Effects of alkyl chain length and solvents on thermodynamic dissociation constants of the ionic liquids with one carboxyl group in the alkyl chain of imidazolium cations

Article abstract of DOI:10.1021/jp501731j

Thermodynamic dissociation constants of the Br??nsted acidic ionic liquids (ILs) are important for their catalytic and separation applications. In this work, a series of imidazolium bromides with one carboxylic acid substitute group in their alkyl chain ([{(CH₂)_nCOOH}mim]Br, n = 1,3,5,7) have been synthesized, and their dissociation constants (pK _a) at different ionic strengths have been determined in aqueous and aqueous organic solvents at 0.1 mole fraction (x) of ethanol, glycol, iso-propanol, and dimethyl sulfoxide by potentiometric titrations at 298.2 K. The standard thermodynamic dissociation constants (pK_a^T) of the ILs in these solvents were calculated from the extended Debye-Huì?ckel equation. It was found that the pK_a values increased with the increase of ionic strength of the media and of the addition of organic solvent in water. The pK_a^T values also increased with the increase of the alkyl chain length of cations of the ILs. In addition, the effect of solvent nature on pK_a^T values is interpreted from solvation of the dissociation components and their Gibbs energy of transfer from water to aqueous organic solutions. ? 2014 American Chemical Society.

Full text of DOI:10.1021/jp501731j

Article
pubs.acs.org/JPCB
Effects of Alkyl Chain Length and Solvents on Thermodynamic
Dissociation Constants of the Ionic Liquids with One Carboxyl Group
in the Alkyl Chain of Imidazolium Cations
Yuehua Chen, Huiyong Wang, and Jianji Wang*
Collaborative Innovation Center of Henan Province for Green Manufacturing of Fine Chemicals, Key Laboratory of Green Chemical
Media and Reactions, Ministry of Education, School of Chemistry and Chemical Engineering, Henan Normal University, Xinxiang,
Henan 453007, People’s Republic of China
ABSTRACT: Thermodynamic dissociation constants of the Brønsted acidic
ionic liquids (ILs) are important for their catalytic and separation applications.
In this work, a series of imidazolium bromides with one carboxylic acid
substitute group in their alkyl chain ([{(CH2)nCOOH}mim]Br, n = 1,3,5,7)
have been synthesized, and their dissociation constants (pKa) at different ionic
strengths have been determined in aqueous and aqueous organic solvents at 0.1 mole fraction (x) of ethanol, glycol, iso-propanol,
and dimethyl sulfoxide by potentiometric titrations at 298.2 K. The standard thermodynamic dissociation constants (pKaT) of the
ILs in these solvents were calculated from the extended Debye−Huckel equation. It was found that the pK values increased with
̈
a
the increase of ionic strength of the media and of the addition of organic solvent in water. The pKaT values also increased with the
increase of the alkyl chain length of cations of the ILs. In addition, the effect of solvent nature on pKaT values is interpreted from
solvation of the dissociation components and their Gibbs energy of transfer from water to aqueous organic solutions.
function of IL concentrations because the determinations were
not performed at constant ionic strength.
1. INTRODUCTION
In the past two decades, room temperature ionic liquids (ILs)
have aroused considerable attention due to their superior
physicochemical properties such as no vapor pressure and
nonflammability under ambient conditions, wide electro-
chemical window, high chemical and thermal stability, large
electrical conductivity, and unusual solubility for many
compounds. Thus, ILs have been widely applied in organic
synthesis,1,2 nanotribology,3,4 biocatalysis,5 energetic materials,6
analytical and separation science,7,8 and so on.
In the present work, we synthesized a series of four
imidazolium bromides with one carboxylic acid substituent
group in the alkyl chain of their cations, i.e., 1-alkylcarboxylic
acid-3-methylimidazolium bromides ([{(CH2)nCOOH}mim]-
Br, n = 1,3,5,7). For the sake of easy understanding, the
chemical structure of these ILs is shown in Figure 1. The pKa
Among these ILs, those based on the imidazolium cation
with carboxylic acid groups in their alkyl chain have been used
to selectively dissolve metal oxides9,10 and metal hydroxides,11
efficiently catalyze the synthesis of cyclic carbonates,12,13 and
significantly stabilize the metal nanoparticle in aqueous
solution.14 Apparently, these applications are related with the
acidic nature of carboxyl group in these ILs. Therefore,
knowledge of acid−base property is very important for their
potential applications. It is known that thermodynamic
dissociation constants are important physicochemical parame-
ters to reflect the acid−base behavior of Brønsted acidic ILs in
solutions. However, it is surprising to find that only two articles
have been published for the determination of pKa values of the
ILs with carboxyl groups in the alkyl chain of their cations. In
this context, Fei et al.15 reported the pKa values of the ionic
liquids ([{(CH2)nCOOH}mim]X, n = 1, 3; X = Cl, BF4,
CF3SO3) and ([{(CH2)nCOOH}2im]X, n = 1,3; X = BF4,
CF3SO3) in aqueous solution at the ionic strength of
approximate 0.01 mol/L by potentiometric titrations. Geng et
al.16 also measured the pKa values of some polycarboxyl
imidazolium ILs (such as [{(HOOCCH2CHCOOH}eim]Br)
in water by the same methodology. These pKa values are a
Figure 1. The structures of the four ionic liquids.
