ATOMIC ENERGY ÄJKÄ L'ENERGIE ATOMIQUE OF CANADA … · RESUME On présente des résultats de...
Transcript of ATOMIC ENERGY ÄJKÄ L'ENERGIE ATOMIQUE OF CANADA … · RESUME On présente des résultats de...
AECL-8370
ATOMIC ENERGY Ä J K Ä L'ENERGIE ATOMIQUEOF CANADA LIMITED 1 ^ 2 ^ DU CANADA, LIMITEE
HEAT CAPACITY DATA FOR SELECTED CESIUM- AND IODINE-CONTAINING
ELECTROLYTES IN WATER AT ELEVATED TEMPERATURES
RESULTATS DE MESURES DE CAPACITE CALORIFIQUE D'ELECTROLYTES A
CESIUM ET IODE DANS DE L'EAU A DES TEMPERATURES ELEVEES
P. P. S. Saluja, J. C. LeBlanc
Whiteshell Nuclear Research Etablissement de recherchesEstablishment nucléaires de Whiteshell
Pinawa, Manitoba ROE1LOSeptember 1985 septembre
ATOMIC ENERGY OF CANADA LIMITED
HEAT CAPACITY DATA FOR SELECTED CESIUM- AND IODINE-CONTAINING
ELECTROLYTES IN WATER AT ELEVATED TEMPERATURES
by
P.P.S. Saluja and J.C. LeBlanc
Whiteshell Nuclear Research EstablishmentPinawa, Manitoba ROE 1L0
1985 September
AECL-8370
RESULTATS DK MESURES DE CAPACITE CALORIFIQUE D1ELECTROLYTES A
CÉSIUM ET IODE DANS DE L'EAU À DES TEMPERATURES ÉLEVÉES
par
P.P.S. Saluja et J.C. LeBlanc
RESUME
On présente des résultats de mesures de capacité calorifique
d'halogénures de césium et d'iodate de césium dans de l'eau à 0,6 MPa et à
une température s'échelonnant de 25 à 100°C. On a évalué les fonctions de
capacité calorifique moléculaire partielle, C ~(T), en appliquant le modèlep,2
d'interaction des ions de Pitzer pour extrapoler les résultats calculés decapacité calorifique moléculaire apparente, $ , à la dilution infinie. Les
c
valeurs de C ~(T) dans l'eau dépendent fortement du sel de césium et de la
température. Pour tous les sels, les fonctions C „(T) présentent des
maxima s'échelonnant de 75 â 93CC. On peut se servir de ces résultats pour
déterminer les diverses propriétés thermodynamiques des systèmes aqueux à
césium et iode à des températures élevées.
L'Energie Atomique du Canada, LimitéeEtablissement de recherches nucléaires de Whiteshell
Pinawa, Manitoba ROE 1L01985 septembre
AECL-8370
HEAT CAPACITY DA1A FOR SELECTED CESIUM- AND IODINE-CONTAINING
ELECTROLYTES IN WATER AT ELEVATED TEMPERATURES
by
P.P.S. Saluja and J.C. LeBlanc
ABSTRACT
The results of heat capacity measurements are presented for cesium
halides and cesium iodate in water at 0.6 MPa and in the temperature range
25 to 100°C. Partial molar heat capacity functions, C ~(T), were evaluated
by applying Pitzer's ion-interaction model for the extrapolation of calcu-
lated apparent molal heat capacity data, <]> , to infinite dilution. The
C „(T) values in water depend strongly on the cesium salt and the tempera-P> -o
ture. For all salts, the C „(T) functions show maxima in the 75 to 93°C
range. These data can be used to determine various thermodynamic properties
of aqueous cesium and iodine systems at elevated temperatures.
Atomic Energy of Canada LimitedWhiteshell Nuclear Research Establishment
Pinawa, Manitoba ROE 1L01985 September
AECL-837C
CONTENTS
Page
1. INTRODUCTION 1
2. GENERAL APPROACH 2
3. EXPERIMENTAL 5
3.1 APPARATUS 5
3.2 MATERIALS AND SOLUTION PREPARATION 6
3.3 OPERATIONAL PROCEDURE 6
4. RESULTS AND DISCUSSION 9
4.1 PRIMARY RESULTS 9
4.2 DATA TREATMENT USING THE ION-INTERACTION MODEL 94.3 PRESSURE COEFFICIENTS OF C° 2(T) 11
4.4 COMPARISON WITH L^ZRATURE DATA 12
5. CONCLUSIONS 15
ACKNOWLE DGEMENT S 15
REFERENCES 16
TABLES 19
FIGURES 27
1. INTRODUCTION
Thermodynamic data for aqueous species at elevated temperatures
are required to understand and model various chemical processes that might
occur in the vicinity of a nuclear fuel waste disposal vault [1,2J. Much
attention has recently been given to cesium-water and iodine-water chemistry
because cesium and iodine are major fission products in used nuclear fuel
[3]. Preliminary analysis of the nuclear fuel waste disposal scheme [4] has
emphasized the importance of limiting the release of 135Cs and 129I to the
environment. Thus, our current interest lies in understanding the inter-
actions of aqueous cesium and iodine species with major constituents of
groundwaters and with various mineral surfaces. Gibbs energy data, G_(T),
at elevated temperatures are needed to calculate solubilities and chemical
speciation of aqueous cesium and iodine systems [2].
The G~(T) can be calculated at any temperature (T) using the fol-
lowing thermodynamic relation:
G°(T) = G°(TR) - S°(TR)(T-TR) - T/ Kf dT + / c" 2(T)dT (1)
The room temperature (T ) data, G (T ), and entropy, S.(T_), are usuallyK. L K ù R
available in the literature, in regularly updated compilations [5]• To—o
evaluate G (T), we need to determine, experimentally, the partial molar heatcapacity function, C 9(T), for the desired chemical species,p ,z
While C . data for a few aqueous cesium and iodine species are
available for room temperature [6,7], very few data [2,5] are available for
elevated temperatures, because measurements at elevated temperatures have
been difficult. Consequently, the available data for aqueous species are
- 2 -
often inconsistent, as discussed by several authors [8,9]. For example,
there are only two sets of literature C data for aqueous CsCl and Csl atP»2
elevated temperatures [10,11]. In 1964, Mitchell and Cobble [10] published
C „ data for aqueous Csl in the 10 to 90°C range, based upon integral heatP> -oof solution measurements. In 1969, RUterjans et al. [11] published C „
data foi aqueous CsCl and Csl in the 30 to 130°C range, based upon adiabatic
calorimetric measurements. Except for the 30 to 50°C range, the two sets of
data do not agree. The data from the adiabatic calorimetry [11] show a
maximum at about 80°C (see Figure 1), whereas the data of Mitchell and
Cobble [10] show no maximum in the 0 to 100°C range.
