The standard molar enthalpy of formation of LnPO4(s) (LnÂ =Â La, Nd, Sm) by solution calorimetry

The standard molar enthalpy of formation of LnPO4(s)(Ln 5 La, Nd, Sm) by solution calorimetry

Deepak Rawat • Smruti Dash

Indian Special Chapter

� Akademiai Kiado, Budapest, Hungary 2012

Abstract The standard molar enthalpy of formation of

LaPO4(s), NdPO4(s), and SmPO4(s) has been determined

using an isoperibol solution calorimeter. The solution

calorimeter vessel was held at 298.15 K. The precipitation

reaction between aqueous solution of rare-earth chloride

(LnCl3(aq.)) and ammoniacal solution of ammonium

dihydrogen phosphate (NH4H2PO4(aq.)) was studied. The

temperature of the calorimeter vessel was measured before,

during, and after the reaction. The enthalpy change due to

precipitation of LaPO4(s), NdPO4(s), or SmPO4(s) from

required solutions was measured at 298.15 K. Using these

values and other auxiliary data from the literature, ther-

mochemical reaction scheme were devised to calculate the

standard enthalpy of formation of each phosphate com-

pound i.e., LaPO4(s), NdPO4(s), and SmPO4(s). The cal-

culated values for LaPO4(s), NdPO4(s), and SmPO4(s) at

298.15 K were found to be -1947.5 ± 3.2, -1938.3 ±

3.6, and -1942.9 ± 3.4 kJ mol-1, respectively.

Keywords Standard molar enthalpy of formation �Isoperibol solution calorimeter � Monazite �Thermochemical cycles � Rare-earth phosphates

Introduction

Safe storage of radioactive waste is a major challenge for the

nuclear industry. At present, nuclear waste are immobilized

in glass matrix. The main weakness of a glassy waste form is

that fundamentally glass is a meta-stable material which

tends to re-crystallize, although this process is very slow.

Crystalline waste forms (ceramics) are expected to be several

orders of magnitude more durable than glasses [1]. Miner-

alogy is a good basis for designing ceramics as it provides

reliable data for the long-term behavior of ceramic matrix in

natural environment. Natural radioactive minerals are the

only source of information for long-term accumulation of

radiation damage in crystals. The zircon, zirconolite, hol-

landite, synroc, brannerite, apatite, and monazite are the

potential minerals for storage of nuclear waste. The present

study is focused on monazite which is the natural light rare-

earth orthophosphate (AXO4, P21/n), where A site is occu-

pied by large cations, nine fold coordinated, such as trivalent

rare-earth elements (RE) or other cations similar in size

(Ca2?, U4?, Th4?) while the X-site contains small, tetrahe-

drally coordinated cations such as P5?. The REO9 coordi-

nation polyhedra can be explained by a combination of a

pentagon and a tetrahedron unit. The REO9 unit shares its

two opposite edges with PO4 unit to form an infinite alter-

nating chain along c-axis [2, 3].

In addition to RE3?, both trivalent (Am3?, Cm3?, Pu3?)

and tetravalent (Np4?, U4?, Th4?) actinides can be incor-

porated in monazite by substitution with RE3? for trivalent

ions and by coupled substitution An4? ? Ca2? = 2RE3?

(An = actinide) for tetravalent ions [4]. The evaluation of

the long-term durability and reactivity of rare-earth materials

require accurate determination of their thermodynamic

properties. Thermodynamic properties should be measured

in different methods to select the true value. The solution

calorimetry offers an effective methodology for the deter-

mination of enthalpies of formation. In this study, enthalpy

of formation of LaPO4(s), NdPO4(s), and SmPO4(s) have

been measured using semi-adiabatic solution calorimeter.

D. Rawat � S. Dash (&)

Product Development Division, Bhabha Atomic Research

Centre, Mumbai 400085, India

e-mail: [email protected]

123

J Therm Anal Calorim

DOI 10.1007/s10973-012-2742-3

Experimental

Chemicals

KCl, NBS-1655 from National Bureau of Standard,

Washington DC, USA was taken for the chemical calibration

of the solution calorimeter. LaCl3�7H2O(s) (purity: 0.999

mass fraction, LEICO Industries, USA), NdCl3�6H2O(s)

(purity: 0.999 mass fraction, Sigma-Aldrich, USA),

SmCl3�6H2O(s) (purity: 0.9999 mass fraction, Sigma-

Aldrich, USA), and NH4H2PO4(s) (purity 0.995 mass frac-

tion, E. Merck, India) were used for the solution calorimetric

work.