values of these ILs at different ionic strengths have been
determined by potentiometric titration in aqueous solution and
in aqueous organic solvent mixtures at 298.2 K due to its
simplicity and accuracy. Ethanol (EtOH), glycol (GL), iso-
propanol (i-PrOH), and dimethyl sulfoxide (DMSO) are
selected as organic additives for of the following reasons: (i)
they have strong ability to solubilize slightly soluble or insoluble
compounds in water and thus have been widely used in
analytical and separation processes; (ii) they are fully miscible
with water but less polar than water; and (iii) they are potential
green solvents based on the classification of Moity and his co-
workers.17 The mole fraction (x) of organic compounds in the
solvent mixtures is kept at 0.1. The thermodynamic dissociation
Received: February 18, 2014
Revised: April 9, 2014
Published: April 10, 2014
© 2014 American Chemical Society
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Table 1. pKa Values of the ILs 1 and 2 in Different Solvents at 298.2 K as a Function of Ionic Strength (I)
IL 1
IL 2
I (mol·L−1
)
H2O EtOH + H2O glycol + H2O i-PrOH + H2O DMSO + H2O H2O EtOH + H2O glycol + H2O i-PrOH + H2O DMSO + H2O
0.10
0.30
0.50
0.80
1.0
2.02
2.06
2.08
2.11
2.14
2.25
2.27
2.32
2.28
2.28
2.32
2.46
4.03
4.05
4.08
4.12
4.15
4.32
4.34
4.39
4.27
4.40
4.45
4.57
2.32
2.40
2.44
2.47
2.51
2.49
2.51
2.54
2.56
2.59
4.32
4.40
4.46
4.48
4.51
4.61
4.63
4.65
4.68
4.71
2.37
2.41
2.36
2.39
2.41
4.44
4.48
4.49
4.53
4.55
1.2
1.5
2.19
4.21
constants (pKaT) have been obtained by nonlinear fitting of the
meter, equipped with a pH combined electrode in a water-
jacketed glass vessel. The temperature in the glass vessel was
kept at 298.2 0.1 K by circulating water from a thermostat.
Purified argon was passed through the solution to reduce the
adverse effect of carbon dioxide during titrations.
extended Debye−Huckel equation. These results have been
̈
utilized to understand the effects of the ionic strength of the
media, the alkyl chain length of the cations, and the nature of
the solvents on the acid−base property of the ILs.
Before determination of the pKa values, calibration of the
electrode system was performed by using the procedure
described in the literature.18 After calibration of the electrode,
the titration of IL solution with aqueous NaOH was conducted
under the same experimental conditions. The dilute IL
solutions prepared in water and in the mixed solvents were
titrated with 0.01 mol/L of NaOH. During the titrations, all the
titrands and titrants had the same ionic strength and solvent
composition. The volume of the titrant and the corresponding
potential of the cell were recorded after a suitable time interval
(around 2−3 min). The pKa values of the ILs were determined
in triplicate, and the standard error was within 0.02 pKa unit.
2. EXPERIMENTAL SECTION
2.1. Materials. N-methylimidazole (Linhai Kaile Chem.
Co., 99%) was distilled before use. 2-Bromoacetic acid (Aladdin
Chem. Co. Ltd., 98.5%), ethyl 4-bromobutyrate (Aladdin
Chem. Co. Ltd., 95%), 6-bromohexanoic acid (Aladdin Chem.
Co. Ltd., 97%), 8-bromooctanoic acid (Alfer Aser, 97%), and
hydrobromic acid (Aladdin Chem. Co. Ltd., 48%) were used
without further purification. Ethanol (Shanghai Chem. Co.),
glycol (Shanghai Chem. Co.), iso-propanol (Tianjin Guangfu
Chem. Co.), and dimethyl sulfoxide (Shanghai Chem. Co.)
were analytical grade and used after drying over 4 A-type
molecular sieves (Shanghai Chem. Co.).
A stock aqueous solution of carbonate-free NaOH (Shanghai
Chem. Co, AR) was prepared and standardized with potassium
hydrogen phthalate (KHP, PT, Aladdin Chem. Co. Ltd.). The
concentration of HCl (Shanghai Chem. Co, AR) was
determined by titration with standard aqueous NaOH solution.
In order to reduce the experimental error, stock solutions of the
ILs were prepared, and their accurate concentrations were
determined by potentiometric titration. Dilute solutions were
prepared by appropriate dilution of the stock solutions. All the
aqueous organic solvents were prepared by weight and the
mole fraction of the organic compounds was fixed at 0.1. The
ionic strength in all the solutions was adjusted with KCl
(Shanghai Chem. Co, AR). Conductivity water with a
conductivity of 1.2 μΩ−1 cm−1 was prepared by redistilling
the deionized water and subsequently boiled to remove CO2
before use.
