number of observed fundamentals in the infrared since the planar model (C2v symmetry) would also have nine fundamentals. Under high resolution the C-N stretching in PF2CN appeared to be composed of a number of lines spaced by 2.3 cm 1. The surprising number of lines observed in the C-N stretching region is probably attributable to "hot" bands as explained in Appendix C. Due to the lack of model compounds, and since the cyanide and isocyanide stretching frequencies fall close together(57, it is not felt that the infrared spectrum of PF2CN can be used to differentiate between the cyanide and isocyanide structure. C. The Characterization of 4-Oxo-bisdifluorophosphine, F2POPF2The interesting new species p-oxo-bisdifluorophosphine was isolated in 73% yield from the reaction of PF2I and Cu20 according to the equation 2PF2I + Cu20 -> F2POPF2 + 2CuI. This new species is the first highly volatile compound containing a P-O-P link. One closely related compound

- 30 - described in the literature(58'59) is diphosphoryltetrafluoride, F2P(O)-O-P(O)F2, which boils at 71~ as compared to -18.3~ (extrapolated from vapor pressure equation) for F2POPF2 Although some difficulty was encountered in obtaining a satisfactory elemental analysis for F2POPF2 by classical techniques, it is felt that a good characterization of the new species has been obtained on the basis of other evidence. The low resolution 31P and 19F nmr spectra of F2POPF2 (R) consist of a 1:2:1 triplet and a simple doublet, respectively; the coupling constants derived from both the 19F and 31p spectra agree closely and are comparable to P-F couplings observed in related compounds. Although a more complete discussion of nmr spectra will follow, the 1:2:1 triplet of the 31p nmr spectrum indicates that each phosphorus is bound directly to two fluorine nuclei. This "elemental analysis" by nmr, when considered with other data presented below, offers no reasonable alternative to the formulation F2POPF2. The mass spectrum of F2POPF2 (Table 6).has the expected fragmentation pattern. Since no peaks appeared at higher than m/e = 154, the parent peak for F2POPF2, the spectrum affords a confirmation of the vapor density molecular weight of 154.2 g./mole. The vapor pressure data for F2POPF2 were obtained and are summarized in Table 7 where they are compared

- 31 - Table 6 Characterization o f PI-Oxo-bisdifluorophosphine A. Vapor density molecular weight. obsvd: 154.2 g./mole calcd: 154.0 g./mole B. Mass spectrum. relative m/e peak height assignment 154 20.0 F2POPF2+ 135 2.6 F2POPF+ 88 2.5 PF3+ 85 3.3 PF2aO 69 100.0 PF2 66 4.3 PFO+ 50 16.8 PF+ 47 27.7 PO0 31 5.9 P 25 0.4 PF2+ 19 1.2 F 16 0.4 0 15.5 0.3 p2+

- 32 - Table 7 Physical Constants of F2POPF2 A. Vapor pressure data. ~C mm (obsvd) mm (calcd) -83.6 13.1 13.3 -78.5 19.5 20.0 -63.0 61.0 60.4 -50.0 141.6 143.1 -45.0 190.5 189.5 B. Equation. log p(mm) = -1300 + 7.981 T C. Trouton constant b.p.(extrapolated) 23533 cal./deg. mole -18.3oC m.p. -132.1 to -131.8~C

- 33 - with values calculated from the vapor pressure equation. Also included in Table 7 are the freezing point, boiling point, and Trouton constant for F2POPF2. Samples of F2POPF2 have been stored in clean glass tubes at 25~C and saturation pressure for one day with less that 1% decomposition. At lower pressures and temperatures the compound is even more stable; however, adsorbant surfaces like asbestos accelerate the disproportionation, i.e., a sample of F2POPF2 placed in contact with asbestos paper at 25~ was nearly 100% decomposed after one day according to the equation F2POPF2 -> PF3 + (POF). The infrared spectrum of gaseous i-oxo-bisdifluorophosphine is displayed in Figure 8. The comparison of the infrared spectrum of F2POPF2 with the Raman spectrum of (59) F2P(O)-O-P(O)F2 shown in Table 8 indicates the presence of a P-O-P oxygen bridge bond in F2POPF2. The salient difference between F2POPF2 and F2P(O)-O-P(O)F2 is the absence of the terminal P-O bonds in the former. Thus, the very intense absorption characteristic of P-O (the 1400 - 1200 cm-1 region)(59,60) is absent in the infrared of F2POPF2; however, absorption bands characteristic of a P-O-P bridged linkage are observed at 976.1 and 682 cm-1. The latter assignment of P-O-P vibrations agrees closely with that made by Corbridge(0) who observed that many species containing a P-O-P link absorb in the 970 - 870 and ~ 700 cm-1 regions. On the other

Figure 8 The Infrared Spectrum of F2POPF2(g) 1~~ frequency (cm- ) 1400 1200 1000 800 600 400

- 35 - Table 8 A Comparison of the Vibrational Spectra* of F2POPF2 and F2P(0)-O-P(O)F2 Infrared Raman ( F2POPF2(g) F2P(O) — P(0)F2()(5) freq. (cm-~) freq. (cm-1) Sc intensity note assgn. & polarization assgn. 1390 p vPO 1370 vPO 1270 dp overtone 1077 w overtone 1083 dp overtone 976.1 vvs br v POP 987 dp asPOP 863.1 vvs 951 dp) 853 vvs br vPF 890 p vPF 842 vvs 855 J 682 m v POP 721 p}P 518 PF2 scissors 519 m 1 PF2 515 m sh scissors 480 POP 440 p PF2 scissors 460 w pPF2 595 1559Q.P31562 pPFa, 359 w TPF2 362 pPF2, 352 wP0~ 340 and 295 torsional 273 dp modes 205 POP 160 p POP 16oo e p scissors * For an explanation of notation see Appendix A.

- 36 - (59) hand, Robinson( assigned absorptions around 700 cm-1 in C12P(O)-O-P(0)Cl2 and F2P(O)-O-P(O)F2 to P-0 wagging motions. Such P-0 wagging motions are precluded in F2POPF2, and since the absorption at 682 cm-' seems too low to be assigned to a P-F stretch and too high to be attributed to any PF2 deformation, it has been assigned to the assymetric P-O-P stretch. Also in accord with the assignments made here Griffiths(6) has made the following assignments for (CF3)2POP(CF3)2: vasPOP = 925 and = 715 cm-'1 The high resolution 19F nmr spectrum of F2POPF2(I) is shown in Figure 9. The spectrum remained unchanged from -80 to 370, and therefore, it seems unlikely that the "unusual" splitting pattern can be attributed to temperature dependent phenomena like exchange or internal rotation. The spectrum is probably best described as a "second-order" A2X4 type. An analysis of this spectrum in terms of an X2AA'X'2 system is found in Appendix D. The phosphorus spectrum (Figure 10) also appears to be consistent with an X2AA'X'2 type, however, analysis is incomplete at this time. The 31p ( = -111 ppm, OPA std) and 19F chemical shifts (6 = -39.9 ppm, TFA std) of F2POPF2 are both in good agreement with the corresponding values for PF3 [6(P) = -105 ppm; b(F) = -42.3 ppm]; this agreement is not surprising if F2POPF2 is considered as two PF3 molecules with a bridging oxygen, fluorine's neighbor

Figure 9 94.1 Me 19F NMR Spectrum of F2POPF2(j) @ 370 1344 cps TFA __ J —--- ^ -------- I —-- AvIV ------------ 50 cps/div 6 = -39.9 ppm

Figure 10 40.4 Mc 31 NMR Spectrum of F2POPF2(I) @ 370 1354 cps - 354 cps I OPA lI I I I I I 50 cps/div 6 = -111 ppm

- 39 - in the periodic chart, replacing a fluorine atom in each PF3 unit. Do Characterization of Tetrafluorodiphosphine, P2F4. The literature(62-64) mentions several attempted preparations of tetrafluorodiphosphine via the fluorination of a tetrahalodiphosphine, P2X4. In the course of this study, P2F4 was isolated in 94% yield when PF2I was shaken with mercury. The equation for the reaction is 2PF2I + 2Hg - P2F4 + Hg2I2. High yields were obtained when the initial pressure of PF2I was around 50 mm; however, in another experiment when the initial pressure of PF2I was raised ca. 450 mm, the main reaction was disproportionation of the PF2I according to 3PF2I - 2PF3 + PI3. When the P2F4 was separated from the unreacted PF2I (see Experimental), a small amount of F2PNMe2 was detected in the unreacted PF2I. Perhaps the presence of F2PNMe2 in the PF2I was a factor which promoted the formation of (45) P2F4 since others(4) have unsuccessfully tried to prepare P2F4 from the reaction of PF2I and Hg. Although the characterization of tetrafluorodiphosphine is not extensive, several pieces of evidence confirm the formula P2F4. A mass balance of products and reactants (see Experimental) shows that two moles of PF2I react with mercury to form one mole of a species having a vapor

- 40 - density molecular weight of 140.1 g./mole; theory for the monomer P2F4 is 138o0 g./mole. The principal ions which appear in the mass spectrum (Table 9) correspond to the fragmentation pattern anticipated for P2F4; the m/e = 138 peak, which corresponds to the P2F4 parent ion, confirms the vapor density molecular weight. Other minor peaks attributed to fragments from F2POPF2 and PHF2 were also observed in the spectrum but were not mentioned in Table 9. Since after P2F4 has been transferred a number of times in the vacuum system the sample is found to contain F2POPF2 and PHF29 it is not surprising that the latter species appear in the mass spectrum. It is felt that the PHF2 and F2POPF2 result from hydrolysis of the P2F4 by trace amounts of water adsorbed on the stopcock grease. Perhaps the equation is 2P2F4 + H20 -2 2PHF2 + F2POPF2. Tetrafluorodiphosphine and diphosphine are probably essentially isostructural and in both cases all nuclei are magnetically active with I = 2-. Since P2H4 has been observed to give "complex" nmr spectra which can be interpreted in terms of an X2AA'X'2 system(65)' it is to be expected that P2F4 should also have X2AA'X'2 nmr spectra. The 19F nmr spectrum of P2F4(1) has been determined at -20~, and as shown in Figure 11, it is indeed complex as is the 31p spectrum (Figure 12). Although analyses of these "second-order" spectra are not complete at this time, it is felt that they will be consistent

- 41 - Table 9 Characterization of Tetrafluorodiphosphine A. Vapor density molecular weight. obsvd: 140.1 g./mole calcd: 138.0 g./mole B. Mass spectrum. relative m/e peak height assignment 138 29.5 P2F4 119 10.3 P2F3~ 88 7.8 PF3+ 69 100.0 PF2 50 26.5 PF+ 31 12.4 P+ 25 2.2 PF2+ 19 3.3 F 15.5 1.2 p2+

Figure 11 94.1 Me NMR Spectrum of -of —-— 1130. cps —--- F2PPF2(9) @ -201 T4T;A~~~~~~~~~~~~~~~~~~~~~~~~~~~~~~4 ro 6 = 37.6 ppm

Figure 12 40.4 Me 31 NMR Spectrum of P2F (g) @ -20Q -C-1- 1, 25 cps' e 1 | - 1 1 25 cp s y ^J(^\A. —^J/VU-y^ I-^A^^^iM^^AJUA. I 10I I i 100 cps/div 5 = -226 ppm

- 44 - with an X2AA'X'2 system. The problem of analysis is discussed in more detail in Appendix D. A number of configurations may be proposed for the F2PPF2 molecule. Some reasonable configurations for F2PPF2 and their respective symmetries are illustrated in Figure 13 along with the infrared and Raman activity predicted for each symmetry. The infrared and Raman spectra of F2PPF2 are displayed in Figures 14 and 15, respectively; the vibrational frequencies and tentative assignments derived from these spectra are shown in Table 10. A comparison of the vibrational spectra predicted for each form of F2PPF2 (Figyure 13) with those actually observed indicates that tetrafluorodiphosphine probably exists in the trans conformation. The latter conclusion was reached as discussed below. The mutual exclusion rule(66) states that in molecules containing a center of symmetry no coincidences should occur between transitions allowed in the infrared and those allowed in the Raman. The near overlap of a few frequencies in the infrared and Raman spectra of P2F4 determined in this study is probably accidental. This accidental overlap may be due to frequency shifts caused by the different physical states of P2F4 used in determining the respective vibrational spectra. Thus, the mutual exclusion rule apparently holds, pointing to either a trans or planar form for the molecule. The appearance of only seven bands in the Raman spectrum, one of which

Figure 13 Some Possible Configurations for F2PPF2 Planar (D2h) Trans (C2h) Cis (C2v) Semi-Eclipsed (C2) Gauche (C2) Symmetry Modes of Vibrational Spectra IR Raman IR Raman IR Raman IR Raman Blu 3 A 3 Au 4 A 4 A1 4 A1 12 B 12 B 2 B2u 2 Blg 3 Bu 2 Bg 3 B1 3 A2 2 B3u 1 29g 2 B2 3 B1'inactive A, 2 B2 uI

PF \ PFFI ~~~~~Fi ~gure 1 ~4 ~0I The Infrared Spectrum of P2F4(g) frequency (cm-), 1000 800 600 400 200

- 47 - l| polarized ~. polarized * Hg line frequency., a I i i I I -, a (cm-1) 0 200 400 600 800 1000 1200 1500 Figure 15 The Raman Spectrum of P2F4() @-20O

- 48 - Table 10 Vibrational Spectra of P2F4* Infrared Raman P2F4 (g) P2F4(C) at -20~ freqo (cm -) freqo (cm-') & intensity note assignment & intensity note p PF2 214 m p 361 m Pas PF2 365 m asPF2 6sPF2 377 s p 378.7 vw PF2I impurity 2'p PF2 403 vw, sh p 417.8 412. 2 vw PF2I impurity 406.8) 6 PFa? 453 w dp? s -as v PP 541 vs p v- PF 803 w, sh dp? s -as 827.8 vvs P v PF 833.9 vvs R as vs PF 825 m p? 839 vvs P v PF 846.9 vvs R as * For an explanation of notation see Appendix A9

- 49 - is probably an overtone, is evidence against the cis, gauche, and semi-eclipsed conformations as possible structures since all twelve vibrations should be active in these cases. Futhermore, the observation of a P-P stretching vibration at 541 cm-1 in the Raman and absence of the corresponding motion in the infrared spectrum would not be expected if the cis, gauche, or semi-eclipsed forms of P2F4 were correct. Finally, a consideration of the number and distribution of polarized Raman lines offers good evidence for the elimination of the planar form of P2F4 from further consideration. Although polarization is possible only in the case of totally symmetric lines, it is not always necessarily observed( Since no polarization was apparent in the P-F stretching region, the totally symmetric P-F stretch must not display a polarization effect. In addition to a P-F stretch, two totally symmetric vibrations are expected for planar P2F4 and three for trans molecule. The observation of three definitely polarized lines in the Raman effect, well outside the P-F stretching region, therefore strongly indicates that the trans form represents the correct structure for P2F4. (67-69) Since N2F4 is known to yield NF2 free radicals 79 it is of interest to investigate the epr spectrum of P2F4o A broad signal was indeed found for all the samples of P2F4 investigated (Figure 16). In the case of the single line observed for neat P2F4(J) the g-value was determined

- 50 - Figure 16 A. EPR Spectrum of P2F4(9) @ 200 12 gauss/div. H46 gauss/div. 2 B. EPR Spectrum of P2FA(g) or P2Fn in CC4 20

