Solubility of Phenylboronic Acid and its Cyclic Esters in Organic Solvents

Article abstract of DOI:10.1007/s10953-020-00988-y

The solubilities of phenylboronic acid, its pinacol ester and azaester in organic solvents (chloroform, 3-pentanone, acetone, dipropyl ether and methylcyclohexane) have been determined experimentally by a dynamic method, in which the disappearance of turbidity was determined by measuring of light intensity using a luminance probe. Phenylboronic acid has high solubility in ether and ketones, moderate in chloroform and very low in hydrocarbon. Pinacol ester and azaester show better solubility than the parent acid in all tested solvents. For pinacol ester differences between particular solvents are small, while for azaester the differences are significant. For both esters the highest solubility is observed in chloroform and the lowest in the hydrocarbon. The results have been correlated by the Wilson, NRTL and Redlich–Kister equations. For the phenylboronic acid better correlation of the data is obtained by polynomials in comparison with the above equations. It is connected with additional acid-anhydride equilibrium in the system. The influence of polarity of the solvents on the solubility is discussed.

Full text of DOI:10.1007/s10953-020-00988-y

Journal of Solution Chemistry
https://doi.org/10.1007/s10953-020-00988-y
Solubility of Phenylboronic Acid and its Cyclic Esters
in Organic Solvents
Paweł Leszczyński1 · Tadeusz Hofman1 · Andrzej Sporzyński1
Received: 20 December 2019 / Accepted: 3 March 2020
© The Author(s) 2020
Abstract
The solubilities of phenylboronic acid, its pinacol ester and azaester in organic solvents
(chloroform, 3-pentanone, acetone, dipropyl ether and methylcyclohexane) have been
determined experimentally by a dynamic method, in which the disappearance of turbid-
ity was determined by measuring of light intensity using a luminance probe. Phenylbo-
ronic acid has high solubility in ether and ketones, moderate in chloroform and very low
in hydrocarbon. Pinacol ester and azaester show better solubility than the parent acid in all
tested solvents. For pinacol ester diferences between particular solvents are small, while
for azaester the diferences are signifcant. For both esters the highest solubility is observed
in chloroform and the lowest in the hydrocarbon. The results have been correlated by the
Wilson, NRTL and Redlich–Kister equations. For the phenylboronic acid better correlation
of the data is obtained by polynomials in comparison with the above equations. It is con-
nected with additional acid-anhydride equilibrium in the system. The infuence of polarity
of the solvents on the solubility is discussed.
Keywords Boronic acids · Boronic esters · Solubility · Redlich–Kister equation · Wilson
equation · NRTL equation
1 Introduction
Arylboronic acids and their derivatives are an important group of compounds due to their
broad applications in organic synthesis, catalysis, supramolecular chemistry, and materi-
als engineering [1–4]. The compounds have been known for about 150 years, but increas-
ing interest is observed after the discovery of new areas of their use. The most important
felds are the synthesis of biaryl compounds in the Suzuki–Miyaura reaction (Nobel 2010)
[5], the molecular receptors of sugars [6], covalent organic frameworks [7] and their use
Electronic supplementary material The online version of this article (doi:https://doi.org/10.1007/s1095
3-020-00988-y) contains supplementary material, which is available to authorized users.
*
Andrzej Sporzyński
spor@ch.pw.edu.pl
1
Faculty of Chemistry, Warsaw University of Technology, Noakowskiego 3, Warsaw 00-664,
Poland
1 3
Vol.:(0123456789)

Journal of Solution Chemistry
as biologically active compounds [8]. Such wide applications require recognition of the
physicochemical properties of these compounds. An important issue is the description of
phase equilibria, including solubility in water and organic solvents. Knowledge of these
data allows correct selection of the solvent for a particular reaction or for purifcation of the
products by crystallization. They are also important for biological tests and formulation of
biologically active compounds.
There is a need to fnd boronic acids that have good water solubility. This feature is
important for the construction of molecular sugar receptors and the use of boronic acids
as biologically active agents. Although it is known that substituents in phenylboronic acids
signifcantly afect the solubility of these compounds in water, there is a surprisingly small
number of solubility data in literature. Even for unsubstituted phenylboronic acid, only
two values of its solubility in water are reported: “The solubility (g/100 g. of water) is
approximately 1.1 at 0° and 2.5 at 25°; the solubility-temperature relationship is linear to at
least 45°” [9]. One can also fnd some generalized and misleading statements in the litera-
ture such as “boronic acids are water soluble” [10] or “since the reactants are insoluble in
water, this reaction is an >on water<reaction” [11]. Only recently, solubilities of several
simple arylboronic acid and benzoxaboroles in water were published [12]. It was found,
that solubility of this class of compounds was low, with the value of 1.9 g/100 g H2O for
phenylboronic acid at 20 ºC and even lower for some substituted compounds. Solubility
of 4-methoxyphenylboronic acid in water as a function of pH was also reported [13]. In
addition, there are some examples of particular compounds well soluble in water. They are
large molecules in which good solubility in water results from the presence of substituents
in the fragment connected to the phenylboronic acid ring [14, 15].
