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QC100 .11573 V7;1;1966 C.1 NBS-PUB-C 1964

/?

NSRDS-NBS 7

PUBLICATIONS High Temperature Properties

and Decomposition of

Inorganic Salts

Part 1. Sulfates

U.S. DEPARTMENT OF COMMERCE

NATIONAL BUREAU OF STANDARDS
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UNITED STATES DEPARTMENT OF COMMERCE • John T. Connor, Secretary

NATIONAL BUREAU OF STANDARDS • A. V. Astin, Director

High Temperature Properties and

Decomposition of Inorganic Salts

Part 1. Sulfates

K. H. Stern and E. L. Weise

Institute for Basic Standards

National Bureau of Standards

Washington, D.C.

Sr**,

NSRDS

"CE D ^

NSRDS-NBS 7

National Standard Reference Data SeriesNational Bureau of Standards 7

(Category 5 — Thermodynamic and Transport Properties)

Issued October 1, 1966

For sale by the Superintendent of Documents, U.S. Government Printing Office Washington, D. C. 20402 — Price 35 cents
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Foreword

The National Standard Reference Data System is a government-wide effort to give to the

technical community of the United States optimum access to the quantitative data of physical

science, critically evaluated and compiled for convenience. This program was established in

1963 by the President’s Office of Science and Technology, acting upon the recommendation of the

Federal Council for Science and Technology. The National Bureau of Standards has been as- signed responsibility for administering the effort. The general objective of the System is to coordinate and integrate existing data evaluation and compilation activities into a systematic, comprehensive program, supplementing and expanding technical coverage when necessary, establishing and maintaining standards for the output of the participating groups, and providing mechanisms for the dissemination of the output as required.

The NSRDS is conducted as a decentralized operation of nation-wide scope with central co- ordination by NBS. It comprises a complex of data centers and other activities, carried on in

government agencies, academic institutions, and nongovernmental laboratories. The independent

operational status of existing critical data projects is maintained and encouraged. Data centers

that are components of the NSRDS produce compilations of critically evaluated data, critical re- views of the state of quantitative knowledge in specialized areas, and computations of useful

functions derived from standard reference data.

For operational purposes, NSRDS compilation activities are organized into seven categories

as listed below. The data publications of the NSRDS, which may consist of monographs, looseleaf sheets, computer tapes, or any other useful product, will be classified as belonging to one

or another of these categories. An additional “General” category of NSRDS publications will

include reports on detailed classification schemes, lists of compilations considered to be Standard

Reference Data, status reports, and similar material. Thus, NSRDS publications will appear in

the following eight categories:

Category

12

3

4

5

6

7

8

Title

General

Nuclear Properties

Atomic and Molecular Properties

Solid State Properties

Thermodynamic and Transport Properties

Chemical Kinetics

Colloid and Surface Properties

Mechanical Properties of Materials

The present compilation is in category 5 of the above list. It constitutes the seventh publication in the new NBS series known as the National Standard Reference Data Series.

A. V. ASTIN, Director.

Ill
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 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 un- doubtedly 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 thermodynamic 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, sulfates 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 — ref- erences 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.

Kinetic data which have been included are

(a) Qualitative and quantitative information on the kinetics of thermal decomposition.

For the convenience of the user a number of special tables are also included; for example,

a listing of compounds in the order of increasing melting points and decomposition pressures

which may prove useful is given.

IV
Contents

Foreword

Preface The sulfates A. Introduction

B. Phase transitions C. Density

D. Decomposition equilibria E. Kinetics of thermal decomposition

Units, symbols, and abbreviations SO3 , SO2 , and O2 Aluminum

Americium

Antimony

Barium

Beryllium

Bismuth

Cadmium

Calcium

Cerium

Cesium

Chromium

Cobalt Copper

Dysprosium

Erbium

Europium

Gadolinium

Galhum

Germanium

Hafnium

Holmium

Indium

Iridium

Iron Lanthanum

Lead

Lithium

Lutetium

Magnesium 20 Manganese 20 Mercury 22 Neodymium 23

Nickel 23 Palladium 23 Plutonium 24 Polonium 24 Potassium 24 Praseodymium 25 Radium 25 Rare Earths 25 Rhodium 26 Rubidium 26 Ruthenium 27 Samarium 27 Scandium 27

Silver 27 Sodium 27 Strontium 29 Terbium 29 Thallium 30 Thorium 30 Thulium 31 Tin 31 Titanium 32 Uranium 32 Vanadium 33 Ytterbium 33 Yttrium 33 Zinc 34 Zirconium 35

Relative thermal stability of some sulfates 35 Arrhenius activation energies for the thermal dissociation of some inorganic sulfates 36

Melting points in increasing order 36

Boiling points 36 References 37

Page

III IV

1

1

12

234

4

56

6

6

789

9

10

11

1111121414

14

14

14

15

15

15

1516

16

17

18

1920

V

High Temperature Properties and Decomposition of Inorganic Salts

Part 1. Sulfates

K. H. Stern and E. L. Weise

The literature dealing with the high-temperature behavior of inorganic sulfates 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, equilibrium constants of reactions, and partial pressures of gaseous components were

tabulated. Auxiliary data on phase transitions, densities, and kinetics of chemical decomposition

have also been included. Key words: Sulfates, thermodynamic functions, decomposition pressures The Sulfates

A. Introduction

In this section the general features of sulfate

decomposition are summarized. In all metallic

sulfates the SO4 ion is a well defined entity in which

the oxygen atoms are arranged tetrahedrally about

the sulfur atom with an S—O bond length of 1.51 A,

0.19 A less than the sum of the single-bond covalent

radii of the atoms. Pauling [96c] 1 has fisted the

various resonance structures which give these

bonds some double-bond character. The tetra- hedral structure of the sulfate ion is distorted to some extent by cations, the distortion increasing

with the polarizing power of the cation. Since the

polarizing power increases with increasing electronegativity and decreasing ionic radius, salts with

small, highly charged cations would be expected to be least stable. For example, Ostroff and Sander- son [90] found that the decomposition temperature

for eleven anhydrous sulfates of divalent cations is a linear function of (r/S) 1

, where r is the crystal

(Goldschmidt) radius and S is the electronegativity.

Similar observations have been made for the alkali

metal sulfates [125].

Few metal sulfates are sufficiently stable at high

temperatures to have a measurable melting point.

All decompose eventually with evolution of SO3 . The latter in turn dissociates to SO2 and O2 . Thus

in a closed system there are always two simultane- ous equilibria:

mso4 ^mo+so3 so3 ^

so2 + 1/202

Equilibrium data for the second reaction are given

in separate tables for calculation of partial pressures

of the 3 gases, and of the total decomposition pressure. The first reaction is only schematically repre1 Figures in brackets indicate the literature references at the end of this paper.

sented above. Although many sulfates do decompose

directly to the oxide containing the metallic element

with unchanged valence, some show oxidation, e.g., FeSCX-»Fe2 C>3 , SnSC^—^SnC^, MnSCb-*Mn3 04 or reduction, e.g., AgzSfX -*' Ag°. Some decompose

through an intermediate oxysulfate or basic salt such as CuO • CuS04 , or even a sequence of intermediates. It is primarily because of these complications that a brief but critical review has been

included at the beginning of each section, describing

the high-temperature behavior of each sulfate.

The literature has been searched through mid1964, and some later references have been included.

The bibliography, beginning with reference 11,

is arranged alphabetically by author. The refer- ences preceding [11] are those critical compilations that have been used here most frequently.

Original publications, the data of which have been

reworked, presented, and referenced in the first ten compilations are not usually fisted separately

in the bibliography.

B. Phase Transitions

1. 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 differ so much that it is not clear whether the same or a dif- ferent transition is being described.

A number of well-defined transitions have been

fisted in NBS Circular 500 [1]. The data on these

and others have been collected and are fisted.

2. Melting Points. Melting points have been

reported for several of the more common sulfates.

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.

1
Melting points are usually obtained by rapid heating of the salt to minimize decomposition, but it cannot always be avoided. The listed values should

therefore be regarded as lower limits.

3. Boiling Points. The dissociation pressures

of sulfates rise so rapidly above the melting point

that measurement of the boiling point is impossible. We have listed only those values calculated

for the alkali metal salts from the law of corresponding states by Lorentz and Hertz [70].

C, Density

The total literature of high-temperature density

measurements <5n inorganic sulfates is very small.

There exists only one set of reliable values for the

liquid salts, those of Jaeger [59] on the alkali metal

sulfates. We have reproduced his original values

without attempting to fit them to an equation. One

reason for the lack of data is that most sulfates

decompose considerably below the melting point

so that a measurement of any physical property

of the pure liquid is inherently impossible.

Most density values of solid sulfates have been

reported for ambient temperatures. Except for a very limited study of PbS04 [80] there appears

to have been no study of the temperature depend- ence for any solid sulfate.

