AN ABSTRACT OF THE THESIS OF George Vance Gritton for the M. S. (Name of student) (Degree) in Chemistry presented on p^,--.,, •;•?. (Major) (Date) Title: STUDIES RELATING TO THE PREPARATION AND PROPERTIES OF PERFLUOROALKYL SUBSTITUTED BORANES Abstract approved: Redacted for privacy Theran D. Parsons The structures of the boron hydrides are surveyed. Aspects of several theories of electron deficient bonding are discussed. Mechanisms for the disproportionation of unsymmetrical boranes are reviewed and the chemical consequences of the ability or lack of ability of the CF - group to participate in three-center bonds are discussed. Possible approaches to the formation of CF-B bonds, including successful and unsuccessful attempts to synthesize such bonds, are examined. Successful synthesis of a new compound, CF B(CH_) , has been accomplished. The preparation reactions are: 1. 2CF-I + Hg +Cd ACd • (CF,)_Hg +Cdl 3 Amalg. 3 2 2 2. BBr3 + 2(CH3)3B -3(CH3)2BBr 3. (CF3)2Hg + 2(CH3)2BBr -2 CF^fCH^ + Hg Br2 The experimental molecular weight of CF~B(CH_)~ has been determined to be 112. The compound is monomeric in the vapor phase. CF B(CHJ is unstable under the conditions of its synthesis. The major identifiable product of the decomposition was determined to be (CH3)_BF, There is evidence that, in addition, a polymer of CF_ is formed. The conclusion is drawn that the CF - group, in the physical and molecular environment involved, is unable to participate in the formation of three-center bonds. Studies Relating to the Preparation and Properties of Perfluoroalkyl Substituted Boranes by George Vance Gritton A THESIS submitted to Oregon State University in partial fulfillment of the requirements for the degree of Master of Science June 196 7 APPROVED: Redacted for privacy Redacted for privacy Chairman of Department of Chemistry Redacted for privacy Dean of Graduate School & Date thesis is presented ^"jJUugULu. °?y /<?6> 7 Typed by Opal Grossnicklaus for George Vance Gritton ACKNOWLEDGMENT The author wishes to express his sincere appreciation of the guidance, encouragement, and patience of Dr. Theran D. Parsons throughout the course of this investigation. The material aid of the National Science Foundation for a portion of this work is acknowledged with appreciation. TABLE OF CONTENTS I. INTRODUCTION 1 II. EXPERIMENTAL 17 Apparatus and Equipment 17 Reagents 19 Preparations and Reactions 24 III. DISCUSSION AND CONCLUSIONS 48 IV. SUMMARY 55 BIBLIOGRAPHY 56 LIST OF FIGURES Figure Page 1 Typical reaction or storage vessel. 17 2 Vapor pressure vs. l/T for CF BMe 32 3 IR spectrum of one fraction of impure Me B from the decomposition of CF BMe . 34 4 IR spectrum of a second fraction of impure Me B from the decomposition of CF BMe . 35 5 IR spectrum of Me BF from the decomposition of CF3BMe2. Z 37 6 Vapor pressure vs. l/T of Me BF from the decomposition of CF BMe . 38 7 IR spectrum of the highly volatile fraction from the decomposition of CF BMe . 39 8 Vapor pressure vs. l/T of an unknown substance resulting from the reaction of CF I with Me N. 41 9 Vapor pressure vs. l/T of an unknown substance resulting from the reaction of CF I with pyrolized LiBH4. 45 10 IR spectrum of an unknown substance resulting from the reaction of CF I with pyrolyzed LiBH . 46 LIST OF TABLES Table Page 1. Constant temperature baths 20 2. Vapor pressure thermometers 20 STUDIES RELATING TO THE PREPARATION AND PROPERTIES OF PERFLUOROALKYL SUBSTITUTED BORANES I. INTRODUCTION Until recently, interest in boron hydride chemistry was primar ily academic and was mainly concerned with the failure of the boron hydrides and some of their derivatives, as well as certain analogous compounds of other metals such as aluminum, gallium, and platinum, to conform with the rules which relate chemical composition with classical valence theory. Within the last 15 years or so, however, the boron hydrides and their derivatives have assumed great practical importance. This surge of interest has been the result of the use of certain boron compounds as high energy fuels, of others as fuel addi tives, and of still others as powerful and selective reducing agents in organic chemistry. Because these compounds depart from classical valence theory, it is apparent that knowledge of their structure and bonding is of prime importance in applications such as those above and in the potential discovery of new applications. It is also possible that a knowledge of the bonding of these compounds may be of help in elucidating the nature of the bonding in other troublesome areas, such as the bonding in intermetallic compounds. For these reasons, it is interesting to trace the development of the theory of the structure and bonding of the boron hydrides and their derivatives. A series of boron hydrides, or boranes, was first character ized by Alfred Stock (46) and his collaborators in the period from 1922 to 1936. His pioneer work in the