Solubilities of triethanolamine hydrochloride ionic liquid in aqueous hydrochloric acid solutions with concentrations from (0 to 8.37) mol·kg-1

Article abstract of DOI:10.1021/je100584w

The (solid + liquid) phase equilibria of triethanolamine hydrochloride ([HTEA]Cl) ionic liquid and aqueous hydrochloric acid solutions with concentrations from (0 to 8.37) mol·kg^-1 were determined by a dynamic method within the temperature range of (285.70 to 344.65) K. Solubilities of [HTEA]Cl were found to increase with increasing temperature for all the investigated solutions. However, solubilities of [HTEA]Cl in aqueous hydrochloric acid decrease with the increment of the concentration of the acid due to the common ion effect. K_SP values of [HTEA]Cl were obtained by use of the solubility data of [HTEA]Cl in water. The modified Apelblat equation was used successfully to correlate experimental data of solubilities in hydrochloric acid. Molar dissolution enthalpy ΔH_Sol of [HTEA]Cl in hydrochloric acid solutions was determined with the newly obtained Apelblat equation parameters.

Full text of DOI:10.1021/je100584w

4434
J. Chem. Eng. Data 2010, 55, 4434–4439
Solubilities of Triethanolamine Hydrochloride Ionic Liquid in Aqueous
†
Hydrochloric Acid Solutions with Concentrations from (0 to 8.37) mol· kg-1
Yan Zhang and Zhibao Li*
Key Laboratory of Green Process and Engineering, Institute of Process Engineering, National Engineering Laboratory for
Hydrometallurgical Cleaner Production Technology, Chinese Academy of Sciences, Beijing 100190, China
The (solid + liquid) phase equilibria of triethanolamine hydrochloride ([HTEA]Cl) ionic liquid and aqueous
hydrochloric acid solutions with concentrations from (0 to 8.37) mol · kg-1 were determined by a dynamic
method within the temperature range of (285.70 to 344.65) K. Solubilities of [HTEA]Cl were found to
increase with increasing temperature for all the investigated solutions. However, solubilities of [HTEA]Cl
in aqueous hydrochloric acid decrease with the increment of the concentration of the acid due to the common
ion effect. KSP values of [HTEA]Cl were obtained by use of the solubility data of [HTEA]Cl in water. The
modified Apelblat equation was used successfully to correlate experimental data of solubilities in hydrochloric
acid. Molar dissolution enthalpy ∆HSol of [HTEA]Cl in hydrochloric acid solutions was determined with
the newly obtained Apelblat equation parameters.
to the best of our knowledge, the solubility data of [HTEA]Cl
have not been measured yet.
Introduction
The chemical and physical properties of ionic liquids have
been vigorously investigated in the past decade. Because ionic
liquids are organic salts, they exhibit many specific character-
istics, including high heat capacities, high densities, extremely
low volatilities, nonflammability, high thermal stability, and a
wide temperature range for the liquid state.1 Hence, ionic liquids
can be applied for catalysis, synthesis, and separation. Ionic
liquids have a large number of possible variations in cation and
anion combinations. With an increasing awareness of how the
structure of an ionic liquid affects its physical properties, the
advantages of ionic liquids as compared to volatile organic
solvents become evident.2 To be able to use ionic liquids
efficiently, it is important to characterize their fundamental
thermodynamic properties and phase behavior in water and other
solvents.3
In this work, triethanolamine hydrochloric with the mass
fraction purity of more than 0.98 was synthesized. Then the
solid-liquid equilibrium (SLE) temperatures of triethanolamine
hydrochloride in water and hydrochloric acid with various
concentrations were determined using a dynamic method. Effects
of temperature and concentration of hydrochloric acid solution
were investigated. The SLE phase diagram was presented. Then
the molar dissolution enthalpy ∆HSol was calculated for these
systems.
Experimental Section
Materials. Analytical reagent (AR) grade triethanolamine was
obtained from Guangdong Xilong Chemical Co., Ltd., China.
Hydrochloric acid with mass fraction of (0.36 to 0.38) and AR
grade ethanol were from Beijing Chemical Works, China. All
chemicals were used as received, and water used was deionized
water.
