Studies on thermochemical properties of ionic liquids based on transition metal

Article abstract of DOI:10.1007/s10973-008-7743-x

A brown and transparent ionic liquid (IL), [C₄mim][FeCl ₄], was prepared by mixing anhydrous FeCl₃ with 1-butyl-3-methylimidazolium chloride ([C₄mim][Cl]), with molar ratio 1/1 under stirring in a glove box filled with dry argon. The molar enthalpies of solution, Δ_s H_m, of [C₄mim][FeCl ₄], in water with various molalities were determined by a solution-reaction isoperibol calorimeter at 298.15 K. Considering the hydrolyzation of anion [FeCl₄]^- in dissolution process of the IL, a new method of determining the standard molar enthalpy of solution, Δ_s H_m ⁰ , was put forward on the bases of Pitzer solution theory of mixed electrolytes. The values of Δ_s H_m ⁰ and the sum of Pitzer parameters: (4β_Fe,Cl^(0)L+4β_{C4 mim,CL} ^(0)L+Φ_{Fe,C4 mim}^L) and (β _Fe,Cl^(1)L+4β_{C4 mim,CL}^(1)L were obtained, respectively. In terms of thermodynamic cycle and the lattice energy of IL calculated by Glasser's lattice energy theory of ILs, the dissociation enthalpy of anion [FeCl₄]^-, ΔH_dis≈5650 kJ mol^-1, for the reaction: [FeCl₄]^-(g) →Fe³⁺(g)+4Cl^-(g), was estimated. It is shown that large hydration enthalpies of ions have been compensated by large the dissociation enthalpy of [FeCl₄]^- anion, Δ_d H_m, in dissolution process of the IL.

Full text of DOI:10.1007/s10973-008-7743-x

Journal of Thermal Analysis and Calorimetry, Vol. 94 (2008) 2, 507–510
STUDIES ON THERMOCHEMICAL PROPERTIES OF IONIC LIQUIDS
BASED ON TRANSITION METAL
*
W. Guan, L. Li, H. Wang, J. Tong and J.-Z. Yang
Department of Chemistry, Laboratory of Green Chemistry, Liaoning University, Shenyang, 110036, P.R. China
A brown and transparent ionic liquid (IL), [C4mim][FeCl4], was prepared by mixing anhydrous FeCl3 with 1-butyl-3-methylimid-
azolium chloride ([C4mim][Cl]), with molar ratio 1/1 under stirring in a glove box filled with dry argon. The molar enthalpies of
solution, DsHm, of [C4mim][FeCl4], in water with various molalities were determined by a solution-reaction isoperibol calorimeter at
298.15 K. Considering the hydrolyzation of anion [FeCl4]– in dissolution process of the IL, a new method of determining the standard
0
0
molar enthalpy of solution, DsHm, was put forward on the bases of Pitzer solution theory of mixed electrolytes. The values of DsHm and
the sum of Pitzer parameters: (4b(F0e),LCl +4bC(0)mLim,Cl +F LFe,C mim) and (bF(1e),LCl +b(C1)mLim,Cl) were obtained, respectively. In terms of
4
4
4
thermodynamic cycle and the lattice energy of IL calculated by Glasser’s lattice energy theory of ILs, the dissociation enthalpy of
anion [FeCl4]–, DHdis»5650 kJ mol–1, for the reaction: [FeCl4]–(g)®Fe3+(g)+4Cl–(g), was estimated. It is shown that large hydration
enthalpies of ions have been compensated by large the dissociation enthalpy of [FeCl4]– anion, DdHm, in dissolution process of the IL.
Keywords: enthalpy of solution, ionic liquid, isoperibol calorimeter, Pitzer’s theory