The solubility of boron compounds in organic solvents is also important because of their
very wide application in organic synthesis, e.g. in the Suzuki–Miyaura reaction. The choice
of solvent signifcantly afects the efciency and course of this reaction [16]. In many reac-
tions boronic acid derivatives, e.g. esters or azaesters, can be used instead of the parent
acids. These derivatives are easier to obtain in a pure state and have a stable composi-
tion in contrast to boronic acids, which may contain diferent amounts of anhydrides. Their
applications in syntheses is a good alternative to boronic acids [17]. They are easy-to-use
compounds, in which protecting groups can be removed from the product after completion
of reaction. Knowledge of the solubility of such derivatives in typical organic solvents is
therefore of great importance. However, there is a lack of such data in the literature.
There are several reasons why reliable solubility data for boronic acids are difcult to
obtain in comparison with common organic compounds. One of them is an easy dehydra-
tion process of boronic acids with the formation of boroxines (Scheme 1).
HO
OH
B
B
3
O
B
O
B
3 H2O
+
O
Scheme 1 Dehydration of phenylboronic acid with the formation of boroxine
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Journal of Solution Chemistry
As a result of this reaction, commercial boronic acids are mixtures of acid and anhy-
dride containing variable amounts of both components. The equilibrium of the dehydration
reaction depends on the substituents in the aryl ring and on the solvent. Due to the signif-
cant diferences in the solubility of the acid and anhydride, it is difcult to obtain repeat-
able solubility data. Another factor that causes difculties in solubility determination of
boronic acids is their ability to form micellar systems [18]. Moreover, experimental deter-
mination of basic thermodynamic parameters, such as melting temperature or enthalpy
of fusion, needed to describe the solid-liquid equilibrium by some models, is difcult for
boronic acids. These compounds, when heated up in an open vessel, frst undergo a dehy-
dration process, often a multistage one, so the melting parameters are actually measured for
the anhydride [19]. As mentioned above, compared to acids, esters and azaesters are stable
compounds with a specifc composition. In the present work we investigated cyclic ester 2
and azaester 3 shown in Fig. 1.
The aim of the present work is the solubility determination of phenylboronic acid 1 as
well as its pinacol- (2) and azaester (3) in common organic solvents and to show possibility
of using typical correlation equations to describe the experimental data.
2 Experimental Section
2.1 Chemicals Used
Phenylboronic acid (Alfa Aesar, ≥98%) (1) was crystallized from water and dried on air.
Pinacol ester (2) (Fluorochem, 98%) was used without further purifcation. Azaester (3)
was synthesized and purifed according to [20] (see Supplementary material). Purity of
1
the compounds was checked by H and 11B NMR (see Supplementary material). Solvents
(chloroform, Chempur, 98.5%; acetone, Chempur, 99.5%; 3-pentanone, Alfa Aesar, 99%;
dipropyl ether, Alfa Aesar, 99%; methylcyclohexane, Sigma-Aldrich, ≥99%) were dried
before use using conventional methods.
2.2 Experimental Procedure
Solubility was determined using a synthetic method described in [21]. A biphasic sample
of known composition (accuracy of 0.00016 of mole fraction) was slowly heated while
stirred rigorously and disappearance of turbidity was taken as the solid-liquid equilibrium
Fig. 1 The compounds investi-
gated in the present work
HO
OH
B
O
O
N
B
B
O
O
1
2
3
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Journal of Solution Chemistry
point. The extent of turbidity was measured by the intensity of a light beam crossing the
measuring cell. An ordinary LED (white light, 20°, 10000 mcd) was used as a source of
light. Light intensity was measured by the LP 471 LUM 2 luminance probe supplemented
by Photo-Radiometer HD2102.1 (DeltaOHM, Italy). The temperature was measured using
a Pt100 resistance probe P 650 (DOSTMANN electronic, Germany) with an accuracy of
0.03 K and stability of 0.01 K. Heating rate was 1 K·h−1. It was checked that reduction
of heating rate down to 0.3 K·h−1 did not infuence signifcantly the determined solubility
temperature. Measured temperatures and light intensities were acquired independently and
transformed into the light intensity versus temperature dependence. This dependence was
correlated by two straight lines and fltered. A typical observed dependence is shown in
Supplementary material (Fig. S4). Standard uncertainty of the solubility temperature was
estimated to be 0.5 or 1 K. The latter value concerns solubilities of azaester in chloroform
or 3-pentanone as for these systems relatively high scattering of the data was observed,
and for the azaester+methylcyclohexane mixture with very low solubility and, as a result,
a very steep solubility curve. The results were correlated using Redlich–Kister, Wilson,
and NRTL models [22–24]. For the compound 2 better correlation was achieved with
polynomials.
