Article
`
`pubs.acs.org/JACS
`
`Protodeboronation of Heteroaromatic, Vinyl, and Cyclopropyl
`Boronic Acids: pH−Rate Profiles, Autocatalysis, and
`Disproportionation
`Paul A. Cox,† Andrew G. Leach,§ Andrew D. Campbell,‡ and Guy C. Lloyd-Jones*,†
`†School of Chemistry, University of Edinburgh, Joseph Black Building, David Brewster Road, Edinburgh EH9 3FJ, United Kingdom
`‡Pharmaceutical Technology and Development, AstraZeneca, Silk Road Business Park, Macclesfield SK10 2NA, United Kingdom
`§School of Pharmacy and Biomolecular Sciences, Liverpool John Moores University, Byrom Street, Liverpool L3 3AF, United
`Kingdom
`*S Supporting Information
`
`ABSTRACT: pH−rate profiles for aqueous−organic protodeboronation of 18
`boronic acids, many widely viewed as unstable, have been studied by NMR and
`DFT. Rates were pH-dependent, and varied substantially between the boronic
`acids, with rate maxima that varied over 6 orders of magnitude. A mechanistic
`model containing five general pathways (k1−k5) has been developed, and
`together with input of [B]tot, KW, Ka, and KaH, the protodeboronation kinetics
`can be correlated as a function of pH (1−13) for all 18 species. Cyclopropyl and
`vinyl boronic acids undergo very slow protodeboronation, as do 3- and 4-pyridyl
`boronic acids (t0.5 > 1 week, pH 12, 70 °C). In contrast, 2-pyridyl and 5-thiazolyl
`boronic acids undergo rapid protodeboronation (t0.5 ≈ 25−50 s, pH 7, 70 °C),
`via fragmentation of zwitterionic intermediates. Lewis acid additives (e.g., Cu, Zn
`salts) can attenuate (2-pyridyl) or accelerate (5-thiazolyl and 5-pyrazolyl)
`fragmentation. Two additional processes compete when the boronic acid and the
`boronate are present in sufficient proportions (pH = pKa ± 1.6): (i) self-/autocatalysis and (ii) sequential disproportionations of
`boronic acid to borinic acid and borane.
`
`■ INTRODUCTION
`Boronic acids are key reagents in synthesis,1 and ubiquitous in
`classic processes such as, inter alia, Suzuki−Miyaura,2 oxidative
`Heck,3 Chan−Evans−Lam,4 and Liebeskind−Srogl couplings,5
`and addition to enones,6 carbonyls, and imines.7 Boronic acid
`decomposition, notably by in situ protodeboronation,1
`compromises
`reaction efficiency, and motifs
`such as 2-
`heteroaryl,8 vinyl,9 and cyclopropyl10 are sometimes trouble-
`some. As a consequence, a range of techniques have been
`developed to mitigate decomposition during coupling:1,11 these
`include highly tuned catalysts,12 the use of additives (e.g., Cu,
`Zn, and Ag salts) sometimes proposed to act by trans-
`metalation,13 masked reagents,14 and slow release of
`the
`boronic acid in situ11 from MIDA boronates15 and trifluor-
`oborates.16
`Given the importance of boronic acids in Suzuki−Miyaura
`coupling, a process that is frequently conducted in aqueous−
`organic solvent media,16c general mechanistic understanding of
`direct aqueous protodeboronation is
`surprisingly limited.
`Moreover, nearly all studies have focused on substituted
`phenylboronic acids.1,12,17 The most detailed investigation was
`reported by Kuivila, who measured the protodeboronation
`kinetics of a series of ArB(OH)2 species (Ar = o,m,p-X-C6H4; X
`= MeO, Me, Cl, and F) in aqueous buffers at 90 °C, with initial
`ArB(OH)2 concentrations in the range 3−5 mM.17c−f By
`
`analysis of pH−rate profiles (between pH 1.0 and 6.7), two
`pathways were identified, Scheme 1. The first was a specific
`acid-catalyzed process (k1), shown to proceed via aromatic
`electrophilic substitution of B by H. The second pathway was
`found to be base-catalyzed, and proposed to proceed via
`hydrolysis (k2) of the boronate anion ([ArB(OH)3]−). This
`latter species is generated in a pH-determined equilibrium
`involving association of water (Ka)18 or hydroxide (Ka/KW)
`with the boronic acid. A key issue is that the kinetics were
`measured by UV−vis
`spectroscopy and could not be
`determined above pH 6.7 due to the onset of UV-dominating
`boronic acid oxidation processes.17f As a consequence of the
`pH being substantially below that required to effect significant
`conversion of
`the boronic acid to the boronate,
`the rate
`constant (k2) was calculated from kobs, using an estimated value
`for Ka. In addition to the elucidation of the two major pathways
`(k1 and k2) for aqueous protodeboronation, Kuivila’s studies
`also identified that electron withdrawing groups, at para or meta
`positions on the aromatic ring, attenuate protodeboronation
`rates, via both pathways. However, while these studies were
`extensive,17c−f they were conducted long before the ascendency
`the Suzuki−Miyaura reaction.2
`of
`In other words,
`the
`
`Received: March 30, 2016
`Published:
`June 29, 2016
`
`© 2016 American Chemical Society
`
`9145
`
`DOI: 10.1021/jacs.6b03283
`J. Am. Chem. Soc. 2016, 138, 9145−9157
`
`PETITIONER NPC EX. 1023
` Page 1 of 13
`
`

`

`Journal of the American Chemical Society
`
`Scheme 1. Kuivila Mechanisms (k1, k2) for ArB(OH)2
`Protodeboronation in Aqueous Acidic (k1)17c and Basic
`(k2)17f Mediaa
`
`involving boronate [ArB(OH)3]−
`aAlso shown is a third pathway,
`deprotonation and C−B protolysis (k3), proposed by Perrin17k for
`substrates with 2,6-disubstitution (X = F, Cl, Br, CF3), and an
`uncatalyzed pathway (k4) involving direct reaction with (autoionized)
`water.