Introduction
and transparent IL, [C4mim][FeCl4], was prepared by
mixing FeCl3 with 1-methyl-3-butylimidazolium chlo-
ride ([C4mim][Cl]), and the molar solution enthalpies of
the ILs, DsHm, in water with various molalities were de-
termined using the on-line solution-reaction isoperibol
calorimeter at T=298.15 K; (2) Since the hydrolyzation
of anion [FeCl4]– in dissolution process of the IL, a new
method of determining the standard molar enthalpy of
Since Wilkes et al. [1] synthesized the air- and wa-
ter-stable ionic liquid (IL) [Emim][BF4] (1-methyl-3-
ethylimidazolium tetrafluoroborate), ILs are receiving
increased attention from industrial and academic com-
munities as a novel class of green benign solvents,
which promise possibly replacing currently used or-
ganic solvents industrially [2–6]. However, air- and wa-
ter-sensitive ILs based on AlCl3 and FeCl3 are still very
attractive for basic electrochemical research, and for es-
pecial catalytic reaction and for nanotechnology be-
cause their acidity can be varied over wide ranges [7].
Recently, Sitze et al. [8] and Zhang et al. [9] also pre-
pared ILs based on FeCl3. They pointed out that [FeCl4]–
was main anion in the ILs.
0
solution, D sHm , was put forward on the basis of Pitzer
0
electrolyte solution theory [13]. The values of D sHm
and
the
sum
of
Pitzer
parameters:
(0)L
(4b
+4b(0)L +FLFe,C mim ) and (bF(1e),LCl +b(1)L ) were
Fe,Cl
C4mim,Cl
C4mim,Cl
4
obtained, respectively. (3) According to thermodynamic
cycle and the lattice energy of IL calculated by Glasser’s
theory [14], the dissociation enthalpy of [FeCl4]– anion,
that is heat for the reaction [FeCl4]–(g)®Fe3+(g)+
4Cl–(g), was estimated.
Most study on ILs has been investigated in elabo-
ration of the synthetic methods, application in catalytic
processes, electrochemistry and other more specialized
fields. However, the physico-chemical properties of ILs Experimental
have not been studied systematically, this holds in par-
Materials
ticular for thermochemical properties. The solution
enthalpy of hydrophilic ILs in water is employed in
Chemicals
many thermodynamic calculations, from thermo-
Anhydrous FeCl3 was purchased from Aldrich, opened
chemistry to solution chemistry and is important for de-
signing chemical engineering processes. However, only
a very restricted number of research on solution
enthalpy of hydrophilic ILs in water is available in the
literature [10–12]. As a continuation of our previous in-
vestigations on ILs, this paper reports that (1) A brown
in the glove box filled with dry argon, and used without
further purification. 1-Methylimidazole (AR reagents)
was distilled under reduced pressure. Acetonitrile,
chlorobutane and ethyl acetate were AR reagents and
distilled. KCl, with purity more than 99.99%, was dried
*
Author for correspondence: jzyanglnu@yahoo.com.cn
1388–6150/$20.00
© 2008 Akadémiai Kiadó, Budapest
Akadémiai Kiadó, Budapest, Hungary
Springer, Dordrecht, The Netherlands

