Download The solubility of Ca(OH) in extremely concentrated NaOH solutions

Cent. Eur. J. Chem. • 10(2) • 2012 • 332-337
DOI: 10.2478/s11532-011-0145-0
Central European Journal of Chemistry
The solubility of Ca(OH)2 in extremely
concentrated NaOH solutions at 25oC
Research Article
Attila Pallagi1, Ágost Tasi1, Attila Gácsi1, Miklós Csáti2,
István Pálinkó3, Gábor Peintler4, Pál Sipos1*
1
Department of Inorganic and Analytical Chemistry,
University of Szeged, Szeged, H-6701, P.O. Box 440, Hungary
Inspectorate for Environment, Nature Conservation and Water
in the Region of Lower Tisza, H-6721 Szeged, Hungary
2
3
Department of Organic Chemistry,
University of Szeged, H-6720 Szeged, Hungary
Department of Physical Chemistry and Materials Science,
University of Szeged, Szeged, H-6720, Hungary
4
Received 8 July 2011; Accepted 14 November 2011
Abstract: The solubility of Ca(OH)2 in aqueous NaOH solutions up to 12.50 M at 25oC has been determined. The solubility data obtained
for NaOH concentrations lower than 3 M was compared with those published in the literature. The solubility of Ca(OH)2 steadily
decreases with the increasing NaOH concentration. The solubility data obtained at a constant ionic strength (I = 1 M Na(Cl,OH))
enabled the determination of the conditional solubility product of Ca(OH)2(s) (lgLCa(OH)2 = – 4.10 ± 0.02). Formation of the hydroxo
complex CaOH+(aq) was invoked to describe the variation of [Ca2+]T with [OH–]T. Its conditional stability constant was found
to be lgKCaOH+ = 0.97 ± 0.02. The experimental protocol employed was proven to be suitable for accurate solubility determinations
in rapidly equilibrating systems comprising of highly concentrated, alkaline solutions and containing analytes in the ppm range.
Keywords: Solubility apparatus • Solubility • ICP AES • Ca(OH)2 • Concentrated electrolytes • Aqueous solutions
© Versita Sp. z o.o.
1. Introduction
The importance of calcium ions in the Bayer process,
which uses highly concentrated base solutions for
digesting bauxite, is well known [1]. Significant amount
of lime is employed during causticization, when the
carbonate content of process liquor is turned into
hydroxide and the solid CaCO3 is removed from the
system. Ca2+ ions are also used in the analytical chemical
practice for the quantitative removal of carbonate from
certain (but not all) concentrated base stock solutions [2].
In cement porewater (which contains Na+, K+, Ca2+ and
OH– ions) the solubility of Ca(OH)2 (portlandite) controls
the concentration of calcium ions [3]. Therefore accurate
knowledge of the equilibrium concentration of Ca2+ (and
the solubility product of Ca(OH)2) in highly concentrated,
aqueous NaOH solutions is of importance on several
counts. Surprisingly, relevant literature data is scarce.
Some papers [4-6] limit their analysis to [NaOH]T ≤ 1 M,
a concentration range far below the industrial relevancy.
NaOH concentrations up to 3.0 and 3.6 M were covered
in the works of Konno et al. [7] and Duchesne and
Reardon [3], respectively. However, the data appears in
contradiction with each other (see below). Moreover, we
believe that the solubility of Ca(OH)2 in solutions with the
industrially important concentration range of 4 M – 20 M
(latter is the solubility of NaOH at room temperature) is
fully unexplored.
