Solution enthalpies of hydrates LnCl3·xH2O (Ln=Ce-Lu)

Article abstract of DOI:10.1016/s0040-6031(98)00389-x

Trichlorides of the lanthanide elements Ln=Ce-Lu form: (a) isotypic hexahydrates LnCl₃·6H₂O with a coordination number (CN) 8 for the Ln³⁺ ions. (b) Two isotypic groups of trihydrates LnCl₃·3H₂O, in the first group Ln=Ce-Dy the CN is 8; the structure of the second group Ln=Er-Lu is unknown. With Ho no trihydrate exists; a dihydrate is formed. (c) Two isotypic groups of monohydrates LnCl₃·H₂O with unknown structure - Ln=Ce-Dy and Ln=Ho-Lu. For all compounds and for anhydrous chlorides LnCl₃ solution enthalpies were measured with an isoperibolic calorimeter. The Δ_solH⁰ values do not depend only on the difference (lattice enthalpies/hydration enthalpies), but also on the state in solution. According to Spedding the CN of the Ln³⁺ ions against water changes from 9 to 8 between Nd and Sm, causing minima in the series of solution enthalpies. Dihydrates LnCl₃·2H₂O are found for Ln=Ce, Pr, Nd, Sm and presumably for Eu and Gd. They are not yet well characterised.

Full text of DOI:10.1016/s0040-6031(98)00389-x

Thermochimica Acta 320 (1998) 1±7
Solution enthalpies of hydrates LnCl3ÁxH2O (LnˆCe±Lu)1
H.J. Seifert*, S. Funke
Inorganic Chemistry, University Gh Kassel, Heinrich-Plett-Str. 40, D-34109 Kassel, Germany
Received 6 May 1998; accepted 29 May 1998
Abstract
Trichlorides of the lanthanide elements LnˆCe±Lu form: (a) isotypic hexahydrates LnCl3Á6H2O with a coordination
number (CN) 8 for the Ln3‡ ions. (b) Two isotypic groups of trihydrates LnCl3Á3H2O, in the ®rst group LnˆCe-Dy the CN is
8; the structure of the second group LnˆEr±Lu is unknown. With Ho no trihydrate exists; a dihydrate is formed. (c) Two
isotypic groups of monohydrates LnCl3ÁH2O with unknown structure ± LnˆCe±Dy and LnˆHo±Lu. For all compounds and
for anhydrous chlorides LnCl3 solution enthalpies were measured with an isoperibolic calorimeter. The ÁsolH0 values do not
depend only on the difference (lattice enthalpies/hydration enthalpies), but also on the state in solution. According to Spedding
the CN of the Ln3‡ ions against water changes from 9 to 8 between Nd and Sm, causing minima in the series of solution
enthalpies. Dihydrates LnCl3Á2H2O are found for LnˆCe, Pr, Nd, Sm and presumably for Eu and Gd. They are not yet well
characterised. # 1998 Elsevier Science B.V.
Keywords: Hydrates of lanthanide chlorides; Preparation; Crystal structures; Solution enthalpies
‡
1. Introduction
which monomeric units [LnCl2(H2O)6] are held
together by the third Cl ions.
For two decades we have been investigating the
formation of ternary chlorides in pseudobinary sys-
tems alkali metal chlorides/LnCl3, where Ln are the
elements from Ce to Lu. Starting compounds for the
preparation of anhydrous rare earth chlorides are their
hexahydrates. For LnˆNd±Lu they crystallize at
ambient temperature from solutions of oxides in
hydrochloric acid. The earlier elements La±Pr form
hepta-hydrates; however, these are already dehydrated
at temperature lower than 508C to the hexahydrates.
The compounds LnCl3Á6H2O are isotypic: they
crystallize in the monoclinic space group P2/n, in
With few exceptions the hexahydrates can be
dehydrated to trihydrates. Monohydrates, in some
cases dihydrates also, were found by thermo-
gravimetry. A comprehensive TG investigation for
all elements Ln was performed in 1965 by
Haeseler and Matthes [1]. They used air and 1:1
air/HCl mixtures as the gas atmosphere. However,
instead of the use of a mixture with HCl, they could
not suppress the formation of oxidchlorides by hydro-
lysis, mainly with smaller ions. These are stronger
Pearson acids and are therefore more liable to reac-
tions with water like LnCl3‡H2OˆLnOCl‡2HCl. As
a consequence the mass loss in TG at temperatures
higher than ca. 808C was interpreted repeatedly as the
formation of compounds like, for example,
DyCl3Á2.5H2O.