`
`importance of detailed study of the base-catalyzed process,
`across the full alkaline pH range (i.e., well above pH 6.7), was
`not yet apparent.
`Indeed,
`it was only rather recently that the kinetics of
`protodeboronation of arylboronic acids have been studied
`under basic conditions. In 2002, Frohn measured protodeboro-
`nation rates of various polyfluorophenyl boronic acids in
`aqueous pyridine, and in aqueous basic methanol, concluding
`that the mechanism involved protolysis of either the boronate
`(i.e., k2) or the conjugate base of the boronic acid.17i In 2003
`Cammidge reported in detail on the effect of various anhydrous
`and aqueous−organic media in the protodeboronation of 2,3-
`difluoro-4-heptyl-6-tolyl boronic acid mediated by CsF,17j
`concluding that aqueous protolysis of
`the corresponding
`boronate was involved. In 2010 Buchwald used calorimetry to
`measure protodeboronation kinetics of a series of substituted
`2,6-difluorophenyl boronic acids (3-F, 3-OBu, 4-F, and 4-H) in
`a biphasic basic aqueous medium (aq K3PO4/THF).12 Perrin
`extended this study,
`including other electronegative 2,6-
`disubstituents, Cl, Br, and CF3,17k which led to the proposal
`of a new, i.e., non-Kuivila type, mechanism involving specific-
`base-mediated protolysis (k3) of the boronate anion ([ArB-
`(OH)3]−). The process was reported to only occur with
`boronic acids bearing a substituent at both ortho-positions (i.e.,
`2,6-disubstitution).17k
`Despite the core role of heteroaromatic boronic acids in
`synthesis and discovery, and the propensity for many to
`undergo protodeboronation, during storage17i and in cou-
`pling,8,12 there is a near-complete absence of the kinetic data
`requisite for their behavior to be compared and contrasted.
`Thus, while it is known empirically, or anecdotally, that certain
`heteroaromatic boronic acids are much more prone to
`protodeboronation than others,1,8,11b it is not clear whether
`overall they behave similarly to substituted phenylboronic acids,
`i.e., displaying the simple acid- and base-catalyzed pH
`relationships (k1, k2) identified by Kuivila, or whether there
`are more complex pH dependencies for some classes of
`heteroaromatic boronic acids, for example involving heterocycle
`basicity, or other pathways, such as the specific-base-mediated
`protolysis (k3) identified by Perrin.17k Indeed, it is not even
`clear for an individual class of heteroaromatic boronic acid
`
`Article
`
`whether extremes of pH (low or high) are to be avoided, or are
`beneficial, in terms of stability.
`For all of the above reasons, we set out to study the intrinsic
`aqueous protodeboronation of a range of heteroaromatic (2−
`17), vinyl (18), and cyclopropyl (19) boronic acids,
`in a
`homogeneous organic−aqueous medium. Herein, we report the
`overall kinetics of
`their protodeboronation, but more
`importantly also show how the resulting pH−rate profiles can
`be simulated and analyzed using a general kinetic model. The
`model extends beyond the basic Kuivila processes (k1, k2), by
`including the Perrin mechanism (k3), plus three additional
`protolysis processes (k2cat, k4, and k5, vide infra), the requisite
`pH-dependent speciation equilibria for boronic acid association
`with water (Ka), and, if required, the protonation state of basic
`heterocycles (KaH). The model can be used in two ways: (i) as
`a general exploratory tool, with manual input of the requisite
`pH, concentrations, rates, and equilibrium constants, or (ii) as a
`means to fit experimental data, through automated numerical
`iteration of rate and equilibrium constants (including Ka and
`KaH), provided that the rate data has been acquired over a
`suitably wide pH range. To assist in application of the model, a
`preconfigured spreadsheet is provided as part of the Supporting
`Information.
