Boronic Acids E-Book

• Herausgeber: Wiley-VCH
• Kategorie: Wissenschaft und neue Technologien
• Sprache: Englisch

Seitenzahl: 981

Veröffentlichungsjahr: 2012

Beschreibung

Following the huge success of the first edition, which has become THE reference source for everyone working in the field, this long-awaited, completely updated edition features almost 50% new content.
The world-renowned chemist Prof Dennis Hall is joined by a select group of top authors to cover all modern aspects of boronic acid derivatives in one comprehensive handbook. The experimental procedures described make for extremely useful reading.
From the reviews of the first edition: "...deserves to be on the bookshelf of all synthetic chemists, whether in discovery or process chemistry."

Leseprobe

Contents

Cover

Related Titles

Title Page

Copyright

Foreword

Preface

List of Contributors

Chapter 1: Structure, Properties, and Preparation of Boronic Acid Derivatives: Overview of Their Reactions and Applications

1.1 Introduction and Historical Background

1.2 Structure and Properties of Boronic Acid Derivatives

1.3 Preparation of Boronic Acids and Their Esters

1.4 Isolation and Characterization

1.5 Overview of the Reactions of Boronic Acid Derivatives

1.6 Overview of Other Applications of Boronic Acid Derivatives

References

Chapter 2: Metal-Catalyzed Borylation of C−H and C−Halogen Bonds of Alkanes, Alkenes, and Arenes for the Synthesis of Boronic Esters

2.1 Introduction

2.2 Borylation of Halides and Triflates via Coupling of H−B and B−B Compounds

2.3 Borylation via C−H Activation

2.4 Catalytic Cycle

2.5 Summary

References

Chapter 3: Transition Metal-Catalyzed Element-Boryl Additions to Unsaturated Organic Compounds

3.1 Introduction

3.2 Diboration

3.3 Silaboration

3.4 Carboboration

3.5 Miscellaneous Element-Boryl Additions

3.6 Conclusion

References

Chapter 4: The Contemporary Suzuki–Miyaura Reaction

4.1 Introduction

4.2 Developments Made in the Coupling of Nontrivial Substrates

4.3 Asymmetric Suzuki–Miyaura Cross-Couplings

4.4 Iterative Suzuki–Miyaura Cross-Couplings

4.5 Conclusions and Future Outlook

References

Chapter 5: Rhodium- and Palladium-Catalyzed Asymmetric Conjugate Additions of Organoboronic Acids

5.1 Introduction

5.2 Rh-Catalyzed Enantioselective Conjugate Addition of Organoboron Reagents

5.3 Pd-Catalyzed Enantioselective Conjugate Addition of Organoboron Reagents

5.4 Conclusions

References

Chapter 6: Recent Advances in Chan–Lam Coupling Reaction: Copper-Promoted C–Heteroatom Bond Cross-Coupling Reactions with Boronic Acids and Derivatives

6.1 General Introduction

6.2 C−O Cross-Coupling with Arylboronic Acids

6.3 C−N Cross-Coupling with Arylboronic Acids

6.4 Substrate Selectivity and Reactivity in Chan–Lam Cross-Coupling Reaction

6.5 C−N and C−O Cross-Coupling with Alkenylboronic Acids

6.6 C−N and C−O Cross-Coupling with Boronic Acid Derivatives

6.7 C−S and C–Se/C–Te Cross-Coupling

6.8 Mechanistic Considerations

6.9 Other Organometalloids

6.10 Conclusion

6.11 Note Added in Proof

Acknowledgment

References

Chapter 7: Transition Metal-Catalyzed Desulfitative Coupling of Thioorganic Compounds with Boronic Acids

7.1 General Introduction

7.2 Boronic Acid-Thioorganic C−S Desulfitative Cross-Couplings Using Catalytic Nickel or Palladium

7.3 Thioorganic C−S Desulfitative Cross-Couplings Using Only Catalytic Copper

7.4 Miscellaneous

7.5 Conclusions

References

Chapter 8: Catalytic Additions of Allylic Boronates to Carbonyl and Imine Derivatives

8.1 Introduction

8.2 Additions to Aldehydes

8.3 Additions to Ketones

8.4 Additions to Imine Derivatives

8.5 Conclusions

References

Chapter 9: Recent Advances in Nucleophilic Addition Reactions of Organoboronic Acids and Their Derivatives to Unsaturated C–N Functionalities

9.1 Introduction

9.2 Recent Advances in the Petasis Borono-Mannich Reaction

9.3 Reactions of N-Acyliminium Ions with Organoboronic Acids and Their Derivatives

9.4 Advances in Metal-Catalyzed Additions with Organoboronic Acids and Their Derivatives

9.5 Conclusions

References

Chapter 10: Asymmetric Homologation of Boronic Esters with Lithiated Carbamates, Epoxides, and Aziridines

10.1 Introduction

10.2 Lithiated Primary Alkyl Carbamates for the Homologation of Boranes and Boronic Esters

10.3 Lithiated Secondary Carbamates for the Homologation of Boranes and Boronic Esters

10.4 Benzylic N-Linked Lithiated Carbamates for the Homologation of Trialkylboranes

10.5 Lithiated Epoxides for the Homologation of Boronic Esters

10.6 Lithiated Aziridines for the Homologation of Boronic Esters

10.7 Conclusions

References

Chapter 11: Organotrifluoroborates: Organoboron Reagents for the Twenty-First Century

11.1 Introduction

11.2 Synthetic Approaches to Organotrifluoroborates

11.3 Elaboration of Organotrifluoroborates via Transformations of Pendant Functional Groups

11.4 Transition Metal-Catalyzed Processes

11.5 Miscellaneous Reactions of Organotrifluoroborates

11.6 Carbon–Carbon Bond-Forming Reactions with Activated Electrophiles

11.7 Conclusions

References

Chapter 12: Borate and Boronic Acid Derivatives as Catalysts in Organic Synthesis

12.1 Introduction

12.2 Nonchiral Boron-Based Catalysis [1]

12.3 Chiral Boron-Based Catalysis

12.4 Conclusion

References

Chapter 13: Applications of Boronic Acids in Chemical Biology and Medicinal Chemistry

13.1 Introduction

13.2 Boronic Acids as Potential Medicinal Agents

13.3 Probes for Detecting Reactive Oxygen Species

13.4 MRI and PET Agents for in vivo Carbohydrate Imaging

13.5 Carbohydrate Biomarker Binding Agents and Sensors

13.6 Conclusions

References

Chapter 14: Boronic Acids in Materials Chemistry

14.1 Introduction

14.2 Linear Boronate-Linked Materials

14.3 Macrocycles and Cages

14.4 Networks

14.5 Summary and Outlook

References

Index

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Foreword

Hydroboration, discovered in 1956, has made organoboranes readily available. This discovery opened the gate to a new continent for the chemical community to explore, develop, and exploit. Mainly in the late-1960s and 1970s, many novel types of carbon–carbon and carbon–heteroatom bond forming reactions of organoboranes were discovered and developed for use in organic synthesis. The palladium-catalyzed cross-coupling of organoboron compounds with organic electrophiles such as organic halides in the presence of base was developed in 1979. Over the past 30 years or so, transition metal-catalyzed cross-coupling reactions of boronic acid derivatives have emerged as one of the most important and widely used organometallic reactions for carbon–carbon and carbon–heteroatom bond formation, and is now regarded as an integral part of any synthetic route toward building complex organic chemicals. The coupling reaction has many advantages: the reactants are readily available, nontoxic, and air- and water-stable, and they react under mild conditions and are amenable to a variety of reaction conditions, including the use of aqueous solvents. Moreover, the inorganic boron byproduct can be easily removed after the reaction. Most important of all, the coupling proceeds with high regio- and stereoselectivity, and is little affected by steric hindrance. The process does not affect other functional groups in the molecule and can thus be used in onepot strategies. In addition, the reaction has proved to be extremely versatile. Consequently, these coupling reactions have been actively utilized not only in academic laboratories but also in industrial processes, like in pharmaceutical and agrochemical industries as well as other industries for the production of liquid crystals and organic LEDs.