Preparation and characterization of LaPO4(s),

NdPO4(s), and SmPO4(s)

The lanthanide phosphates were prepared by solution

chemistry route. In this method, aqueous solution of rare-

earth chloride (REECl3(aq.)) and ammoniacal solution of

ammonium dihydrogen phosphate (NH4H2PO4(aq.)) were

taken in 1:1 mole ratio. Addition of ammoniacal solution of

NH4H2PO4(aq.) to separate aqueous solutions of LaCl3(aq.),

NdCl3(aq.), and SmCl3(aq.) with continuous stirring resulted

in precipitates of respective rare-earth phosphates. The

resulting precipitates were separately filtered and washed in

distilled water until the filtrate was neutral. The precipitates

thus obtained were dried using infra-red lamp, followed by

furnace heating at 1,000 K for 8 h in air. The product was

analyzed by XRD, which was recorded in a STOE, Germany,

X-ray diffractometer using the monochromatic CuKa radi-

ation and a nickel filter. Patterns were recorded in the 2hrange from 10� to 70� at 298 K. XRD patterns of LaPO4(s),

NdPO4(s), and SmPO4(s) were well matched with the

reported pattern of LaPO4(s) (JCPDS file no: 83-651),

NdPO4(s) (JCPDS file no: 83-654), and SmPO4(s) (JCPDS

file no: 83-655) [5], respectively. Accurate mass of the

sample was a prerequisite for calorimetric work. Rare-earth

chlorides and ammonium dihydrogen phosphate contain

water of crystallization. In order to determine the water of

crystallization in lanthanum chloride, neodymium chloride,

samarium chloride, and ammonium dihydrogen phosphate,

thermal gravimetric analysis (TG) of these compounds were

carried out with Netzsch STA 409 PC, NETZSCH-Gerate-

bau GmbH, Germany. Prior to the studies, the instrument

was calibrated for temperature using the melting points of

In, Sn, Bi, Zn, Al, Ag, Au, and Ni and mass calibration was

done using CaC2O4�2H2O(s).

TG plots of lanthanum chloride, neodymium chloride,

samarium chloride, and ammonium dihydrogen phosphate

were recorded at the heating rate of 2 K min-1 under

flowing argon (20 ml min-1) in the temperature range

from 303 to 980 K (Figs. 1–4). Figure 1 shows loss of

seven water molecules in case of lanthanum chloride.

However, Figs. 2 and 3 show loss of six water molecules

for neodymium chloride and samarium chloride. Figure 4

does not show loss of any water molecule for

NH4H2PO4(s) but continuous mass loss due to evolution of

NH3(g) near 473 K was observed.

Instrumentation

The heat of solution experiments were performed in ther-

mometric precision solution calorimeter model 2225,

Sweden. The outer part of the calorimetric system was a

stainless steel cylinder mounted in the TAM-III bath,

which has an oil bath with stability ±1 lK. The actual

calorimeter consisted of thin-walled Pyrex-glass reaction

vessel fitted with a thermistor for sensing temperature and

300 350 400 450 500 550–12

–10

–8

–6

–4

–2

0

Temperature/K

Mas

s lo

ss/m

g

LaCl3.7H2O(s)

Fig. 1 Thermal behavior of LaCl3�7H2O(s)

300 350 400 450 475 500

Temperature/K425375325

–30

–25

–20

–15

–10

–5

0M

ass

loss

/mg

NdCl3.6H2O(s)

Fig. 2 Thermal behavior of NdCl3�6H2O(s)

D. Rawat, S. Dash

123

a heater for calibration and equilibration. The resistance of

thermistor was measured with a Wheatstone bridge, which

was contained in the cylindrical backbone. The calori-

metric vessel was attached to the cylindrical backbone

which also contained motor for stirrer, which could hold

ampoule containing sample. The calorimeter was provided

with SolCal program to control experimental data acqui-

sition, graphical data presentation, and data analysis. The

temperature offset from the bath was measured accurately

as a function of time by the SolCal program. The tem-

perature change of the calibration and sample runs were

chosen nearly as identical as possible.