Solubility is an important property for ionic liquid and has
attracted increased attention. Solubility behavior determines the
method of supersaturation generation as well as the yield of
crystallization and plays an important role in industrial processes
and in theoretical research.4,5 The use of solid + liquid equilibria
(SLE) phase diagrams rather than a trial-and-error approach
would greatly reduce the process development time and cost.6
However, because ionic liquids are relatively new, experimental
measurements of phase behavior are relatively limited. The
deficiency of solubility data could hinder the development of
ionic liquid as solvent and catalyst.
Synthesis of Triethanolamine Hydrochloride. The synthesis
of [HTEA]Cl was performed in a 1 L double-jacket glass reactor,
and the temperature was controlled by a water circulator as
shown in Figure 1.7 A standard volume (250 mL) of trietha-
nolamine of 5 mol ·L-1 located in the reactor (a) was kept at
temperature 293.15 K with circulating water (b). Solution
temperature was monitored with a thermometer (c). Agitation
was provided by an agitator (d) with the stirring speed of 200
rpm. Hydrochloric acid (114.6 mL) was added by titration with
a pump (e) over 60 min. The slurry was filtered, washed with
ethanol, and finally dried in an oven at 323.15 K to remove all
volatile components, water, and excessive acid. The mass
fraction purity of [HTEA]Cl, determined by titration with silver
nitrate solution as the titrant and potassium chromate as the
Triethanolamine hydrochloride ([HTEA]Cl) (C6H15NO3 ·HCl,
CAS No. 637-39-8) is a kind of ammonium ionic liquid and
can be used for detecting elements of tin and antimony or as
the catalyst in the production of resin. This ionic liquid can be
produced by the reaction of triethanolamine and aqueous
hydrochloric acid. Its solubilities in water and hydrochloric acid
solutions are important for the production process. However,
1
indicator, was above 0.98. Its purity was also verified by H
NRM and 13C NRM analyses. The X-ray powder diffraction
(XRD) pattern of the synthesized sample is shown in Figure 2.
Solubility Measurement. Solid solubility was determined
using a dynamic method.8,9 A jacketed glass vessel with a
† Part of the “Sir John S. Rowlinson Festschrift”.
* Corresponding author. E-mail: zhibaoli@home.ipe.ac.cn. Phone: 86-10-
62551557.
10.1021/je100584w  2010 American Chemical Society
Published on Web 08/18/2010

Journal of Chemical & Engineering Data, Vol. 55, No. 10, 2010 4435
Table 1. Experimental Mole Fraction Solubilities of [HTEA]Cl in
Water and KSP Values Calculated by Equation 3 with Solubility
Data
T/K
x1
m1
KSP
285.70
292.47
298.37
303.45
308.65
311.05
316.05
318.35
322.00
326.65
333.05
339.25
344.65
0.0335
0.0399
0.0462
0.0525
0.0587
0.063
0.0699
0.0747
0.0815
0.0884
0.0998
0.1108
0.1215
1.926
2.309
2.691
3.078
3.464
3.735
4.175
4.485
4.930
5.387
6.159
6.923
7.684
3.71
5.33
7.24
9.48
12.00
13.95
17.43
20.12
24.30
29.02
37.93
47.92
59.04
Results and Discussion
Figure 1. Experimental setup used in the synthesis of [HTEA]Cl.
Solubility of [HTEA]Cl in Water. The solubilities of [HTEA]-
Cl in water were measured with the procedure mentioned above
at the temperature range of (285.70 to 344.65) K. The results
are listed in Table 1 and presented in Figure 4. It was found
that the solubility of [HTEA]Cl in water increases sharply with
increasing temperature over the investigated temperatures. As
can be seen in Figure 4, the solubility increases to more than
3.7 times when the temperature increases from (285.7 to 344.65)
K.
Figure 2. XRD pattern of synthesized [HTEA]Cl.
Figure 4. Mole fraction solubility (x1) of [HTEA]Cl in water.