*Corresponding author. Fax: +49 561 8044010.
1Dedicated to Professor Antonius Kettrup on the occasion of his
60th birthday.
0040-6031/98/$19.00 # 1998 Elsevier Science B.V. All rights reserved
P I I S 0 0 4 0 - 6 0 3 1 ( 9 8 ) 0 0 3 8 9 - X

2
H.J. Seifert, S. Funke / Thermochimica Acta 320 (1998) 1±7
To avoid such misinterpretations we used another
2.3. Solution calorimetry
strategy when investigating the dehydration of the
hydrates of cerium (III) chloride [2]: CeCl3ÁxH2O
was heated in a drying cabinet and greater than
908C in an HCl stream to temperatures increasing
all 10 h at 58C. The samples ± about 10 g in ceramic
boats ± were weighed at each step to control the loss
of water. If the weight loss was as high as calculated
for the formation of an expected lower hydrate, the
composition of the sample was controlled by X-ray
diffraction (XRD). Thus, we could ®nd out if either
a new phase was formed or if we had a mixture of
the two phases. This decision was possible by com-
paring the pattern with that of analogous hydrates
of other lanthanide chlorides. Thus, we could pre-
pare the existing hydrates in amounts suf®cient for
measuring solution enthalpies. These measurements
are described in this paper; the ÁsolH0 values are
discussed depending on the ionic radii of the Ln3‡
ions.
The heats of solution were measured using an
isoperibolic underwater calorimeter ®rst described
in 1978 [4]. The solvent H2O was slightly acidi®ed
with HCl (10 2 mol l 1). In a silver vessel of 1.3 l
volume samples of 2±4 g in thin-walled glass
ampoules are cracked under water at the start of the
measurement. Thus, with a dissolution rate 1:15 000
virtually ideal solutions are formed. The change of
temperature ÁT against the surroundings ± a thermo-
stat at ca. 258C with a temperature constancy of
2Â10 48C ± generated by the solution enthalpy is
measured with a thermopile. The inevitable heat loss is
compensated according to Dickinson: The corrected
temperature change ÁT* is obtained by setting F1ˆF2
in Fig. 1. The calibration was done by Joule's heat
giving a temperature change ÁTE. The heat of solution
QL is ®nally calculated from the equation
QL
QE
ˆ
ÁTÃ ÁTE
For the solution enthalpy of KCl we found
1
17.67Æ0.05 kJ mol from the literature [5].
2. Experimental
2.1. Preparation of hexahydrates LnCl3Á6H2O
Lanthanide oxides (99.9%, Fa. Heraeus, Hanau)
were dissolved in concentrated hydrochloric acid
and evaporated to saturation at ca. 508C. Under cool-
ing with ice the chloride±hexahydrates were precipi-
tated with gaseous HCl, sucked off with a glass sinter
crucible, washed with ether and dried at ca. 508C.
They were identi®ed by comparing their X-ray pat-
terns with calculated intensities in s.g. P2/n, using the
site parameters of PrCl3Á6H2O [3]. The correct water
content was controlled by thermogravimetry (Ther-
mobalance TG 750, Stanton Redcroft), using dehy-
dration to the trihydrate. (Chemical analysis for Ln
and Cl was not applied because the results are not
suf®ciently sensitive.)
3. Results
In Table 1 the measured solution enthalpies are
compiled. Each value is the mean of two measure-
ments with the deviation for the last digit in brackets.
Below these values the preparation temperatures at
which the weight loss of ca. 10 g hydrate has become
greater than ca. 0.05 g in 10 h are given. Then, the
sample was kept at this temperature till the weight loss
reached the calculated value. Experience had proved
that at temperatures higher than ca. 708C the forma-
tion of oxide chlorides by hydrolysis became effective.