Recent advances in flow microcalorimeters [12-15] allow improved
precision of C data and make measurements for aqueous solutions at ele-
vated temperatures easier. Thus, we determined the heat capacities of fis-
sion products in water at temperatures up to 100°C. We applied Pitzer's
ion-interaction model [8,9,16-21] to our data and obtained partial molar
heat capacities as a function of temperature, C ~(T). In this report, we
present measured heat capacities, C (T,m) (where m is the molality of the
solution in mol «kg"1), calculated apparent molal heat capacities, <j> (T,m),
and calculated C 9(T) functions for four aqueous cesium electrolytes, CsF,
CsCl, Csl and CsI03, up to 100°C and at a constant pressure of 0.6 MPa.
2. GENERAL APPROACH
As shown in Equation (1), various thermodynamic properties for—o
aqueous chemical species can be calculated from the C -(T) function inp> -o
water, in combination with room-temperature data [5]. The C 9(T) function
is identical to the apparent molal heat capacity at infinite dilution, <b (T)c
at m = 0. However, we can carry out heat capacity measurements and deter-
mine <j> only in the finite concentration range, and a theoretical, or empir-
ical, approach is required for extrapolating $ (T,m) to infinite dilution.
- 3 -
A successful approach for extrapolating <f> (T,m) to infinite dilu-
tion is the application of the ion-ion interaction model of Pitzer [8,16].
Pitzer's model includes the usual Debye-HUckel term for long-range electro-
static interactions, important in dilute solutions. Furthermore, the model
includes a virial coefficient term for short-range chemical interactions
encountered in concentrated solutions, e.g., dielectric saturation, solva-
tion, électrostriction and ion association effects [22]. Recent applica-
tions of this model are (1) the extrapolation of experimental $ (T,m) data
for aqueous NagSO^ to infinite dilution [13], (2) the calculation of high-
temperature thermodynamic properties, e.g., osmotic and activity coeffi-
cients, from isopiestic measurements [20], and (3) the calculation of
thermodynamic properties of multi-component aqueous electrolyte solutions at
25°C [18].
tion for the apparent molal heat capacity, * , of an electrolyte-water
Pitzer's ion-interaction model [8,16] yields the following equa-
he appi
(MX-H2O) system:
*c = ~Cl,2 + V|ZMZX |AJ 1 2b — 2 V V
where
B ,
V ßMX + T ( 3 )
a 1
The quantities in Equation (2) are defined as follows: v is the sum of the
number of cations, v .., and anions, v„_, for electrolyte MX (v = 2 for 1:1
electrolyte); Z and Z are the charges on the cation and anion; A is theM A J
Debye-HUckel slope for heat capacity as a function of temperature, and was
obtained from the tables of Bradley and Pitzer [23]; I is the ionic
strength, I » J Z c.Z? where c is the concentration of solution in mol»dm"
(for 1:1 electrolyte, I = c in mol'dm"3); and R is the gas constant. The
ßj_. , ß^. and <!_. are temperature- and pressure-dependent virial coef-
ficients. These are related to short-range interionic forces and to the
indirect forces arising from the solvent. The b and a are constants whose
values were somewhat arbitrarily fixed by Pitzer [8,16] at 1.2 and
2.0 mol • kg , respectively. These values of b and a were based on the
success of the model in the treatment of a large body of recent experimental
data for solutions at elevated temperature and pressure [13,16-18,21]. The(1)J J
higher order terms, ß and C„„, which are important only at concentra-
tions above 1.0 mol'kg"*, a r e n o t required for the present <(> (T,m) data.
Therefore, the simplified Pitzer's equation [8,16] for this work can be
written as
2v vNow, two procedures are discussed for applying Pitzer's equation
[8,16] to our 4> (T,m) data to obtain the C „ value at a single temperature,c -o p>
as well as the temperature-dependent C 0(T) function. Firstly, an applica-
tion of Equation (4) to the $ (m) data at a single temperature provides
isothermal values of C° „ and the ion-ion interaction parameter,This fitting procedure is a useful check on the precision of C measure-
p,2
ments at a constant temperature. The standard deviations of these iso-
thermal fits show any significant errors in the individual data points re-
sulting from either operator or instrumental errors. A second procedure is
to apply Pitzer's equation to the complete set of <)> (T,m) data for a given
MX-H2° system. The data set usually consists of 25 to 30 experimental
points in the 25 to 100°C temperature range and in the concentration range
0.05 to 1.0 mol*kg~* (or to the solubility limit, whichever is lower). The
temperature dependence of C _, and also the ion-interaction parameter,
ß™ , are adequately described by the functional form F(T)[C° „(T) orToll p'2
ß ^ X(T)] = A/T + B + CT. Evaluation of these six temperature-dependent
parameters (A , B , C ; A-, B and C ) requires a large number of experi-
mental heat capacity data points with a high degree of precision,±1 x 10-1* J'K-^g"1 (±0.01%) in C „ measurements. The C data should
Pi2 p,2
also be obtained over a wide range of temperature and concentration. Next,
we discuss the modifications made to the instrumentation and peripheral
devices to achieve the required precision of ±0.01% in the C „ data.P»2
- 5 -
3. EXPERIMENTAL
3.1 APPARATUS
In the last decade, flow calorimetry [12-15] has been established
as an excellent tool for thermodynamic studies on aqueous solutions at, or
near, room temperature. The fast response and the high sensitivity that
make flow calorimetry a powerful tool at room temperature have proven
equally valuable at higher temperatures and pressures. Our elevated-
temperature C flow microcalorimeter and flow densimeter system, shown
schematically in Figure 2, are modifications of the room-temperature version
of the Picker design [12]. The major difference from the original design is
the use of corrosion-resistant Pt-10% Rh, instead of stainless steel, for
the tubing of the calorimeter cell, so that the cell tubing is not corroded
by the salt solutions at elevated temperatures. Other modifications have
also been made to the sample injection system, the back pressure regulation
and the delay line between the twin cells.