- 51 - to be 2.00129(Figure 16A). More complex spectra were observed for neat P2F4(g) and P2F4 in CCl4; the spectra appeared to be the same for the gas and solution and fine structure was noted on the high-field line (Figure 16B)o It should be noted that the signals observed for P2F4 were weak; because of the high sensitivity of the epr spectrometer, the observed signals may be due to impurities. Eo Characterization of Difluorophosphine, PHF2o The parent member of the difluorophosphine homologous series, PHF2, was prepared by the "coupling-reaction" represented in the equation PF2I + HI + 2Hg -> PHF2 + Hg2I2o Nitrogen has been known to form haloamines for some time(70 72) The literature also mentions the attempted characterization of halophosphines, but no real halophos(73- 74) phines were ever isolated (73' ) Therefore, PHF2 is the first fully characterized halophosphine. Recently, however, Blaser and Worms(75'76) reported the isolation of an halide of hypophosphorus acid; when hypophosphorus acid was dissolved in excess liquid HF at -78o and the solution then allowed to warm slowly to room temperature, they obtained PH2F3o From a similar reaction between HF and phosphorus acid they isolated PHF4 which is not as stable as PH2F3 and decomposes at room temperature to give HF and PFs, irreversiblyo The recent Russian liter(77) ature also reports a new class of hydrogen-containing

- 52 - aryl phosphorus fluorides which were isolated during attempts to prepare fluorophosphines from aryldichlorophosphines and potassium bifluoride: ArPC12 + 2KHF2 - 2KC1 + HF + ArPF3H. In contrast to difluoroamine, NHF2, difluorophosphine never underwent violent disproportionation during the course of the present investigation. It is a stable, colorless gas which can be maintained at 25~ and saturation pressure for 5 hr. with less that 5% decomposition. At lower pressures and temperatures it is even more stable. Although some difficulty was encountered in obtaining a complete elemental analysis, the formula PHF2 is supported unequivocally by other data such as the hydrogen analysis, the vapor density molecular weight, and the mass spectrum, which are shown in Table 11. The proton, fluorine, and phosphorus spectra (Figures 17, 18, and 19, respectively) alone unequivocally confirm the PHF2 formula. The vapor density molecular weight is consistent with that for the monomer PHF2 and is confirmed by the mass spectrum(Table 11) which had no peaks at higher m/e values than 70 and displayed the fragmentation pattern anticipated for PHF2. The proton nmr spectrum (Figure 17) shows a doublet, due to P-H spin-spin coupling, with each member of the doublet split into a lo 21 triplet by spin-spin coupling

- 53 - Table 11 Characterization of Difluorophosphine A. Hydrogen analysis. obsvd 1.46 calcd for PHF2a 1 43 B. Vapor density molecular weight. obsvd~ 70.8 g./mole calcd~ 70.0 g./mole C. Mass spectrum. relative m/e peak height assignment 70 100.0 PHF2 69 52.4 PF2' 51 84.3 PHF+ 50 27.5 PFE 34.5 o.8 PF22' 32 3.9 PH+ 31 16.0 P 25.5 1.0 PHF2+ 25 3.0 PF2+ 20 o.6 HF+ 19 4.6 F 16 1.2 PH2+ 16.5 2.5 p2"

Figure 17 H —- 60 Me H NMR Spectrum of PHF2(l) @ -200 | __JFH p = 41.7 cps JpH = 182.4 cps SMS 6- -7.65 ppm

Figure 18 94.1 Me 19F NMR Spectrum of PHF2() @-20 540 cps HPF * 1135 cps.PF TFA 6 = 42.2 ppm

Figure 19 31 40.4 Mc P NMR Spectrum of PHF2(1) @ -20~ H LI ILS I ~ JpH = 184 cps OPA JPF = 1143 cps] A 6 = -224 ppm

- 57 - of the hydrogen with two equivalent fluorine nucleio The P-H coupling constant is 182.4 cps and the F-H coupling constant is 41.7 cps. The P-H coupling constant is discussed further below. The fluorine nmr signal (Figure 18) is also split into a doublet by coupling with a single phosphorus nucleus' and each member of the original doublet is again split into a doublet by coupling with the proton attached to phosphorus. The P-F coupling constant was determined to be 1135 cps. The JFH value of 40 cps is consistent with the value of 4107 cps found in the proton nmr spectrum. The fluorine nmr signal for PHF2 occurs 42.2 ppm upfield from TFAo The phosphorus nmr spectrum (Figure 19) shows the expected 1201 triplet (JpF = 1143 cps) of doublets (JpH = 184 cps) centered 224 ppm downfield from OPAo It is interesting to compare the coupling constants derived from the nmr spectra of difluorophosphine with those of related compounds for which the molecular geometry is known as shown in Table 12o Examination of the Table 12 Comparison of Coupling Constants and Geometry PHF2 PH3 PF3 F2PNMe2 JpF(cps) 1135 1441(51)94(2) Jp-(cps) 182.4 18o (65a) /_ FPF? O 104~(52) 92.5 ( 79) / XPH? 9.5~ 0(78)

- 58 - table reveals that JPH for phosphine and difluorophosphine are almost identical. Also, the JpF value for difluorophosphine is closer to that found for dimethylaminodifluorophosphine than that found for trifluorophosphineo Since PH(80) and PF(5) coupling constants correlate qualitatively with molecular geometry, it would appear that PHF2 is pyramidal with the P-substituent bond angles close to 90~o Stated in terms of the a-bonding hybridization of phosphorus, difluorophosphine apparently exhibits p - hybridization with the lone-pair electrons on phosphorus having a high degree of "s-character"o The vapor pressure data for PHF2, tabulated in Table 14 indicated that it is less volatile than either PH3 or PF3o A comparison of freezing points, boiling points, and Trouton constants for PHF2, PH39 and PF3 as shown in Table 13, further substantiates the difference in volatilityo Since difluorophosphine freezes and boils Table 13 Comparison of Physical Properties Trouton Consto bopo (~C) fopo (0C) (calo/degomole) PFs(47) -101o8 -151 5 23.1 PH3(47) -87 8 -13358 18o8 PHF2 -64 6 -12402 24o7 to -12306

- 59 - Table 14 Physical Constants for PHF2 A. Vapor pressure data. ~C mm (obsvd) mm (calcd) -129.4 2.7 2,8 -112.0 20.6 19.7 -100.9 56.4 55~5 -84.4 206.0 206.7 -78.5 308.2 313.7 B. Equation. -1126 log p(mm) = 2 + 8.280 C. Trouton constant b.p.(extrapolated) 24.70 cal./deg. mole -64.6~C m.p. -124.2 to -123.6~C

- 60 - at higher temperatures than either PH3 or PF3, and likewise, displays a higher entropy of vaporization (Trouton constant), it would appear that difluorophosphine is associated in the condensed states. It is yet to be determined whether association occurs through a'P-F —-H —-P or a,P —-H -P- type interaction. The difference in the proton chemical shifts for liquid PHF2 (-7.65 ppm) and gaseous PHF2(-1.3 ppm)* seems to corroborate the presence of association in the liquid(77). The corresponding difference in the chemical shifts of liquid PH3 (-1.76 ppm) and gaseous PH3 (-1.3 ppm)(78) is even smaller. The larger difference for PHF2 probably points to a stronger association in liquid PHF2 than is present in liquid PHs. This implies that the P-H linkage is less protonic in PH3 than in PHF2. The infrared spectrum of gaseous PHF2 (Figure 20) shows the six fundamentals expected for either the planar (C2 symmetry) or the pyramidal (Cs symmetry) PHF2 moleculeo Examination of the spectrum, however, reveals that several of the absorption bands have definite shapes * While discussing chemical shift of PHF2 vapor, it should be noted that the spectrum was very different from that of the liquid. It was very complex and the exact nature of the spectrum was difficult to determine since the low proton concentration presented difficulties with saturation of the signal by the radiofrequency field. Nevertheless, the center of the multiplet was determined to be 1.3 ppm downfield from TMS (internal), and when cooled and liquified, the sample displayed the simple doublet of triplets observed previously for PHF2(.)

Figure 20 Infrared Spectrum of PHF2(g) W1 II \2600D 220 * 0 1 2 0 0 1000 800 U 600D 4 00G frequency (cm - ) -- -- -- -- - -\ - - ---------------- 2600 2200 1200 1000 800 6.00 400

- 62 - characteristic of vibration-rotation interaction. These band-outlines were used as an aid in making tentative vibrational assignments (Table 15) Since the P-H stretching motion and F-P-F deformation motion are separated from other absorptions which might complicate the interpretation of their band-outlines, their band-outlines were compared with those predicted for the planar or pyramidal geometries for PHF2o The analysis, which is discussed in detail an Appendix B, is consistent with a pyramidal geometry for PHF2 and inconsistent with a planar configuration, thus confirming the same conclusion drawn from the nmr coupling constants observed for liquid PHF2. When compared with the P-H stretching frequencies of phosphine (v PH = 2327, as PH = 2421 cm-1)(83), the s as P-H stretching frequency observed for difluorophosphine (2240o5 cm-') indicated that the P-H bond is probably more labile in difluorophosphine. The infrared spectrum of solid PHF2 was also determined and is displayed in Figure 21. The tentative vibrational assignments for both solid and gaseous PHF2 are given in Table 15. When comparing the infrared spectra of solid and gaseous difluorophosphine, it is interesting to note the pronounced shift of the P-H stretching motion to a higher frequency in the solid, contrary to what would be expected if hydrogen bonding occurred in the solido The lowering of the P-F stretching

- 63 - Table 15 Infrared Spectra* of PHF2 PHF2(g) PHF2(s) freq. (cm-'1) reqeq (cm1) & intensity note assignment & intensity note 2251 s R, br 2240o5 vs Q v PH 2317.0 vs 2233 s P, br 1015.7 vs 6bPH 100805 vs? 980 w 1008 s br as PH 968.0 s as 957.7 vw? 838.3 vs vas PF 825.7 s 851.4 vvs R 794 m sh v PF 825.3 vvs P 782.8 vvs br 367 w R, br 6 FPF 371.3 vs 348 w P, br * For an explanation of notation see Appendix A.

Figure 21 The Infrared Spectrum of PHF2(s) __al' iii v A 2600 2200 1200 1000 800 600 400 frequency (cm l)

65 frequencies in the solid is probably indicative of interaction through the fluorine atoms resulting in a diminution of the inductive effect of fluorine and a concomitant strengthening of the P-H bond so that the P-H stretching frequency approaches that usually observed in phosphines. Some physical properties relating to the N-H and P-H bonds of NHF2 and PHF2 are similar when compared to those of NHs and PH3. The proton nmr chemical shift (5 = ca. -704 ppm)* and N-H stretching (vNH = 3i93 cm1 )( observed for difluoroamine are shifted relative to ammonia 6 -4.8 ppm(8; vN = 3-46, pp8338, aNH = 3414 cm (6) in the same manner as the corresponding values for difluorophosphine are shifted relative to phosphine. It will be interesting to observe how the chemistry of the N-H and P-H bonds in difluoroamine and difluorophosphine parallel one another. IIIo Some Chemistry of the New Difluorophosphineso Ao General. The reactions of only two of the newly prepared difluorophosphines have been studied very extensively - they are difluoroiodophosphine and difluorophosphine. It was suspected that the P-I bond in PF2I would prove labile enough to be cleaved selectively, and * The nmr data for NHF2 are very qualitative; however, from the data Kennedy and Colburn(70) it is estimated that 6 = approximately the shift of benzene = -7.4 ppm (TMS std.)o

- 66 - therefore, PFaI would be a useful and reactive intermediate with which one could introduce the difluorophosphino-moiety into new molecules. As mentioned previously, difluorophosphine indeed has been used for the preparation of a number of interesting new ligands. Since the preparative methods involving PF2I are adequately discussed in previous sections, and since the chemistry of PF2I was investigated only for its synthetic utility, further discussion is not warranted. The reactions of the "parent-difluorophosphine", PHF2, have been investigated in somewhat more detailo Difluorophosphine has been found to undergo reactions which may be classified as 1) Solvolysis, 2) Adduct Formation, and 3) Base Displacement. The latter typereactions are discussed below and some new species are briefly characterized. Bo Solvolysis of Difluorophosphine. 1) Hydrolysis. The hydrolysis of PHF2 in 40% NaOH at 100~ was found to yield nearly one mole of H2 for every mole of PHF2 taken. Initially, it was thought that the equation might be PHF2 + H20 - PF2(OH) + H29 however, P-F bonds are known to be cleaved readily in (35) alkaline solution,5) and no doubt a more likely reaction is PHF2 + 2NaOH + H20 - P(OH)3 + H2 + 2NaFo

- 67 - Although the evolution of hydrogen in basic solution can be characteristic of hydridic M-H linkages as in the silanes(87), it is felt that the P-H bond in difluorophosphine is protonic. If the P-H hydrogen in PHF2 is hydridic9deuterium should be observed in the hydrogen collected after hydrolysis of PDF2o When the basic hydrolysis was conducted with a mixture of PHF2 and PDF2, a mass-spectral analysis of the liberated hydrogen indicated that nodeuterium was present. Thus the hydrolysis of difluorophosphine is felt to proceed with retention of the P-H bond according to the following schemes PHF2 + 2NaOH > HP(OH)2 + 2NaF 1H20 0 HP(OH)2 + H2o The latter scheme demonstrates that the phosphorus in PHF2 is formally in the +1 oxidation state and that the H2 results from the oxidation of HP(OH)2 by water. not from the attack of Ha2 on a hydridic P-H bond~ When excess difluorophosphine is treated with water vapor, a different hydrolysis occurs as indicated by the isolation of SiF4o This hydrolysis might be PHF2 + H20 -> HP(OH)2 2HF, followed by HF + (glass walls of tube) - SiF4o Such a secondary scheme may account for the non-quantitative recovery of H2 from the basic hydrolysis of PHF2o

- 68 - 2) Aminolysiso Since F2PNMe2 and FP(NMe2)2 can be prepared from PF3 and HNMe2 by aminolysis (42) it was felt that analogous reactions with difluorophosphine might proceed as shown 2HPF2 + 3HNMe2 - 2FHPNMe2 + Me2NHo2HF HPF2 + 3HNMe2 - HP(NMe2)2 + Me2NHo2HFo When equal amounts of HNMe2 and PHF2 were mixed as gases, a cloud of white solids did form and an unstable, volatile species which may have been FHPNMe2 was isolated. The infrared of the latter species (Figure 22B) is very similar to that of F2PNMe2 (Figure 22A) except that bands which can be ascribed to P-H stretching and deformation motions are present near 2240 and 930 cm-', respectively. Also, disregarding the peaks at 890 and 860 cm-' which are due to PF3, it can be seen that the peaks in the 900 - 650 cm- region are shifted somewhat relative to F2PNMe2, After removal of the PF3, the remaining fraction was found to have a vapor density corresponding to a molecular weight of 94 go/mole (theory for FHPNMe2 = 95 go/mole) Treatment of PHF2 with larger amounts of dimethylamine (HNMe2/PHF2 = 2o1, 4l1) led to an even different product, perhaps HP(NMe2)2o As can be seen from the IR-spectrum shown in Figure 22C, there is no evidence for a P-F bond in the latter species but P-H stretching and deformation motions are evident around 2250 and 930 cm-. However, the supposed HP(NMe2)2 decomposed fairly rapidly to HNMe2 and brown, oily solids; the highest

- 69 - Figure 22 Infrared Spectra A. F2PNMe2(g) B. FHPNMep(g) 2 I PF3 C. HP(NMe2)2(g) D. HNMe2(g) _,, I,,,,-, 3000 2000 1500 1000 900 800 700 frequency (cm- )

