J. of Supercritical Fluids 35 (2005) 106–110
Solubility of crystalline alkali metal iodides in supercritical ammonia夽
Germán Sciaini a , Ernesto Marceca a , Roberto Fernández-Prini a,b,∗
a
b
INQUIMAE/DQIAQF, Facultad de Ciencias Exactas y Naturales, Universidad de Buenos Aires,
Ciudad Universitaria, Pabellón II, Buenos Aires 1428, Argentina
Unidad Actividad Quı́mica, Comisión Nacional de Energı́a Atómica, Av. Libertador 8250, Buenos Aires 1429, Argentina
Received 28 June 2004; received in revised form 29 November 2004; accepted 20 December 2004
Abstract
The solubilities of crystalline NaI, KI and CsI in supercritical ammonia (SCA) were determined for temperatures around 420 K and different
ammonia densities smaller than the critical density using the spectral charge-transfer-to-solvent (ctts) band of iodides. The optical absorptivity
of the ctts band was also determined with the potassium salt and shown to be density independent. The results show that the solubility of the
alkali metal iodides exhibit a very simple dependence on fluid density; this implies that over the density range we have explored, the first
solvation shell of the ion pairs is mostly complete even at the lowest studied density.
© 2005 Elsevier B.V. All rights reserved.
Keywords: Supercritical ammonia; Solubility; Charge-transfer-to-solvent; Alkali metal iodide; Solvation
1. Introduction
The use of solvated electrons as a means of treating hazardous chemicals and materials [1] is a very interesting procedure for destruction of chemical weapons and other substances of high risk. Solvated electrons are usually produced
by dissolution of alkali metals in ammonia. However, another
possibility of obtaining solvated electrons [2] is to generate
them by irradiation of iodides dissolved in different solvents,
especially in supercritical ammonia (SCA), thus allowing to
couple the medium reaction capacity with a fine control of
the solvent density.
In the course of our research programme to study the properties of solvated electrons in SCA, it became necessary to
understand the solvation of alkali metal iodides dissolved in
SCA. It was clear that, as a first step, it was necessary to determine the solubility of the salts in SCA over a range of fluid
densities. Recently, Brandt et al. [3] have reported the results
of a study of the phase diagrams of binary mixtures of NaI
夽 RFP and EM are members of Carrera del Investigador (CONICET).
∗
Corresponding author. Tel.: +54 11 67727175; fax: +54 11 67727121.
E-mail address: [email protected] (R. Fernández-Prini).
0896-8446/$ – see front matter © 2005 Elsevier B.V. All rights reserved.
doi:10.1016/j.supflu.2004.12.013
and KI with NH3 ; they described the features of phase equilibria and also gave the values of salt concentration along the
binary critical line. However, they explored a thermodynamic
region that is far removed from the one we were interested
in, moreover, no solubility data for the iodides were given.
In this work, we report the measurement of KI solubility at temperatures close to 423 K and densities in the range
4.3–10.0 mol dm−3 , and of NaI and CsI at 418.4 K over a
more limited density range. The method employed to determine solubilities was based on the measurement of the height
of the charge-transfer-to-solvent (ctts) UV peak of each iodide due to the UV-excitation of the dissolved alkali metal
iodide ion pair. It was also necessary to determine the molar
absorptivity of the ctts band of iodide under our experimental
conditions.
2. Experimental
Crystals of NaI, KI and CsI (pa) were used to prepare the
solutions and NH3 was AGA 99.99%. The strong increase of
the solubility of the alkali metal iodides with NH3 density was
studied in more detail for KI; in this case, it was necessary to
employ two different high-pressure high-temperature optical
G. Sciaini et al. / J. of Supercritical Fluids 35 (2005) 106–110
cells having sapphire windows. The two high p–T cells were
heated electrically and the temperature was maintained at the
desired value by means of a PID control system with a Pt
thermometer. One cell had an optical path of 20.4 mm and
the other of 2.08 mm; both cells had magnetic stirrers, which
were operated from outside the cells. The internal volumes
of both cells were about 3 cm3 .