There are several methods for determining solubilities
of (crystalline) solid compounds in aqueous solutions,
which vary both in accuracy, precision and convenience
[8]. The simplest way to determine solubility includes
* E-mail: [email protected]
332
Unauthenticated
Download Date | 6/14/17 5:55 PM
A. Pallagi et al.
the addition of an appropriate excess amount of solute
into the solvent in a container and gentle agitatation
until equilibrium is reached and then sampling for
analysis [9]. During the solubility measurements, special
attention must be paid to the accurate temperature
control, to the establishment of the equilibrium, to the
appropriate mixing conditions (to avoid attrition), and
to the problems associated with rapidly re-equilibrating
systems. Recently, we have built a simple and robust
solubility apparatus designed to meet the criteria
described above. The apparatus has been tested and
verified via solubility measurement of Ca(COO)2•H2O
at 20.00 ± 0.02oC, Further experimental protocol
enabled us to determine the solubility of Ca(OH)2 in
I = 1 M Na(Cl,OH) mixtures, conditional stability constant
of CaOH+aq and the solubility product of Ca(OH)2.
Additionally, solubility of pure NaOH solutions up
to 20 M base concentrations was determined.
The results of these measurements are reported below.
2. Experimental procedure
2.1. Materials
Concentrated NaOH (~20 M) stock solutions were
prepared from Millipore MilliQ water and a.r. grade
NaOH (Hungaropharma, 99% purity) and their carbonate
content was minimised as described previously [2]. The
density of the solution was determined picnometrically
and the exact concentration of the stock solution was
determined by acid-base titration, following gravimetric
dilution of the base solution. The NaOH solution was
stored in an airtight, caustic-resistant Pyrex bottle.
Solutions with various NaOH concentrations were also
prepared via gravimetric dilution.
NaCl (99.9%+ grade, Spektrum-3D product), CaCl2
(a.r. grade, Reanal product) and (NH4)2(COO)2 (a.r. grade,
Spektrum 3D product) were used as received. CaO (a.r.
grade, Reanal product) was heated overnight to 900oC
to remove CaCO3, and was stored in a desiccator, in
vacuo. The phase purity was analysed with powder x-ray
diffractometer (XRD) (Philips PW1710 instrument, Cu
Kα radiation), and compared with literature data [10].
2.2. Apparatus
Solubility measurement was performed with a
multiposition magnetic stirrer with 15 stirring positions
(VELP Multistirrer 15). The stirring rate was adjustable to
10 grades, with slowest setting being ca. 50 rpm. For that
purpose, a custom-designed, one-piece multiposition
water-jacketed glass pot was fitted. The pot was capable
of accommodating 15 individual, sealable polyethylene
vessels with a volume of 50-100 mL. The pot is thermally
equilibrated with a Julabo F12-MB thermostat.
The stability of the temperature in the various
thermostated positions was tested at nominal
temperature 25oC with a mercury-in-glass thermometer
with ±0.01oC precision (operational in the 22-27°C
region). Thermal stability was checked at three different
positions of the thermostated bath: closest to the water
inlet, in the middle, and at the farthest from the water
inlet. We have found that at the nominal temperature
the fluctuation was smaller than ± 0.02°C over 48 hours
at each position and that the mean temperature was
identical.
2.3. Methods of testing
To check instrument performance, test measurements
were performed by saturating distilled water with calciumoxalate-monohydrate (Ca(COO)2•H2O). This compound
was chosen, because the equilibrium concentration
of calcium in a solution in equilibrium with solid
CaCOO2•H2O (2.096 ppm at 20.00oC) [11] was in the
middle of the expected range of calcium concentrations
in our target solutions. Also, the preparation of calciumoxalate-monohydrate was straightforward [12]. Solutions
containing crystalline Ca(COO)2•H2O were allowed to
equilibrate for 6-18 hours. The initial and the equilibrium
solid phase were analysed with powder XRD after the
solubility experiments. The supernatant was withdrawn
using a WHATMAN, Anotop 25 (0.02 μm) syringe filter
and it was analysed spectroscopically (see below).
Based on the results of six independent measurements,
the concentration of Ca2+ in the saturated solution was
found to be 2.093±0.021 ppm, which is in agreement
with literature data.