- 70 - observed molecular weight was 96 g./mole (calcd. for HP(NMe2)2 = 120 go/mole), but as can be seen by comparing the IR of HNMe2 (Figure 22D) with that for the molecular weight sample (Figure 22C), some dimethylamine was probably present in the HP(NMe2)2o Since in the case of both FHPNMe2 and HP(NMe2)2 the supposed species readily decomposed to brown, oily, intractable solids and HNMe2S they were not studied further. The liberation of dimethylamine may indicate that either an intra- or intermolecular hydrogen transfer from the phosphours to the dimethylamino-substituent promotes the decomposition. Since Noth and Vetter(26) have found (Me2N)2PHoBH3 to be reasonably stable, perhaps diborane could be used to remove FHPNMe2 and HP(NMe2)2, in the form of their respective borane adducts, from the systems mentioned above. Co PHF2~Lewis-Acid Adduct Formation. 1) Reaction of HI and PHF2o In addition to the desired. product, an unstable species believed to be PF2H~HI was isolated during the preparation of difluorophosphineo The latter hydrogen iodide adduct of PHF2 was characterized by its chemical and physical similarity to an equimolar mixture of HI and PHF2o A mass-balance of products and reactants was consistent with the formula PF2HoHI (see Experimental) When allowed to warm to room temperature the adduct apparently split out HF which reacted with the glass in the vacuum system to form SiFt(

- 71 - PF3, and leave yellow, iodine-containing solids behind; the same behavior was noted for an equimolar mixture of PHF2 and HI. Because of the aforementioned decomposition, it was difficult to obtain a vapor density or to compare the stability of PF2HoHI with PH3~HIo However, the infrared spectrum was obtained on a fast scan with the Perkin-Elmer 137 (Figure 23). The spectrum is markedly different from that of difluorophosphine and appears to confirm the presence of a volatile difluorophosphonium iodide "ion-pair", (PH2F2)I. If the difluorophosphonium ion is assumed to be tetrahedral (C2v symmetry), eight IR-active bands would be expected for this moiety. Only six lines were observed in the spectrum of PF2H'HI, but the PF2 deformation and rocking motions no doubt occur below the range of the instrument used. In addition to those lines assigned to the difluorophosphonium moiety, a number of very weak bands were noted in the 1500 - 1250 cm-1 region. The latter bands may be due to impurities or overtones and combinations, however, this region is characteristic of bridging-hydrogen atom motions which might be observed because of P- e-I type-bonding. Perhaps the latter type of bridge bond would be more evident with a weaker acid like HC1o The observed IRfrequencies of PF2H~HI and their tentative assignments are listed in Table 16. Because of its instability, the mass spectrum of PF2H'HI was difficult to obtain and often only SiF4 and

P F3 I Figure 23 The Infrared Spectrum of PF2H*HI(g) frequency (cml ) 2500 2000 1500 1200 1000 900 800 700 I..1 I -. I I _ I

- 75 - Table 16 Infrared Spectrum* of PF2HoHI freqo (cm-") tentative & intensity note assignment 2551 w vasPH 2483 w vsPH 1450 vw bridge 1375 vw hydrogen 1300 vw motions 1015 vs b HPH s 999 vs coHPH 980 m sh 890 w r PF3 865 m) impurity 835 vvs R 825 vvs Q v PF 815 vvs P 785 w R 768 vs Q as PF 755 w P For an explanation of notion see Appendix A.

- 74 PF3 were observed; however, a number of spectra did show very prominent peaks at m/e = 89, 71, 70, 69, 51, 509 325 31, 20, and 19 which may be assigned to PHF3s PH2F2s PHF2, PF2s PHF, PF, PH, P, HF, and F unipositive ions, respectivelyo The m/e = 89 peak may result from the combination of HF and PF2 but another prominent peak found at m/e - 76 cannot be explained unless perhaps it is due to F,9 which seems unlikely. The important feature of the mass spectrum, nevertheless, is the presence of peaks consistent with the PH2F2 ion. 2) Reaction of B2H6 and PHFo2 A very interesting adduct is formed in 82% yield when diborane and. difluorophosphine are mixed as gases at 25~0 PHF2 + 2-B2H6 - PHF2~BH3 Pure difl-orophosphine borane has been maintained for 22 hr. at 250 and a pressure of ca. 600 mm with no evidence for dissociation or decomposition,in stark contrast to PH3~BH3 and PF3sBH3 which are both highly dissociated under comparable conditions. The stable new borane adduct may also be prepared by displacement of PH3 or PF3 from their respective borane adducts; such base dispacement reactions are discussed completely in Section IV and. in the Experimentalo Although classical methods were not used to obtain an elemental analysis, the formula PHF2 BH3 is supported unequivocally by other physical methods such as'9F, 1B1 and'H nmr spectroscopy and mass spectroscopy.

- 75 - The density of the vapor corresponds well with that expected for the monomer PF2HoBH3 (Table 17). The mass spectrum (Table 17) displays no appreciable peaks at m/evalues higher than 84 and thus confirms the vapor density molecular weight, however, some features should be noted in the fragmentation patterno The presence of peaks at m/e = 71, 52 might be due to PF2H2+ and PFH2+ ions, respectively, which form by recombination of a proton and the appropriate fragmento Such. recombination peaks are rather common in high-pressure mass spectroscopy(8) and have also been observed in conventional mass spectro(89) scopy' Peaks are displayed at m/e = 49, 48 which also might be attributed to recombination to form "1BF2, 1~BF2, unipositive ions~ The peak at m/e = 30 is better assigned to 10BHF+ than to 11BF+since no peak appeared at m/e = 29 ('-BF)o Other assignments appear straightforward but are not listed in Table 17 because of the multitude of isotopic permutations and combinationso The vapor pressure data for difluorophosphine borane are compared with the values calculated from the vapor pressure equation (Table 18); the extrapolated boiling point and Trouton constant are also giveno Since the stability of PF2HoBH3 appeared'ganomolous" when compared to PH3SBH3 and PF3~BH39 it was thought that perhaps it might exhibit an unusual structureo The proton, boron, and fluorine nmr spectra of PHF2aBH3 are what might be termed "textbook examplesn of first-order nmr spectra

- 76 - Table 17 Characterization of Difluorophosphine Borane A. Vapor density molecular weight. obsvd: 83.9 g./mole calcd: 83.8 g./mole B. Mass spectrum. relative relative m/e peak height m/e peak height 84 19.7 49 7.8 83 69.8 48 1.9 82 100.0 44 1.0 81 22.9 43 3.2 80 1.8 42 1.9 71 3 o7 41 0.o 70 13.7 35 3.2 69 46.1 33 28.6 63 1.6 32 17.9 62 3.0 31 63.2 61 9.4 30 14.2 60 2.1 19 o.6 52 1.0 13 31.5 51 36.7 12 23.4 50 25.9

- 77 - Table 18 Physical Constants of PHF2'BH3 A Vapor pressure data. ~C mm (obsvd) mm (calcd) -78.5 5.2 4,9 -63.9 14.3 15o6 -45.6 53.4 54.3 -37o5 88.8 88.6 -36.5 96.4 93.9 -31.9 122.7 121.5 -23.6 191.8 190.4 B. Equation. log p(mm) = -1407 + 7.917 T C. Trouton constant b.p. (extrapolated) 23.04 cal./deg. mole 6.2~C

- 78 - and show unquestionably that difluorophosphine borane is characteristic of the usual BH3 adduct with a normal P-B bond; the nmr spectra provide no evidence for structures such as F H H FP'B/ 9 [F2P][BH.], or [F2PH2][BH2]o F/ /\H The proton nmr spectrum of difluorophosphine borane is indicative of two hydrogen environments in the molecule. The basic spectrum consists of a 1:1:1:1 quartet and a doublet of 1l2:1 triplets centered o078 and 7.68 ppm downfield from TMS, respectively (Figure 24)o The quartet (JBH = 103 cps) is attributed to the borane hydrogen atoms which are split by the L1B nucleus (I = 3) o Each member of the borane quartet exhibits further splitting as shown in Figure 25. The spin-spin splitting pattern in the latter figure can be attributed to coupling of the borane-hydrogens with the two equivalent 19F nuclei (I = ~) to give a 1:2:1 triplet (JFPBH = 26o0 cps) Each member of the triplet is split further into a doublet of doublets by interaction of the borane hydrogens with phosphorus (I = a) and the single phosphine hydrogen; the respective coupling constants are J - 17.5 cps and JHPBH = 4.0 cpsO The doublet of triplets in the proton nmr spectrum of PHF2~BH3 (Figure 25) is assigned to the phosphine hydrogen (JpH = 467 cps), the basic phosphine doublet being split into a 12:1 triplet by the two equivalent

l l Figure 24 1~00 Me 1 NMR Spectrump ~~~~~~of ~~~TMS PHF2.BH3() @ 37~ jFPH = 55.2 cps JpH = 467 cps BH = 103 cps 6(PH) = -7.68 ppm 6(BH) = -0.78 ppm TMS

Figure 25 H NMR of PHF2 BH3(g) ONE MEMBER OF BORANE QUARTET 0 JPBH = 4.0 cps l J I B l 2 JpBH = 17.5 cps JFPBH = 26.0 cps

- 81 - fluorine nuclei (JFPH = 55~2 cps) Careful examination of the 1:201 triplet (Figure 26) shows that each component consists of a 1:3:4:4:4:4:4:453:1 ten-line multiplet. This multiplet can be attributed to coupling of the phosphine hydrogen atom with the 1B nucleus to give a 11oo:1o quartet (JPH = 8.0 cps) each component of which is further split into a 1:353:1 quartet by spin-spin interaction with the three equivalent borane hydrogens (JBH = 4.0 cps); the latter signals overlap to give the observed ten-line multiplet The ratio of borane hydrogen to phosphine hydrogen 1H nmr signals was 3:1 for PHF20BH3. A:353o:1 quartet of doublets is displayed 60o4 ppm upfield from TMB in the 11B nmr spectrum of PHF2~BH3 (Figure 27). The quartet pattern conclusively demonstrates the presence of the BH3 moiety in the moleculeo Other boron hydride fragments might display a 1:1:lo1 quartet in the proton nmr spectrums but only a borane nucleus bonded to three equivalent hydrogens would be expected to display the observed 1:3:1ol quartet (JHB = 102 cps) in the "'B spectrum. The doublet splitting of each member of the quartet is attributed to the directly bonded 3 P nucleus (Jp = 48.6 cps) The basic L9F nmr spectrum consists of a doublet centered 21o5 ppm downfield from TFA (Figure 28). The doublet is indicative of spin-spin interaction of the fluorine nuclei with the directly bonded phosphorus

Figure 26 H NMR Spectrum of PHF2.BH3(I) ONE MEMBER OF PHOSPHINE DOUBLET 1 1,1 " 1 1 11,1 1 1,1 1 1 1 55. J'BP = 8.0 cps HI BPH = 4.0 cps JFPH = 55.2 cps l I~~~~~~~~~~~~~~~~~~~~~~~~~~

Figure 27 32.1 Mc 1B NMR Spectrum of PHF2.BH3(9) @ 370 ~TMB LT] T9 T ~ T -JPB = 48.6 cps = 60. JBH 102 cps 6 = 60.4 ppm

Figure 28 94.1 Mc 19F NMR Spectrum of PHF2BH3(1) @ 370 00 j jPF = 1151 cps [U JHBPF = 26-5 cps TFA HPF = 54.5 cps - p 6 = -21.5 ppm

- 85 - (JFF = 1151 cps) Each member of the doublet consists of a 1o3o4o4o3ol sextet which can be attributed to the overlap of a doublet of 135:3:1 quartets. The latter doublet splitting (JHPF = 54.5 cps) is due to the phosphine hydrogen, and the 1:03o31 quartet to the three equivalent borane hydrogens (JHBPF = 26.5 cps) The coupling constants derived from the proton, boron, and fluorine nmr spectra are compared in Table 19 where it can be seen that they agree within experimental error. Some of these coupling constants are discussed further when the borane adducts of phosphine and trifluorophosphine are compared with difluorophosphine borane (Section IV) Table 19 Coupling Constants for PHF2~BH3'H nmr 11B nmr 19F nmr JH w103 102 - JFPBH 26.0 26.5 JPBH 17.5'HPBH 4 — JPB -- 48.6 - F —- 1151 rPH 467 -- JFpH 55~2 -- 54o5 JBPH 80

- 86 - In summary, the nmr results show the structure of PHF2~BH3 to be F, H (Figure 29) F w P-P -B H, Structure H/ H — of PHF2oBH3 where free rotation exists about the P-B bond and averages out the molecular environments so that all the nmr spectra are easily interpretable with first-order spin-spin coupling rules; this case is similar to that observed by Shoolery(90) for Me2PH-BH3. The structure for PHF2~BH3 (Figure 29) deduced from nmr spectroscopy has Cs symmetry and should exhibit 18 IR-active fundamentals - six vibrational frequencies associated with each end of the molecule, considered as a free species with Cs symmetry, and six arising as a consequence of the P-B bond. Therefore, the PHF2 and BH3 moieties each contribute four symmetric and two asymmetric motions, while three symmetric and three asymmetric motions are due to the presence of a P-B bond; the eighteen fundamental vibrations and their associated symmetries are summarized in Table 20. Although accurate vibrational assignment is difficult without complimentary Raman data and isotopic substitution, under high resolution a number of bands could be associated with boron by the appearance of a higher frequency shoulder arising from the presence of natural abundance'~B; tentative assignments and vibrational frequencies for gaseous PHF2oBH3 are listed in Table 21, and the spectrum is shown in Fig. 30

- 87 - Table 20 Fundamental Vibrations* for PHF2 BH3 Internal Motions Consequence of of PHF2 BH3 P-B bond class vsPH (V5) v BH3 (vl) vPB (V7) class s s v PF2 (v8) v asBH3 (V2) pBH3 (vwo) b6 PH (V) b BH3 (V3) pPF2 (v,,) 5 5 b6FPF (v9) bsHBH (v4) A class asPF2 (V15) vasBH3 (v12) TPB (v'a) class as as _oPH (v14) coBH (v13) cBHs (v17) cPF2 (8is) * For an explanation of notation see Appendix A.

- 88 - Table 21 The Infrared Spectrum* of PHF2 BH3(g) freq. (cm-1) freq. (cm-1) & intensity note assgn. & intensity note assgn. 2475.5 m 1B V2, V12 904.5 vs V8 2473.0 m 10B V1 903.0 vs P V15 2464.2 vs lBt v2, V12 820 vw? 2462.2 vs 11B Vi 732.3 w 1~B Wlo 2440 m R 729.7 w 10 B 17 2423.7 vs Q 1, 2422. vs Q 5 728.2 w? 2406 m p 723.2 s 11B vlo 1120 w V713 720.8 s 11B 7V17 1035 m R /6 1031 w V4 577 w 10B 7 1026.9 m 10B V3 567 w 11B 7 1023.2 vs Qt 76 389.2 m vl6 1021.8 m 11B V3 3579.8 m 1013 m P 374 m 9 374.7 m 978.5 w 714 230 w vs1 9211o vs R v15 225 w 71l 912.8 vvs Q V15 * For an explanation of notation see Appendix A and Table 20. t Appeared to be split under high resolution.