GUAN et al.
in a vacuum oven at 130°C for 6 h prior to use. THAM
the tested isoperibol calorimeter the molar enthalpies
of solution of [C4mim][FeCl4] with various molalities
in water were measured.
(tris-(hydroxymethy)aminomethane), GR grade reagent,
was dried in a vacuum oven before use. Double-distilled
water was used.
Results and discussion
Preparation of 1-methyl-3-butylimidazolium chloride
[15]
The molar enthalpy of solution
1-methyl-3-butylimidazolium chloride ([C4mim][Cl])
was synthesized by refluxing the 1-methylimidazole
with a large excess of the chlorobutane for 24 h. The ex-
cess chlorobutane was removed by evaporation and
crude product was recrystallized from the mixed solvent
of (acetonitrile+ethyl acetate). The resulting white pre-
cipitate was isolated by filtration and then dried in vacuo
for 24 h. The melting point is 66–68°C. Analysis of
The values of molar solution enthalpy of
[C4mim][FeCl4] with various molalities in water at
298.15 K are listed in Table 1. From Table 1, the dis-
solution process of ionic liquid [C4mim][FeCl4] is
exothermic. The molar enthalpies of solution of the IL
increase with the decrease of molalities.
0
The standard molar solution enthalpy, D sHm , of the IL
1
[C4mim][Cl] by H-NMR resulting in a spectrum is in
good agreement with the literature [15].
Preparation of IL [C4mim][FeCl4]
The solution reaction of IL, [C4mim][FeCl4], is
[C4mim][FeCl4](l)+WH2O(l)®
[C4mim]+(aq)+Fe3+(aq)+4Cl–(aq)
All glasswares that contacted the IL were cleaned in
hot dilution nitric acid and rinsed repeatedly in dou-
bly deionized water, then were baked dry in 393 K
oven and stored in desiccators before use.
where W is number of kilogram of water. The solution
enthalpies of IL [C4mim][FeCl4], DsH, may be ex-
pressed:
FeCl3 was added slowly with stirring to a small
glass vial containing the equal molar [C4mim][Cl] in a
glove-box filled with dry argon, then the brown and
transparent ionic liquid compounds, [C4mim][FeCl4],
was obtained [8]. Analysis of the product by Raman
scattering gave a spectrum identical to that for
[C4mim][FeCl4] in literature [8].
D sH =n1H1 +n2 (H +H
Fe+
+4H )-
Cl –
[C4mim]+
0
0
n1H1 -n2H[C mim][FeCl
(1)
4 ](l)
4
where Hi and ni are partial molar enthalpy and mole
number of component i, subscripts 1 and 2 mean sol-
0
vent and solute, respectively, H1 and HBMIFeCl mean
0
4 (l)
enthalpy of pure solvent and pure IL [C4mim][FeCl4],
respectively. Equation (1) may be rewritten as
0
D sH =L+n2 (H +H
0
0
+4H )-
Determination of the molar enthalpies of solution of ILs
Fe+
[C4mim]+
Cl –
0
n2H[C mim][FeCl
On the basis of other calorimetric apparatus [16, 17]
and our previous work [18], an isoperibol calorimeter
was constructed. The construction of the calorimeter
and the detailed experimental procedure have been
described elsewhere [12, 18].
(2)
4 ](l)
4
0
where L=H–H is relative thermodynamic function of
solution. As the concentration m approaches zero, we
have
The performance and accuracy of the calorimet-
ric system was tested by measuring the molar
enthalpy of solution of KCl in water and THAM in
0.1 mol dm–3 HCl(aq) at 298.15 K. The mean molar
enthalpies solution are 17542±31 J mol–1 for KCl and
–29794±21 J mol–1 for THAM which are good in con-
formity with the corresponding published data:
0
lim m ®0 (D sH/n2 )=(H +H
0
0
+4H )-
Fe+
[C4mim]+
Cl –
0
H[C mim][FeCl =D sHm
0
(3)
4 ](l)
4
0
where D sHm is the molar enthalpy of solution of IL at
infinite dilution, that is, standard molar enthalpy of
solution. At finite concentrations we therefore have
0
17536±9
J
mol–1 for KCl [19, 20] and
D sHm =D sH/n2 =D sHm +L/n2
(4)
–29765±10 J mol–1 for THAM [20, 21]. Then, using
Table 1 Values of molar solution enthalpy of [C4mim][FeCl4] with various molalities at 298.15 K
102m/mol kg–1
0.2784
0.2995
0. 3393
0. 4001
0. 4017
0. 5010
DsHm/J mol–1
102m/mol kg–1
DsHm/J mol–1
–62.06
0. 5970
–59.07
–61.01
0. 6044
–58.42
–60.47
0.8042
–58.50
–59.93
0. 9999
–58.00
–59.79
1.500
–59.41
2.998
–57.51
–57.30
508
J. Therm. Anal. Cal., 94, 2008