The densities of solid salts are generally determined 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 generally higher than experimental values by as much as 10 percent. Since most crystallographers do not indicate the temperature for which

the density is calculated, we have preferentially

listed experimental values when these were available. However, in some cases experimental values

seem to depend on the method of preparation, for

some unknown reason. 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 significance to figures

beyond the third. This is particularly true because

the units of the reported density values are fre- quently in doubt. In many cases they are not

stated at all; in others it is not clear whether the

unit is g/cm3 or g/ml; in still others they are specific gravity values and may be referred to water

at 4 °C or at some other temperature. The dif- ference 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, respectively)

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.

D. Decomposition Equilibria

As a measure of the stability of each sulfate, we

have calculated wherever possible, thermodynamically, the partial pressures, and total pressure,

of the SO3, SO2, and O2 decomposition products

that would be expected at equilibrium at various

temperatures in a closed system containing in- itially only the sulfate. This is the type of system

in which decomposition pressure measurements

are made by the static method.

For these calculations, four items of information

are needed:

(1) Knowledge of the course of the decomposition. This is described under each sulfate. Where specific references are not given, the description is an abstract from such sources as the handbooks of inorganic chemistry of Gmelin, Pascal,

and Mellor.

(2) High-temperature thermal data such as those

tabulated in reference [2] 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 sulfates themselves.

Items (3) and (4) are tabulated as “Thermodynamic

Data” under each sulfate to show the data used

and to point out the gaps in existing data.

It has usually been most convenient to perform

the calculations through the medium of “free

energy functions” (fef) as discussed in reference [4], pp. 166-9. Where the high-temperature thermal

data are available in the form of enthalpy and

entropy increments (Ht — H^qs and St~ Sms), these

are combined with 298.15 °K entropies to get the

free energy functions:

fef : fO I/O IT o no

Lrr CZ 298 n T [1 298 (St -^ 298 ) C°

*^298

where T is the temperature of interest in °K. Then

for a given T°K, the fefs 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 free energy change attending the reaction is

AG£= TAfef +

from which the equilibrium constant for the reaction

is obtained through the relation:

AGt =— RTlnK.

For the case of simplest stoichiometry:

MS04 ^ MO + S03 K= PSo

3 2
the activity, here taken as equal to partial pressure,

of SO3 in atm is equal to K. SO2 and O2 appear by

decomposition of SO3 :

Pso.Pl'2 S03 5*S02 + V202 K' = —^—-• PTM3 When both equilibria are attained simultaneously:

Pso2 + Po2 = ( 3

/2)(2) 1/3 (KK'Y13

and

PS02 = 2/z{PS02+ Po2

) and

Po2 = y%{Pso2 + po2

) since

Pso2 — 2Por

This is true only in an initially evacuated system

where all the gases originate from the MSO4 . When PSo3 is very much less than Pso2

, the reaction

is often written in the literature as:

MSO4 MO + SO2 I/2O2

for which the equilibrium constant is KK'

If the metallic element undergoes oxidation during the decomposition, equilibrium constants are obtained through free energy functions as before,

but calculation of partial pressures from them is more complicated because the relation, PS02 = 2P02 , no longer holds. The general method is shown for

the case of MnSfX —» Mn3 04 in the section on MnS04 . In the tables, these functions, constants, and

pressures are given at 100 degree intervals. For

other temperatures interpolation will be most ac- curate with fef s, calculating pressures from them.

For the three gases, the high-temperature thermal

data are derived from spectroscopic observations,

and are available to higher temperatures than

needed here. But for the solids, the data are obtained from calorimetric measurements which must

be discontinued either at the limit of the apparatus

or at the temperature at which decomposition becomes noticeable during the time of the experiment.

Thus the calculated pressures would not exceed

about 10~3 atm, or less. We have therefore extrapolated, at the fef level, for several hundred degrees

or to about 1 atm total decomposition pressure.

To go beyond 1 atm would involve fugacities; cal- culation of these would involve uncertain extrapolations and is therefore not justified. Extrapolating more than a few hundred degrees becomes

increasingly dangerous, particularly because un- reported phase transitions might change the slope

of a fef versus T plot.

E. Kinetics of Thermal Decomposition

By far the largest number of sulfate decomposition studies are concerned with practical applications, usually with the liberation of S03 and SO2 from naturally occurring sulfates, e.g., CaS04 and

Na2 S04 , at the lowest possible temperature. Since

certain metal oxides are catalysts for the decomposition, experiments are commonly carried out with

these as additives, i.e., with impure materials;

moreover, the experimental conditions are rarely

described adequately. This latter problem also

afflicts several of the studies on pure materials.

Kinetic studies of sulfate decomposition can be

conveniently divided into isothermal and nonisothermal ones. Chief among the latter are thermogravimetric analysis (TGA) and differential thermal

analysis (DTA). TGA is particularly useful for

determining the stability ranges of compounds

whose knowledge is necessary for analytical purposes, but gives relatively little information con- cerning the kinetics of the decomposition. DTA

is useful in determining phase transition temperatures and can be used to determine enthalpies

and activation energies. The accuracy with which

these quantities can be obtained is still rather poor,

however [83].

Since the rate of sulfate decomposition under

isothermal conditions depends on a rather large

number of experimental conditions, these must be

carefully controlled. Thus, for the specific case

of a metal sulfate having only a single stable valence

state which decomposes directly to the oxide — the

simplest possible case — the following two reactions

must be considered:

MS04 = M0 + S03 (1) S03 = S02 + 1/2 02 . (2) Assuming that the oxide formed in (1) is not a

catalyst for the decomposition the following variables affect the reaction rate: (a) temperature, (b)

gas phase composition; this may be variable if, for

example, the system is initially pumped out and the

rate followed by measuring the pressure change,

or it may be kept constant by passing a gas over

the decomposing sample. In the latter case the

reaction rate may depend on (c) gas composition —

inert, or some mixture involving the components

of the system, and (d) flow rate; (e) since the re- action occurs at a solid-solid interface, sample

size and shape may be significant. However, most

studies indicate that the oxide formed in (1) is sufficiently porous to the gases that the rate cal- culated for a unit area of interface is independent

of depth of penetration.

In the description of the decomposition kinetics

of the individual sulfates we discuss only those

studies, rather few in number, which meet at least

216-963 0- 66—2 3
some of the criteria outlined above. In any case, however, these discussions are to be regarded

only as guides to the behavior to be expected from

a given compound. Any numerical values given

are to be used in this sense and certainly do not

have the validity attached to thermodynamic

parameters.

To illustrate the present state of knowledge in

this field, we have summarized in a separate table

all the activation energies for decomposition which

are discussed under the individual compounds.

It is clear that the variety of results obtained, in- cluding the great differences in Ea obtained by

different investigators for the same reaction,

makes any discussion of these results premature.

They are included only to stimulate further work

in this field. Units, Symbols, and Abbreviations

J = joule

cal = thermochemical calorie =4.1840 J P = pressure; 1 atm = 101325 N m

-2 = 1013250

dyn cm-2 T{°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)

R = gas constant = 8.3143 J mol-1 deg-1 = 1.98717 cal mol-1 deg-1 H° = standard enthalpy

AH

f°= standard heat of formation

S° = standard entropy

G° = standard Gibbs (free) energy

G

O II o T 298.15 tet =lree energy tunction = — TGA = thermogravimetric analysis

DTA = differential thermal analysis.

S03 , S02 , and 02 The SO3 produced on decomposition of a sulfate

dissociates partially to S02 and 02 . Data for this

secondary equilibrium are presented in tables A,

B, and C.

The free energy functions (fef) for SO3, S02 , and

02 (table A) are calculated from data in reference

[6] using entropies from [6]. The entropies for

these gases tabulated in [6] and [9] supersede

those in [1] and [3], The heats of formation for these gases in [6] are based on those in [1], differing by only 0.01 kcal/mol

for S02 and 0.02 for S03 . They are part of a set of enthalpies consistent with those of the sulfates

and oxides in [1]. The equilibrium constants of

table B, calculated from these enthalpies of the

gases, are used later with metal sulfate data whereever the latter are taken from [1]. Reference [9], the first part of a revision of [1], has a small, but significant change in the heat of

formation of S03 . Otherwise the data there for

these gases differ only slightly from the data of [6]

as shown in the table immediately below. Table

C presents a new set of equilibrium values incorporating the effect of these changes. In the future

table C should be used for calculations with the

forthcoming revised data for sulfates. Table C

should also be used instead of table B for any independent purpose.

Thermodynamic Data (298.15 °K

)

s

D

J mol-1 cal mol-1 Refer- deg-1 deg- 1 ences

so3 256.6671 61.3449 6 256.65 61.34 9 so2 248.0974 59.2967 6

248.1, 59.30 9, 10 o2 205.05 49.0070 6

205.00 48.996 9

AHf°

kj mol-1 kcal mol-1 Refer- ences

S03 — 395.

2

6 -94.47 6 -395.72 -94.58 9 so2 -296.8.5 -70.947 6, 10 -296.83 - 70.944 9 o2 0 0

Dissociation of SO3 S03(g) «=* S02(g)+ 1/2 02 (g) AW^g 15 = 98.408 kj, 23.52 kcal, reference [6] data 98.893 kj, 23.636 kcal, reference [9] data.