Figure 30 The Infrared Spectrum 1of o PHF2.BH3(g) frequency(cm~ ) 26002200 1200 1000800600I400 2 6 00 2 200 12~00 1 0 00 8010 60O'0 400

- 90 - Do Base Displacement by Difluorophosphine. 1) Displacement of CO from Ni(CO)40 This preliminary study demonstrated that PHF2 does displace CO from Ni(CO)4 to form (PHF2)Ni(CO)3o PHF2 + Ni(CO)4 - (PHF2)Ni(C0)3 + CO However, the reaction of nickel tetracarbonyl and difluorophosphine is complicated by side reactions which yield brown, oily, non-volatile substances. After 3 days reaction time at 0~ when the initial PHF2/Ni(CO)4 ratio was 5:1, an amount of gas equaling the amount of PHF2 consumed was isolated; after purification by fractional condensation the fraction displayed the infrared spectrum shown in Figure 31. The following discussion shows that the spectrum is consistent with the formula (PHF2)Ni(CO)3. Because of the small amount of sample isolated further characterization was not possible If PHF2 is assumed to be a point-mass, L. and nickel is considered to be tetrahedral, then the various possible difluorophosphine-substituted nickel carbonyls and their respective symmetries can be deduced. as shown in Table 23, which also gives the number of C-0 stretching frequencies expected for each degree of substitution. From a comparison of Table 23 with the spectrum (Table 22 and Figure 31), it is immediately apparent that the new species cannot be the tris- or tetrakis-difluorophosphine derivative since the C-O stretching region is too complexo Still the four

Figure 31 The Infrared Spectrum H Of (PHF2)Ni(CO)3(g) I I I 1 2 400 2nnn/\_,________,________,_________frequency (cm 6)___I 24___?n'1200 __ oo 10oo ____600 __ 400___

- 92 - Table 22 The Infrared Spectrum* of (PHF2)Ni(CO)3 freq. (cm-') tentative & intensity assignment 2331.8 vs vPH 2110.6 m vCO 2087.2 vs vCO 2052.0 vvs vCO 2016 w 2-0PH 1035.9 m overtone or 1030 m combination 1019.9 vs bPH 982 w cPH 853.3 vvs vPF 836.0 vvs vPF 457 m 6NiCO 443 m b6NCO 421 w coNiC0 408 s bPF2 * For an explanation of notation see Appendix A.

- 93 - Table 23 Symmetry and IR-active C-0 Stretching for LxNi(CO)4-x, x = 14 C3v C2v 3v Td L L 0 L C ~I ~ \ L I No -co N N.co L L N 00o'\c c. L \ IR- two two one active anone vCO1s a2 & e a1 & b2 a. bands in this region are more than expected for either of the two remaining possibilities; however, if L is not symmetric about the C3 axis as assumed above in Table 23, and as is the case for PHF2, the degeneracy of the asymmetric C-0 stretch in LNi(CO)3 would be removed and a total of three C-0 stretches might be expected~ The fourth absorption in the C-0 region may be due to an overtone enhanced in intensity by Fermi-resonance(66) The infrared provides further evidence for the monosubstituted species when the P-H stretching region is also considered. The compound (PHF2)2Ni(C0)2 would be expected to show symmetric and asymmetric P-H stretching frequencieso Even under high resolution there was evidence for only one absorption in the P-H region, just as expected for (PHF2)Ni(CO)3; tentative assignments for the infrared spectrum are listed in Table 220

- 94 - When equal amounts of difluorophosphine and nickel tetracarbonyl were mixed at 250, the displacement of CO appeared to be more extensive as evidenced by the infrared spectra of the reaction products which displayed considerable complexity in the P-H and. C-0 stretching regionso However, before characterization of the products could be completed the sample was inadvertantly losto Nevertheless, the experiments mentioned above show that difluorophosphine displaces CO from nickel tetracarbonyl under much milder conditions than those reported for PF3 (22b) to effect the corresponding displacement(22b) 2) Displacement of PH3 and PF3o Experiments were conducted in which difluorophosphine was found to displace phosphine and trifluorophosphine from their respective borane adducts, nearly quantitatively~ Typically, after about a day at -78~ and an additional day at -45~, difluorophosphine displaced PF3 from PF3 BH3 according to the equation PF3~BH3 + PHF2 - PHF2~BH3 + PF3s With an excess of trifluorophosphine borane the PHF2 was completely consumedo In a similar experiment involving an excess of PH3~BH3, after - day at 0~ the difluorophosphine reacted completely, displacing an equal amount of PH3 as shown in the equation PH3~BH3 + PHF2 -> PHF2~BH3 + PH3o The latter two base displacements were effected with

- 95 - a stoichiometric deficiency of PHF2; yet the difluorophosphine was completely consumedo If the original adducts formed stronger coordinate P-B bonds than PHF2, under these conditions unreacted difluorophosphine would be present; therefore, PHF2 is a stronger base than either PH3 or PF3. The relative base strengths of PH3 and PFs towards borane are established by experiments and observations discussed in the next section. IVo Relative Base Strengths of PH3s PHF2, and PF3 Towards BH3. A. General. The relative base strengths of phosphine, difluorophosphine, and trifluorophosphine toward borane would best be measured by the enthalpy change, AH, for the dissociation of the adducts in the gas phase into the component donor and acceptor molecules. L~BH3 L + 2 B2H6; L = PH39 PHF23 PF3 2102) Such thermodynamic data have been obtained for PFs3BH3(10 and conceivably could be for PHF2-BH3, if it dissociates at higher temperatures without decomposition; however, PH3~BH3 is essentially nonexistent in the vapor, and such a study is precluded in this case. Nevertheless, both base-displacement reactions and the equilibrium constant for adduct dissociation, K = - [LoBHs-] have been used eviously as a measue of atve base sengh The previously as a measure of relative base strengtho The

- 96 - present study qualitatively relates the results to the equilibrium involved in adduct formation and finds that the relative base strengths towards borane is PHF2 > PF3 > PH3. So that apparent anomalies in adduct stability cannot be attributed to unusual structural differences, a necessary prerequisite for a study of the relative stability of the coordinate link in a series of Lewis acid-base adducts is the assurance that the adducts be essentially "isostructuralK o Therefore, before the evidence establishing the relative order of coordinating ability for PF3, PHFa2 and PH3 is discussed, the structural similarity of their respective boranes will be shown by nmr spectroscopyo Bo Structures of the PHF2, PH3g and PF3 Borane Adducts. 1) Difluorophosphine Borane. An examination of the nmr spectra of PHF2~BH3 as presented in Section III C-2 clearly indicates that the structure is H H F —-P —B o FO \H This structure is consistent with the usual borane adduct, and, as will be seen in the following discussions it is basically no different from the known structures of (91) (92) PH3sBHs39 and PF3sBH3 92 2) Phosphine Borane. In 1940 Gamble and Gilmont prepared "diborane diphosphineE by mixing diborane and phosphine at low temperatures(90) On the basis of

97 - rudimentary chemical evidence an analogy was drawn to the diammoniate of diborane and a structure analogous to that accepted at that time for B2H6~2NH3 was proposed, (PH4) (H3BPH2BH3) o No molecular weight data were available to support the postulated ionic dimero Some of the reactions of (H2PBH3)n can be interpreted best in terms of a single monomeric formulationo For example, trimethylamine displaces PH3 quantitatively to (93) give H3BN(CH3)3(9 and a kinetic study of the reaction of B2H6 and PH3 by Brumberger and Marcus(94) suggested the monomeric representationo Since the original structural postulates were presented, a new model for B2H6~2NH3 has been accepted(95) in place of the earlier ammonium type of solid, (NH4 +)(HsBNH2BH3-) but neither the new diammoniate model nor the monomeric model for (H3BPH3)n rationalized the known reaction with ammonia as easily as did the early dimeric formulation proposed by Gamble and Gilmont0 NH3 + (PH4+)(H3BPH2BH3-) -. (NH4+)(HsBPH2BH3-) + PH3o Recently, an x-ray crystallographic study by McGandy(91) showed that (H3PBH3)n is indeed monomeric in the solido Nevertheless, phosphine borane was investigated further by nmr, infrared, and Raman spectroscopyo The LLB nmr spectrum of molten HsBPH3 at 370 (Figure 32) consists of a 13o3 1 quartet of doublets centered 6008 ppm upfield from TMBo The three magnetically equivalent protons split the boron signal into a quartet

Figure 32 32.1 Mc lB NMR Spectrum of PH3.BH3() @ 370 III' I TMB 5 = 60.8 ppm U -_____.A___________________I JP B = 27 cps ~~'~/~' ~~~~Jg = 103 cps

99 (JBH = 103 cps); then each member of the quartet is split into a doublet by coupling with the phosphorus (I = 1) (JBp = 27 cps). While the 1:53031 quartet of doublets would be expected for both models, H3BPH3 and PH4(H3BPH2BH3)s the coupling constant, J = 103 cps, compares well with BH. the values (JB = 100 cps) obtained by Gilje(96) for CH3PH2BH3 and the values (JBH = 96 cps, a = 55.6 ppm upfield from TMB) obtained by Phillips, Miller and Muetterties(97) for HP(CH3)2BH3o The LH nmr spectrum of H3PBH3(I) at 370 (Figure 33) shows only two clear groups of hydrogen signals indicating the presence of two hydrogen atom environments in the molecule. A 1:1olll quartet is found centered 0.53 ppm downfield from TMS while a doublet is displayed 4.31 ppm downfield from TMSo The quartet at 0.53 ppm is attributed to the borane hydrogen atoms split into a Il1:1:l quartet by the "lB (I = ) o The value of JBH (104 cps) is in good agreement with the value derived from the "B spectrum (103 cps) and with that found in PH2Me-BH3 (100 cps) The doublet at 4o31 ppm due to coupling of phosphorus with hydrogen shows a P-H coupling constant of 372 cps which checks well with the value of 370 cps reported by Gilje(96) for H2CH3PBHao Careful examination of each component of the phosphine doublet reveals a 1:53ol quartet (JHPBH = 8 cps) which would be expected from coupling of the three borane hydrogens with the hydrogens attached to phosphine. Coupling with the borane

Figure 33 60 Mc H NMR Spectrum of PH3-BH3(g) @ 370 H 0 iu~ ------------------— L J HPBH= 8 cps ui LLUL JPH 16 cps 3PH = 372 cps BH 16 6(BH) =-0.53 ppm BH = 104 cps 6 (PH) =-4.31 ppm TMS

- 101 - hydrogens rather than with the boron (I = 3) is indicated by the 135o3ol line intensities in the quartet. Boron coupling would give 1o1ol1 line intensitieso Such H-P-B-H coupling was also observed by Shoolery(90) in an earlier nmr study of Me2HPBH3, where the coupling constant JHPBH was listed as 6 cps. The spin-spin splitting pattern of hydrogens attached to phosphorus is consistent with the H3BPH3 model but is not consistent with the formula (PH4+)(H3BPH2BH3) o In the latter representation each member of the P-H doublet should be split into a septet of intensities 16. 6150300156olo Even if the outer member of the septet were obscured by low signal intensity, the pattern would be clearly different from the 1:3e3:1 pattern observedo Under higher resolution each member of the borane quartet was seen to consist of a 103l404[31l six-line multiplet as would be expected for two overlapping 1i030ol quartets. One can identify a doublet arising from spin-spin coupling between the borane hydrogens and the phosphorus nucleus of H3BPH3 (JpBH = 16 cps) with each member of the doublet split into overlapping 1l3035o quartets by the three phosphine hydrogens (JHPBH = 8cps) This spin-spin splitting pattern is completely consistent with the formula H3BPH3. It is not consistent with the formula (PH4 )(H3BPH2BH3 ) where each component of the P-B-H doublet would be split into a 1o 21 triplet instead of the quartet by two hydrogens(rather than three)

- 102 - attached to the phosphorus. Finally, peak intensities in the proton spectra are in good. agreement with predictions of the formula H3BPH30 Boron in the sample is about 80% 11B and 20% 10Bo The 1B with nuclear spin of 2 gives rise to the quartets observed. aboveo The 10B with a nuclear spin of 3 gives rise to a septet which is not detectable above the background and which is not included in the 11B quartet intensities. Phosphorus with only a single magnetic nucleus does not split the P-H signal beyond the doublet. Thus the compound H3BPH3 should show a ratio of Hboron to Hhosphorus of 008 to 1.0 or o.80o The observed ratio in the spectrum is 0.8. The foregoing spectra indicate conclusively that molten (H3BPH3) is the simple acid-base monomer, HaBPH3, not the compound (PH4+) (H3BPH2BH-) The x-ray study of McGandy(9l) first and conclusively showed, that phosphine borane is monomeric in the solid state at 25~o The infrared and Raman data presented here (Figures 34-36 and Table 24) are also consistent with the monomeric structure for PH3~BH3(s) at -l80~o The monomer should display an ethane type configuration with C3v symmetry for which eleven infrared active motions would be expected - four vibrational frequencies associated with each of the -MH3 moieties, considered as free molecules with C3v symmetry, and three arising as a consequence of the P-B bond; a fourth skeletal frequency, the

Fiqure 34 The Infrared Spectrum of PH3.BH3(s) frequency (cm- )..I. M I.I I I',I 2600 2200 1 200 1 000 800 600 400

- 104 - Fig.ure 35 The Raman Spectrum of PH3.BH3(s) 400 600 800 1000 1200 1500 2000 2400 * Hg line frequency(cm- )

- 105 - Figure 36 The Raman Spectra Of PH3* BH3(i:) 11 polarized 1 polarized * Hg line freQuency(cm ) 0 200 400 600 1000 1500 2000 2400 a I,. a. a.. I * * a

Table 24 Vibrational Spectra* of PH3sBH3 Infrared Raman freq. (cm-') assignment freq. (cm1) freq. (cm-1), intens. & intensity note & symmetry & intensity & polarization (solid -1O50) (solid -80~) (liquid 35) 2421.7 2426.0 s split in crystal v9, vPH, e 2416 w, sh t 2399.1 s V7, vBH, e t t 2392.0 vs vl, vBH, a, 2401 vs 2399 vs, p 2361.5 s V3, vPH, ai 2363 s 2358 s, p 2272.8 w 2 V5 2263.8 w combination t 2262 w, p 1141.1 m v5s, PH3, al 1144 w 1135 m 1103.1 o108 ~w split in crystal vll, PH3, e 1102 m 1102 m, dp 1077.5 w 10 1068.9 m "B v2, 6BH3, a: 988.2 s ~B 968.7 vs "lB v8, bBH3, e t 989 m, dp 829.5 w split in crystal, 825.6 w overlap of 11B vio, pBH3, e 817 vw 822o5 w and IoB motions. 576.6 vw 10B 563.7 w 1"B V4, vPB, al 572 m 551 m, p 447.1 w V12, pPHs3 e 451 w 436 m * For an explanation of notation see Appendix Ao 8 Very weak band observed, but determination of position is difficult. t Absorption corresponding to this assignment probably masked by nearby peaks.