`the
`The model provides the basis for quantification of
`dominant protodeboronation processes occurring for different
`boronic acid species, at different pH and substrate concen-
`trations. Thus,
`the impact of pH on the kinetics allows
`identification of new mechanistic regimes, and these mecha-
`nisms can then be explored in more detail using kinetics,
`isotopic labeling, effects of additives, and computation. Overall
`the study has facilitated the following: (i) classification of the
`reactivity imparted by 16 different heterocyclic structures (2−
`17) between pH 1 and 13; (ii) elucidation and investigation of
`new protodeboronation mechanisms and a competing dis-
`proportionation process; (iii) preliminary details on origins of
`the (de)stabilizing effect of additives such as Zn and Cu salts on
`some heterocyclic boronic acids; and (iv) identification of
`substrate-specific pH-stability zones, in which even notoriously
`unstable boronic acids, e.g., 2-pyridyl, can be stable for a few
`hours at 70 °C. This information will aid a more informed
`the preparation, storage,17i and
`choice of conditions for
`application of boronic acids in synthesis,1 as well as a means
`to induce their deliberate19 protodeboronation.
`■ RESULTS AND DISCUSSION
`1. Protodeboronation via Kuivila Mechanisms (k1 and
`k2). In preliminary studies we confirmed that the protodeboro-
`nation of a simple para-substituted phenyl boronic acid could
`be satisfactorily analyzed in situ by 1H/11B NMR under aqueous
`conditions. With p-anisyl boronic acid (1),17c,f protodeborona-
`tion kinetics were determined in water at 90 °C, without a
`malonate buffer.17c,f The aqueous association constant (pKa 1 =
`9.10; 90 °C) was determined by 11B NMR pH titration.
`Control experiments confirmed that basic solutions of 1
`became pale brown in color, with large increases observed in
`UV−vis absorption spectra, as reported by Kuivila.17f However,
`the NMR spectra of such samples were unaffected: there was
`no sign of mequinol (p-hydroxy anisole),
`the anticipated
`product of oxidation of 1, or indeed anything other than the
`time-average signal from [1/1OH] and the protodeboronation
`products, anisole and boric acid. The trace quantities of
`oxidative side product(s) are thus intensely UV-active, and
`possibly polymeric.
`
`9146
`
`DOI: 10.1021/jacs.6b03283
`J. Am. Chem. Soc. 2016, 138, 9145−9157
`
`PETITIONER NPC EX. 1023
` Page 2 of 13
`
`

`

`Journal of the American Chemical Society
`
`Using HCl and KOH to explore the acid (pH 1−3) and base
`(pH 11−13) regimes, protodeboronation kinetics were
`analyzed by nonlinear regression of the exponential decays
`observed for [1/1OH]. The second-order rate constants (k1 and
`k2) are given in Table 1, entry 1. Kuivila’s value for the limiting
`
`Table 1. Protodeboronation 1 and 2 at 90 °C
`
`k2 M−1 s−1
`k1 M−1 s−1
`a
`pKa
`ArB(OH)2
`entry
`3.9 × 10−8b
`0.68 × 10−4b
`9.10
`1
`1
`8.4 × 10−6f
`1.1 × 10−4e
`9.60d
`2c
`1
`1.4 × 10−8b
`3.3 × 10−6b
`8.91
`2
`3
`a11B NMR pH titration. b11B NMR at pH 1−3 and 11−13. cFrom refs
`17c and f. dEstimated in ref 17f. eExtrapolated from data in ref 17c.
`fFrom kobs at pH 6.7 (25 °C, uncorrected) and estimated Ka.
`
`rate constant at high pH (k2, entry 2) was obtained by pH−rate
`extrapolation and is 2 orders of magnitude too large.17k This
`arises from the conflation of an overestimated pKa for 1 (9.60),
`with an uncorrected pH (25 °C)20 for the kobs determination at
`90 °C, and reinforces the value of full pH range rate profiling.
`Moving to the protodeboronation of heterocycles, 3-
`thienylboronic acid (2) was chosen for initial studies, on the
`basis of its solubility, relative stability, and low basicity. Second-
`order rate constants (k1 and k2; Table 1, entry 3) were
`determined under the same conditions (50 mM, H2O, 90 °C)
`as for p-anisyl boronic acid 1. Within the limits of the pH range
`explored (pH 1−13), there was no detectable contribution by
`the base-catalyzed boronate mechanism (k3), or direct reaction
`of the boronic acid with H2O (k4; Scheme 1),21 although both
`mechanisms (k3 and k4) were found to be important with some
`heterocycles, vide infra.