As a consequence only the dehydration of hexahy-
drates up to NdCl3Á6H2O could be performed in air,
using a drying cabinet.
2.2. X-ray powder patterns
One example for this method: From 10.00 g
HoCl3Á6H2O, 8.56 g HoCl2Á2H2O will be formed.
The temperature of a perceptible degree of dehydra-
tion was 708C. This procedure had to be performed in
an HCl atmosphere. After 30 h the formation of the
dihydrate was complete; this was controlled by an
XRD record.
Powder patterns were taken at ambient temperature
with a Philips PW 1050/25 goniometer equipped with
a proportional counter and a vacuum attachment.
During exposure (Cu Kꢀ radiation) the samples were
kept under a helium atmosphere.

H.J. Seifert, S. Funke / Thermochimica Acta 320 (1998) 1±7
3
Fig. 1. Temperature±time diagram for an exothermic dissolution reaction in an isoperibolic calorimeter.
The preparation of monohydrates in HCl could be
made without dif®culty because their range of exis-
tence is ca. 508C. For instance, HoCl3Á2H2O is dehy-
drated to HoCl3ÁH2O at 100±1208C. The formation of
anhydrous HoCl3 begins above 1808C. However, it is
necessary to increase the temperature slowly to pre-
vent partial hydrolysis. Thus, the dehydration of 10 g
HoCl3Á6H2O to HoCl3 takes altogether 2 weeks.
The proof that for HoCl3, no trihydrate exists, is
given in Fig. 2. An X-ray pattern of a preparation with
the composition (HoCl3‡3H2O) reveals a superposi-
tion of the peaks from HoCl3Á2H2O and HoCl2Á2H2O.
The dihydrate itself has another peaks-motif than the
trihydrates of TbCl3Á3H2O and ErCl3Á3H2O, respec-
tively.
In the same way it could be proved that for dys-
prosium the lower hydrates DyCl3Á3H2O and
DyCl3Á1H2O exist, but no hydrate with 2.5 H2O [1].
DyCl3Á3H2O belongs to an isotypic group of trihy-
drates ranging from CeCl3Á3H2O to TbCl3Á3H2O. It
crystallizes in Pnma with Zˆ4. The structure consists
of [LnCl4/2Cl(H2O)3] chains where two Ce3‡ ions are
connected via two chloride ions.
the intensity calculation are taken from CeCl3Á3H2O
[2].
A special problem was the further dehydration of
the trihydrates LnCl3Á3H2O with LnˆCe±Gd. Accord-
ing to the TG measurements dihydrates should exist.
We could con®rm this for the elements LnˆCe, Pr, Nd
and Sm; for Eu and Gd we obtained no unequivocal
results. The prepared dihydrates are very hygroscopic,
give only poor X-ray patterns and reveal a high
tendency to lose water under the formation of mono-
hydrates. Thus, the temperature range for their pre-
paration is only ca. 58C.
From the literature only solution enthalpies of
anhydrous chlorides and their hexadrates are known,
compiled in [6,7]. It must be pointed out that our
values for the hexahydrates are 1.0±1.5 kJ mol 1 less
negative than those from the literature.
For completeness the solution enthalpies for the
existing heptahydrates LnCl3Á7H2O (LnˆCe, Pr) at
ambient temperature were also measured. These com-
pounds were prepared by precipitation with HCl gas
from saturated solutions at 08C, washing with ether
and drying in a N2 stream at 258C.
Measured values in kJ mol
1
Ê
Ê
Lattice parameters: aˆ12.22 A; bˆ8.56 A;
:
CeCl3Á7H2O:
Ê
cˆ6.68 A. hkl values and intensities for the re¯exes
25.3(2) ([8]: 28.9) PrCl3Á7H2O: 26.1(1) ([8]:
up to 208 are given in Table 2; site parameters for
29.5)

4
H.J. Seifert, S. Funke / Thermochimica Acta 320 (1998) 1±7
Table 1
1
Solution enthalpies ÁsolH0 in kJ mol and preparation temperatures
LnCl3Á6H2O
LnCl3Á3H2O
LnCl2Á2H2O
LnCl3ÁH2O
LnCl3
138.9(8)
Ce
Pr
39.4(1)
/508
38.1(1)
558/708
36.5(1)
/758
33.2(2)
558/758
34.1(2)
/858
36.5(1)
558/858
37.5(2)
608/908
39.4(5)
ꢀ658/908
41.2(2)
858/958
42.8(2)
808/958
43.9(4)
/958
89.3(5)
858/758
85.4(1)
908/758
91.9(8)
/858
93.5(4)
958/958
97.9(4)
/1208
100.7(8)
758/1208
107.5(4)
/>1308
107.6(2)
1208/>1208
Not ex.