2.4. Analytical measurements
Ca2+ determinations were made with a Thermo’s IRIS
Intrepid II ICP-OES spectrometer. The instrument was
externally calibrated with a calibration solution series
prepared from ICP Multi element standard solution XXIII
made by CertiPUR. The calcium concentration of each
sample was determined at three different wavelengths:
315.887 nm, 317.933 nm and 393.366 nm. The
calibration curve was fitted to the target concentration
range (0.5–100 ppm) and the parameters were set to
obtain the best available fit. The effect of the NaOH
matrix on the obtained Ca2+ concentrations was checked
with standard addition technique and was found to be
comparable to those caused by the dilution. Therefore,
the precision of the Ca2+ concentration measurements
can be estimated as ± 2%.
333
Unauthenticated
Download Date | 6/14/17 5:55 PM
The solubility of Ca(OH)2 in extremely
concentrated NaOH solutions at 25oC
4000
(a)
1500
3500
3000
I (a.u.)
1000
500
2500
I (a.u.)
The time required for reaching equilibrium
in NaOH/CaO systems was first checked. Supernatants
were withdrawn with a PALL Acrodisc Supor syringe filter
(0.45 mm) from representative mixtures in every 2 hours
for 24 hours. We found, that 6 hours was the sufficient
time for the equilibrium to be established under these
experimental conditions. This is in accordance with
results described in literature [7]. However, other cited
sources [3] suggested 24 hours of equilibration times for
these systems. Accordingly, every mixture was allowed
to equilibrate for at least 24 hours.
During the measurements, calculated amount of
ignited CaO (at least in ten-fold excess) was added to
the solutions. The XRD of the ignited solid compound
showed the phase purity of CaO (Fig. 1a). Also, based
on the comparison of the XRD traces with literature
data [10], this solid material turned out to be Ca(OH)2
(portlandite) after 24 hours contact with aqueous
alkaline solutions (Fig. 1b). The lack of diffraction peaks
characteristic to various crystalline forms of CaCO3
(i.e., calcite and aragonite) indicates the absence of
significant amounts of carbonate. Such diffraction peaks
were seen in the XRD trace of the unignited CaO (insert
in Fig. 1a) found at 2Θ = 23.10o, 29.50o, 36.02o, 39.5o
43.20o and 48.54o). This finding is important in the sense
that CaCO3 might cause significant variations in the
solubility of Ca2+ in alkaline solutions, in particular under
superambient conditions [7].
The solubility data obtained for pure NaOH solutions
is shown in Fig. 2. and in Table 1.
The equilibrium concentration of Ca2+ steadily
decreases with an increasing concentration of NaOH
This indicates a lack of formation of higher, soluble,
stepwise Ca(OH)n2–n hydroxo complexes (n > 2) even
at the highest concentration of the base, as their
formation would cause steady increase in the solubility
of Ca(OH)2 with increasing concentration of the base.
Samples with [NaOH]T > 12.50 M were also prepared.
Ca-determinations with ICP AES were hindered by the
high viscosity of these solutions and also by the very
low total concentrations of Ca2+ in these supernatants.
For these samples, equilibrium [Ca2+]T in the range of
less than 0.3 ppm (10 μM) can be estimated. The results
are in a reasonable agreement with some published data
[3], but are significantly higher than those presented
in other cited sources [7]. Close inspection of data in
[7] reveals that multiplying them by two eliminates
the discrepancies and brings all the three data sets
together.
Further measurements were performed with mixtures
containing NaOH and NaCl at [OH–]T + [Cl–]T = 1.00 M
(const.) and saturated with Ca(OH)2 (Table 1). It has been
shown in a previous paper (Fig. 3 in [13]), that the mean
activity coefficient of ions, γ± is very close to a constant
in mixtures at 1 M ionic strength. Therefore, conditional
formation of constants, that is, solubility products or
ion association constants, can be reasonably estimated
from data obtained in the presence of such mixtures.
2000
0
10
20
1500
30
40
2 Θ (deg)
50
60
1000
500
0
10
15
20
25
30
35
40
45
50
55
60
2 Θ (deg)
5000
(b)
4500
4000
3500
3000
I (a.u.)