The sample solutions were contained in a stainless steel cylinder
having a Teflon liner and piston assembly. The solutions were injected into
the flow calorimeter by using nitrogen as a driving gas. The use of a
Teflon liner prevented corrosion and eliminated the possible release of
impurities from the cylinder to the sample solutions. The Teflon also
promoted reproducible movement of the piston, allowing maintenance-free
performance and rapid cleaning. A fine capillary (0.005 cm I.D.) was used
at the sample outlet of the microcalorimeter to regulate the flow rate of
the sample. The delay line was maintained at, or near, the calorimeter bath
temperature, rather than at ambient temperature. The flow densimeter was
also maintained at the calorimeter bath temperature. Simultaneous measure-
ments of the volumetric heat capacity and density provided the heat capa-
city, C , in J "K~ l«g~ 1. Thus, both apparent molal heat capacities, <t>
P c(J*K~*«mol"1) and molal volumes, $ (cm3#mol~^), were determined at the same
temperature and pressure during a single experiment.
- 6 -
3.2 MATERIALS AND SOLUTION PREPARATION
The water used as a reference solvent and also to prepare solu-
tions for C measurements was distilled and deionized using a Milli-Q water
system. The water was degassed for one hour at about 50cC immediately
before use. This prevented formation of gas bubbles in the twin calorimeter
cells.
The chemicals used in this work were obtained from Alfa Products,
except for CSIO3, which was obtained from Atomergic Chemetals Corporation.
The manufacturer listed the purity of the CsF, CsCl, Csl and CsIO3 as 99.9%,
99%, 99.9% and 99.5%, respectively. This was not confirmed by us. All
chemicals were used without further purification. The salts were dried
overnight in a vacuum oven at about 127°C and then transferred directly into
a desiccator containing anhydrous CaSO^. The solutions prepared from degas-
sed water were transferred to evacuated air-tight cylinders. In most cases,
solutions were used within two days of preparation. In any case, C results
were unaffected by aging of the sample solutions, and all duplicate deter-
minations agreed to within ±1 x 10"1* J'K"1^"1.
3.3 OPERATIONAL PROCEDURE
For each heat capacity measurement, the circulating ethylene gly-
col bath was first regulated to within 0.002°C of the desired temperature in
the 25 to 100°C range. The water from the nitrogen-gas-driven injection
system was then allowed to flow through the densimeter, and the working and
reference cells of the microcalorimeter, at the same volumetric flow rate.
To achieve a steady water-water baseline, W/W, both cells were heated with
an equal and constant base power, W /2. This base power causes a small
increment, AT, detected by the thermistors downstream on both the cells.
- 7 -
The experimental run was begun by introducing the sample solution
into the working cell by rapidly switching the valve on the injection system
from the water reservoir to the solution reservoir. When the sample solu-
tion displaced the -rater in the working cell, it caused the temperature
detected at the thermistor to change from the original increment AT. This
AT perturbed the W/W baseline, and the measured S/W signal, AW, was the
amount of power necessary to counterbalance this perturbation. This compen-
sation power, AW, was automatically supplied to the. heater of the working
cell heater by the servo-mode circuit of a thermal detector. The power, AW,
added to the base-power, W /2, caused the temperature of the working cell
thermistor to return to the value that gave the original W/W baseline
(see Figure 3). When the sample solution also displaced water from the
reference cell, so that the solution flowed through both cells, the autom-
ated servo circuit removed this compensation power, AW, from the base power.
Once again, a new solu'ion-solution, S/S, baseline was observed. This
baseline was identical to the W/W baseline observed at the start of the
experimental run. The valve on the calorimeter inlet was switched again to
liquid water. Water displaced the solution from the; working cell while the
solution was still flowing through the reference cell. In this case, a new
compensation power, AW1, was removed from the base power, bringing the
temperature of the working cell thermistor back to the value that gave the
S/S baseline prior to re-admitting water. Once water displaced the solution
from the reference cell, servo operation added power back to the working
cell heater to bring the thermistor temperature back to the value that gave
the W/W baseline at the start of the run. Measurements of both compensation
heating powers, AW and AW1, were then used to calculate the volumetric heat
capacity, o2 (in J«K~1»cm~3), of a sample solution using the following
equations:
a2 = ° (1 ~ w~ } and 02 = a S F (5)
1 o l (1 + f-)o
where o is the volumetric heat capacity of water at the operating tempera-
ture and pressure.
- 8 -
Solution heat capacities, C . (in JMC^'g" 1), were calculated by
combining these volumetric heat capacities, a2, with the measured densities,
d2, at the operating temperature and pressure, using the following
equations:
AW
- yr > ando
where C and d are the heat capacities and densities, respectively, of the
solvent (water) at the calorimeter mean temperature and operating pressure.
These were obtained from the tables of Kell [24].
As in other calorimetric experiments, a separate electrical-
calibration experiment was required to determine AW, AWr and W . This was
achieved by applying the appropriate compensation current to the Zener
diode, or heater, of the working cell while water was flowing through both
the cells. This electrical-calibration experiment simulated the ffl signal
obtained during the sample run. W was determined by multiplying the cur-
rent to the working cell heater by the voltage of the Zener diodes. For
high precision, the calibration signal was chosen to be nearly equal to the
experimental AW or AW' signals obtained for the sample solution.
The precision of a heat capacity determination depends on factors
such as (1) the precision of the measurements of the base power (W ) and
the two compensation powers (AW and AW'), (2) the temperature difference
between the working and reference side of the twin cell calorimeter, (3) the
temperature regulation of the thermostating bath, and (4) the error in de-
termining the molalities and the densities of the solutions. On the basis
of the reproducibility of our experimental data, the overall precision is
±1 x lO"4 JpK~l'g~l for the heat capacity determination and 1 x 10" •" g -cm" 3
for the density determination [21].