Today, the use of boronic acid derivatives and their applications continue to evolve with many new findings reported during the past decade. For example, new reactions, catalysts and ligands have been developed. Increasingly, industry is seeking to use more environment-friendly processes. These often require ingenious solutions to which Suzuki coupling is well suited. We can expect to see many more interesting versions of the coupling reactions and other applications of boronic acids in the future.

Hokkaido University, Sapporo, Japan                                                                      Akira Suzuki

Preface

From the first isolation of a boronic acid by Frankland in 1860 to the report of their palladium-catalyzed cross-coupling with carbon halides by Suzuki in 1979, advances in the chemistry and biology of boronic acids have been few and far between. The early 1980s announced a drastic turn. In the past two decades alone, numerous breakthroughs have been reported. From Miyaura's discovery of rhodium-catalyzed couplings to alkenes and aldehydes to the commercialization of Velcade™, the first boronic acid drug used in human therapy, new reactions and applications of boronic acids have been reported at a spectacular rate. As seen in Figure P.1, the number of publications focused on boronic acid derivatives has increased exponentially, elevating boronic acids to a new status, that of an essential class of organic compounds. The attribution of the 2010 Chemistry Nobel Prize for palladium-catalyzed cross-coupling reactions, shared by Professor Akira Suzuki, cements the importance of boronic acids in this revolutionary class of C−C bond forming processes.

Figure P.1 Number of publications focused on boronic acids over time (note that only those publications carrying the word “boronic” in their titles were included). Hatched bars indicate the volume of publications reported after publication of the first edition of Boronic Acids.

This sudden rise in the usefulness and popularity of boronic acids necessitated a comprehensive book on their synthetic and biological applications. In just a few years working in the field of boronic acid chemistry, I had quickly come to regret the absence of a specialized book on this topic. Thus, I could not turn down an opportunity to help fulfill this need and lead such a project that led to the first edition of Boronic Acids in 2005. I was most fortunate to assemble a select group of experts, who literally included legends in the field. The successful result of this project was a popular handbook containing 13 chapters that covered all modern aspects of boronic acid derivatives. All efforts were made to achieve a comprehensive coverage of the field, with particular emphasis on topics of great interest to a large audience of synthetic organic, organometallic, and medicinal chemists. A quick look at Figure P.1 is sufficient to justify the need for an expanded, two-volume second edition only 6 years later. The recent period of 2005–2009 following the publication of the first edition of Boronic Acids shows a continuous, exponential burst of research activity around boronic acids, with the year 2010 showing no signs of stagnation. Clearly, it has become difficult to keep up with the literature on boronic acids, and it is anticipated that the second edition of Boronic Acids will be of invaluable assistance. In the past 5 years alone, impressive new advances have been made in the use of boronic acids in molecular recognition, chemical biology, materials science, and catalysis. Compared to the first edition, this second edition consists of chapters containing entirely new material, replacing previous chapters that have seen less advances in recent years, while other chapters have been completely revised and updated.

Our understanding of the structure and properties of boronic acids, their important ester derivatives, and other parent compounds such as trifluoroborate salts, is described in Chapter 1. In the past, the limited number of methods for the preparation of boronic acid derivatives had long impeded their use as synthetic reagents. This has changed drastically, and Chapter 1 describes modern methods for the preparation of all types of boronic acid derivatives, including several useful tables of examples. It also provides an overview of their synthetic, biological, and medicinal applications. One of the latest advances in the preparation of boronic acids, the use of transition metal-catalyzed borylation of C−X and C−H bonds, is discussed in Chapter 2. In the same manner, Chapter 3 describes metalloborylation reactions of unsaturated compounds. Much has happened in the development of new conditions and catalysts to expand the scope of transition metal-catalyzed C−C bond formation processes using boronic acids. Chapter 4 describes the most recent advances in the Suzuki–Miyaura cross-coupling reaction. A few years ago, rhodium(I) complexes were found to catalyze the addition of boronic acids to enones and aldehydes. These discoveries have now flourished into highly efficient catalytic enantioselective processes that can afford functionalized products with over 99% optical purity. These impressive advances are reviewed in Chapter 5. The copper-catalyzed coupling of boronic acids with heteroatom functionalities, such as phenols, amines, and amides, is yet another recent synthetic application that has contributed to the emergence of boronic acids as a popular class of reagents. This new and useful process, described in Chapter 6, has already become firmly established in the synthesis of natural products and in medicinal chemistry research. Chapter 7 describes the Liebeskind–Srogl cross-coupling reaction between boronic acids and thioorganic compounds. Already a workhorse in the synthesis of polypropionate compounds, the addition of allylboronates to carbonyl compounds and imine derivatives is still getting increasing attention as a result of new modes of catalytic activation highlighted in Chapter 8. The important discovery that boronic acids add to imine derivatives and iminium ions, even in a three-component fashion, has been exploited in a number of synthetic applications and progress in this area is reviewed in Chapter 9. Described in Chapter 10 is a new twist to the seminal Matteson homologation of boronic esters, using novel reagents such as lithiated carbamates, epoxides, and aziridines. Chapter 11 presents an overview of the synthetic applications of organotrifluoroborates, boronic acid derivatives that have become a very popular class of stable and efficient “go-to” reagents in cross-coupling chemistry. Boronic acids and several of their ester derivatives can serve as stable and mild Lewis acids, and this unique property has inspired the development of catalysts for several reaction processes, including asymmetric transformations; this topic is reviewed in Chapter 12. Boronic acids have long been known to bind and inhibit the action of certain classes of proteolytic enzymes. This important topic, as well as the applications of boronic acids in chemical biology, is discussed in Chapter 13 along with other emerging therapeutic applications. Finally, Chapter 14 presents an impressive overview of all the recently uncovered potential of boronic acids in materials science. From the rich contents of this book, it is clear that the spectacular rise of boronic acids as a class of compounds may have just begun. It is hoped that the second edition will contribute to generating more work and continue to attract more researchers to the field.

The success of a book project relies heavily on the involvement of several dedicated individuals. I would like to thank all authors and coauthors who have generously agreed to contribute a chapter. Their expertise and professionalism were invaluable assets to this ambitious project. Grateful acknowledgements are also offered to the Wiley-VCH editorial staff, in particular to Elke Maase and Renate Doetzer in the first edition, and Bernadette Gmeiner for the newer, second edition. For their valued support in various stages of editing this book I am also indebted to thanking Jack Lee, Jin-Yong Lu, Ho-Yan Sun, Hongchao Zheng, and Nitin Vashisht.