In order to determine the molar absorptivity of the iodide
ctts band in SCA, subsaturated KI solutions were prepared;
this procedure required a very careful operation to obtain
good reproducible results. We introduced a known amount of
KI into the cell having the longest optical path in the form of
KI solutions in ethanol (pa), the solvent was carefully evaporated at reduced pressure, then ammonia was introduced
in the cell at room temperature using a manual pump, the
amount of added NH3 was such that it gave a density of
about 20 mol dm−3 at the desired temperature. The solution
of KI in NH3 thus prepared had a known concentration of
ca. 10−4 mol dm−3 ; supercritical conditions were achieved
by heating the cell gently while stirring its contents. Following this procedure, no salt precipitation occurred during the
heating period. The density of the supercritical solution at
the desired temperature was then carefully reduced by distillation until a density of about 9 mol dm−3 was attained;
this value was high enough to prevent salt precipitation. The
small loss of solute during the distillation was determined
by redissolving the NH3 extracted from the cell in a small
vessel containing water. Finally, the spectra of the solutions
in SCA were recorded at constant KI concentration and different ammonia densities obtained by NH3 addition (over the
range 9–19 mol dm−3 ). The cell was cooled down to ambient temperature condensing a solution of KI in liquid NH3 .
The optical absorption of this solution was measured at the
ccts maximum (265 nm), and using the molar absorptivity reported by Shapira and Treiner for this condition [4], we calculated the salt concentration in the cell. The amount of KI
found in the cell agreed with the one calculated by substracting the amount of iodide in the aqueous solution contained
in the small vessel from the mass of KI originally introduced
into the cell.
The solubilities of the alkali metal iodides were determined putting into the cells a small amount of crystals large
enough to saturate the solutions, they were evacuated to eliminate all traces of moisture. The cells were then heated to the
desired temperature, which could be fixed within ±0.2 K;
finally NH3 was introduced with a manual pump. It was
observed that the heat loss of the cells generated a small
temperature difference between the set-point value and the
actual temperature of the solutions inside the cells, which
depended on their design. Thus, a final temperature calibration was performed measuring the vapour pressure of pure
ammonia at different temperatures and comparing them with
those obtained from the equation of state. The temperature
calibration was carried out for each cell because their design was not the same, so precautions were taken to avoid
an undue influence of this difference. The equation of state
107
for pure ammonia used in the present work was that proposed by Tillner-Roth et al. [5]. These authors give the following values of the critical parameters of NH3 , which we
used in the present work: Tc = 405.4 K, pc = 11.333 MPa and
ρ1c = 13.218 mol dm−3 .
Although for KI, the solubility measurements in each cell
were performed at temperatures of 419.5 and 426.0 K, respectively, all the measured data could be used as a unique
series assigned to the average temperature of 423 K if plotted
against fluid density. This is supported by the fact that thermodynamic properties of dilute supercritical solutions, when
represented as function of fluid density, exhibit very weak
temperature dependence. This fact was noted previously in
other systems [6,7]; we shall expatiate this point below. The
value of the fluid density was calculated using the equation
of state [5] for the solvent and the experimental p, T values
as input.
The spectra were recorded with a Shimadzu UV-3101PC
in the range 210–350 nm to follow the ctts band; the resolution
employed was of 1 nm. We also measured the NIR spectra of
the solvent at different fluid densities up to 1500 nm corresponding to the combination-tone frequencies [8] using the
same spectrophotometer, we shall describe below the use we
made of the NIR bands of NH3 .
The spectra of the solutions did not change with time
even when solutions were kept in the cells for several
days.
3. Results
Solubility measurements were carried out at a reduced
temperature greater than 1.03 so that any effects of local density inhomogeneities in the neighborhood of the solute were
minimized [9]. Moreover, we have shown previously [10] that
these have only minor consequences on the value obtained
for the solubility. Hence, no local density inhomogeneities
are expected to affect our results.
The spectra of the alkali metal iodides in SCA showed
that the positions of the ctts band maxima (NaI, 261 nm; KI,
263 nm; and CsI, 266 nm) were not affected by changes in
fluid density over the explored range. Fig. 1 gives a representative example of the KI ctts band in SCA.
To determine the molar absorptivity of the ctts band,
we employed subsaturated solutions of KI in SCA at
418.2 K, the average of four values over the density range
9.6–18.6 mol dm−3 was εmax = 650 ± 30 m2 mol−1 . We assessed the effect of temperature on the molar absorptivity of the solutions increasing the temperature of the solutions to 453.0 K and determining again the molar absorptivity of KI solutions in SCA; at this temperature, we obtained
εmax = 705 m2 mol−1 showing that over a 30 K interval, the
change was only slightly bigger than the uncertainty of ε
at 418.2 K. Consequently, the difference of molar absorptivity at the temperatures used to measure the solubilities (cf.
Table 1) is less than 2%. The overall uncertainty of the mea-
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G. Sciaini et al. / J. of Supercritical Fluids 35 (2005) 106–110
Fig. 1. The ctts band of KI dissolved in SCA at 426 K and 9.5 mol dm−3 .
sured solubilities was estimated as 6%. The same value of
molar absorptivity was used for all the iodides.