`3-Thienyl boronic acid 2 is less susceptible to aromatic
`electrophilic substitution (k1)22 than 1, but the boronates (1OH
`and 2OH) are of similar reactivity (k2, Table 1, entries 1 and 3).
`Computational studies on this process identified rate-limiting
`C-protonolysis of the boronate (2OH, Figure 1, upper structure)
`by water. As found for all of the boronates studied, the three
`hydroxyl groups in 2OH preferentially adopt a distinctive
`triskelion conformation. Hydrogen bonding of the boronate by
`incoming water results in the C−B bond stretching such that, in
`the transition state (TS), the 3-thienyl carbanion is midpoint in
`its translation from the Lewis acid (B(OH)3) to the Brønsted
`acid (OH).
`2. Autocatalyzed Protodeboronation (k2cat). If proto-
`deboronation of 2 arises solely by the Kuivila mechanisms (k1,
`k′2), (k′2 = k2[H2O]) the empirical rate equation predicts a
`simple pH−log kobs rate profile, Figure 2. Specifically, for k′2
`(see solid line, Figure 2), the rate should rise and then reach a
`plateau as the pH extends above the pKa of the boronic acid 2
`(pKa = 8.91 at 90 °C). The data deviates from this theoretical
`curve, with the deviation being greatest at pH 8.9. This
`deviation was found for many of the boronic acids studied, vide
`infra, and indicates that there is an additional protodeborona-
`tion process that augments k′2, but with a distinctly different
`pH profile. A pronounced concentration dependence was
`
`Article
`
`Figure 1. DFT (M06L/6-311++G**)23−27 transition state structures
`for protonolysis of 3-thienylboronate 2OH by water (k2, top) and by
`boronic acid (2, k2cat, bottom) to generate 2H and [B(OH)4]−.
`Distances are shown in Å and free energies are given in kcal/mol.
`
`Figure 2. Effect of pH and concentration on the rate of
`protodeboronation of thienyl boronic acid 2. Dashed lines: kobs =
`((k′2 + k2cat[B]tot)/(1 + 10(pKa−pH))) + ((k1/10pH)/(1 + 10(pH−pKa)))
`values as Table 1, entry 3, plus k2cat = 6.2 × 10−5 M−1 s−1; for solid
`line, k2cat = 0. Inset shows rate profile in D2O (red), using an x-axis
`scale of pD + ΔpKW to account for the change in water autoionization
`(KW).30 pKa 2 (11B NMR pH titration, 90 °C): 8.91 (H2O) and 9.68
`(D2O).
`
`noted, with the deviation from k′2 becoming greater as the
`boronic acid concentration is raised (compare 0.05 to 0.40 M,
`Figure 2). As the maximum deviation occurs when pH = pKa
`(8.91), where the proportions of [RB(OH)3]− and RB(OH)2
`are identical, analogous to that of a bimolecular Job-plot,28 this
`suggests self-catalysis (k2cat). To account for the approximately
`pseudo-first-order kinetics (vide infra), the product (B(OH)3)
`needs to be a similarly effective autocatalyst,29 as the boronic
`acid is a self-catalyst.17l This was confirmed by addition of 350
`mM 10B(OH)3 to the protodeboronation of 50 mM 2. At pH
`
`9147
`
`DOI: 10.1021/jacs.6b03283
`J. Am. Chem. Soc. 2016, 138, 9145−9157
`
`PETITIONER NPC EX. 1023
` Page 3 of 13
`
`

`

`Journal of the American Chemical Society
`
`8.90, 11B NMR analysis afforded the same pseudo-first-order
`rate constant as for 400 mM 2 alone ([B]tot = 400 mM in both
`cases). Using an augmented rate equation containing k2cat
`allows satisfactory kinetic simulations (dashed lines in Figure
`2). DFT transition state energies confirmed that water can be
`replaced by boronic acid 2 (Figure 1, lower structure) or boric
`acid as proton sources, and KIEs determined from reactions in
`D2O (Figure 2, inset) confirm rate-limiting proton-transfer.
`3. B−OH-Catalyzed Disproportionation. After some
`preliminary investigation with heterocyclic boronic acids, a
`mixture of H2O/1,4-dioxane (1:1) at 70 °C was found to
`provide the best combination of substrate solubility (50 mM)
`while remaining fully homogeneous at the wide range of base,
`salt, buffer, additive, and acid concentrations required to
`usefully explore the pH range (1−13). With higher initial
`concentrations of boronic acid, some substrates displayed small
`deviations from pseudo-first-order protodeboronation kinetics
`when carefully monitored at pH = pKa ± 0.1.