Edmonton (Alberta, Canada)                                                                                      Dennis Hall

May 2011

List of Contributors

Varinder K. Aggarwal University of Bristol School of Chemistry Cantock's Close Bristol BS8 1TS UK

Robert A. Batey University of Toronto Department of Chemistry Davenport Research Laboratories 80 St. George Street Toronto, Ontario, M5S 3H6 Canada

Guillaume Berthon-Gelloz Syngenta Crop Protection Münchwilen AG Schaffhauserstrasse 4332 Stein Switzerland

Tim G. Elford Department of Chemistry University of Alberta Edmonton, Alberta, T6G 2G2 Canada

Ethel C. Garnier-Amblard Emory University Sanford S. Atwood Chemistry Center 1515 Dickey Drive Atlanta, GA 30322 USA

Dennis G. Hall 4-010 Centennial Centre for Interdisciplinary Science Department of Chemistry University of Alberta Edmonton, Alberta, T6G 2G2 Canada

Tamio Hayashi Kyoto University Graduate School of Science Department of Chemistry Sakyo Kyoto 606-8502 Japan

Tatsuo Ishiyama Hokkaido University Graduate School of Engineering Division of Chemical Process Engineering Kita 13, Nishi 8 Sapporo, Hokkaido 060-8628 Japan

Ludivine Jean-Gérard University of Pennsylvania Department of Chemistry 231 South 34th Street Philadelphia, PA 19104-6323 USA

Patrick Y.S. Lam Bristol Myers Squibb Pharmaceutical Co. Discovery Chemistry Princeton, NJ 08543-5400 USA

John J. Lavigne University of South Carolina Department of Chemistry and Biochemistry 631 Sumter Street Columbia, SC 29208 USA

Lanny S. Liebeskind Emory University Sanford S. Atwood Chemistry Center 1515 Dickey Drive Atlanta, GA 30322 USA

Jie Liu University of South Carolina Department of Chemistry and Biochemistry 631 Sumter Street Columbia, SC 29208 USA

Norio Miyaura Hokkaido University Graduate School of Engineering Division of Chemical Process Engineering Kita 13, Nishi 8 Sapporo, Hokkaido 060-8628 Japan

Gary A. Molander University of Pennsylvania Department of Chemistry 231 South 34th Street Philadelphia, PA 19104-6323 USA

Nanting Ni Georgia State University Department of Chemistry and Center for Biotechnology and Drug Design Atlanta, GA 30302-4098 USA

Toshimichi Ohmura Kyoto University Graduate School of Engineering Department of Synthetic Chemistry and Biological Chemistry Katsura, Nishikyo-ku Kyoto 615-8510 Japan

Michael G. Organ York University Department of Chemistry 4700 Keele Street Toronto, Ontario, M3J 1P3 Canada

Joshua N. Payette The University of Chicago Department of Chemistry 5735 South Ellis Avenue Chicago, IL 60637 USA

Jennifer X. Qiao Bristol Myers Squibb Pharmaceutical Co. Discovery Chemistry Princeton, NJ 08543-5400 USA

Timothy R. Ramadhar University of Toronto Department of Chemistry Davenport Research Laboratories 80 St. George Street Toronto, Ontario, M5S 3H6 Canada

Michinori Suginome Kyoto University Graduate School of Engineering Department of Synthetic Chemistry and Biological Chemistry Katsura, Nishikyo-ku Kyoto 615-8510 Japan

Cory Valente Northwestern University Department of Chemistry 2145 Sheridan Road Evanston, IL 60208 USA

Binghe Wang Georgia State University Department of Chemistry and Center for Biotechnology and Drug Design Atlanta, GA 30302-4098 USA

Matthew P. Webster University of Illinois School of Chemical Sciences Chemistry Roger Adams Lab, Room 237, Box 90-5 600 S Mathews Urbana, IL 61801 USA

Hisashi Yamamoto The University of Chicago Department of Chemistry 5735 South Ellis Avenue Chicago, IL 60637 USA

Chapter 1

Structure, Properties, and Preparation of Boronic Acid Derivatives

Overview of Their Reactions and Applications

Dennis G. Hall

1.1 Introduction and Historical Background

Structurally, boronic acids are trivalent boron-containing organic compounds that possess one carbon-based substituent (i.e., a C−B bond) and two hydroxyl groups to fill the remaining valences on the boron atom (Figure 1.1). With only six valence electrons and a consequent deficiency of two electrons, the sp2-hybridized boron atom possesses a vacant p-orbital. This low-energy orbital is orthogonal to the three substituents, which are oriented in a trigonal planar geometry. Unlike carboxylic acids, their carbon analogues, boronic acids, are not found in nature. These abiotic compounds are derived synthetically from primary sources of boron such as boric acid, which is made by the acidification of borax with carbon dioxide. Borate esters, one of the key precursors of boronic acid derivatives, are made by simple dehydration of boric acid with alcohols. The first preparation and isolation of a boronic acid was reported by Frankland in 1860 [1]. By treating diethylzinc with triethylborate, the highly air-sensitive triethylborane was obtained, and its slow oxidation in ambient air eventually provided ethylboronic acid. Boronic acids are the products of a twofold oxidation of boranes. Their stability to atmospheric oxidation is considerably superior to that of borinic acids, which result from the first oxidation of boranes. The product of a third oxidation of boranes, boric acid, is a very stable and relatively benign compound to humans (Section 1.2.2.3).

Figure 1.1 Oxygen-containing organoboron compounds.

Their unique properties and reactivity as mild organic Lewis acids, coupled with their stability and ease of handling, are what make boronic acids a particularly attractive class of synthetic intermediates. Moreover, because of their low toxicity and their ultimate degradation into boric acid, boronic acids can be regarded as “green” (environment-friendly) compounds. They are solids, and tend to exist as mixtures of oligomeric anhydrides, in particular the cyclic six-membered boroxines (Figure 1.1). For this reason and other considerations outlined later in this chapter, the corresponding boronic esters are often preferred as synthetic intermediates. Although other classes of organoboron compounds have found tremendous utility in organic synthesis, this book focuses on the most recent applications of the convenient boronic acid derivatives. For a comprehensive description of the properties and reactivity of other classes of organoboron compounds, interested readers may refer to a selection of excellent monographs and reviews by Brown [2], Matteson [3], and others [4–8]. In the past two decades, the status of boronic acids in chemistry has gone from that of peculiar and rather neglected compounds to that of a prime class of synthetic intermediates in their own right. The attribution of the 2010 Chemistry Nobel Prize for palladium-catalyzed cross-coupling reactions to Professor Akira Suzuki and other pioneers recognized the great importance of boronic acids in this revolutionary class of C−C bond forming processes. In the past 5 years, impressive advances have been made in the use of boronic acids in molecular recognition, materials science, and catalysis. The approval of the anticancer agent Velcade®, the first boronic acid-containing drug to be commercialized (Section 1.6.5), further confirms the growing status of boronic acids as an important class of compounds in chemistry and medicine. This chapter describes the structural and physicochemical properties of boronic acids and their many derivatives, as well as modern methods for their preparation. A brief overview of their synthetic and biological applications is presented, with an emphasis on topics that are not covered in other chapters of this book.

1.2 Structure and Properties of Boronic Acid Derivatives

1.2.1 General Types and Nomenclature of Boronic Acid Derivatives

The reactivity and properties of boronic acids highly depend upon the nature of their single variable substituent, more specifically, on the type of carbon group (R, Figure 1.1) directly bonded to boron. In the same customary way employed for other functional groups, it is convenient to classify boronic acids into subtypes such as alkyl-, alkenyl-, alkynyl-, and arylboronic acids.

When treated as an independent substituent, the prefix borono is employed to name the boronyl group (e.g., 3-boronoacrolein). For cyclic derivatives such as boronic esters, the IUPAC RB-1-1 rules for small heterocycles (i.e., the Hantzsch–Widman system) are employed along with the prefix “boro.” Thus, saturated five- and six-membered cyclic boronic esters are, respectively, named as dioxaborolanes and dioxaborinanes. For example, the formal name of the pinacol ester of phenylboronic acid is 2-phenyl-4,4,5,5-tetramethyl-1,3,2-dioxaborolane. The corresponding nitrogen analogues are called diazaborolidines and diazaborinanes, and the mixed nitrogen–oxygen heterocycles are denoted by the prefix oxaza. Unsaturated heterocycles wherein the R group and the boron atom are part of the same ring are named as boroles.