We made use of the NIR spectrum of SCA to verify our
calculation of fluid density because it is the most important
thermodynamic variable. A typical NIR spectrum is illustrated in Fig. 2 for ρ1 = 9.5 mol dm−3 . The peak used for this
purpose was the one having λmax = 1032 nm. This band was
chosen because its optical absorption varied between 0.2 and
0.04 in the two cells and its position did not change over the
experimental fluid density range. The difference of absorption between the maximum of this peak and the absorption
at 1070 nm (indicated by arrows in Fig. 2) that corresponds
to a valley was plotted against ρ1 , resulting in a nearly linear
plot. All the other NIR bands had a very low or a very large
optical absorption in one of the two cells and the comparison
was not possible.
The solubilities of KI in SCA, c2 , are reported in Table 1
and those of NaI and CsI in Table 2. The results are plotted in
Fig. 3 as ln c2 against fluid density for all the measurements.
It is clear that the values obtained for the solubilities of KI in
each cell are very consistent with each other; it is noteworthy
Fig. 2. The NIR spectra of NH3 at the same temperature and density as in
Fig. 1. The dashed arrows indicate the positions of wavelengths 1032 and
1070 nm (cf. text).
Table 2
Solubility of NaI(cr) and CsI(cr) in supercritical ammonia at different fluid
densities at 418.4 K
Salt
p (MPa)
ρ1 (mol dm−3 )
103 c2 (mol dm−3 )
NaI
8.18
9.70
10.34
10.99
11.63
3.3
4.4
4.9
5.6
6.4
0.0060
0.0173
0.0309
0.0580
0.118
CsI
10.97
11.54
12.08
12.83
5.6
6.3
7.1
8.7
0.0075
0.0151
0.0332
0.133
that the six data points at intermediate densities belong to
measurements in different cells and they agree very well.
Fig. 3 shows that the solubilities of NaI and KI are very similar
and appreciably larger than that of CsI.
Table 1
Solubility of KI(cr) in supercritical ammonia at different fluid densities
ρ1 (mol dm−3 )
103 c2 (mol dm−3 )
T (K)
p (MPa)
419.5
9.7
10.3
11.0
11.6
12.1
4.3
4.8
5.5
6.2
6.9
0.0136
0.0252
0.0524
0.114
0.200
426.0
11.2
11.8
12.6
13.2
13.6
14.0
14.2
14.4
5.2
5.8
6.7
7.6
8.2
9.0
9.5
10.0
0.0398
0.0694
0.153
0.319
0.514
0.92
1.22
1.73
Fig. 3. ln c2 against the density of SCA. KI: () 419.5 K; (䊉) 426.0 K; NaI:
(,); CsI: ().
G. Sciaini et al. / J. of Supercritical Fluids 35 (2005) 106–110
4. Discussion
The values obtained for ln c2 were fitted with simple equations; for sodium and cesium iodides, we used a linear fit because the number of data points was small. On the other hand,
the data for KI plotted on Fig. 3 was fitted to the following
simple quadratic equation:
ln(c2 (423 K)/mol dm−3 )
= −16.996 + 1.5876ρ1 − 0.05288ρ12
(1)
It is interesting to compare the values of c2 at zero density, c2 (ρ1 → 0), obtained using the fitting equations, with the
concentration of the pure iodides due to their vapour pressure,
c2∗ (ρ1 = 0), i.e. the enhancement of the solubility of the salts
produced by the solvent. Cogin and Kimball [11] determined
the vapour pressure of alkali metal halides above 690 K and
proposed equations for the vapour pressures as function of
temperature that allowed us to calculate the vapour pressures
at our experimental temperatures. We have verified that the
value calculated for KI at 423 K is about a half of the value
given by JANAF [12], so that the equations proposed by Cogin and Kimball give a fair estimation of pvap at our experimental temperatures. For the three iodides, we used the values
given by Cogin and Kimball.
The enhancement factor (c2 /c2∗ ) is equal to the reciprocal of
the fugacity coefficient of the solute at infinite dilution φ2∞ ,
the values of the enhancement factor when ρ1 → 0, which
we denote as the limiting enhancement factor, are reported in
Table 3. It should be noted that the value for CsI has a larger
uncertainty due to the fact that there are only four solubility
data points for this salt. Their values are a clear indication
of the importance of the solvation of alkali metal iodides by
NH3 even at densities that are about a third of the critical
density.