The heat energy for electrical calibration during the

experiment was supplied to the reaction vessel with the

help of heater. A known precision resistance (20 X) was

mounted in series with the heater. By measuring the volt-

age across this precision resistance, the actual value of

current passing through the heater can be determined. The

voltage across the heater was measured close to the calo-

rimetric reaction vessel to minimize the power dissipation

due to the internal resistance of the electrical circuit. These

two values, the actual current and the voltage were multi-

plied to give the electrical power that was dissipated as

heat inside the vessel. The sample containing ampoule was

equilibrated in the calorimeter containing solvent. The

calorimetric temperature was recorded as a function of time

before the ampoule was broken. The reaction was initiated

by breaking the ampoule containing the sample. The

reaction could be observed as an instantaneous change in

temperature, the system was calibrated electrically before

and after each reaction. The smallest heat exchange with

the environment during the reaction and the heat arising

from stirring were mathematically adjusted using the

information obtained from the baseline temperature chan-

ges before and after the reaction.

Ampoule break experiment

The enthalpy of reaction was measured in ampoule break

experiment in which known quantity of solute (either

KCl(s) or LnCl3�xH2O(s)) was loaded into the calori-

metric ampoule. Calorimetric vessel contained 25 g of

double distilled water in case of KCl(s) as solute and

0.02 mol dm-3 ammoniacal solution of NH4H2PO4(aq.) in

case of LnCl3�xH2O(s) as solute. The ammoniacal solution

was prepared in 25 g of double distilled water. The mole

ratio of solute to solvent present in calorimetric vessel was

more than 1:2,500. The temperature of the calorimeter was

recorded as a function of time before the reaction took

place. The reaction was initiated by breaking the ampoule

containing solute and an instantaneous change in temper-

ature was measured. The measured temperature change for

KCl(s) was due to enthalpy of solution at infinite dilution

and that for rare-earth chlorides were due to both enthalpy

of solution at infinite dilution and enthalpy of precipitation

of LnPO4(s) (Ln = La, Nd, Sm). The enthalpy of reaction

was calibrated electrically before and after each reaction.

Enthalpy of formation

When a reaction occurs in the calorimetric vessel, the

enthalpy of reaction can be calculated from the following

expression:

�Q ¼ C � DT ;DrH�m 298:15 Kð Þ ¼ Qreac=n soluteð Þ

¼ C � DTcorr=n soluteð Þ ð1Þ

where, Q = heat evolved or absorbed, C = water equiva-

lent of the calorimetric vessel, DT = temperature change

during the process, DrHm� (298.15 K) = standard molar

enthalpy of reaction, n(solute) = number of moles of

solute, and DTcorr = corrected temperature. The water

300 350 400 450

Temperature/K

425375325

–2

0

–4

–6

–8

–10

–12

–14

Mas

s lo

ss/m

g

SmCl3.6H2O(s)

Fig. 3 Thermal behavior of SmCl3�6H2O(s)

300 400 500 600 700 800 900 1000

Temperature/K

Mas

s lo

ss/m

g

–30

–35

–25

–20

–15

–10

–5

0NH4H2PO4(s)

Fig. 4 Thermal behavior of NH4H2PO4(s)

Standard molar enthalpy of formation of LnPO4(s)

123

equivalent of calorimeter was determined by passing

known amount of current for a fixed period through the

electrical heater submerged in the solution. The rise in

temperature of the calorimeter was determined accurately.

Using the values of the temperature rise (DT) and the

amount of heat input the water equivalent of the calorim-

eter was determined. For each experiment, the water

equivalent of the calorimeter was determined before and

after the ampoule break experiment and the average value

of the water equivalent was used for the calculations.

Results

The calibration of the solution calorimeter was carried out

with KCl(s) and its enthalpy of solution was measured for

the following reaction:

KCl sð Þ þ 1H2O lð Þf g ¼ KCl aq:ð Þ ð2Þ

The experimental values of the enthalpy of dissolution

were used to calculate the enthalpy of solution at infinite

dilution. From the present study DsolHm? (298.15 K) of KCl

was calculated to be 17.23 ± 0.52 kJ mol-1. The detail of

the calibration experiment and a typical calibration plot

were given in Table 1 and Fig. 5, respectively. After the

calibration experiment, the ampoule break experiment for

LaCl3�7H2O(s), NdCl3�6H2O(s), and SmCl3�6H2O(s) were

successively performed in the calorimeter at 298.15 K. The

results of the enthalpy of reaction for these rare-earth

chlorides are given in Table 2. In each ampoule break

experiment the solvent was diluted ammoniacal solution of

ammonium dihydrogen phosphate. The thermochemical

reaction scheme to derive standard molar enthalpy of

formation of LaPO4(s), NdPO4(s), and SmPO4(s) from

precipitation reaction at T = 298.15 K are given in

Tables 3, 4, and 5, respectively. In all the experiments

the mass of solvent, double distilled water was taken as

25 g. A strict control of the stoichiometries in each step of

the calorimetric cycle was made, with an objective that the

dissolution of the reactants gave the same composition as

those of the products.

