AN ABSTRACT OF THE THESIS OF RONALD EDWIN LERCH for the Ph. D. in Chemistry (Degree) (Major) Date thesis is presented April 25, 1966 Title LOW TEMPERATURE THERMAL ANALYSIS OF SOME BORON HALIDE COMPLEXES WITH WEAK LEWIS BASES Abstract approved Major professor Thermal analysis was used to investigate the following seven systems: H2S -B2F4, SOF2 -B2F4, COC12 -B2F4, SF4 -B2F4, Xe -BF3, NF3 -BF3, and NF3 -BC13. In two of the seven systems compound formation was observed. In the system, H2S -B2F4 the compounds H2S B2F4 and H S B2F4 has a melting 2H2S' B2F4 are formed. The compound 2 ± point of -94, 9 0, 3° C while the compound 2H2S B2F4 melts at - 100. 3 O. 3° C. The system has three eutectics at 16. 2, 35. 2 and 50.5 mole percent B2F4 and at -107.4 t 0. 3° C, -100.4 t 0. 3° C, and -94, 9 ± 0. 3° C, respectively. In the system, SF4 -B2F4 the compound SF4 B2F4 is formed. The SF4 B2F4 slowly decomposes at room temperature. Two ± eutectics are found in the phase diagram at 3 t 2 and 98 2 mole percent B2F4 and at -116.3 ± 0. 5° C and -59.8 ± 0. 5° C, respectively. 24 mole per- The system, SOF2 -B2F4 has a single eutectic at ± corn- cent B2F4 and a temperature of -141. 7 Q. 5o C. The two ponents are completely miscible. four mole The system, COC12 -B2F4 has a single eutectic at percent B2F4 and a temperature of -132. 5 f O. 5o C. The two components are completely miscible. BF3 is The compounds Xe and BF3 are immiscible although in liquid slightly soluble in liquid Xe. Xe is completely insoluble BF3. NF3 is The compounds NF3 and BF3 are immiscible although NF3 slightly soluble in liquid BF3. The solubility of BF3 in liquid was not investigated. NF3 being The compounds NF3 and BC13 are immiscible with in very slightly soluble in liquid BC13. The solubility of BC13 liquid NF3 was not investigated. SF4. B2F4 BF3 is shown to displace B2F4 from the complex in 93% yield. BF3 is a stronger Lewis acid than either BC13 or B2F4 compound toward the Lewis bases SOF2 and COC12 on the basis of acid formation. No conclusion can be drawn about the relative strengths of BC13 and B2F4 relative to these two bases. Lewis acid Based on acid displacement, BF3 is a stronger than B2F4 relative to the Lewis base SF4. Toward the Lewis base H2S, the order of increasing Lewis acidity appears to be B2F4 < BF3 < BC13, B2C14. This is a qualitative order based on the height of the maximum between eutectics in the phase diagrams. Some difficulty was encountered with the synthesis of B2F4. Possible reasons for these difficulties are discussed. The generally weaker acidity of B2F4 as compared to BF3 is accounted for on the basis of electronegativity and steric con- siderations. LOW TEMPERATURE THERMAL ANALYSIS OF SOME BORON HALIDE COMPLEXES WITH WEAK LEWIS BASES by RONALD EDWIN LERCH A THESIS submitted to OREGON STATE UNIVERSITY in partial fulfillment of the requirements for the degree of DOCTOR OF PHILOSOPHY June 1966 APPROVED: Associate Professor Hof Chemistry In Charge of Major Chairman of Department of Chemistry Dean of Graduate School Date thesis is presented zs Typed by Judy E. Lerch ACKNOWLEDGEMENT to Dr. T. D. The author wishes to express his appreciation during Parsons for his guidance, encouragement and inspiration thank his wife, Judy, the course of this work. He also wishes to and Dr. Robert C. Olsen for their encouragement. TABLE OF CONTENTS Page INTRODUCTION 1 Boron Halides 2 Boron Subhalides 7 Experimental Methods 12 Lewis Base Strengths 18 Research Objectives 20 EXPERIMENTAL 22 Apparatus and Technique 22 Vacuum Technique 22 Special Apparatus 23 Preparation of Materials 31 Boron Trifluoride 31 Diboron Tetrafluoride 32 Boron Trichloride 40 Sulfur Tetrafluoride 40 Thionyl Fluoride 41 Nitrogen Trifluoride 41 Xenon 41 Hydrogen Sulfide 42 Phosgene 42 Procedure 43 Thermocouple Calibration 49 RESULTS AND DISCUSSION 52 The System, Hydrogen Sulfide -Diboron Tetrafluoride 52 The System, Thionyl Fluoride -Diboron Tetrafluoride 56 The System, Phosgene- Diboron Tetrafluoride 57 The System, Sulfur Tetrafluoride -Diboron Tetrafluoride 63 Displacement of Diboron Tetrafluoride from Sulfur Tetrafluoride with Boron Trifluoride 67 The System, Xenon -Boron Trifluoride 68 The System, Nitrogen Trifluoride -Boron Trifluoride 71 The System, Nitrogen Trifluoride -Boron Trichloride 75 Lewis Acid Strengths of Diboron Tetrafluoride and Boron Trifluoride 77 Further Work Suggested by this Study 78 SUMMARY 80 BIBLIOGRAPHY 83 LIST OF FIGURES Figure Page 1 Simple phase diagram 14 2 Phase diagram showing compound formation 15 3 Portable vacuum apparatus 26 4 Freezing point cell 29 5 Typical cooling curve 45 6 Selected cooling curves for the system, H2S -B2F4 46 7 Selected cooling curves for the system, H2S -B2F4 47 8 Selected cooling