The phase transitions and thermal stability of the samples were studied using diferen-
tial scanning calorimetry (DSC). The DSC data were obtained using Q200 scanning cal-
orimeter (TA Instruments). Heat fow calibration was performed using standard sapphire
samples. The temperature and cell constant were calibrated with high purity by mass frac-
tion (0.99999) indium. Samples were placed in aluminum Tzero unhermetic pans. An empty
pans was used as the reference. Data analysis was carried out using the TA Universal Anal-
ysis application. Three cycles of heating, cooling and heating at 10 K·min−1 for each scan
were carried out in the fow of nitrogen.
3 Results and Discussion
For the solubility measurements of boronic compounds 1–3, four types of organic solvents
were chosen, namely: chloroalkane (chloroform), cyclic alkane (methylcyclohexane), ether
(dipropyl ether) and ketone (acetone or 3-pentanone). The results are shown in Figs. 2, 3
and 4 and raw experimental data are given in Table S1 of Supplementary material.
The best solubility of phenylboronic acid is observed in dipropyl ether. These data
confrm that linear or cyclic ethers are the best solvents for the synthesis or reactions of
boronic acids. Lower solubility is observed for the ketone. The solubility in chloroform is
relatively good, and the slope of this curve suggests the possibility of using this solvent for
the crystallization of boronic acids. Very low solubility in methylcyclohexane confrms the
usefulness of hydrocarbons to remove organic impurities in the synthesis of boronic acids.
Attempts to determine the solubility of the acid in 3-pentanone were unsuccessful due to
the discrepancy between data. However, it was possible to obtain reproducible results in
the case of acetone.
The curves shown in Fig. 2 show an unusual course: they intersect with each other.
Moreover, the dispersion of results is much larger than for other compounds. There are
several reasons of such behavior. Phenylboronic acid undergoes an equilibrium dehydra-
tion reaction (Scheme 1). This means that not only two components (solute–boronic acid,
and solvent) are present in the system, but also boroxine and water. The completely dif-
ferent nature of the acid and anhydride causes signifcant diferences in their solubility,
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Journal of Solution Chemistry
Fig. 2 Solubility temperatures of phenylboronic acid (1) as a function of its mole fraction (x1) in saturated
solution of the solvents listed on the chart. Solid lines: calculated by Eq. 2
Fig. 3 Solubility temperatures of pinacol ester of phenylboronic acid (2) as a function of its mole fraction
(x1) in saturated solution of the solvents listed on the chart. Solid lines: correlations obtained with the Wil-
son equation. Ideal solubility is indicated by a dotted curve
depending on the type of solvent. Moreover, the miscibility of water being formed with
some solvents, e.g. acetone, in contrast to its very low solubility in others, e.g. cyclohex-
ane, causes a considerable complexity of the system studied.
Solubilities of pinacol ester of phenylboronic acid are shown in Fig. 3.
The esterifcation of the boron group with pinacol signifcantly increases solubility
compared to the parent acids: very high solubility was observed for all solvents used even
at low temperatures. This indicates the possibility of using pinacol esters of boronic acids
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Journal of Solution Chemistry
Fig. 4 Solubility temperatures of azaester of phenylboronic acid (3) as a function of its mole fraction (x1)
in saturated solution of the solvents listed on the chart. Solid lines: correlations obtained with the Wilson
equation. Ideal solubility is indicated by a dotted curve
to carry out reactions in any type of organic solvents. The diferences between particular
solvents are small. The highest solubility was observed in chloroform and the lowest in
methylcyclohexane.
Solubility of azaester of phenylboronic acid is shown in Fig. 4.
Compared to pinacol ester, for azaester there are signifcant diferences between indi-
vidual solvents: the compound shows very good solubility in chloroform slightly weaker
in 3-pentanone, but much lower in methylcyclohexane and dipropyl ether. The tendency is
similar to that of pinacol esters: better solubility appears in polar solvents.