1.2.2 Boronic Acids

1.2.2.1 Structure and Bonding

The X-ray crystal structure of phenylboronic acid (1, Figure 1.2) was reported in 1977 by Rettig and Trotter [9]. The crystals are orthorhombic, and each asymmetric unit was found to consist of two distinct molecules, bound together through a pair of O−H−O hydrogen bonds (Figure 1.3a and 1.3b). The CBO2 plane is quite coplanar with the benzene ring, with a respective twist around the C−B bond of 6.6° and 21.4° for the two independent molecules of PhB(OH)2. Each dimeric ensemble is also linked with hydrogen bonds to four other similar units to give an infinite array of layers (Figure 1.3c). The X-ray crystallographic analysis of other arylboronic acids like p-methoxyphenyl boronic acid 2 [10] and 4-carboxy-2-nitrophenylboronic acid (3, Figure 1.2) [11] is consistent with this pattern. The structures of heterocyclic boronic acids such as 2-bromo- and 2-chloro 5-pyridylboronic acids (4 and 5) were reported [12]. Although the boronate group has a trigonal geometry and is fairly coplanar with the benzene ring in structures 1, 2, 4, and 5, it is almost perpendicular to the ring in structure 3. This observation is likely due to a combination of two factors: minimization of steric strain with the ortho-nitro group, and because of a possible interaction between one oxygen of the nitro group and the trigonal boron atom. Based on the structural behavior of phenylboronic acid and its propensity to form hydrogen-bonded dimers, diamond-like porous solids were designed and prepared by the crystallization of tetrahedral-shaped tetraboronic acid 6 (Figure 1.2) [13]. With a range of approximately 1.55–1.59 Å, the C−B bond of boronic acids and esters is slightly longer than typical C−C single bonds (Table 1.1). The average C−B bond energy is also slightly smaller than that of C−C bonds (323 versus 358 kJ/mol) [14]. Consistent with strong B−O bonds, the B−O distances of tricoordinate boronic acids such as phenylboronic acid are fairly short, and lie in the range of 1.35–1.38 Å (Table 1.1). These values are slightly larger than those observed in boronic esters. For example, the B−O bond distances observed in the X-ray crystallographic structures of the trityloxymethyl pinacolate boronic esters (e.g., 7 in Figure 1.2) are in the range of 1.31–1.35 Å (Table 1.1), and the dioxaborolane unit of these derivatives is nearly planar [15]. The X-ray crystallographic structure of cyclic hemiester 8 (Figure 1.2) was described [16]. Like phenylboronic acid, this benzoxaborole also crystallizes as a hydrogen-bonded dimer, however without the extended network due to the absence of a second hydroxyl group. The cyclic nature of this derivative induces a slight deviation from planarity for the tricoordinate boronate unit, as well as a distortion of the bond angles. The endocyclic B−O bond in 8 is slightly longer than the B−OH bond. This observation was attributed to the geometrical constraints of the ring, which prevents effective lone pair conjugation between the endocyclic oxygen and the vacant orbital of boron. The unique properties and reactivity of benzoxaboroles along with their preparation were recently reviewed [17].

Figure 1.2 Boronic acid derivatives analyzed by X-ray crystallography.

Figure 1.3 Representations of the X-ray crystallographic structure of phenylboronic acid. (a) ORTEP view of a dimeric unit. (b) Structure of the dimeric unit showing hydrogen bonds. (c) Structure of the extended hydrogen-bonded network.

Table 1.1 Bond distances from X-ray crystallographic data for selected boronic acid derivatives (Figure 1.2)

In order to complete boron's octet, boronic acids and their esters may also coordinate basic molecules and exist as stable tetracoordinated adducts. For example, the X-ray crystallographic structure of the diethanolamine adduct of phenylboronic acid (9, Figure 1.2) [18] confirmed the transannular B−N bridge long suspected from other spectroscopic evidence such as NMR [19, 20]. This dative B−N bond has a length of 1.67 Å (Table 1.1), and it induces a strong Nδ+−Bδ− dipole that points away from the plane of the aryl ring. This effect was elegantly exploited in the design of a diboronate receptor for paraquat [21]. Chelated boronic ester 10 presents characteristics similar to that of 9 [22]. Trihydroxyborate salts of boronic acids are discrete, isolable derivatives that had not been characterized until recently [23]. The sodium salt of p-methoxyphenyl boronic acid 11 was recrystallized in water and its X-ray structural elucidation showed the borate unit in the expected hydrogen bonding network accompanied with the sodium cation coordinated with six molecules of water. In principle, the boron atom in tetrahedral complexes can be stereogenic if it is bonded to four different ligands. Hutton and coworkers recently reported the first example of one such optically pure complex stereogenic at boron only [24]. Stable complex 12 (Figure 1.4) was made through a chirality transfer process described in Section 1.2.3.6. When tetracoordinated such as in structures 9 –11 [23] (Figure 1.2), the B−O bond length of boronic acids and esters increases to about 1.43–1.48 Å, which is as much as 0.10 Å longer than the corresponding tricoordinate analogues (Table 1.1). These markedly longer B−O bonds are comparable to normal C−O ether bonds (1.43 Å). These comparisons further emphasize the considerable strength of B−O bonds in trigonal boronic acid derivatives. Not surprisingly, trigonal B−O bonds are much stronger than the average C−O bonds of ethers (519 versus 384 kJ/mol) [14]. This bond strength is believed to originate from the conjugation between the lone pairs on the oxygens and boron's vacant orbital, which confers partial double bond character to the B−O linkage. In fact, it was estimated that formation of tetrahedral adducts (e.g., with NH3) may result in a loss of as much as 50 kJ/mol of B−O bond energy compared to the tricoordinate boronate [25].

Figure 1.4B -Chiral tetrahedral boronate 12 and model compounds for boron hypercoordination.

In rare instances where geometrical factors allow it, boronic acid derivatives may become hypervalent. For example, the catechol ester 13 (Figure 1.4) was found by X-ray crystallographic analysis to be pentacoordinated in a highly symmetrical fashion as a result of the rigidly held ether groups, which are perfectly positioned to each donate lone pair electrons to both lobes of the vacant p-orbital of boron [26]. The boronyl group of this two electron–three atom center is planar, in a sp2 hybridization state, and the resulting structure possesses a slightly distorted trigonal bipyramidal geometry. According to DFT calculations, the bonding is weak and ionic in nature [26b]. The corresponding diamine 14, however, behaved quite differently and demonstrated coordination with only one of the two NMe2 groups [27].