2.2. Synthesis of the ILs. All the ILs, except for IL 2, were
synthesized in acetonitrile or tetrahydrofuran through the
reaction of N-methylimidazole with excess bromine substituted
carboxylic acid under reflux at about 65 °C. After the reaction,
solvents were removed by rotary evaporation, and then the
crude products of the ILs 1, 3, and 4 were washed with
acetonitrile, acetone, and diethyl ether, respectively. The IL 2
was prepared in a similar procedure described in the
literature.15 These ILs were dried under vacuum at 50 °C for
3−4 days in the presence of P2O5. The melting points of the
ILs 1, 2, 3, and 4 were 182.4, 99.6, 109.5, and 124.6 °C,
respectively, as measured by a NETZSCH 204 F1 differential
scanning calorimetry. 1H NMR spectra of the ILs were
determined by a Bruck AV-400 spectrometer.
3. RESULTS AND DISCUSSION
3.1. Dissociation Constants and Thermodynamic
Dissociation Constants of the ILs. Dissociation equilibrium
of the ILs in solution can be described in terms of the following
single process:
R+(CH2)nCOOH + SH ⇌ R+(CH2)nCOO− + SH+2
(1)
+
where SH represents a solvent, and SH2 denotes a solvated
proton, and R+(CH2)nCOOH and R+(CH2)nCOO− refer to the
acid and conjugate base forms of the ILs, respectively.
According to the above equilibrium, the dissociation constant
can be expressed by
Ka = [R+(CH2)nCOO−][SH+2 ]/[R+(CH2)nCOOH]
(2)
At any point of the titration, the mass and charge balance are
expressed by the equations:
+
−
C
= CBr
[R (CH2)nCOOH]Br
= [R+(CH2)nCOOH]
+ [R+(CH2)nCOO−]
(3)
(4)
[Na+] + [SH+2 ] + [R+(CH2)nCOOH] = [S−] + [Br−]
In these equations, C stands for the total concentration of each
IL at any moment of titration. From eqs 2, 3, and 4, we have
pKa = −lg[SH+2 ] − lg{([Na+] + [SH+2 ] − [S−])
2.3. Determination of the Dissociation Constants. All
titrations of the aqueous ILs solution and calibration of the
electrode were performed on an ORION 720A digital pH
/(C − [Na+] − [SH+2 ] + [S−])}
(5)
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Table 2. pKa Values of the ILs 3 and 4 in Different Solvents at 298.2 K as a Function of Ionic Strength (I)
IL 3
IL 4
I (mol·L−1
)
H2O EtOH + H2O glycol + H2O i-PrOH + H2O DMSO + H2O H2O EtOH + H2O glycol + H2O i-PrOH + H2O DMSO + H2O
0.10
0.30
0.50
0.80
1.0
4.36
4.75
4.77
4.82
4.68
4.91
4.94
4.99
4.56
5.00
5.03
5.06
4.84
5.24
5.28
5.19
4.39
4.42
4.45
4.48
4.50
4.71
4.76
4.80
4.82
4.85
5.03
5.05
5.07
5.10
5.13
4.60
4.62
4.64
4.66
4.69
4.87
4.94
4.97
4.99
5.03
5.22
5.24
5.26
5.28
5.31
4.86
4.89
4.98
5.00
5.04
5.08
5.10
5.31
5.33
5.36
1.2
1.5
Table 3. pKaT Values of the ILs in Different Solvents at 298.2 K, and the b Value in eq 13 and Their Standard Error (RSD)
IL 1
IL 2
IL 3
IL 4
solvent
H2O
pKaT
b
RSD (%)
pKaT
b
RSD (%)
pKaT
b
RSD (%)
pKaT
b
RSD (%)
1.93
2.16
2.14
2.15
2.35
0.089
0.5
1.9
1.2
1.1
1.4
3.93
4.15
4.20
4.27
4.47
0.10
1.0
2.1
1.4
1.2
1.2
4.26
4.56
4.64
4.78
4.89
0.08
1.4
1.8
1.2
1.7
1.2
4.47
4.72
4.90
5.11
5.09
0.06
1.0
1.7
glycol + H2O
EtOH + H2O
i-PrOH + H2O
DMSO + H2O
0.14
0.08
0.06
0.07
0.15
0.08
0.07
0.07
0.10
0.07
0.05
0.07
0.11
0.033
0.04
0.62
1.6
0.04
1.8
Table 4. pKaT Values of the ILs and Some Carboxylic Acids in Aqueous Solution at 298.2 K
pKaT
pKaT
pKaT
pKaT
IL 1
1.93
4.75
2.82
2.82
IL 2
3.93
4.81
0.88
4.52
IL 3
4.26
4.88
0.62
IL 4
4.47
4.89
0.42
acetic acid
ΔpKaT
2-chloroacetic acid
butyric acid
ΔpKaT
4-chlorobutyric acid
hexanoic acid
ΔpKaT
octanoic acid
ΔpKaT
pKaT = pKa × lg γ
− lgγR+(CH ) COO− − lg γ
SH+2
When the titration was carried out at the pH value below 7
against a solution of 0.01 mol/L of IL, the value of [S−] can be
neglected.19 Thus,
R+(CH2)nCOOH
2
n
(9)
+
−
where γR (CH ) COO can be regarded as 1, and in the electrode
2
n
pKa = −lg[SH+2 ] − lg{([Na+] + [SH+2 ])
calibration, the lg γ(SH value is already contained in the E0′
+
2
)
/(C − [Na+] − [SH+2 ])}
term based on eq 7, and is thus included in the values of pKa. In
this case, we have
(6)
During the titration of each IL solution, the unknown values of
lg[SH+2] could be obtained by the following equation:
pKaT = pKa + lg γ
