Thermodynamic data (298.15 K)

Thermal expansion coefficient:

s jot = por wi (20°-70°) = 25.5]
FeCOs: Siderite.. 2808 with
Jmol deg-* | cal molt | References pore 10)=917 ° 9, Vol. I
lee -
CoC0, a5 m2 20 Thermodynamic data (298.15 K)
CoO 52.93, 12.65 10
Co.01 102.5 24.5 10 s
AHP J mol~ deg~' | cal mol-! deg! | References
kJ mol“! | kcal mol-* | References FeCO; 92.9 22.2 1,10
CoCOs — 722.6 =172.7 1,10 FeO (Fe.170) 57.49 13.74
CoO — 238.9 —51.1 10 FesOu 146.4 35.0 5
Co304 —905.0 —216.3 10 Fe.0s 87.4 20.9 5
auf
kJ mol-* kcal mol-* Keterences
Copper FeCOs 740.6 —177.0 | 270-3 unpubl.
. . FeO (Feosn0)} —266.5 —63.7 1
No high-temperature thermodynamic data are — Fes0x —17. — 267.0 1,10
available and decomposition pressures of pure — Fes0s ~822.15 —196.5 1,10
CuCO; have not been measured. The temperature
at which decomposition of the minerals, azurite
and malachite, becomes noticeable varies, not sur- Lead

prisingly, with the rate of heating [119].

Thermodynamic data (298.15 K)

°
J mol! deg? | cal molt deg! | References
CuCO, 87.9 21.0 10
Cu 42.63 10.19 5,6, 10
Cu,0 92.38 40.42 22.08+0.10 Ola
Auf
KJ molt keal mol-! | References
Cuco, | 595.0 142.2 10
Cud "| —155.77+0.63 | —37.23+0.15 Oa
CuO 170.83 1.26 | —40.83+0.30 Na
Iron

Anhydrous FeCOs can be prepared by heating
the 1/2 hydrate in CO, at high pressure (40 to 50
atm) near 300 °C. The thermal decomposition
always leads to a higher-valent oxide, since the
CO. liberated immediately oxidizes the FeO to
Fe;0, [22]. The reaction may also be written as

3 FeCOs=Fes04+2 CO2+ CO.

The oxidation occurs so fast that even in vacuum
no FeO is formed.

Density of FeCOs
Phase T(K) d References
c, siderite 293.9 “851 16

16

In a careful study of PbCO, decomposition,
Centnerszwer, Falk, and Awerbuch [36] deter-
mined the following successive equilibria at | atm

CO2:

74° 286°
PbCOs <=> 3 PbO -5 PbCO; <—>

0 412°

3
PbO - PbhCO; <—> 2 PbO - PhCO; <—> PbO.

In a kinetic study of PbCO3; decomposition, these
authors (31) found an induction period which was
accelerated by small amounts of H20, and a first-
order decomposition for the main course of the
reaction independent of moisture content.

Glasner and Hodara [57] have reexamined the
basic carbonates and find 2 PhO-3 PbCOs to be
the carbonate of lowest oxide content. They list
the following pressures:

PbCO; 2 PbO-PbCO; |5 PbO -6 PbCO;
°C 222 271 |240 = 246 = 320) 290-380
P (mm) 72 690) 75 81 690! 101 690
PbO -PbCO;| 3 PbO 2 PbCOs
°C 350-390 370 430
p (mm) 123 690 150 690

Kelley and Anderson [3] have used the enthalpy
data of Marshall and Bruzs [93] to calculate the
thermodynamic functions for the reactions
(1) 2 PbCO;= PbO - PhCO; + CO,
AH(c: 20,230+ 4.23 T—8.21 x 10-87?
AG(cal) = 20,230 — 9.74 log T+8.21
X 10-87? — 13.60 T



If Boldyrev’s arguments are correct, it would
seem that the study of reversible decomposition
offers the best chance for a quantitative study of the
relation between the various effects discussed above
and the rate of decomposition. This is particularly
true for the main course of the reactions which is
much less affected by the defect structure than is
the initial nucleation.

Boldyrev and Medvinskii [27] have attempted
a quantum-mechanical interpretation of the above
decomposition scheme by considering changes in
the electron energy levels which correspond to the
various types of decomposition, primarily in the
language of band theory. The theory is not yet suffi-
ciently well developed to permit quantitative pre-
dictions, although it provides a rational framework
for the classification scheme. In terms of this
theory, intra-ionic decompositions involve very
local changes in the electron levels, i.e., “excitons
of small’ radius.” The authors predict that these
reactions will in the future be describable by energy
level diagrams.

C. Suggestions for Experimental
Work

In view of the preceding discussion, it is clear
that considerable care in experimental design and
execution is required if meaningful thermodynamic
and kinetic data on the decomposition of inorganic
salts are to be obtained. For the measurement of
equilibrium pressures, the requirements are rather
simple: (a) the measured temperature must be the
temperature of the sample; this requires placing
the measuring device close to the sample. (b) The
measured pressure must be the equilibrium pres-
sure; precautions against pseudo-equilibria require
that the same pressure is obtained when approached
from above and below. (c) The condensed phases
must be well characterized in order for the measured.
Pressure to be thermodynamically meaningful. In
the case of a solid decomposition, both the reactant
and product should be examined for evidence of
solid solution formation. If a liquid phase forms,
pressure measurements are probably of very limited
usefulness since not only the composition of this
phase, but the component activities must be known
in order to calculate the equilibrium constant of
the reaction. The characterization of the solid
phases is also important for another reason. In
cases where the reactant can exist in more than
one crystalline modification, it is possible that a
form other than the most stable one becomes
“frozen” in’a metastable state at the temperature
of the decomposition. The decomposition pressure
and the reaction Gibbs energy calculated from it
will be different for the two forms and hence the
reaction to which AG® corresponds will not be prop-
erly identified. Hence AG? will .have its correct
equilibrium value only if the solid phases are the
most stable ones. (d) Just as the vapor pressure of
small drops is greater than that of large ones, so
the decomposition pressure of small crystals is

10

larger than that of large ones. Experimental evi-
dence for this phenomenon exists [36A]. For equi-
librium measurements, relatively large (> 10-? em)
crystals should probably be used, or the pressure
studied for several particle sizes large enough for
the pressure to be independent of particle size.

Since the measurement of meaningful decom-
position pressures is difficult, work on the deter-
mination of high-temperature heat capacities would
be appropriate. A perusal of the tables in this mono-
graph shows that only about half of the carbonate
decomposition pressures can be calculated from
existing thermodynamic data. Most of the remainder
(Cu, Fe, Pb, Zn) could be calculated if these Cp
values were available. Co and Ni also lack reliable

2. values.

“fhe proper execution of kinetic studies is more
difficult since more factors determine the rate of
decomposition than determine the equilibrium.

Since these factors have already been discussed
in the previous section, we only emphasize here
their consequences for the proper experimental ap-
proach. We assume that in a chemical study the
main interest lies in elucidating the chemical kinetic
steps, i.e., the reaction mechanism. Therefore it is
desirable to ensure that purely physical steps, such
as diffusion of the gas and heat transfer to the re-
action site, not be rate-limiting. In planning a ki-
netic study, the following factors should then be
considered:

(a) The significant temperature is that at the re-
action site, ic., at the phase boundary between the
reacted and unreacted material. Placement of the
temperature-measuring device is thus of critical
importance. Particularly in very endothermic or
exothermic reactions, temperature gradients within
the furnace and even within the sample may be
large.

(b). In order to minimize temperature gradients
within the sample, to ensure that the rate is not
controlled by the diffusion of gaseous products and
by the rate of heat transport between sample and
surroundings, the sample should be small and con-
sist of a thin layer of finely divided material unless
the study is primarily concerned with nucleus forma-
tion on the surface, in which case a large sample of
known geometry and perhaps surface structure,
e.g., a single crystal, may be preferable. If the same
rate is attained with samples of different thickness,
temperature gradients are probably absent.

(c) Depending on the mechanism of the reaction,
the particle size of the sample may or may not affect
the rate. Although the theoretical models discussed
in section B2 predict such dependence, experi-
mental evidence for both dependence and non-
dependence exist. It is therefore advisable to use as
narrow a size distribution as possible in each experi-
ment and to measure the rate of several such
distributions with the total sample size held fixed.

(d) Particularly for readily reversible reactions,
the observed decomposition rate depends on the
difference between the equilibrium pressure, Pe,



Nickel with the data in this work to give the partial pres-
sures of all the vapor species.

There is no evidence for the existence of an-

hydrous NiCOs, although hydrates are well known. Density of KeCOs
These lose H,O on heating up to 470 K, but some
CO, is probably also lost before all the H:O is re- Phase Tew) a
moved [41, 52]. In addition, the NiO product seems
to lose oxygen rather easily since metallic Ni has
been found in NiO formed by the decomposition © 293 2.330 gml 9, Vol. IT
of NiCOs below 700 K [114]. Stos.is of NiCO; has © 288 2.29 emi” 9, Vol. I
been calculated from low temperature Cp measure- 11908 | 19922 ¢em-s | 75
ments to be 85.4 J mol! deg [80]. 1185.3 1.8904
iss |
. i E
Potassium 1206.6 1.8824
1213.5 | 1.8778
The free energy functions of K2O in [10a] are 1220.0 1.8749
based on estimated Cy and S%yg values. In the 1966 1223.6 1.8735
revision [10b] fef values of KzCOs are listed to tags | Laero
2500 K, but decomposition pressures cannot be 1245.3 118640
calculated above 1100 K because fef data for K20 1257.4 | 1.8584
are lacking above this temperature. K2O is suffi- 1270.0 | 1.8527
ciently unstable for its dissociation to Kg) and Oz 1285.5 | heHer
to be significant. Equilibrium constants for this —
reaction are given in [10a] and can be combined ‘The liquid data are fitted by

Transitions of KxCOs

Phase 1K) | AH As
change

KJ J
mol | mol-! deg-*

References

e,IV>c, Il] 523
e, Ic, II 701
e,ll>e,I 895,
eI>1 1171+2 27

(75, 77, 111, 112

Thermodynamic data (298.15 K)

&
J cal
mol-ideg-! | mol-1deg-! | References
K.COs 155.5 37.17 10b
K,0 94.1 22.5 10b
AHP
KJ mol! kcal mol“! | References
K,COs —1150.2 274.9 10b
K,0 —363.2 = 86.8 10b

20

d=2,4141—0.4421 x 10 T (1180-1280 K).


High Temperature Properties and Decomposition of Inorganic Salts
Part 2. Carbonates

K. H. Stern! and E. L. Weise”

‘The literature dealing with the high-temperature behavior of inorganic carbonates has been
critically reviewed. Free energy functions, of reactants and products of the decomposition reactions
were calculated and have been tabulated from 298 K up to as high a temperature as possible. Free
energy functions and equilibrium constants of reactions were tabulated. Auxiliary data on phase
transitions, densities, and kinetics of thermal decomposition have also been included. ‘The literature
of the endothermic decomposition kinetics of solids, as it applies to carbonates, has been reviewed.

Key words: Carbonates; thermal decomposition; thermodynamic functions.

The Carbonates

A. Introduction

1. The Structure of the CO; Ion
The carbonate ion is a well-defined entity in which
the three oxygens are arranged in the same plane as
the carbon atom at a distance of 1.30+0.01 A and
with bond angles of 120°. This value of the C—O
distance, calculated by Pauling [103] on the basis
of the resonance structures

a}:

0: 0:
I |
c c

an

oN
:0: :0-+ +Q 70:

is in good agreement with experimental measure-
ments, on calcite [45,117], ranging from 1.29 to
1.31 A. The resonance energy is 176 kJ (42 kcal)
per mole.

2. General Features of Carbonate
Decomposition

All carbonates decompose with the evolution of
COz, the decomposition product being either the
corresponding oxide or a basic carbonate, the
latter decomposing with further evolution of CO: as
the temperature is raised. The oxides of some metals
are unstable in the same temperature range as the
corresponding carbonates so that the decomposi-
tion occurs as the simultaneous reactions

MCO;= MO-+ CO.g)
MO=M(g)+ 1/2 Ox(g).

1 Blectrochemistry Branch, Naval Research Laboratory, Washington, D.C. 20390.
for Basic Standards.
hhington, D.C. 20231. Formerly of the NBS Institute for

Basic Standards.

Consequently the vapor phase contains CO2, M, and
O». We have not calculated the oxide equilibria but
have indicated under the particular element whether
it is important, such as for,the higher atomic weight
alkali metals. Thus the melting points of pure
carbonates are generally not measurable directly,
though they may be estimated by extrapolation of
phase diagrams or by rapid heating, particularly
under high CO: pressures, which minimizes decom-
position, The dissociation, CO,—-CO+1/2 Oz,
is too slight at temperatures relevant for carbonate
decomposition to need taking into account (see
below).

As pointed out in out previous paper [121],3 the
thermal stability of a polyatomic anion decreases
with increasing polarizing power of the associated
cation, since polarization of the anion structure
leads to the distortion and consequent weakening
of the intra-ionic bonds.

This argument applies to the carbonates. There
is, however, some question as to how the polarizing
power of the cations is to be measured. For example,
Alekseenko [13] noted that the decomposition
temperatures (Peo, atm) of carbonates with 18
electrons in their cation outer shell were lower than
those of carbonates with cations of comparable
size containing 8 electrons. Ostroff and Sanderson
[98] suggested that an anion is most stable when it
has the greatest control over its valence electrons
and is in a nonpolarizing environment. Since the
polarizing power of the cations is inversely propor-
tional to their size and since the electronegativity
measures the power of an atom or ion to attract
electrons, Ostroff and Sanderson suggested that
anion stability should increase as some power of
(r/S), where S is the Sanderson stability ratio, a
quantity linearly related to other electronegativity

® Figures in brackets indicate the literature references at the end of this paper.