Due to electronegativity differences (B = 2.05, C = 2.55) and notwithstanding the electronic deficiency of boron, which is compensated by the two electron-donating oxygen atoms (see above), the inductive effect of a boronate group should be that of a weak electron donor. The NMR alpha effect of a boronate group is in fact very small [28]. On the other hand, the deficient valency of boron and its size relatively similar to that of carbon have long raised the intriguing question of possible pi-bonding between carbon and boron in aryl- and alkenylboronic acids and esters [29]. NMR data and other evidence, such as UV and photoelectron spectroscopy and LCAO-MO calculations, suggest that B−C pi-conjugation occurs to a moderate extent in alkenylboranes [30–32], and is even smaller in the case of the considerably less acidic boronate derivatives. A thorough comparative study of NMR shift effects, in particular the deshielding of the beta-carbon, concluded to a certain degree of mesomeric pi-bonding in the case of boranes and catechol boronates [28]. For example, compared to analogous aliphatic boronates, the beta-carbons of a dialkyl alkenylboronate and the corresponding catechol ester are deshielded by 8.6 and 18.1 ppm, respectively. In all cases, the beta-carbon is more affected by the boronate substituent than the alpha-carbon, which is consistent with some contribution from the B−C pi-bonded form B to give resonance hybrid C (Figure 1.5). X-ray crystallography may also provide insights into the extent of B−C pi-bonding. The difference in B−C bond distances for arylboronic acids (Table 1.1) is significant enough to suggest a small degree of B−C pi-bonding. The B−C bond distance (1.588 Å) in the electron-poor boronic acid 3, which is incapable of pi-conjugation because it has its vacant p-orbital placed orthogonally to the pi-system of the phenyl ring, is expectedly longer than that of phenylboronic acid (1.568 Å). Interestingly, the B−C bond of 2 stands at 1.556 Å, suggesting only a minimal contribution from the mesomeric form E (Figure 1.5). On the other hand, the B−C bond distance of 1.613 Å in the diethanolamine adduct 9 (Table 1.1), where the boron vacant orbital is also incapacitated from B−C pi-bonding, is 0.045 Å longer than that of free phenylboronic acid 1. In so far as bond length data correlate with the degree of pi-bonding [33], this comparison is consistent with a small B−C pi-bonding effect in arylboronic acids and esters (i.e., hybrid form F in Figure 1.5). This view is further supported by chemical properties such as substituent effects on the acidity of arylboronic acids (see Section 1.3.8.3) and chemical shifts correlations [34]. Likewise, B−C pi-bonding is also present in alkenylboronic acids and esters, but this effect must be weak in comparison to the electron-withdrawing effect of a carbonyl or a carboxyl group. For instance, alkenylboronic esters do not readily act as Michael acceptors with organometallic reagents in the same way as the unsaturated carbonyl compounds do [35]. On the other hand, the formal electron-withdrawing behavior of the boronate group manifests itself in cycloadditions of dibutylethylene boronate with ethyldiazoacetate [36] and in Diels–Alder reactions where it provides cycloadducts with dienes like cyclopentadiene [37] and cyclohexadiene, albeit only at elevated temperatures (about 130 and 200 °C, respectively) [38, 39]. The higher reactivity of ethylene boronates as dienophiles compared to ethylene has been rationalized by MO calculations [29], but their reactivity stands far from that of acrylates in the same cycloadditions. In fact, more recent high-level calculations suggest that the reactivity of alkenylboronates may be mainly due to a three-atom–two-electron center stabilization of the transition state rather than a true LUMO-lowering electron-withdrawing mesomeric effect from the boronate substituent [40]. Another evidence for the rather weak electron-withdrawing character of boronic esters comes from their modest stabilizing effect on boronyl-substituted carbanions, where their effect has been compared to that of a phenyl group (see Section 1.3.8.3).

Figure 1.5 Limit mesomeric forms involving B−C pi-overlap.

1.2.2.2 Physical Properties and Handling

Most boronic acids exist as white crystalline solids that can be handled in air without special precautions. At the ambient temperature, boronic acids are chemically stable and most display shelf stability for long periods of time (Section 1.2.2.5). Alkyl-substituted and some heteroaromatic boronic acids, however, were shown to have a limited shelf stability under aerobic conditions [41]. Boronic acids normally do not tend to disproportionate into their corresponding borinic acid and boric acid even at high temperatures. To minimize atmospheric oxidation and autoxidation, however, they should be stored under an inert atmosphere. When dehydrated, either with a water-trapping agent or through coevaporation or high vacuum, boronic acids form cyclic and linear oligomeric anhydrides such as the trimeric boroxines already mentioned (Figure 1.1). Fortunately, this behavior is usually inconsequential when boronic acids are employed as synthetic intermediates. Many of their most useful reactions (Section 1.5), including the Suzuki-Miyaura cross-coupling, proceed regardless of the hydrated state (i.e., free boronic acid or anhydride). Anhydride formation, however, may complicate analysis, quantitation, and characterization efforts (Section 1.4.3). Furthermore, upon exposure to air, dry samples of boronic acids may be prone to decompose rapidly, and it has been proposed that boronic anhydrides may be initiators of the autoxidation process [42]. For this reason, it is often better to store boronic acids in a slightly moist state. Presumably, coordination of water or hydroxide ions to boron protects boronic acids from the action of oxygen [42, 43]. Incidentally, commercial samples tend to contain a small percentage of water that may help in their long-term preservation. Due to their facile dehydration, boronic acids tend to provide somewhat unreliable values of melting points (Section 1.4.3.1). This inconvenience and the other above-mentioned problems associated with anhydride formation explain in large part the popularity of boronic esters and other derivatives as surrogates of boronic acids (Section 1.2.3.2).

The Lewis acidity of boron in boronic acids and the hydrogen bond donor capability of their hydroxyl groups combine to lend a polar character to most of these compounds. Although the polarity of the boronic acid head can be mitigated by a relatively hydrophobic tail as the boron substituent, most small boronic acids are amphiphilic. Phenylboronic acid, for instance, was found to have a benzene–water partition ratio of 6 [44]. The partial solubility of boronic acids in both neutral water and polar organic solvents often complicates isolation and purification efforts (Section 1.4). Evidently, boronic acids are more water soluble in their ionized form in high-pH aqueous solutions and can be extracted more readily into organic solvents from aqueous solutions of low pH (see Section 1.2.2.4).

1.2.2.3 Safety Considerations

As evidenced by their application in medicine (Chapter 13), most boronic acids present no particular toxicity compared to other organic compounds [45]. Small water-soluble boronic acids demonstrate low toxicity levels, and are excreted largely unchanged by the kidney [46]. Larger fat-soluble boronic acids were found to be moderately toxic [46–48]. At high doses, boronic acids may interact promiscuously with nucleophilic enzymes and complex weakly to biological diols (Section 1.2.3.2.3). Boronic acids present no particular environmental threat, and the ultimate fate of all boronic acids in air and aqueous media is their slow oxidation into boric acid. The latter is a relatively innocuous compound, and may be toxic only under high daily doses [49]. A single acute ingestion of boric acid does not even pose a threatening poisoning effect to humans [50] unless it is accompanied by other health malfunctions such as dehydration [51].

1.2.2.4 Acidic Character

By virtue of their deficient valence, boronic acids possess a vacant p-orbital. This characteristic confers them unique properties as a mild class of organic Lewis acids capable of coordinating basic molecules. When doing so, the resulting tetrahedral adducts acquire a carbon-like configuration. Thus, despite the presence of two hydroxyl groups, the acidic character of most boronic acids is not that of a Brnsted acid (i.e., oxyacid) (Equation 1.1, Figure 1.6) but usually that of a Lewis acid (Equation 1.2). When coordinated with an anionic ligand, the resulting negative charge is formally drawn on the boron atom, but it is in fact spread out on the three heteroatoms.

Figure 1.6 Ionization equilibrium of boronic acids in water.