The enthalpy of solution values of DH1 (Table 3), DH11

(Table 4), and DH21 (Table 5) for the reaction

LnCl3 sð Þ þ 1H2O lð Þf g ¼ LnCl3 aq:ð Þ ð3Þ

can be calculated from the addition of reactions (4) and (5):

LnCl3�xH2O sð Þ þ f1H2O lð Þg ¼ LnCl3 aq:ð Þ þ xH2O lð Þð4Þ

LnCl3 sð Þ þ xH2O lð Þ ¼ LnCl3�xH2O sð Þ ð5Þ

where, Ln = La, Nd, or Sm, x = 7 for LaCl3(s), x = 6

for NdCl3(s) and SmCl3(s). The concentration of the

NH4H2PO4(s) was in excess compared to the amount of rare-

earth chlorides used in each precipitation reaction; it is

assumed that the whole of the LnCl3(s) reacted to form the

respective LaPO4(s), NdPO4(s), and SmPO4(s) precipitate.

The enthalpy of reaction: 2HCl(aq.) = 2HCl(g) ? {?H2O

(l)} (DH9, DH19, and DH29) reported in the literature has been

taken for the calculation of enthalpy of formation of LnPO4

(Ln = La or Nd or Sm) using thermochemical cycle. In this

study, it has been assumed that the enthalpy of reaction:

Table 1 Data on ampoule break experiment of KCl(s) in water

Compound Wt. of sample/mg Tinitial Off./mK DTcorr/mK C/J K-1 Q/J DsolHm� (298.15 K)/kJ mol-1 Mean DsolHm

� /kJ mol-1

KCl(s) 14.6 165.213 -28.822 116.612 3.361 17.16 17.23 ± 0.52a

14.7 142.625 -30.598 116.773 3.573 18.12

17.0 146.841 -33.101 117.792 3.899 17.10

15.1 153.452 -29.126 116.563 3.395 16.80

16.5 159.756 -31.843 117.857 3.753 16.96

Data given are: Q reaction (J), initial temperature offset Tinitial Off. (mK), Tcorr (mK), and C (average total heat capacity of the calorimetric

vessel; J K-1)a The molar mass of KCl(s) = 74.55 g mol-1

0 200 400 600 800 1000 1200

Time/sec

Tem

pera

ture

/mK

120

130

140

150

160

170

180

190

calibration-1

calibration-2

Breaking of ampoule

Fig. 5 Calibration plot with KCl(s)

D. Rawat, S. Dash

123

Table 2 Data on ampoule break experiments of {LaCl3�7H2O(s), NdCl3�6H2O(s), and SmCl3�6H2O(s) ? Soln A} (Soln A = ammoniacal

NH4H2PO4(s) ? ?H2O(aq.))

Reactants Wt. of

sample/mg

TinitialOff./mK DTcorr/mK C/JK-1 Q/J DrHm� (298.15 K)/

kJ mol-1Mean DsolHm

� /

kJ mol-1

LaCl3�7H2O(s) ? Soln A 19.2 92.154 32.109 116.351 -3.736 -72.2 -74.7 ± 2.2a,b

19.8 83.908 34.768 117.062 -4.070 -76.2

18.0 63.813 31.470 116.745 -3.674 -75.8

NdCl3�6H2O(s) ? Soln A 20.4 77.669 47.428 116.917 -5.545 -97.5 -94.8 ± 2.4a,b

20.0 38.431 44.565 117.752 -5.248 -94.1

19.5 25.007 42.476 118.807 -5.046 -92.8

SmCl3�6H2O(s) ? Soln A 18.0 31.283 49.241 117.345 -5.778 -117.0 -115.3 ± 2.1a,b

14.9 38.152 39.157 117.845 -4.614 -112.9

17.5 33.145 47.355 117.549 -5.566 -116.0

a Uncertainty is twice the standard deviationb Molar enthalpy of solution based on the molar mass of 371.372, 358.691, and 364.501 g mol-1 for LaCl3�7H2O(s), NdCl3�6H2O(s), and