(2) PbO-PbCO,=2 PhO+COs
AH(cal)=20,910-+4.23 T—8.21 10-87?
AG*(cal) = 20,910 —9.74 T log T
+8.21 x 10-87?— 12.967.

iIn studying the decomposition of PbCOs in vac-
uum as a function of temperature and particle size,
Kadlets and Dubinin [81] found that, in con-
trast to the work of Centnerszwer and Awerbuch,
the induction period was absent if the sample was
brought to the reaction temperature very rapidly.
The rate (—dm/dt) was proportional to the sur-
face area. The expected conformity with eq (10)
was observed. Rate constants, expressed in g
s-tcm~*, were independent of particle size. Hence
the decomposition of PbCOs; conforms to the con-
tracting sphere model with rapid surface coverage
by nuclei. The kinetic results are summarized
by the relation

k(g cm~*s~!)= (0.5 + 0.5) exp (— 41,500 kcal/RT)

in the range 236 to 335 °C.

These results were confirmed by a similar study
in nearly the same temperature range [115]
for which E=40.2 keal was found. The reaction
rate with freshly prepared and with ground sam-
ples differed probably because of changes in the
surface structure.

Density of PbCOs

Phase T(K) d References

c 275.3 6.533 16
c 299 6.582 (x-ray) 123a
Thermodynamic data (298.15 K)
s

[x mot deg-* | cal mol deg | References
PbCOs 131.0 313 2
PbO (yellow) 68.70 16.42 2
PbO (red) 66.5 15.9 2
PbO-PbCOs 20s | 488 2

AHP

kJ mol-? | keal mol! | References
PbCOs = 699.1 167.1 2
PbO (yellow) | —217.3 —51.94 2
PbO (red) 219.0 = 52:34 2
PbO-PDCO, | 816.7 =195.2 2

Lithium

The melting point of LigCOg is still uncertain.
Two recent measurements [75,°111] are consid-
erably lower than the value of 1008 K given in (1),

17

but differ from each other by 6°. In both of these
studies, the salt was protected from decomposition
by being kept under a pressure of CO, high enough
to prevent weight changes. AHm values given by
[77] and [112] differ by 7 kJ.

Equilibrium pressures of CO, above liquid LizCO3
were measured by Janz and Lorenz [75], using a
dynamic method in which the temperature was
varied until weight changes (losses on heating, gains
on cooling) were noted. The resulting pressures are
larger by several orders of magnitude than K values
determined from thermodynamic data. This ap-
parent discrepancy can be accounted for by con-
sidering that K=(a:yolatco,)Pco, and therefore
(415,0/414.c0,) =K/Pco. For example, at 1000 K we
calculate X 10-8, whereas Poo, ~ 1.9X 10-2,
ie., Qtizo/@tisco; ~ 10-4. This is not unreasonable
since LigO almost certainly dissolves in molten
LigCOs. It is for this reason that we have relied,
whenever possible, on thermodynamically calcu-
lated equilibrium constants.

Density of LizCOs
Phase T(K) d References

ec 290.7 2.111 9, Vol.
1 1012.2 1.8246 75

1027.4 1.8189

1029.7 1.8190

1052.1 1.8119

1069.7 1.8036

1082.7 1.7987

1105.1 1.7910

1120.1 1.7843

The liquid data are fitted by

d= 2.2026 — 0.3729 X 10-*T (1010-1120 K).

Transitions of LigCOs

Phase change TK) AH kJ mol! | References
em] 993 i
999 42. 15
1008 1
Thermodynamic data (298.15 K)
s
Jmol! deg-? | cal mol-t deg! '} References
LisCOs 90.17 21.55 10b
Li 37.89 9.056 10a
aur
KJ mol-* kcal mol! References
LisCOs — 1216.04 — 290.64 0b
LiO — 598.7 143.1 10a



Transitions of Na:CO3

Phase change T® AH References
KJ mol
ec, IVe, UI 629 1
¢, He, IL 759 1
el>e,T 891 1
11 1129 +2 31 75, 77, 112
Thermodynamic data (298.15 K)
s
Jmol-* | cal mol-t deg- | References
deg"
Na,COs 138.7 33,17 10b
Na,O 75.27 17.99 10b
Ane
KJ mol-! | kealmol-! | References
Na:COs 1130.7, = 270.26 10b
Na,O 415.5 99.4 10b
Decomposition of Na:COs
Na:COx)= NazOie) + COxa
AH ia.15= 321.4 kJ, 76.8) keal
T Na;O Reaction
fef Afef AG? log K K
J deg? J deg Kd
298.15 75.27 ~150-11 276.62 3.45 x 10-#
400 ~78.19 ~150.0; 261.36 7.40 x 10-%
500 — 83.82 149.6, 246.56 1.75 x 10
600 90.51 148.8; 232.08 6.25 K 10%
700 96.98 147.55 218.12 5.29 X 10-17
800 — 103.58 145.8, 204.71 4.30 10-4
900 —110.03, —144.3; 191.44
1000 —116.26 = 143.03 178.34
41100 — 122.30 141.72 105.47
1200 — 147.25 —140.3 152.98
1300 ~154.13 ~ 139.8; 139.57
1400 — 160.65, = 139.32 126.31
1500 — 166.85 — 138.75 113.19
1600 ~172.74 138.2 99.77
1700 ~178.35 — 137.69 87.30
1800 — 183.70 =1B7.4 74.51
1900 — 188.82 136.61 61.81
2000 —193.73 — 136.05 49.21
2100 —198.44 = 135.5; 36.68
2200 202.97 135.05 24.23
2300 — 207.33 134.53 11.84
2400 211.52 134.05 0.454
Strontium 850 °C and was finished at 1175 °C. The rate for

In contrast to the numerous studies of CaCO;
and MgCOs, the decomposition of SrCOs has re-
ceived very little attention. Wanmaker and Radiel-
ovic [128] studied the decomposition in air thermo-
gravimetrically and report that it commenced at

fine and coarse powders was nearly the same.
Up to 950 °C the decomposition rate was constant,
at higher temperatures a change from zero-order
to first-order kinetics occurred as the reaction
progressed beyond a=0.5.

Lander [86] observed considerable hysteresis
in measuring the orthorhombic-hexagonal transi-


Decomposition of MgCOs
MgCOsy=MeOiert COnw
Ailing s= 116.9 kJ, 27.95 keal

T MgO MgCO; Reaction
fof fef Afef AG? log K K
. J deg J deg Ky
298.15 = 26.94 -174.72 64.85, = 13.361 4.36 X 10-4
400 — 28.45 —174.60 47.10 —6.151 7.07107
500 — 3153 174.23 29.82 -3.115 7.67x 10-4
600 35.09 ~173.23 13.00 ~ 1132 7.39 X 10
700 — 38.75 —172.90 —4.088 +0.305 2.02
800 —42.3; —172.00 —20.66 +1349 22.35
900 — 45.87 —171.02 — 36.98 + 2.146 140,
1000 = 49.2, 169.97 —53.03 +2.770 539.
Manganese Density of MnCO;
The thermal decomposition of MnCOs; is com-
plicated by the easy interconvertibility of the Phase T(K) d References
various manganese oxides. When MnCOs; is
decomposed in O,-containing atmospheres, the 293 | 3125 emt | 9, VoLT
oxide formed depends on the temperature. In © ‘ 125 gl” Vol.
addition, a number of nonstoichiometric oxides ° “odechrosite 294.4 | 3.638 16
are also formed [24], the particular oxide formed -
depending primarily on the temperature and Thermodynamic functions (298.15 K)
partial pressure of oxygen. In the absence of
oxygen, the decomposition product is MnO, at s°
least up to 300 or 400°C [43]. Above this tempera-
ture range, MnO may be oxidized by the COz J j
i cal
formed according to [94] smol-deg- | motSteg-+ | References
MnO + 1/3 CO2= 1/3 Mn304+ 1/3 CO.
ae MnCO. 85.8 20. 10
The decomposition can thus also be thought of M0 59.71 tea : 10
= Mn;0.(a) 154 36.8 99
3 MnCOs=Mn30, +2 CO2+ CO. MnsO 1 368 2
The actual temperature at which such trans- MnO, 53.05 12.68 5
formations occur therefore depends on the efficiency
of COz removal and on the partial pressure of Os. an
The thermodynamics of many reactions involving
manganese compounds have been given by Mah 7 7
[90a]. The predominent mode of decomposition KJ mol keal mol References
of MnCO; is MnCO;=MnO+CO, and we have
therefore given equilibrium constants only for MnO; —94,96 5
this reaction. Our results agree very closely with MnO 385.1 6, 10
Mah’s. Thermodynamic functions for other oxides MusOu(a) 1386-5 of,
are listed in [11 and 90a]. The kinetics follows yng." 2568 9, 10
the contracting sphere model.
Decomposition of MnCO3
MnlOs¢ InOie) + COne
SH G00.15= 116.32KJ, 27.80 kal
T MnO MnCO3 Reaction
fef fef Atef AG log K K
Jdeg- Jeg" kJ
298.15 59.71 = 187.56 60.39 —10.580 2.63 x 10-4
400 61:53 = 187.47 41:32 = 5.396 4.02 x10-*
500 — 64.98 = 187.13 22.75 —2.377 4.20 x 10-*
600 68.95 103.9; ~ 186.6, 4.345 0.378 4.19 10-1
700 -297 112.39 = 185.9 = 13.815 +1031 10.74

19


Calcium

More kinetic studies of the thermal decomposition
of CaCO; have been reported than for any other
carbonate, and perhaps for any other salt. Many
of the general features of carbonate decomposition
described in the Introduction have been elucidated
from CaCOx decomposition and will not be repeated
here.

The reaction CaCO3=CaO+ CO is reversible.
Decomposition occurs without formation of inter-
mediate products. For compacted powders the
contracting sphere model applies [72]. A compari-
son of experiments done with various particle sizes
and with single crystals indicates that even the
reaction order may vary, depending on which mech-
anism is rate-determining at each phase of the re-
action. For example, in a careful analysis of
composition versus distance in a rectangular block
of powder, Kappel and Huttig [82] found different
rate laws applying at varying distances from the
surface. Such apparent changes in mechanism can
easily occur if different processes, such as heat
transport to the reaction site or escape of CO2,
become rate-determining at various stages of the
decomposition. In decomposing a cube of CaCO3
with thermocouples imbedded in it, it was found
that the center of the cube was colder by from 15
to 85° than the nominal furnace temperature [18].
The authors suggest from this and other evidence
that the decomposition rate is proportional to the
heat absorbed. Such a mechanism would most
likely not be applicable to thin layers of fine-grained
powders.

Many values of the activation energy have been
reported, nearly all near 167 kJ (40 kcal)/mol,
a value approximately equal to the standard
enthalpy of decomposition.

The aragonite— calcite transition has been
studied as a function of temperature and pres-
sure [37, 90]. At ambient pressure, the transition
occurs near 730 K [62, 122]. In addition, a transi-
tion from vaterite to calcite occurs in the range
620 to 670 K, with small impurity concentrations
producing large changes in the transition tempera-
ture [122].

Density of CaCOs
Phase T(K) d References
c, calcite 273 | 2.771 emit 9, Vol. Il
c, calcite 298 | 2.713 15
c, calcite. 299 | 2.711 (x-ray) 123a
¢, aragonite 297 | 2.930 15

‘Thermal expansion coefficient:

| +0

CaCOs;: Calcite . .

ile

CaCOy: Aragonite. . { 20°—

jax(2° to 81°) = 25.1353

.01182, || opt. axis;

| jand — 5.5782 + 0.00138¢,
jt opt. axis

26.21 +0.0160(¢—40),
pt. axis; and =—5.40

+ 0.0087(¢—40), opt.

axis
|a(50° to 60°)= 1.447
| afe=34.60+-0.0337¢
of? = 17.19 + 0.0368"
ag? = 10.16 + 0.0064¢

ow |

9, Vol. TIT

Transitions of CaCOs

Phase Change

T(K) References

¢ (aragonite c (calcite) 728 62
¢ (vaterite)—> c (calcite) 623-673 122
Thermodynamic data (298.15 K)
5°
[J mol-! deg-! |cal mol-! deg-'] References
CaCO.
(aragonite) 88.7 21.2 1,5
CaCO; (calcite)! 92.9 22.2 5,6
Cad 39.7 95 5,6
AHf
kJ mol~! | Keal mol-! | References
CaCOy
(aragonite) —1207.08 | ~288.49 1
CaCO; (calcite) | — 1206.87 | —288.45 1
cao — 035.09 151.79 o

14


The second limiting case arises when the rate
of linear propagation of the reacting interface
is so great that each particle decomposes as soon
as it has a nucleus. For this case, Hume and Colvin
give

Rate=koM:
dN|dt=koN: ap

where N;, is the uumber of particles at time t.

iLe.,

Integration from No to NM: gives

1n(N:dNo) = kot (12)
a first-order reaction. This behavior is most likely
to be approached by very small particles.

In general the kinetics of decomposition are
more complex than in the two limiting cases just
outlined. Thus, if nucleation is not extremely
rapid, the number of nuclei will increase in time.
Each of these nuclei will then begin to grow at a
different time. Some of them will spread over
sites at which nucleation would have occurred, at
some future time. ‘These sites are referred to as
“phantom nuclei” by Jacobs and Tompkins. The
problem has also been studied in great detail by
Mampel [92]. Both Jacobs and Tompkins and
Mampel assume the rate of nuclei formation to
be first order,

= kyNoeTM! (13)

dt

where No is the total number of possible sites.
For the induction period Mampel finds a ¢* law
for a, a relation derived by Jacob and Tompkins
for random nucleation and short times, and by
Fischbeck and Spingler [51] for the case of no
overlap, a constant rate of formation of nuclei, and
a large total number of nuclei.
For large radii (R>ket), Mampel finds

ket
R
an eq (10) given previously for the contracting

sphere case. For small radii the fraction decomposed
is

=1—(1—-a)'*

X=1-ce*#

where ¢ is independent of t; but does depend on
R in a rather complicated way. The rates are
dependent on particle size and both small and large
radii give lower rates than intermediate radii.

