Enthalpies of reaction of calcium chloride and sodium oxalate in an aqueous NaCl solution

Article abstract of DOI:10.1134/S0036023609120328

The heats of mixing of dilute aqueous solutions of calcium chloride and sodium oxalate with additions of 1-5 wt % NaCl at 298.15 K and the heats of dilution of calcium chloride solutions were measured. Increasing the sodium chloride content in a solution

Full text of DOI:10.1134/S0036023609120328

ISSN 0036-0236, Russian Journal of Inorganic Chemistry, 2009, Vol. 54, No. 12, pp. 2027–2030. © Pleiades Publishing, Inc., 2009.
Original Russian Text © A.V. Kustov, M.B. Berezin, N.L. Smirnova, A.F. Syshchenko, B.D. Berezin, V.N. Trostin, 2009, published in Zhurnal Neorganicheskoi Khimii, 2009,
Vol. 54, No. 12, pp. 2109–2112.
PHYSICAL CHEMISTRY
OF SOLUTIONS
Enthalpies of Reaction of Calcium Chloride and Sodium
Oxalate in an Aqueous NaCl Solution
A. V. Kustova, M. B. Berezina, N. L. Smirnovaa, A. F. Syshchenkob,
B. D. Berezina, and V. N. Trostina
a Institute of Solution Chemistry, Russian Academy of Sciences, Akademischeskaya 1, Ivanovo, 153045 Russia
b Belarusian State University, Minsk, Belarus
E-mail: kustov@isuct.ru
Received November 27, 2008
Abstract—The heats of mixing of dilute aqueous solutions of calcium chloride and sodium oxalate with addi-
tions of 1–5 wt % NaCl at 298.15 K and the heats of dilution of calcium chloride solutions were measured.
Increasing the sodium chloride content in a solution noticeably increases the time of precipitation of calcium
oxalate. A fine precipitate of CaC2O4 formed in solutions containing 3 and 5 wt % NaCl is difficult to remove
from the parts of a calorimeter cell. The enthalpies of precipitation of CaC2O4 depend slightly on the content
of the “background electrolyte,” whereas the enthalpies of dilution, owing to ion association, significantly
decrease in magnitude and become positive in a 5% NaCl solution. The “standard” enthalpy of precipitation in
water, determined by extrapolation of the experimental values to the zero concentration of the background elec-
trolyte, differs noticeably from the enthalpy of precipitation in water.
DOI: 10.1134/S0036023609120328
Development of new methods of conservative treat- tive conclusion that the increased salt level of urine
ment of urolithiasis calls for comprehensive study of would be favorable for the formation of harder concre-
the formation of the major components of uroliths, tions [7, 8].
namely, calcium salts insoluble in water and multicom-
In the present work, we studied specific features of
ponent aqueous systems modeling, to an extent, the
the formation of CaC2O4 solid in sodium chloride solu-
conditions of in vivo formation of mineral composites.
tions and determined the enthalpy of precipitation as a
These processes have still been little studied and,
function of the salt level.
although some data on the influence of various factors
on the formation of insoluble calcium salts are cur-
rently available [1–8], the mechanism of formation of
uroliths is still unclear.
EXPERIMENTAL
The heats of reaction of calcium chloride and
sodium oxalate in an aqueous NaCl solution were mea-
sured on an isothermal-shell variable-temperature calo-
rimeter with automatic data collection and processing
[9]. A calorimeter vessel (~70 cm3) was equipped with
a mercury seal four-bladed stirrer, a heater, and a tem-
perature gage (a thermistor with a resistance of ~7.4 kΩ
at 25°C). The stirrer also served as the ampoule holder.
The reaction was initiated by brief pressing of the stir-
rer shaft so that the ampoule was broken against the
bottom of the calorimeter vessel. Measurements were
taken by a comparative method, the electric calibration
time and temperature rise being selected to be as close
as possible to those observed in the experiment. The
thermistor resistance was measured by a standard tem-
perature measuring instrument. These data were
received and processed by a computer. The temperature
in the range 10–65°C in a thermostat in the region of
location of the calorimeter cell was maintained with an
Previously [1, 2], we studied the reaction of calcium
chloride and sodium oxalate in electrolyte and nonelec-
trolyte systems partially modeling the biological condi-
tions of formation of uroliths and found that the nature
of the additive noticeably influences the precipitation
process. In particular, in water and 3% solutions of urea
and tetrabutyl- and cetyltrimethylammonium bro-
mides, precipitation is rapid and the nascent flaky pre-
cipitate is presumably a crystal hydrate mixture [4, 5].
In 3 wt % Trilon B solutions, there is no precipitation at
all because of the endothermic complexation of the cal-
cium cation with the ethylenediaminetetraacetate ion;
therefore, Trilon B is used for both inhibiting precipita-
tion and dissolving the already existing concretions [1].
In a sodium chloride solution, the formation of
CaC2O4 solid was noticeably hindered, which could be
interpreted as partial inhibition of precipitation of cal-
cium oxalate-based concretions. However, the CaC2O4
precipitate formed under these conditions is fine and accuracy of no worse than 0.001°ë. The sensitivity of
very compact, which made it possible to draw a tenta- the measuring circuit was 0.00001°ë (0.003 J). The cal-
2027

