j1
`
`1 S
`
`tructure, 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 CB 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 carbox-
`ylic 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).
`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
`
`Boronic Acids: Preparation and Applications in Organic Synthesis, Medicine and Materials, Second Edition.
`Edited by Dennis G. Hall.
`Ó 2011 Wiley-VCH Verlag GmbH & Co. KGaA. Published 2011 by Wiley-VCH Verlag GmbH & Co. KGaA.
`
`

`

`2j 1 Structure, Properties, and Preparation of Boronic Acid Derivatives
`
`R B
`
`R''
`
`R'
`
`R B
`
`OH
`
`R'
`
`R B
`
`OH
`
`OH
`
`HO B
`
`OH
`
`OH
`
`borane
`
`borinic acid
`
`boronic acid
`
`boric acid
`
`R
`
`BO
`
`R
`B
`
`O
`boroxine
`(cyclic boronic anhydride)
`
`OR'
`
`R B
`
`OR'
`boronic ester
`(R' = alkyl or aryl)
`
`O
`B
`
`R
`
`Figure 1.1 Oxygen-containing organoboron compounds.
`
`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 CC 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 VelcadeÒ, 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
`
`

`

`1.2 Structure and Properties of Boronic Acid Derivativesj3
`
`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 diazabor-
`inanes, 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
`OHO hydrogen bonds (Figure 1.3a and b). The CBO2 plane is quite coplanar with
`the benzene ring, with a respective twist around the CB 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
`
`NO2
`
`B
`
`OH
`
`OH
`
`3
`
`B
`
`OH
`
`OH
`
`X
`
`N
`4 X = Br
`5 X = Cl
`
`B
`
`OH
`
`OH
`
`HO2C
`
`X
`
`C
`
`1 X = H
`2 X = OMe
`
`B
`
`OH
`
`OH
`
`4
`
`6
`
`OO
`
`B
`
`H
`
`8
`
`MeO
`
`B
`
`OH
`OH Na
`OH
`
`11
`
`Ph3COCH2B
`
`O
`
`O
`
`7
`
`O
`B
`
`O
`
`HN
`
`O
`
`10
`
`Me
`
`CH3S
`
`O O
`
`B NH
`
`9
`
`Figure 1.2 Boronic acid derivatives analyzed by X-ray crystallography.
`
`

`

`4j 1 Structure, Properties, and Preparation of Boronic Acid Derivatives
`
`H
`
`O
`
`H
`O
`
`BA
`
`r
`
`B A
`
`r
`
`O
`
`H
`
`O
`
`H
`H
`
`B
`
`H
`
`OO
`
`H H
`
`(b)
`
`O O
`
`H
`
`B
`
`O
`
`H
`O
`
`BA
`
`r
`
`B A
`
`r
`
`O
`
`H
`
`O
`
`H
`
`H
`
`O
`
`H
`O
`
`r
`
`BA
`
`r
`
`B A
`
`O
`
`H
`
`O
`
`H
`
`(a)
`
`(c)
`
`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.
`
`(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
`, the CB bond of
`
`(Figure 1.2) [13]. With a range of approximately 1.55–1.59 A
`boronic acids and esters is slightly longer than typical CC single bonds (Table 1.1).
`The average CB bond energy is also slightly smaller than that of CC bonds (323
`versus 358 kJ/mol) [14]. Consistent with strong BO bonds, the BO distances of
`tricoordinate boronic acids such as phenylboronic acid are fairly short, and lie in the
`
`range of 1.35–1.38 A
`(Table 1.1). These values are slightly larger than those observed in
`boronic esters. For example, the BO bond distances observed in the X-ray crystal-
`
`

`

`1.2 Structure and Properties of Boronic Acid Derivativesj5
`
`Table 1.1 Bond distances from X-ray crystallographic data for selected boronic acid derivatives
`(Figure 1.2).
`
`Compound
`
`BC (A
`)
`
`
`
`BO1 (A)
`
`
`
`
`
`BO2 (A)
`
`
`
`
`
`BX (A)
`
`
`
`
`
`1
`2
`3
`4
`5
`7
`8
`9
`10
`11
`
`1.568
`1.556
`1.588
`1.573
`1.573
`1.560
`1.494
`1.613
`1.613
`1.631
`
`1.378
`
`1.365
`1.363
`1.362
`1.316
`1.408
`1.474
`1.438
`1.492
`
`1.362
`
`1.346
`1.357
`1.352
`1.314
`1.372
`1.460
`1.431
`1.487
`
`1.666
`1.641
`1.471
`
`Reference
`
`[9]
`[10]
`[11]
`[12]
`[12]
`[15]
`[16]
`[18]
`[22]
`[23]
`
`lographic structures of the trityloxymethyl pinacolate boronic esters (e.g., 7 in
`
`Figure 1.2) are in the range of 1.31–1.35 A
`(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 benzox-
`aborole 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 BO bond in 8 is slightly
`longer than the BOH 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].
`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 BN bridge long suspected from
`other spectroscopic evidence such as NMR [19, 20]. This dative BN bond has a
`(Table 1.1), and it induces a strong Nd þBd
`
`dipole that points away
`length of 1.67 A
`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 character-
`istics 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
`
`