`This deviation had already been observed with thienyl
`boronic acid 2 in H2O at 90 °C, vide supra, and together with
`tests for boryl exchange at the aryl ring, Scheme 2, indicated
`that processes in addition to k2cat occur when both [RB-
`(OH)3]− and RB(OH)2 are present at high concentration (pH
`= pKa).
`
`Scheme 2. 10B/11B Exchange: Reaction of 11B(OD)3 (0.2 M)
`with [10B]-2OD (0.2 M) To Generate Small Quantities of
`[11B]-2/[11B]-2OD and 10B(OD)3 (10/11B NMR, D2O, 90 °C)a
`
`aRegioisomer 3 protodeboronates without significant disproportiona-
`tion.
`
`11B NMR analyses of this process indicated other transient
`minor (≤5%) species
`to be present, and the pH-shift
`dependency of one of these species suggested it to be a
`borinate,31 [(3-thienyl)2B(OH)2]−, generated by disproportio-
`nation of 2/2OH. Regioisomeric 2-thienyl boronic acid (3) was
`found to undergo much faster protodeboronation (k2 and
`k2cat) than 3-thienyl 2, with no significant disproportionation.
`In contrast, 2-furyl boronic acid (4, 1:1 d8-dioxane/D2O, 0.5
`equiv of KOD; pD = pKa) underwent disproportionation faster
`than protodeboronation when the concentration was raised to
`0.4 M, Figure 3.
`The temporal evolution (1H NMR) indicated a two stage
`process, initially giving difurylborinic acid (R2B(OD)), which
`disproportionates further to give the trifurylborane (R3B, Figure
`3). Over a period of days, all species protodeboronate, directly
`or indirectly, to give boric acid and furan (4D).33 DFT studies
`investigated a number of mechanisms for the aryl transfer
`between the boron centers. These included an analogue of the
`autocatalysis mechanism, Figure 1, TS 2OH (k2cat), in which the
`furyl anion transfers to the Lewis acid B, rather than Brønsted
`
`Article
`
`Figure 3. Upper: temporal concentration data for disproportionation
`of 2-furyl boronic acid 4 (1:1 d8-dioxane/D2O, 0.5 equiv of KOD, 70
`°C). Circles: data. Solid lines: simulation (see SI for details). 2H1-furan
`(RD = 4D) is volatile under the reaction conditions and was not
`monitored. Lower: transition states (DFT) for disproportionation to
`difurylborinic acid and trifurylborane.32
`
`acid H, in an H-bonded [4 + 4OH] intermediate. However, the
`computed barrier for transfer to B (22.5 kcal/mol) is greater
`than that for H (20.3 kcal/mol), and protodeboronation would
`dominate if solely these isomeric transition states are operative.
`An alternative process was therefore considered in which a
`cyclic boroxine-ate complex (Figure 3, X = furyl or OH, upper
`structure) facilitates aryl-migration across the ring. The rate-
`limiting barriers for aryl migration are computed to be 19.1−
`19.4 kcal/mol for 2-furyl 4OH and 19.5−19.9 kcal/mol for 3-
`thienyl 2OH.
`Accurately computing free energies in solution for very
`different processes is challenging; nonetheless, the calculations
`suggest that migration dominates over protodeboronation for
`2-furyl 4, while the opposite is the case for 3-thienyl 2. The
`transition state structures suggest an electronic and steric
`component to this preference. Mulliken charges indicate the
`nonmigrating 2-furyl bears a larger negative charge than the
`equivalent group for 3-thienyl. Further,
`the preferred
`conformation of
`the transition state places the migrating
`group in close proximity to the other aromatic ring; larger rings
`or those with C−H groups adjacent to the boronic acid group
`will inevitably involve steric clashes, even more so when X =
`aryl. The subsequent step leading to triarylborane requires
`reaction of a borinic with a boronic acid and thus cannot
`involve a cyclic boroxine. Instead it is computed to proceed via
`a transition state involving a dimer (Figure 3, lower structure).
`The free energy barrier (+18.9 kcal/mol) is comparable to that
`for the first step, and thus consistent with the observed
`
`9148
`
`DOI: 10.1021/jacs.6b03283
`J. Am. Chem. Soc. 2016, 138, 9145−9157
`
`PETITIONER NPC EX. 1023
` Page 4 of 13
`
`

`

`Journal of the American Chemical Society
`
`Article
`
`Figure 4. pH−rate profiles (I−IV) for pseudo-first-order protodeboronation (kobs/s−1) of boronic acids 2−19 in 1:1 H2O/1,4-dioxane at 70 °C.
`Reactions analyzed in situ by 11B NMR were conducted in quartz NMR tubes to avoid [B(OH)4]− release from borosilicate NMR tubes. Circles:
`experimental data. Solid lines: simulation using model (Figure 5) with data from Table 2. Rates below log kobs = −7 are not modeled.