A. Free energies, reference [6] data

T S03 S02 o2 Reaction

fef fef fef Afef AG°

°K J J J J kj

298.15 -256.67 -248.10 -205.05 -93.956 + 70.37 400 -258.75 -249.71 -206.20 -94.065 60.79 500 -262.89 -252.86 -208.42 -94.182 51.30 600 -267.76 -256.52 -210.94 -94.232 41.88 700 -272.84 -260.31 -213.51 -94.219 32.47 800 -277.91 -264.06 -216.02 -94.157 23.10 900 -282.87 -267.70 -218.45 - 94.056 13.77 1000 -287.67 -271.22 -220.77 -93.935 + 4.48 1100 -292.28 -274.58 -222.99 -93.801 -4.77

1200 -296.71 -277.81 -225.11 -93.650 -13.97

1300 - 300.96 -280.90 -227.13 -93.504 -23.14

1400 -305.03 -283.86 -229.06 -93.353 -32.30

1500 - 308.94 -286.69 -230.90 -93.199 -41.38

1600 -312.69 -289.40 -232.67 -93.048 -50.46

1700 -316.30 -292.01 -234.36 -92.893 -59.50

1800 -319.77 -294.51 -235.99 -92.743 -68.53

1900 -323.11 -296.93 -237.55 -92.600 -77.53

2000 -326.32 -299.25 -239.07 - 92.458 -86.53

4
B. Equilibrium constants, reference [6] data

T log K K

°K

298.15 - 12.329 4.69 X 10- 13 400 -7.938 1.154 x 10~ 8 500 -5.358 4.39 X10-6 600 -3.646 2.26 XlO-4 700 -2.423 3.78 X IO-3 800 -1.508 3.11 XlO"2 900 -0.798 0.1592 1000 -.234 .583 1100 + .226 I.683 1200 .608 4.06 1300 .930 8.51 1400 1.205 16. Cb 1500 1.441 27.6 1600 1.647 44.4 1700 1.828 67.3 1800 1.989 97.5 1900 2.131 135.2 2000 2.260 182

C. Free energies and equilibrium constants, reference [9] data

T Afef AG°t log a: K

°K J kj

298.15 -93.964 + 70.88 -12.417 3.83 XlO-13 400 - 94.073 61.26 -7.999 1.00 x 10- 8 500 -94.190 51.80 -5.411 3.88 x 10- 6 600 -94.240 42.35 -3.687 2.06 x 10-4 700 -94.228 32.93 -2.457 3.49 x 10- 3 800 -94.165 23.56 -1.538 2.90 x lO- 2 900 -94.065 14.23 -0.826 1.49 x 10- 1 1000 -93.943 + 4.95 -.259 5.51 x 10- 1 1100 -93.809 -4.30 + .204 1.60 1200 -93.659 -13.50 • 5875 3.87 1300 -93.512 -22.67 .911 8.15 1400 -93.362 -31.82 1.187 15.4 1500 -93.207 -40.92 1.425 26.6 1600 -93.056 -50.00 1.632 42.9 1700 -92.902 -59.04 1.814 65.2 1800 -92.751 -68.06 1.975 94.4 1900 -92.609 -77.07 2.119 131.5 2000 -92.466 -86.04 2.247 177

Aluminum

The kinetics of the decomposition have been

studied by Pechkovsky and Ketov [102] and Warner

and Ingraham [138]. No intermediate oxysulfate

has been reported, the decomposition proceeding

directly to the oxide

Al2 (S 04)3 = AI2 O3 T- 3 SO2 + 3/2 O2 .

Between 650 and 950 °C the rate of decomposition,

as measured by weight loss, is independent of

flow rate of an inert gas, indicating that the product layer is porous to the flow of gases. Under

these conditions k= 1.14 X 10 10 g cm-2 min -1exp

(—268,000/^7*). With increasing oxygen content

of the flow gas the reaction rate decreases somewhat, but the effect is not very great. Between 21

and 100 percent O2 the rate is virtually independent

of oxygen content.

Decomposition pressures reported in the litera- ture do not represent true equilibria. Kelley

[1946] and Warner and Ingraham [137] noted that

y — AI2O3 is formed initially, which changes slowly

to the stable a—AI2O3.

Free energy functions for a — AI2O3 and AI2

(S04)3 are calculated from data in [2]. Those

calculated from data in [6] are only slightly different.

Density of AhiSOi)^

Phase T°K d Refer- ences

c ambient 2.71 89

Transitions of AhiSOffc

Phase

change

c —> c None known, 298-1100 °K

c —> 1 . Decomposes

Thermodynamic Data (298.15 °K)

S°

J mol-1 cal mol-1 Refer- deg-1 deg-1 ences

A12 (S04 )3 239 57.2 1,3,8 a —

12 03 50.96 12.18 3, 4, 6, 8

AHf°

kj mol-1 kcal mol 1 Refer- ences

A12 (S04 )3 -3435 -821.0 1,8 a — A12 03 - 1669.8 -399.09 1

5
Decomposition of Al2 (SOfh Antimony

(1) Al2 (S04 yc)^ Al2 03 (a) + 3 S03(g) 579.6* kj, 138.5 kcal

(2) Al2(S04 )3(c)5=i Al2 03(a) + 3 S02 (g) + 3/2 02 (g) A//|91i . 15 = 874.8T kj, 209.1 kcal.

A. Free energies

T AI2O3 AblSO-ib Reaction (1)

fef fef Afef AG°t

°K J J J kj

298.15 -50.96 -239.32 -581.64 + 406.22 400 - 54.29 -250.54 -579.99 347.65 500 -61.13 -273.26 -576.54 291.37 600 - 69.20 -300.12 -572.36 236.23 700 - 11.65-0 -328.15 -568.03 182.05 800 -86.15 -356.31 -563.58 128.78 900 -94.35 -383.71 -559.25 76.32 1000 - 102.34 -410.20 -555.14 + 24.51 1100 -110.00 -435.60 -551.25 -26.72

B. Equilibrium constants

T log Ki log K2 k2

°K

298.15 -71.165 6.84 x IO

- 72 -108.152 7.05 x 10- 109 400 - 45.394 4.04 x 10-46 -69.208 6.19 x IO

-70 500 -30.437 3.66 x IO

- 31 -46.511 3.08 x 10“47 600 - 20.564 2.73 x 10-21 -31.502 3.15 x IO

-32 700 - 13.583 2.61 x 10- 14 -20.852 1.41 x 10-21 800 - 8.408 3.91 x 10-9 - 12.932 1.17 x 10- 13 900 -4.429 3.72 x 10-5 -6.826 1.49 x 10- 7 1000 - 1.280 5.25 x 10-2 -1.982 1.04 x 10-2 1100 + 1.269 18.6 + 1.947 88.5

c. Equilibrium pressures (atm)

T S02 02 S03 Total

°K

298.15 1.17 x 10-24 5.83 x 10-25 1.90 X 10-24 3.65 x 10"24 400 5.26 x 10- 16 2.63 x lO" 16 7.39 X lO" 16 1.53 x 10- 15 500 5.82 X 10- 11 2.91 x 10- 11 7.15 X 10~ n 1.59 x 10-'° 600 1.26 x 10-7 6.29 x 10-8 1.40 x 10-7 3.29 x 10-7 700 2.93 x 10~5 1.46 x 10-5 2.97 X 10-5 7.36 x 10“5 800 1.69 X 10-3 8.42 X 10"4 1.58 x 10-3 4.10 x HP3 900 5.50 X 10"2 2.75 x 10-2 5.74 X lO"2 1.40 x 10- 1 1000 4.57 X 10- 1 2.29 X10- 1 3.74 x 10- 1 1.06 1100 3.41 1.71 2.65 7.77

Americium

Am (ill): Arri2 (S 04)3 does not melt, but decomposes with oxidation to Am02 . A TGA curve

[50a] showed a plateau for the anhydrous salt ex- tending to 990 °K. Conversion to Am02 was

complete by 1070 °K.

On heating, Sb2 (S 04)3 decomposes with partial

oxidation: oxysulfates exist, about which little

is known.

There are no thermodynamic data above 298 °K.

Thermodynamic data (298.15 °K)

s°

J mol-1 deg-1 cal mol-1 deg-1 Refer- ences

Sb2 (S04)3 Sb2 04 127.2 30.4 9, 3, 8 Sb2 05 125. 29.9 9, 3, 8 Sb4 C>6 (II, cubic) 220.9 52.8 9 Sb4 C>6 (I, orthorhombic) 246.o 58.8 9, 3, 8

A///°

kj mol-1 kcal mol-1 References

Sb2 (S04 )3 -2402.5 -574.2 9 Sb2 04 -907.5 -216.9 9 Sb2 Os -971.9 -232.3 9 Sb4 C>6 (II, cubic) - 1440.5 -344.3 9

Sb4 06 (I, orthorhombic) -1417.1 -338.7 9

Barium

BaS04 melts without obvious decomposition.