±
108.6(6)
118.1(8)
1408/1208
112.6(2)
1558/>1208
116.1(4)
/1508
137.0(2)
1308/>1508
144.3(1)
/>1508
149.5(3)
1508/1608
155.9(3)
ꢀ2008/1808
157.6(2)
2308/1908
166.1(9)
2108/>1908
167.3(1)
2108/>1908
170.1(6)
/>1908
1008/858
145.8(5)
149.5(4)
164.7(2)
171.5(3)
180.3(5)
188.8(4)
197.1(3)
209.4(9)
211.9(1)
212.2(7)
212.9(8)
211.6 (6)
1088/908
112.2(3)
/1008
Nd
Sm
Eu
Gd
Tb
Dy
Ho
Er
1158/1108
1108/
1258/
135.1(9)
1358/>1208
112.7(1)
(1158)/>1208
115.1(3)
/>1208
119.5(3)
/>1208
1308/
Tm
Yb
Lu
45.9(2)
/1008
47.2(5)
/1008
175.4(4)
/>1908
172.7(3)
/>190
111.3(7)
/>120
(
TG values [1] cursive.
4. Discussion
working with a gas stream of controlled H2O pressure
are also possible, but are rather unwieldy.
4.1. The existence of chloride±hydrates
As described for the dehydration of HoCl3Á6H2O,
the existence or non-existence of a lower hydrate was
insured by X-ray patterns. This can be done exactly
when the structures of all involved phases are known,
so that peak intensities can be calculated. According to
our ®ndings monohydrates exist for all lanthanide
elements. The compounds LnCl3ÁH2O with LnˆCe±
Tb are isotypic; a second group of isotypic monohy-
drates ranges from DyCl3Á6H2O to LuCl3Á6H2O. For
trihydrates again two groups exist. The structures for
the compounds LnClÁ3H2O with LnˆCE±Tb were
recently determined in our group [2]. The structure
of the isotypic group LnˆHo±Lu is not yet known.
Trihydrates with Ho and also Y do not exist; here
dihydrates are formed. The issue of the existence
of other dihydrates has already been discussed in
Section 3.
The main issue is how to prepare chloride±hydrates
with the exact amount of hydrate water. Neither the
method of dehydration by heating at ambient pressure
or in vacuum nor TG measurements are equilibrium
methods, because the H2O pressure at an applicated
temperature is disturbed by spreading to an unlimited
volume either by a gas stream or by an applied
vacuum. Thus, one has to control the dehydration
process by continued weighing. It is appropriate to
choose the preparation temperature as low as possible.
Thus, this temperature will be in general lower than
that found by `rapid' method thermogravimetry. How-
ever, it must be pointed out, that equilibrium processes
are also possible: TG measurement in labyrinth cru-
cibles work quasiisobar and preparative methods

H.J. Seifert, S. Funke / Thermochimica Acta 320 (1998) 1±7
5
Table 2
X-ray powder pattern for DyCl3Á3H2O
#/grad hkl
Iexp.
Ical.
7.25
200
64
100
29
17
23
22
20
33
24
17
23
21
84
100
34
14
23
29
18
32
18
14
28
18
7
7.56
8.36
101
011
210
020
211
220
121
311
221
112
410
031
411
022
122
420
312
231
421
8.39
10.28
11.10
12.62
12.78
13.82
14.26
14.77
15.54
16.92
41
7
14
5
17.36
18.00
18.14
18.50
19.25
15
23
28
32
20
11
21
24
13
Cu Kꢀ, #ˆ0-208; s.g. Pnma.
be expressed as ÁsolH0ˆ Álat.H0‡Áhyd.H0, where
Á
lat.H0ˆlattice enthalpy, Áhyd.H0ˆhydration enthalpy
Both energies vary with 1=rꢁLn3‡†. Thus, with
decreasing radii in the series Ce3‡)Lu3‡ a greater
affordanceofvaporisingthecrystallattice(endothermic
Á
Á
lat.H0) will be more or less compensated by a gain in
hydH0.