Figure 3. Experimental setup for the solubility measurement of [HTEA]Cl
with a dynamic method.
volume of 250 mL was used in this system as can be seen in
Figure 3. The solvent of known composition was put into the
vessel (a), and the system was maintained at a certain temper-
ature T using a thermostat (b). A known mass of [HTEA]Cl
was added into the solvent. A magnetic stirrer was used to
provide vigorous agitation to the sample (c). Some time later,
more weighted [HTEA]Cl was added if the last trace of salts
was observed to disappear. The mixture of solute and solvent
was heated very slowly (<1 K·h-1 near the equilibrium
temperature) with continuous stirring. A condenser (d) was used
in the system. The temperatures at which crystals disappeared,
detected visually, were measured with a thermometer (e). The
temperature precision was ( 0.1 K. All the chemical regents
were prepared by weighting the pure components with an
uncertainty of ( 0.001 g.
Figure 5. Values of log KSP of [HTEA]Cl in water. 9, log KSP obtained by
experimental solubility data of [HTEA]Cl in pure water with eq 3; solid
line, log KSP values correlated with eq 4.

4436 Journal of Chemical & Engineering Data, Vol. 55, No. 10, 2010
Table 2. Experimental Mole Fraction Solubilities of [HTEA]Cl (1)
in Aqueous Hydrochloric Acid Solutions with Different
Concentrations of HCl (2)
The solubility data of [HTEA]Cl in water were used to
calculate a fundamental thermodynamic parameter, KSP, the
solubility product of the salt. Solubility equilibrium of [HTEA]-
Cl in water can be described by the following dissolution
reaction
x1
T/K
x1
T/K
x2 ) 0.0179 (1.01 mol·kg-1
)
x2 ) 0.0367 (2.12 mol·kg-1
)
0.0375
0.0435
0.0494
0.0552
0.0610
0.0667
0.0723
0.0779
0.0834
0.0888
0.0941
0.0994
296.55
301.30
305.95
310.65
315.00
318.00
321.30
325.95
329.05
332.10
335.35
338.65
0.0305
0.0366
0.0427
0.0486
0.0545
0.0603
0.0661
0.0718
0.0774
0.0840
0.0895
296.35
301.85
306.85
311.85
316.05
319.70
324.55
327.80
331.35
335.15
338.55
[HTEA]Cl T [HTEA]+ + Cl-
(1)
The solubility product of [HTEA]Cl, KSP,[HTEA]Cl, is expressed
as the following equation
KSP,[HTEA]Cl ) a[HTEA]+aCl
(2)
-
+
+
-
where a[HTEA] and aCl are the activity of cation [HTEA] and
anion Cl- in solution. To make it simple, the activities can be
replaced by concentrations of ions as follows
x2 ) 0.0550 (3.23 mol ·kg-1
)
)
)
x2 ) 0.0748 (4.49 mol ·kg-1
)
0.0192
0.0255
0.0317
0.0377
0.0438
0.0497
0.0556
0.0614
0.0671
0.0727
0.0783
289.15
296.55
302.65
308.25
313.15
317.25
321.30
325.95
329.65
333.55
336.65
0.0194
0.0257
0.0320
0.0381
0.0442
0.0502
0.0561
0.0619
0.0677
0.0734
292.85
300.25
306.85
312.75
317.60
321.95
326.35
330.55
334.35
337.75
KSP,[HTEA]Cl ) m[HTEA]+mCl
(3)
-
+
+
-
where m[HTEA] and mCl are the concentration of cation [HTEA]
and anion Cl- in molality. The values of KSP,[HTEA]Cl calculated
with the solubility data in water are also listed in Table 1. These
KSP values were correlated with an empirical equation as shown
below7,10
x2 ) 0.0925 (5.85 mol·kg-1
x2 ) 0.1163 (7.31 mol·kg-1
)
0.0212
0.0281
0.0349
0.0416
0.0482
0.0547
0.0611
0.0675
0.0737
0.0799
295.75
304.85
311.10
316.55
321.95
326.95
331.35
335.45
339.05
342.45
0.0195
0.0258
0.0320
0.0382
0.0443
0.0503
0.0562
0.0620
0.0678
0.0735
0.0792
294.85
303.57
310.95
315.85
320.95
326.25
330.35
334.55
338.55
341.75
345.25
10686.419
log10 KSP,[HTEA]Cl ) -114.625 +
+ 0.391T -
T
0.0004162T2 (4)
where T is the temperature in Kelvin. The results are well fitted
as can be seen in Figure 5.