A generalized equation has been derived by
Erofeyev [47] in terms of probability theory. In
his treatment

a=1—exp (= [‘ear)
0

where p is the probability that a molecule will
react in the time interval dt.

(14)

The general solution of (14) for thermal decom-
position is
a=1—exp (—kt*) (15)
where the constant n depends on the shape of the
nucleus and the number of electrons necessary
for the formation of a stable nucleus. For example,
cylindrical ‘nuclei (the centers of formation are
edges or surface cracks) give n=3, flat nuclei
give n=4. Thus the evaluation of n, most easily
carried out by transforming (15) into
in [—In (1—a)]=Inkt+n Int (16)
and plotting the left side of (16) versus In ¢, gives
some information about the nucleation mechanism.

(b) Surface Area Changes During Decomposition

Since the molar volume of the product oxides
is less than that of the corresponding carbonates,
decomposition results in a porous structure through
which the CO. readily escapes. Although not
directly related to the decomposition kinetics,
the resulting increase in surface area during decom-
position is of some interest in the production of
“active solids” and has been treated by several
authors. Accordingly, we briefly review it here.

The oxide is first formed at sites on the carbonate
lattice and is therefore in a strained state. Gregg
[61] describes these as small crystallites or micelles.
This strained pseudo-lattice will tend to recrystallize
into a stable oxide lattice with a resulting decrease
in surface area. Three mechanisms for this re-
crystallization are generally recognized. At tempera-
tures T/Tm <0.2, where Tm is the melting point
of the oxide, adhesion predominates. This mech-
anism results from contact between parts of the
micelles which are atomically contiguous. For
0.2 <T/Tm< 0.35, surface diffusion predominates.
Here. mobile atoms or ions in the surface layer
leave their position on the lattice to migrate to
sites of lower energy, i.e., to stable oxide lattice
positions.

In DTA studies of carbonate decomposition [109]
small exothermic peaks (200-1000 cal/mol) were
observed at temperatures 40 to 80 °C above the
main endothermic decomposition peak, consistent
with such an oxide lattice rearrangement. The
formation of the surface mobile species from the
active species on the original lattice is considered
by IIschner [71] to be slower than the prior formation
of the active species and the subsequent incorpora-
tion of the mobile species into the stable oxide
lattice. Above T/Tm=0.5 sintering, i.e., lattice
or bulk diffusion, is the most ‘significant process
leading to decreased surface area. Thus as decom-
position proceeds, there are basically two processes
leading to changes in surface area: the formation
of the metastable oxide produces an increase in
surface, and the formation of stable oxide decreases



References

{1] Rossini, F. D., Wagman, D. D., Evans, W. H., Levine, S.,
and Jaffe, J., Selected values of chemical thermody
namic properties, NBS Circ. 500 (U.S. Government
Printing Office, Washington, D.C., 20402, 1952).

[2] Wagman, D. D., Evans, W. H., Halow, I., Parker, V. B.,
Bailey, S. M., and Schumm, R. H., Selected values of
chemical thermodynamic properties, NBS Tech. Note
270-3 (1968) (An advance issue of a revision of the
tables of [1] containing the first 34 elements in the
standard order of arrangement, Some selected, but not
yet published, values are also included.).

[3] Kelley, K. K., and Anderson, C. T., Contributions to the
data on theoretical metallurgy. IV. Metal Carbonates—
Correlations and Applications of ‘Thermodynamic
Properties. U.S. Bur. Mines Bull. 384, 73 pp. (1935).
Reissued as part of Bull. 601 (1962).

[4] Kelley, K. K., Contributions to the data on theoretical
metallurgy. XIII. High temperature heat-content, heat
capacity, and entropy data for the elements and in-
organic compounds, U.S. Bur. Mines Bull. 584., 232 pp.

(1960).

{5] Kelley, K. K., and King, E. G., Contributions to data on
theoretical metallurgy. XIV. Entropies of the elements
and inorganic compounds, U.S. Bur. Mines Bull. 592,
149 pp. (1961).

Lewis, G. N., Randall, M., Pitzer, K. S., and Brewer, L.
Thermodynamics, 2d ed. (McGraw-Hill Book Co., New
York, N.Y., 1961).

[7] Wagman, D., Evans, W., Halow, I, and Dudley, R., Sec-
tion 4, pp. 170-216 in American Institute of Physics
Handbook, 1963.

[8] McBride, B. J. Heimel, S., Ehlers, J. G., and Gordon, S.,
Thermodynamic properties to 6000 °K for 210 sub:
stances involving the fist 18 elements, NASA SP-3001,
(Washington, D.C., 1963).

[9] International Critical “Tables (McGraw Hill Book Co.,
New York, N.Y., used for some densities in Vol. I, IL,

1928).

[10] Kubaschewski, 0., and Evans, E. LL., Metallurgical
‘Thermochemistry, 3d ed. (Plenum Press, New York and
London, 1958).

[102] JANAF Thermochemical Data compiled and calculated

by the Dow Chemical Co., Thermal Laboratory, Mid-
land, Mich. (1965).
(b) First Addendum 1966.
(c) Second Addendum 1967.
[11] Coughlin, I. P., Contributions to the Data of Theoretical
letallurgy. Heats and Free Energies of Formation
of Inorganic Oxides. U.S. Bur. Mines Bull. 584 (1954).

[12] Adami, L. H., and Conway, K. C., U.S. Bur. Mines, Rept.

Invest. No. 6822, 7 pp. (1966) CA 65, 9830 (1966).

[6]

[13] Alekseenko, L. A., Sbornik Nauch Rabot. Akad. Nauk.

Beloruss. SSR, Inst. Khim. 36 (1956).

[14] Ambrozhii, M. N., Luchnikora, E..F., and Sidorova, M. I.
Zhur, Neorg. Khim. 5, 366-71 (1960), Engl. Transl

. 176.

15. Amicrson, C. T., J. Am. Chem. Soc. 56, 340-2 (1934).

16] Anderson, C. T., J. Am. Chem. Soc, 56, 894 (1934).

[17] Andrussov, L., Z. Phys. Chem. 115, 273-88 (1925).

[18} Aufiredic, J. P., and Vallet, P., C. R, Acad. Sci., Paris,
Ser. C 263, No. 10, 652-5 (1966).

[19] Baker, E. H., J. Chem, Soc. 2525 (1962).

[20] Barro, P. and Bonnefair,L., Bull. Soc. Chim., France S76
(196i

(21 Batuccas T., Rancaio, A ., Ibarz, J., An. Soc. Esp. Fis.

Quim, 27, 290 (1929)
[22] Berg, L. G., and Buadow, K. A., Zhur. Neorg. Khim. 6,
2003-8 (1961).

[23] Berg, L. G., and Buzdov, K. A., Zh. Neorgan. Khim. 7,
1773-8 (1962).

26

[24] Bergstein, A., and Vintera, J., Coll. Czech. Chem. Com-
muns, 22, 884-95 (1957).

[28] Bischoff, F., Zanorg. Allgem. Chem. 250, 10 (1942);
ibid. 262, 18 (1950).

[26] Boldyrer, V. V., Kinetika i Katalio 1, 203 (1960).

[27] Boldyrev, V. V., and Medvinskii, A. A., Kinetika i Kataliz

963)
[28] Bretsznajder,

4, 481 (1 .
S., and Cibor, E., Bull. Acad. Polon. Sei.

Chim. 13, 709-12, 713-16, 765-9 (1965).

[29] Britton, H. T. S., Gregg, S. J:, and Winsor, G. W., Trans.
Faraday Soc. 48, 63-69 (1952).

[30] Buzdov, K. A., Zhur. Neorg. Khim. 11, 1251 (1966).

[31] Centnerszwer, M., and Awerbuch, A’, Z. Phys. Chem.
123, 127-33 (1926).

[32] Centnerszwer, M., and Bruzs, B., Z. Phys. Chem. 114,
237-50 (1924).

[33] Centnerszwer, M., and Bruzs, B., Z. Phys. Chem. 115,
365-76 (1925).

[34] Centnerszwer, M., and Bruzs, B., Z. Phys. Chem. 119,
405-18 (1926).

[35] Centnerszwer, M., and Bruzs, B., Z. Phys. Chem. 123,
111-26 (1926).

[36] Centnersewer, M., Falk, G., and Awerbuch, A., Z. Phys.
Chem, 115, 29-53 (1925).

[36a] Centnerszwer, M., and Krustinsons, J., Z. Phys. Chem.
130, 187 (1927), 132, 185 (1928).

[37] Clark, S. P., Jr., Am. Mineralogist 42, 564-6 (1957).

[38] Cremer, E., Allgauer, K., and Aschenbrenner, W., Radex
Rundschau 1953, 494-502.

[39] Cremer, E., and Nitsch, E., Z. Elektrochem. 66, 697

(1962).
[40] Cremer, E., and Nitsch, W., Z. Elektrochem. 66, 697-708
1962)

[41] Deren, J., Haber, J., and Sloczynski, J., Bull. Acad.
Polon. Sci., Ser. Sci. Chim. 9, 245-52 (1961).

[42] mei pull Soc. Chim. (3) 19, 34 (1898); Bull. Soc.

Min. 20, 195 (1897).

[43] Dolimore,'D., and Tonge, K. H., Proc. 5th Symp. Reac-
tivity of Solids, 1964, pp. 497-! “508 (Elsevier Publ. Co.,
Amsterdam, The Netherlands, 1965).

and 6 C., Anal. Chim.
ta 20, 263-7 (1959).

N., J. Am. Chem. Soc. 59, 1380 (1937).
, Talanta 12, 883-92

(4s) wafer
[46] Bde, L Liptay, G., and Cal,

1965).
a Eiticen B. V., Dokl. Akad. Nauk. SSSR 52, 511 (1946).

Ferrari, Al, and Colla, C., Atti Aecad. Lincei 10, 594
(1929).
[49] Fischbeck, K., and Schnaidt, K., Z. Elektrochem. 38,
99-205 (1932).

[50] Fischbeck, K., and Schnaidt, K., Z. Elektrochem. 38,
649 (1932).

[51] Fischbeck, K., and Spingler, H., Z. f. Anorg, u. allgem.
Chem. 241, 209 (1939).

52] Francois, J., Compt. Rend. 230, 2183-4 (1950).

53) Gafner, G., Trans. Faraday Soc. 55, 981 (1959).

54] Garner, W. E., Ed., Chemistry of the Solid State (London:
Butterworths, 1955).

[55] Gattow, G., Pure Appl. Chem. 2, 121-3 (1961).

[56] Ginzbure, D. M., J. Gen. Chem. USSR, 26, 1103-5 (1956).

[57] Glasner, 'A., and Hodara, 1., Bull. Res. Council (Israel)

5 00 (1958).
[58] Glasner, A., Levy, E., Steinberg, M., and Badenheimer,
, Talanta 11, 403-14 (1964).
{39 Goldschmidt, V., Ber. 60, 1263 (1927).
Gordeeva, V. A., Zhabrova, G. M., and Kushnerer, N. Y.,

Dokl. Akad. Nauk. SSSR, 129, 835-8 (1959).

[61] Gregg, S. J., J. Chem. Soc. 3940 (1953).

62] Gruver, R. M., J. Am. Ceram. Soc. 33, 171-4 (1950).

(63] Hashimoto, H., Nippon Kagaku Zasshi. 82, 1605 (1961).



instances we have listed the thermal expansion
coefficients, where these were listed in’ Vol. III of
[9]. The quantities given are:

—— > at £°5 = lta th Ate.
7 at t°; or a= 7A over the range At*.

A=10 40 (resp. 10° 5). i.e., 10°X the rate of
change of a with

t, at t° (resp. over At®).

5. Decomposition Equilibria

As a measure of the stability of each carbonate,
we have calculated the equilibrium constants and
pressures of CO, for each decomposition reaction.
Equilibrium pressures may either be obtained from
direct measurement in which the carbonate and
oxide are allowed to equilibrate with CO2 at various
temperatures or they may be calculated by standard
thermodynamic procedures.

Direct measurements are frequently unreliable
since in many cases “pseudo-equilibria” appear, i.¢.,
different constant pressures are obtained, depending
on whether the equilibrium is approached from
above or below [131]. If the same pressure is ob-
tained by both approaches, this possibility is mini-
mized and we have listed a few such values in cases
where the necessary thermodynamic values were
not available. Whenever possible we have calcu-
lated the equilibrium properties by thermodynamic
methods.

For these calculations, four items of information
are needed:

(1) Knowledge of the course of the decomposition.
For the carbonates this is generally quite simple,
particularly if the carbonate and oxide coexist as
pure solids. In that case the course of the reaction
is described by the overall stoichiometry, e.g.,
MCOs=MO+CObp, and the equilibrium pressure
of CO: is a function of temperature only.

Several carbonates, e.g., lead carbonate, do not
decompose directly to the oxide but rather to a basic
carbonate such as PbO -PbCO3. These cases are
described in the individual sections, but equilibrium
pressures can usually not be calculated for the basic
carbonates because the necessary thermodynamic
data are largely lacking.

(2) High-temperature thermal data such as those
tabulated in reference [4] in the form of enthalpy
and entropy increments above 298.15 K, for each
compound.

(3) Heats of formation at 298.15 K.

(4) Entropies at 298.15 K.