Free energy functions for BaS04 and BaO are cal- culated from data in [2].

Density of BaS04

Phase T d References

°K

c, II 288 4.499 7

Transitions o/BaS04

Phase change T AH AS References

°K kj mol~' J mol~' deg~ x c, II ->c, I 1422 1, 5

c, 1-^ 1. 1620 40.6 25 5

6
Thermodynamic data (298.15 °K) C. Equilibrium pressures (atm)

s°

J mol-1 cal mol 1 References

deg-1 deg-1

BaS04 132.2 31.6 1, 3, 8 BaO 70.3 16.8 1, 3, 4, 8

AHf

kj mol-1 kcal mol-1 References

BaS04 - 1465.2 -350.2 1, 8 BaO -558.1 -133.4 1

Decomposition o/BaS04

(1) BaS04(c) BaO(c) + S03 (g) A//£98 15 = 51 1.9i kj, 122.35 kcal

(2) BaS04(c) «=* BaO(c) + S02(g) + 1/2 02(g) \H298.15 = 610.32 kj, 145.87 kcal.

A. Free Energies

T BaO BaS04 Reaction (1)

fef fef Afef AG°t

°K J J J kj

298.15 -70.29 -132.21 -194.74 + 453.84 400 -72.11 - 136.48 - 194.38 434.17 500 -75.77 - 145.02 - 193.64 415.09 600 -79.91 - 154.98 -192.70 396.31 700 -84.18 - 165.23 -191.80 377.65 800 -88.39 - 175.35 - 190.95 359.15 900 -92.38 - 185.06 - 190.20 340.74 1000 -96.23 - 194.35 -189.55 322.38 1100 -99.91 -203.13 - 189.06 303.93 1200 - 103.39 -211.58 -188.51 285.68 1300 - 106.69 -219.62 - 188.03 267.48

B. Equilibrium constants

T log Kx Ki log K2 k2

°K

298.15 -79.506 3.12 x 10-80 -91.835 1.46 x IO-92 400 - 56.693 2.03 x IO-57 -64.631 2.34 x lO"65 500 -43.361 4.36 x 10-44 -48.719 1.91 x IO-49 600 - 34.499 3.17 x IO-35 -38.145 7.16 x IO-39 700 -28.178 6.64 x 10-29 -30.601 2.51 x IO-31 800 -23.448 3.57 x 10-24 -24.956 1.11 X10-25 900 -19.775 1.68 X 10-20 -20.574 2.67 x 10-21 1000 - 16.838 1.45 x 10- 17 -17.072 8.47 x 10- 18 1100 - 14.430 3.72 x 10-15 - 14.204 6.25 x 10- 15 1200 - 12.435 3.67 x 10- 13 -11.827 1.49 x 10- 12 1300 -10.746 1.80 x 10- 11 -9.816 1.53 x 10- 10

The weight loss of BeS04 measured in a current

of air has been studied by Marchal [75]. Decomposition rates below 883 °K were too slow to be

measurable. At 973 °K the weight loss was a

linear function of time up to a composition corre- sponding to the basic salt 5BeO • SO3 . Beyond

this composition the rate was considerably higher.

Marchal’s directly measured equilibrium decomposition pressures, 863 to 1103 °K, are higher than

those calculated here for decomposition to BeO.

However, others have not been able to confirm the

formation of an intermediate oxysulfate, noting

rather that the last traces of S03 are retained

tenaciously.

Free energy functions for BeS04 are calculated

from data in [126], and for BeO from data in [134].

A better value for the heat of formation of BeO at 298.15 °K than that used in the calculations is — 598.7 kj, —143.1 kcal [4, 6, 8, 10], but the older

value from (1) is used because the BeS04 value

depends on it. The calculated decomposition

pressures above 900 °K are based on an extrapolation of the free energy function (fef) of BeS04 .

Density of BeS04

Phase T d Refer- ences

c °K

ambient 2.54 g/cm3(crystallographic) 49

c 298 2.44a 7

Transitions of BeS04

Phase Refer- ences

c — c None observed. 298-900 °K

c — 1. Decomposition

7
Thermodynamic data (298.15 °K) C. Equilibrium pressures (atm)

5°

J mol-1 cal mol-1 Refer- deg-1 deg-1 ences

BeS04 77.91 18.62 126 BeO 14.1 3.37 1, 3, 4, 8

AHf°

kj mol-1 kcal mol-1 Refer- ences

BeS04 -1196.6 -286.0 1, 8 BeO -610.9 - 146.0 1 (see text)

Decomposition of BeS04

(1) BeS04(c) «=* BeO(c) + S03(g)

, 5 = 190.6 kj, 45.5s kcal

(2) BeS04 (c)«=* Be0(c) + S02(g) + l/2 02(g) A//298. ls = 289 o kJ, 69.0? kcal.

A. Free energies

T BeO BeS04 Reaction (1)

fef fef Afef AG°

°K J J J kj

298.15 -14.10 -77.91 - 192.86 + 133.08 400 -15.22 -92.59 -181.37 118.03 500 -17.56 - 106.36 -174.10 103.03 600 -20.43 -119.45 -168.73 89.34 700 -23.48 -131.84 - 164.49 75.44 800 -26.58 - 143.64 - 160.85 61.90 900 -29.63 - 155.02 - 157.49 48.84

(1000) -32.61 -166.57 -153.71 36.87

(1100) -35.49 -177.61 -150.16 25.40

(1200) -38.28 - 188.74 - 146.25 15.08

T S02 o2 S03 Total

°K

298.15 2.18X 10-24 1.09 x 10-24 4.85 x 10-24 8.12 x 10-24 400 3.42 X 10- 16 1.71 x 10-16 3.87 X 10-16 9.00 x lO"16 500 2.08 x 10-11 1.04 X 10- 11 1.53 X10-11 4.65 x 10-11 600 3.06 X 10-8 1.53 X lO-8 1.67 X lO-8 6.25 x lO-8 700 5.40 x lO"6 2.70 x lO-6 2.35 x lO-6 1.05X10-5 800 2.52 X 10-4 1.26 X 10-4 9.10 X 10-5 4.69 x 10-4 900 4.76 X 10-3 2.38 x lO-3 1.46 x lO"3 8.60 x lO-3 (1000) 4.58 X 10-2 2.29 x 10-2 1.19 X lO"2 8.05 X lO"2 (1100) 2.80 x 10-1 1.40 X10-1 6.22 X lO"2 4.82 x 10- 1 (1200) 1.17 0.585 0.221 1.98

Bismuth

Bi2 (S04) 3 is sufficiently stable at 650 °K for ignition to constant weight. Noticeable decomposition

begins about 25° above this temperature, yielding one or possibly a sequence of two oxysulfates. In open air the product Bi2 C>3 • S03 persists at least as high as 1250 °K. The final product is expected to be Bi2 03 . A complication is the low melting point

of this oxide, 1090 °K [5].

Density of Bi2 (S04)3

Phase T d Refer- ences °K

c 288 5.08 16

Transitions o/Bi2 (S 04)3

Phase

change

c — c No information. c-* 1. Decomposes.

Thermodynamic data (298.15 °K)

S°

B. Equilibrium constants

J mol-1 cal mol-1 Refer- deg"1 deg" 1 ences T log Ki Ki log K2 k2 Bi2 (S04)3 Bi2 C>3 151.5 36.2 1, 3, 9, 8 °K

298.15 -23.314 4.85 x 10-24 -35.643 2.28 X lO"36 400 -15.412 3.87 X 10- 16 -23.350 4.47 X lO"24 AHf

500 -10.815 1.53 x lO" 11 -16.173 6.71 X 10- 17 600 -7.777 1.67 x lO-8 -11.423 3.78 X 10- 12 700 - 5.629 2.35 X lO-6 -8.052 8.87 X 10-9 kj mol-1 kcal mol-1 Refer- 800 -4.041 9.10 x 10-5 -5.549 2.83 X lO"6 ences

900 -2.835 1.46 x lO"3 -3.634 2.32 X 10-4 (1000) -1.926 1.19 X lO-2 -2.160 6.92 x lO-3 (1100) -1.206 6.22 X lO"2 -0.980 1.05 XIO"1 Bi2 (S04 )3 - 2544.3 -608.1 9, 8

(1200) -0.656 2.21 X 10-1 -0.048 8.95 X 10- 1 Bi2 03 -573.88 -137.16 8

8
No calculation of decomposition pressures is possible. High-temperature data exists only for

Bi2 03 [2]. Cadmium

Solid phase transitions and decomposition products of CdSC>4 have been studied recently by two

groups of workers [57, 95], both of whom give re- views of the confused earlier literautre. On heating, CdS04 undergoes at least one polymorphic

transformation, probably two, but there is still disagreement on the transition temperatures. The

highest temperature form melts near 1300 °K.