SmCl3�6H2O(s), respectively

Table 3 Reaction scheme for the standard molar enthalpy of formation of LaPO4(s) at 298.15 K

Reactions Enthalpy DrHm/kJ mol-1

LaCl3(s) ? {?H2O(l)} = LaCl3(aq.) DH1

LaCl3(aq.) ? NH4H2PO4(aq.) = LaPO4(s) ? NH4Cl(aq.) ? 2HCl(aq.) DH1 ? DH2 -74.7 ± 2.2

La(s) ? 3/2Cl2(g) = LaCl3(s) DH3 -1071.1 ± 1.5a

1/2N2(g) ? 3H2(g) ? P(s) ? 2O2(g) = NH4H2PO4(s) DH4 -1452.5 ± 1.2b

NH4Cl(s) = 1/2N2(g) ? 2H2(g) ? 1/2Cl2(g) DH5 314.9 ± 0.4b

2HCl(g) = H2(g) ? Cl2(g) DH6 184.6 ± 0.01b

NH4Cl(aq.) = NH4Cl(s) ? {?H2O(l)} DH7 -14.5 ± 0.01b

NH4H2PO4(s) ? {?H2O(l)} = NH4H2PO4(aq.) DH8 16.3 ± 1.2b

2HCl(aq.) = 2HCl(g) ? {?H2O(l)} DH9 149.5 ± 0.01b

La(s) ? P(s) ? 2O2(g) = LaPO4(s) DH10 -1947.5 ± 3.2

DH10 = DH1 ? DH2 ? DH3 ? DH4 ? DH5 ? DH6 ? DH7 ? DH8 ? DH9

a Ref. [21]b Ref. [22]

Table 4 Reaction scheme for the standard molar enthalpy of formation of NdPO4(s) at 298.15 K


NdCl3(s) ? {?H2O(l)} = NdCl3(aq.) DH11

NdCl3(aq.) ? NH4H2PO4(aq.) = NdPO4(s) ? NH4Cl(aq.) ? 2HCl(aq.) DH11 ? DH12 -94.8 ± 2.4

Nd(s) ? 3/2Cl2(g) = NdCl3(s) DH13 -1041.8 ± 2.0a



2HCl(g) = H2(g) ? Cl2(g) DH16 184.6 ± 0.01b




Nd(s) ? P(s) ? 2O2(g) = NdPO4(s) DH20 -1938.3 ± 3.6


a Ref. [21]b Ref. [22]


123

{2HCl(aq.) = 2HCl(g) ? {?H2O(aq.)} is same as that of

2HCl(aq.) = 2HCl(g) ? {?H2O(l)}; because, in actual

experiment {?H2O(aq.)} is a solution of {?H2O(l) ?

NH4?(aq.) ? Cl-(aq.)}. Since the mole fraction of NH4

? and

Cl- ions are infinitely small compared to that of H2O(l)

in H2O(aq.), the error involved due to this assumption is

negligible. Applying Hess’s law, the enthalpy of formation

was calculated for LaPO4(s) {-1947.5 ± 3.2} kJ mol-1,

NdPO4(s) {-1938.3 ± 3.6} kJ mol-1, and SmPO4(s)

{-1942.9 ± 3.4} kJ mol-1 using enthalpy data from Tables 3, 4,

and 5, respectively.

Discussion

The experimental value of enthalpy of solution of KCl(s) at

infinite dilution measured in this study (17.23 ± 0.52 kJ mol-1)

agrees well with the value 17.21 ± 0.01 kJ mol-1 reported

by Venugopal et al. [6] and that of N.B.S. value (17.241 ±

0.018 kJ mol-1) [7]. The results obtained for dissolution of

KCl(s) established the reliability of calorimetric measurements.

The enthalpy of formation values of LaPO4(s),

NdPO4(s), and SmPO4(s) measured in this study are com-

pared with that reported in the literature [8–19] and tabu-

lated in Table 6.

The vaporization study of LaPO4(s) was carried out by

Rat’kovskii et al. [8] using mass spectrometry combined

with Knudsen effusion method. Henderson et al. [9] reported

enthalpy of formation of LaPO4(s) measured by Marinova

et al. [10] using LaCl3 ? H3PO4 reaction calorimetry.

Ousoubalyev et al. [11] also reported enthalpy of formation

of LaPO4(s), NdPO4(s), and SmPO4(s) by calorimetry.