These data are required for the decomposition
products as well as for the carbonates themselves.
Items (3) and (4) are tabulated as “Thermodynamic
Data” under each carbonate to show the data used
and to point out the gaps in existing data.

Tt has usually been most convenient to perform
the calculations through the medium of “free

energy functions” (fef) as discussed in reference
[6], pp. 166-9. Where the high-temperature thermal
data are available in the form of enthalpy and
entropy increments (Hr—Hg and S7—Sfs),
these are combined with 298.15 K entropies to get
the free energy functions:

fof = Gi Hise Hi Hes (52 34.) — Si

where T is the temperature of interest in K. Then for
a given T K, the fef’s for decomposition products and
sulfate are added and subtracted, in the same
manner as heats of formation in a thermochemical
equation, to yield a Afef for the decomposition
reaction. The heat of reaction at 298.15 K is obtained
from the heats of formation. Then the Gibbs energy
change attending the reaction is

AG7= TAfef + Ais

from which the equilibrium constant for the reaction
ais obtained through the relation:

RT InK.

AGt=
For the case of simplest stoichiometry:

MCO;=MO+CO: — K=Peo,.
In some cases the decomposing carbonate and the
oxide product may form a solid or liquid solution. In
that case the equilibrium constant

v — amo

KR awcos 0
where ayo and ayco, are the oxide and carbonate
activitics in the condensed phase and the pressure
is now dependent on the solution composition. Such
information is not always provided. We would there-
fore caution against uncritically accepting measured
decomposition pressures as meaningful thermo-
dynamic data unless the condensed phases have
been examined, for example, by x-ray diffraction
of the solids, or direct chemical. analysis.

Many of the older decomposition pressure
measurements were used by Kelley and Anderson
[3] in the calculation of thermodynamic properties
of reactions and compounds. When no newer data
were available, we incorporated their results into
this monograph. The thermodynamic functions
in the tables were calculated in 1966 and the litera-
ture survey is reasonably complete up to that date.
In some cases newer 298 K values have been in-
troduced when these represented major changes
from existing values or when this was necessary to
be consistent with the ongoing revision of [2]. The
values at higher temperatures were then revised
accordingly.


scales. As a measure of stability they chose the
“lowest temperature at which decomposition could
be noted,” and found that for several sulfates this
temperature varied as (r/S)1/?.

In order to define stability in more quantitative
terms, it would seem that a thermodynamic measure
would be less ambiguous. If one examines the
thermodynamic functions for the thermal decom-
position of the carbonates, it is evident that the rea-
son for differences in stability are due to AH® since
AS° values are nearly the same for all the carbonates
and arise primarily from the formation of a gaseous
product. Since AH° values at 298 K and those at
higher temperatures increase in a parallel manner,
the 298 K values are an adequate measure of sta-
bility. It has been shown by a semi-empirical method
[121a] that AH%,, is a linear function of r¥?/Z*,
where r is the cation radius and Z* is the effective
nuclear charge, as calculated from Slater’s rules

cs

Hyg (kcal)

1 L L 1 Ll
02 03 0.4 Os 06 o7
FicuRE 1. Decomposition enthalpies as a function of cation r'"/Z*.

[119a]. The plot is shown in figure 1. Although this
correlation is probably not unique, it does illustrate
the kind of connections to be expected between
atomic and thermodynamic properties.

3. Phase Tran:

a. Solid Transitions. The literature on solid
transitions is confusing. Several systems of naming
them exist and differences in the reported values of
the transition temperatures frequently are so great
that it is not clear whether the same or a different
transition is being described.

number of well-defined transitions have been
listed in NBS Circular 500 [1]. The data on these
and others have been collected and are listed.

b. Melting Points. Melting points have been
reported for a few of the more stable carbonates.
We have selected what appear to be reliable values.
It should be noted, however, that the decomposition
which occurs below the melting point makes a

measurement of the pure compound difficult and
that the solid phase may contain some oxide. Melting
points are usually obtained by rapid heating of the
salt to minimize decomposition, but it cannot always
be avoided. Another technique is to measure the
melting point under a high CO, pressure. The listed
values should therefore be regarded as lower limits.

c. Boiling Points. All carbonates decompose so
extensively in the solid or liquid range that boiling
points cannot be measured.

4, Density

The total literature of high-temperature density
measurements on inorganic carbonates is very small.
There exists only one set of reliable values for the
liquid salts, those of Janz and Lorenz [75] on the
alkali metal carbonates. We have reproduced their
original values and also their best fit to a linear
equation. On reason for the lack of data is that most
carbonates decompose considerably below the
melting point so that a measurement of any physical
property of the pure liquid is inherently impossible.
Except for a few limited studies below 100 °C, there
appear to be no studies of the temperature de-
pendence over a wide range of temperature in the
solid state.

The densities of solid salts are generally de-
termined either directly by pycnometry, or are
calculated from crystallographic (x-ray diffraction)
data. In the latter method the atomic packing must
be assumed. Since no allowance is made for crystal
imperfections, calculated densities are frequently
higher than experimental values. Since most
crystallographers do not indicate the temperature
for which the density is calculated, we have pref-
erentially listed experimental values when these
were available. An exception is the careful x-ray
work of Swanson and co-workers [123, 124] which
has been listed. Units for x-ray densities are
g/cm-*, In some cases experimental values seem to
depend on the method of preparation, probably
because of differences in the defect structure. Thus,
although the density of a given sample can certainly
be determined to four or five significant figures, we
would caution against attributing too much signif-
icance to figures beyond the third. This is par-
ticularly true because the units of the reported
density values are frequently in doubt. In many
cases they are not stated at all; in others it is not
clear whether the unity is g/cm* of g/ml; in still
others they are specific gravity values and may be
referred tu water at 4 °C or at some other tem-
perature. The difference between milliliters and
cubic centimeters is so small (the ratio is 1.000028)
as to be negligible. However, the difference in the
density of water between 4 and 20 °C is in the third
significant figure (1.00000 and 0.99823 g/ml, re-
spectively) and no longer negligible for the better
measurements.

We have indicated the units as reported by the
original authors. When no units are given, it can
generally be assumed that g/ml is meant. In a few



from calcite
from aragonite Ain

Decomposition of CaCO:

CaCOyg= Cat COny
Affins = 118.28 kJ, 42.61 keal
78.45 KJ, 42.65 keal.

A. Calcite
T CaO CaCOs Reaction
fef fef Afef AG log K K
J deg-' Jdeg-' J deg
298.15 —39.7 92.9 = 160.5
400 ~415 —96.4 = 160.3
500 —44.9 = 103.3 ~159.7
600 —48.7 —14 = 158.9
“+ 700 52.7 119.8 — 158.1
800 —56.6 —128.2 = 157.2
900 —60.3 = 136.4 ~156.3
1000 -63.85 144.2 = 155.4
1100 ~67.2 —151.8 = 154.5
1200 =105 = 159.0 = 153.6
B. Aragonite
T CaO CaCOs Reaction
fet fef Afet AG? log K K

298.15 39.7 —88.7 —164.7 129.3 —22.661 2.19 x 10-8

400 =41 =92.0 = 164.6 112.6 =14:707 1.97 x 10-8

500 44.9 =98.7 — 164.3 96.30 ~10.060 8.71 10-"

600 = 48.7 — 106.5 —163.9 80.12 = 6.975 1.06 x 10-7

Cesium -Cobalt

Very little information on Cs2COy is available.
The decomposition pressure has been measured
[88]. Decomposition becomes noticeable near
880 K (2.6 x 10-8 atm at 883 K) and reaches 0.207
atm at 1453 K.

Transitions of Cs2COs

Information on the thermal decomposition of
CoCOs is scarce. A number of basic oxides have
been reported [85], but their existence cannot be
regarded as proved. In a more recent study [30],
no intermediates were reported, but the decom-
position pressure exhibited an unexplained break
in the log p versus 1/T plot.

Phase change T(K) References in view of several density values in the 4.1-4.2
Tos in range, the value of 2.818 for spherocobaltite given
e7l in [9; Vol. I] is probably in error.
Thermodynamic data (298.15 K)
s
J mol-! deg-!_ | cal mol-! deg"! | References
Cs,COs
Cs,0
Density of CoCOs
AHP
Phase d 1K) References
k Jmol kcal mol-* References
c 4.07 116
Cs:COs -1118.s — 267.4 1, 10 ¢ 4.24 (x-ray) 48
Cs,0 —318. -75.9 10 c (spherocobaltite) | 4.214 (x-ray) 299 123



kinetics of the remainder of the reaction depends on
the rate at which the interface progresses into the
crystal.

We next: consider various mechanisms for each
of the stages in the decomposition.

(1) Nucleus Formation. Jacob and Tompkins
consider only structure-sensitive nucleation, which
occurs at definite sites in the lattice where the
activation energy is least, such as lattice defects
and dislocations. The rate of nucleus formation
thus depends both on the defect density and on the
activation energy.

Jacobs and Tompkins distinguish two cases; (1)
the decomposition of a single molecule leads to the
formation of a nucleus. The probability of this
unimolecular decomposition is

hy=y exp (AG{/RT) (2).

where y is the frequency of lattice vihrations and
AG? is the activational Gibbs energy for nucleus
formation. The rate of nucleus formation at time ¢
then is

dN/dt=k,No exp (— Aut)
or @)
N=Nol[(1—exp (—/it)]

where No is the total number of potential nucleus
forming sites. The possible loss of nucleus-forming
sites through their ingestion by growing nuclei is
neglected. In the early stages of the reaction and
especially for large AGZ and consequently small ky

N= ksNot 4)

so that the number of nuclei increases linearly with
time.

(2) A power law results from two possible mech-
anisms: (a) a stable nucleus is formed in a bimolecu-
lar combination of two active intermediates, and
(b) several decompositions are required to form a
stable nucleus. The first of these possibilities may
require the surface migration of several B mole-
cules until they coalesce into a stable nucleus.
Some aspects of this phenomenon have been dis-
cussed by MacDonald [89].

(2) Growth of Nuclei. As each nucleus is formed,
it grows generally hemispherically, i.e., it spreads
across the surface of the crystal and into its interior.
Since the laws for the rate of formation of nuclei
are different from those for their growth, the overall
decomposition rate is likely to be complicated if the
two rates are comparable in magnitude. This is
particularly true if the complications resulting from
overlapping nuclei are taken into account. As nuclei
grow, they ingest or overlap sites which would
otherwise yield nuclei. Also, as nuclei grow, they
impinge on one another and cease to grow along the

line of contact. Both of these effects result in rates
different from those predicted from models without
these refinements.

The simplest decomposition kinetics, commonly
observed for carbonates, result from a model in
which nucleation and subsequent surface growth
are extremely rapid. The rate of decomposition is
then determined by the progression of the reaction
interface into the crystal. Rate laws have been
derived for various geometries, but the one for
spherical particles is most useful for carbonate
decomposition and has been given in several equiv-
alent forms by various authors. In this derivation
it is assumed that the reaction interface proceeds
at constant velocity into the crystal, i-e., for a con-
stant interfacial area, the rate of decomposition is
a constant. For a spherical particle the reaction
proceeds along a constantly diminishing interface.
For such particles of initial radius, Ro, the fraction
decomposed at time ¢ is

4/3 R§— 4/3x(Ro— kat)? _ i-(-

« 4/303
at? , Kae
3 RS ®)
and.
da 3k. 6k3t | 3kst®
ar Ry RRR ©

These equations were first derived by Hume and
Colvin [65].

The same, but superficially different, equation
was derived by Fischbeck and Schnaidt [50] in a
form now frequently used.

da
aah (0)
where A is the total area of the reacting interface.

For a spherical particle A=s3(1—a)*, where ky
depends on Ro, and hence

Fe Ka). ®

In this form the equation is frequently referred to
as the 2/3 power law. In terms of mass change it
can be written (20)

-dn,
dt

m2ls (9)

where k again depends on Ro. Equation (8) is usually
given in integrated form

—(l-a)", (10)

The dependence of rate on particle size can be
included by the substitution kz = k2/Ro.


Decomposition of KzCOs

SER eat
Tr K.0 K.COs Reaction
fef fef fet AG} log K K
Ideg-! JIdeg-! Ay
298.15 94.1 ~152.3 348.1 ~ 60.987 1.03 x 10-8
400 915 ~152.5 332.5 — 43.420 3.80 10-4
500 = 104.2 = 152.9 317.1 33.123 7.53 x 10-4
600 -1118 =153.1 301.6 — 26.260 5.49 x 10-27
700 -119.5 —153.1 286.3 = 21.365 4:31 10-2
800 271 —153.0 271.1 17.703 1.98 10-18
900 - 134.5 ~152.7 256.1 14.864 1.37 10-8
1000 —141.7 —152.2 241.3 = 12.603 2.50 10-1
1100 —148.5 -151.7 226.7 10.763 1.73 x 10-1
Rare Earths Thermodynamic data (298.15 K)
The existence of the anhydrous carbonates at s
ambient temperatures has not been proved, al- 1-! deg |cal mol" deg’ | Ref
though hydrates are well known [126]. Such hy- J mol" deg” | cal mol deg” | References
drates decompose on heating, losing water at first, BCs ia BD 2 (anpubl,)
and then forming a series of basic carbonates, but 024, :
it is not certain whether the initial loss of water is Aue
also accompanied by some decomposition of the
carbonate ion [14]. However, anhydrous carbonates kJ mol kcal mol References
of La, Ce, Pr, Nd, Sa, Eu, and Gd, have been re- —_Rb.COs =1128. —269.6 1,10
ported by Head and Holley [64], who carefully de- Rb:0 — 330. BY 1,10
composed the hydrates by TGA and removed the
theoretical amount of H,0.
The formation of carbonates also results from .
the thermal decomposition of oxalates, but these Silver

seem to be formed as basic carbonates or in mix-
tures with oxides [58, 70, 101]. No thermodynamic
data are available.