curves for the system, H2S -B2F4 48 9 Thermocouple calibration curve 50 10 The system, H2S -B2F4 55 11 The system, SOF2 -B2F4 59 12 The system, COC12 -B2F4 62 13 The system, SF4 -B2F4 66 14 The system, Xe -BF3 70 15 The system, NF3 -BF3 74 LIST OF TABLES Table Page I VAPOR PRESSURE EQUATIONS OF COMMON REAGENTS 24 II COMMON COLD BATHS 25 III INTERPRETATION OF SYMBOLS USED IN FIGURES 3 AND 4 27 IV DATA FOR THE SYSTEM, H2S -B2F4 54 V DATA FOR THE SYSTEM, SOF -B2F4 58 VI DATA FOR THE SYSTEM, B2F4 -COC12 61 VII DATA FOR THE SYSTEM, B2F4 -SF4 65 VIII DATA FOR THE SYSTEM, BF3 -Xe 69 IX DATA FOR THE SYSTEM, BF3 -NF3 73 X DATA FOR THE SYSTEM, BC13 -NF3 76 LOW TEMPERATURE THERMAL ANALYSIS OF SOME BORON HALIDE COMPLEXES WITH WEAK LEWIS BASES INTRODUCTION For many years, boron containing compounds, especially boron halides, have served as classical examples of Lewis acids. When a Lewis base is mixed with the boron compound, a molecular addition compound may form. These molecular addition corn- pounds (Lewis adducts) may be very stable or may exist only under carefully defined conditions, This research deals with interactions of boron halides and weak Lewis bases at low temperature investi- gated by means of thermal analysis. The ability of simple boron compounds to complex has been known since the first molecular addition compound between BF3 and NH3 was prepared by Gay Lussac in 1809. The development of any theoretical understanding of the coordination process, however, was to follow more than a century later. In 1923, G. N. Lewis (48) first presented the fundamental concepts of his electronic theory of acids and bases in which an acid is defined as a molecule, radical, or ion capable of accepting a pair of electrons to form a coordinate bond, and a base is any molecule, radical, or ion capable of donating an electron pair to form a coordinate bond. This general theory was later discussed 2 and extended by both Lewis (47) and Luder (49). With this theory, many hitherto unrelated phenomena could be explained. Boron Halides The boron halides represent a group of typical Lewis acids. Interest in the relative electron acceptor properties (and thus acid strength) of these boron halides dates back to the original introduc- tion of the electronic theory by Lewis. Tr. aditionally, the Lewis acid strength is thought to increase in the order BBr3 < BC13 < BF3(51). This idea is probably due to the fact that far more com- plexes of BF3 than of other boron halides have been characterized. More recently it has become apparent that the order of relative acid strength is dependent upon the particular base that is used for coordination. Using pyridine and nitrobenzene as reference bases, Brown and Holmes (18) have shown that the acceptor properties of the boron halides increase in the order BF3 <BC13 < BBr3. This same apparent reverse order probably occurs for AsH3 also since Stieber (67) and Stock (68) report the formation of the adducts H3As BC13 and H3As BBr3, respectively, whereas Martin and Diel (54) report that H3As BF3 does not form even below -100° C. Similarly, 1blrnes (44) reported that the adduct C13P BBr3 exists, but the PC13 adducts to BF3 and BC13 do not exist. In discussing the relative acceptor properties of the boron 3 halides, it is necessary to consider the electron deficiency of the boron atom in its particular molecular environment. According to the concepts of modern valence theory, the bonding electrons of boron in BX3 occupy three sp2 hybrid orbitals at angles of 120o, p orbital perpendicu- leaving a vacant or perhaps partially vacant z lar to the plane containing the three óbonds. Thus, in terms of the octet rule the electronic environment of the boron atom is in- complete, and in the presence of suitable electron donor species, the boron atom will act as an electron pair acceptor. The electron p orbital is of special density in the region of the normally vacant z importance. The electron distribution in the cT bonds is mainly a function of the size and electronegativity of the halogen. Since electronega- tivity increases and size decreases in the order I, Br, Cl, F, the electron deficiency on the boron atom in BX3 should increase in the order [BI3] < BBr3 < BC13 < BF3. Pauling (61, p. 317 -320) has suggested that the normally vacant p orbital of boron is at least partially filled by means of z intramolecular back 'rr bonding utilizing non -bonding p electrons of the halogens. The marked shortening of the boron -fluorine bond in BF3 led Pauling to suggest the following resonance: 4 S+ X: :X: =X: S: -G-3- B s -E----- B \\ ¡ 1 .///BS- II. + II O \ X X iXj XS X/ XU+ Due to the increasing amount of p -p overlap, the back it bonding would increase along the series BBr3, BC13, BF3.