field stimulated others, not ably H. I. Schlesinger, to enter this area of inorganic chemistry around 1930. Schlesinger and his co-workers developed new and improved methods of production of the boron hydrides, characterized many new compounds related to them, and made numerous structural studies. By 1942 boron hydride chemistry was considered of suffi cient importance to warrant an extensive review by Schlesinger and Burg (42). The hydrides of boron and their related compounds are in a class characterized by electron-deficient bonding. This term is applied to those compounds which contain insufficient valence elec trons to permit all of the adjacent atoms to be held together by simple two-atom electron-pair bonds. Other examples of compounds of this typeare Al (CHJ,, Al CI,, and [ (CH3) Pt] . In a wider sense, the term "electron-deficient" is used to designate those compounds in which the number of available orbitals of approximately equal energy is greater than the number of normally available electron pairs. This definition allows the inclusion of compounds such as (CH-KB and BF., in which the p orbital of the valence shell of boron is nominally vacant. Diborane, B_H/? may be considered as an example of such 2 6 compounds. Each boron atom can contribute three bonding electrons and each hydrogen atom, one. There is available for bonding, then, a total of 12 electrons, or six electron pairs. This is not sufficient to bond together the eight atoms with "normal" two-electron bonds. Soon after the determination of the molecular formula of diborane by Stock, several structural theories were proposed (42). The first, however, that met with any degree of success was that of the one-electron bond proposed by Sidgwick (43, p. 103). This was later expanded by Pauling (29, p. 23 9-244) to include other resonance structures, including a no-electron bond. Bauer's interpretation of electron-diffraction data (4, 5), which precluded bridged structures such as Dilthey (10) had earlier proposed, strongly supported the essentially ethane-like structures upon which these bonding theories were based. As further experimental data were recorded, these structures were abandoned. The compounds were found to be diamagnetic, a finding which was at variance with any structure postulating a one- electron bond. Infra-red spectroscopy studies by Price (32, 33) were instrumental in establishing a bridged structure to be the correct configuration of diborane. More precise electron diffraction studies by Hedberg (18) aided in definitely establishing this structure, and in removing the objections raised by the previous work of Bauer. 4 An ethylene-type structure with two protons buried in a double bond connecting the two boron atoms was proposed by Pitzer (31), It was felt, however, that proposals concerning the bonding of dibor ane should be applicable to compounds such as the dimers of alumi num chloride and trimethyl aluminum as well, and here Pitzer's concept was not well received, for "methylated double bonds" seemed quite strange, while "halogenated double bonds" strained credulity. In 194 7 Rundle published a paper on "half-bonds" applied to electron deficient bonding (34). The proposed half-bonds used one orbital and one electron pair to form two different bonds. His work was expanded and continued in two subsequent papers (35, 36). This quantum-mechanical treatment led to the very important conclusion that "normal" structures which leave one or more low-energy orbitals vacant are unstable with respect to electron-deficient bonding with half-bonds in which all of the low-energy orbitals are used. In his view, the tendency to use all low-energy orbitals is the underlying principle of electron-deficient bonding. In 1954 Eberhardt, Crawford, and Lipscomb (11, 22) presented a similar treatment, though in greater detail, of the boron hydrides and coined the term The pitfalls of boron hydride structure theory are entertain ingly pointed out in this quotation from Eberhardt et al. ; "We have even ventured a few predictions, knowing that if we must join the ranks of boron-hydride predictors later proved wrong, we shall be in the best of company. " "three-center bond. " They pointed out that in a normal covalent bond two atoms use two orbitals, one centered on each atom, which interact to form one bonding and one antibonding orbital. If two electrons are present, they will fill the bonding orbital to produce a normal covalent bond. Similarly, three atoms may provide three orbitals which combine to form one bonding and two antibonding orbitals; two electrons may fill this bonding orbital to form a three- center bond. Again it is pointed out that for a given atom the forma tion of a three-center bond which uses all available orbitals results in a lower energy state, and hence is a more stable system, than does the formation of normal two-center bonds which leave a vacant orbital.