Similarly, the existence of Pu(COs), as an inter-
mediate in the decomposition of Pu(C20,.)2 has been
reported [78].

Rubidium

Very little information is available. The melting
point listed is 38° higher than the previously ac-
cepted value. However, it has been confirmed by a
recently determined value of 1143 K [46].

Transitions of RbxCOs

Phase change T(K) Reference

1146

m

col

21

The thermal decomposition of AgsCO; is com-
plicated by the fact that the decomposition product,
AgoO, decomposes further (Ags0 > 2 Ag+1/2 Os,
for thermodynamic data see [100a]). Therefore
weight loss experiments can only be carricd out in
the relatively narrow temperature region, < 250 °C,
where AgsCO3 decomposes but Ag,O does not.
Alternatively, at higher temperatures the reaction
can be followed by absorbing CO, and Op in separate
traps and measuring their respective weight changes

[81].

The kinetics are most consistent with a model of
rapid nucleation followed by contracting sphere
kinetics [120]. However, the actual rates are ex-
tremely dependent on the method of preparation.
Kadlets and Dubinin [81] studied this effect by
preparing AgeCO; by two methods—one by pre-
cipitation from concentrated solutions which pre-
cipitated rapidly, the other by precipitation from


UNITED STATES DEPARTMENT OF COMMERCE «¢ Maurice H. Srans, Secretary
NATIONAL BUREAU OF STANDARDS ¢ Lewis M. Branscomn, Director

High Temperature Properties and

Decomposition of Inorganic Salts

Part 2. Carbonates

K. H. Stern

Electrochemistry Branch, Naval Research Laboratory
Washington, D.C. 20390

and

E. L. Weise
U.S. Patent Office, Washington, D.C. 20231

NSRDS-NBS 30

Nat. Stand. Ref. Data Ser., Nat, Bur. Stand. (U.S.), 30, 32 pages (November 1969)
CODEN: NSRDA

Issued November 1969

or sale by the Superintendent of Documents, U.S. Government Printing Office
Washington, D.C. 20402 (Order by SD Catalog No. C 13.4830), - Price 45 cents


4. Activation Energy

The activation energy E”, as defined by the usual
Arrhenius equation

k= Ae-E RT

is usually at least as great as the thermodynamic
enthalpy for the reaction, AH, though it is fre-
quently greater. When E*=AH, the reaction is
often described as occurring “without activation.”
The question as to whether E*= AH for the endo-
thermic carbonate decompositions was first con-
sidered by Fischbeck and Schnaidt [50]. In
comparing experimental values of E* and AH, they
found approximate agreement, but the scatter of
the then available data prevented rigorous com-
parison.

Shortly thereafter Zawadzki and Bretsznajder
[130] showed that for a reversible reaction‘ of the
type A(s= Bi) + Ci, the experimental value of E*
for the forward reaction, as determined from the
temperature dependence of k, must necessarily in-
crease with rising pressure and will have its mini-
mum value in the absence of the backward reaction,
i.e., in vacuum. This dependence was verified for
the decomposition of CaCO3. CdCOs3, and AgeCOQs.
In vacuum the temperature dependence of k was
the same as that of the equilibrium pressure pe,
ie., E*=AH. The dependence of E* on the pres-
sure has recently been reexamined by Pavlyu-
chenko and Prodan [104], who distinguish two
cases: (a) if E* is determined from the temperature
dependence of the rate, with the pressure p held
constant at the various temperatures, then E* in-
creases as p— pe. (b) If, however, the ratio p/pe is
held constant as the temperature is changed, then
E* will be independent of the actual numerical value
of the ratio and will have the same value as in
vacuum. This conclusion was reached by a theo-
retical argument and verified by data on the de-
composition of CdCOs. Since condition (a) is the
more usual experimentally, particularly in the older
literature, it is not surprising that widely varying
values of E* for the same reaction have been re-
ported [121] by authors who failed to recognize
the dependence of E* on pressure.

The atomic mechanism responsible for mono-
molecular reactions, including thermal decomposi-
tions, was first discussed by Polanyi and Wigner
[105]. Their model assumes that decomposition
occurs when, due to energy fluctuations in the bonds
of the molecule, the bond strength is exceeded; or
more precisely, that the bond energy “‘resides in
harmonic vibrations and that decomposition occurs
when their amplitude is exceeded.” The resulting

“For this type of reaction reversibility implies that the pressure of the gas C is the
‘equilibrium pressure P. and that infinitesimal changes in this pressure will produce the
appropriate shifts in the direction of the reaction, ie., to the left for p > Pe, and to the
tight for p <Pe. Moreover, the solid phases must be the most stable forms, rather than
metastable phases,

expression for the first-order Polanyi-Wigner rate
constant is

k=v exp (-E/RT) (22)

where v is the atomic frequency of vibration and E
is the activation energy. Thus a molecule having
energy E would leave the surface.

Shannon [118] has analyzed in some detail the
theory as it applies to the thermal decomposition
of solids. He finds that of the 31 reactions for which
he compared experimental rate constants with
those calculated from the Polyani-Wigner equation,
only a third show order-of-magnitude agreement.
In Shannon’s view, this lack of agreement stems
from neglect of rotational and other vibrational
degrees of freedom.

Shannon’s is probably the most ambitious attempt
to date to calculate a rate constant for the thermal
decomposition of a solid from absolute reaction rate
theory. The resulting equation is of the same form
as eq (22), but v is replaced by a partition function
ratio,

*

AT
ka $ exp (CE/AT) (23)

where Q* is the complete partition function for the
activated complex excluding that for the reaction
coordinate and Q is the complete partition function
for the reactant. Detailed calculations of the par-
tition functions were carried out for the decomposi-
tion of CaCO; and MgCOs, using known values of
the rotational and vibrational partition functions
and several models for the activated complex. The
best agreement with experiment was obtained from
a model in which the carbonate ion in the lattice is
free to rotate and the CO, molecule leaves directly
from the surface of the solid, rather than from a
mobile surface layer.

5. The Effect of Method of Preparation on
Rate

Various authors have‘recognized that the rate of
thermal decomposition may depend on the method
of preparation and the thermal “history” of the
sample, but systematic studies in this area are in-
herently difficult and few have been carried out.
For example, even if the effect of surface structure
on the rate of nucleation is recognized, the char-
acterization of the surface in quantitative terms
and the reproducible preparation and systematic
variation of surface and bulk properties is as yet.a
largely unsolved problem. Additional problems
arise if impurities present in the samples at very
low concentrations, such as H20, catalyze the de-
composition [108]. In that case, the resulting
change in mechanism will affect not only the rate
but also the activation energy.

If the carbonate is a single crystal, the reaction
will propagate from the surface into the interior.
In polycrystalline materials the reaction may ad-


the area. Consequently the “surface area” versus
“fraction decomposed” plot exhibits a’ maximum
before decomposition is complete. Nicholson [97]
‘has developed a mathematical model for this
process which agrees well with experimental
observations. If the purpose of the decomposition
is the preparation of an “active solid,” the reaction
should not be carried to completion in order to
achieve the desired maximum surface area.

3. The Effect of Pressure on Rate

When the pressure of CO, in a carbonate-oxide
system is equal to the equilibrium pressure, pe, no
net reaction occurs. When p< pe, the thermody-
namic driving force favors oxide formation; con-
versely, when p> pe, carbonate formation is favored.
In actual systems the favored reaction may not
occur, however, because kinetic factors prevent it.
Particularly when p is not too far from pe, the reac-
tion may not proceed because some rate-limiting
process, such as nucleus formation, is proceeding
too slowly. The resulting spurious equilibria [131]
give rise to hysteresis effects, i.e., decomposition
stops for some p<pe, recombination stops for
P> Pe. It is for this reason that we have largely
relied on thermodynamic methods for the calcula-
tion of equilibrium pressures. When the carbonate
and oxide are present as pure phases, the value of
Pe is completely unambiguous, i.e., at a given tem-
perature pe=Ky, the equilibrium constant for the
reaction MCO3;=MO+ COs. However, when the
condensed phase consists of a solid or liquid solu-
tion, pe no longer has a unique value since

a
K=2000

= and
@nco, Pe

Pe now depends on the solution composition. Since
the rate of decomposition is generally some function
of pe—p, we consider only systems without solution
formation. We have also omitted discussion of the
effect of “foreign” gases, i.e., gases other than
COs, on the rate. Such gases, e.g., O2, H2O, fre-
quently function as catalysts or inhibitors, but
these effects are difficult to relate to well-defined
properties of the system [25, 66, 68]. In order to
study the dependence of reaction rate on CO2z
pressure, it is necessary that this pressure remains
fixed throughout the course of the decomposition
while more CO; is evolving. Since the significant
value of the pressure is that at the reaction interface
it is possible that in a closed system where one
depends on diffusion to remove the excess COs, the
pressure at the interface will exceed the nominal
value, particularly if the sample size is large and
the reaction is rapid. The effect of a pressure
gradient on the reaction rate has been investigated
[20] by the application of Fick’s law. Both for slow
and for fast reactions, the rate is proportional to
Pe—po, where po is the pressure at the reaction
site but the proportionality constants differ in the
two cases. One way to avoid the problem of pressure
gradients is to use sweep gases made from a known

mixture of CO: and some inert gas. Independence
of decomposition rate from flow rate is generally
taken as evidence that the reaction is not diffusion-
controlled, e.g., [69].

If the surface area is constant, the decomposition
rate is proportional to pe—p, i.e.,

v=k(pe—p).

Such a relation was also found to hold for both
Ppineirical pellets of calcite and powdered CaCOg

Cremer and Nitsch [39], in studying the decom-
position of CaCOs, found that for samples which
followed a 2/3 rate law the pressure dependence of
the rate (in m4/%t~!) was given by

van (> a) (7)

A similar equation has been derived [69] from a
two-step model in which the CaO is assumed to be
formed first in some “active” state CaO* which
occupies a constant fraction of the surface. The
reactions are then

k
CaCO; = Ca0*+ CO,

hy

1-6 0 P

k
Ca0* = CaO
ia

both of which are reversible. The rate

_ — kokiks —Phokoks
Re—duldt= Tit kg the 18)
where the constant ko is introduced to convert the
surface area @ to surface concentration in molecules
per square centimeter, to account for the roughness
factor, and for the fact that only a portion of the
total surface sites may be potentially reactive.
Equation (18) can be written as

_1-cP,
~VP+D (19)
When P=0,  R=Ro=1/D.
When P=P., R=0 and C=1/P.
Hence
_ =P Pe.
R=BP+URy (20)
When 1/Ro > BP
1 1
R=gp-pe en

which is essentially the equation given by Cremer
and Nitsch [39].


[64] Head, E. L., and Holly, C. E., in Proc. 3d Conf. Rare
Earth Research (Gordon and Breach, New York, publ.
1964), pp. 51-63.

{és Hume, J., and Colvin, J., Phil. Mag. (7), 8, 589 (1929).

Huttig, G. F., and Herrmann, E., Kolloid Z. 92, 9 (1940).

[67] Huttig, C. F., Meller, A., and Lehmann, E., Z. Phys.

Chem. BL9, 1 (1932).

(68] Huttig, G. F.,” and Strial, K., Z. Anorg. Allgen. Chem.
209, 249 (1932).

[69] Hyatt, E. P., Cutler, I. B., and Wadsworth, M. E., J. Am.
‘Ceram. Suc. 41, 70-4 (1958).

[70] Ulner, E., Z. Chem. 6, 190-1 (1966).

[71] Ischner, B., Z. Metallkunde 55, 153 (1964).

[72] Ingraham, T. R., and Marier, P., Can. J. Chem. Eng. 41,
170-3 (1963).

[73] Jacobs, P. W. M., and Tompkins, I. C., Ch. 7 in “Chem-
istry of the Solid State,” W. E. Garner, Ed. Butter-
worths, 1955).

[74] Jaffray, J., and Martin, P., Compt. Rend. 236, 1755-6
(1953).

[75] Janz, G. J., and Lorenz, M. R., J. Electrochem. Soc. 108,
1052 (1961).

(16) Kadlete."O.,’and Dobinin, M. M., Invest. Akad. Nauk.
SSSR, Otdel Khim. Nauk, 1031 (1958).

(77) Janz, G; J. Neuenschwander, B., and Kelly, F. J., Trans.
Faraday Soc. 59, 841-5 (1963).

[78] Jenkins, I. L., and Waterman, M. J., J. Inorg. Nucl. Chem.

, 131-7 (1964).

[79] Kadlets, 0. and Dubinin, M. M., Izvest. Akad. Nauk.
SSSR. Otdel Khim. Nauk. 1031-6 (1958).

[80] Kostryukov, V. N., and Kalinkina, I. N., Zh. Fiz. Khim. 38,
780-1 (1964).

[81] Kadlets, 0., and Dubinin, M. M., Invest. Akad. Nauk.
SSSR, Otd. Khim. Nauk. 390 (1961).

[82] Kappel, H., and Huttig, G. F., Kolloid Z. 91, 117-34

(1940).

[83] Kapustinskii, A. F., and Stakhanova, M. S., Dokl. Akad.
Nauk. SSSR 57, 575 (1947).

[84] Khlapova, A. Doklady Akad. Nauk. SSSR 116,

979-82 (1957).

85] Krustinsons, J., Z. Elektrochem. 39, 193 (1933).

86] Lander, J. J., J. Chem. Phys. 17, 892-90 (1949),

87] Langmuir, I., J. Am. Chem. Soc. 38, 2263 (1916).

{88] Lebeau, P.,’Compt. Rend. 136, 1256-7 (1903); 137,
1255 2903).

[89] MacDonald, J. Y., Trans. Faraday Soc. 47, 860 (1951).