Solubility of [HTEA]Cl in HCl Solutions. The same proce-
dure was used to measure solubilities of [HTEA]Cl in hydro-
chloric acid solutions with concentrations from (1.01 to 8.37)
mol·kg-1 within the temperature range of (289.15 to 345.25)
K. The results are given in Table 2 and presented in Figures 6a
and 6b. A 3-D diagram of solubilities of [HTEA]Cl in aqueous
hydrochloric acid solutions is shown in Figure 7. It can be seen
that the solubility values are dependent on the system temper-
ature and the concentration of HCl. The solubilities of [HTEA]Cl
increase with increasing temperature in all solvents over the
investigated temperatures. However, the common ion effect
makes it decrease with increasing concentration of aqueous
hydrochloric acid.11,12
x2 ) 0.1309 (8.37 mol·kg-1
0.0196
0.0260
0.0323
0.0385
0.0447
0.0507
0.0567
0.0626
0.0684
0.0742
296.35
304.90
311.55
317.35
322.90
328.15
332.55
336.65
340.45
344.30
cal
where n is the number of experimental points; x is the
solubility calculated by eq 5; and x is the experimental value
1,i
Figures 8a and 8b show that the logarithm of the mole fraction
solubility x1 determined in this work plotted against the inverse
temperature shows good linearity. Therefore, we decided to
correlate the solubility data by the modified Apelblat equation13,14
exp
1,i
of mole fraction solubility. The values of a, b, and rmsd for the
solubilities of triethanolamine hydrochloride in hydrochloric acid
solutions are given in Table 3. It can be observed that the
correlated solubilities agree well with the experimental values.
This indicates that the modified Apelblat equation is a successful
description of the solubilities of [HTEA]Cl in hydrochloric acid
solutions.
b
T/K
ln x1 ) a +
(5)
where x1 and T are mole fraction solubility of solute and absolute
temperature, respectively, and a and b are empirical parameters.
The values of the two parameters were evaluated by nonlinear
least-squares method. The root-mean-square deviation (rmsd)
is defined as
The solubilities of [HTEA]Cl in hydrochloric acid solutions
were predicted at T ) (300, 310, 320, 330, 340, and 350) K
using the modified Apleblat equation and parameters from Table
3. Hence, the SLE phase diagram of [HTEA]Cl in aqueous
hydrochloric acid solutions was constructed by the calculated
solubility data as shown in Figure 9. It can be seen in Figure 9
that each curve indicates one boundary of the crystallization
compartments, as an important feature of the phase diagram.6
The crystallization region can be found by the curves in the
n
1/2
1
rmsd )
(xcal - xexp 2
)
1,i
(6)
∑
1,i
[
]
n i)1

Journal of Chemical & Engineering Data, Vol. 55, No. 10, 2010 4437
Figure 6. Mole fraction solubility of [HTEA]Cl (x1) in aqueous hydrochloric
acid solutions with various concentration (x2). (a) 9, x2 ) 0.0179; b, x2 )
0.0367; 2, x2 ) 0.0550; 1, x2 ) 0.0748. (b) Solid triangle pointing right,
x2 ) 0.0925; solid triangle pointing left, x2 ) 0.1163; [, x2 ) 0.1309.
Figure 8. Plots of ln x1 versus 1/T for mole fraction solubility of [HTEA]Cl
(x1) in aqueous hydrochloric acid solutions with various concentrations (x2):
(a) 9, x2 ) 0; b, x2 ) 0.0179; 2, x2 ) 0.0367; 1, x2 ) 0.0550. (b) Solid
triangle pointing right, x2 ) 0.0748; solid triangle pointing left, x2 ) 0.0925;
O, x2 ) 0.1163; 0, x2 ) 0.1309. The line is the values calculated with the
modified Apelblat equation.