1.2.2.4.1 Complexation Equilibrium in Water and Structure of the Boronate Anion

Boronic acids are more soluble in aqueous solutions of high pH (>8). Although the acidic character of boronic acids in water had been known for several decades, it is only in 1959 that the structure of the boronate ion, the conjugate base, was elucidated. In their classical paper on polyol complexes of boronic acids [52], Lorand and Edwards demonstrated that the trivalent neutral form, likely hydrated, is in equilibrium with the anionic tetrahedral species (Equation 1.2, Figure 1.6) and not with the structurally related Brnsted base (i.e., the trivalent ion shown in Equation 1.1). The first X-ray crystallographic structure of a trihydroxyboronate salt has been reported recently (11 in Figure 1.2) [23]. It is this ability to ionize water and form hydronium ions by “indirect” proton transfer that characterizes the acidity of most boronic acids in water. Hence, the most acidic boronic acids possess the most electrophilic boron atom that can best form and stabilize a hydroxyboronate anion. The acidity of boronic acids in water has been measured using electrochemical methods as early as the 1930s [53–55]. Values of pKa are now measured more conveniently by UV spectrophotometry [56] and NMR spectroscopy. Phenylboronic acid, with a pKa value of 8.9 in water, has an acidity comparable to a phenol (Table 1.2). It is slightly more acidic than boric acid (pKa 9.2). With the pKa values as shown in Table 1.2, the relative order of acidity for the different types of boronic acids is aryl > alkyl. More values can be found elsewhere [57]. For para-monosubstituted aromatic boronic acids, the relationship between the pKa and the electronic nature of the substituent can be described with a Hammet plot [57]. Bulky substituents proximal to the boronyl group can decrease the acid strength due to steric inhibition in the formation of the tetrahedral boronate ion. For example, ortho-tolylboronic acid is slightly less acidic than its para-isomer (pKa 9.7 versus 9.3, Table 1.2) [62]. This difference was explained in terms of F-strain in the resulting ion (Equation 1.7, Figure 1.7) [67]. As expected, the presence of electron-withdrawing substituents in the aryl group of arylboronic acids increases the acid strength by a fairly significant measure [53, 55, 60, 68]. For example, the highly electron-poor 3-methoxycarbonyl-5-nitrophenyl boronic acid was attributed a pKa value of 6.9 [63]. Exceptionally, ortho-nitrobenzeneboronic acid [61] is much less acidic than its para-isomer [60] (pKa 9.2 versus 7.1, Table 1.2) presumably due to internal coordination of one of the nitro oxygens that prevents the complexation of a hydroxyl anion [55]. Perhaps one of the most acidic of all known boronic acids, with a pKa of approximately 4.0, 3-pyridylboronic acid 15 exists mainly as a zwitterion in water (Equation 1.7, Figure 1.7) [64]. Similarly, arylboronic acids of type 16 (Equation 1.7), which benefit from anchimeric participation of the ortho-dialkylaminomethyl group, display a relatively low value of pKa of about 5.2 [66]. In this case, the actual first pKa is that of ammonium ion deprotonation and formation of the putative tetrahedral B−N ate adduct 16. The latter form was shown to exist in organic solvents, but in water and other hydroxylic solvents, complex 17 forms through a water-insertion mechanism [69]. The application of boronic acids of type 16 in the aqueous recognition of saccharides is briefly discussed in Chapter 13. Fluoride ions also form strong dative bonds with boron, and it has been noted long ago that boronic acids dissolved in aqueous solutions of hydrofluoric acid are very difficult to extract into organic solvents unless the fluoride is precipitated out [70].

Table 1.2 Ionization constant (pKa) for selected boronic acids.

Boronic acid, RB(OH)2pKaReferenceBoric acid, B(OH)39.0[58]Methyl10.4[58]Phenyl8.9[59]3,5-Dichlorophenyl7.4[59]3,5-Bis(trifluoromethyl)phenyl7.2[59]2-Methoxyphenyl9.0[57]3-Methoxyphenyl8.7[59]4-Methoxyphenyl9.3[60]4-Carboxyphenyl8.4[56]2-Nitrophenyl9.2[61]4-Nitrophenyl7.1[60]4-Bromophenyl8.6[59]4-Fluorophenyl9.1[59]2-Methylphenyl9.7[62]3-Methylphenyl9.0[62]4-Methylphenyl9.3[62]3,5-Dimethylphenyl9.1[59]3-Methoxycarbonyl-5-nitrophenyl6.9[63]2-Fluoro-5-nitrophenyl6.0[57]3-Pyridyl 154.0, 8.2[64]3-Benzyl-3-pyridylium4.2[57]8-Quinolinyl4.0, 10[65]2-(R1R2NCH2)phenyl (e.g., 16)5.2–5.8[66]

Figure 1.7 Ionization equilibrium of special boronic acids.

Boronic acids display Brnsted acidity (cf. Equation 1.1, Figure 1.6) only in exceptional cases where the formation of a tetrahedral boronate adduct is highly unfavorable. For example, coordination of hydroxide ion to boron in heterocyclic boronic acid derivative 18, to form 19B, would break the partial aromatic character of the central ring (Equation 1.6, Figure 1.7). Indeed, based on NMR and UV spectroscopic evidence, it was suggested that 18 acts as a Brnsted acid in water and forms conjugate base 19A through direct proton transfer [71]. A small number of other boronic acids are suspected of behaving as Brnsted acids due to the same reasons [72].

1.2.2.4.2 Bimolecular Lewis Acid–Base Complexation under Nonaqueous Conditions

As evidenced by the high pH required in the formation of boronate anions, boronic acids and most dialkyl esters are weak Lewis acids. This behavior is in sharp contrast with trialkylboranes, which form strong adducts with phosphines, amines, and other Lewis bases [73]. Apart from the formation of boronate anions, discussed in the previous section, very few examples of stable intermolecular acid–base adducts of boronic acids (esters) exist. It has been known for a long time that aliphatic amines and pyridine can form complexes in a 1 : 3 amine:boronic acid stoichiometry [74]. Combustion analyses of these air-stable solids suggested that two molecules of water are lost in the process, which led the authors to propose structure 20 (Equation 1.7, Figure 1.8). Much later, Snyder et al. used IR spectroscopy to demonstrate that these 1 : 3 complexes rather involved the fully dehydrated boroxine 21 [75]. Boronic esters are generally weak Lewis acids but catechol boronates are quite acidic, and provided that cooperative effects are exploited, bimolecular complexes with fluoride anions and amines have been reported [76–78]. The B−F bond strength is a key factor in these complexes as other halide salts do not form similar adducts. As suggested by NMR spectroscopic studies, an ortho-phenyldiboronic ester 22 showed cooperative binding of two amine molecules in putative complex 24 (Equation 1.8, Figure 1.8) [79]. Other diboronate receptors were found to bind to diamines selectively using the two boron centers for B−N coordination [80–82]. Catechol esters and other cyclic five-membered boronic esters with sp2 centers are more acidic as complexation to form a tetrahedral boron atom relieves strain. The concept of strain has recently been exploited in the design of a receptor with photoswitchable Lewis acidity [83]. Pyridine complexation studies by NMR spectroscopy showed that bisthiophene boronate receptor 25 is more acidic in its closed cross-conjugated form 26 compared to the less strained, open form 25 (Equation 1.9).

Figure 1.8 Bimolecular Lewis acid–base complexes with boronic esters.

1.2.2.5 Chemical Stability

1.2.2.5.1 Ligand Exchange and Disproportionation

Several favorable factors contribute to the stability of boronic acids and their esters. Substitution of the carbon-containing group of boronic acids with other substituents is a slow process, and B−C/B−O bond metatheses to give the corresponding disproportionation products (trialkylborane, borinic acid, or boric acid) are thermodynamically unfavorable [25]. This redox disproportionation is rather used to transform borinic esters into boronic esters [84]. Similarly, thermodynamic considerations make the exchange of the hydroxyl substituents of boronic acids with other ligands quite unfavorable. Substitution with most alcohols or diols to form boronic esters usually requires dehydration techniques in order to drive the reaction forward (Section 1.2.3.2.1). In general, from the B−X bond energy values of all possible boronic acid derivatives (RBX2), it can be said that free boronic acids remain unchanged when dissolved in solutions containing other potential anionic ligands [24]. The only type of B−X bond stronger than a B−O bond is the B−F bond. Chemical methods to accomplish this type of exchange and other B−O bond derivatizations are described in Sections 1.2.3.7 and 1.2.3.8.