Ushakov et al. [12] and Helean et al. [13] reported enthalpy

of formation of LaPO4(s), NdPO4(s), and SmPO4(s) using

Table 5 Reaction scheme for the standard molar enthalpy of formation of SmPO4(s) at 298.15 K


SmCl3(s) ? {?H2O(l)} = SmCl3(aq.) DH21

SmCl3(aq.) ? NH4H2PO4(aq.) = SmPO4(s) ? NH4Cl(aq.) ? 2HCl(aq.) DH21 ? DH22 -115.3 ± 2.1

Sm(s) ? 3/2Cl2(g) = SmCl3(s) DH23 -1025.9 ± 2.0a



2HCl(g) = H2(g) ? Cl2(g) DH26 184.6 ± 0.01b




Sm(s) ? P(s) ? 2O2(g) = SmPO4(s) DH30 -1942.9 ± 3.4


a Ref. [21]b Ref. [22]

Table 6 A comparison of the standard molar enthalpies of formation values of LaPO4(s), NdPO4(s), and SmPO4(s) with that reported in the

literature

Authors Techniques DfHm� LaPO4(s)/

kJ mol-1DfHm

� NdPO4(s)/

kJ mol-1DfHm

� SmPO4(s)/

kJ mol-1

Rat’kovskii et al. [8] Mass spectrometry with Knudsen effusion method -1913 ± 10 – –

Helean et al. [13] Oxide melt solution calorimetry at 975 K -1970.7 ± 1.8 -1968.4 ± 2.3 -1965.7 ± 2.4

Popa et. al. [19] Recalculated from experiment results of [11] -1969.7 ± 1.9 -1967.9 ± 2.5 -1965.8 ± 2.9

Poitrasson et al. [16] From the temperature dependence of solubility – -1928 –

Wood and Solubility – -1932 –

Williams-Jones [17]

Williamson et al. [18] Estimated data -1955.2 ± 2.1 -1960.4 ± 2 –

Cetiner et al. [15] From the temperature dependence of equilibrium constant -1947 -1907 ± 4 -1962

Marinova and Yaglov [14] Solubility products and estimated entropies of LnPO4 -1962 ± 8 -1941 -1933

Marinova et al. [10] Acid solution calorimetry -1949 ± 4 – –

Ousoubalyev et al. [11] Calorimetry -1942 -1930 -1925

This study Solution calorimetry at 298.15 K -1947.5 ± 3.2 -1938.3 ± 3.6 -1942.9 ± 3.4

D. Rawat, S. Dash

123

oxide melt solution calorimetry at 975 K. Calorimetric

measurements were performed in a Calvet-type twin

microcalorimeter in sodium molybdate (3Na2O�4MoO3)

solvents at 975 K. Marinova and Yaglov [14] derived DfHm�

(298.15 K) of LaPO4(s), NdPO4(s), and SmPO4(s) from

solubility products and estimated entropies of LnPO4(s)

(Ln = La, Nd, Sm). Cetiner et al. [15] calculated enthalpies

of the reaction (6):

LnPO4 sð Þ þ 3Hþ ¼ Ln3þ þ H3PO4 aq:ð Þ ð6Þ

from the temperature dependence of the solubility of

LaPO4(s), NdPO4(s), and SmPO4(s) using Van’t Hoff

relation. Helean et al. [13] also reported the DfHm�

(LnPO4,s,298.15 K) calculated from the acid solubility

data of Marinova and Yaglov [14]. Poitrasson et al. [16]

derived DfHm� (NdPO4,s,298.15 K) from the temperature-

dependant solubility data, and Wood and Williams-Jones

[17] also derived this value from solubility. Estimated

enthalpy of formation value at 298.15 K for LaPO4(s) and

NdPO4(s) were reported by Williamson et al. [18].

Enthalpy of formation value of LaPO4(s) (Ln = La, Nd,

Sm) were also recalculated by Popa et al. [19] using

experimental results of Helean et al. [13].

The DfHm� (LaPO4,s,298.15 K) measured in this study is

matching excellently with that measured by Marinova et al.

[10] using LaCl3 ? H3PO4 reaction calorimetry and that

calculated from the temperature dependence of equilibrium

constant by Cetiner et al. [15]. The estimated data of

Williamson et al. [18] also showed reasonable agreement.

However, DfHm� (LaPO4,s,298.15 K) value by Rat’kovskii

et al. [8] from mass spectrometry is *30 kJ mol-1 higher

and that by Ushakov et al. [12] and Helean et al. [13] from

oxide melt solution calorimetry is *20 kJ mol-1 lower

than that measured in the present study.