According to the extended Debye−Huckel equation:23
R+(CH2)nCOOH
(10)
(11)
lg[SH+2 ] = (E − E0′)/s
̈
(7)
lg γ
= −A I /(1 + Ba I ) + bI
̊
R+(CH2)nCOOH
where E is a measured electrode potential in the titration, s and
E0′ are the slope and the formal electrode potential, which
could be obtained by a linear correlation between electrode
where a is the ion-size parameter, and A and B are Debye−
̊
Huckel parameters, which are related with the dielectric
̈
+
potential E recorded and [SH2 ] calculated on the basis of the
constant (ε) of the solvents and the thermodynamic temper-
ature (T) in the following way:24
titration of HCl with NaOH. By using a series of volume and
concentration data of titrant and titrand, the value of pKa has
been calculated, and the results are listed in Tables 1 and 2. It
can be seen that pKa values increase with the increase of ionic
strength. Such a trend is identical with that of the protonated
triethanolamine,20 but different from monocarboxylic acids.21,22
Therefore, it seems likely that the ILs with carboxyl group in
their alkyl chain of cations could not be simply regarded as the
traditional carboxylic acids.
A = 1.8246·106/(εT)3/2
Thus, eq 10 can be transformed to eq 13:
pKa = pKaT − A I /(1 + Ba I ) + bI
B = 50.29/(εT)1/2
(12)
(13)
̊
The dielectric constants of the mixed solvents studied in the
present work were obtained by interpolation from the
experimental data reported by Åkerlof25 and Khoo.26 Thus,
In fact, the thermodynamic dissociation constant is only
dependent on temperature and can be formulated by
̈
the values of pKaT were calculated by using nonlinear least-
squares regression. It was found that using the a value of 10 nm
̊
KaT = Ka × γ
× γ /γ
SH+2 R+(CH2)nCOOH
R+(CH2)nCOO−
gave the maximum correlation coefficient (R) and the
minimum standard error. Therefore, all the pKaT values of
the ILs in water and in the aqueous organic solvents were
(8)
where γi represents activity coefficient of the species i. Taking
the logarithms on both sides of eq 8, this equation would be
changed into
calculated with a =10 nm, and the results were collected in
̊
Table 3.
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Table 5. Dielectric Constant (ε) of Different Aqueous Organic Solvents (x = 0.1) and the Values of ΔGt0, ΔGt0(H+), and
Δd{ΔGt0[R+(CH2)nCOO−] − ΔGt0[R+(CH2)nCOOH]} from Water to the Mixed Solvents at 298.2 K (in KJ·mol−1)
IL1
IL2
IL3
IL4
solvent
dielectric constant (ε)
ΔGt0(H+)
ΔGt0
Δd
ΔGt0
Δd
ΔGt0
Δd
ΔGt0
Δd
glycol + H2O
EtOH + H2O
i-PrOH + H2O
DMSO + H2O
70.4
65.7
59.4
76.4
−0.86
−2.37
−3.74
−6.65
1.31
1.20
1.26
2.40
2.17
3.57
5.00
9.05
1.26
1.54
1.94
3.08
2.12
3.91
5.68
9.73
1.71
2.17
2.97
3.60
2.57
4.54
6.71
10.2
1.43
2.45
3.65
3.54
2.29
4.79
7.39
10.2
ΔGt0 = ΔGt0(H+) + ΔGt0[R+(CH2)nCOO−]
− ΔGt0[R+(CH2)nCOOH]
3.2. Effect of Cationic Alkyl Chain Length on the
Dissociation Constants of the ILs. From Tables 1−3, we
can clearly see that at a given solvent and a given ionic strength,
the values of pKa and pKaT increase with the increase of alkyl
chain length of the ILs. Similar trend has been reported for the
dissociation constants of saturated fatty acids in aqueous
solution.27 This result can be interpreted from the fact that the
longer alkyl chain may produce inductive effect, resulting in the
enhanced electrostatic attraction between oxygen and hydrogen
in carboxyl group and the decreased ability to release proton. It
is interesting to note from Table 4 that the pKaT values for the
ILs in water are smaller than those of carboxylic acids with the
same alkyl chain length. This certainly results from the presence
of an imidizolium ring in the ILs, the positive charge of which
reduces the electrostatic interaction of the oxygen atom with
hydrogen atom in carboxyl group, thus increasing acidity of the
ILs. With the increase of alkyl chain length, such an effect
becomes smaller, which leads to a decreased difference in pKaT
value (ΔpKaT) between the IL and the corresponding
carboxylic acid in water, as indicated by the data in Table 4.