2028
KUSTOV et al.
Table 1. Enthalpies of mixing of Na2C2O4 and CaCl2 solutions (ΔmixH), enthalpies of dilution of CaCl2 solutions (ΔdilH), and
enthalpies of precipitation of CaC2O4 (ΔprH) in aqueous NaCl solutions at 25°C
n(Na2C2O4)a × 104,
n(NaCl) × 104,
m(CaCl2)fin,
mol/kg
–ΔmixH,
–ΔmixH,
–ΔdilH,
mol
mol
kJ/mol
kJ/mol
kJ/mol
1 wt % NaCl
3.398
3.460
3.381
1.484
1.220
1.456
0.00211
0.00171
0.00206
23.82
23.73
23.35
1.42
1.46
1.43
22.40
22.27
21.92
ΔprH (av.) = –22.0 0.28b
2 wt % NaCl
3.469
3.903
3.962
1.292
1.338
1.357
0.00182
0.00191
0.00190
23.06
23.30
22.96
1.07
1.07
1.07
21.99
22.23
21.89
ΔprH (av.) = –22.04 0.21
3 wt % NaClc
0.00142
0.00189
0.00198
0.00225
3.393
3.470
3.547
3.848
1.004
1.329
1.419
1.589
22.61
22.46
22.15
22.47
0.71
0.68
0.67
0.65
21.90
21.78
21.48
21.82
ΔprH (av.) = –21.75 0.18
5 wt % NaCl
3.612
3.624
3.955
3.927
1.177
0.7296
1.256
1.148
0.00166
0.00102
0.00176
0.00161
20.57
20.00
20.22
20.45
–0.14
20.71
20.08
20.37
20.58
–0.08
–0.15
–0.13
ΔprH (av.) = –20.44 0.27
a
Notes: The first and second columns, respectively, list the number of sodium oxalate and calcium chloride moles in the solution before
b
c
mixing. Errors are expressed as the doubled standard mean deviation. Taken from [1]. Note. The enthalpies of dilution were
found by linear extrapolation of the data in Table 2. The enthalpy of precipitation in water is –20.37 0.11 and –20.56
0.08 kJ/mol according to {1, 2], respectively.
orimeter performance was checked by measuring the main and final periods of the process were recorded.
enthalpies of solution of potassium chloride and 1-pro- The time of the mixing reaction was 3–15 min, depend-
panol in water at 25°C. The results demonstrated com- ing on the content of the background electrolyte; there-
plete agreement between our data and the most reliable fore, to determine the change in temperature corrected
literature data [9].
for heat exchange, the modified Dickinson method was
used, in which the linear segments of the temperature
versus time curve are extrapolated to the mean temper-
ature of the main period so that the areas cut off by the
vertical line under and over the experimental curve are
equal.
Calcium chloride (pharmacopeial grade, from Rea-
khim) was recrystallized from absolute methanol and
vacuum dried at 250°C to constant weight. Sodium
oxalate and chloride (chemically pure, Reakhim) were
not additionally purified and only vacuum dried at 90°C
to constant weight. Solutions were prepared by intro-
ducing exact weights (the error of weighing was
1 × 10–4 g) of the dried compounds into a specified
water volume.
RESULTS AND DISCUSSION
The enthalpies of mixing of solutions of sodium
oxalate and calcium chloride are presented in Table 1.
As in [1, 2], all measurements were taken from dilute
solutions in order the systems under consideration were
studied under conditions similar to the in vivo ones, on
the one hand, and the solubility product of calcium
oxalate was achieved and the enthalpy of precipitation
could be determined experimentally, on the other hand.
The experiments were carried out as follows. A por-
tion of ~0.1–0.3 mL of an aqueous CaCl2 solution (m =
0.51967 mol/kg) was poured into a thin-walled glass
ampoule. The ampoule was placed into the calorimeter
cell filled with a solution containing 1–5% sodium
chloride and Na2C2O4 (m ≈ 0.005 mol/kg solvent).
Then, the calorimeter was assembled and placed into a
thermostat where it was kept for 30–50 min at the
It follow from Table 1 that the enthalpies of mixing
ΔmixH are negative and decrease in magnitude with
experiment temperature. The specific heat of the calo- increasing the content of sodium chlorine in a solution.
rimeter was determined by means of electric calibration This decrease in the enthalpy constitutes ~20% in going
immediately before the experiment was started. Then, from 1% to 5% NaCl solution. Changes in the enthalpy
the initial period was recorded, the ampoule was broken of dilution of calcium chloride are more significant
against the bottom of the calorimeter vessel, and the (here, an NaCl solution is treated as an individual sol-
RUSSIAN JOURNAL OF INORGANIC CHEMISTRY Vol. 54 No. 12 2009