Table 3. Parameters of the Modified Apelblat Equation (Equation
5) for [HTEA]Cl (1) in Aqueous Hydrochloric Acid Solutions with
Different Concentrations of HCl (2)
x2
a
b
rmsd
0.000
4.236
4.547
5.106
5.975
6.026
5.906
5.715
5.627
-2179.482
-2313.932
-2538.412
-2857.758
-2909.892
-2883.372
-2842.659
-2826.393
0.001071
0.001055
0.000918
0.001302
0.000805
0.004044
0.003508
0.000638
0.0179
0.0367
0.0550
0.0748
0.0952
0.1163
0.1309
of dissolution enthalpy can indicate solid-liquid interfacial
energy, which is a significant physical property influencing
nucleation and growth of crystals.15,16
Figure 7. 3-D diagram of solubility of [HTEA]Cl in aqueous hydrochloric
acid solutions.
On the basis of the principle of solid/liquid equilibrium, when
a solid is dissolved in a solvent and forms an ideal solution,
the chemical potential of the solute i can be expressed as17
SLE phase diagram, which is of utility for the synthesis process
of [HTEA]Cl.
Calculation of Dissolution Enthalpy. Dissolution enthalpy
∆HSol is a relevant thermodynamic measure for the degree of
interaction of the solvent and the solute molecules. The value
*
µi ) µi + RT ln xi
(7)

4438 Journal of Chemical & Engineering Data, Vol. 55, No. 10, 2010
where µi, µ*, xi, R, and T stand for the chemical potential of
i
the pure solid, chemical potential of the solute in ideal solution,
the mole fraction in the solution, the gas constant, and absolute
temperature, respectively. However, for nonideal solutions, the
equilibrium relation can be rearranging as
*
µi ) µi + RT ln(γixi)
(8)
where γi is the activity coefficient of solute in solution.
Rearranging, we obtain
*
µi
µi
ln(γixi) ) -
+
(9)
RT
RT
Figure 10. Dissolution enthalpy ∆HSol versus the concentration of
hydrochloric acid x2.
In accordance with the phase rule, the temperature and pressure
can be varied independently. When the pressure is held constant,
the partial derivative of activity to T can be obtained
b
T
∂(a +
∂T
)
∂ ln xi
∂T
b
T2
)
(
)
) -
(13)
(
)
p
P
∂ln(γixi)
∂T
Hi - hi
RT2
)
(10)
(
)
P
By combining eq 12 and eq 13, we obtain
Hi is the partial molar enthalpy of the component in the ideal
solution, and hi is its enthalpy per mole as the pure solid, both
referring to the temperature T. the equation may therefore be
rewritten as
∆HSol
RT2
b
T2
-
)
(14)
(15)
∂ ln(γixi)
∂T
∂ ln γi
∂T
∂ ln xi
∂T
∆HSol
RT2
Finally, the dissolution enthalpy can be expressed as
)
+
)
p
(
)
(
)
(
)
p
p
(11)
∆HSol ) -bR
where ∆HSol is the dissolution enthalpy of the component i. If
γi has low temperature dependence and can be neglected, eq
11 can be simplified to
The values of ∆HSol determined by eq 15 are presented in
Figure 10. It can be seen that ∆HSol increases with increasing
concentration of aqueous hydrochloric acid when x2 is lower
than 0.075. However, it remains constant with the value of about
23.3 kJ ·mol-1 when x2 is more than 0.075.
∂ ln xi
∂T
∆HSol
RT2
)
(12)
(
)
p
Conclusions
The solubilities of [HTEA]Cl in water and hydrochloric acid
with concentrations from (1.01 to 8.37) mol ·kg-1 have been
successfully determined by a dynamic method. Solubility of
[HTEA]Cl in hydrochloric acid increases with increasing
temperature. However, the common ion effect makes the
solubilities of [HTEA]Cl in hydrochloric acid decrease with
increasing concentration of hydrochloric acid solutions. The
molality-based solubility product (assuming activity coefficients
of unity) was obtained by correlating the solubility data of
[HTEA]Cl in deionized water. Experimental data of solubilities
in aqueous hydrochloric acid were correlated by the modified
Apelblat equation, and the calculated solubilities showed good
agreement with the experimental data. On the basis of the two-
parameter Apelblat equation, molar dissolution enthalpy ∆HSol
of [HTEA]Cl in aqueous hydrochloric acid was calculated.
Substituting eq 5 for ln xi, the following equation can be obtained
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JE100584W