- 107 - P-B torsional motion of a2 symmetry, is inactive. The Raman effect for this model should also display the same eleven fundamentals, If the structures 1) [PH4][H2P(BH3)2], as originally proposed by Gamble and Gilmont(93), or 2) [H2B(PH3)2][BH4]I analogous to the structure of the "diammoniate of diboranel' were present in the solid, the infrared and Raman spectra of both these structures should differ significantly from those of the C3 monomero If the -BH3 groups in 1) and the -PH3 groups in 2) are considered as point masses and all angles assumed tetrahedral, each molecule contains ions of Td and C symmetry on the basis of which skeletal frequencies can be predicted. The Td moieties should display two IR-active motions and four fundamentals in the Raman effect. Likewise, the C2v moieties would be expected to give rise to eight infrared frequencies and nine Raman. Internal motions of the -MH3 groups would increase the complexity of the observed spectrum in the regions of hydrogen frequencies. Thus, it is apparent that the infrared and Raman spectra of either 1) or 2) should differ significantly. In particular, the spectra should display complexity in areas characteristic of P-B and -MH3 motions. The infrared and Raman spectra of the solid have been found to be similar, relatively simple, consistent with the C3v model, and easily related to the Raman spectrum of the liquido Accordingly, it is felt that this

- 108evidence complements the crystallographic study(9) which showed unequivocally that solid."diborane diphosphine" is monomerico The observed infrared and Raman frequencies are listed in Table 24. No evidence for (PH4 ) or (BH4) is obtained upon comparison of these spectrao Only one P-B motion could be assigned and the spectra are relatively simpleo 3) Trifluorophosphine Borane. Taylor and Bissot(98) found that the Raman effect for liquid PF3~BH3 at -80~ was consistent with the ethane configuration for the molecule (C vsymmetry) A recent microwave study(9) has established the molecular parameters of trifluorophosphine borane. The nmr spectra of PF3oBH3 were determined in the present study and also clearly indicate an ethanetype configuration for PF3SBH3 in which free rotation about the P-B bond exceeds the nmr time constant, averaging out the magnetic environments of the fluorine and hydrogen nuclei. The boron nmr spectrum of liquid PF3~BH3 at 37~ (Figure 37) gives definite evidence for a BH3 group directly bonded to phosphorus. It consists of a 103o3:1 quartet of doublets centered 66.6 ppm upfield from TMBo The basic quartet (JBH = 106 cps) affirms the direct bonding of three hydrogens to the boron, while the doublet splitting (JpB = 39 cps) of each member of the 1o303.1 quartet is due to the directly bonded phosphorus nucleuso A close examination of each peak reveals incipient fine

Figure 37 32.1 Mc 1BNNMR Spectrum of PF3.BH3(g) @ 37 TMB ~I__________________ // a JFPB = 6 cps JpB = 39 cps 6 = 66.6 ppm J=H = 106 cps I111 Illll~~~~~~B

Figure 38 60 Mc 1H NMR Spectrum of PF3.BH3() @ -40~ ~: 36 cps J1OBH 36cpI I I I I 1 I I I IIIJ~ 1I I I I JI = 108 cps Y FPBH BH TMpBH = 18 cps TMS 6= 0 ppm

- 112 - visible members of the septet, the central and outermost members. The spacing between the nuclear energy levels in a magnetic field is directly proportional to the magnetogyric ratio, 7i, and the perturbing magnetic field(8 ) in the case of spin-spin splitting the magnetic field within the molecule itself and the 7i of the perturbing nucleus. Thus, since the only variable altered upon substituting 10B for 11B is 7i, the J(llB):J('1B) ratio should be equal to the ratio 7(l"B):7(l'B) as is the case.* J(lB)/J(l=B) = = 3.000 and y(B) = [1 / = 2.6880 1.8006 2 986 (10B) l, 1 / T 10B 3/2 3 The ratio of the 11B-H signal to the 10B-H signal is also in good agreement with the ratio calculated from the natural abundances of "1B and'~B. The observation of 1OB-H spin-spin coupling is usually precluded by the nuclear-quadrupole moment of the 10B nucleus which can interact with fluctuating electric-field gradients to provide a rapid spin relaxation mechanism and consequent broadening of the nmr signal.t The sharpening of a signal normally expected to * The expressions were evaluated using the values given in reference 81. t The uncertainty principle AEAt h n can be used to estimate the order of magnitude of broadening. Since AE = hAv, the uncertainity principle implies that the uncertainity in the frequency of absorption is l/(27At) o Thus the line width measured on a frequency scale, owing to spin relaxation, will be inversely proportional to the relaxation time (8l)

- 113 - be broadened by quadrupole interaction can be attributed to the lack of a field gradient at the nucleus, thus inhibiting the relaxation mechanism as in the case of the (99) (97) highly symmetrical ammonium( or borohydride( ions. Therefore, the observation of 10B-H spin-spin coupling in PF3~BH3 could also be attributed to the lack of an appreciable field gradient at the boron nucleus, or stated differently, a highly symmetrical electron distribution. A basic doublet attributed to P-F spin-spin coupling (JPF = 1406 cps) is observed centered 19o9 ppm downfield from TFA in the'1F nmr spectrum of PF3~BH3 (Figure 39) If the values of JFP and JF derived from the proton FPBH FPB and boron nmr spectra (18 and 6 cps, respectively) are used to determine the expected fine structure pattern for each member of the doublet, a 13-line multiplet with intensity ratios of 1:1:1:40353o6o3.3o4o1:1~ 1 and spacings of 6 cps is predicted as shown in Figure 40. Although the expected 13-line. multiplet could not be resolved, a pattern which shows the prominent features of the expected multiplet was observed, ioe., a strong central maximum and progressively weaker peaks displaced approximately 18 and 36 cps on either side of the center on the outer peaks (Figure 39) Although the vibrational spectra of trifluorophosphine borane have been investigated rather extensively by Taylor and Bissot(98), and Wyma(100), the spectrum of PFs3BH3 vapor has not been reported in detail(102); intermolecular

high field member of doublet H TFA JpF = 1406 cps A i 6 = -19.9 ppm Figure 39'' 36 cps/div 94.1 Mc 19F NMR Spectrum of PF-BH (9) @ -20~ _3' 3

- 115 - Figure 40 Expected 19F nmr Pattern for Each Member of P-F Doublet in PF3oBH3 spacing = 6 cps 1.0101.043535603553o4ollol overlap JFPB =6 cps |:Lolol quartets JFPBH = 18 cps I J l 1:353:1 quartet forces such as liquid association or crystalline-lattice forces would not perturb the valence force field in the vapor state. In obtaining the infrared spectrum of PF3oBH3 vapor during this study, dissociation was minimized by maintaining the vapor in equilibrium with the liquid which was maintained at about -126~o The spectrum of PF3oBH3 vapor (Figure 41 and Table 25) was easily discernable from those spectra due to the small amounts of PFs and B2H6 which did form. As shown in Table 25, which compares the gas-phase infrared spectrum to the Raman spectrum of PF3~BH3(Q) at -80~ determined by Taylor and Bissot(9 ) there are considerable shifts in some of the B-H

! -7 B2HPPF 3 The Infrared Spectrum | of PF3 PF3'BH3(g) frequency (cm-1) 0 2 i I 0 6 4I l 2600 2200 1200 1000 800 600 400

- 117 - Table 25 Vibrational Spectra* of PF30BH3 Infrared (g) Ramant (e) freq. (cm-') assignment freq. (cm-1), intens. & intensity note & symmetry & polarization 2464.5 m vBH,e 1 oB 2462.'7 m vBH, a 2453.1 s vBH,e 2455 vs, dp 2451.3 s B vBH, ai 2585 vs, p 1131.7 w R 1116.6 w Q 6BH3,e 1117 s, dp 1105.4 w P bBH3, a 1077 w, p 958.9 vs R 9513. vvs Q vPF,e 957 m, dp 943.4 vvs R,P 9313. vvs Q vPF, a, 944 m, p? 922 vs,sh P 693 vw pBH3,e 697 vw 619.1 w 10PB B 607 s, p 609.5 m 1B 452.2 w R 448.4 w Q? 441.8 m Q 0PF3,al 441 m, p 433.0 w P 370.2 wbr bPF3,e 370 vw pPFs,e 197 m, dp * For an explanation of notation see Appendix A. t The spectrum of Taylor and Bissot 98), with the exception of many overtone and combination bands.

- 118 - vibrational frequencies. The PF3 rocking frequency fell below the range of the Beckman IR-12, and the symmetric BH3 deformation assigned to the absorption at 1077 cm in Raman could not be observed in the infrared of the gas - perhaps it has shifted to higher frequency and is one of the lines observed around 1120 cm-o. It also appears that the symmetric PF3 deformation Q-branch at 441.8 cm- in the infrared (441 cm-' in the Raman) shows another Q-branch, probably due to a 10B isotopic shift, at 448.4 cm-'o Co Phosphine Borane and Trifluorophosphine Borane Systems. Difluorophosphine is a markedly stronger base towards BH3 than either PF3 or PH3 as demonstrated by the base displacement reactions described in Section III D-20 In order to determine whether phosphine or trifluorophosphine is a stronger base towards borane, the base displacement reactions described below were investigated. 1) Phosphine and Trifluorophosphine Borane. When PF3~BH3 was treated with twofold excess of PH3 at -35~ to 00, after about one day 87% of the PF3 had been displaced according to the equation PH3 + PF3~BH3 -> PF3 + PH3~BH3o However, other experiments showed that under nearly identical conditions when equal amounts of PF3oBH3, PH3 and PF3 were taken, the displacement only proceeded 48% to completiono Thus, there appears to be an equilibrium

- 119 - situation among the following processes. PF3~BH3(g) = PF3(g) + 2B2H6(g) 2 PH3(g) + -B2H6(g) P PH3~BH3(s) From the various experiments a value was obtained for the equilibrium constant for the displacement (p in atm.) K(O0) = -— [p(PF3).- 107 to 2.2xl0-5 [p(PH3)][p(PF3 ~BH3) ] The fact that PF3 is displaced to any appreciable extent is felt to be a consequence of an unusually high lattice energy for PH3oBH3(s) which drives the equilibrium in the favor of its formation. Therefore, the displacement reaction is not a true measure of the P-B bond strength in the above case. Evidence which indicated the weakness of the P-B coordinate link in PH3-BH3 is cited belowo 2) Observations Concerning PH3oBH3 and PF3BSH3 Gamble and Gilmont(93) did not mention the volatility of (101) PH3~BH39 however, Schlesinger and Burg noted that the compound could be sublimed almost undecomposed in a high-vacuum; vapor density measurements are precluded by the rapid decomposition of P.H3oBH3 vapor. The present study confirmed the latter observations and found that when the sublimate condenses into a cold tubes it does so before mercury, ie.o, the black mercury ring appears in regions of the tube which are colder than those where the white ring of PH3-BH3 forms, implying that the vapor pressure of PH3oBH3 at 25~ is lower than that of mercury. The low vapor pressure of PH3aBH3 is also indicated by

- 120 - the fact that the vapors comprising the dissociation pressure of PH3~BH3(s) are entirely composed of PH3 and B2H63 at least with the usual means of detection such as vapor density and infrared spectroscopy. Therefore, PH3~BH3 apparently exists for appreciable lengths of time only in the molten (viz. nmr spectra) and solid states. This leads to the postulate that the P-B bond is very weak and that the driving force leading to the formation of PH3-BH3 in the condensed states is an unusually large "lattice energy:o Perhaps this "lattice energy" is due to (H ) —-(H-) "hydrogen-bonding" between the acidic phosphine hydrogens and the hydridic borane hydrogens. (93) In the original study (9 of PH3'BH3 the data shown in Table 26 were given for the equilibrium dissociation pressures of the solid. The dissociation occurs according to the equation PH3~BH3(s) = PH3(g) + 2 B2H6(g)9 Table 26 Equilibrium Data for PH3 BH3 Dissociation ~C pressure93) (mm) K* 25 1400 o96200 0 200 o05190 -11 71 o01100 -21 11.00067 * p in atmo

- 121 - and therefore, the equilibrium constant (p in atm.) K = [p(PH3)][p(B2H6)]2 for the dissociation can be obtained by using the law of partial pressures to solve for the equilibrium pressures of PH3 and B2H6 [viz. P(B2H6) = - p(total), etc. ] The equilibrium constants calculated from the data of Gamble and Gilmont(93) are tabulated in Table 26 and log K is plotted vs. 1/~K in Figure 42. The slope of this plot corresponds to a AH of approximately 18 kcal/mole for the dissociation* PH3~BH3(s) PH3(g) + o-B2H6(g) Recently Burg and Fu(102) reported the heat of dissociation for PF3 BH3, PF3~BH3(g) = PF3(g) + -B2H6(g)9 to be AH = 11087 kcal/mole. In order to compare the latter value with that for PH3~BH3(s) as a measure of the P-B bond strength, it is necessary to subtract the heat of sublimation for PH3~BH3(s) = PH3 BH3(g) from the heat of dissociation for PH3~BH3. Although the * The data of Gamble and Gilmont (93are felt to be accurate since the value for the equilibrium constant at 0~ given by Brumberger and Marcus(9) (lo14x106 mm3 or 0.05090 atmo2) corresponds well with the value derived from the 3 data of G & Go (0.05190 atmo.2) Since there is only a small amount of data, however, the calculation is indeed rather crude but does give an estimate of the energetics involved in P-B bond formation for PH3~BH3o

- 122 - 10~. -1 0 10-' -2. 10-3 3.3 3.5 3.7 3.9 103/O K Figure 42 Plot of log Kdissoc vs. 1/ K for PH~3.BH3(s) — -PH3(g) + 1/2 B2H6(g)

- 123 - heat of sublimation for phosphine borane is not known*, it is expected to be substantially greater than the 6.5 kcal/mole difference between the AH of dissociation for PF3 BH3(g) and PH3.BH3(s). In other words, compensating for the heat of sublimation of PH3sBH3(s), the thermodynamic data show that the P-B link in PF3oBH3 is stronger than in PH3oBH3. The observations mentioned above, as well as the existence of PF3 BH3 in the vapor and the non-existence of PH3'BH3 in the vapor, imply that the coordinate bond in PH3oBH3 is weaker than that in PF3~BH3; this conclusion, considered along with the displacement of PF3 and PH3 from their respective boranes by PHF2, to give gaseous PHF2'BH3, shows that the relative base strength toward borane decreases in the order PHF2 > PF3 > PH3; a consideration of the equilibrium involved in adduct formation leads to the same order as is indicated below. Since no diborane of difluorophosphine can be detected in equilibrium with PHF2'BH3 at 250, it can be said that the dissociation constant (pressure in atm.) 1 K (PHF2'BH3) = [P(PHF2)][P(B2H6)2 << [p(PHF2 BH3) ] another good indication of a strong P-B bond. ~ Gmelin1 ~ does list a value for the heat of sublimation of PH3oBH3 (15 kcal/mole), however, a personal communication with Dr. R. A. Marcus, the reference cited in Gmelin, indicated that he had not determined the value nor knew of any such determination.

- 124 - The dissociation constant (pressure in atm.) 1 K (PF3'BH3) = [p(PF3) ][p(B2H6) ] [p(PF3'BH3)] was estimated earlier to be nominally 1 at 25~0 Estimation of this constant during the course of this study yielded comparable values which Burg and Fu(102) recently confirmed. Although it was not possible to observe the dissociation constant of PH3-BH3 directly, it is estimated that K (PH3 BH3) = [P(PH3)][p(B2H6) ]2 > [p(PH3~BH3)] at 25~o Observations show that the vapor pressure of PH3'BH3 is very low at 250, approximately that of mercury (~10-3 mm); since the corresponding dissociation pressure of PH3~BH3 is rather large (1400 mm)(9), a very large dissociation constant is indicated. The coordinate bond in PH3oBH3 must be very weak. The qualitative arguments advanced above show that the dissociation constants increase in the order K (PHF2-BH3) < K (PFaoBH3) < K (PH3-BH3) which corresponds to an ordering of donor ability towards borane of PHF2 > PF3 > PH3s Do Conclusions. The previous discussion has established the relative base strength of PHF2, PF3, and PH3 towards BH3. It follows that one would expect the P-B stretching frequencies in these borane adducts to increase directly

- 125 - with the stability of the P-B bond; however, since the vibrational assignments for these molecules are still tentative, and since the extent of coupling would be undetermined without a complete normal-coordinate analysis, no discussion along these lines will be attempted hereo* An examination of the nmr data (Table 27) obtained for the series H3BoPHxF3_x (x = 0,1,3) shows a striking correlation between JpB and the strength of the P-B bond established here [JpB(cps): 49 (PHF2aBH3) > 39 (PF3SBH3) > 27 (PH3~BH3)]. A trend in JpB is also evident for the series PHxMe3-xoBH3 (x = 0,1-3) [JpB(cps)~ 64 (PMe3oBH3) (10l) 50 HMe9 (PHMe3)97 > 4 (PH BH3)(105) > 27 (PH3~BH3)]; this trend follows the generally accepted order of adduct stability~ However, the correlation of JPB with P-B bond strength does not necessarily hold from one acid series to another since JpB(cps) = 174 (PMe3~BF3)(104) > 64 (PMe3oBH3) (10 ) where Graham and Stone(7) have shown by displacement that with PMe3, borane forms a stronger P-B coordinate link than boron trifluoride. Intuitively, it would seem that such a convenient'handle" as the magnitude of the P-B coupling constant should be related to bond order and thus the strength of the coordinate linko However, the terms which contribute to the observed * A vibrational analysis of PF3~BH3 has nearly been completed by Dro Ro Co Taylor, University of Michigan. The results give a P-B force constant of about 2o38x105 dynes/cm for PF3~BH3. A simple diatomic molecule approximation of the P-B force constant in PH3~BH3 indicates that it is approximately l.8x105 dynes/cmo

- 126 - Table 27 NMR Data* for PF3 BH3 PHF2 BH3 PH3 ~ BH3 (cps (cps) (cps) JBH 107 103 104 JpF 1406 1151 pH -- 467 372 PH JpB 39 49 27 4 8 JHBPH JFPBH 18 26 JFPBH Jpp~T 1655 u FPH JpH 18 17 16 PBH JBPF 6 1H nmr (TMS std) (ppm) (ppm) (ppm) 6(BH) 0 -0.78 -0.53 6(PH) — 7.68 -4.31 9F nmr (TFA std) b(PF) -19.9 -21.5 "lB nmr (TMB std) (:B) 66.6 60.4 60.8 * For an explanation of notation see Appendix A.