ENTHALPIES OF REACTION
2029
Table 2. Enthalpies of dilution of CaCl2 solutions from 0.51697 mol/kg to m(CaCl2)fin in NaCl solutions of various concen-
tration at 25°C
m(CaCl2)fin,
–ΔdilH,
m(CaCl2)fin,
mol/kg
–ΔdilH,
m(CaCl2)fin,
–ΔdilH,
mol/kg
kJ/mol
kJ/mol
mol/kg
kJ/mol
1 wt % NaCl
0.001905
0.004156
2 wt % NaCl
0.00144
0.00357
5 wt % NaCl
1.44
1.24
1.10
0.93
0.00101
0.00297
0.0047
–0.07
–0.27
–0.47
vent). Their values (Table 2) not only decrease in mag- activity coefficients, as well as with the appearance of a
nitude but also change their sign. Dilution of a calcium
chloride solution with water is accompanied by heat
release (~–3 kJ/mol [1]), whereas dilution with a 5%
NaCl solution leads to heat absorption (0.1–
0.2 kJ/mol). It is evidently due to the appearance in the
considerable amount of background electrolyte ions so
that, in the reaction zone, the Ca2+ and C2O42– ions turn
out to be separated by sodium and, especially, chloride
ions. The fine precipitate that slowly forms in these sys-
aqueous solution of an electrolyte with common anion, tems is tightly adsorbed on the cell surface and is very
namely, chloride ion. As the concentration of the latter
increases, the competition between the ions for the sol-
vent and the tendency for association of Ca2+ and Cl–
ions become stronger. These processes are presumably
responsible for the change of the sign at ΔdilH.
difficult to remove mechanically from the latter. It is
likely that, in solutions of sodium chloride with a con-
centration of 3 wt % or higher, the CaC2O4 precipitate
is not a crystal hydrate. Inasmuch as Na+ and Cl– ions
constitute a major part of the mineral content of urine
To calculate the change in enthalpy in the course of
precipitation of Ca2C2O4 (ΔprH), the enthalpy of dilution
of a calcium chloride solution from 0.51697 mol/kg to
the final concentration m(CaCl2)fin should be subtracted
from ΔmixH:
T, °C
24.99
0
3
ΔprH = ΔmixH – ΔdilH.
24.98
These values determined experimentally are listed in
Table 2. The enthalpies of mixing and dilution (Tables
1 and 2) decrease in magnitude with an increase in the
NaCl content in water; however, the rates of these
changes are different. In a 1% NaCl solution, the
decrease in ΔdilH as compared to that in pure water is
~1.5 kJ/mol. The ΔmixH values in water and the salt
solution are nearly the same (see data in Table 1 and in
[1]). This leads to an increase in the exothermic charac-
ter of precipitation in solutions with low NaCl contents.
A further increase in the concentration of the back-
ground electrolyte brings about a decrease in heat
release in the course of mixing and dilution. However,
ΔdilH decreases at a slower rate, which leads to that the
ΔprH values in a 5% NaCl solution and in water become
the same. This, however, in no way means that the
mechanism of precipitation in these systems is identi-
cal. Figure 1 shows the temperature versus time curves
of mixing of calcium chloride and sodium oxalate in
water and in aqueous NaCl solutions. In water and in a
1% NaCl solution, precipitation is rapid. The resulting
flaky precipitate is presumably a mixture of the CaC2O4 ·
H2O and CaC2O4 · 2H2O crystal hydrates, which consti-
tute the major part of oxalate calculi [4]. As follows
from Fig. 1, a further increase in the background elec-
trolyte concentration significantly increases the reac-
tion time. This can be associated with some increase in
the solubility of calcium oxalate owing to a decrease in
24.97
2
24.96
1
24.95
24.94
24.93
5
500
1000
1500
2000
2500
τ, s
Fig. 1. Temperature vs. time curves of mixing of an aqueous
CaCl solution with an Na C O solution in water and in
2
2 2 4
aqueous NaCl solutions of various concentrations at 25°C:
(0) water, (1) 1% NaCl, etc. For clarity the curves are shown
with a shift along the ordinate axis; their linear portions
refer to the initial and final periods of the process.
RUSSIAN JOURNAL OF INORGANIC CHEMISTRY Vol. 54 No. 12 2009