In addition, to qualitatively evaluate the ability of electron-
withdrawing of the imidizolium ring, the pKaT values of 2-
chloroacetic acid and 4-chlorobutylic acid in aqueous solution
are also included in Table 4 for comparison. It can be seen that
pKaT values of the ILs 1 and 2 are respectively smaller than
those of 2-chloroacetic acid and 4-chlorobutyric acid. This
suggests that electron-withdrawing ability of the positive charge
on the imidizaolium ring is much stronger than that of the
chlorine atom in the chlorinated carboxylic acids.
3.3. Solvent Effects on Dissociation Constants of the
ILs. It is evident from Tables 1−3 that pKaT and pKa values of
the ILs in aqueous organic solutions become greater than that
in water under the same conditions. A similar trend has been
reported for monocarboxylic acids such as acetic acid, benzoic
acid, and others.28 This indicates that acid−base property of the
ILs changes with addition of organic solvents in water, and the
dissociation of the acidic ILs is not favorable in the presence of
organic solvent. It is known that the change in thermodynamic
dissociation constants results from the difference in the
thermodynamic properties of the species participated in
dissociation reaction in water and in mixed solvents. Thus,
the difference in pKaT value of the ILs in two different solvents
is closely related to the standard Gibbs energy of transfer, ΔGt0,
for the dissociation reaction of the ILs:29
(15)
where ΔGt 0 (H+ ), ΔGt 0 [R+ (CH2 )n COO− ], and
ΔGt0[R+(CH2)nCOOH] stand for the standard Gibbs energies
of transfer for H+, R+(CH2)nCOO− and R+(CH2)nCOOH from
water to the mixed solvents. Among these terms, the ΔGt0
values can be obtained from our standard dissociation constant
values according to eq 14, and the ΔGt0 (H+) values have been
calculated from the experimental data of Khoo26 and Wells30,31
by interpolation. Then the difference between
ΔGt0[R+(CH2)nCOO¯] and ΔGt0[R+(CH2)nCOOH], Δd, can
be calculated, and the results are given in Table 5. The negative
ΔGt0 (H+) values indicate that the proton is more stable in the
mixed solvents than in water. This can be understood from the
much higher values of Kamlet−Taft β parameter of DMSO (β
= 0.76), i-PrOH (β = 0.95), EtOH (β = 0.77) and glycol (β =
0.52) than that of water (β = 0.18),32 which has been
established as a measure of the hydrogen bond accepting ability
or the hydrogen bond basicity of liquid solvents and a higher β
parameter value indicates a stronger hydrogen bond basicity of
the solvent. However, the positive ΔGt0 values show that the
solvent effect on the proton does not account for the increase
of pKaT values in the aqueous organic solution, and the solvent
effect on the charged ions R+(CH2)nCOOH and its conjugated
base R+(CH2)nCOO¯ plays a dominant role in the standard
Gibbs energy of transfer for the dissociation reaction of the ILs.
Unfortunately, the exact values of transfer energies for both of
them are not available in literature. Therefore, we have to
qualitatively demonstrate the solvent effect in thermodynamic
dissociation constants.