- 9 -
4. RESULTS AND DISCUSSION
4.1 PRIMARY RESULTS
The C 9(T,m) values for CsF, CsCl, Csl and CsIO3 are presented in
the third column of Tables 1 to 4, respectively. These values were used to
calculate the apparent molal heat capacities, 4> (T,m), from the following
equation:
1000[C (T,m)-• (T,m) - M C (T,m) + Eii E (7)c i. p,Z m
where M2 is the molecular weight of the dissolved solute. The 0 (T,m)
values are given in the fourth column of Tables 1 to 4. Based on the pre-
cision of ±1 x 10"1* J'lC^-'g-1 in the C ,(T,m) measurements, the precision
of the calculated if (T,m) values is ±1 J »K~ 1«mol~ ̂ at 0*1 «mol «kg and dfl.l
J#K~^»mol~^ at 1.0 mol'kg"^. Sich precision in <j> (T,m) values is acceptable_ c
for obtaining C values as a function of temperature, because even an
approximate C ~(T) function can yield quite accurate values of equilibrium
constants of reactions Involving desired cesium species [20,21].
4.2 DATA TREATMENT USING THE ION-INTERACTION MODEL
Isothermal C „ values for CsF, CsCl, Csl and CsIO3 were obtained
I the $ (m) data at a
equation for 1:1 electrolytes:
by fitting the $ (m) data at a single temperature to the simplified Pitzer's
T7T
- 10 -
These values are listed as the top entry in Table 5. However, we are mainly—o
interested in C „(T) values as a function of temperature. This requires
fitting an entire set of $ (T,m) data, as a function of temperature and
concentration., for extrapolation to infinite dilution.
The temperature-dependent equation for $ (T,m) can be represented
by rewriting Equation (8) for 1:1 electrolytes as follows:
A (T)<t>c(T,m) = C° 2(T) + ~ 2 ~ ln(l+1.2Ä) - 2RT2m ß ^ V ) • (9)
The A (T) values were obtained from the tables of Bradley and
Pitzer [23]. The temperature dependence of C and of the ion-interaction
parameter, ß™ , were fitted to the following equations:
AcC p 2(T) = — + Bc + CcT (10)
vwhere A , B , C , A_, Bg and Cß are constants. These constants were ob-
tained by fitting Equations (9), (10) and (11) to our experimental values of
<t> (T,m) for each electrolyte. Subroutines DEM1NG and EQUATION, written by
Pitzer et al. [8,16], were used for this purpose. The values of the six
constants for CsF, CsCl, Csl and CsIO3 are given in Tables 6 and 7. All
isothermal C values, calculated from Equation (10), are within
±l.t> J«K *mol~^ of the experimental values. The latter were obtained by
fitting Equation (8) to the <)> (m) data at a single temperature (see
Table 5).
For all cesium salts in water, the C „-temperature plot shows a
maximum in the 25 to 100°C temperature range (see Figure 4). Using the
derived C ~(T) functions, we calculated the temperature, T , , ofp,2x i- » max, he
- 11 -
maximum heat capacity and the maximum C „(T , ) value at T , . Thesep,2 max,he max,he
two quantities are given in Table 8. The shape of the C ?(T) function
depends on the nature of the ionic species in solution. Factors such as
ionic charge-size ratio, available H-bonding sites, and structure-making or
structure-breaking characteristics of the ionic species contribute to the—o
specific C „-temperature profile [21].
4.3 PRESSURE COEFFICIENTS OF C p(T)
The pressure dependence of the heat capacity, the enthalpy H , and
the Gibbs energy of a species can be calculated from the temperature depen-
dence of the partial molar volumes, V_(T), using the following standard
thermodynamlc relations:
(12)
V°(T) - Tfe(T)
(13)
(14)
We applied the following equation, derived from Pitzer's model
[8,16], to our volumetric data, <t>.r(T,m), calculated from about 30 density
measurements, for each cesium elsctrolyte-water system as a function of
temperature to get V (T):
V°(T)A (T)- ~ - lnU+1.2,40 + 2RTm ßj£ (T> (15)
- 12 -
where A (T) represents the temperature-dependent Debye-HUckel slope for
volume and was obtained frooi the tables of Bradley and Pitzer [23].
ß (T) is the ion-ion interaction parameter for electrolyte MX as a func-
tion of temperature. The temperature-dependence of V„(T) was adequately
described by Equation (16), which is analogous to Equation (10), used to
describe the temperature dependence of C „:
AV°(T) = ^ + Bv + Cv T (16)
The constants A , B and C are given in Table 9 for CsF, CsCl, Csl and
CsIOt>. The Dressure coefficients of C „(T), calculated using Equation
(12), are listed in Table 10 for the cesium halide-water and the CsIO 3~
water systems.
4.4 COMPARISON WITH LITERATURE DATA
Our <(> (T,m) results at 25°C and 0.6 MPa for the CsF-H2U, CsCl-H2O
and CsI-H20 systems are compared in Figure 5 with the results of Fortier et
al. [6] and Desnoyers et al. [7]. The literature results were obtained at
25°C and 0.1 MPa, and therefore w.re extrapolated to 0.6 MPa, using the
pressure coefficients of C 2^T) (Table 10) determined in our work. The
agreement between the two sets of data is very good (see Figure 5).
—oOur room-temperature C „ values for CsF, CsCl and Csl at 0.6 MPa
and those of Fortier et al. [6] and Desnoyers et al. [7] at 0.1 MPa are
given in Table 11. The literature C . data, extrapolated to 0.6 MPa, using
our measured pressure coefficients of C (Table 10), are only about
1 J'K^'mol"1 less negative than C° „ results obtained in the present work.
Therefore, the agreement between the two sets of data at 25°C and 0.6 MPa is
very good.
- 13 -
Figure 6 shows our C „ values, from Equation (10), for CsCl at
elevated temperatures and the C „ values of Rüterjans et al. [11] • The
literature results are based on only three sets of high-concentration C „p,Z
data for aqueous CsCl solutions, obtained from adiabatic calorimetricmeasurements. The agreement in the 25 to 70°C temperature range is reason-
—oable. Also, both sets of data show maxima in the C „-temperature plot.