[90], MacDonald, G. J. F., Am, Mineralogist 41, 744 (1956).

{90a] Mah, A. D., US. Bur. Mines, Rept. Invest. 5600 (1960).

{91} Mah, 'A. D., U.S. Bur. Mines, Rept. Invest. 617) (1963).

[91a] Mah, A. B., Pankratz, L. B., Weller, W. W., and King,
E. C., U.S. Bur. Mines, Rept. Invest. 7026 (1967).

[92] Mampel, K., Z. Phys. Chem. A187, 43,235 (1940).

[93] Marshall, A’ L., and Bruzs, B., J. Phys. Chem. 29, 1184
(4925).

[94] Mohanty, 8. R., and Ray, M. N., Indian J. Chem. 3, 37

94a] Motil, J. Phys. Chem. 59, 139 (1955).
95] ara K., Yamagata Daigaku Kiyo, Shizen Kagaku 6,
1

(1964),

[96] Narsimhan, G. V., Chem. Eng. Sci. 16, 7 (1961).

(4 Rein 'D., Trans. Faraday Soc. 61, 990 (1965).

[98] Ostroff, A. G., and Sanderson, R., J. Inorg. Nucl. Chem.
9, 45 (1959).

[99] Otto, E., J. Electrochem. Soc. 111, 88 (1964).

[100] Otto, E., J. Electrochem. Soc. 112, 367 (1965).

(100a] Otto, E., J. Electrochem, Soc. 113, 643 (1966).

[1011 Pannetier, G., Nataf, J., and Dereigne, A., Bull. Soc.
Chim. France 1965 (2), 318-21

[102] Pauling, L., The Nature of the Chemical Bond, 3d Ed.
(Cornell Univ. Press, Ithaca, N.Y., 1960), p. 93.

ret

27

[103] Pauling, L., The Nature of the Chemical Bond (Comell
Univ. Press, Ithaca, N.Y., Third Ed., 1960), pp. 282-283.

[104] Pavlyuchenko, M. M., and Prodan, E. A., 5th International
‘Symposium on Reactivity of Solids, G. M. Schwab, ed.
(Elsevier Publ. Co., Amsterdam, The Netherlands,
1965), pp. 407-21.

05] Pola, m., and Wigner, E., Z. Phys. Chem, A139, 439

[106] Popov, M. M., and Sincbure, D. M., J. Gen. Chem. USSK
26, 1107-17 (1956).

[107] Prodan, E. A., and Pavlyuchenku, M. M., Geterogennye
Khim. Reaktsii 1961, 79-91.

[108] Prodan, E. A., and Paviyuchenko, M. M., Geterogennye
Nauk. Belorussk, SSR 1965, 64-79.

[109] Rao, C. N; R., Yoganarasimhan, S. Rand Lewis, M. P.,

[110] Rassa, M., Compt. Rend. 236, 1673-5 (1953); Ann. Chim.
(12) '8, 755 (1953).

112] Rolin, M., and Recapet, J. M., Bull. Soc. Chim. France
1964, 2504-10.

[114] Rufov, Yu. N., and Markova, Z. A., Kinetika i Kataliz 6,
731-2 (1965).

[iio] Saslawaki, IL, Z- Anots. Chem. 146; 819 0929)

[117] Sass, R. L., Vidale, R., and Donohue, j., Acta Cryst. 10,
Shannon, R. D., Trans. Faraday Soc. 60, 1902 (1964).
Simpson, D. R., Fisher, R., and Libsch, K., Am. Mineral-

[19a] Slater, J. C., Phys. Rev. 36, 57 (1930).

[120] ‘Spencer, W. F., and Topley, B., J. Chem. Soc. (London)

[121] Stern, K: H., and Weise, E. L., High-Temperature Proper-

ities of Inorganic Salts, Part 1, Sulfates, NSRDS-NBS 7
966).
[122] Subba Rao, G. V., Natarajan, M., and Rao, C.N.R.,J. Am-
Ceram. Soc. 51, 179 (1968).
patterns, Nat. Bur. Stds. Circular 539.
(a) with Fuyet, R. K., Part 2 (1953).
(c) with Gilfrich, N. T., Cook, M. L, Stinchfield, R., and
Parks, P. C., Part 8 (1959).
Part 10 (1960).
[124] Swanson, H. E., Morris, M. C., Stinchfield, R. P., and
patterns, Nat. Bur. Stds. Monograph 25, Part 1 (1962).
[125] Tranguard, A.. Lacassagne, C.,.Boinon, M.I., and Capella,
L,

[126] Vickery, R_C., Chemistry of the Lanthanons (Academic
Press, New York, 1953), pp. 217-219.

~ (1932).

[128] Wanmaker, W. L., and Radielovic, D., Sth Intern. Symp.

[129] Young, D. A., Decomposition of Solids (Pergamon Press,

New York, N.Y.

KhimReaktsii, Inst. Obshch. i Neorgan. Khim. Akad.
Can. J. Chem. 38, 2359-62 (1960).
a Reisman, A., J. Am. Chem. See. 80, 3558-61 (1958).
[113] Rose, A. J., Bull. Soc. Franc. Min. 71, 16-108 (1948).
{115} Samal, G. I., Geterogennye Khim. Reaktsii 1961, 92-8.
pe 567 (1957).
119
ogist 49, 1111-14 (1964).
1929, 2633-50.
{221a] ieee K. H., J. Chem. Educ., in press.
[123] Swanson, H. E., et al., Standard xRay diffraction powder
(b) with Cilfrich, N. T., and Cook, M. L., Part 7 (1957).
(d) with Cook, M.'L, Evans, E. A., and deGroot, J. H.,
Evans, E. H., Standard xRay diffraction powder
C. R. Acad, Sci., Paris. Ser. C264, 1111-4 (1967).
[127] Vorlander, D., Hollatz, J., and Fischer, J., Ber. 65, 536
Reactivity Solids, Munich 1964, 529-39 (Pub. 1965).
[130] Zawadzki, J., and "pretender, S., Z. Elektrochem. 41,
21 (1935).

7

131) Zara, 7 J, aa Bretsnajder, S., Trans. Faraday Soc.

[132] Zawadzki, M. J., and Szamborska, W., Bull. Intern. Acad.
Polon. Sci., ‘Gane sci, Math. Nat., Ser. A., 1940-46,
27-31 (1948).

[133] Zhabrova, G. M., and Gordeeva, V. A., Izvest. Akad.
Nauk. SSSR, Otdel Khim. Nauk. 1959, 596-9.

US. GOVERNMENT PRINTING OFFICE: 1969 OL~357-056


B. Kinetics of Thermal Decomposition

1. General

Although the literature on the thermal decomposi-
tion kinetics of carbonates is probably more exten-
sive than that of any other class of compounds,
none of the data are of “reference data’” quality, as
that term is commonly used. One reason for this
situation is that it does not yet seem to be possible
to prepare duplicate samples of any inorganic solid
salt which are identical in all the properties that
may determine the rate of decomposition, e.g., the
density of dislocations.

Nevertheless, the kinetics of decomposition are
of sufficient interest, both for theoretical and for
practical reasons, that it seems worthwhile to
include the results of at least the more carefully
done studies. In this section we summarize those
features common to carbonate decompositions and
list factors affecting the rate of decomposition. The
discussion is essentially limited to solids, since most
experimental and theoretical work has been in this
area. Kinetics of the individual compounds are
treated in their respective sections.

2. Mechanism of Endothermic Solid State
Decomposition

Many authors have derived equations for the rate
of decomposition, A(s) > B(s) +C(g), from theo-
retical models. The aim of all these studies is to find
functional relations between some measurable
quantity of the system, such as weight and time,
which can be tested against experimental results.
Although the usual caveat against taking agreement
between the predictions of the model and experi-
mental results as evidence for correctness of the
mechanism of the model applies, nevertheless such
agreement, together with other experimental evi-
dence, e.g., microscopic examination, is frequently
very helpful in at least narrowing the possibilities.

In this section some of the models which have
been proposed are briefly summarized. Greater
detail can be found in books by Garner [54] and
Young [129] and in the original references.

Most models try to account for the shape of the
experimental a(=fraction decomposed) versus
time t curves. These are generally S-shaped with a
relatively slow induction period, a steep portion
during which the bulk of the material decomposes,
and a slow final section. Experimentally determined
curves differ primarily in the relative duration of the
induction period. This induction period is commonly
identified with the formation of nuclei of the new
phase B imbedded in A; these nuclei grow, both by
spreading over the, surface and into the bulk of the
decomposing particle. After the surface is covered
by B, this phase grows into the bulk of the material.
Decomposition thus occurs only at the reactant-
product phase boundary, as has been argued by
Langmuir from phase rule considerations [87]. It
seems to be generally agreed that the evolution of
gas C during endothermic solid decomposition is

not a rate-limiting step in the process, i.e., the prod-
uct B is sufficiently porous to permit the gas to
escape readily. However, if the gas is generated in
the body of the crystal, e.g., as a result of ionizing
radiation, and must diffuse out along grain bound-
aries, the rate of decomposition may depend mark-
edly on the diffusion constant of the gas. This case
has been discussed by Gafner [53].

‘Although the measured decomposition rate
of carbonates is not rate-limited by diffusion
of the gas out of the crystal, it is clear that in
the case of readily reversible decompositions,
such as those of the carbonates, the measured
rate will depend on the difference between the
equilibrium pressure pe and the actual pressure p
at the reaction interface. Thus, although the func-
tional dependence of a on t may be independent
of pe—p, the actual rates are not. It is largely for
this reason that it is not possible to tabulate “stand-
ard reference data” rate constants, since the pres-
sure p is not usually known. Instead one finds
decompositions carried out in streams of various
inert “sweep gases” which may or may not remove
product gas from the reaction interface, or in
vacuum. In the latter case, the question as to
whether the manometrically measured pressure
equals the interface pressure needs to be answered.
Small sample size and a porous oxide favor this
condition.

(a) Kinetic Equations

The most complete classification of solid decom-
positions has been worked out by Jacobs and
Tompkins [73]. When the decomposition begins at
some points on the A lattice where the local energy
is most favorable, small fragments of B are im-
bedded in A. These may at first retain the lattice
parameters of A and will therefore possess some
strain energy as a result of the deformation. Jacobs
and Tompkins show that the Gibbs energy accom-
panying the formation of a fragment of B containing
m molecules is

AC= am — bm co)

where a is proportional to the strain energy and b is
the negative of the bulk Gibbs energy per molecule.
AG passes through a maximum at m=m* when a
fragment has the critical size to be in equilibrium
with its surroundings. Smaller fragments are un-
stable and revert to A; larger ones are stable nuclei
and grow. Whether decomposition proceeds from
the growth of a few nuclei or whether many small
nuclei are formed depends on the relative magni-
tudes ‘of the activation energies for nucleus forma-
tion (AG¥) and the growth stage (AG¥). If AG? < AGZ,
growth of existing nuclei predominates over forma-
tion of new ones. If AG¥ ~ AGZ, many small nuclei
are formed, none of which grow to visible size. In
that case the induction period is much reduced.
The whole surface then nucleates rapidly and the


Decomposition of LisCOs

LisCOxe)TM LitQat COxy
AHfng 5" 223.84kI, 53.49 keal

T LiO LisCOs Reaction
fef fef Afef AG? log K K
J deg J deg iy
298.15 90.17 =161.35 175.69 —30.780 1.66 x 10-%
400 = 94.16 —161.14 159.34 — 20.808 1.56 x 10-21
500 —102.1, — 160.65 143.47 14.988 203 10-18
600 —111.8, ~ 159.82 127.90 =11.135 7.33 x 10-%
700 122.50 ~ 158.39 112.93 8.427 3.75 X10-*
800 — 133.5; —156.50 98.60 6.438 3.65 X 10-7
900 — 144.3 — 154.67 84.60 —4.910 1.23 x 10>
1000 — 192.95 -114.76 109.04 — 5.695 2.02 x 10-*
1100 208.71 —112.42 100.13 4.755 1.76 x 10-*
1200 — 214.06 — 110.16 91.60 — 3.988 1.03 « 10-4
1300 — 224.0; — 108.00 83.40 3.351 4.46 X 10-4
1400 — 233.5. 105.96 75.46 2.815 1.53 x 10-*
1500 242.71 — 104.03 67.75, 2.359 4.37x 10
1600 — 251.50 — 102.22 60.25 —1.967 1.08 x 10-*
1700 — 299.94 — 100.50 52.94 — 1.027 2.36 X 107%
1800 — 268.04 —98.90 45.78 —1/329 4.69 10-2
1900 275.85 122.25 8.473 +0.233 71
2000 — 283.34 121.16 18.53 +0.484 3.05,
Magnesium Density of MgCOs
In vacuum, MgCO, decomposes to MgO at Phase T(K) d References
temperatures as low as 320 °C, the solids exhibiting
: . c 294.6 2.980 15
x-ray lines corresponding to MgCO; and MgO oe 5037 9, Vol. 1
only [60]. In earlier work, Centnerszwer and c (magnesite) | 298 3.009 (x-ray) 123b

Bruzs [32, 33] had reported the following series of
equilibria at 1 atm CO, (in °C).

373° 442°
MgCO; <> MgOMgCO;<—> 3
469°
MgO-MeCO; <> 4 MgO0+CO,

These equilibria are based entirely on pres-
sure measurements and weight loss data. Until
structural evidence for the existence of these com-
pounds is produced, their existence must be re-
garded as doubtful.

The decomposition follows the “rapid nuclea-
tion and contracting sphere” kinetics with an
activation energy of 150 kJ (36 kcal), somewhat
higher than the enthalpy of decomposition [29,
38, 95].