- 127 - value are manifold, can vary in sign, and involve the entire electronic configuration of the molecule(81); therefore, it is still to be determined whether this correspondence of JpB to P-B bond strength can be justified on other than empirical baseso Similar reasoning leads to the conclusion that no meaningful correlations can be forwarded for the coupling constants which operate over a number of bonds; however, JpF' JpH' and JBH have been related by some investigators to molecular sterochemistry (or orbital hybridization) As also observed by Heitsch(101) in a more extensive nmr study of boron-containing acid-base adducts, there appears to be no correlation between the 11B chemical shift and adduct stability (Table 27). Phosphorus chemical shifts on the other hand have provided some interesting correlations which will be discussed in the next section in relationship to the nature of the P-B coordinate linko Vo The Nature of the P-B Bondo Since phosphine and trifluorophosphine form borane adducts which are highly dissociated at 25~, the stability of difluorophosphine borane appears anomolous At least a two parameter model is necessary to rationalize the presence of the unexpected maximum in base strength at PHF2 rather than the expected monotonous increase in base strength upon progression from PH3 to PHF2 to PF3o The

- 128 - qualitative arguments which might expain the anomolous base strength of PHF2 are legion and unfortunately appear to be just beyond the realm of predictive utility. It must be emphasized that more tenable predictions and theories will result only after more quantitative work has been completed in this field. Nevertheless, some models will be discussed. One model which might be used to explain the unusual base strength of PHF2 towards borane might be presented as follows. Some authors(7'911l17-19) maintain that in compounds of borane with phosphines, an H3 group orbital of w-symmetry can overlap with a vacant dv-orbital on the phosphorus ligand. Thus, they believe that the P-B coordinate link might have a dv-pv component supplementing the classical o-bond and thereby enhancing the stability of the adduct. This description of "back-bonding" is parallel to the picture of methyl hyperconjugation except that methyl hyperconjugation leads to an increase in charge separation, whereas w-bonding from the borane group would reduce the charge separation existing in the dative a-bond. Thus, the existence of "unusual" compounds such as OCoBH3 and F3PoBH3 has been attributed to supplementary w —bonding(6 18) In general, fluorine bonded to phosphorus should increase its ability to w-bond through a removal of negative charge from the phosphorus with a concommitant lowering of the 3d7-electron energy levels

- 129 - and greater overlap with substituent v-orbitals(lo6); therefore, difluorophosphine should show less v-bonding propensity than trifluorophosphine, while certainly such bonding would be minimal for phosphine. Thus, on the basis of the aforementioned w-bonding arguments alone there is no a priori reason to suspect that PHF2-BH3 would be more stable than PF3~BH30 The stability of difluorophosphine borane relative to PF3~BH3 must then be due to a substantial increase in the a-bond strength. The a-bond contribution is supposedly essentially nonexistent in PF3-BH3 and very small in PH3oBH3 but increases just enough in PHF2oBH3 to effect a synergistic balancing of a- and v-contributions to adduct stability. The progression of these contributions might be pictured as shown in Figure 43, however, it is emphasized that this curve was arrived at after the fact and a priori considerations based on a combination of a- and w-bonding probably would not have arrived at this conclusion. On the basis of bond lengths and heats of formation VanWazer(07) suggested that many four-coordinate phosphorus compounds form multiple bonds which help to reduce the positive charge on the central phosphorus atomo However, if such a reduction in charge separation is necessary for adduct stability, in the fluorophosphine boranes reduction of the positive charge through multiple bonding from the borane group seems far less reasonable to this author than reduction by fluorine which has

- 130 - Figure 43 Hypothetical Sigma- and Pi-Bonding Contributions to Adduct Stability overal 0 _T o —-...k\ // \ non-bonded electrons available for back-donation~ In fact, some authors(l16'108) feel that even with the tri-coordinate phosphorus in PF there is considerable degree of multiple bonding which arises in order to compensate for the drift of negative charge through the a-system to the PF3 PPFl hW PI non-bonded electrons available for back-donationo In fact, some authors^lOElO8 feel that even with the tr-coordinate phosphorus in PP3 there is considerable degree of multiple bonding which arises in order to compensate for the drift of negative charge through the s-system to the highly electronegative fluorine atomso Theoretically, the degree of multiple bonding should increase upon addition of a fourth substituent since the dative bond formation implies an increase in the formal positive charge on phosphorus. If indeed back-coordination is essential for adduct stability, the inability of BH3 to back-coordinate would also help to explain why PH3~BH3 is unstable, for since there are no non-bonded electron pairs in coordinated

- 131 - PH3, certainly it cannot be expected to reduce the separation of charge through back-donation (Figure 44) On the other hand in the fluorophosphine boranes the fluorine atoms have non-bonded electron pairs available for multiple bonding (Figure 44), Figure 44 A Possible Back-Donation Mechanism ~ e "~-,,.4-,' e,* ~ 1 *; 11 s — O ] 4 H:P: B MI;:F Pi B:I P - F- CF::P36s::HI4 etc,:F: H'F H-3P -B3 3FP- BH3 no back-doncation bac k - don+at on The "back-bonding" hypothesis has also been forwarded to explain the "reverse order" of coordination(71719) shown by the borane group (P>N, S>O); however, this "reverse order" has recently been shown to be a function of the strength of the reference Lewis-acid(20) Thus, the coordinating ability of phosphines and other similar polarizable ligands like sulfides is strongly dependent on the field strength of the reference acid, ioeo the acid o: high field strength forms a stronger coordinate link with the more polarizable ligand.

- 132 - If a model that neglects w-bonding contributions is used, the observed base strength towards borane, PHF2 > PF3 > PH3, can be related to the polarizability of the electron pair on phosphorus. The arguments are similar to those forwarded by Weaver and Parry(09) to explain the base strengths of the alkyl amines and the alkyl phosphines. The lone electron-pair on phosphine has a high degree of "s-character"' and is tightly held by the three closely bonded protons; thus, the pair has a relatively low polarizability and is not available for the formation of a strong coordinate bond. Since fluorine is bonded at a greater distance than hydrogen and is larger than hydrogen, the substitution of fluorine for hydrogen should have a loosening effect on the lone-pair, albeit the effect would be smaller than for the isoelectronic methyl group because of the electronegativity of the fluorines and the relatively small change in bond distance (for PF3, P-F = 1.55 A52); PH3, P-H = 1.42 A(78)) The electronegative fluorines would also be expected to distort the loosened P-F bonding electrons as well as the phosphorus lone-pair, thus shifting negative charge away from the phosphorus and reducing its a-bond base strength. In order to fit the observed trend this model requires that the loosening due to the removal of one hydrogen is large and that the tightening by the fluorines becomes effective only in the case of PF3. The two opposing effects can be pictured as shown in Figure 45. In line with this

- 133 Figure 45 Hypothetical Contributions of Fluorine and Hydrogen to Phosphorus Lone-Pair Polarizability </ "\ oX~verall hydrogen'0" L _ c a - I \ /fl fluori \ne t I \ PF3 PtH, P5zF P-3 model Weaver and Parry(109) have attributed the ncreasing base strength in the series PH3 through PMe3 to a loosening effect on the electron pair as protons are replaced by methyl groups. The series PH3 through PF3 differs importantly in that fluorine is more electronegative than a methyl group. Again the latter trends (Figure 45) are presented after the fact and probably such a marked loosening of the phosphorus lone-pair would not have been anticipated upon replacement of a fluorine in PF3 with a hydrogen. A mechanism which might cause a loosening effect when one F is replaced by an H is presented belowo An interaction between the protonic hydrogen and the non-bonded electron pairs of the fluorine atoms might

- 134 - account for the anomolous base strength of difluorophosphine towards borane. Such an interaction can be pictured as reducing the effective positive charge of the hydrogen and thus permitting a distortion of the P-H bonding electrons towards phosphorus and a loosening of the phosphorus lone-pair electrons. In phosphine no such interaction is possible and the protons are very effective in contracting the lone-pair and rendering it essentially unavailable for dative bond formationo Extrapolation of this model to the as yet unknown fluorophosphine, PH2F, would predict a species with greater base strength than PH3, probably comparable to that of PF39 but not as great as PHF2 since one fluorine could not as effectively attenuate the positive charges of two protonso Although there is no unequivocal evidence for such an H —-F interaction in difluorophosphine, the P-H stretching frequency is low (2240.5 vs 2327, 2421 cm-' for PHs83)) which in itself is indicative of a reduction in the electron density of the P-H bond; a preliminary vibrational analysis of PHF2 implies that an appreciable interaction between P-H and P-F motions may existo An accurate determination of the molecular parameters for PHF2 will provide an interesting test of the latter model; the proposed interaction would be expected to give rise to short H —-F distances, and an F-P-F angle more acute than those in PF3. The H-P-F angle would probably approximate 90~ but the size of the fluorines should preclude H-P-F

- 135 - angles much smaller than the H-P-H angles in PH3. Previously, the correlation of JPB with the strength of the coordinate bond in the methyl phosphine boranes and the fluorophosphine boranes was noted (Table 27) The 31p chemical shifts for the series PHXF3-X (x = 0,1,3) (Table 28) tend toward lower field with increasing base strength if BH3 is used as the reference acid; a similar trend is also evident for the methyl phosphines PHxMe3-x (x = 0,1,2,3) (Table 28). There appears to be general agreement(110-115) that 31p shifts are related to the electrostatic potential produced at the nucleus by the electrons. Such potential is a result of diamagnetic shielding which may be: 1) enhanced by multiple bonding, 2) attenuated by a "second-order paramagnetic term' due to electronic asymmetry around the nucleus, or 3) decreased due to the ionic nature of the bonds induced by electronegative substituents. The trend in the methyl phosphines is probably related primarily to a reduction of the electrostatic potential at the nucleus resulting from replacement of the hydrogens by the larger methyl groups. Such substitution gives rise to a transition from a tight, highly-screening environment around phosphorus to a more diffuse electronic environment providing less shieldingo On the other hand, the factors contributing to the 31P chemical shifts of PF3 and PHF2 are felt to be more numerous and complex. Since, in direct contradiction to predictions based on electronegativity alone, the PF3

Table 28 NMR Data* for PHF2 PHF2 BHs PF3 PF3 ~BH3 PH3 PH3 BH 31P nmr (OPA std) b(ppm) -224 -171 -105t -107 246t 115 A (ppm)t +53 -2 -133 19F nmr (TFA std) b (ppm) 42.2 -21.5 -42.3t -19.9 JPF (cps) 51134 1151 1441(51) 14o6 JPH (cps) 182.4 467 182.2(65a) 72 PMe(112) PHMe2112) PHMe 112) PH3 31p nmr (OPA std) 6 (ppm) 62 98.5 163.5 246t * For an explanation of notation see Appendix A. t Determined during this study, however, spectra not shown, P-B P P A (b - P) where 6 = chemo shift of free ligand, P-B and 6 = chemo shift of borane adducto

- 137 - signal appears at higher field than any of the other phosphorus trihalides, a significant shielding of the phosphorus nucleus due to multiple bonding has been (110) proposed. Since hydrogen is less electronegative than fluorine replacement of a fluorine with hydrogen in going from PF3 to PHF2 would be expected to give a Psubstituent bond with less ionic charactero The replacement of hydrogen by fluorine should also decrease the degree of multiple bonding due to a concommitant raising of the phosphorus 3d-electron energy levelso Therefore, in comparison with PF3, it would be difficult to rationalize the low field shift of PHF2 on the basis of ionic and multiple-bonding differences unless one of these effects is much larger than the othero The latter should enhance the phosphorus shielding while the former decreases it. The low field resonance of difluorophosphine might best be attributed to a marked attenuation of the diamagnetic shielding by the'second-order paramagnetic contribution" which is significant in this case because of the electronic asymmetry around phosphoruso It is tempting to try to relate the phosphorus chemical shifts mentioned above (Table 28) to the polarizability of the phosphorus lone-pair and therefore the relative base strength within both series towards boraneo The comparison seems tenable in the case of the methyl phosphines but may be fortuitous in the case of PHFa2 PF3 and PH3 in the light of the many variables which can

- 138 - affect the chemical shift. Additional data are needed to test the correlation. When a phosphine ligand forms a coordinate bond with a Lewis-acid like borane, there usually appears to be significant decrease in the shielding of the phosphorus due to the attraction of the lone-pair electrons by the acid - what might be termed a rehybridization of the phosphorus bonding orbitals or a loosening of the lonepair electrons occurs. Therefore, a shift in the phosphorus resonance of the coordinated phosphine to lower field relative to the free ligand would be expected, ioeog P-B P P the quantity defined as A = [P -B b] (where = chemo P-B shift of free phosphine-ligand and 6 = chemo shift of borane adduct) would be negative. Meriwether and Leto ) found this to be true for the great majority of nickelcarbonyl-phosphine complexes they studied. No such data could be found for other phosphine complexes, but it is interesting to anticipate the results in the case of the (109) methyl phosphine boranes. Weaver and Parry( contend that the lone-pair in PH3 is tight and relatively hard to polarize while for PMe3 the polarizability is greater and the lone-pair more diffuseo However, the arguments of Weaver and Parry('09) do not describe the change in the phosphorus lone-pair upon coordinate bond formationo Since coordination of BH3 to PH3 requires a considerable:drawing-out" of the lone-pair, the phosphorus deshielding effect should be large, ioeOo a large negative Ao

- 139 - Similarly, increasingly smaller changes in deshielding are anticipated upon progression through PH2Me to PHMe2 and PMe3 because of the greater ease (Smaller lone-pair perturbation) with which the P-B dative bond formation occurs. Thus, the negative A-values are predicted to become increasingly smaller in going from PH3SBH3 to PMe3oBH3. During the present work the 31p nmr spectra (Figure 46) and the quantity A were determined for PH3-BH39 PF3'BH3S and PHF2-BH3. The A-values do not show a regular decrease in negative value with increasing P-B bond strength as predicted above for the methyl phosphine series. However. the large negative A-value for PH3oBH3 (A = -133 ppm) is indicative of the large change in hybridization and loosening of the lone-pair required for coordination in the case of PH3; this change from -p to -sp is also implied upon comparison of the P-H coupling constants of PH3 and PH3~BH3 (Table 28), the increase in JpH indicating more "s-character" in the bonding-orbitals. Since the lone-pair in PF3 is probably slightly more diffuse that in PH3, and since it must protrude somewhat, smaller changes in hybridization would be anticipated on coordination to BH3. IThe small. negative A-value for PF3oBH3 (A = -2 ppm) indeed implies a slight change which is substantiated by the similarity in the P-F coupling constants for PF3 and PF3oBH3 (Table 28). The positive A-value for PHF2oBH3 (A = 53 ppm) is somewhat surprising. It might be attributed by some to a

Figure 46 40.4 Mc 31P NMR Spectra of PHxF3 x.BH3(x=O,1,3) ift X *3-x 3( PHF2BH3 @ 37o OPA Jp= 468c ps 1..I..1= -171 ppm JPF: 1149cps PH3'BH3 @ 370 0 OPA J 113 ppm..... ~ I JpH: 366cps "F3 ~ ~ PF3.B3 D 37~ PF3 PF3 OPA: -17 ppm J PF y

- 141 - compensation of the phosphorus deshielding through backdonation from the BH3. Alternatively, the electronic symmetry about a four-coordinate phosphorus should be greater than about a three-coordinate one. Hence the positive A-value for PHF2oBH3 could be attributed to the existence of a relatively small paramagnetic contribution to the phosphorus chemical shift in PHF2oBH3 as opposed to a large paramagnetic contribution for difluorophosphine itself. Upon coordination, a large change in the electron environment of the phosphorus in PHF2 is also implied by comparison of the 19F chemical shifts for the free and coordinated PHF2 ligand (Table 28). Extrapolating from the theory of Saika and Slichter(116) for 19F chemical shifts, this shift of the fluorine signal to lower field upon coordination of PHF2 points to a change to more covalent P-F bonds in the borane adduct. The fluorine environments must be similar in both PHF2oBH3 and PF3~BH3, judging from the closeness of their chemical shifts (Table 28). There is also other evidence of a change in hybridization upon coordination of PHF2 to BH3 since the P-H coupling constant changes from 183 to 467 cps, indicative of more "s-character" in the P-H bond after coordination

- 142 - VI. Summary. The new mixed phosphorus halide, PF2I, was characterized and used as a reactive starting material in two new synthetic routes to difluorophosphine ligands via metathetic and coupling reactions. The new species, PF2CN, F2POPF2a P2F4, and PHF2, were characterized and some of the chemistry of PHF2 investigated. Difluorophosphine was found to form an interesting borane adduct PHF2~BH3 which is more stable than either PH3-BH3 or PF3-BH3. Further study showed that the relative base strength towards borane decreases in the order PHF2 > PF3 > PH3. Rationale for the order was presented. Spectroscopic studies which yielded infrared, Raman, and'H, 9F, llB, and 31P nmr data for many of the compounds were also completed and discussed in relationship to structure and bonding.