`
`Figure 5. General model for heterocycle protodeboronation (Z = basic nitrogen), individual pH−log kobs profiles, overall rate equation based on the
`three-state speciation [neutral (X), N-protonated (XH+), and boronate (XOH)], and example combined pH−log kobs profiles (i, ii, iii).
`
`transient accumulation of diarylborinic acid. Using these
`models for
`reversible furyl
`transfer, with an irreversible
`
`9149
`
`DOI: 10.1021/jacs.6b03283
`J. Am. Chem. Soc. 2016, 138, 9145−9157
`
`PETITIONER NPC EX. 1023
` Page 5 of 13
`
`

`

`Journal of the American Chemical Society
`
`Article
`
`Table 2. Equilibrium and Rate Constants Employed in the pH Simulation (See Overall Rate Equation in Figure 5) for
`Protodeboronation of Heteroaromatic, Vinyl, and Cyclopropyl Boronic Acids 2−19 in 1:1 H2O/1,4-Dioxane at 70 °C
`
`a
`pKaH
`
`b
`
`e
`
`log k′4
`
`log k′5
`
`log k′3
`log k′2
`c
`log k2cat
`log k1
`pKa
`RB(OH)2
`entry
`≤−5.32d
`−4.21
`−6.01
`−5.95
`11.04
`2
`1
`≤−3.07d
`≤−3.54d
`−3.92
`−4.44
`10.38
`3
`2
`≤−4.12d
`−3.32
`−4.98
`−3.71
`10.29
`4
`3
`≤−1.93d
`≤−2.07d
`−2.34
`−3.41
`9.64c
`5
`4
`−3.72
`−3.27
`≤−6.02d
`−3.16
`11.61
`6
`5
`−2.62
`−2.00
`−4.01
`−5.51
`10.45
`7
`6
`≤−4.75d
`−3.44
`−5.04
`≤−4.82d
`8.18
`8
`7
`≤−5.59d
`−3.60
`−6.05
`≤−3.10d
`9.76
`9
`8
`≤−5.78d
`−3.78
`−5.93
`≤−3.76d
`8.94
`10
`9
`≤−6.60d
`−3.95
`−6.24
`≤−5.64d
`9.90
`11
`10
`≤−2.21d
`≤−1.58d
`−2.02
`≤−2.60d
`8.41c
`12
`11
`≤−3.94d
`−3.49
`−3.87
`≤−4.22d
`8.82
`13
`12
`≤−3.93d
`−3.19
`−4.35
`≤−4.58d
`9.11
`14
`13
`≤−3.74d
`≤−0.67d
`≤−4.26d
`≤−0.74d
`10.76c
`15
`14
`≤−6.04d
`≤−2.33d
`≤−6.72d
`≤−1.36
`9.54c
`16
`15
`≤−4.82d
`≤−1.84d
`≤−6.32d
`≤−2.06d
`8.49c
`17
`16
`≤−5.73d
`−3.90
`−6.02
`−5.57
`11.21
`18
`17
`≤−5.91d
`−5.01
`−6.01
`−5.95
`12.72
`19
`18
`apKaH in parentheses determined by 1H NMR pH titration at 25 °C; pKaH at 70 °C from iterative fitting of rate data (unless stated). bpKa by 11B
`NMR pH titration at 70 °C (unless stated). cpKa from iterative fitting of rate data. dThis value is not required for satisfactory simulation; greater
`values induce ≥5% change in the sum of square errors between simulation and data across the overall pH profile. epKaH fixed at empirical offset
`(−0.46 units) based on entries 5, 8, 11−16. fDue to the pKaH of 17 being <−1.06, this value is a substantial overestimate.
`
`1.26 (1.70)
`0.04e (<0.50)
`0.14e (<0.60)
`3.60 (4.22)
`3.36e (3.82)
`2.95e (3.41)
`1.62 (1.85)
`1.08 (1.62)
`1.36 (2.00)
`3.52 (3.86)
`3.16 (3.60)
`−1.06e (<−0.6)
`
`−6.48
`−6.60
`−6.38
`−5.44
`−5.71
`−5.71
`−6.89
`−1.83
`−3.78
`−4.28
`−1.60
`−2.41
`−2.75
`
`≤−5.68d
`≤−4.61d
`≤−4.83d
`≤−6.70d
`−5.77
`−7.07
`≤−4.22d
`≤−5.31d
`≤−5.84d
`≤−4.24d
`≤−4.52d
`≤−0.06d
`
`protodeboronation, predominantly via the boronic acid, a good
`fit to the experimental data was achieved by kinetic simulation
`(solid lines, Figure 3).
`4. General Model for Protodeboronation Kinetics.