Fig. 2. Measured X-ray peaks for HoCl3Á6H2O, HoCl3Á2H2O and a
sample with the composition (HoCl3‡3H2O).
Hydration enthalpies are the sum of the hydration of
In the literature other hydrates are also described,
like NdCl3Á5H2O and NdCl3Á4H2O [9]. As discussed
in an earlier paper [2], we could not con®rm their
existence.
the Ln3‡ ions and Cl ions, whereby the second will
be constant for all compounds. According to Spedding
et al. [10] the coordination number of the Ln3‡ ions
against H2O in aqueous solution is 9 for La3‡±Nd3‡
and 8 for Gd3‡±Er3‡. Between these groups a `dis-
placement equilibrium' exists for Pm3‡, Sm3‡ and
Eu3‡. They derive their ®ndings from partial molar
volume data, calculated from speci®c gravimetry
measurements of LnCl3 solutions. On the other hand,
4.2. Dependence of ÁsolH0 on ionic radii
As the thermodynamic circle for the solution of
anhydrous chlorides shows, solution enthalpies can

6
H.J. Seifert, S. Funke / Thermochimica Acta 320 (1998) 1±7
This, however, must be the `displacement equilibrium'
as predicted by Spedding, over the whole range of
Ce3‡ to Sm3‡, otherwise a distinct falling-off would
exist. This is in correspondence with the results of
Helm et al. [12] obtained by the neutron scattering
®rst-order difference technique with isotypic substitu-
tion. Kinks in the lines for the three- and monohy-
drates, might depend on change in crystal structure
between Dy and Er for LnCl3Á3H2O and between Dy
and Ho for LnCl3ÁH2O.
This `displacement effect' is less distinct for anhy-
drous chlorides. As already pointed out by Burgess
and Kijowski [7], ``the general trend of increasing
negative enthalpies of solution going from LaCl3 to
LuCl3 indicates that hydration enthalpies, which for
the 3‡ -cations increase in this direction as ionic radii
decrease, dominate increasingly over lattice enthal-
pies.'' The fact, that from HoCl3 to LuCl3, solution
enthalpies are almost constant means, that the term
( Álat.H0 ‡ Áhyd.H0) has become constant: lattice
enthalpy and Áhyd.H0 (Ln3‡) change in the same
degree.
4.3. Quantitative estimations
Gunsilius et al. [13] have calculated with single
crystal dates at ambient temperature the stable mod-
i®cation of TbCl3 (UCl3-type) and the Madelung part
of lattice energy as 4520 kJ mol 1. With a Born
coef®cient nˆ10 ± mean of 3 Cl (nˆ9) and one
Ln3‡ (nˆ12) ± the factor (1 1/n) (see textbooks of
solid state chemistry) is 0.9 and the lattice energy
ELˆEMad ERepˆ4067 kJ mol 1. The van der Waals
energy will be somewhat greater than that of BaCl2
(110 kJ mol 1) [14]), so that the lattice energy is more
than 4200 kJ mol 1. Thus anion polarisation is not
considered.
Fig. 3. Solution enthalpies in dependence on the atomic number.
all hexahydrates LnCl3Á6H2O [11] have the same
‡
crystal structure with a coordination [Ln(H2O)6Cl2] ;
their mole volumina, which are in a ®rst approxima-
tion a measure for lattice enthalpy, decrease continu-
ously from Ce3‡ to Lu3‡ (Fig. 3). That is, the energy
to overwhelm these lattice forces becomes larger, and
if this effect would be the only one the solution
enthalpies should become more positive, and less
exothermic.