`

`6j 1 Structure, Properties, and Preparation of Boronic Acid Derivatives
`
`O
`
`N
`
`O
`
`Ph
`
`B
`O
`
`O O
`B
`
`MeO
`
`OMe
`
`Me2N
`
`O O
`B
`
`NMe2
`
`12
`
`13
`
`14
`
`Figure 1.4 B-Chiral tetrahedral boronate 12 and model compounds for boron hypercoordination.
`
`
`
`Section 1.2.3.6. When tetracoordinated such as in structures 9–11 [23] (Figure 1.2),
`the BO bond length of boronic acids and esters increases to about 1.43–1.48 A
`
`,
`
`longer than the corresponding tricoordinate analogues
`which is as much as 0.10 A
`(Table 1.1). These markedly longer BO bonds are comparable to normal CO
`ether bonds (1.43 A
`). These comparisons further emphasize the considerable
`strength of BO bonds in trigonal boronic acid derivatives. Not surprisingly, trigonal
`BO bonds are much stronger than the average CO 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 BO 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 BO bond energy compared to the tricoordinate boronate [25].
`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 DFTcalculations, 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 (B¼ 2.05, C¼ 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 13C 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 BC 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 13C NMR shift effects,
`
`

`

`1.2 Structure and Properties of Boronic Acid Derivativesj7
`
`R
`
`β
`
`α
`
`A
`
`B(OR')2
`
`R
`
`B(OR')2
`
`B
`
`RO
`
`B(OH)2
`
`RO
`
`B(OH)2
`
`E
`D
`Figure 1.5 Limit mesomeric forms involving BC pi-overlap.
`
`δ –
`B(OR')2
`
`R
`
`δ+
`
`δ+
`RO
`
`C
`
`F
`
`δ –
`B(OH)2
`
`
`
`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 BC pi-bonded form B to give resonance hybrid C (Figure 1.5). X-ray crystal-
`lography may also provide insights into the extent of BC pi-bonding. The difference
`in BC bond distances for arylboronic acids (Table 1.1) is significant enough to
`suggest a small degree of BC pi-bonding. The BC bond distance (1.588 A
`) 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
`). Interestingly, the BC bond of 2
`
`longer than that of phenylboronic acid (1.568 A
`
`, suggesting only a minimal contribution from the mesomeric form
`stands at 1.556 A
`E (Figure 1.5). On the other hand, the BC bond distance of 1.613 A
`
`in the
`diethanolamine adduct 9 (Table 1.1), where the boron vacant orbital is also incapac-
`itated from BC pi-bonding, is 0.045 A
`
`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 BC 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 11B chemical shifts correlations [34]. Likewise, BC 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 organ-
`ometallic 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 ethyldiazoace-
`tate [36] and in Diels–Alder reactions where it provides cycloadducts with dienes like
`cyclopentadiene [37] and cyclohexadiene, albeit only at elevated temperatures (about
`
`C, respectively) [38, 39]. The higher reactivity of ethylene boronates as
`130 and 200
`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
`
`

`

`8j 1 Structure, Properties, and Preparation of Boronic Acid Derivatives
`
`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).
`
`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 Structure and Properties of Boronic Acid Derivativesj9
`
`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 Brønsted
`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.
`
`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 Brønsted 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
`
`(1.1)
`
`(1.2)
`
`+ H3O+
`
`+ H3O+
`
`O O
`
`H
`
`R B
`
`R B
`
`OH
`OH
`OH
`
`R B
`
`R B
`
`OH
`
`OH
`
`OH
`
`OH
`
`+ H2O
`
`+ 2 H2O
`
`Figure 1.6 Ionization equilibrium of boronic acids in water.
`
`