-o p>
However, the temperature of maximum C „ from the literature data [11] is
85°C, as opposed to 77°C obtained in the present study. The bottom curve in
Figure 6 shows the deviation in the two sets of data, which increases from
4 J'K-i-mol-1 at 70°C to 20 J'K-^mol"1 at 130°C. The following factors may
account for the deviation. Firstly, adiabatic calorimetric experiments
require vapour space corrections, because both liquid and vapour phases of
the sample coexist during the course of C measurements [11,25]. The
large vapour space corrections at higher temperatures, and also the uncer-
tainties associated with their determination, decrease the precision of the
measured C data [11]. The present flow calorimetric method does notP»2
require vapour space corrections. Secondly, only three <f (m) data points,derived from C „ measurements on concentrated solutions (0.498, 1.084 and
P.2
1.606 mol«kg *)> were used to extrapolate to infinite dilution by an empiri-
cal equation, which is neither theoretically valid at higher concentrations
nor consistent with the Debye-HUckel equation at the lower end of the con-
centration scale. For example, Rüterjans et al. [11] stated that the ex-
perimental slopes in the 0.498 to 1.606 mol«kg"1 concentration range differ
considerably from the predicted theoretical slopes of the Debye-Hlickel
theory. Thus, these extrapolations gave only approximate values of C „.
We also compared our present C „ results for Csl with two inde-
pendent sets of data published in 1964 [10] and in 1969 [11]. Figure 7
shows that, up to 77°C, all three sets of data are in reasonable agreement.
Our C . values are about 10±3 J-K^'raol"1 less negative than the C° „P,2 ° p,2
values of Mitchell and Cobble [10], obtained from the temperature dependence
of the integral heat of solution, AH , and about 13 ±2 J'K"1^©!"1 less nega-_ s
tive than the C „ values of Rüterjans et al. [11], obtained from adiabatic
calorimetric measurements. Furthermore, the data of Rüterjans et al. [11]
and our results show a maximum at nearly the same temperature, 80°C, in the
- 14 -
C „-temperature plot. In contrast, Mitchell and Cobble [10] did not ob-—o
serve a maximum in the C .-temperature plot;, at least up to 100°C. These
authors [10] concluded that, if such a maximum does exist, it is at 100°C or
higher.
The discrepancies between our G „ data and those of RUterjans etP»2
al. [11] can be attributed to the uncertainties associated with the vapour
space corrections of the latter data and the extrapolation procedure, as
discussed above, for aqueous CsCl. The discrepancies with the C „ data
obtained from the integral heat method [10] can be attributed to the
combined uncertainties in the determination of AH at infinite dilution ands
those propagated during the differentiation of AH with respect tos
temperature. The AH , obtained from a linear extrapolation to infiniteS
dilution, has a range of values at a given temperature. To choose a single
value from this range of values, the purely empirical assumption was made
that the change in AH with concentration (corrected for Debye-Hückel
effects) is a smooth function of temperature.
In addition to the above data, there are three literature values—o
for C „ of Csl at room temperature [6,7,26,27]. Two of these (-126 and
-209 J'K^'raol""1) [26,27] are based upon adiabatic C „ measurements. Fur-—o '
thermcrej the C . values were derived by the extrapolation of the data for
concentrated Csl solutions using purely empirical equations. Thus, these
values possibly have large errors and do not provide a meaningful comparison
with our data. Our C , value of -144 J»K~1»mol~1 at 25°C agrees well with
a recent value of -145 J»K~1 «mol"* of Fortier et al. [6] and Desnoyers et
al. [7], who used a room-temperature version [12] of our flow microcalori-
meter system.
In the case of CsF and CsIO3, no literature data are available for
comparison.
- 15 -
5. CONCLUSIONS
We have obtained C „ values as a function of temperature for
cesium and iodine electrolyte-uater systems that are important in nuclear
fuel waste disposal and nuclear safety. The precision of the present
C 9(T,m) data is ±1 x 10"1* J'K'^g"1 as compared with ±0.01 J'K~1'g~l for
P> -othe literature data on aqueous cesium and iodine systems [10,11]. Our C „
values agree well with the existing data in the 25 to 70°C temperature
range. For the current nuclear fuel waste assessment, we recommend the useof the experimental C 9(T) function obtained in this study until standard
p,z
reference tables, incorporating these and other literature results and
recommended by an international body, become available. We have recently
applied the data in this report to determine precise equilibrium constants
of the iodine disproportionation reaction up to a temperature of 300cC.
ACKNOVfLEDGEMENT S
The data analysis was carried out at Lawrence Berkeley Laboratory,
Berkeley, California. I am indebted to Professor Kenneth S. Pitzer for his
invitation to carry out part of this work, and to Dr. Ramesh C. Phutela for
his help with the computations.
- 16 -
REFERENCES
1. R.S. Dixon and E.L.J. Rosinger (editors), "The Canadian Nuclear FuelWaste Management Program 1983 Annual Report," Atomic Energy of CanadaLimited Report, AECL-7811 (1984).
2. D.F. Torgerson, N.H. Sagert, D.W. Shoesmith and P. Taylor (editors),"Underlying Chemistry Research for the Nuclear Fuel Waste ManagementProgram," Atomic Energy of Canada Limited Report, AECL-7786 (1934).
3. H.E. Flotow, P.A.G. O'Hare and J. Boerio-Goates, "Heat Capacity from 5to 350 K and Thermodynamic Properties of Cesium Nitrate to 725 K,"J. Chem. Thermodynamics 13, 477-483 (1981).
4. K. Mehta, "Radionuclides VTiich Merit Detailed Attention", in "Chemistryand Geochemistry - Proceedings of the Thirteenth Information Meeting ofthe Nuclear Fuel Waste Management Program," Atomic Energy of CanadaLimited Technical Record, TR-201*, pp. 231-236 (1982).
5. D.D. Wagman, W.H. Evans, V.B. Parker, R.H. Schumm, I. Halow,S.M. Bailey, K.L. Churney and R.L. Nuttal, "NBS Tables of ChemicalThermodynamic Properties. Selected Values for Inorganic and C^ and C 2
Organic Substances in SI Units," U.S. National Bureau of Standards.Supplements to J. Phys. and Chem. Reference Data 11(2), 1-392 (1982).
6. J.-L. Fortier, P.-A. Leduc and J.E. Desnoyers, "Thermodynamic Proper-ties of Alkali Halides. II. Enthalpies of Dilution and Heat Capacitiesin Water at 25°C," J. Solution Chem. 323-349 (1974).
7. J.E. Desnoyers, C. de Visser, G. Perron and P. Picker, "Re-examinationof the Heat Capacities Obtained by Flow Microcalorimetry. Recommenda-tion for the Use of a Chemical Standard," J. Solution Chem. 5, 605-616(1976).