The measured enthalpy of formation value of NdPO4(s) is

matching excellently with that of Marinova and Yaglov [14]

derived from the solubility products. It also agrees reason-

ably well with that measured by Ousoubalyev et al. [11] by

calorimetry and that by Poitrasson et al. [16] and Wood and

Williams-Jones [17] from solubility measurements. How-

ever, DfHm� (NdPO4,s,298.15 K) measured in this study is

higher by *25 kJ mol-1 than that reported by Helean et al.

[13] from oxide melt solution calorimetry and that calculated

by Williamson et al. [18] and Popa et al. [19].

The DfHm� (SmPO4,s,298.15 K) value measured in

this study is matching reasonably well with that of

Ousoubalyev et al. [11] from calorimetry and that of

Marinova and Yaglov [14] from the solubility product

determination. However, DfHm� (SmPO4,s,298.15 K) by

Helean et al. [13] from oxide melt solution calorimetry, that

calculated by Popa et al. [19] and that reported by Cetiner

et al. [15] from solubility measurements are *20 kJ mol-1

lower than that measured in this study.

The enthalpy of formation of LaPO4(s), NdPO4(s), and

SmPO4(s) at 298.15 K from respective component oxides

can be calculated for reaction (7) using Eq. (8):

1=2 Ln2O3 sð Þ þ 1=2 P2O5 lð Þ ¼ LnPO4 sð Þ ð7Þ

Df�oxH�m LnPO4; sð Þ ¼ DfH�m LnPO4; s; 298:15 Kð Þ

� 1=2 DfH�m Ln2O3; s; 298:15 Kð ÞÞ

� 1=2 DfH�m P2O5; l; 298:15 Kð Þ

ð8Þ

Df-oxHm� (LnPO4,s,298.15 K) is calculated by taking the

value of standard enthalpy of formation of LnPO4(s)

(Ln = La, Nd, Sm) measured in this study and that reported

in the literature by different research groups [10, 11, 13–17].

These Df-oxHm� (LnPO4,s,298.15 K) values are plotted

against ionic radii of trivalent rare-earth ion (Ln3? = La3?

(103 pm), Nd3?, (98.3 pm) and Sm3? (95.8 pm) [20]) in

Fig. 6. The plot shows that Df-oxHm� (LnPO4,s,298.15 K)

calculated from oxide melt solution calorimeter [13] is lower

than that of the present study as well as other literature data.

However, different studies show a trend that Df-oxHm�

(LnPO4,s,298.15 K) value increases as the ionic radii of the

trivalent rare-earth ion decreases. This indicates that more

amount of heat is released (per mole) on formation of

LaPO4(s) than that of NdPO4(s) and SmPO4(s) from their

component oxides. Hence, LaPO4(s) is more stable than

NdPO4(s) which is more stable than SmPO4(s).

Conclusions

The enthalpy of formation of LaPO4(s), NdPO4(s), and

SmPO4(s) were measured using isoperibol solution calo-

rimeter at 298.15 K. Normally the enthalpy of formation of

104 102

Ionic radii of trivalent rare-earth ion(Ln3+)/pm

100 98 96 94–400

–380

–360

–340

–320

–300

–280

–260

–240

La3+

Nd3+Sm3+

Marinova et al.[10]

Ousoubalyev et al.[11]

Helean et al.[13]

Marinova and Yaglov [14]

Cetiner et al.[15]

Poitrasson et al.[16]

Wood anf Williams-Jones et al. [17]

This Study

Δ f-o

xH°

m(L

nPO

4,s,

298.

15 K

)/kj

mol

–1

Fig. 6 A plot of Df-oxHm� (LnPO4,s,298.15 K) as a function of ionic

radii of Ln3?. It compares Df-oxHm� (298.15 K) of LaPO4(s),

NdPO4(s), and SmPO4(s) measured in this study with that reported

in the literature


123

thermally stable substance is measured by high-tempera-

ture reaction calorimeter. However, this study measures the

enthalpy of formation of insoluble ceramics like rare-earth

phosphates at room temperature using solution calorimeter.

This study demonstrates successful application of this

technique for the enthalpy of formation measurements and

also elucidates energetic trends in rare-earth phosphates.

The thermochemical cycle suggested in this study for rare-

earth phosphates can be used for measurement of enthalpy

of formation of actinide phosphates.