`With suitable conditions established for analysis (50 mM
`substrate, 1:1 H2O/1,4-dioxane, 70 °C) the pseudo-first-order
`reactions of boronic acids 2−19 were
`kinetics (kobs) of
`measured as a function of pH (1−13), Figure 4, profiles I−
`IV, see SI for full details. Rates were pH-dependent, and varied
`substantially between the boronic acids, with rate maxima that
`varied over 6 orders of magnitude (half-lives ranging from
`seconds to weeks). For some substrates, specific pH ranges
`reduced the protodeboronation rates to such an extent that, for
`reasons of accuracy, they were omitted from the log(kobs)−pH
`analyses. An arbitrary threshold of log kobs ≥ 7.0 (half-life ≤ 80
`days) was set for data inclusion (see dashed lines in profiles I−
`IV, Figure 4).
`To analyze the pH-dependency of the protodeboronation
`reactions of heterocyclic boronic acids, we developed a general
`model, Figure 5. Aiming to keep as minimal a model as
`required, we began with the Kuivila processes, and added
`further processes, when necessary, as
`the analysis of
`heterocyclic boronic acids 2−17 evolved. To simplify the
`discussion, the overarching model is presented in advance of
`the analysis of the protodeboronation characteristics of 2−17.
`As many of the heterocyclic systems studied are basic, the
`model
`includes, when appropriate,
`the pKaH of
`the N-
`
`in addition to the pKa
`
`protonated form of the heterocycle,
`for aqueous association at boron.
`The kinetics are determined by a 3-fold speciation of the
`boronic acid (X): as an N-protonated form (XH+; only for 6−
`17), a neutral form (X), and a boronate form (XOH). Specific
`protodeboronation processes occur from the three speciation
`states. For the neutral form (X), there is the acid-catalyzed
`Kuivila process (k1), and a pH-independent direct reaction with
`water (k4). For the latter, although this is not kinetically
`differentiated in the model, basic heterocyclic boronic acids can
`engage in a pre-equilbrium (K4) with autoionized water to
`generate a zwitterionic adduct (XZW). For the boronate form
`(XOH),
`there are the base-catalyzed Kuivila process (k2),
`concentration-dependent autocatalysis (k2cat) occurring with
`rate maximum when pH = pKa, and the Perrin mechanism
`involving base-catalyzed protonolyis (k3). For
`the latter,
`although this is not kinetically differentiated in the model, the
`boronate can engage with hydroxide in a pre-equilbrium (K3)
`to generate the dianion (XO‑). Finally, for the N-protonated
`form (XH+), this being distinct from the zwitterion (XZW) due
`to the presence of a boronic acid, not a boronate, there is a
`direct protodeboronation by water (k5).
`As
`indicated graphically beside the protodeboronation
`mechanisms in the model (Figure 5), each of the six processes
`(k1, k2, k2cat, k3, k4, k5) has a distinct pH−log kobs profile. It can
`be instructive to consider hypothetical combinations of selected
`pH−log kobs
`relationships; examples (i,
`ii, and iii) are
`
`9150
`
`DOI: 10.1021/jacs.6b03283
`J. Am. Chem. Soc. 2016, 138, 9145−9157
`
`PETITIONER NPC EX. 1023
` Page 6 of 13
`
`

`

`Journal of the American Chemical Society
`
`preconfigured in the spreadsheet provided in the SI. By
`mathematical combination of all six steps (see SI for full
`derivation) and calculation of the three-state speciation, an
`“overall rate equation” (Figure 5, center) can be generated. The
`equation allows analysis of
`the empirical rate (kobs) as a
`function of pH and boronic acid concentration, using up to
`nine constants: pKa, pKaH, k1, k′2, k′2cat, k′3, k′4, k′5, pKw, in
`which processes indicated by the term k′ contain amalgamated
`constants (e.g., K, k, and [H2O]). For nonbasic boronic acids,
`pKaH is nominally set to −5 to preclude XH+ speciation.
`Values for pKa and pKaH can be determined independently,
`or via the pH−log kobs simulation. Initial pKaH values for 6−17
`were determined by 1H NMR pH titration (1:1 H2O/1,4-
`dioxane, 25 °C), Table 2. An empirical correction for
`temperature (ΔpKaH 25−70 °C = −0.46) was determined
`from the pH−rate profile simulation. Generally, the B(OH)2
`unit slightly decreases the basicity relative to the parent
`heterocycle.34 For example, 2-, 3-, and 4- pyridyl boronic acids
`(15, 9, 10) have pKaH values (25 °C) of 3.86, 4.22, and 3.82,
`compared to that of 4.38 for pyridine, see SI. The pKa values for
`most of the boronic acids were measured by 11B NMR pH
`titration (1:1 H2O/1,4-dioxane, 70 °C); the pKa of the more
`reactive species (5, 12, 15−17) were estimated via the pH−log
`kobs simulation.
`5. pH−Rate Profiles Analysis of Boronic Acids 2−19.
