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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