However, slow decomposition to 2 CdO • CdS04 is observed well below this temperature. The final

product of decomposition is CdO. Marchal [76]

measured decomposition pressures in the range

1053 to 1303 °K, but her experimental difficulties cast doubt on their value. They cannot be calculated thermodynamically because no high temperature data exist for any of these compounds.

Density o/CdS04

Phase T d Refer- ences

c °K293 4.692 (20/4) 116

c 297 4.691 7

Transitions of CdSCh

Phase change T Refer- ences

c — c No information

c * 1. 1273

Calcium

CaS04 decomposes directly to CaO. The supposed existence of an intermediate oxysulfate and

discrepant values for the melting point of CaS04

are explained [87] by a eutectic between CaO and

CaS04 near 1638 °K, with less than 20 percent CaO.

The increase in pressure in a static system was

studied by Briner, Pamm, and Paillard [21], but

their data are not sufficiently precise to yield kinetic

parameters.

Of the various equilibrium decomposition pres- sure measurements, those of Tschappat and Piece

[131] agree very closely with those calculated

thermodynamically.

The free energy functions of CaS04 and CaO are calculated from data in [2]. The decomposition

pressures in the range 1500 to 1800 °K depend on

an extrapolation of the free energy function of

CaS04 .

Density of CaSC>4

Phase T d Refer- ences

°K

c, II 298 2.960 ±0.005 14-5

Transitions of CaSC>4

Phase change T AH AS Refer- ences °K kj mol-1 J mol-1 deg-1

c, II—> c, I c, I —» 1. 1453

see text 21 14.2 5

Thermodynamic data (298.15 °K) c, II (or a) is the mineral anhydrite. A metastable 03) form

exists to a few hundred degrees above ambient temperature.

Thermodynamic data (298.15 °K

J mol-1 cal mol-1 Refer- deg-1 deg-1 ences

CdS04 123. 29.4 3, 8 CdO 54.8 13.1 1,3,8

AHf°

S°

J mol-1 cal mol-1 Refer- deg" 1 deg-1 ences

CaS04(II) 106.7 25.5 1, 3, 4, 8 CaO 39. 7 9.5 1, 3, 4, 8

kj mol-1 kcal mol-1 Refer- ences

CdS04 -928.8 — 222.0±0.3 11 -926.17 -221.36 1, 8 CdO -254.6 -60.86 1 -256.! — 61.2 ± 0.2 11, 8

AHf°

kj mol-1 kcal mol-1 Refer- ences

CaS04 - 1432.7 -342.42 1, 8 CaO -635.5 -151.9 1, 8

9
Decomposition of CaSO. C. Equilibrium pressures (atm)

(1) CaSO,(c)sfs CaO(c) + S03(g) 1H* -401.* kj, 96.0, kcal T S02 o2 so3 Total

(2) CaSCMc)^ CaO(c) + S02 (g) + 1/2 02 (g) ^298 , 5 = 500.3s kj, 119.59 kcal. °K

298.15 3.49 X 10~49 1.75 x 10-49 3.11 x lO"61 5.24 X lO"49 A. Free energies 400 2.59 X 10-34 1.30 x 10-34 2.56 x lO"43 3.89 x 10"34 500 1 31 x 10-25 6 57 x lO-26 7.67 x 10-33 1.97 X 10~25 600 8.19X 10-20 4.09 x 10-20 7.33 x 10-26 1.23 x lO" 19 T CaO CaSO-i Reaction (1) 700 1.11 x lO" 15 5.54 x 10- 16 6.90 x 10-21 1.66 x 10 15 800 1 34 x 10-12 6.72 x 10- 13 3.55 x lO" 17 2.02 x 10- 12 fef fef Afef AC°T 900 3.22 x 10- 10 1.61 x 10- 10 2.57 x 10- 14 4.83 x 10- 10 1000 2.48 x 10-8 1.24 x 10-8 4.74 x 10- 12 3.73 x 10-8 1100 8.39 x 10~7 4.19 x 10-7 3.23 x 10- 10 1.26 x HP6 °K J J J kj 1200 1.52 X 10-5 7.62 x 10-6 1.04 x 10-8 2.29 x 10-5 298.15 -39.75 - 106.69 -189.72 + 345.39 1300 1.73 x lO"4 8.64 X 10-5 1.89 x 10-7 2.60 x 10-4 400 -41.51 -110.83 - 189.42 326.18 1400 1.35 X lO"3 6.73 x 10-4 2.18 x 10-6 2.02 x 10-3 500 -44.89 -118.74 - 189.04 307.44 (1500) 7.78 x lO"3 3.89 x 10~3 1.76 x 10-5 1.17 x 10-2 600 -48.74 - 127.82 - 188.68 288.74 (1600) 3.55 x 10-2 1.78 x 10-2 1.07 x 10-4 5.34 x 10-2 700 -52.72 -137.36 -188.20 270.20 (1700) 1.34 x 10- 1 6.70 x 10-2 5.15 x lO"4 2.02 x 10- 1 800 -56.61 - 147.03 - 187.50 251.96 (1800) 4.34 x 10- 1 2.17 x 10-’ 2.08 X 10-3 6.53 x 10- 1 900 — 60 29 — 156 73 — 186 43 234 18 1000 -63.85 - 166.36 -185.16 216.81 1100 -67.24 -175.85 - 183.67 199.91 1200 -70.46 - 185.06 -182.11 183.43 1300 -73.51 -194.01 - 180.46 167.36 1400 -76.44 -202.76 -178.72 151.75

(1500) - 79.24 -211.25 - 176.94 136.55

(1600) -81.92 -219.45 -175.16 121.70

(1700) -84.52 -227.32 -173.50 107.01 cerium

(1800) -86.99 -234.81 -171.95 92.45 For high-temperature properties, see “Rare

Earths.”

Density

Phase T d Refer- ences

Ce2(S04)3 (c) Ce(S04 )2 (c)

°K

294 3.912 17

Thermodynamic data ( 298.15 °K)

S°

B. Equilibrium constants r

T log Ki Ki log K2 k2

J mol-1 deg

cal mol 1 deg

Refer- ences

°K Ce2(S04 )3 298.15 -60.507 3.11 x 10-61 - 72.836 1.46 X 10-73 Ce(S04 )2 400 -42.592 2.56 x 10-43 -50.530 2.95 x 10-51 Ce2 03 500 -32.115 7.67 x lO"33 -37.473 3.37 x 10-38 Ce02 600 -25.135 7.33 x 10-26 -28.781 1.66 x 10-29 700 -20.161 ' 6.90 X10-21 -22.584 2.61 X 10-23 800 - 16.450 3.55 x 10- 17 -17.958 1.10 x 10- 18 AHf° 900 — 13 590 2 57 x 10-14 — 14 389 4 08 x 10~ 15 1000 -11.324 4.74 x 10- 12 -11.558 2.77 x 10- 12 1100 -9.491 3.23 x 10- 10 -9.265 5.43 x 10- 10 kj mol-1 kcal mol-1 Refer- 1200 -7.984 1.037 x 10-8 -7.376 4.21 x 10-8 ences

1300 — 6 724 1 888 x 10~7 — 5 794 1 61 x 10~6 1400 -5.662 2.18 x 10-6 -4.457 3.49 x lO"5 (1500) -4.755 1.76 X10"5 -3.314 4.85 x 10-4 Ce2 (bU4 )3 (1600) -3.972 1.067 x 10-4 -2.325 4.73 x 10-3 Ce(S04 )2 (1700) -3.288 5.15 x 10-4 -1.460 3.47 x lO"2 Ce2 03 (1800) -2.683 2.08 x lO"3 -0.694 2.02 x 10- 1 Ce02 IO88.7 260.2 8

10
Cesium

CS2SO4 is stable even beyond the melting point.

According to Jaeger [59], vaporization is first noticeable at 1600 °K, and is very fast by 1710 °K. Spitsyn

and Shostak [125], in maintaining Cs2 S04 for 2 hr

at each of several temperatures, found no weight

loss at 1073 °K, 0.2 percent loss at 1173 °K, and

0.87 percent loss at 1273 °K, 13.9 percent loss at 1473 °K, attributed entirely to vaporization.

No high-temperature thermal data are available

for CS2SO4 and CS2O from which equilibrium deDensity 0/CS2SO4

Phase T d Refer- ences

Solid °K

293 4.246 133 298 4.243 7 333 4.222 132 Liquid 1309 3.037 59 1336 3.018 59 1378 2.988 59 1438 2.937 59 1494 2.889 59 1547 2.841 59 1604 2.787 59 1645 2.743 59

1696 2.690 59

1743 2.636 59 1803 2.566 59

Transitions of CS2 SO4

Phase change T AH AS Refer- ences

c, II ->c, I °K933

kj mol~l J mol-1 deg-1 5

c, I-> 1. 1277 40.2 31.4 5

Thermodynamic data (298.15 °K)

S°

J mol 1 cal mol-1 Refer- deg-1 deg-1 ences

Cs2 S04(II) Cs2 0

AHf°

kj mol-1 kcal mol 1 Refer- ences

Cs2 S04 (II) - 1420.0 -339.38 1, 8 CS2 O -317.6 -75.9 1, 8

composition pressures could be calculated. Entropies have not been reported for either compound.