8. K.S. Pitzer, J.C. Peiper and R.H. Busey, "Thermodynamic Properties ofAqueous Sodium Chloride Solutions," J. Phys. Chem. Ref. Data 13, 1-102(1984).
9. A. Kumar, G. Atkinson and R.D. Howell, "Thermodynamics of ConcentratedElectrolyte Mixtures. II. Densities and Compressibilities of AqueousNaCl-CaCl2 at 25°C," J. Solution Chem. 11_, 857-870 (1982).
10. R.E. Mitchell and J.W. Cobble, "The Thermodynamic Properties of HighTemperature Aqueous Solutions. VII. The Standard Partial Molal HeatCapacities of Cesium Iodide from 0 to 100°C," J. Amer. Chem. Soc. 86_,5401-5403 (1964).
11. H. RUterjans, F. Schreiner, U. Sage and Th. Ackermann, "Apparent MolalHeat Capacities of Aqueous Solutions of Alkali Halides and AlkylAmmonium Salts," J. Phys. Chem. 73, 986-994 (1969).
- 17 -
12. P. Picker, P.-A. Leduc, P.R. Phillip and J.E. Desnoyers, "Heat Capacityof Solutions by Flow Microcalorimetry," J. Chem. Thermodynamics 3^631-642 (1971).
13. P.S.Z. Rogers and K.S. Pitzer, "High-Temperature Thermodynamic Proper-ties of Aqueous Sodium Sulfate Solutions," J. Phys. Chem. 85, 2886-2895(1981).
iA. D. Smith-Magowan and R.H. Wood, "Heat Capacity of Aqueous SodiumChloride from 320 to 600 K Measured with a New Flow Calorimeter," J.Chem. Thermodynamics I3j 1047-1073 (1981).
15. R.H. Busey, H.F. Holmes and R.E. Mesm«r, "The Enthalpy of Dilution ofAqueous Sodium Chloride to 673 K Using a New Heat-Flow and Liquid-FlowMicrocalorimeter. Excess Thersnodynamic Properties and Their PressureCoefficients," J. Chem. Thermodynamics _16_, 343-372 (1984).
16. K.S. Pitzer, "Thermodynamics of Electrolytes. I. Theoretical Basisand General Equations," J. Phys. Chem. ]]_, 268-277 (1973).
17. K.S. Pitzer and J.J. Kim, "Thermodynamics of Electrolytes. IV.Activity and Osmotic Coefficients for Mixed Electrolytes," J. Amer.Chem. Soc. 96^ 5701-5707 (1974).
18. C.E. Harvie and J.H. Weare, "The Prediction of Mineral Solubilities inNatural Waters: The Na-K-Mg-Ca-Cl-SO^-HjO System from Zero to HighConcentration at 25°C," Geochim. Cosmochim. Acta 44_, 981-997 (1980).
19. K..S. Pitzer, "Thermodynamics of Unsymmetrical Electrolyte Mixtures.Enthalpy and Heat Capacity," J. Phys. Chem. j37_, 2360-2364 (1983).
20. H.F. Holmes and R.E. Mesmer, "Thermodynamic Properties of AqueousSolutions of the Alkali Metal Chlorides to 250°C," J. Phys. Chem. &7_,1242-1255 (1983).
21. P.P.S. Saluja, "Thermodynamic Data for Selected Electrolytes atElevated Temperatures," J. Nucl. Mater. 130, 329-335 (1985). Alsoavailable as AECL-8447.
22. P.P.S. Saluja, "Environment of Ions in Aqueous Solutions - Chapter 1"j_n_MTP International Review of Science, Series Two, Volume 6 (Electro-chemistry), J.O'M. Bockris (ed.), Butterworths and Co. Ltd., Loadon,1976, pp. 1-51.
23. D.J. Bradley and K.S. Pitzer, "Thermodynamics of Electrolytes. 12.Dielectric Properties of Water and Debye-HUckel Parameters to 35O°Cand 1 kbar," J. Phys. Chem. 83, 1599-1603 (1979).
24. G.S. Kell, "Thermodynamic and Transport Properties of Fluid Water -Chapter 10," _in Water, A Comprehensive Treatise, Volume 1, The Physicsand Physical Chemistry of Water, F. Franks (ed.), Plenum Press, NewYork, 1973, pp. 363-416.
- 18 -
25. A.J.B. Crvickehank, Th. Ackermann and P.A. Gigue1re, "Heat Capacity ofLiquids and Solutions Near Room Temperature," in_Experimental Thermo-dynamics, Volume I (Calorimetry of Non-Reacting Systems),J.P. McCullough and D.W. Scott (eds.), Plenum Press, New York, 1968,pp. 421-535.
26. M. Eigen and E. Wicke, "Ionenhydratation und Spezifische WärmeWässringer Electrolyt-lösungen," Z. Elektrochem. Angew. Physik.Chem. 55_, 354-363 (1951).
27. A.F. Kapustinskii, I.I. Lipilina and O.Ya. Samoilov, "Cold Calorimeterfor the Thermochemlcal Study of Solutions with a Thermometric Sensi-tivity of 0.00005°: Determination of the Specific Heat of CesiumIodide Solutions with a Precision of 0.03%," J. Chim. Phys. 54, 343-347(1957).
* Unrestricted, unpublished report available from SDDO,Atomic Energy of Canada Limited Research Company,Chalk River, Ontario KOJ 1J0.
- 19 -
TABLE 1
APPARENT MOLAL HEAT CAPACITIES (<|>_) AND HEAT CAPACITIES (C _)_
OF CsF IN AQUEOUS SOLUTIONS AS A FUNCTION OF TEMPERATURE, AT 0.6 MPa
Temperature
(°C)
25
50
75
100
Molality, m
(mol «kg"1)
Water (0.0)*
0.031850.095850.209300.658021.06028
Water (0.0)
0.031850.095850.209300.658021.06028
Water (0.0)
0.031850.095850.209300.658021.06028
Water (0.0)
0.031850.095850.209300.658021.06028
CP,2
(J'K-^g-1)
(4.1779)
4.15364.10584.02443.73263.5136
(4.1790)
4.15584.11034.03253.75193.5395
(4.1911)
4.16824.12324.04623.76933.5593
(4.2143)
4.19114.14574.06823.79023.5791
i
(J'J-T^mol-1)
(75.26)
-132-128.5-122.1-109.75-92.82
(75.28)
-97.1-92.4-87.4-79.15-65.49
(75.50)
-85.8-82.1-77.7-68.46-55.22
(75.92)
-91.8-86.0-80.1-68.78-55.42
*Values in the parentheses are from Reference 24.