IONIC LIQUIDS BASED ON TRANSITION METAL
2 1/2
Y=DsHm/2RT –(1/2m)(4IAH/1.2)ln(1+1.2I )=
(0)
According to thermodynamics, the relationship
a0 -m(4b
+4b(0)
+FFe,[C mim] )-
between the relative enthalpy, L, and excess Gibbs
ex
free energy, G , is
Fe,Cl
[C4mim],Cl
4
4m(b(F1e),LCl +b(1)L
)g(x)
(17)
[C4mim],Cl
2 ex
L=–T [¶(G /T)/¶T]p,m
(5)
where Y is the extrapolate function, which can be ob-
tained by experimental data. a0 in the Eq. (17) is de-
fined by following equations, respectively.
where the subscript m and p indicate constancy of
composition and pressure for the partial derivatives,
respectively. It is useful to recall that the excess Gibbs
free energy for a mixed electrolytes by Pitzer equa-
tions [13]. We shall write out for the mixture system
of {BMI+(aq)+Fe3+(aq)+4Cl–(aq)} without the terms
of triple-ionic interaction:
0
a0=D sHm /2RT
2
(18)
The experimental data of molar solution
enthalpies of [C4mim][FeCl4] was fitted to Eq. (17),
and the values of a0=–0.047042,
(4b(F0e),Cl +4b(0)
+F )=0.063383
ex
G /WRT=f(I)+2ScSamcmaBca+2ScSc’mcmc’Fcc’ (6)
and
+b(1)L )=0.096257 and standard deviation of
C4mim,Cl
Fe,C4mim
(1)L
(b
Fe,Cl
C4mim,Cl
where subscripts, c and a, mean cation and anion, re-
the fit, s=2.8·10–4, and correlation coefficient, r=0.97,
for [C4mim][FeCl4] are obtained, respectively. The
value of standard molar solution enthalpy,
spectively, other symbols are defined by following Eqs.
1/2
f(I)=–(4IAF/1.2)ln(1+1.2I )
(7)
(8)
D sHm =–69.5 kJ mol–1 was calculated using Eq. (18).
0
Bca=b(c0a) +bc(1a) g(x)
In comparison with D sHm =–76.6 kJ mol–1 of
0
2
[C2mim][FeCl4] in [22], we can draw a conclusion
that each mole of methylene group contributes to the
standard molar enthalpy of solution of [Cxmim]
[FeCl4] (x=2, 4) is about 3.55 kJ mol–1 since the
hydrophobicity of hydrocarbonic chain.
g(x)=2[1–(1+x)exp(–x)]/x
(9)
1/2
x=2I
(10)
For a mixture of {[C4mim]+(aq)+Fe3+(aq)+
4Cl–(aq)}, Eq. (6) may be simplified. Because
m
Fe3+
=m
=m, m =4m, Eq. (6) changed to
[C4mim]+ Cl –
Eq. (11)
G /WRT =-(4IAF /1.2)ln(1+1.2I 1/ 2 )+
ex
(0)
2m (4bFe,Cl +4b(0)
+FFe,[C mim] )+
2
[C4mim],Cl
4
(1)
8m (bFe,Cl +b(1)
)g(x)
(11)
2
[C4mim],Cl
One may now proceed to obtain L by differentia-
ex
Fig. 1 The schematic diagram of thermodynamic cycle for
tion of (G /WRT)
hydration process of [C4mim][FeCl4]
L=-(WRT )[-(4IAH /1.2)ln(1+1.2I 1/ 2 )+
2
–
Dissociation enthalpy of [FeCl4] (g), DdHm
2m (4bFe,Cl +4b(0) L
+FLFe,[C mim] )+
(0) L
2
[C4mim],Cl
4
Figure 1 is the thermodynamic cycle for the reaction
(1) L
8m (bFe,Cl +b(1)
)g(x)
(12)
2
[C4mim],Cl
of solution of [C4mim][FeCl4](l) in water at infinite
0
dilution. In Fig. 1, D sHm is standard molar solution
where m is molality, I means ionic strength, R is gas
constant, other symbols are defined by following
equations:
enthalpy
–4435 kJ mol–1 [23] is hydration enthalpy of Fe3+,
DH
=–201 kJ mol–1 [11] is hydration enthalpy of
of
[C4mim][FeCl4](l),
DH
=
Fe3+
[C4mim]+
(0)L
MX
(0)
=(¶b /¶T )P
[C4mim]+ and DH =–376 kJ mol–1 [23] is hydration
b
b
(13)
(14)
(15)
MX
Cl –
enthalpy of Cl–, UPOT is crystal energy, 420 kJ mol–1,
(1)L
MX
(1)
=(¶b /¶T )P
MX
which was calculated by Glasser’s theory [14]:
FLFe,[C mim] =(¶FFe,[C mim] /¶T )P
UPOT=1981.2(r/M)1/3+103.8 kJ mol–1
(19)
4
4
where M and r are molar mass and density of ionic
liquid [C4mim][FeCl4], respectively, DdHm is the dis-
sociation enthalpy of anion [FeCl4]–(g) and may be
calculated in terms of the thermodynamic cycle in
Fig. 1:
AH=(¶AF/¶T)P
(16)
where AH is Debye-Hückel parameter of enthalpy and
its value was taken from literature [13]. Substituting
Eq. (12) into (4) and rearranging, the working equa-
tion to determine Pitzer’s parameters was obtained:
0
D d Hm =D sHm -U POT -DH
-DH -4DH
Fe3+ Cl –
»
[C4mim]+
5650 kJ mol–1
J. Therm. Anal. Cal., 94, 2008
509

GUAN et al.
11 W. Guan, H. Wang, L. Li and J.-Z. Yang, Thermochim.
The result shows the cause of low solution
enthalpy of the IL, that is, the large hydration enthal-
pies of ions are made compensation [24, 25] for the
large dissociation enthalpy of anion [FeCl4]–(g). In
comparison of DdHm (5650 kJ mol–1) in this work with
5659 kJ mol–1 obtained by Zhang et al. [22], they are
in good agreement within experimental error.
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Acknowledgements
This project was supported by NSFC (20773056).
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Received: June 6, 2006
Accepted: January 4, 2008
OnlineFirst: September 20, 2008
DOI: 10.1007/s10973-008-7743-x
510
J. Therm. Anal. Cal., 94, 2008