Chromium

Cr (il). CrS04 is known only in solution and as hydrates and double salts, all easily oxidized by

air. The monohydrate on heating yields C^OsCr (ill). Cr2 (S04)3 is stable to 573 °K, but decomposes not far above that temperature, ulti- mately to Cr2 C>3. The course of the decomposition

is unknown. Oxysulfates have been prepared in

ways other than by pyrolysis of the anhydrous

sulfate.

Density of Cr2(S 04 )3

Phase T d Refer- ences

c °K

298 2.706 81

Transitions of Cr2(S 04)3

Phase

change

c —> c No information. c —* 1. Decomposition.

Thermodynamic data (298.15 °K)

s°

J mol-1 cal mol-1 Refer- deg-1 deg-1 ences Cr2 (SC>4)3 Cr2 03 8I.I7 19.4 1, 3, 8

AHf°

kj mol-1 kcal mol-1 Refer- ences

Cr2 (S04 )3 Cr2 03 -1128.4 -269.7 1, 8 Cobalt

The decomposition of C0SO4 has been studied by

Pechkovsky [100, 103] in a dynamic system at a flow rate of 3 1/hr between 1023 and 1223 °K. In- crease of SO2 in the carrier gas inhibited the decomposition. The decomposition is linear in time

with no evidence for an oxysulfate, i.e., a = kt. In k

is a linear function of 1/T between 1113 and 1133

°K, the slope yielding on activation energy of 319 kj

(76.2 kcal).

216-963 0-66—3

11
In an initially evacuated static system the product is CoO, though C03 O4 often appears in TGA

runs in air. Decomposition pressures have been

measured by Warner [136] over the range 950 to 1170 °K. The results are compared graphically

with those in the earlier literature. The measured

pressures are not tabulated, but equations derived

from them are given for AG° and log PS03 as functions of temperature, valid only over the experimental range. From these we have calculated

Afef, extrapolating for 900 and 1200 °K. Equilibrium constants and decomposition pressures are calculated from these Afef values.

The entropy of C0SO4 listed by [1] on which Afef

at 298.15 °K depends, is quite incertain, although

close to what one would estimate. Later critical

compilations [3, 8] do not list an entropy.

Density of C0SO4

Phase T d Refer- ences

°K a 298 3.86 g/cm3 25 P 298 3.770 g/cm3 109

These densities are crystallographic. Pycnometric values are doubtful because the high-temperature /3-form can be frozen in. Materials used

may be mixtures.

Transitions of C0SO4

Phase change T AH AS Refer- ences

°K kj mol_1 J mol_1 deg~ x

a —

>

/3 705 ±8 6 . 8 ± 0.6 9.6±0.8 109

Thermodynamic data (298.15 °K)

S°

J mol-1 deg-1 cal mol-1 deg-1 Refer- ences

C0SO4 113.4 27.1 1 CoO 52.97 12.66 3, 8 C03 O4 102.5 24.5 3, 8 AHf°

kj mol-1 kcal mol-1 Refer- ences

C0SO4 -887.0 — 212.0±0.4 11 CoO -238.9 -57.1 19, 8 -878. - 210. 1 C03O4 - 905.0 -216.3 8

Decomposition of C0SO4

(1) CoSO,(c) CoO(c) + S03 (g) A ,5 = 252.4 kj, 60.32 kcal

(2) CoSO,(c) <=* CoO(c) + S02(g) + 1/2 02 A^98 .„ = 351.3 kj, 83.96 kcal.

A. Free energies, reaction (1)

T Afef AG£

°K J kj

298.15 - 196.2 193.89

(900) - 188.3 82.9a 1000 -187.4 64.

9

4 1100 - 186.6 47.1,

(1200) -185.8 29.4«

B. Equilibrium constants

T log Ki A, log K2 k2

°K

298.15 -33.966 1.08 x IO-34 -46.383 4.14 X IO

-47 (900) -4.813 1.54 x 10-5 -5.639 2.29 x 10-« 1000 -3.392 4.06 X 10-4 -3.651 2.23 X 10-4 1100 -2.237 5.79 x 10-3 -2.033 9.27 X 10-3 (1200) -1.282 5.22 X 10-2 -0.694s 2.02 x 10-1

C. Equilibrium pressures (atm)

T S02 o2 S03 Total °K

298.15 1.5 X IO

-31 7.5 x 10"32 1.1X10-34 2.3 x IO-31 (900) 2.2 X 10-4 1.1 X 10-4 1.5 X 10-5 3.4 x 10-4 1000 4.6 X 10-3 2.3 x 10-3 4.1X10-4 7.4 X 10-3 1100 5.6 X 10-2 2.8 x 10-2 5.8 x 10-3 8.9 X 10-2 (1200) 4.3 x 10-1 2.2 X 10-1 5.2 X 10-2 7.0 xlO"1

Copper

Cu (i). CU2SO4 oxidizes in air. Cu (il). There is general agreement that the first product of the decomposition of CuS(>4 is an oxysulfate, CuO • CuS04 . In kinetic studies, the temperature range of this reaction has been reported

as 1092 to 1133 °K [61] and 1053 to 1123 °K [128],

possibly because of differences in the gas phase

composition. The rate of decomposition decreases

with increasing oxygen content of the flow gas and

increases with increasing SO3 content because

these gases also displace the S03 = SO2 + 1/2 02 equilibrium in opposite directions.

The decomposition of both sulfate and oxysulfate

is zero-order. For the former an Arrhenius activation energy of 212 kj (50.6 kcal) was obtained between 1053 and 1123 °K, and for the decomposition

of the oxysulfate Ea = 346 kj (82.8 kcal) between

1133 and 1193 °K [128]. However, somewhat dif- ferent values of 238 kj (56.8 kcal) and 269 kj (64.3

kcal) for these reactions have also been reported

12
[103]. Although the values for the sulfate decomposition agree fairly well, those for the oxysulfates are quite different, possibly because the

Russian workers ran this decomposition at a lower

temperature, 1033 to 1073 °K.

Two sets of measurements [85, 117], in fair agreement, have been made of the equilibrium decomposition pressures, 948 to 1093 °K, in the systems

CuS04 CuO • CuS04 + gases

CuO • CuS04 ^ CuO+ gases

where the pressures in the first system are higher

than in the second. For lack of data on the oxysulfate, it is not yet possible to calculate these

pressures thermodynamically. We have, however, calculated pressures for the system

CuS04 ^ CuO 4- gases

Thermodynamic data (298.15 °K)

s°

J mol-1 cal mol-1 Refer- deg"1 deg-1 ences CuS04 113.4 27.1 1 (see text) CuO 42.63 10.19 3, 4, 8 Cu2 0 93.7 22.4 3, 8

A///

kj mol-1 kcal mol-1 Refer- ences

CuS04 -769.86 - 184.00 1, 8

(see text) CuO -155.4 -37.1 1,8

Cu2 0 -1667. -39.84 1, 8

using high-temperature thermal data from [2] and

extrapolating the free energy function of CuS04 above 900 °K. These pressures are, as they should

be, intermediate between those for the two actual

systems. Inasmuch as the latter pressures differ

by less than an order of magnitude, the calculated

pressures given in the tables that follow will serve

as a rough approximation to the actual systems,

particularly outside the narrow temperature range

of the measurements.

The entropy of CuS04 is uncertain. [3] lists 25.3 ±1.0 cal mol-1 deg-1 as “suggested” in [60],

We estimate 26.9. The heat of formation of CuS04 has recently been redetermined [11] as — 183.4 ±0.6

kcal mol-1 (767.3 kj mol-1

), but the value from [1] was retained in the calculations because it is part

of a self-consistent set.

Density of C11SO4

Phase T d Refer- ences

? °K

298 3.542 (pycnometric) 81 a 298 4.113 g/cm3(crystallographic) 110 P 453 3.923 g/cm3(crystallographic) 110

Transitions of C11SO4

Phase

change

T AH AS Refer- ences

a—> /3 °K

9 No data No data

c —> 1 decomposes

Decomposition 0/C11SO4

(1) CuS04(c) «=* CuCXc) + S03(g)

,5 = 219.45 kj, 52.45 kcal

(2) CuSCMc) CuO(c) + S02 (g) + 1/2 02(g) A#29s.i 5 = 317.86 kj, 75.

7 kcal.