3.1. Measurements in pure NaOH solutions
3.2. Measurements in Na(Cl,OH) mixtures
2500
2000
1500
1000
500
0
10
15
20
25
30
35
40
45
50
55
60
2 Θ (deg)
Figure 1. a:
XRD trace of the CaO sample ignited at 900oC
overnight (insert: the same sample before ignition);
b: XRD trace of the equilibrium solid phase obtained
after 24 hours of contact with an alkaline solution.
1.50
[Ca 2+]T (mM)
3. Results and discussion
1.00
0.50
0.00
0
2.5
5
7.5
10
12.5
[NaOH]T (M)
Figure 2.
The equilibrium concentration of Ca2+, [Ca2+]T as
a function of the total concentration of NaOH, [NaOH]T
in an aqueous solution saturated with solid Ca(OH)2, at
25.00 ± 0.02oC. ○: present work. ♦: [3]; ----: [7].
334
Unauthenticated
Download Date | 6/14/17 5:55 PM
A. Pallagi et al.
Table 1. Solubility
of Ca(OH)2 in aqueous NaOH solutions at varying and at constant ionic strength, and at 25 oC. As solubility data were not
tabulated in [7], the values shown here are approximations.
Present study
Varying ionic
strength
Constant ionic
strength,
I=1M
Na(Cl,OH)
Literature data
[OH–]T
(M)
[Ca2+]T
(mM)
[OH–]T
(M)
[Ca2+]T
(mM)
Ref.
0.444
0.547
0.643
0.745
0.848
0.950
0.991
1.497
1.996
2.488
3.001
5.000
7.500
10.000
12.500
1.6430
1.4080
1.1740
0.9860
0.8830
0.8080
0.7230
0.5450
0.4600
0.3290
0.3010
0.1200
0.0445
0.0347
0.0143
0.481
0.500
0.600
1.26
0.80
0.67
[3]
[7]
[7]
0.800
0.54
[7]
0.978
1.000
1.939
2.000
3.000
3.580
0.65
0.42
0.42
0.23
0.19
0.23
[3]
[7]
[3]
[7]
[7]
[3]
0.100
0.150
0.200
0.250
0.300
0.350
0.400
0.450
0.500
0.550
0.600
0.650
0.700
0.750
0.800
0.850
0.900
0.950
14.680
8.390
5.710
4.440
3.270
2.620
2.210
2.100
1.560
1.610
1.410
1.210
1.110
1.020
1.010
0.893
0.904
0.811
As it can be seen in the data shown in Table 1, the
equilibrium concentration of Ca2+ smoothly decreases
with the increasing concentration of the base, which the
trend is qualitatively similar to that obtained for NaCl-free
systems. However, the solubility of Ca(OH)2 increases
with added NaCl, which is similar to the observations
made in the literature by others [14].
There is no agreement in literature regarding the
identity of water soluble calcium species in equilibrium
with Ca(OH)2(s) [3,14-16]. If Ca2+(aq) is the only calcium
species in the solution [3,14], the solubility product of
Ca(OH)2 can be calculated as
(1)
From Fig. 3, it can be seen that the product [Ca ]
– 2
[OH
]T increases systematically as [OH–]T increases; the
T
extent of this variation is almost one order of magnitude
within this data series. Given, that the change in γ±,
when chloride ions are exchanged to hydroxide ions,
is only a few percent [13], activity coefficient variations
are most unlikely to encompass this kind of systematic
changes. However, other literature sources [15,16]
suggest, that in alkaline solutions, beside Ca2+(aq), the
formation of CaOH+(aq) is also significant. Therefore,
in systems containing NaOH solution and CaO as the
initial solid phase, the following chemical equilibria can
be considered:
CaO(s) + H2O
Ca(OH)2(s)
(2)
Ca(OH)2(s)
Ca2+ + 2 OH
(3)
Ca2+ + OH–
CaOH+
(4)
2+
Based on the XRD of the equilibrium solid phase, CaO
is fully converted to Ca(OH)2. The stability constant of
the complex CaOH+ is:
335
Unauthenticated
Download Date | 6/14/17 5:55 PM
The solubility of Ca(OH)2 in extremely
concentrated NaOH solutions at 25oC
1.00
(5)
T
From these data the solubility product of Ca(OH)2,
LCa(OH)2, and the stability constant of the hydroxo
complex CaOH+, KCaOH+ can be calculated as follows.