Experimental Io General Procedures. The compounds investigated were handled in a highvacuum system. High-vacuum techniques such as pressure volume measurements, vapor density determinations, etc. are adequately discussed elsewhere(ll7' 118) All stopcocks and standard-taper joints were greased with ApiezonNo In order to minimize the contact of reactive materials with mercury, manometers were isolated from the system with stopcocks. Non-volatile materials were handled in an atmosphere of dry nitrogen. IIo Starting Materials. Dimethylaminodifluorophosphine was made in a manner (29) similar to that described by Schmutzler by the fluorination of dimethylaminodichlorophosphine with NaF in a suspension of tetramethylene sulfone. Hydrogen Iodide was prepared as outlined in the literature(19) by the reaction of iodine and 1,2,3,4tetrahydronaphthalene. Dimethylamine was obtained from Matheson Co., Inco and passed through a -78~ U-tube before use. Diborane was kindly supplied by Callery Chemical Co.o it was purified by fractional condensation. Phosphine was prepared by the pyrolysis of phosphorus acid(l20) at 200-250~o - 143 -

- 144 - Phosphine Borane was formed at -105~ from the combination of liquid diborane and phosphine(93). When the desired amount of product had formed the excess reactants were pumped from the tube which was maintained at -78~. Trifluorophosphine was obtained from the Ozark Mahoning Co. and purified by passing it through a -160~ trap before use. Trifluorophosphine Borane of high purity was donated by Ro J. Wyma of Dr. R. C. Taylorts research group. Nickel Tetracarbonyl did not require further purification as obtained from the Matheson Co.o, Inc. In general, volatile materials were checked for purity before use by vapor pressure, molecular weight, and/or infrared spectroscopy. Cuprous Oxide (Baker & Adamson reagent grade) and Cuprous Cyanide (Baker & Adamson technical grade) were used without further purification. Mercury was triply distilled in Dr. R. Co Taylor's laboratories. III. Preparation and Reactions of PF2I. A. Preparation. In a typical experiment difluoroiodophosphine, PF2I, was prepared by separately condensing 12.76-mmoles of F2PNMe2 and 25.52-mmoles of HI into a 1000 cc reaction bulb equipped with a stopcock. The bulb was then allowed

- 145 - to warm slowly to 25~. Reaction was indicated by the formation of clouds of pale red-yellow solids. The bulb was allowed to stand at 25~ for 15 min. after the clouds had settled and then the volatile products were separated by fractional condensation (-126~0 -1960) The desired difluoroiodophosphine (11.96-mmoles) was retained at-126~ while PF3 (0.25-mmole) passed through the -126~ trap and was held at -196~. It is important to use exact stoichiometric amounts of HI and F2PNMe2 if pure PF2I is desired since it is difficult to separate PF2I and F2PNMea, and since excess HI appears to promote the decomposition of PF2I in the liquid state. A 244.6 mg. sample of PF2I was hydrolyzed in 25 ml. of 40% NaOH and the solution analyzed for P, F, and I by Huffman Laboratories, Inc., Wheatridge, Colo.; results are given in the Discussion. B. Reaction of PF2I and Cu20 - Synthesis of F2POPF2. A 40 cc reaction tube equipped with a stopcock was charged with "O.5 go Cu20 (3.5-mmoles), attached to the vacuum system, evacuated, and PF2I (1.07-mmoles) condensed into the tube. The contents of the tube were allowed to warm to 25~ and cooled to -196~ several times. Reaction was evidenced by the formation of tan solids. Finally, the volatile products were removed from the vessel by pumping through a series of traps held at -112, -145, and -196~o It was necessary to heat the tube to -150~ with a hot-air gun to remove the last trace of products. Retained

- 146 - at -112~ was 0.05-mmolecf an unidentified material; PF3 (0.04-mmole) was stopped at -196~o The desired F2POPF2 (0.39-mmole) corresponding to 73% yield based on the amount of PF2I taken was removed from the -145~ trap. The x-ray powder photograph of the solids remaining in the tube was consistent with a mixture of CuI and Cu20o The preparation was also carried out on a considerably larger scale with somewhat smaller yields. C. Reaction of PF2I and CuCN - Synthesis of PF2CNo A reaction tube (70 cc) was charged with -1 go (ll-mmoles) of CuCN by shaking the powder through a piece of polyethylene tubing which had been inserted into the vessel through the 4 mm vacuum stopcocko The tube was placed on the vacuum system and o139-mmoles of PF2I were condensed onto the powder. After repeatedly warming to 25~ and cooling to -1960 the volatile products were removed from the tube and separated by fractional condensation (-95, -126, -196~). It was necessary to heat the reaction vessel to -150~ with a hot-air gun to remove the last trace of product. The -196~ trap contained 0o31mmole of PF3 while an undetermined amount of unreacted PF2I was held at -126~. Cyanodofluorophosphine (O044mmole) was retained at -95~ which corresponds to a 32% yield based on the PF2I taken.

- 147 - D. Reaction of PF2I and Mercury - Synthesis of P2F4. Difluoriodophosphine (2.39-mmoles) was condensed into a 1000 cc bulb equipped with a stopcock and containing 2 cc of mercury. The bulb and contents were slowly warmed to 25~0 removed from the vacuum system, and shaken. A discoloration of the mercury and formation of yellow and gray solids indicated reaction. After 4 hr. of shaking, the bulb was attached to the system and the volatile contents removed by pumping through U-tubes maintained at -112, -126, and -196~o The -196~ fraction was 0.10mmole of PF3 while 0.30-mmole of unreacted PF2I (contained undetermined amount of F2PNMe2) was retained at -112~. The desired P2F4 (0.97-mmole) slowly passed through the -112~ trap and was held at -126~. This corresponds to a 93% yield based on the amount of PF2I which reacted. Another experiment was conducted in which the initial PF2I pressure was higher. Difluoroiodophosphine (0.84mmole) was shaken with mercury in a 35 cc reaction tube. After 21 hrs. the volatiles were investigated and found to consist entirely of PF3 (0.52-mmole). Perhaps the additional reaction time in the latter experiment precluded the observation of P2F4, or the PF2I disproportionated under influence of mercury at the higher pressure. E. Reaction of PF2I, HI and Mercury - Synthesis of PHF2. A 2 cc sample of mercury was placed in a 70 cc reaction tube which was equipped with a stopcock and

- 148 - standard taper joint. The tube was evacuated and 3.36mmoles of PF2I and 3.36-mmoles of HI were condensed into the tube at -196~o After warming to 25~ the tube was shaken for 1.5 hr. The tube was then opened to the Toepler system through two -196~ traps and O.17-mmole of H2 (identified by gas density) was recovered. The products condensable at -196~ were passed through traps at -140, -160, and -196~. A 0.72-mmole sample of PF3 was trapped at -196~ and a 0.47-mmole sample of an unstable compound assumed to be PHF2oHI was retained at -140~. This latter compound decomposed upon warming to room temperature to yield PF3, SiF4, and yellow solids which contained iodine. A 1.85-mmole sample of the desired PHF2 was retained in the trap at -160~. Another run using 1.52-mmoles of each reagent gave 0.84-mmole of PHF2, 0.30-mmole of PF3, 0.10-mmole of H29 and 0.18mmole of the unstable species retained at -140~ as products. A satisfactory mass balance is obtained for each run if one assumed that PF3 was formed by 3PF2I - 2PF3 + PI3 and that the unstable material stopped at -140~ was PHF2oHI. Run 1: mmoles in: P, 3036; F, 6072; H, 3.36; mmoles out: P, 3.40; F, 6.80; H, 3.13. Run 2: mmoles in: P, 1.52; F, 3.04; H, 1.52; mmoles out: P, 1.47; F, 2.94; H, 1.40. The x-ray powder pattern of the gray solids formed during the reaction was consistent with that of Hg2I2o A sample of difluorophosphine (0.81-mmole) was

- 149 - pyrolyzed over chunks of uranium metal. The hydrogen evolved (0.41-mmole) was identified by gas density. IV. Reactions of PHF2. A. Hydrolysis of PHF20 Difluorophosphine (0.49-mmole) was condensed into a tube containing 10 ml. of 40% NaOH and then the tube was sealed off and held at 100~ for 2 hrs. After opening the tube 0.44-mmole of H2 was recovered in the Toepler system. In a similar experiment involving PDF2 the hydrogen evolved on hydrolysis was found to contain no deuterium by mass spectroscopic analysis. In a separate experiment, water vapor (0.32-mmole) was allowed to mix with PHF2 vapor (0.43-mmole) at 25~ in a 500 cc bulb. The only volatile recovered after 15 min. of reaction was silicon tetrafluoride (0.17-mmole)o B. Aminolysis of PHF20 The reaction vessel consisted of two bulbs (50 and 500 cc) separated by a stopcock so that the reactants could be introduced into each bulb separately and then mixed as gases by opening the stopcocko In one experiment 0.56-mmole of PHF2 was condensed into the 50 cc bulb and 0o36-mmole of HNMe2 was introduced into the 500 cc bulb. Upon mixing the gaseous reactants at 250, white clouds formed which turned pale yellow-brown upon settling. After 10 min. the volatile products were fractionated

- 150 - (-145, -196~); a trace amount of PFs was isolated at -196~ and an unstable species thought to be FHPNMe2 (0.16-mmole) remained at-145~. A preliminary characterization of the latter species is presented in the Discussion Section. In two other experiments the HNMe2/PHF2 ratio was increased and the major product produced appeared to be HP(NMe2)2 as shown in the Discussion Section. In one case 1.08-mmoles of PHF2 were mixed with 1.95-mmoles of HNMe2 in the same manner as described in the previous experiment. Again a trace amount of PF3 was separated by passage through a -135~ into a -196~ trap. Stopped at -1355 was 0.93-mmole of an unstable mixture which was difficult to cleanly free of excess dimethylamine; however, a rough separation could be obtained by fractional condensation at -95 and -135~, the supposed HP(NMe2)2 being retained at -95~O Similar results were obtained when 0.55-mmole of PHF2 was mixed with 2.08mmoles of HNMe2. Because of the instability of the species further experiments were not undertaken. Co Reaction of PHF2 and B2H6. Diborane (1.94-mmoles) and difluorophosphine (1.92mmoles) were condensed into a 150 cc vessel attached to a manometer and allowed to warm to 250. After 4 hr., when the pressure in the vessel had become constant at about 3 its calculated initial value, the products were separated by fractional condensation (-135, -196~) Retained in the

- 151 - -1355 trap was PHF2'BH3 (1.57-mmoles). A mixture of B2H6 and PF3 (1.35-mmoles) was stopped at -196~. The amount of PF3 was estimated to be about 0.20-mmole from the infrared spectrum of the mixture. A small amount of yellow solid and a trace of noncondensable gas were also formed during the reaction. Do Reaction of PHF2 and Ni(C0) 4 Nickel tetracarbonyl (0.43-mmole) and difluorophosphine (2.00-mmoles) were separately condensed into a 40 cc reaction tube equipped with a stopcocko The reactants were warmed to and held at -23~ for 3 hrs. The rate of evolution of noncondensable gas increased after warming the tube to 0~ where it was maintained for 17 additional hours, and finally after 18 hrs. at 25~ the evolution of CO had ceased and the products were recovered by fractional condensation (-112, -196~). Carbon monoxide (0o32-mmole identified by gas density) was collected in the Toepler system while 1.86-mmoles of unreacted PHF2 were retained at -196~. A difluorophosphine derivative of Ni(C0)4 thought to be (PHF2)Ni(C0)3 was recovered from the -112~ fraction (0o15-mmole). Characterization of the latter species is mentioned in the Discussiono The ratio of reactants was varied in another experiment in which 2,05-mmoles of Ni(C0)4 and l.86-mmoles of PHF2 were allowed to react at 25~ for 2 hrs. in a 40 cc tube. Carbon monoxide (1o34-mmoles) was displaced in this timeo Fractionation (-112, -195~) showed that all the PHF2

- 152 - had reacted. An undetermined mixture of difluorophosphine substituted nickel carbonyls was retained at -112~ (the sample was inadvertantly lost). Passing the -112~ trap was an appreciable amount of SiF4 (ca. 0.5-mmole). In this and in the previous experiment, appreciable amounts of black, non-volatile oils remained behind. Thus, it is felt that side-reactions also occurred. V. Base Displacement Reactions. Ao PHF2 and PF3~BH3. In a typical experiment, 059-mmole of PHF2 was condensed into a 40 cc tube with 0o83-mmole of PF3"BH3s The tube was maintained at -78~ for 25.5 hrs. after which time 0.34-mmoles of PF3 was isolated (57% yield). Everything was condensed back into the tube and the reaction allowed to proceed for 24.5 additional hours at -45~ before the products were finally separated by fractional condensation (-126, -160, -196~). Trifluorophosphine (0.58 —mmole) passed through the -160~ trap, while PF3-BH3 (0.22-mmole) was retained at -160~. Difluorophosphine borane (0.57-mmole, 97% yield) was retained at -126~o The best separation was effected when distillation through the -126~ U-tube was rapid. Upon prolonged standing at -126~, PHF2~BH3 slowly passes to colder regionso Bo PHF2 and PH3oBH3o Phosphine borane (1o93-mmoles) and difluorophosphine