It is noted from Table 5 that the values of
{ΔGt0[R+(CH2)nCOO−] − ΔGt0[R+(CH2)nCOOH]} are
increasingly positive in the aqueous organic solutions with
the sequence: DMSO > i-PrOH > EtOH > GL. The positive
values indicate that values of ΔGt0[R+(CH2)nCOO−] are
greater than those of ΔGt0[R+(CH2)nCOOH] when the two
species transfer from water to the mixed solvents. In other
words, compared to the cationic form R+(CH2)nCOOH, the
ampholyte R+(CH2)nCOO− is stabilized more strongly in
water. Actually, due to the presence of two opposite charges,
R+(CH2)nCOO− can be treated as a dipole ion similar to amino
acid. Previous investigations showed that solvent property had a
significant impact on the existence form of amino acids. Hughes
et al.33 reported that the ratio of dipolar ion to uncharged form
of the amino acids decreased from 104−105 in water to 2−40 in
DMSO, suggesting that water is a favorable solvent for dipolar
ion of the amino acids. They ascribed this difference to the
stronger solvation of the carboxylate anion in water than in
T
ΔGt0 = 2.303RT[pKaT(S) − pKa (W)]
(14)
In the above equation, pKaT(S) and pKaT(W) are the standard
thermodynamic dissociation constants of ILs in mixed solvent
and in water, respectively. According to this equation, if all
participation of the solvent in eq 1 is disregarded, ΔGt0 can be
represented by the following equation:29
DMSO. Dog
̆
an et al.34 also reported that this ratio is in the
range of 103.5−105 in aqueous ethanol mixtures (20−80%
ethanol by volume), and the microscopic protonation constant
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(k22) of amino acids increases due to the weaker solvation of
dipolar ions by ethanol. From these investigations, it is
reasonable to suggest that the solvation of dipolar ion of
amino acids is mainly determined by the solvation of the
carboxylate anion because water or alcohol may provide
stronger hydrogen bonding to stabilize these anions. However,
it should be emphasized that unlike H+ solvation, solvation of
carboxylate anion is dependent on the hydrogen bond donating
capacity of the solvents, which can be characterized by the
Kamlet−Taft α parameter. It is known that the α parameter for
DMSO, i-PrOH, EtOH, glycol, and water is 0, 0.76, 0.83, 0.90
and 1.17, respectively.32 The greater the α parameter, the
stronger the hydrogen bond donating capacity of the solvent.
Therefore, it is clear that the solvation strength of carboxylate
anion by the studied solvents decreases in the order H2O > GL
> EtOH > i-PrOH > DMSO. This indicates that Gibbs energies
for the transfer of carboxylate anion and thus R+(CH2)nCOO−
from water to the studied solvents are positive, and the
ΔGt0[R+(CH2)nCOO−] values should decrease in the order
DMSO > i-PrOH > EtOH > GL.
4. CONCLUSION
In this work, four imidazolium bromides with one carboxylic
acid substitute group in their alkyl chain have been synthesized,
and their dissociation constants and thermodynamic dissocia-
tion constants have been determined in water and aqueous
organic solvents at 0.1 mole fraction of ethanol, glycol, iso-
propanol, and dimethyl sulfoxide by potentiometric titrations at
298.2 K. It was found that the values of pKa increased with the
increase of ionic strength of the media, the addition of organic
solvent in water, and the alkyl chain length of cations of the ILs.
The acidity of these ILs in water is stronger than that of
carboxylic acids with the same alkyl chain length, and the acid−
base property of the ILs studied here is analogous to that of
protonated amino acid. From the viewpoint of chemical
thermodynamics, the solvation of the charged ions
R+(CH2)nCOOH and its conjugated base R+(CH2)nCOO¯ is
predominant for the solvent effect in the thermodynamic
dissociation constants of the ILs. These findings suggest that
the dissociation constants of these ILs can be tuned by
changing the alkyl chain length of cations of the ILs, using
appropriate additives such as alcohol and DMSO and/or
modulating of ionic strength of the media to meet the needs of
practical application.
According to the previously reported ΔGt0(i) values for the
transfer of cations and anions from water to aqueous organic
solvents, it is generally known that the ΔGt0(i) is positive for
anions, and negative for most of cations.35 The ΔGt0(H+)
values presented in Table 5 confirm this observation. Similarly,
we may infer that the ΔGt0[R+(CH2)nCOOH] values would be
negative for the transfer of R+(CH2)nCOOH from water to the
mixed solvents. Based on the study of Popovych,36 cations in all
polar solvents are solvated by interaction with the negative end
of the solvent dipoles to form coordination bonds with the lone
pair electrons from oxygen, nitrogen, and sulfur atoms. For a
given cation, the electron density at the oxygen atom of organic
solvents is mainly responsible for the magnitude of cation
solvation. DMSO and i-PrOH molecules have similar molecular
structures, and their difference lies in the SO group in
DMSO and the C−OH group in i-PrOH. The existence of the
SO group results in a higher electron density on the oxygen
atom in DMSO compared with the C−OH group in i-PrOH
molecules. Therefore, solvation of R+(CH2)nCOOH species in
DMSO−water would be stronger than that in aqueous i-PrOH
solution, thus leading to greater −ΔGt0[R+(CH2)nCOOH]
values for the transfer from water to aqueous DMSO solvent.
Due to the inductive effect of alkyl group of alcohol molecules,
the electron density at the oxygen atom decreases in the order
AUTHOR INFORMATION
■
Corresponding Author
*E-mail address: Jwang@htu.cn. Phone:+86-373-3325805; Fax:
+86-373-3329030.
Notes
The authors declare no competing financial interest.
ACKNOWLEDGMENTS
■
This work was supported by the National Natural Science
Foundation of China (Grant No. 21273062, 21203057, and
21133009).
REFERENCES
■
(1) Wasserschein, P.; Welton, T. Ionic Liquids in Syntheses; VCH-
Wiley: Weinheim, Germany, 2003.
(2) Hallett, J. P.; Welton, T. Room-temperature ionic liquids:
Solvents for synthesis and catalysis. 2. Chem. Rev. 2011, 111, 3508−
3576.