The above solubility data were described by typical GE-models, namely: the Wilson,
NRTL and Redlich–Kister equations. To perform these correlations, it was necessary to
know basic thermodynamic parameters, i.e. melting temperatures and enthalpies of fusion.
These data for the investigated compounds were determined by the DSC method (see Figs.
S5–S7 in supplementary material) and are collected in Table 1.
Contrary to the esters 2 and 3, which melt congruently without decomposition, boronic
acid dehydrates into boroxine during heating according to Scheme 1. The temperature of
about 220 °C reported in the literature as the melting point of phenylboronic acid, actually is
temperature of melting of triphenylboroxine. The data given in Table 1 for 1 were estimated
from the DSC diagram assuming that the frst peak at ca. 90 °C corresponds to the melting of
this compound, while that at ca. 108 °C corresponds to the dehydration. Due to the complex-
ity of acid-containing systems, the correlation of data using a polynomial is most justifed.
Table 1 Melting temperatures
Phenylboronic Pinacol ester (2)
acid (1)
Azaester (3)
and enthalpies of fusion of
investigated compounds
ΔHm/kJ·mol−1
T m/ oC
14.90a
90.04a
13.26 0.1
26.17 0.5
44.83 0.4
131.03 0.5
aEstimated value
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Journal of Solution Chemistry
Correlation curves for the compounds 1–3 are shown in Figs. 2, 3 and 4. Additionally, ideal
solubility given by Eq. 1 with γ1=1 was determined for the compounds 2 and 3.
The model description of the solubility data is based on the solubility equation which
directly results from the condition of phase equilibrium, i.e. equality of chemical potentials
of a compound existing in solid and liquid phase. For a system in which pure solid (1) is in
equilibrium with liquid mixture and additionally if (i) no transition points in the solid phase
appear, (ii) a solute is stable up to the melting temperature, and (iii) the enthalpy of melting is
independent of temperature, the solubility is expressed by the following general equation.
(
)
[
(
)]
1
1
R ln x1ꢀ1 x1, T = −ΔHm1
−
(1)
T
Tm1
where x1, γ1, ΔHm1, and Tm1 are the mole fraction, activity coefcient, enthalpy and tem-
perature of melting of the component (1), respectively. It should be noted that this equation
cannot be used for the solubility of phenylboronic acid since condition (ii) is not fulflled
for this case. If the activity coefcient is equal to one, the ideal solubility is obtained which
is drawn in Figs. 3 and 4 as a dotted line. For the pinacol ester + {3-pentanone or chloro-
form} systems, rather low negative deviations from ideality are observed what manifests
itself by solubilities slightly higher than ideal ones. Usually, such behavior is explained
by attractive interactions between unlike molecules higher than those occurring in pure
components. The character of deviations from ideality for other solvents is not obvious
as they are less signifcant but probably they are slightly positive. For the systems with
azaesters as solute this trend is repeated with the deviations from ideality being more sig-
nifcant and diferences between various solvents considerably greater. The positive devia-
tions from ideality increase from dipropyl ether to methylcyclohexane and negative ones
increase from 3-pentanone to chloroform. In the frst two systems relatively weak inter-
actions between the solute and the solvent are expected because of the non-polar or only
slightly polar character of the solvents. For the last one, a strong interaction in the form of
hydrogen bond between chloroform and azaester is probably responsible for the observed
high solubility. In these hydrogen bonds chloroform’s molecule acts as a hydrogen bond
donor.
In the approach expressed by Eq. 1, a model used to describe solubility data concerns in
fact the dependence of the solute activity coefcients as a function of concentration and tem-
perature. Since this dependence results directly from the excess Gibbs energy (GE), it is the
function which defnes the model and, as a consequence, the solubility equation. As the model
parameters are being adjusted to the experimental data, they have mainly correlational charac-
ter and their physical signifcance is restricted rather to qualitative considerations. The follow-
ing GE models were used to represent solubility data:
The Redlich–Kister equation [22]:
n
ꢀ
GE = x1x2
ai(x1 − x2)i
(i=0)
where n is the order of the GE/x1x2 polynomial, that is n+1 determines the number of
model parameters.