A. Free energies

T CuO CuS04 Reaction (1)

fef fef Afef AG°t

°K J J J kj

298.15 -42.63 -113.39 -185.92 + 164.02 400 — 44.45s -117.47 -185.74 145.16 500 -47.82 - 125.27 -185.44 126.73 600 -51.76 - 134.26 - 185.25 108.30 700 -55.81 -143.59 -185.06 89.91 800 -59.79 - 152.88 - 184.82 71.60 900 -63.64 - 162.09 - 184.42 53.47

(1000) -67.32 -171.29 - 183.69 35.76

(1100) -70.92 - 180.50 - 182.70 18.48

B. Equilibrium constants

T log Ki Ki log K2 k2

°K

298.15 - 28.734 1.85 X 10-29 - 41.063 8.65 x 10-42 400 - 18.954 1.11 x 10-19 - 26.892 1.28 x 10-27 500 - 13.238 5.78 X 10-14 - 18.596 2.54 X 10- 19 600 - 9.427 3.74 x 10-10 - 13.073 8.45 x 10-14 700 - 6.708 1.96 X 10-7 - 9.131 7.40 x 10- 10 800 - 4.674 2.12 x 10-5 - 6.182 6.58 x 10-7 900 - 3.103 7.89 x 10-4 - 3.902 1.25 xlO-4

( 1000) - 1.868 1.355 X 10-2 - 2.102 7.91 x 10"3

( 1100) — 0. 8772 1.325 X 10-1 - 0.651 2.23 XlO" 1

13
C. Equilibrium pressures (atm) Europium

T S02 o2 so3 Total

°K

298.15 5.31 x 10-28 2.66 x 10-28 1.85 x lO" 29 8.15 X 10"28 400 1.49 x 10“ 18 7.44 x 10- 19 1.11 x 10- 19 2.34 x 10“ 18 500 5.05 x 10- 13 2.52 x 10-13 5.78 x 10- 14 8.15 X 10- 13 600 2.43 x 10-9 1.21 x 10-9 3.74 x 10- 10 4.01 x 10-9 700 1.03 x 10-6 5.15 x 10-7 1.96 x 10-7 1.74 x 10-6 800 9.53 x lO-5 4.76 x 10-5 2.12 x 10“5 1.64 x 10-4 900 3.16 x 10-3 1.58 x 10-3 7.89 x lO'4 5.52 x 10-3

( 1000) 5.00 x 10-2 2.50 x 10-2 1.36 x 10-2 8.86 X 10-2

( 1100) 4.64 x 10- 1 2.32 x 10- 1 1.33 x 10- 1 8.28 x 10- 1

For high-temperature properties, see “Rare

Earths.”

Density

Phase T d References

EuS04(c) °K

293 4.989 79 EuSC>4(c) 298 4.981 79 Eu2 ( 804)3 ( 0 ) 297 3.997 17

Dysprosium

For high-temperature properties, see “Rare

Earths.”

Gadolinium

For high-temperature properties, see “Rare

Earths.”

Density o/Dy(S04)3

Phase T d References

°K

c 297 3.801 17 Erbium

For high-temperature properties, see “Rare

Earths.”

Density o/E^SCEh

Phase T d References

°K

c 297 3.851 17

Thermodynamic data (298.15 °K)

s°

kj mol 1 kcal mol 1 References

Er2 (S04 )3 Er2 03 - 1897.9 - 453.6 8

Thermodynamic data ( 298.15 °K)

S°

kj mol-1 kcal mol 1 References

Gd2 (S04)3 Gd2 03 - 1815.9 - 434.0 8

Gallium

Ga2(S04)3 decomposes, without melting, to Ga2 03 . Marchal [76] measured equilibrium decomposition

pressures over the range 833 to 973 °K, from which

it appears that this compound is less stable than

A12 (S04)3 . Pressures cannot yet be calculated

thermodynamically.

Density of Ga2 (S 04)3

Phase T d References

°K

c 296 3.239 17

14
Thermodynamic data (298.15 °K) Thermodynamic data (298.15 °K)

s° s°

J mol 1 deg"1 cal mol-1 deg" 1 References J mol-1 deg"1 cal mol-1 deg” 1 References

Ga2 (S04 )3 Ga2 03 (/3) 84.9 20.3 Hf(S04 )2 3, 8 Hf02 59.33 14.18 3, 4, 8

AHf° AHf°

Ga2 (S04 )3 Ga2 03 (/3)

kj mol-1 kcal mol-1 References kj mol-1 kcal mol 1 References

-1079. -258

Hf(S04 )2 1, 8 Hf02 -1113.1 -266.0 4, 8

Germanium

Density of Ge(S04)2

Phase T d References

°K

c 295 3.92 121

Thermodynamic Data (298.15 °K)

S°

J mol 1 cal mol 1 References

deg-1 deg"1

Ge(S04 )2 Ge02 (c, soln.) 55.

7 13.21 3

AHf°

kj mol 1 kcal mol 1 References

Ge(S04 )2 Ge02 (c) -539.7 -129.0 5

Hafnium

Von Hevesy and Cremer [135] attempted to measure equilibrium decomposition pressures of Hf(S04)2

in the range 823 to 923 °K, but were not sure of

obtaining true equilibrium. Decomposition was

direct to Hf02 , not through an intermediate oxysulfate, and the sulfate appeared to be more stable

than Zr(SC>4 )2 .

Holmium

For high-temperature properties, see “Rare

Earths.”

Density o/Ho2(S04 )3

Phase T d References

°K

c 297 3.830 17

Thermodynamic data (298.15 °K)

s°

J mol-1 cal mol-1 References

deg- 1 deg- 1

Ho2 (S04 )3 Ho2 03

A///°

kj mol 1 kcal mol-1 References

Ho2 (S04 )3 Ho2 03 - 1881.1 -449.6 8

Indium

Seward [122] measured equilibrium decomposition pressures of In2 (S 04)3 over the range 918 to 1093 °K, from which it appeared to be more stable

than A12 (S04 )3 . The product was In2 C>3 without an

oxysulfate intermediate, contrary to indications in the early literature.

A kinetic study [63] at temperatures above 1073 °K

shows that the initial rate of decomposition is nearly

15
constant, but decreases with time [63]. The degree

of decomposition is given by

a = 1 — e~

ktH

which can also be written in the logarithmic form

ln[— ln(l — a:)] = rc Inf + lnA:. From the linear plot of the left-hand side versus In t the constants n and k are obtained, n increases

linearly from 1.0 at 1123 °K to 1.6 at 1273 °K while k

increases logarithmically from 0.021 to 0.410 between these temperatures. The values of n indicate

that decomposition proceeds from nuclei in two

dimensions.

Density of In2(S04 )3

Phase T d References

°K

c 298 3.438 7

Thermodynamic data (298.15 °K)

S°

In2 (S04)3 In2 C>3

J mol-1 deg"1 cal mol-1 deg"1 References

AHfO

kj mol-1 kcal mol-1 References

In2 (S04)3 2908 . -695 1 In2 (>3 -930.9 -222.5 1 -926.8 -221.5 8

Iridium

Ir (ill). Ir2 (S04)3 is known with certainty only as

a hydrate, oxidizing in air. Ir (iv). Ir(S04)2 has been reported.

Iron

The sulfates of both Fe (il) and Fe (ill) decompose

without melting.

The decomposition of Fe2(S04)3 to Fe2 03 is reversible, without the formation of an intermediate oxysulfate. FeS04 decomposes irreversibly, with

oxidation, to Fe2 C>3, through intermediates that have

not been identified with certainty, though FeO is often assumed.

Neumann and Heintke [84] measured decomposition pressures by the static method for both FeS04 (633-862 °K) and Fe2 (S04)3 (793-973 °K). They

include comparative tables of all the values from

the earlier literature. At particular temperatures,

pressures are higher for FeS04 than for Fe2(S04)3. With FeS04 , irreversibility prevented approaching

pressures from two directions. Later, Warner and

Ingraham [137] published values for Fe2 (S04)3 (903-

997 °K), but compared their results only with the erroneous data of Wohler, Pliiddemann, and Wohler

[144]. They agree better with Neumann and

Heintke; plots of the two sets of pressures against

temperature would cross near 960 °K. These pres- sures cannot be calculated thermodynamically

because high-temperature thermal data exist only

for the oxides [2].

Kinetics

FeS04 When FeS04 is heated above 823 °K, the FeO

formed is immediately oxidized by the oxygen or SO3 liberated in the reaction [84, 101, 115]. The

reactions occurring are 2FeS04 = 2FeO + 2S03 (1)

S03 = S02 + l/202 (2)

2FeS04 = Fe2 03 + S02 + S03 (3) corresponding to the net reaction

2FeS04 = Fe2 03 + S02 + S03 . (4)

If the SO3 remains in the system the following reaction may occurFe2 03 + 3S03 = Fe2 (S04)3 . (5) However, this reaction is slow and may be suppressed by removing the SO3.

As expected, an increase of 02 content in the

sweep gas inhibits the decomposition.

When the weight loss was measured in air [104]

between 953 and 1003 °K the decomposition fol- lowed the Roginsky equation

1 — (1 — a)113 = a = kt

when a is the degree of decomposition. From the

linear temperature dependence of In k on 1/T an Arrhenius activation energy of 253 kj (60.4 kcal) was

obtained. Since, at any temperature in the above

range, the decomposition pressure of FeS04 is greater than that of Fe2(S04)3 the rate-determining

step for the decomposition of the former is the decomposition of the ferric salt. The decomposition

16
appears Densities to occur exclusively on the surface of the

original material, rather than in the interior of the

grains. Only in the final phases of the decomposition does the diffusion of gases through the oxide

product appear to be rate-limiting.