Acknowledgements Authors are thankful to Dr. K. L. Ramakumar,

Director, Radiochemistry and Isotope Group and Shri S. G. Kulkarni,

Head, Product Development Division for their keen interest in this

study. Authors are also thankful to Dr. G. A. Rama Rao and Shri Neeraj

Gupta of Product Development Division and Dr. R. Mishra, Chemistry

Division for thermo-gravimetric studies and Dr. K. Krishnan of Fuel

Chemistry Division for XRD analysis.

References

1. Weber WJ, et al. Radiation effects in crystalline ceramics for

the immobilization of high-level nuclear waste and plutonium.

J Mater Res. 1998;13:6.

2. Clavier N, Podor R, Dacheux N. Crystal chemistry of the mon-

azite structure. J Eur Ceram Soc. 2011;31:941–76.

3. Ni Y, Hugues JM, Mariano AN. Crystal chemistry of the mon-

azite and xenotime structures. Am Mineral. 1995;80:21–6.

4. McCarthy GJ, White WB, Pfoetsh DE. Synthesis of nuclear waste

monazites, ideal actinides hosts for geologic disposal. Mater Res

Bull. 1978;13:1239.

5. PCPDFWIN version 2.2, JCPDS-ICDD, 2001.

6. Venugopal V, Shukla NK, Sundaresh V, Roy KN, Prasad R, Sood

DD. Thermochemistry of caesium iodide and caesium chromate.

J Chem Thermodyn. 1986;18:735–8.

7. The Enthalpy of Solution of SRM 1655 (KCI) in H20 Nat. Bur.

Stand. Certificate for SRM 1655 (KCl). US Dept. of Commerce,

Washington, March 1981.

8. Rat’kovskii IA, Botylin BA, Novikov GI. Dokl Akad Nauk

Beloruss SSR. 1973:232.

9. Henderson P. Rare earth element geochemistry. Elsevier; 1984.

p. 510.

10. Marinova LA, Glibin VP, Volkov AI. Rassh Tezisi Dokl Tbilissi

Meznierba. 1973;71.

11. Ousoubalyev I, Batkybekova M, Yousoupov V, Kydynov M.

(1975) Les Proprietes thermochimiques des phosphates et des

iodates des elements des terres rares. 4e Conference ternationale

de Thermodynamique Chimique. 217–23.

12. Ushakov SV, Helean KB, Navrotsky A. Thermochemistry of

rare-earth orthophosphates. J Mater Res. 2001;16:2623–33.

13. Helean KB, Navrotsky A. Oxide melt solution calorimetry of

rare-earth oxides. J Therm Anal Calorim. 2002;69:751–71.

14. Marinova LA, Yaglov VN. Thermodynamic characteristics of

lanthanide phosphates. Rus J Phys Chem. 1976;50:477.

15. Cetiner ZS, Scott A, Wooda T, Christopher H, Gammons H. The

aqueous geochemistry of the rare earth elements, Part XIV. The

solubility of rare earth element phosphates from 23 to 150 �C.

Chem Geol. 2005;217:147–69.

16. Poitrasson F, Oelkers E, Schott J, Montel JL. Experimental

determination of synthetic NdPO4 monazite endmember solu-

bility in water from 21�C to 300�C: implications for rare earth

element mobility in crustal fluids. Geochim Cosmochim Acta.

2004;68:2207–21.

17. Wood SA, Williams-Jones AE. The aqueous geochemistry of the

rare-earth elements and yttrium 4. Monazite solubility and REE

mobility in exhalative massive sulfide-depositing environments.

Chem Geol. 1994;115:47–60.

18. Williamson MA, Huang JC, Putman RL. Fundamental thermo-

dynamics of actinide-bearing mineral waste forms. US: Depart-

ment of Energy; 2000. p. 14.

19. Popa K, Konings RJM. High-temperature heat capacities of

EuPO4 and SmPO4 synthetic monazites. Thermochim Acta.

2006;445:49–52.

20. Greenwood NN, Earnshaw A. Chemistry of the elements. 2nd

edn. Butterworth-Heinemann. 1997. p. 1233.

21. Wagman DD et al. NBS Tech. Note 270-7. Selected values of

chemical thermodynamic properties. 1973.

22. Pedley JB, Kirk A, Seilman S, Heath LG. Computer analysis of

thermo chemical data, CATCH tables. Brighton: University of

Sussex; 1972.

D. Rawat, S. Dash

123

Documents

The standard molar enthalpy of formation of LnPO4(s) (LnÂ =Â La, Nd, Sm) by solution calorimetry