`Using the rate equation in Figure 5, the pH−rate profiles for
`2−19 were simulated by automated iteration of rate and
`equilibrium constants, minimizing the sum square error (SSE)
`between predicted and observed data across the full profile (pH
`1−13). The SSE-minimized fitting constants are provided in
`Table 2. In all cases (2−19), only a subset of the six pathways
`(k1, k′2, k2cat, k′3, k′4, k′5) were required for satisfactory
`simulation (solid lines through data, Figure 4). The constants
`that are not required for simulation, but are in principle feasible,
`are reported as threshold values (≤) that induce a ≤5% change
`in the SSE.35
`Nonbasic Heterocycles (2−5). The thienyl (2, 3) and furyl
`(4) boronic acids (Figure 4, profile I) required only the Kuivila
`processes (k1, k′2)17c,k and autocatalysis (k2cat) for simulation.
`The 2-pyrrole boronic acid 5 required an additional pH-
`independent process (k′4)21 to fit the data between pH 3 and 6.
`This process is slow enough (k′4, half-life > 24 days) to be
`consistent with a water autoionization mechanism.36 Higher
`reactivity of 2- versus 3-thienyl and furyl boronic acids has been
`noted before, but only for acid-catalysis (k1).22 In all cases (2−
`5), the base-catalyzed process (k′2) is more efficient than the
`acid (k1), and the rates rise substantially through the series;
`above pH 11, 2-pyrrolyl 5 has a half-life of less than 3 min.
`Basic Heterocycles (6−17). The rate data obtained for 2-
`pyridyl boronic acids 15−17 (profile II) show a near inverse
`pH−rate profile compared to that of the nonbasic heterocycles
`2−5 (profile I) and required just a single term (k′4) for
`simulation. Maximum rates are attained when speciation
`disfavors the boronate (15−17OH; high pH) and pyridinium
`(15−17H+; low pH) forms. Thus, for the least basic 17 (pKaH
`<−0.6), the protodeboronation is not detectably attenuated by
`acid, even at pH 1. The much less reactive 5-pyrimidyl (8) and
`3-pyridyl (9) systems required base-catalysis (k′2 + k2cat) in
`addition to the neutral mechanism (k′4) for simulation. The 5-
`pyrazolyl and 5-thiazolyl boronic acids (13−15) were highly
`reactive, requiring both k′4 (neutral) and k′2 (basic) pathways
`for satisfactory simulation, with rates attenuated at pH below
`their pKaH (profile III). The 4-pyridyls (10, 11) are of similar
`
`Article
`
`they still
`in that
`reactivity to 3-pyridyl (9), but differ
`protodeboronate effectively at a pH substantially below their
`pKaH (profile III). This effect was satisfactorily simulated by
`including a direct (H2O-mediated) protodeboronation of the
`conjugate acid (k′5); the 4-pyridyl systems (10, 11) were the
`only species requiring this. The two remaining basic hetero-
`cycles, 4-pyrazolyl (6) and 4-isoxazolyl (7), gave the most
`complex profiles, requiring acid (k1), base (k′2/k2cat), base-
`catalyzed boronate (k′3), and neutral (k′4) pathways (profile
`IV; Figure 4). These were the only examples requiring the
`Perrin mechanism (k′3), a process that in Ar−B(OH)2 systems
`is proposed to require 2,6-disubstitution.17k
`Vinyl and Cyclopropyl Boronic Acids (18−19). Strongly
`acidic or basic solutions were required to effect any significant
`protodeboronation of the vinyl (18) and cyclopropyl (19)
`boronic acids (profile IV). Even at pH ≥ 11, the half-lives are
`weeks. A small selection of simple alkyl boronic acids (Me, c-
`Bu, and c-Hex) were also investigated. The extremely slow
`reactions (half-lives of months, see SI) made it difficult to
`clarify, by 1H/13C/11B NMR analysis, if protodeboronation was
`the major pathway of decomposition, rather than, for example,
`oxidation.
`6. Protodeboronation Mechanisms for Basic Hetero-
`cycles. To aid rationalization of the diverse range of pH
`for basic heterocycles 6−17, protodeboronation
`profiles
`mechanisms were explored by 11B NMR, DFT calculations
`(M06L/6-311++G**;23 see SI for details), and by testing the
`effect of additives.
`Zwitterionic Water Adducts (XZW). 2-Pyridyl boronic acids
`are notorious
`for
`their
`susceptibility to protodeborona-
`tion,1,8,11−13 as is found for 15−17. However, only a single
`neutral process (k′4) is required for simulation of the full pH−
`rate profile, indicating that the classic Kuivila-type acid- and
`base-catalyzed mechanisms (k1 and k′2) are negligible
`processes. Indeed H+/OH− act as powerful protodeboronati