‘Thermal expansion coefficient:

MgCOs:Magnesite{

21.3 with A@20°—70°)
=33.9||; and a =5,99
with A@0°—70°)}=24.3 1

9, Vol. I

Thermodynamic data (298.15 K)

18

s°
J molt deg~| cal mol~! deg-| References
MgCO; 65.86 15.74 10b
MgO (periclase) 26.94 6.44 10b
Ane
kJ mol* keal mol-" References
MgCOs —1117 —265.7 10b
MgO (periclase) = 601.2 143.7 10b



‘tion, but even the higher of his two temperatures

is 13° lower than the value given below.

Thermodynamic data (298.15 K)

Density of StCOs s
Phase 1K) d References mol-! deg- | cal mol-! deg-! | References
SrCO: 97.1 23.2 1,5
c,strontianite 298 3.736 15 ° i
ciprecipitated salt 3.585 2 50 544 18.0 10
Transitions of StCOs np
Phase change T(K) | AHf | References KJmot! | kealmol-' | References
k}-mol
. SrCO3 = 1220.1 ~291.6 12
c,orthorhombic— c,hexagonal| 1198 17 1, 19, 86 _
Chexagonal—> cyeubie | 1770 19 S10 592.0 TLS 1
eo 1770 1
Decomposition of StCOs
SrCOx)= SrOj + COs
AHisq.1s = 234.57 kJ, 56.05 keal
T SrO SrCOs Reaction
fef fef Afef AG? log K K
J deg J deg" J deg AJ
298.15 54.4 97.1 171.0 183.7 — 32.18
400 —56.3 — 100.6 —170.8 166.4
500 59.8 — 107.6 170.4 149.5
600 — 63.8 115.7 — 169.8 132.8
700 —68.0 = 124.5 —169.2 116.2
800 -72.0 = 132.2 = 168.7 99.75
900 76.0 =140-2 —168.1 83.3;
1000 —79.7 147.5 — 167.5 67,
1100 —83.3 =155.4 —167.0 51.0;
1200 —86.7 —162.¢ — 166.2 35.20
1300 —89.9 170.3 — 164.3 21.12
1400 —93.0 179.6 162.5 7.24
1500 95.9 — 186., — 160.7 6.36
Thallium Thermodynamic data (298.15 K)
There is disagreement about the course of the s
decomposition. According to TGA studies [44] 1 dee — dee
0 ti GA > J mol~ deg~! | cal mol-! deg! | Ref
TlCOs is stable well into the liquid range, up to —— eee
645 K, and then decomposes to a basic salt. Both ThCOs 155.2 37.1 2
Gattow [55] and Rossa [110] find decomposition 70 +126. 80. 2
directly to TO, without intermediates. AHP
Density of TCOs KJ mol! keal molt | References
Phase T() d References ThCOs —700.0 — 167.3 2
TLO —178.7 = 42.7 2
© 293 | 7.91 9, Vol. I
cy monoclinic 7.24 (X-ray) 125
Transitions of TCO. .
ransitions of TlCOs Zine
Phase |T(K) | AH As References
Change
7 - dep The decomposition of ZnCOs has been studied
kJ mol-* | J mol-* dep-* 5
C,—e,1| 591 pmo 127 by several authors. Huttig, Meller, and Lehman
sI>1 546 18 36 1 [67] used natural ZnCOs, smithsonite. Bretsznaider
536 125 and Cibor [28] used synthetic material. In neither

24


Contents

Page Paes
Foreword. iit C. Suggestions for experimental work........ 10
iv Units, symbols, and abbreviations.

COz, CO, and O2..

2. General features of carbonate decom-
position...
3. Phase transitions.
4, Density...
5. Decomposition equilibri
B. Kinetics of thermal decomposition
1. General...........::0:005 ”
2. Mechanism of endothermic solid state
decomposition...
(a) Kinetic equations.
(b) Surface area changes during de-
composition...........
3. The effect of pressure on rat
4. Activation energy.
5. The effect of method of preparation
6.
7.

Thallium.
Zine...
References.

. The effect of heat transfer on rat
. The influence of crystal defect struc-
ture on the rate of decomposition...



and the actual pressure, p. Therefore, only limiting
rates (p= 0) are meaningful for the forward reaction.
Measurement should therefore be carried out at a
sufficiently low pressure for the mean free path
in the gas to be of the order of the pore size. If
an inert sweep gas is used to remove gaseous
products, it must be carefully controlled so that the
product is really removed from the reaction
interface.

(e) Since heat transfer to the sample is more
difficult in a vacuum, special care must be taken to
provide good thermal contact between sample and
heat source. The sample. should be kept from
decomposing until the desired temperature is
reached by maintaining the pressure above the
equilibrium value and then rapidly decreasing it to
zero at the beginning of the reaction.

(f) The method of preparation may affect the rate
by determining:

(1) The crystal habit. Samples may be crystalline
or amorphous. In some cases the crystal structure
may depend on the method of preparation, e.g.,
high-temperature modifications may be frozen on
by quenching the molten salt.

(2) Particle size (cf. (c) above).

(3) The nature and concentration of substances
adsorbed on the crystal surface. These may function
as catalysts or inhibitors for the decomposition.

(4) The defect structure. If decomposition starts
at dislocations on the lattice, the rate of decomposi-
tions may be anomalously high if the defect density
is high.

It is primarily the difficulty of reproducibly prepar-
ing well-characterized samples which accounts for
the present lack of meaningful rate constants. The
measurement of rate constants for solid decomposi-
tions which are as significant and reproducible as
those for gas reactions will require great care, both
in the preparation of materials and in the execution
of experiments.

Units, Symbols, and Abbreviations

=joule

hermochemical calorie= 4.1840 J

ressure; 1 atm= 101325 N m= 1013250
dyn cm=

(K)= temperature in degrees Kelvin, defined in
the thermodynamic scale by assigning 273.16
K to the triple point of water (freezing point,
273.15 K=0 °C)

=gas _constant=8.3143 J mol!

= 1.98717 cal mol-! deg

J
c
P
T

deg?

HH tandard enthalpy

AHf?= standard enthalpy of formation
S° =standard entropy

ce tandard Gibbs (free) energy

fef =free energy function= © — Mi.

TGA = thermogravimetric analysis
DTA = differential thermal analysis.

357-056 O-69—2

ll

CcO,, CO, and O,

The CO: produced by the decomposition of a
carbonate dissociates partly to CO and O2. How-
ever, at temperatures for which most of the car-
bonate decomposition equilibrium pressures have
been calculated, the decomposition of CO, is very
slight. We have therefore given the decomposition
pressures in terms of CO, only. However, for those
wishing to refine the calculations still further, we
have calculated the free energy functions and equi-
librium constant for the reaction CO2= CO + 1/2
Ong).

Entropy and enthalpy values at 298.15 K are taken
from [2]. Free energy functions above 298.15 K are
based on [10] for CO, and O» and on [8] for CO.

Thermodynamic data (298.15 K)

s
Jmol! deg-t | cal mol- deg-! | References
COz 213.64 51.06 2
co 197.56 47.219 2
02 205.03 49.003 2
AHP
kJ mol“ kcal mol? References
CcOz 393.51 94.051 2
co 110.52 26.416 2
O2 0 0 2
Dissociation of COz
COx)=COip + 1/2 Ong
AH 5= 282.98 KJ, 67.635 keal
A. Free Energies
T CO; co oO: Reaction
fel fet fet Afef Az
K J deg* J deg is]
298.15 | —213.63 205.04 257.21
400 —218.15 = 206.90 248/39
500 — 218.15 208.41 239.51
600 — 221.62 —210.93 230.60
700 = 225.14 —213.50 221-86
800 —228.81 216.02 213.07
900 232.36 218.44 204.31
1000 ~- 235.75 220.77 195.59
1100 239.03 222.99 186.90
1200 — 242.18 225.10 178:
1300 245.19 227.12 169.64
1400 — 248.10 2291 161.06
1500 250.90 — 230.89 152.52
1600 253.58 = 232.66 144.00
1700 — 256.20 — 234.36 135.52
1800 — 258.70 — 235.98 127.07
1900 = 261.11 — 237.55 118.65
2000 — 263.43 — 239.06 110.26
2100 — 265.68 —240.51 101.89
2200 — 267.86 —241.91 93.55
2300 — 269.96 — 243.22 85.23
2400 —272.01 — 244.58 76.93



Decomposition of BaCOs

Bat.O4iq)= BaQio + COng
Ans 15 269.2 KI. 64.35 keal

T BaO BaCO; Reaction
K fef fef Afef AG? log K K

50 Jdeg" Jdeg* Ideg Ly

298.15 = 70.3 =112.1 -11.8 218.0 — 38.196 0.36 x 10-2

400 72.1 —115.7 —115 200.6 6.32 x 10-2"

500 158 =122.9 -171.0 183.7 6.39 x 10-2

600 -79.9 ~131.1 —1104 167.0 2.89 x 10-18

700 — 84.2 ~139.7 169.7 150.5 5.91 x 10-

800 —88.4 M483 168.9 134.1 175 x 10-9

900 —92.4 = 156.7 —168.0 118.0 1.42. 10-7
1000 96.2 — 164.8 167.1 102.1 4.63 x 10-*
1079* 91 171.10(a) — 166.39 89.70 4.55 X 10°

171.088) | —166.37 89.72 4.54 x 10-8
1100 —99.9 13.1 165.8 86.8; 7.54 X 10-5
1200 103.4 — 182.25 163.3 73.26 6.46 x 10-*
1241" —104.7 = 185.88, 7) — 162.3 67.8; 1.39 x 10-2
1300 — 106.7 191.1 — 100.8 60.29 3.81 x 10-3
1400 ~ 109.9 — 199.6 — 158.4 414. 1.69 x 10?
1500 1129 207.7 — 156.0 35.1; 5.96 x 107
1600 —115.8 — 215.5 153.9 23.05 177x104
(1700) —118.6 = 222.9 ~151.8 111, 4.55 x 10>
(1800) ~121.3 230.0 | 150.0 —0:791 1.05
(4900) —123.9 — 236.6 — 148.4 — 12.65 2.23
Beryllium tion. The kinetics follow the contracting sphere

Normal anhydrous BeCO; probably does not
exist at ambient temperatures. Dehydration of the
hydrate at elevated temperatures probably forms
basic carbonates [83]. The value given for AHf? of
BeCOs is calculated from a measured value vf AH
for the reaction BeO+CO2=BeCO; [83] and
the current value of BeO to be published in a
continuation of [2].

Thermodynamic data (298.15 K)

3
Jmol deg-* [eal mol* deg+| References
BeCOs
BeO 4.15 3.38 2
AHF
kJ molt kcal mol! | References
BeCOs = 1025 —245 2
BeO —145.0 2
Cadmium

The rate of decomposition has been measured by
-Centnerszwer and Bruzs [34] in a CO, atmosphere
and by Prodan and Pavlyuchenko [107] as a func-
tion of CO: pressure. The decomposition yields
CdO without the formation of intermediates. In
vacuo the activation energy is 36 kcal. The rate of
decomposition depends on the method of prepara-

model [50, 65]. Experimental dissociation pressure
measurements have been reviewed by Kelley and
Anderson [3], who selected the measurements of
Andrussow [17] for their analysis. For the reaction
CdCO;=CdO+ CO; they obtain

AG*cal) = 23,200 — 10.46 T log T+8.71
x 10-°7?— 13.80 T.
The density must be regarded as very uncertain

since there are no 20th-century: measurements and
DeSchulten reports a value of 4.960 [42].

Density of CACOs
Phase | T(K) | Density} References
c 293 | 4.25—g mi ¥, Vol.

Thermodynamic data (298.15 K)

J mol-! deg-? eal mol! deg! | References
CaCO; 92.5 22.1 2
dO 54.8 13.1 2

AHF

kJ mol"! deg-1| keal mol! deg!| References
CaCO, —750.6 -179.4 2
cd —258.1 61.7 2

13


Preface

During the past decade or two, inorganic salts, particularly in the liquid state, have assumed increasing importance
in a variety of applications. They are useful as reaction media, in metallurgical processes, and in electrochemical power
sources such as fuel cells and thermal batteries.

Experimentalists and theoreticians have found molten salts an interesting subject for study since these ionic fluids
offer an unusual opportunity for the study of short-range ionic interactions in the liquid state.

So far, attention in this field has largely been focused on the alkali and alkaline earth halides since their stability at
high temperatures is well known. The only decomposition which they can undergo is dissociation to the elements. The
extent of this dissociation can be easily calculated from existing thermodynamic compilations.

For most inorganic salts the situation is more complicated. In many cases the decomposition reactions are not well
defined and high-temperature thermodynamic and kinetic data are either lacking or scattered through the literature. Thus,
although the study of many salts would undoubtedly prove interesting and useful, they have received little attention because
in many cases not even the range of thermal stability is known.

The purpose of the present publication is to alleviate this situation by publishing in concise form thermodyhamic
and kinetic data relevant to the high-temperature behavior of important classes of inorganic salts. For the present, data
in this work are restricted to anhydrous compounds with monatomic cations and oxyanions containing one element besides
oxygen. Each volume in this series will deal with compounds of one anion, carbonates in the present volume.

Thermodynamic information has been heavily stressed because it is most important for dealing with the stability
problem, and because it constitutes the bulk of what is available. Whenever possible we have used in our calculations
data already critically evaluated by others—references are given in the appropriate places—but we have also included
qualitative and semiquantitative information on the assumption that some knowledge is better than none. However, in
these cases we have tried to warn the reader by posting appropriate “Beware!” signs.

‘Thermodynamic variables which have been included are:

(a) Phase transition temperatures above 298.15 K, except those at high pressure, together with the corresponding
enthalpies and entropies.

(b) Equilibrium constants and decomposition pressures, as well as relevant free-energy functions from 298.15 K
to as high a temperature as data exist. AHf and S° values of reactants and products at 298.15 K from which
the above functions are calculated are also given.