(3) Asencio, R. A.; Cranston, E. D.; Atkin, R.; Rutland, M. W. Ionic
Liquid Nanotribology: Stiction Suppression and Surface-Induced
Shear Thinning. Langmuir 2012, 28, 9967−9976.
i - P r O H
>
E t O H
>
G L . T h u s , v a l u e s o f
ΔGt0[R+(CH2)nCOOH] are increasingly negative in the order
GL < EtOH<i-PrOH < DMSO. On the other hand, considering
the acidic nature of R+(CH2)nCOOH, its solvation in different
solvents can also be estimated by the Kamlet−Taft β parameter
of the solvents, and the same conclusion can be obtained for
the transfer of R+(CH2)nCOOH from water to aqueous organic
solutions. In this way, the {ΔGt0[R+(CH2)nCOO−] −
ΔGt0[R+(CH2)nCOOH]} values in different aqueous solutions
listed in Table 5 can be qualitatively interpreted.
(4) Werzer, O.; Cranston, E. D.; Warr, G. G.; Atkin, R.; Rutland, M.
W. Ionic liquid nanotribology: Mica−silica interactions in ethyl-
ammonium nitrate. Phys. Chem. Chem. Phys. 2012, 14, 5147−5152.
(5) van Rantwijk, F.; Sheldon, R. A. Biocatalysis in ionic liquids.
Chem. Rev. 2007, 107, 2757−2785.
(6) Singh, R. P.; Verma, R. D.; Meshri, D. T.; Shreeve, J. M. Energetic
Nitrogen-Rich Salts and Ionic Liquids. Angew. Chem., Int. Ed. 2006, 45,
3584−3601.
(7) Vidal, L.; Riekkola, M. L.; Canals, A. Ionic liquid-modified
materials for solid-phase extraction and separation: A review. Anal.
Chim. Acta 2012, 715, 19−41.
At this stage, it seems reasonable to conclude that
{ΔGt0[R+(CH2)nCOO−] − ΔGt0[R+(CH2)nCOOH]} values
not only compensate the ΔGt0(H+) term but also determine
the sign of the total ΔGt0 term. The interaction of both
R+(CH2)nCOOH and its conjuated base with solvent molecules
in the mixed solvents are largely responsible for the observed
changes of pKaT values with the addition of organic solvents in
water.
(8) Tang, B. K.; Bi, W. T.; Tian, M. L.; Row, K. H. Application of
ionic liquid for extraction and separation of bioactive compounds from
plants. J. Chromatogr. B 2012, 904, 1−21.
(9) Nockemann, P.; Thijs, B.; Pittois, S.; Thoen, J.; Glorieux, C.; Van
Hecke, K.; Van Meervelt, L.; Kirchner, B.; Binnemans, K. Task-specific
ionic liquid for solubilizing metal oxides. J. Phys. Chem. B 2006, 110,
20978−20992.
4634
dx.doi.org/10.1021/jp501731j | J. Phys. Chem. B 2014, 118, 4630−4635

The Journal of Physical Chemistry B
Article
(10) Nockemann, P.; Thijs, B.; Parac-Vogt, T. N.; Van Hecke, K.;
Van Meervelt, L.; Tinant, B.; Hartenbach, I.; Schleid, T.; Ngan, V. T.;
Nguyen, M. T.; Binnemans, K. Carboxyl-Functionalized Task-Specific
Ionic Liquids for Solubilizing Metal Oxides. Inorg. Chem. 2008, 47,
9987−9999.
(11) Nockemann, P.; Van Deun, R.; Thijs, B.; Huys, D.; Vanecht, E.;
Van Hecke, K.; Van Meervelt, L.; Binnemans, K. Uranyl Complexes of
Carboxyl-Functionalized Ionic Liquids. Inorg. Chem. 2010, 49, 3351−
3360.
(12) Sun, J.; Han, L. J.; Cheng, W. G.; Wang, J. Q.; Zhang, X. P.;
Zhang, S. J. Efficient Acid-Base Bifunctional Catalysts for the Fixation
of CO2 with Epoxides under Metal- and Solvent-Free Condition.
ChemSusChem. 2011, 4, 502−507.
(13) Xiao, L. F.; Lv, D. W.; Wu, W. Brønsted Acidic Ionic Liquids
Mediated Metallic Salts Catalytic System for the Chemical Fixation of
Carbon Dioxide to Form Cyclic Carbonates. Catal. Lett. 2011, 141,
1838−1844.
water with acetone, isopropanol, glycerol or methanol. J. Chem. Soc.,
Faraday Trans. 1 1974, 70, 694−704.
(31) Wells, C. F. Ionic solvation in water + cosolvent mixtures. Part
9.Free energies of transfer of single ions from water into water +
ethanol mixtures. J. Chem. Soc., Faraday Trans. 1 1984, 80, 2445−2458.