The limiting enhancement factors are related to the
changes in Gibbs energy for the solvation of the ion pairs,
which are the dominant species in our solutions as discussed
below; the relationship is:
c2 (ρ1 → 0)
∞
solv G = −RT ln
c2∗ (ρ1 = 0)
= +RT ln Φ∞
2 (ρ1 → 0)
(2)
The values obtained for solv G∞ at 420 K were −99.6,
−92.7 and −87.2 kJ mol−1 for NaI, KI and CsI, respectively.
The available thermodynamic information only allowed us
Table 3
Limiting enhancement factors of alkali metal iodides in SCA
NaI
KI
CsI
T (K)
1019 c2∗ (ρ1 = 0)
(mol dm−3 )
418
423
418
0.83
1.23
5.8
108 c2 (ρ1 → 0)
(mol dm−3 )
20
4.16
4
− ln φ2∞
28.51
26.55
24.96
109
to estimate that the KI ion pair in NH3 at 423 K is solvated
by 8–12 NH3 first-neighbor solvent molecules. There is very
little information for the other two alkali metal iodides, so
we shall centre our discussion on KI, nevertheless, the same
behaviour is expected for the other salts. The almost linear
dependence of ln c2 with ρ1 is a clear indication that the
first solvation shell of the KI ion pair is almost unaffected by
changes in the NH3 density over our experimental range. This
behaviour suggests that the solvated ion pairs, the dominant
species in our solutions, do not alter their average molecular
structures in the density range explored, hence, it is interesting to enquire about it.
Hnizda and Kraus [13] calculated the association constant
for KI in liquid ammonia at 239 K from conductivity measurements; the value was 246 in the molarity scale. On the
other hand, Buback and Harder [14] have studied the dielectric behaviour of ammonia, including supercritical states. The
dipole moment of NH3 in the gas phase is 4.94 × 10−30 C m,
NH3 molecules interact with each other through their dipoles
and also participate in intermolecular hydrogen bonding. The
results of Bubak and Harder [14] show that the dielectric constant in the solutions studied by Hnizda and Kraus was 25,
while it is between 1.4 and 3.5 at 420 K over the density range
we covered in the present work, thus confirming that the ion
pairs are the dominant species in our solutions. This agrees
with the shift in the wavelength of the ctts peak with cation
size as noted above, which strongly suggests that the ion pair
is involved in the transition [15].
Evans et al. [16] report the energy values for the formation
of M+ –NH3 as −122, −88 and −70 kJ mol−1 for sodium,
potassium and cesium, respectively, and −31 kJ mol−1 for
I− –NH3 . Tongraar et al. [17] have studied the solvation of
K+ and I− ions in liquid ammonia by molecular dynamic
simulations; they calculated the energy of solvation of the
individual ions with different number of NH3 molecules in
the first solvation shell at 240 K and equilibrium vapour pressure. They found that the energy involved in the formation
of the singly solvated ions K+ –NH3 and I− –NH3 was −80.7
and −18.4 kJ mol−1 , respectively, values that are somewhat
smaller than the experimental ones [16]. The study of Tongraar et al. [17] gives for the equilibrium distances between
the ions and the nitrogen of NH3 0.29 nm for potassium ion
and 0.42 nm for iodide. They also reported that the radial
distribution functions obtained by simulation show that the
cation is surrounded by about 8 NH3 and iodide by 13 very
loosely bound solvent molecules. Energies of solvation decrease somewhat with the number of NH3 molecules in the
first solvation shell due to repulsion between them, so that
the average interaction energy per NH3 is smaller but close
to the values given above.
The M+ –NH3 interaction energy being so much greater
than that in I− –NH3 , suggests that the MI ion pair is more
strongly solvated around the cation than around the anion,
in agreement with the much sharper first peak observed in
the radial correlation functions of the individual ions [17].
There is evidence that an asymmetric solvation of cation and
110
G. Sciaini et al. / J. of Supercritical Fluids 35 (2005) 106–110
anion of an ion pair also occurs in water clusters [18], so it is
reasonable to expect a similar asymmetric solvation, though
less marked, in SCA.
The relative values of φ2∞ (ρ1 → 0) show that the solvation
free energy of the ion pair is more negative for sodium and
decreases as the size of the cation increases. This agrees with
the view that the absolute value of the cation solvation energy
per NH3 is large diminishing with the cation size, and that
of the anion is smaller. A more detailed calculation is not
possible at present since many factors that should be taken
into account are not known.