2030
KUSTOV et al.
range. Second, even if a linear dependence were
obtained, the use of equations from [10] shifts the curve
in Fig. 2 down. Hence, the extrapolated value of the
enthalpy of precipitation will be 3–4 kJ/mol lower than
ΔprH in water [1, 2], which is also shown in Fig. 2. In
our opinion, a major reason for the phenomena
observed is that sodium chloride, in this case, is not an
indifferent electrolyte with respect to the process. It has
an effect on the association of Ca2+ and Cl– ions, thus
significantly changing the enthalpies of dilution and
mixing, and changes the time of precipitation and the
form of the nascent precipitate. A crystal hydrate is
mainly deposited in water, whereas calcium oxalate
precipitating in an NaCl solution is presumably not a
crystal hydrate. Hence, the ΔprH values in water and
aqueous NaCl solutions can hardly be compared.
–20.5
–21.0
–21.5
–22.0
–22.5
ACKNOWLEDGMENTS
This work was supported by the Russian Foundation
for Basic Research (project no. 08-02-12053-ofi).
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tional to m, is subtracted from the heat of the process,
also gives no way of obtaining a linear dependence
since its variation is rather small in this temperature
RUSSIAN JOURNAL OF INORGANIC CHEMISTRY Vol. 54 No. 12 2009