The Wilson equation [23]:
(
)
(
)
GE∕RT = −x1 ln x1 + x2ꢀ12 − x2 ln x2 + x1ꢀ21
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Journal of Solution Chemistry
(
)
(
)
(V20)
(V10)
(V10)
(V20)
Δg12
Δg21
ꢀ12
=
exp
ꢀ21
=
exp
RT
RT
and the NRTL equation [24]:
[
]
G21Δg21
G12Δg12
GE = x1x2
+
x1 + x2G21 x2 + x1G12
where
Δg12 = g12 − g22
(
)
(
)
Δg12
Δg21
,
Δg21 = g12 − g11
,
G12 = exp −ꢀ
,
G21 = exp −ꢀ
RT
RT
The ai, α, Δg12 and Δg21 parameters are independent of x1 and T. The latter two
parameters are interpreted as residuals between energy of interaction between like and
unlike molecules. The nonrandomness parameter α has been fxed to 0.2. The ratio of
/
molar volumes of pure components V0 V10 which appears in the Wilson equation was
assumed to be equal to one, because of2difculties in the estimation of the molar volume
of a pure solute in the hypothetical liquid state. Nevertheless we do not recognize con-
tributions as signifcant, since the Δg12 and Δg21 parameters can incorporate the volume
ratios while being adjusted to the experimental data. The Wilson and the NRTL equa-
tions possess a certain theoretical background in spite of some crude simplifcations
included. The Redlich–Kister equation does not have any theoretical support as it is
a simple power series. However, it has one practical advantage over other models—a
variable number of adjustable parameters which makes the model fexible and being
able to describe even complicated dependencies. The adjusted parameters with the
standard deviations are given in Table S2 of the Supplementary Material. The number
of the Redlich–Kister parameters was established in such a way to guarantee the data
reproduction with standard deviations close to the experimental uncertainty, i.e. 1 K for
the azaester + {chloroform, or 3-pentanone, or methylcyclohexane} systems and 0.5 K
for other mixtures. Two adjustable parameters turned out to be sufcient except for the
azaester + chloroform system, for which three parameters were necessary. All the mod-
els give fairly similar results, although some diferences are noticeable. Surprisingly, the
Redlich–Kister equation, in spite of its mainly empirical character, gives standard devia-
tions that are most consistent with experimental uncertainties. The superiority of the
Redlich–Kister equation in the description of the azaester+ chloroform system results
from three adjustable parameters against two for other models. If the two-parameter
Redlich–Kister equation is used, all models give almost the same standard deviations.
The solubilities in systems with phenylboronic acid (1) as a solute cannot be
described by Eq. 1., since this compound decomposes into anhydride before reaching
its temperature of melting. As the thermodynamic characterization of its decomposi-
tion is unknown, only purely empirical equation can be applied. These solubilities were
described by polynomial in the form:
n
ꢀ
T =
aix1n
(2)
i=0
The polynomial order was established in a similar way as the number of the
Redlich–Kister parameters although a strict procedure could not be applied because of
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Journal of Solution Chemistry
Fig. 5 Solubility of the compounds 1–3 at Tr = 0.89 (i.e. 50.09 oC for 1, -6.76 oC for 2, 86.57 oC for 3). The
values were calculated from the correlation equations
a limited number of experimental data. The values of adjusted parameters together with
standard deviations are shown in Table S2 of the Supplementary Material.
In order to compare the solubility of investigated compounds in particular solvents,
solubility values for the same reduced temperature (Tr = T/Tm1). The equilibrium tem-
peratures vs. relative permittivity of the solvent [25], calculated for Tr = 0.89 using the
correlation equations, are shown in Fig. 5.
For the phenylboronic acid 1 solubility increases with the increasing value of relative
permittivity with the exception of dipropyl ether. The same tendency is for both esters
(2 and 3) with the exception of chloroform.
4 Conclusions
Phenylboronic acid reveals high solubility in ether and ketone and very low in hydro-
carbon. Pinacol ester and azaester show better solubility than parent acid in all tested
solvents. For pinacol ester diferences between particular solvents are small, while for
azaester the diferences are signifcant. For both esters the highest solubility is observed
in chloroform and the lowest in hydrocarbon. Generally, the solubility positively corre-
lates with the polarity of solvents although some exceptions can be noted.
For the phenylboronic acid better correlation of the data is obtained by polynomi-
als in comparison with the Wilson, NRTL or Redlich–Kister equations, what can be
explained by an additional equilibrium reaction with the formation of anhydride. For the
esters, the Redlich–Kister equation gives standard deviations which are the most con-
sistent with experimental uncertainties.
Acknowledgements Support of Warsaw University of Technology is kindly acknowledged.
Funding The funding was provided by Wydzial Chemiczny Politechniki Warszawskiej (Grant No.
504/04113/1020/44.000000)
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Journal of Solution Chemistry
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