5. Conclusion
The maxima of the ctts bands of the alkali metal iodide
ion pairs dissolved in ammonia and the absorptivity coefficient did not change with density over the experimental
range of supercritical NH3 densities covered in this study,
i.e. 3.5–10.0 mol dm−3 . The values obtained for the solubilities of the alkali metal iodides were similar for Na+ and K+
and greater than those observed for CsI over the experimental ρ1 range. The enhancement factors at infinite dilution,
obtained from salt solubility and vapour pressure of the salts,
are very large implying that there is an important solvation of
the ion pairs. Moreover, the observed simple dependence of
ln c2 on ρ1 indicates that solvation of the ion pairs is already
very pronounced at the lowest experimental densities studied
in this work.
Acknowledgements
The authors are grateful to ANPCyT for partial economic
support. G.S. thanks CONICET for a graduate student scholarship.
References
[1] G.S. Pearson, R.S. Magee, Critical evaluation of proven chemical
weapon destruction technologies, Pure Appl. Chem. 74 (2002) 187.
[2] C. Frischkorn, M.T. Zanni, A.V. Davis, D.M. Neumark, Electron
solvation dynamics of I− (NH3 )n clusters, Faraday Disc. 115 (2000)
49.
[3] F.S. Brandt, P.M.A. Broers, T.W. de Loos, High pressure phase
equilibria in the system ammonia + potassium iodide, +sodium iodide, +sodium bromide, and +sodium thiocyanate: solid–liquid–vapor
and liquid–liquid–vapor equilibria, Int. J. Thermophys. 22 (2001)
1045.
[4] D. Shapira, A. Treiner, Charge-transfer-to-solvent spectra in liquid
ammonia, J. Phys. Chem. 70 (1966) 305.
[5] R. Tillner-Roth, F. Harns-Watzenberg, H.D. Bähr, Eine neue Fundamentalgleichung für Ammoniak, in: Proceedings of the 20th DKVTagung, vol. II, Heidelberg, Germany, 1993, p. 167.
[6] D.P. Fernández, G. Hefter, R. Fernández-Prini, The solubility of
solids in near-critical fluids. VI. CHI3 in CO2 revisited, J. Chem.
Thermodyn. 33 (2001) 1309.
[7] J.L. Alvarez, R. Fernández-Prini, M.L. Japas, Aqueous nonionic solutes at infinite dilution: thermodynamic description, including the
near-critical region, Ind. Eng. Chem. Res. 39 (2000) 3625.
[8] J.O.P. McBride, R.W. Nicholls, The vibration-rotation spectrum of
ammonia gas I, J. Phys. B 5 (1972) 408.
[9] R. Fernández-Prini, Study of local density enhancement in nearcritical solutions of attractive solutes using hydrostatic hypernetted
chain theory, J. Phys. Chem. B 106 (2002) 3217.
[10] G. Sciaini, E. Marceca, R. Fernández-Prini, Intermolecular solventsolute energies for thermodynamic and spectroscopic properties of
solutes in near-critical solvents, Phys. Chem. Chem. Phys. 4 (2002)
3400.
[11] G.E. Cogin, G.E. Kimball, The vapor pressure of some alkali halides,
J. Chem. Phys. 16 (1948) 1035.
[12] M.W. Chase, NIST-JANAF thermochemical tables, J. Phys. Chem.
Ref. Data, Monograph No. 9, ACS-AIP-NIST, 1998.
[13] V.F. Hnizda, C.A. Kraus, Properties of electrolytic solutions.
XXXIX. Conductance of several salts in ammonia at −34◦ by a
precision method, J. Am. Chem. Soc. 71 (1949) 1565.
[14] M. Bubak, W.D. Harder, The static dielectric constant of ammonia
to high pressures and temperatures. II. Representation of expeirmental data via Kirkwood–Fröhlich’s equation, Ber. Bunsenges. Phys.
Chem. 81 (1977) 609.
[15] M.J. Blandamer, M.F. Fox, Theory and application of charge transfer
to solvent spectra, Chem. Rev. 70 (1970) 59.
[16] D.H. Evans, R.G. Keese, A.W. Castleman, The association of ammonia with halide ions in the gas phase, J. Chem. Phys. 86 (1987)
2927.
[17] A. Tongraar, S. Hannongbua, B.M. Rode, Molecular dynamics simulations of potassium ion and an iodide ion in liquid ammonia, Chem.
Phys. 219 (1997) 279.
[18] G.H. Peslherbe, B.M. Ladany, J.T. Hynes, Structure of NaI ion pair
in water clusters, Chem. Phys. 258 (2000) 201.