The mass balance equation relating to [Ca2+]T (assuming
that the equilibrium concentration of Ca(OH)2 in the
solution phase, [Ca(OH)2(aq)] is negligible)
(6)
[Ca2+]T[OH-]T2 x103 (M3)
K CaOH +
CaOH + 
=
Ca 2+  ⋅ OH − 
0.75
0.50
0.25
0.00
0.00
0.20
0.40
0.60
0.80
1.00
1.20
[OH-]T (M)
Figure 3. The
product [Ca ]T[OH–]T2 as a function of the [OH–]T
in Na(Cl,OH) mixtures with constant ionic strength (1.0 M)
and at 25.00 ± 0.02oC, where [Ca2+]T is the equilibrium
concentration of Ca2+ in an aqueous solution saturated
with solid Ca(OH)2 and [OH–]T is the total analytical
concentration of hydroxide ions.
This can be combined with the equilibrium expressions
to give
2+
18.0
(7)
16.0
14.0
12.0
[Ca 2+]T (mM)
When [Ca2+]T. [OH–]T2 is plotted vs. [OH–]T, a straight
line is obtained, when all assumptions are shown valid.
Note, that if the species of CaOH+(aq) do not form (i.e.,
KCaOH+ = 0), then the left hand side of the equation is
a constant and independent of [OH–]T. Fig. 3 shows that
this is not the case. In principle, LCa(OH)2 and KCaOH+
can be extracted from the intercept and slope. Eq. 7
exhibits a reasonable linearity (correlation coefficient
is ca. 0.988) even at very large substitution of Cl– for
OH–. This is true under the assumption, that the mean
activity coefficient is constant. However, the system in
this form of interpretation is very ill-conditioned, as the
intercept is very sensitive to relatively minor errors in
Ca-determinations. Therefore. the constants obtained
this way may only be considered as first approximations.
he solubility product and the stability constant obtained
from Eq. 7 are: lg LCa(OH)2 = – 4.0 and lg KCaOH+ =
0.90.
These data was also processed with the aid of
the ZITA suite of computer programs [17]. The results
of these fittings are shown in Fig. 4. The data set was
attempted to be fitted under the assumption that only
Ca2+(aq) would be present as water soluble calcium
species. As seen in Fig. 4, this assumption fails to
describe the experimental points (see dotted line in
Fig. 4.) However, when the formation of the complex
CaOH+(aq) is also included in the model, excellent fit it
obtained. The average relative difference between the
observed and calculated data points was 0.8%. The
formation constant of the CaOH+(aq) species was found
to be KCaOH+ = 0.97 ± 0.02 and the solubility product
of Ca(OH)2 was computed as lg LCa(OH)2 = – 4.10
± 0.02. Inclusion of Ca(OH)2(aq) in the model did not
10.0
8.0
6.0
4.0
2.0
0.0
0.00
0.20
0.40
0.60
0.80
1.00
[OH-]T (M)
Figure 4. The
total concentration of calcium, [Ca2+]T in solutions
in equilibrium with Ca(OH)2(s) in Na(Cl,OH) mixtures
with
constant
ionic
strength
(1.0
M)
and
at 25.00 ± 0.02oC, as a function of [OH–]T. Circles
represent experimental data. The dotted line was
obtained via assuming the presence of Ca2+(aq), as the
only water soluble calcium species and with an optimized
solubility product of LCa(OH)2 = (1.7 ± 0.1)×10–4 M3.
The solid line was calculated via including CaOH+(aq)
with KCaOH+ = 9.3 ± 0.5 M–1 and with a solubility product
of LCa(OH)2 = (7.9 ± 0.2)×10–5 M3.
improve the fit further and using this species in place
of the CaOH+(aq) yielded unacceptable fit. From this, it
seems unlikely, that Ca(OH)2(aq) is formed in sufficient
quantities in these systems.