- 20 -
TABLE 2
APPARENT MOLAL HEAT CAPACITIES (<)> ) AND HEAT CAPACITIES (C „)c »- p,2-OF CsCl IN AQUEOUS SOLUTIONS AS A FUNCTION OF TEMPERATURE, AT 0.6 MPa
Temperature
(°C)
25
50
75
100
Molality, m
(mol«kg"1)
Water (0.0)*
0.040240.104830.205440.421460.591120.799971.02638
Water (0.0)
0.11140.21200.40680.59420.78851.0189
Water (0.0)
0.11140.21200.40680.59420.78851.0189
Water (0.0)
0.11140.21200.40680.59420.78851.0189
CP,2
(J.K-i.g-1)
(4.1779)
4.14414.09134.01223.85253.73733.60443.4711
(4.1790)
4.09104.01483.87643.75193.63133.4987
(4.1911)
4.10374.02893.89193.76893.64973.5178
(4.2143)
4.12654.05093.91333.79023.67043.5389
c
(J-K-i-mol-1)
(75.26)
-142.3-137.3-131.1-123.5-116.2-110.1-104.24
(75.28)
-101.2-98.6-91.2-87.1-83.24-78.64
(75.50)
-93.7-86.8-80.3-76.0-72.16-68.55
(75.92)
-93.4-88.7-81.1-75.61-71.84-67.06
* Values in parenthesis are from Reference 24.
- 21 -
TABLE 3
APPARENT MOLAL HEAT CAPACITIES ($_) AND HEAT CAPACITIES (C }_
OF Csl IN AQUEOUS SOLUTIONS AS A FUNCTION OF TEMPERATURE, AT 0.6 MPa
Temperature
(°C)
25
50
75
100
Molality, m
(mol»kg"1)
Water (0.0)*
0.098430.19640.41530.61150.79690.9907
Water (0.0)
0.098430.18300.39140.41530.56840.79690.9907
Water (0.0)
0.094340.18300.39140.56840.82421.0090
Water (0.0)
0.098430.19640.41530.61150.79690.9907
CP,2
(J.K-Lg-1)
(4.1779)
4.06123.95223.72913.55143.39803.2522
(4.1790)
4.06593.97413.76473.74293.60503.41713.2724
(4.1911)
4.08373.98763.78003.62113.41353.2780
(4.2143)
4.10163.99583.77843.60373.45343.3089
c
(J'K-i-mol-1)
(75.26)
-130.5-122.4-111.8-101.84-95.83-89.44
(75.28)
-92.7-87.2-80.4-77.64-73.24-68.28-64.91
(75.50)
-77.5-76.0-68.25-62.02-56,60-53.30
(75.92)
-79.3-74.4-67.94-62.25-57.60-54.21
* Values in parenthesis are from Reference 24.
- 22 -
TABLE 4
APPARENT MOLAL HEAT CAPACITIES (<t> ) AND HEAT CAPACITIES (C 2
OF CsIO3 IN AQUEOUS SOLUTIONS AS A FUNCTION OF TEMPERATURE, AT 0.6 MPa
Temperature
(°C)
25
50
75
100
Molality, m
(mol«kg"1)
Water (0.0)*
0.005010.009870.025550.044910.054690.06850
Water (0.0)
0.009870.019940.043760.06556
Water (0.0)
0.009870.019940.043760.06556
Water (0.0)
0.005170.011650.033290.039760.051160.05924
CP,2
(J-K-i'g-1)
(4.1779)
4.17094.16424.14314.11734.10444.0864
(4.1790)
4.16604.15314.12294.09585
(4.1911)
4.17834.165754.13614.10945
(4.2143)
4.20764.19944.17204.164054.150054.14025
c
(J'K-iinol-1)
(75.26)
-113.4-106.3-90.7-82.0-80.6-77.9
(75.28)
-34.8-20.53-11.76-8.34
(75.50)
-5.66+8.42+17.44+18.73
(75.92)
-0.8+13.7+15.0+17.9+21.6+23.6
* Values in parenthesis are from Reference 24c
- 23 -
TABLE 5
PARTIAL MOLAR HEAT CAPACITIES (C ,,) OF SELECTED CESIUM ELECTROLYTES
IN WATER AS A FUNCTION OF TEMPERATURE, AT OJ} MPa
Electrolyte-HpOSystem
CsF-1120
CsCl-H20
CsI-H2O
CsIO3-H2O
25°C
-138*-140**
-147-149
-143-144
-111-113
f ° c T .*p,2
50°C
-105-105
-114-116
-104-105
-37-40
75°C
-96-95
-107-106
-93-92
-8-5
100°C
-104-104
-112-113
-97-98
+2-1
*Top entry is the experimental C° ̂ value, obtained by extrapolation of the4>c(m) data to infinite dilution, using Equation (8).
**Bottom entry is the calculated C° 2 value, using the temperature-dependentEquation (10) for C° 2(T) and the constants Ac, Bc and Cc from Table 6. Theagreement between the top and the bottom entries is good, within the experi-mental uncertainty of ±2 J»K~^»mol~^- Since Equation (10) is based on alarge data set and has a convenient form for application, it is recommendedfor obtaining £° ̂ values as a function of temperature and for other thermo-dynamic calculations.
- 24 -
TABLE 6
COEFFICIENTS A , B AND C FOR THE TEMPERATURE-DEPENDENT
C°PARTIAL MOLAR HEAT CAPACITY FUNCTION, C° „(T), AT 0.6 MPa
Electrolyte-H20System
CsF-H20
CsCl-H20
CsI-H20
CsIO3-H2O
Ac
(xlO~5)
-6.38443%0.421
-5.95739%O.2O0
-6.63000%0.311
-9.69236%0.596
Bc
(xlO~3)
3.57106%0.255
3.30377%0.121
3.67497%0.188
5.29662%0.362
Cc
-5.26322%0.382
-4.87752%0.183
-5.34924%0.284
-7.23131%0.547
TABLE 7
COEFFICIENTS A , B AND C FOR THE TEMPERATURE-DEPENDENT
ION-INTERACTION PARAMETER, ß5„°)J(T), AT 0.6 MPa1 rlX
Electrolyte-H20System
CsF-H2O
CsCl-H2O
CsI-H20
CsIO3-H2O
- 1
-1
- 1
- 3
Aß
(xlO"1)
•13618±0.43
.3885510.18
.95746+0.28
.31827±1.03
6.