1.2.2.5.2 Atmospheric Oxidation

A significant thermodynamic drive for C−B bond oxidation results as a direct consequence of the large difference between B−O and B−C bond energies (Section 1.2.2.1). Heats of reaction for the oxidative cleavage of methylboronic acid with water and hydrogen peroxide are −112 and −345 kJ/mol, respectively [25]. Yet, fortunately for synthetic chemists, oxidative cleavage of the B−C bond of boronic acid derivatives with water or oxygen is a kinetically slow process, and most boronic acids can be manipulated in ambient air and are stable in water in a wide range of pH. This is particularly true for aryl- and alkenylboronic acids, and in general, samples of all types of boronic acids tend to be significantly more stable when moist (coordination of water to boron likely acts as a protection) (Section 1.2.2.2) [42, 43, 85]. Exceptionally, the highly electron-poor arylboronic acid 4-carboxy-2-nitrophenylboronic acid was reported to undergo slow oxidation to the corresponding phenol when left in aqueous basic solutions (pH 9) [11]. On the other hand, basic aqueous solutions of alkylboronate ions were claimed to be highly tolerant of air oxidation [42]. Free alkylboronic acids, however, are quite prone to a slow atmospheric oxidation and variable amounts of the corresponding alcohols may form readily when dried samples are left under ambient air. Likewise, solutions of arylboronic acids in tetrahydrofuran devoid of stabilizer may turn rapidly into the corresponding phenols. The propensity of alkylboronic acids to undergo autoxidation depends on the degree of substitution, with primary alkyl substituents being less reactive than the secondary and tertiary alkyl substituents, respectively [85]. More potent oxidants such as peroxides readily oxidize all types of boronic acids and their corresponding esters (Section 1.5.2.1). This propensity for oxidation must be kept in mind while handling boronic acids.

1.2.2.5.3 Protolytic Deboronation

Most types of boronic acids are highly resistant to protolysis of the C−B bond in neutral aqueous solutions even at high temperatures. For example, p-tolylboronic acid was recovered unchanged after 28 h in boiling water [86]. Aqueous protodeboronation can become problematic at higher temperatures; p-tolylboronic acid was completely deboronated to toluene after 6 h under pressure at 130–150 °C [86]. Deboronation of arylboronic acids can be effected quite readily in highly acidic or basic aqueous solutions [87]. In particular, ortho-substituted and especially electron-poor arylboronic acids are notorious for their propensity to protodeboronate under basic aqueous conditions, a process that can be exacerbated by exposure to light [64]. Consequently, competitive deboronation may plague some reactions employing boronic acids as reagents like the Suzuki–Miyaura cross-coupling reaction (Section 1.5.3.1), which requires basic conditions often at high temperatures. Under acidic aqueous conditions, however, the more electron-rich arylboronic acids tend to deboronate faster [87, 88]. For example, p-carboxyphenylboronic acid was found to be more tolerant than phenylboronic acid to the highly acidic conditions of ring nitration under fuming nitric acid and concentrated sulfuric acid [89]. Certain heteroaromatic boronic acids with the boronyl group next to the heteroatom (α-substituted) are notoriously prone to protodeboronation, but they can be stabilized as tetrahedral adducts (Section 1.2.3.3) [41, 90]. The effect of acid, temperature, and ring substitution of arylboronic acids on the kinetics of electrophilic protolytic deboronation with strong aqueous acid has been studied by Kuivila and Nahabedian [91]. A relatively complex behavior was found, and at least two possible pH-dependent mechanisms were proposed. In contrast to their behavior with aqueous acids, most arylboronic acids and esters appear to be very resistant to nonaqueous acids, as evidenced by their recovery from reaction processes using strong organic acids. For example, a phenolic methoxymethyl ether was deprotected with a 2 : 1 CH2Cl2/CF3CO2H (TFA) mixture that left intact a pinacol boronic ester functionality [92]. Likewise, free arylboronic acids have been shown to tolerate, at ambient temperature, similar organic acid conditions that effect cleavage of t-butoxycarbonyl groups (Equation 1.10) [93]. On the other hand, a report emphasized that arylboronic acids can be protodeboronated thermally without added acid by prolonged heating in refluxing ethereal solvents [94].

(1.10)

In contrast to arylboronic acids, early reports document the great stability of alkylboronic acids under aqueous acidic solutions. For example, a variety of simple alkylboronic acids were unaffected by prolonged heating in 40% aqueous HBr or HI [42]. Like arylboronic acids, however, deboronation is observed in hot basic aqueous solutions [85]. Alkenylboronic esters undergo protonolysis in refluxing AcOH [95], and alkynylboronic acids were reported to be quite unstable in basic aqueous solutions (Section 1.3.5).

All types of boronic acids can be protodeboronated by means of metal-promoted C−B bond cleavage, and these methods are described separately later in this chapter (Section 1.5.1).

1.2.3 Boronic Acid Derivatives

For the sake of convenience in their purification and characterization, boronic acids are often best handled as ester derivatives where the two hydroxyl groups are masked. On the other hand, transformation of the hydroxyl groups into other substituents such as halides or borate salts may also provide an increase in reactivity necessary for a number of synthetic applications. The next sections describe the most important classes of boronic acid derivatives.

1.2.3.1 Boroxines (Cyclic Anhydrides)

Boroxines are the cyclotrimeric anhydrides of boronic acids. Their properties and applications have been reviewed recently [96]. By virtue of boron's vacant orbital, boroxines are isoelectronic to benzene, but it is generally accepted that they possess little aromatic character [97]. Several theoretical and experimental studies have addressed the nature and structure of these derivatives [96]; in particular, the X-ray crystallographic analysis of triphenylboroxine confirmed that it is virtually flat [98]. Boroxines are easily produced by the simple dehydration of boronic acids, either thermally through azeotropic removal of water or by exhaustive drying over sulfuric acid or phosphorus pentoxide [42]. These compounds can be employed invariably as substrates in many of the same synthetic transformations known to affect boronic acids. Interest in the applications of boroxines as end products has increased in the past decade. Their use has been proposed as flame retardants [99] and as functional materials (see Chapter 14) [100]. The formation of boroxine cross-linkages has been employed as a means to immobilize blue light-emitting oligofluorene diboronic acids [101]. Samples of boroxines, which may also contain oligomeric acyclic analogues, were found to be sensitive to autoxidation when dried exhaustively (Sections 1.2.2.2 and 1.2.2.5.2). A study examined the thermodynamic parameters of boroxine formation in water (Equation 1.11) [102]. Using NMR spectroscopy, the reaction was found to be reversible at room temperature, and the equilibrium constants, relatively small ones, were found to be subject to substituent effects. For example, boroxines with a para-electron-withdrawing group have smaller equilibrium constants. This observation was interpreted as an outcome of a back-reaction (i.e., boroxine hydrolysis) that is facilitated by the increased electrophilicity of boron. Steric effects also come into play, as indicated by a smaller K-value for ortho-tolylboronic acid compared to the para-isomer. Variable temperature studies provided useful thermodynamic information, which was found consistent with a significant entropic drive for boroxine formation due to the release of three molecules of water.

(1.11)

1.2.3.2 Boronic Esters

By analogy with carboxylic acids, the replacement of the hydroxyl groups of boronic acids by alkoxy or aryloxy groups provides esters. By losing the hydrogen bond donor capability of the hydroxyl groups, boronic esters are less polar and easier to handle. They also serve as protecting groups that can mitigate the particular reactivity of boron–carbon bonds. Most boronic esters with a low molecular weight are liquid at room temperature and can be conveniently purified by distillation. Exceptionally, the trityloxymethyl esters described above are crystalline solids [15]. A selection of the most commonly encountered boronic esters is shown in Figure 1.9. Many of these esters are chiral and have also been used as inducers in stereoselective reactions discussed in Section 1.3.8.4. In addition, a number of macrocyclic oligomeric esters have been described [103].

Figure 1.9 Common types of boronic esters.