(c) Densities at 298.15 K and above.

Since the kinetics of carbonate decomposition have been more extensively studied than those of any other class of
inorganic salts, an examination of the relevant literature has revealed many regularities as well as delineated more sharply
our arcas of ignorance. For example, although it is not yet possible to control all the factors determining rate constants,
the conditions for which activation energies are well-defined, reproducible quantities can now be stated rather precisely.
Rather than cite the bulk of the extensive literature in this field, much of which is applicable to very special conditions,
we have contented ourselves with a brief review of carbonate decomposition kinetics in which the common features rather
than the peculiarities of individual compounds are stressed. If such peculiarities are of special interest, they are described
in the appropriate section. However, our bibliography is not intended to be exhaustive. In the preparation of this monograph
we have consulted many papers not referenced here. Our aim has been to describe for the nonspecialist those features
of the decompositions which are generally agreed on, and to illustrate these with representative references of good
quality work.



very dilute solution so that the salt first formed as
a supersaturated solution. Although both prepara-
tions exhibit the same x-ray diftraction pattern, the

Transitions of AgsCOs

Phase change

second preparation decomposes much more rapidly ¢+¢ No information
than the first, an effect which the authors attribute to = c>1 Decomposes
higher defect density. Direct evid is lacking,
a higher defect density. irect evidence is lacking, Thesmodynamie data (298.15 K)
Ss
Free energy functions of AgeCOs and AgsO are
calculated from data in [4]. Jmol-! deg-! |cal mol~! deg! | References
AgsCOs 167. 40.0 2
‘AgO 121s 29 5
Density of AgsCOs auf
Phase T®) d References kJ mol-? keal mol“! | References
ce 293 | 6.077 gml-* 9, Vol. I AgsCOs —505.8 120.9 2
c 298 6.131 (x-ray) 124 AgO 31.05, 7.42 2
Decomposition of AgsCOs
AgsCOxa= ABO + COs
(OHise 581.291 kJ, 19.429 keal
Ag.O AgiCOs Reaction
fef fef Afef ACy log K K
J deg J deg-* J deg ky
(298.15 121.3 — 107.4 — 107.6 ars: — 5.487 3.26 x 10-°
350 —122.2 168.7 167.5 22.66 —3.381 4.16 x 10-4
400 —124.0 171.9 —167.3 14.38 —1.878 1.33 x 10
450 — 126.5 -176.0 ~167.0 61.4, —0.713» 1.85 x 10+
500 —129.2 — 180.7 — 166.6 —2.08, +021, 1.63
(650) — 132.0 —185.6 = 166.2 10:09 +0.958; 9:09
Sodium quite high. Consequently both the observed pres-

The literature on the phase transitions is some-
what contradictory. The transitions at 629 and 759 K
but not the one at 891° were recently observed by
Reisman, as was the melting point. Two other tran-
sitions, at 593 and 723 K were observed by Ginzburg
[56] and Popov [106], and very similar values also
by Jaffrey and Martin [74].

AHm values of [77] and [112] differ by 5 kJ.
We have listed the average value. According to
Khlapova [84], who studied the polymorphism of
NazCO; by DTA, the actual transition temperature
will vary somewhat, depending on sample prepara-
tion and treatment, in some cases by 10 to 15°.
If this is so, the solid-solid transitions would re-
quire careful reinvestigation.

Decomposition pressures were measured by
Janz and Lorenz [75], but the same comments ap-
ply to them as were made for LigCO3. The diff-
culties encountered in measuring equilibrium pres-
sures above the liquid salt by the effusion method
have been described in great detail by Motzfeld
[94a]. He showed that above the melting point
CO:z, Na, and O2 vaporize from the melt, but that
the activation energy for evaporation of CO2 is

22

sure and the steady-state composition of the melt
depend on the size of the effusion orifice.

In this paper we have not calculated the partial
pressures of the above three gaseous species, but
they can easily be calculated from the equilibrium
constants for the dissociation of the oxide given in
[10a]. and the table in this work.

Density of NazCOs
Phase T d References
c 293 2.533 g mi-! 9, Vol. I
1 11379 | 1.9685 gem-* cc)
1146.2 1.9666
1165.5 19970
11788 | 1.9508
1184.7 1.9477
1223.3 | 1.9288
12332 | 1.9264
12448 | 19211
1260.3 1.9142
1277.0 1.9080

The liquid data are fitted by
d=2.4797 — 0.4487 x 10-7 (1140 — 1280 K)


Foreword

‘The National Standard Reference Data System provides effective access to the quantitative data of physical science,
critically evaluated and compiled for convenience, and readily accessible through a variety of distribution channels. The
System was established in 1963 by action of the President’s Office of Science and Technology and the Federal Council

for Science and Technology, with responsibility to administer it assigned to the National Bureau of Standards.

The System now comprises a complex of data centers and other activities, carried on in academic institutions and other
laboratories both in and out of government. The independent operational status of existing critical data projects is main-
tained and encouraged. Data centers that are components of the NSRDS produce compilations of critically evaluated data,
critical reviews of the state of quantitative knowledge in specialized areas, and computations of useful functions derived
from standard reference data. In addition, the centers and projects establish eriteria for evaluation and compilation of
data and make recommendations on needed improvements in experimental techniques. They are normally closely associated
with active research in the relevant field.

The technical scope of the NSRDS is indicated by the principal categories of data compilation projects now active or
being planned: nuclear properties, atomic and molecular propertics, solid state propertics, thermodynamic and transport
properties, chemical kinetics, and colloid and surface properties.

The NSRDS receives advice and planning assistance from the National Research Council of the National Academy of
Sciences-National Academy of Engineering. An overall Review Committee considers the program as a whole and makes
recommendations on policy, long-term planning, and internativnal collaboration. Advisory Panels, each concerned with
a single technical area, meet regularly to examine major portions of the program, assign relative priorities, and identify
specific key problems in need of further attention. For selected specific topics, the Advisory Panels sponsor subpanels
which make detailed studies of users’ needs, the present state of knowledge, and existing data resources as a basis for
recommending one or more data compilation activities. This assembly of advisory services contributes greatly to the guidance
of NSRDS activities.

‘The NSRDS-NBS series of publications is intended primarily to include evaluated reference data and critical reviews
of long-term interest to the scientific and technical community.

Lewis M. Branscoms, Director.


case was the material pure ZnCOs. There is also
some indication that the decomposition is not en-
tirely reversible [113, 132]. It is generally agreed
that over most of the concentration range, contract-
ing sphere kinetics are followed, but the details
of the induction period are not certain. Various
values have been reported for the activation energy
and it is not yet certain whether the value at zero
pressure is equal to the enthalpy. of the reaction.

Density of ZnCOs

Phase T«) d References
© 293 | 4.44 9, Vol. I
, smithsonite 296 | 4.348 16
¢ (smithsonite) 298 4.333 (x-ray) A2z3e

Transitions of ZnCOs

Phase change T(K)

Thermodynamic data (298.15 K)

s°
Jmol“ deg~ | cal mol-! deg! | References
ZnCOs 82.4 19.7 2
ZnO 43.64 10.43 2
AHf?
kcal mol-* References
ZnCOs 194.26 2
ZnO — 83.24 2

25


ditionally propagate along grain boundaries [63].

A systematic study of the effect of sample con-
stitution on decomposition rate was carried out
by Cremer and Nitsch [40]. They decomposed four
calcite crystal samples at 850°C and CO» pressures
ranging from 120 to 220 torr. The CaO product
exhibited increasing grain size with increasing pres-
sure. The four CaO samples were then reacted
with CO, at 850 °C and 420 torr to produce four
CaCOs samples with similarly increasing grain size.
When these four samples were then decomposed
under identical conditions (850 °C, 200 torr CO,),
the rate of decomposition increased in inverse
order of crystallite size. The disappearance of the
induction period with decreasing crystal size was
particularly noticeable. These results are explained
in terms of more rapid nucleus formation at lower
pressures, leading to smaller CaO crystals and
consequently a larger number of active sites in the
CaCO; subsequently formed.

It is evident from these results that the course
of the decomposition is affected by the microstruc-
ture of the sample and that the reproducible prep-
aration and characterization of nominally identical
materials is an important prerequisite for the under-
standing and systematization of decomposition
kinetics.

6. The Effect of Heat Transfer on Rate

An endothermic solid-state decomposition pro-
ceeds as the produce-reactant interface advances
into the interior of the sample. The interface can,
however, advance only if the necessary heat of re-
action is supplied to it. Since the source of heat is
outside the sample, the rate of heat transfer to the
interface may become rate determining if the in-
herent rate of reaction is greater than the rate of
heat transport. The rate at which heat is transported
depends not only on the properties of the product
through which heat must be transported, but also
on the general experimental arrangement. It is for
this reason that one expects and finds the literature
to be conflicting.

For example, Narsimhan [96] has derived an equa-
tion for the rate of thermal decomposition based
on the idea that the entire particle rapidly reaches
the decomposition temperature so that there are no
internal temperature gradients. All of the heat
reaching the reaction interface is used up in the
decomposition so that the reaction rate depends on
the rate of heat transport. For both spherical and
cvylindiival particles, good agreement with some
literature values is found. Under these conditions
the chemical steps are no longer rate-determining.

It can thus be seen that in any solid-state reaction
whose enthalphy is appreciable, the establishment,
maintenance, and measurement of the temperature
at the reaction interface is a complicated problem.
Thus, careful studies of the decomposition of
PbCOs [76] and several other inorganic salts [133]
have shown that the induction period may disappear
if the entire sample is rapidly brought to the reaction

temperature. There is therefore now some question
whether the commonly reported induction periods
in the decomposition of so many substances are
artifacts of the experimental arrangement. Con-
siderable care in the experimental design is certainly
required to ensure a constant and known tempera-
ture at the reaction interface throughout the course
of the decomposition.

7. The Influence of Crystal Defect Structure
on the Rate of Decomposition

The idea that the defect structure of a solid
reactant affects its rate of decomposition seems
to be generally accepted; but of all the factors
influencing the kinetics, this one is the most diffi-
cult to characterize quantitatively. The literature
on the subject has been reviewed by Boldyrev
[26], who then classified decompositions according,
to which aspects of the defect structure needed to
be considered in particular cases. Because of its
general interest, his classification is reproduced
here.

Effect of crystal defects on the rate of thermal decomposition
of solids (26)

Decomposition proceeds through
Breaking of bonds within
cationic and anionic
Crystal defects lattice constituents Electron
affecting the rate transfer
Reversible | Irreversible | from anion
decompo- | decompo- | t cation
sition sition
Change of habit
Growth figures, vicinals| } Effect
macro-fissures........|
Dislocations, and Effect
groups of them........ Effect
Impurity inclusions No
in lattice ... effect
onic defects No
Electronic defects effect

According to this scheme, carbonate decomposi-
tion constitutes the simplest type, since it is a re-
versible reaction in which only intra-ionic bonds
are broken. Therefore, only the first two types of
defects affect the rate of decomposition.

The main distinction between reversible and ir-
reversible decompositions is the more complex
mechanism of the latter, whereas reversible de-
composition proceeds through only onc stcp, such
as the breaking of a C—O bond in the CO; ion.
Since these reactions start on the surface, a change
in the crystal habit which affects the surface area
ratio of more/less reactive faces will necessarily
affect the rate, particularly since such a change of
crystal habit also changes the ratio of surface/bulk
ions. Growth figures and macrofissures have their
greatest effect on the initial rate since nucleation
commonly begins at the most reactive surface sites
and these are usually defects in the structure.


B. Equilibrium constants

T(K) log K K
298.15, — 45.062 8.67 10-*
400 — 32.427 3.74 10-88
500 25.021 8.53 x 10-2"
600 —20.082 8.28 x 10-1
700 2.79 x 10-87
800 1.22% 10-14
900 1:39 10"
1000 6.08 x 10-"
1100 1.33 x 10-®
1200 1.74 x 10-*
1300 1.53 10-7
1400 9.79 x 10-7
1500 4:89 x 10-8
1000 1.99 x 107°
1700 6.85 x 10-5
1800 - 2.06 x 10-4
1900 - 5.41 10-4
2000 - 1.32 x 10-8
2100 - 2:92 x 10-3
2200 - 6.01 x 10-
2300 - 1.16 x 10-2
2400 - 2.12 10-*

Transitions of BaCOs

Barium

The effect of the BaCOs transitions at 1079 and
1241 K on the calculated equilibrium pressure is
virtually nil. Calculations above 1600 K were made
by extrapolating the fef values of the individual
compounds.

The solid-solid transitions of BaCOs3 exhibit
pronounced hysteresis, i.e., different temperatures
are obtained, depending on whether the transition
is approached from above or below [86]. The melting
point can only be measured under high CO:
pressure. AH and AS for the transitions are taken
from (1).

Deusity oy BaCOs

Phase T(K) d References
c 293 4.287 15,
c(witherite) 299 4,308 (x-ray) 123a

c (cubic) 1348 3.889 (x-ray) 123¢

Phase change T(K) AH As References
kT mol“! | J molt deg-*
c.orthorhombic (y)— c,hexagonal (8) tore}
corthorhombie (y) <c hexagonal (8) 1023, 149 137 1, 8
c,hexagonal (8)—>c,cubic (a) 120} 3 4,86
c,hexagonal (8) <c,cubie (a) 1228) " ,
ela) 1 1653 T

Thermodynamic date (298.15 K)

5°
J mol! deg-! | cal mol~! deg-'} References
BaCO,(, II,
witherite) 12.1 26.8 1,10
BaO 70.3 16.8 1,5,6
KJ mol-? keal mol“! | References
BaCO, —244.7 12
BaO —582.0 o1

12