(32) Kamlet, M. J.; Abboud, J. L. M.; Abraham, M. H.; Taft, R. W.
Linear solvation energy relationships. 23. A comprehensive collection
of the solvatochromic parameters, π*, α, and β, and some methods for
simplifying the generalized solvatochromic equation. J. Org. Chem.
1983, 48, 2877−2887.
(33) Hughes, D. L.; Bergan, J. J.; Grabowski, E. J. J. Amino acid
chemistry in dipolar aprotic solvents: Dissociation constants and
ambident reactivity. J. Org. Chem. 1986, 51, 2579−2585.
̆
(34) Dogan, A.; Kılıc,̧ E. Tautomeric and microscopic protonation
equilibria of some α-amino acids. Anal. Biochem. 2007, 365, 7−13.
(35) Bates, R. G.; Covington, A. K.; Jones, P. Hydrogen-Bonded
Solvent Systems; Taylor and Francis: London, 1968.
(36) Popovych, O.; Tomkins, R. P. T. Nonaqueous Solution Chemistry;
John Wiley & Sons, Inc.: Hoboken, NJ, 1981.
(14) Zhang, H.; Cui, H. Synthesis and Characterization of
Functionalized Ionic Liquid-Stabilized Metal (Gold and Platinum)
Nanoparticles and Metal Nanoparticle/Carbon Nanotube Hybrids.
Langumir 2009, 25, 2604−2612.
(15) Fei, Z. F.; Zhao, D. B.; Geldbah, T. J.; Rosario, S.; Dyson, P. J.
Brønsted acidic ionic liquids and their zwitterions: Synthesis,
characterization and pKa determination. Chem.Eur. J. 2004, 10,
4886−4893.
(16) Geng, W. G.; Li, X. H.; Wang, L. F.; Duan, H. L. Acidic
Characterization of Polycarboxyl Imidazolium Ionic Liquids. Acta Phys.
Chim. Sin. 2006, 22, 230−233.
(17) Moity, L.; Durand, M.; Benazzouz, A.; Pierlot, C.; Molinier, V.;
Aubry, J. M. Panorama of sustainable solvents using the COSMO-RS
approach. Green Chem. 2012, 14, 1132−1145.
(18) Gans, P.; O’Sullivan, B. GLEE, a new computer program for
glass electrode calibration. Talanta 2000, 51, 33−37.
(19) Albert, A.; Serjeant, E. P. The Determination of Ionization
Constants; Wiley: New York, 1984.
(20) Brandariz, I.; Castro, P.; Montes, M.; Penedo, F.; Sastre de
Vicente, M. E. Equilibrium constants of triethanolamine in major
seawater salts. Mar. Chem. 2006, 102, 291−299.
(21) Kilpatrick, M.; Eanes, R. D. The Dissociation Constants of Acids
in Salt Solutions. II. Acetic Acid. J. Am. Chem. Soc. 1953, 75, 586−587.
(22) Barriada, J. L.; Brandariz, I.; Sastre de Vicente, M. J. Acid-base
equilibria of monocarboxylic acids in various saline media: Analysis of
data using Pitzer equations. J. Chem. Eng. Data 2000, 45, 1173−1178.
(23) Harned, H. S.; Geary, C. G. The Ionic Activity Coefficient
Product and Ionization of Water in Barium Chloride Solutions from 0
to 50°. J. Am. Chem. Soc. 1937, 59, 2032−2035.
̀ ́
(24) Muinasmaa, U.; Rafols, C.; Bosch, E.; Roses, M. Ionic equilibria
in aqueous organic solvent mixtures the dissociation constants of acids
and salts in tetrahydrofuran/water mixtures. Anal. Chim. Acta 1997,
340, 133−141.
(25) Åkerlof, G. Dielectric constants of some organic solvent−water
̈
mixtures at various temperatures. J. Am. Chem. Soc. 1932, 54, 4125−
4139.
(26) Khoo, K. H. The standard potentials of silver−silver halide
electrodes and ion solvation in dimethyl sulfoxide−water mixtures at
25 °C. J. Chem. Soc. A 1971, 2932−2936.
(27) Everett, D.H.; Landsman, D. A. The Thermodynamics of
Ionization of Some Fatty Acids. Proc. R. Soc. London A 1952, 215,
403−415.
(28) Azab, H. A.; Ahmed, I. T.; Mahmoud, M. R. Potentiometric
determination of the dissociation constants of some monocarboxylic
acids in various hydroorganic media. J. Chem. Eng. Data 1995, 40,
523−525.
(29) Bates, R. G. Solute−solvent interactions and acid−base
dissociation in mixed solvent systems. J. Electroanal. Chem. 1971, 29,
1−19.
(30) Wells, C. F. Ionic solvation in water + Co-solvent mixtures. Part
2.Free energies of transfer of single ions from water into mixtures of
4635
dx.doi.org/10.1021/jp501731j | J. Phys. Chem. B 2014, 118, 4630−4635