7 .
10
8 .
Bß
(xlO-1*
25271±2
6772611
.6820+1
26813±3
)
.41
.08
.73
.06
- 8 .
-10
-14
0 .
Cß
( xlO 7)
63628+3
.643011
.589012
0
.59
.60
.57
- 25 -
TABLE 8
CALCULATED MAXIMUM PARTIAL MOLAR HEAT CAPACITY, C° (MAX),
AND TEMPERATURE, T , , OF MAXIMUM C° „—- -* max, he-* — p,2
Electrolyte-H 2 0System
CsF-H2O
CsCl-H2O
CsI-U2O
CsIO3-H2O
Tmax,he
(K)
348 ±2
350 ±1
353+2
367+2
C° (MAX)P>2
(J 'K-1 'mol-1)
-94+0.2
-101+0.1
-97+0.1
6.3+0.1
TABLE 9
COEFFICIENTS A , B AND C FOR TEMPERATURE-DEPENDENTv» v v
PARTIAL MOLAR VOLUMES, V^(T), AT 0.6 MPa
Electrolyte-li20System
CsF-ll2O
CsCl-ll.2O
CsI-H2O
CsIO3-H2O
AV
-3.24548+0.599
-2.83486±0.169
-2.93638±0.128
-O.83943±O.O4O
BV
( x 10 ~ )
2.2098910.364
2.10442 ±0.102
2.22208 ±0.078
0.75071±0.030
CV
-0.30874+0.055
-0.2553±0.015
-0.22268 ±0.012
0.0
- 26 -
TABLE 10
PRESSURE COEFFICIENTS OF C „ AT 25"C EVALUATEDZ£—V,2
FROM PARTIAL MOLAR VOLUMES, Vn(T), AS A FUNCTION OF TEMPERATURE
Electrolyte-H20System
CsF-H2O
CsCl-H2O
CsI-H2O
CsIO3-H2O
â / ^ a t 2 5°C
(J *K~ 1#mol~ ^MPa" L)
0.73
0.64
0.66
0.17
TABLE 11
C°PARTIAL MOLAR HEAT CAPACITIES, C° „, AT 25°Cp,tJ
COMPARED WITH LITERATURE DATA [6,7]
Elect rolyte-H20
CsF-H2O
CsCl-H2O
CsI-H2O
C° - at 0.6 MPap,2
(PRESENT WORK)
-139±3
-149+2
-144+2
C° ,. a t 0 .1 MPaP»2
(LITERATURE DATA)
-136+2
-149 ±2
-145 ±2
C° „ a t 0 .6 MPap,2
(EXTRAPOLATEDLITERATURE DATA)
-135+3
-149 ±2
-145+2
- 27 -
- 8 0
- 9 0
- 170
-180
-190
-200
\ i i i i r
Mitchell and Cobble [ 10]
Mitchell and Cobble t 10](Smoothed Values)
Rill erJons el ol. [ 11 ]
_L _L _L _LO 20 40 60 80 100 120 140
TEMPERATURE (°C)
FIGURE 1: Comparison of Existing Data [10,11] for the Partial MolarHeat Capacities, C° 2, as a Function of Temperature forAqueous Csl '
- 28 -
Cylinder
Teflon-LinedSomple Cylinders
Test Solution !_»
PickerCirculating
Bath
-Insulated Fluid/Transfer Lindes •; PressureGauge
Waste
Bleed, Vibrating TubePort --y.--
Flow Densimeter !Id
MicrocalorimeterBypass — :
DensimeterDisplay Unit
FlowRegulator
\Thermostated FlowMicrocalorimeter
X-Y Recorder
FIGURE 2: High-Temperature Heat-Capacity Flow Microcalorimeter System
w/w J
s/w
^ s/s
w/s
w/w
TIME
FIGURE 3: Calorimeter Output versus Time During a VolumetricHeat Capacity Measurement
- 29 -
0W
-80
-90
-100
-110
-120
-130
"" -140
'°o<1 -15 G
-160
-170
-180
-190
\\
I
IT
M
CsCsCs
1FCI
\ \ x
\ \
\
1
20 40 60 80 100 120 140 160 180
TEMPERATURE (°C)
FIGURE 4: Temperature Dependence of the Partial Molar Heat Capacities,Cp 2(
T)» of Cesium Halides in Water at 0.6 MPa. • , A and • areisothermal 5° 2 values from present work
- 30 -
100
10
A Literature [ 6 , 7 ]• Present work
CsF 25 °C, 0.6MPa
0 O.I 0.2 0.3 0.4 0.5 0.6 0.7 0.8Molaliiy (mol. kg" )
FIGURE 5: Apparent Molal Heat Capacity, $ , as a Function of Molality, n,for Aqueous Cesium Hallde Solutions: A Comparison with the 25°CLiterature Data of Fortier et a l . [6] and Desnoyers et a l . [7]
FIGURE 6: Temperature Dependence of the Partial Molar Heat Capacities,Cp 2(T), of CsCl in Water from Equation (10): A Comparison withthe Applicable Literature Data [11]. Circles tare isothermalC° 2 values from present work.
- 32 -
-50
-100
oE
I O
-150
-200
Eigen and Wicke [ 26 ]
Kapusiinskii et al. [ 27 ]
Mitchell and Cobble [ 10 ]
Rufer jans et al. C i l ]
Desnoyers et a l . [ 7 ]
Present Work
•TA
0 20 40 60 80 100 120
TEMPERATURE (°C)
140 160
FIGURE 7: Temperature Dependence of the Partial Molar Heat Capacities,C° 2(T), of Csl in Water from Equation (10): A Comparison withthé Applicable Literature Data [6,7,10,11,25,26]. Circles,»,are isothermal C° j values from present work .
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