- 153 - (0o91-mmole) were condensed into a 43 cc reaction tube. The tube was maintained at -23~ for 5.5 hrs. after which time the products were recovered; it was found that displacement had not proceeded to an appreciable extent. Everything was condensed back into the tube and maintained at 0~ for 12 additional hrs. The tube was then cooled to -78~ and the volatiles removed through U-tubes held at -126, -160, and -1960. A very small amount of H2 (Oo02-mmole) passed through the -196~ trap which held a 0.96-mmole mixture of PH3 and PF3 plus a trace of B2H6. This mixture weighed 34.4 mg. which indicates that it contained ca. 0o04-mmole PF3 and ca. 0.91-mmole PH3. All of the PHF2 had been consumed as indicated by the fact that nothing was retained at -160~o The -126~ trap held 0.84-mmole of PHF2IBH3o The solids left in the reaction tube were composed of a small amount of a non-volatile yellow species and 0.97-mmole of PH3~BH30 The latter decomposed on warming to room temperature to give a mixture (1o46-mmoles) of B2H6 and PH3. When the displacement of PH3 from PH3oBH3 by PHF2 was attempted at lower temperatures and/or higher pressures, it was complicated by an increased amount of disproportionation as indicated by the formation of appreciable amounts of PF3 and non-volatile yellow solids. Under these conditions the displaced PH3 appeared to react with free PHF2 to give an unstable species which has not been characterized completely but which does display

- 154 - spectral characteristics similar to those attributed to PHF2*HI. Co PH3 and PF3sBH3. Trifluorophosphine borane (1.00-mmole) was condensed into a 43 cc reaction tube equipped with a stopcocko Phosphine (1.82-mmoles) was also added to the tube which was then warmed to -78~ for 4 hrso, however, the reaction proceeded only slowly at this temperature. After an additional 10 hrs. at -35~, separation of the products showed that Oo56-mmole of PF3 had been displaced. All the products and reactants were returned to the tube which was then maintained at 0~ for 16 additional hrs. Finally, the products were separated (-112, -160, -1960) A mixture (191-mmoles) of PF3 and PH3 (trace of B2H6) which weighed 109.4 mg. passed the -160~ trap. Retained at -160~ was 0l13-mmole of PF3-BH3s The -112~ trap contained a white solid characteristic of PH3oBH3s A mass balance of products and reactants indicates that 87o of the PF3 was displaced under these conditions according to the equation PF3 BH3 + PH3 - PF3 + PH3sBH3s A similar experiment was conducted, but trifluorophosphine was added to the system in order to test its effect on the extent of displacement. Equal amounts (Oo99-mmole) of PF3-BHs, PHs and PF3 were held at -35~ for 10 hrs.; separation of the products and reactants

- 155 - indicated that 27% of the PF3 was displaced. Everything was returned to the reaction tube and maintained at 0~ for an additional 20 hrs. Separation yielded 2.17-mmoles of a mixture (weight 153.9 mg.) of PF3 and PH3 (trace of B2H6), plus 0.51-mmole of PF3sBH3. Thus only 48i of the PF3 was displaced under these conditions.

Appendix A Spectroscopic Notation and Procedures Infrared Spectra. Unless otherwise noted the spectra were obtained with a Beckman IR-12 which operates in the 4000 - 200 cm 1 range. The precise vibrational frequencies listed in the Tables were determined by recording each band under high dispersion. A gas cell with a 75 mm path length and CsI windows which were held in place against 0-ring seals by atmospheric pressure was used for volatile substances. The infrared spectra of solids were observed at ca. -180o in transmission through a thin film sublimed onto the CsI window of a cold cell similar to that described by Wagner and Hornig( ) Raman Spectra. Pure liquids were distilled into 2 mm i.do capillary tubes fitted with flat ends suitable for transmission of the scattered light; the PH3sBH3(~) sample was prepared in situ as the solid (mop. 33~) and warmed to 35~ for the Raman spectrum; the unreacted diborane and phosphine were removed before the capillary tube was sealed. For solid PH3~BH3 the sample was also prepared as just described in a 5 mm o.do pyrex tube and transferred to a cell designed especially for the determination of the Raman effect of solids at -180~o The spectra were graciously determined by C. Fo Farran; further description - 156 -

- 157 - of experimental details and equipment design may be found (122} in his thesis(12). The Raman spectra displayed in this text are microdensitometer traces of the photographic plates. Notation for Vibrational Spectra. The following abbreviations and symbols are used in the Tables: v, stretching; 6, deformation; c, wagging; p, rocking; T, torsional; as, asymmetric; s, symmetric; p, polarized; dp, depolarized. For intensities and band shapes: br, broad; sh, shoulder; v, very; s, strong; m, medium; w, weak. P, Q, and R refer to band shapes reminicent of typical P, Q, and R branches arising from vibrational-rotation interaction. NMR Spectra. The spectra were determined on Varian Associates NMR Spectrometers. Most of the 1H nmr spectra were obtained on the A-60. The spectra of other nuclei were determined by coupling the HR-100 with the appropriate radiofrequency unit i.e., lB (32.1 Mc), 19F (94,1 Me), and 3 P (40.4 Mc). In a few cases where the chemical shifts of non-equivalent hydrogens overlapped, the hydrogen spectra were obtained at 100 Me so that the signals were separated sufficiently for easy interpretation. Conditions such as sample temperature and radiofrequency are included with the individual spectrao Samples were prepared by condensing the pure substances into 5 mm o.do semi-micro nmr tubes (NMR Specialities, Inco, New Kensington, Pao)

- 158 - and sealing the tube under vacuum. However, in the case of PH3~BH3 the solid was prepared in situ from stoichiometric amounts of PH3 and B2H6; it melted at 33~ without evidence of significant decomposition. Signal separations were measured by calibrating the sweep with side bands of a known frequency except in the case of the proton spectra determined on the A-60 where the sweep is internally calibrated. Chemical shifts were obtained by tube interchange. NMR Notationo Coupling constants, J, are expressed in cycles'per second, cps; their assignment is indicated by the appropriate subscript, i.e., JPBH refers to the interaction between phosphorus and hydrogen nuclear spins across two bonds, the P-B and the B-H bonds. Chemical shifts, 6, were calculated according to the equation(8) H - H sample reference Hreference and are reported in parts per million, ppm. Thus, relative to the appropriate standard, low field shifts are negative and high field shifts are positive. The following abbreviations are used for chemical shift standards TIMS = tetramethylsilane, H nmr; TFA = trifluoroacetic acid, 19F nmr; TMB = trimethylborate, "B nmr; OPA = 85% orthophosphoric acid, 31P nmr. Mass Spectrao A Consolidated Electrodynamics Model 21-103 B Mass Spectrometer operating at 70 eV was used

- 159 - to obtain the spectra. Unstable samples were maintained at -196~ until just before injection into the spectrometer. In general the spectra were not determined below m/e - 12 - 14. For the spectra given in the tables all peak heights were adjusted relative to the most intense peak which was taken as 100. EPR Spectra. The spectra were graciously determined by Dr. J. Gendell on his Varian Instrumento The signals observed for P2F4 were fairly broad, 6 - 60 gausso Samples were contained in 3 mm o.d. quartz or pyrex tubes.

Appendix B Band Shape Analysis For Difluorophosphine Assuming the following molecular parameters 0 PF = 1.52 A / FPF = 90 0 PH = 1.52 A F FPH = 1000 the principal moments of inertia for PHF2 were calculated with the aid of a computer program kindly supplied by C. Fo Farran; the calculated values are given below in g. cm2xlO40: IA IB IC 22.79 45.95 64.66 The principal axes fall roughly as shown in Figure 47. Figure 47 Principal Axes of PHF2 C A I I B —— P —- H) -F- --— B view Vi w along A, along C ~~~~c- 16on C - 160 -

- 161 - The moments of inertia are essentially independent of 5, the FPH angle, and are not changed markedly by alternation of the FPF angle between 90 and 1100. Thus, difluorophosphine appears to be an asymmetric top (IA f IB f IC). Three types of band outlines, or hybrid combinations of these, are observed for asymmetric top molecules; they are classified as A, B, and C types depending along which (66,123,124) principal axis the vibration occurs(66'123l ). In Figure 48 the expected band outlines for vsPH and 6 FPF are summarized with respect to type (A, B, or C) for both the planar and pyramidal geometries of PHF2o The band Figure 48 Band Shape and Type for v PH 6 FPF s s planar PHF2 similar to smilvnar to B bronches 3 branches pyramidal PHF2 I' m/lar to similar to C ( bracnc hes branche$s or C-B hybred B observed dl

- 162 - shapes actually observed in the infrared spectrum of gaseous PHF2 (Figures 20 &48) are consistent with the pyramidal geometry. If PHF2 were planar, the P-H stretching motion would not be expected to show a strong central maximum typical of a Q-branch. The other fundamental vibrations of PHF2 give rise to bands which overlap, complicating such straightforward analysis as presented above.

Appendix C The C-N Stretching Vibration in PF2CN Using the following molecular parameters 0 PF = 1.52 A / FPF = 95 0 PC = 1.78 A j FPC = 93~ 0 CN = 1.15 A j PCN = 180~ the principal moments of inertia for PF2CN were calculated on the computer. The values are given below in g. cm2xlO40 along with the respective rotational constants in cml. IA B IC 64.17 146.58 177.87 A B C 0.44 0.19 0.16 The C-N stretching frequency for PF2CN (Figure 49) displays an interesting series of lines spaced 2.3 cm-' which were at first attributed to vibration-rotation interaction, however, this spacing is too large to be consistent with the moments of inertia listed above for pyramidal PF2CN. In fact, additional calculations show that the spacing is too great to be attributed to the planar or any other reasonable structure for PF2CN. Irrespective of a planar or pyramidal configuration for PF2CN, the calculations also indicated that the molecule - 163 -

- 164 - Figure 49 C-N Stretching Vibration in PF2CN(g) frequency 2240 2220 2200 2180 2260 1cm-)

- 165 - is probably an asymmetric top with the C-N stretch occurring along the axis of least moment of inertia (A axis). Therefore, the C-N stretch should show a strong central Q-branch (A-type band for asymmetric top(66,123124) ). The band observed in the C-N region appears to be the superposition of a typical PQR-structure and several additional Q-branches to the low frequency side of the main Q-branch. The additional Q-branches are probably due to "hot" bands(124) which arise when an appreciable number of molecules exist in a vibrationally excited state. Such a vibrationally excited state may be populated by a low frequency vibration. By roughly fitting the intensities of the "hot" bands (each successive "hot" Q-branch was estimated to have 0.42 the absorbance of the previous) to a Boltzmann distribution the energy of the low frequency vibration responsible for the "hot" bands was estimated as 180 cm-.

Appendix D Analysis of X2AA'X'2 Systems Harris(5) has recently derived expressions for the line positions and relative intensities in the X spectrum of an XnAA'X'n nuclear magnetic system when Jx = 0 (Jx is the coupling between nuclei of X and X' types). The expressions are valid only for nuclei of spin 2. Analysis of the 19F nmr spectrum of F2POPF2 was possible using the expressions of Harris; the calculations are outlined below. Unfortunately, expressions applicable to the 31p nmr spectrum of F2POPF2 were not outlined by Harris. Lack of time precluded the application of the less explicit expressions of Lynden-Bell(65b) which do apply to the 31p nmr spectrum. According to Harris for an XnAA'X'n system there may be a total of (2n + 1) pairs of lines symmetrically placed about vx, where vx is the unperturbed resonance frequency of X and Xl. For F2POPF2, 2n + 1 = 5, and three pairs of lines were observed in the fluorine nmr (Figures 9 and 50); the calculations outlined below show that one of the remaining pair has zero intensity and the other remaining pair was unresolved under other lines. If the expressions of Harris are valid in the case of F2PPF2, up to five pairs of lines would also be expected in the 19F nmr spectrum. In fact, at least - 166 -

- 167 - 19 Observed 9F NMR of F2POPF2 4 4-J L.j _) Si( -() = 1368 cps ------ Si(2) = 1372 cps So(1) = 1376 cps - A _2'4 cps- A 3 N —---- N = 1344 cps 4 Figure 50 Calculated 9F NMR of F2POPF2

- 168 - 13-pairs of lines were observed for F2PPF2 (Figure 11). Therefore, it is apparent that the assumption of Jx = 0 is invalid for F2PPF2 and the expressions derived by Harris(25) will not work in this case. Again lack of time prevented the application of other expressions which do not make the assumption that Jx = 0. Undoubtedly, the 31p and 19F nmr spectra of F2PPF2 are amenable to analysis in terms of an X2AA'X'2 nuclear spin system. The expressions used in the following are given by Harris. For F2POP'F'2 it was noted that the innermost doublet (Figure 50) contained half of the total intensity and thus it was assigned the doublet separation N = IJpF + JpOPF1 = 1344 cps, where JPF = JP'F' and JpOPF = JPOP'F' The doublet of separation N was then mentally discarded from the spectrum. The smaller separation between one member of the inner pair and one member of the outer pair of the remaining lines, 4 cps, then gave JA (Figure 50)9 where JA = Jpp,- The distance between the inner pair, Si(l) ( Figure 50), of the four lines is given by Si(l) = [L2 - JA2]2 - JAI = 1368 cps. Since IJAI = 4, then L2 = (13568 + 4)2 + (4)2 L = 1372 cps where L is the parameter IJpF - JP'OPFI Thus, the parameters N, L, and JA have been obtained from the observed spectrum and will now be used to construct the corresponding theoretical spectrum for an X2AA'X'2 system.

- 169 - The most intense doublet in the spectrum has the separation N and one-half of the total intensity [that is, relative intensity 2(2n - 1) = 8]. The other 2n = 4 pairs of lines share the remaining intensity, however, some of the lines may have zero intensity. In the left hand column below the separations of the two pairs of inner and two pairs of outer lines are calculated using the expressions of Harris(25). The corresponding line intensities are calculated in the right hand column. The calculations are a function of X which designates the transition involved; X = 1, 2 for n = 2. The limiting intensities of the lines also vary with Xand are given by n1 r(nC )(nC ) n L r r-X r=X where C is the appropriate binomial coefficient. Thus r the limiting intensities for X= 1, 2 are given by~ for Z = 1 sum = - ((2C Co2C) + (2 C2 2C1)] 1 (2-1 + 2-.12) = 3 for = 2 sum = (202C22Co) = 1 It is also necessary to weight each limiting intensity as a function of g, where g = [X(x - 1)L2 + JA2] ([X2L2 + JA2][(X- 1)2L2 + JA2]}i

- 170 - Thus for X = 1, g 0, and for X= 2, g = 1. In the calculations below, line separations are designated by So(X) or Si(X), and intensities by Zo(C) or Zi(X). The subscripts o and i refer to outer and inner, respectively. Separations Intensities 1 lim. So(l) = [L2 + JA2]i + JA Zo() = (1 - g){in ((1372)2 + 42}) + 4 = (1 - 0o)(3 = 1376 cps = 3 So(2) = (22L2 + JA2) Zo(2) = -( g)( + (L2 + JA2)} -- - i){i} = 2-1372 + 1372 = 0 = 4116 cps 1 lim. Si(l) = L2 + JA2 JA Zi) = (l + g)(n = 1372 - 4 = (1 + 0) (3 = 1368 cps (2L2 + zj()2+ lim. Si(2) = [22L2 + JA} Zi(2) =1 + g tens - (L2 + JA)}2 (1 + 1){l} = 21372 - 1372 =1 = 1372 cps

- 171 - The calculated spectrum is shown as the mirror image of the observed spectrum in Figure 50. From a comparison of the calculated and observed spectra, it is apparent that the doublet of separation Si(2) remained unresolved in the observed spectrum but did contribute its intensity to the Si(l) and So(l) doublets. Therefore, each member of the doublet of doublets has one-half the integrated intensity of each component of the central doublet. From the parameters L, N, and JA' the following coupling constants were obtained for F2POP'F'2 JPF = 1358 cps PF JP'OPF = -14 cps Jpp, = 4 cps JFPOPTFI = 1 where JpF and JpF are only relative in sign and rigorously cannot be distinguished; however, JPF certainly must be the larger. Also the sign of J, is undetermined. PPI

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