`

`10j 1 Structure, Properties, and Preparation of Boronic Acid Derivatives
`
`Table 1.2 Ionization constant (pKa) for selected boronic acids.
`
`Boronic acid, RB(OH)2
`
`Boric acid, B(OH)3
`Methyl
`Phenyl
`3,5-Dichlorophenyl
`3,5-Bis(trifluoromethyl)phenyl
`2-Methoxyphenyl
`3-Methoxyphenyl
`4-Methoxyphenyl
`4-Carboxyphenyl
`2-Nitrophenyl
`4-Nitrophenyl
`4-Bromophenyl
`4-Fluorophenyl
`2-Methylphenyl
`3-Methylphenyl
`4-Methylphenyl
`3,5-Dimethylphenyl
`3-Methoxycarbonyl-5-nitrophenyl
`2-Fluoro-5-nitrophenyl
`3-Pyridyl (15)
`3-Benzyl-3-pyridylium
`8-Quinolinyl
`2-(R1R2NCH2)phenyl (e.g., 16)
`
`pKa
`
`9.0
`10.4
`8.9
`7.4
`7.2
`9.0
`8.7
`9.3
`8.4
`9.2
`7.1
`8.6
`9.1
`9.7
`9.0
`9.3
`9.1
`6.9
`6.0
`4.0, 8.2
`4.2
`4.0, 10
`5.2–5.8
`
`Reference
`
`[58]
`[58]
`[59]
`[59]
`[59]
`[57]
`[59]
`[60]
`[56]
`[61]
`[60]
`[59]
`[59]
`[62]
`[62]
`[62]
`[59]
`[63]
`[57]
`[64]
`[57]
`[65]
`[66]
`
`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 11B NMR spectroscopy. Phenyl-
`boronic 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 1.3, 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
`
`

`

`1.2 Structure and Properties of Boronic Acid Derivativesj11
`
`CH3
`
`B(OH)2
`
`2 H2O
`
`CH3
`
`B
`
`OH
`OH
`OH
`
`H2O
`
`HN
`
`N
`
`B
`
`OH
`
`OH
`
`15
`
`+ H3O
`
`B
`
`OH
`OH
`OH
`
`HO
`
`NHR1R2
`OH
`
`B
`
`OH
`
`pH < 5
`
`pKa 1
`
`NHR1R2
`OH
`
`B
`
`OH
`
`H2O
`
`16
`
`R1R2
`N
`H
`O H
`B OH
`OH
`
`17
`
`pH >12
`
`pKa 2
`
`NR1R2
`OH
`OH
`OH
`
`B
`
`OH
`
`HO
`
`B
`
`– H2O
`
`OH
`
`BO
`
`OH
`
`BO
`
`18
`
`H2O
`
`OH
`
`O
`
`BO
`
`19A
`
`OH
`OH
`
`BO
`
`19B
`
`Figure 1.7 Ionization equilibrium of special boronic acids.
`
`(1.3)
`
`(1.4)
`
`(1.5)
`
`(1.6)
`
`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 1.4,
`Figure 1.7) [64]. Similarly, arylboronic acids of type 16 (Equation 1.5), 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 BN 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 mecha-
`nism [69]. The application of boronic acids of type 16 in the aqueous recognition
`
`

`

`12j 1 Structure, Properties, and Preparation of Boronic Acid Derivatives
`
`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].
`Boronic acids display Brønsted 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 11B NMR and UV
`spectroscopic evidence, it was suggested that 18 acts as a Brønsted 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 Brønsted 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 BF bond strength is a key factor in these complexes as
`other halide salts do not form similar adducts. As suggested by 1H 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 BN
`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 1H NMR spec-
`troscopy 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).
`
`1.2.2.5 Chemical Stability
`
`1.2.2.5.1 Ligand Exchange and Disproportionation Several favorable factors contrib-
`ute 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 BC/
`BO bond metatheses to give the corresponding disproportionation products
`
`

`

`1.2 Structure and Properties of Boronic Acid Derivativesj13
`
`(1.7)
`
`(1.8)
`
`(1.9)
`
`OP
`
`h
`
`Ph
`
`B B
`
`O O
`
`21
`
`Ph B
`
`RNH2
`
`BPh(OH)
`
`BPh(OH)
`
`O O
`
`20
`
`Ph B
`
`RNH2
`
`3 PhB(OH)2 + RNH2
`or
`pyridine
`
`Ph
`
`Bn
`
`Ph
`
`O
`
`NH
`
`H
`
`O
`B
`
`B
`O
`
`H
`
`HN
`
`O
`
`Ph
`
`O
`
`Ph
`
`Bn
`
`BnNH2
`
`Ph
`
`O
`B
`
`B
`O
`
`H
`
`HN
`
`O
`
`BnNH2
`
`Bn
`
`Ph
`
`Ph
`
`23
`
`Ph
`
`24
`
`Ph
`
`B
`
`O
`
`O
`
`312 nm
`
`>434 nm
`
`Ph
`
`Ph
`
`Ph
`
`O
`
`O
`
`O
`
`B B
`
`O
`
`22
`
`Ph
`
`Ph
`
`B
`
`O
`
`O
`
`Ph
`
`S
`25 (open, less acidic)
`
`S
`
`Ph
`
`S
`
`Ph
`Ph
`26 (closed, more acidic)
`
`S
`
`Figure 1.8 Bimolecular Lewis acid–base complexes with boronic esters.
`
`(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. Substi-
`tution 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
`BX 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 BX bond stronge