1.2.3.2.1 Stoichiometric Formation in Nonaqueous Conditions

The preparation of boronic esters from boronic acids and alcohols or diols is straightforward (Equation 1.12, Figure 1.9). The overall process is an equilibrium and the forward reaction is fast with preorganized diols, and particularly favorable when the boronate product is insoluble in the reaction solvent. The backward process (hydrolysis) can be slowed to a practical extent by using bulky diols such as pinanediol or pinacol. Otherwise, ester formation can be driven by azeotropic distillation of the water produced using a Dean–Stark apparatus or, alternatively, with the use of a dehydrating agent (e.g., MgSO4, molecular sieves, etc.). The use of mechanochemistry (i.e., solvent-less grinding) has been reported for the preparation of cyclic esters by condensation of certain diols with aliphatic and aromatic boronic acids [104]. Boronic esters can also be made by transesterification of smaller dialkyl esters like the diisopropyl boronates, with distillation of the volatile alcohol by-product driving the exchange process. In the case of cyclic esters made from the more air-sensitive alkylboronic acids, an alternate method involves treatment of a diol with lithium trialkylborohydrides [105]. Likewise, cyclic ethylboronates were prepared by reaction of polyols with triethylborane at elevated temperatures [106]. One of the first reports on the formation of boronic esters from diols and polyols, by Kuivila et al., described the preparation of several esters of phenylboronic acid by reaction of the latter, in warm water, with sugars like mannitol and sorbitol and 1,2-diols like catechol and pinacol [107]. The desired nonpolar boronic esters precipitated upon cooling the solution. Interestingly, cis-1,2-cyclohexanediol failed to provide the corresponding cyclic ester and the authors rationalized this observation on the basis of the unfavorable diol geometry of the substrate. Thus, although the two diols are not oriented in the same plane in the chair conformation (Equation 1.13, Figure 1.10), they can adopt such a favorable orientation only in the boat conformer, which is thermodynamically unfavorable [107]. Under anhydrous conditions (i.e., refluxing acetone), phenylboronic esters of cis-1,2-cyclopentanol and cis-1,2-cyclohexanol can be isolated [108]. The trans-isomers, however, failed to give a 1 : 1 adduct, and based on elemental analysis and molecular weight determinations, rather gave 1 : 2 adducts such as 45 (Equation 1.14). The existence of a seven-membered trans 1 : 2 adduct of a glucopyranoside was recently demonstrated by NMR spectroscopy [109]. This behavior can be explained in terms of the large energy required for the trans-diol to adopt a coplanar orientation, which would increase ring strain and steric interactions between axial atoms. The marked preference for the formation of boronic esters from cyclic cis-diols was exploited in the concept of dynamical combinatorial chemistry, using phenylboronic acid as a selector to amplify and accumulate one out of nine possible dibenzoate isomers of chiro-inositol that exist under equilibrating conditions through base-promoted intramolecular acyl migration (Equation 1.15) [110]. The relative thermodynamic stability of several boronic esters was examined by comparing the equilibrium composition of products in the transesterification of 2-phenyl-1,3,2-dioxaborolane with various diols by NMR spectroscopy in deuterated chloroform (Figure 1.11) [111]. Rigid, preorganized diols like pinanediol 39 provide the most robust esters and it was also found that six-membered esters are generally more stable than the corresponding five-membered boronates (i.e., 29 versus 28). Presumably, the stabilizing effect of B−O conjugation via overlap of boron with oxygen lone pairs is geometrically optimal in the larger rings. Diethanolamine boronic esters (43, Figure 1.9) represent a useful class of boronic acid derivatives [112]. Other N-substituted derivatives were characterized [113]. The presence of internal coordination between the nitrogen lone pair and boron's vacant orbital constitutes a unique structural characteristic of these tetrahedral derivatives. This coordination makes the hydrolysis reaction less favorable and even stabilizes the boron atom against atmospheric oxidation. Diethanolamine boronic esters can be conveniently formed in high yields, often without any need for dehydration techniques, as they tend to crystallize out of solution. These adducts are solids, often crystalline, with sharp melting points, and can thus be used for purifying and characterizing boronic acids, as well as in the chemical protection of the boronyl group toward various transformations (see Section 1.3.8.6). The concept of internal coordination in diethanolamine esters has been exploited in the development of the DEAM-PS resin for immobilization and derivatization of boronic acids (Section 1.4.2.1).

Figure 1.10 Specific examples of boronic ester formation with cyclic diols.

Figure 1.11 Relative thermodynamic stability in a series of boronic esters.

1.2.3.2.2 Hydrolysis and Cleavage

From a thermodynamic standpoint, the stability of B−O bonds in boronic acids and their ester derivatives is comparable (Section 1.2.2.1). Consequently, hydrolysis, in bulk water or even by simple exposure to atmospheric moisture, is a threatening process when handling boronic esters that are kinetically vulnerable to the attack of water. In fact, hydrolysis is very rapid for all acyclic boronic esters such as 27 (Figure 1.9) and for small unhindered cyclic ones such as those made from ethylene or propylene glycol (28 and 29) and tartrate derivatives 36 [114]. Catechol esters 35 are another class of popular derivatives as they are the direct products of hydroboration reactions with catecholborane (Section 1.3.4.4). Due to the opposing conjugation between the phenolic oxygens and the benzene ring, these derivatives are more Lewis acidic and are quite sensitive to hydrolysis. They are stable only in nonhydroxylic solvents and are not compatible with silica chromatography [115]. In the hydrolytic cleavage of catechol boronic esters, it is often necessary to carefully monitor the pH and buffer the acidity of the released catechol.

In contrast, hydrolysis can be slowed down considerably in the case of hindered cyclic aliphatic esters such as the C2-symmetrical derivatives 37 [116] and 38 [117], pinacol 30 [107], pinanediol 39 [118], Hoffmann's camphor-derived diols (40 and 41) [119], and the newer 42 [120] (Figure 1.9). Indeed, many of these boronic esters tend to be stable to aqueous workups and silica gel chromatography. The robustness of the esters of trans-1,4-dimethoxy-1,1,4,4-tetraphenyl-2,3-butanediol 42 was demonstrated in its applications as a protecting group for alkenylboronic acids [120]. The resulting alkenylboronic esters are tolerant of a wide variety of reaction conditions (Section 1.3.8.6). Unfortunately, the bulky boronic esters 39 –42 are very robust to hydrolysis, and their conversion back to boronic acids is notoriously difficult. The removal of the bulky pinanedioxy group in boronates 39 exemplifies the magnitude of this particular problem. It is generally not possible to cleave a pinanediol ester quantitatively in water even under extreme pH conditions. It can be released slowly (over several days) and rather ineffectively by treatment with other rigid diols in chloroform [121]. Cleavage of various pinanedioxy boronates has been achieved by transborylation with boron trichloride [22, 121–125], which destructs the pinanediol unit, or by reduction to the corresponding borane using lithium aluminum hydride (Equations 1.16 and 1.17, Figure 1.12) [126]. Both of these derivatives can be subsequently hydrolyzed to afford the desired boronic acid. More recently, mild approaches have been developed to convert the robust DICHED, pinacol, and pinanediol esters into difluoroboranes or trifluoroborate salts (Equation 1.18, Figure 1.12) [127, 128]. The latter can then be hydrolyzed to the corresponding boronic acids using various methods (Section 1.2.3.8) [128, 129]. Two-phase transesterification procedures with polystyrylboronic acid [130] or with phenylboronic acid have been described, but the latter is only applicable to small, water-soluble boronic acids [131]. Most of these procedures, such as the BCl3-promoted method, were applied to the particular case of pinanediol esters of α-acylaminoalkylboronic acids [22, 125]. Using such a substrate, 46, an oxidative method allowed the recovery of free boronic acid 47 in good yield from a periodate-promoted cleavage that destructs the pinanediol unit or by using the biphasic transesterification method in hexanes/water (pH 3) (Equations 1.19 and 1.20, Figure 1.12) [132]. The cleavage of methoxyphenyl-substituted pinacol-like boronates 31 (