336
Unauthenticated
Download Date | 6/14/17 5:55 PM
A. Pallagi et al.
The solubility product of Ca(OH)2 (portlandite) has
been researched in the literature [15] (lg LCa(OH)2 = – 5.19
at I = 0 M ionic strength). Our conditional constant at
I = 1 M is higher, which means, that in presence of NaCl
(up to 1 M concentration) the solubility of portlandite
increases [14]. The value of the formation constant of
CaOH+(aq) has been established within values found in
the literature at various ionic strengths (lg LCaOH+ = 0.64
at I = 3 M and 1.30 at I = 0 M).
4. Conclusions
We were successful in constructing an apparatus enabling
measurement of the solubilities of rapidly equilibrating
systems. It was used in determining the solubility of
Ca(OH)2 in a wide range of NaOH concentrations,
including under extremely basic conditions. Problematic
values in the literature have been corrected and it was
shown that these and related data could be obtained
under industrially relevant conditions. Data obtained at
constant ionic strength can be described in terms of the
formation of the complex CaOH+(aq) – the existence
of which has been disputed by some authors, while
supported by others. Spectroscopic evidence to confirm
or refute the existence of CaOH+(aq) is currently being
sought and the results of such experiments will be
published elsewhere.
Acknowledgements
This research was supported by a grant from the
Hungarian Science Foundation (OTKA 83889). The
financial help is highly appreciated.
References
[1] S.P. Rosenberg, J.D. Wilson, A.C. Heath, Light
Metals 19 (2001)
[2] P. Sipos, G.T. Hefter, P.M. May, Analyst 125, 955
(2000)
[3] J. Duchesne, E.J. Reardon, Cement Concr. Res.
25, 1043 (1995)
[4] N. Fratini, Ann. Chim. Appl. 39, 616 (1949)
[5] S. Diamond, Il Cemento 74, 149 (1977) (In Italian)
[6] T. Yuan, J. Wang, Z. Li, Fluid Phase Equil. 297, 129
(2010)
[7] H. Konno, Y. Nanri, M. Kitamura, Powder Technol.
123, 33 (2002)
[8] P. Sipos, G.T. Hefter, P.M. May, Rev. Sci. Inst. 70,
1481 (1999)
[9] W.J. Mader, L.T. Grady, in A. Weissberger,
B.W. Rossiter (Eds.), Physical Methods of
Chemistry. Part V: Determination of Thermodynamic
and Surface Properties (Wiley, New York, 1971)
[10] Q. Huang, O. Chmaissem, J.J. Capponi,
C. Chaillout, M. Marezio, J.L. Tholence, A. Santoro,
Physica C: Superconductivity 227, 1 (1994)
[11] B.C. Nelson, P.C. Uden, G.F. Rockwell, K.M. Gorski,
Z.G. Aguilera, J. Chrom. A 771, 285 (1997)
[12] A. Lakatos, K. Schrantz, E.A. Enyedi, A. Dombi,
Mennyiségi analitikai gyakorlatok, (Szegedi
Egyetemi Kiadó, Szeged, Hungary, 2008)
(In Hungarian)
[13] P. Sipos, M. Schibeci, G. Peintler, P.M. May,
G. Hefter, Dalton Trans. 1858 (2006)
[14] C. Christrov, N. Moller, Geochim. Cosmochim. Acta
68, 3717 (2004)
[15] R.A. Smith, A.E. Martell, Critical Stability Constants
(Plenum Press, New York, 1976)
[16] C.F. Baes, R.E. Mesmer, The Hydrolysis of Cations
(Wiley Interscience, New York, 1976)
[17] G. Peintler ZITA, A Comprehensive Program
Package for Fitting Parameters of Chemical Reaction
Mechanisms, Versions 2.1–5.0 (Department of
Physical Chemistry, University of Szeged, Szeged,
Hungary, 1989-2001).
337
Unauthenticated
Download Date | 6/14/17 5:55 PM