However, the plot for all hydrates in Fig. 3 shows
that from Sm3‡ to Lu3‡ the solution enthalpies
become continuously more exothermic ± they are
determined by the increasing gain in energy by
increasing hydration enthalpies. From Ce3‡ to Nd3‡
the coordination number decreases from 9 to 8 which
causes a decrease in produced hydration enthalpy.
This value must be compared with the experimental
solH0ˆÁlatH0
value
Á
in
the
equation
Á
hydH0(Tb3‡) 3ÁhydH0(Cl ). Hydration enthalpies
are taken from Morss [15] who has derived them using
solution enthalpies of the cubic elpasolites
Cs2NaLnCl6: for Tb3‡ˆ 3604 kJ mol 1; for Cl
1
ˆ 363.2 kJ mol
.
With
Á
solH0(TbCl3)ˆ 188.8
1
1
kJ mol
one obtains ÁlatH0(TbCl3) kJ mol
ˆ‡189 3604 1090ˆ 4505.
1
The difference of 300 kJ mol must be related to
the polarisation effect. Thus the solution enthalpy is

H.J. Seifert, S. Funke / Thermochimica Acta 320 (1998) 1±7
7
the difference of two big numbers, 4505 kJ mol 1 and
Á
for calculating solution enthalpies from these quan-
tities.
References
hydH0gesˆ 4694 kJ mol 1 ± not a favourable situation
[1] G. Haeseler, F. Mattes, J. Less-Common Met. 9 (1965)
133.
[2] G. Reuter, H. Fink, H.J. Seifert, Z. Anorg. Allg. Chem. 620
(1994) 665.
An estimation can be made for the lattice enthalpies
of TbCl3Á3H2O and TbCl3Á6H2O. Taking into account
[3] J. Chen, O. Xu, L. Li, J. Wu, G. Xui, Acta Crystallogr. C47
(1991) 1074.
that the CN for Tb2‡ in solution is 8, the energy for
1
forming one Tb3‡±H2O bond is 450 kJ mol
.
[4] G. Thiel, H.J. Seifert, Thermochim. Acta 14 (1982) 363.
È
[5] S. Sunner, I. Wadso, Acta Chem. Scand. 13 (1969) 97.
[6] H. Bommer, E. Hohmann, Z. Anorg. Allg. Chem. 248 (1941)
373.
Trihydrate:
[Tb(H2O)8]3‡; (ÁhydH0)eff
[TbCl4/2Cl(H2O)3]‡5H2O
)
1
ˆ
2250 kJ mol
1
ÁlatH0 (TbCl3Á3H2O) kJ mol ˆ ‡ 108 2250
[7] J. Burgess, J. Kijowski, Advanc. Inorg. Chem. Radiochem. 24
(1981) 57.
1090 ˆ 3232
‡
Hexahydrate: [TbCl2Cl(H2O)6]
ˆ
‡
2H2O
1
)
[8] R.J. Hinchey, J.W. Cobble, Inorg. Chem. 9 (1970) 917.
[9] G.J. Kipouros, R.A. Sharma, J. Less-Common Met. 160
(1990) 85.
[Tb(H2O)8]3‡; (ÁhydH0)eff
900 kJ mol
1
ÁlatH0(TbCl3Á6H2O) kJ mol
ˆ
‡38
900
[10] F.H. Spedding, M.J. Pikal, B.O. Ayers, J. Phys. Chem. 70
(1966) 2440.
1090 ˆ 1952.
To make such calculations for the second group of
trihydrates and monohydrates, their crystal structures
have to be determined.
[11] E.J. Gaeber, G.H. Conrad, S.F. Duliere, Acta Crystallogr. 21
(1966) 1012.
[12] L. Helm, D.H. Powell, A.E. Merbach, Swiss Neutron News 5
(1994) 16.
[13] H. Gunsilius, H. Borrmann, A. Simon, W. Urland, Z.
Naturforsch. B43 (1988) 1023.
Acknowledgements
[14] T.E. Brackett, E.B. Brackett, J. Phys. Chem. 69 (1965)
3611.
This work was supported by the Deutsche For-
schungsgemeinschaft and the Fonds der Chemischen
Industrie. Their help is gratefully acknowledged.
[15] L.R. Morss, J. Phys. Chem. 75 (1971) 392.