Fe2 (S04)3 The decomposition of Fe2(S04)3 is zero-order

between 973 and 1073 °K. Activation energies of

83 kj (19.9 kcal) [138] and 154 kj (36.8 kcal) [127] for

the reaction have been reported, the latter value for

vacuum conditions, the former in a stream of dry N2 . When the gas phase has a constant composition of

S03 , S02 , and 02 , the reaction rate is a linear function of Pe(S02 ) — p(S02 )

, i.e., the driving force is the

Pe(S02 ) difference between the actual and the equilibrium

pressure of S02 . In air between 953 and 1053° Roginsky’s equation

holds [104], i.e., the degree of decomposition is a

linear function of time. This implies that the decomposition is not autocatalytic and that the new

phase is instantaneously formed on the surface of

the original material. The interphase then migrates

into the material. During the later stages the rate

is controlled by the diffusion of product gases

through the solid oxide. The logarithm of the rate constant k in the Roginsky equation

1-a-ot)113 ^ a = kt

is a linear function of 1 IT. The Arrhenius activation energy calculated from the slope is 289 kj

(69.1 kcal), quite different from the other values

reported.

No explanation is evident for the wide range of

activation energies reported by different authors

since one would not expect a radical change in

mechanism to be produced by a change of experimental conditions. At 1000 °K, Warner and Ingraham give for the reaction

Phase T d References

°K

FeS04 (c) 288 3.346 130 Fe2 (S04 )3(c) ambient 3.10 89

Thermodynamic data (298.15 °K)

S°

J mol-1 cal mol-1 References

deg-1 deg-1

FeS04 107.3 25.7 3, 8 Fe2 (S04)3 FeO(Fe0 .947O) 563.6 13.74 3 Fe3 04 146.4 35.0 1, 3 Fe2 C>3 87.4 20.9 3

AHf°

kj mol-1 kcal mol 1 References

FeS04 -922.6 -220.5 1, 8 Fe2 (S04 )3 FeO(Fe0 .947O) -266.5 -63.7 1 -264.4 -63.2 8 Fe3 04 -1117. -267.0 1, 8 Fe2 03 -822.15 -196.5 1, 8

Lanthanum

Fe2 (S04)3 = Fe2 03 + 3S03 AG = 526.510 — 501.03 T (in joules)

and for the reaction

Fe2 (S04 )3 = Fe2 03 + 3S02 T- 3/2 02 AG = 194,165 — 184.55 T (in joules).

Since the equilibrium SO3 — S02 + 1/2 02 is established relatively slowly, the rate determining step is more likely to involve the expulsion of an SO3 molecule from the crystal lattice. All of the reported

activation energies are much less than A// — 526 kj

(125.8 kcal), but AH per mole of SO3 is 526/3= 176

kj (42 kcal) which is at least in the right range.

For high-temperature properties, see

‘ Earths.”

Density o/La2(S 04)3

Phase T d References

°K

c 294 3.829 17

Transitions 0/La2 (SO4)3

Phase

change

c —*• 1. Decomposition

Rare

17
Thermodynamic data (298.15 °K)

5°

J mol-1 cal mol-1 References

deg deg-1

La2 (S 04 )3 La2 03

AHf°

kj mol-1 kcal mol-1 References

deg

La2 (S04 )3 L&2 O3 - 1801.2 -430.5 8

Lead

Pb (il). Because PbS04 decomposes noticeably below the melting point, this value is difficult to determine reliably. The ultimate decomposition

product is PbO. Various oxysulfates have been

prepared in ways other than by pyrolysis of PbSC>4 . Some have congruent melting points and form

eutectics with PbO or PbS04 ; for a summary and

references see [96b] and for thermodynamics,

see [60]. [1] lists PbO • PbS04 , mp 1243 °K,

2PbO • PbS04, mp 1223 °K. The former is the

mineral lanarkite. Hoschek [57] found three

intermediates before PbO in the pyrolysis of PbS04 , the two oxysulfates above and 4PbO • PbS04 . However, Terem and Akalan [127] in a kinetic study

found no evidence for intermediates, and gave an

Arrhenius activation energy of 384.9 kj (92.0 kcal)

for the zero-order reaction

PbS04 +* PbO + S03 . No equilibrium decomposition-pressure measure- ments are available. Nichols and White [88]

measured pressures by a dynamic method at four

temperatures in the range 1035 to 1177 °K. From

high-temperature thermal data in [2] we have cal- culated pressures for decomposition direct to PbO.

If decomposition is indirect, observed pressures

will be somewhat higher.

Pb (iv). The compound Pb(S04)2 has not been

prepared free of H2SO4.

Density o/PbSC>4

Phase T d Refer- ences

°K

c, II 288 6.2907 80 298 6.2866 80

308 6.2824 80 318 6.2780 80

Transitions o/PbS04

Phase change T AH AS Refer- ences

c, II—» c, I c, I-» 1 .

°K

1139

1360

kj mol~x

17.0 40

J mol~l deg- 1 14.9

29. 7

1, 5

1, 5

Thermodynamic data (298.15 °K

)

S°

J mol 1 deg" 1 cal mol 1 deg-1 Refer- ences

PbS04 (II) 147.3 35.2 1, 3, 8 PbO (yellow) 67.4 16.1 3, 10 PbO (red) 65.3 15.6 3, 10

AHf°

kj mol-1 deg" 1 kcal mol-1 deg" 1 Refer- ences

PbS04 (II) -918.39 -219.50 1 , 8 PbO (yellow) -217.86 -52.07 1 , 10 PbO (red) -219.24 -52.40 1 , 8, 10

18
Decomposition o/PbSO. C. Equilibrium pressures (atm)

(1) PbSO< (c)?± Pb0(c) + S03<g) to vellow PbO: A l5 = 30o.3s kj. 72.98 kcal

to red PbO: AtfJs 15 = 303.9t kj, 72.65 kcal

(2) PbSOdc) ;=* PbCKc)+ S02(g)+ 1/2 02 lgl to yellow PbO: A , 5 = 403.76 kj. 96.50 kcal

to red PbO: ^^.,5 = 402.38 kj. 96.17 kcal

Free energies

T PbO (yellow) PbS04 Reaction (1)

fef fef Afef AG°t

°K J J J kj

298.15 -67.36 - 147.28 -176.75 + 252.65 400 -69.25 -151.42 -176.57 234.72 500 -72.76 -159.41 -176.24 217.23 600 - 76.82 -168.70 -175.88 199.82 700 -81.04 - 178.32 -175.56 182.45 800 -85.19 - 188.03 -175.07 165.29 900 -89.29 -197.74 - 174.42 148.37 1000 -93.22 -207.36 - 173.53 131.82 1100 -96.99 -216.81 -172.45 115.65

(1200) — 101.55 -226.02 - 172.24 98.66

(1300) -106.94 -234.89 -173.01 80.43

(1400) -111.96 -243.34 -173.66 62.23

(1500) -116.61 -251.29 - 174.26 43.96 PbO (red)

298.15 -65.27 - 174.66 + 251.89 400 -67.24 -174.56 234.14 500 -71.06 -174.44 216.75 600 -75.27 -174.33 199.37 700 -79.58 -174.10 182.10 800 -83.83 -173.71 165.00 900 -87.95 -173.08 148.19

B. Equilibrium constants

To PbO (yellow)

T log Ki Ki log K2 k2

°K

298.15 -44.261 5.48 x 10~45 — 56.590 2.57 x 10-57 400 -30.648 2.25 x 10-31 -38.586 2.59 x 10-39 500 -22.692 2.03 x 10~23 -28.050 8.91 x 10~29 600 - 17.394 4.04 x 10" 18 -21.040 9.12 x lO"22 700 -13.614 2.43 x 10~ 14 - 16.037 9.18 x 10~17 800 - 10.792 1.61 x 10"11 -12.300 5.01 x lO" 13 900 -8.610 2.46 x 10-9 -9.409 3.90 x 10" 10 1000 -6.885 1.30 x 10~7 -7.119 7.60 x 10~8 1100 -5.491 3.23 x 10-® —5.265 5.43 x lO'6 (1200) -4.294 5.08 x 10-5 -3.686 2.06 x 10-4 (1300) -3.231 5.88 x 10-4 -2.301 5.00 x 10-3 (1400) -2.322 4.76 x 10~3 -1.117 7.64 x lO-2 (1500) -1.531 2.94 x lO-2 -0.090 8.13 x 10-1

To PbO (red)

298.15 -44.128 7.45 x 10-45 — 56.457 3.49 x 10-57 400 - 30.573 2.67 X 10-31 -38.511 3.08 x 10-39 500 -22.642 2.28 x lO"29 -28.000 1.00 x lO-29 600 — 17.355 4.42 x 10-18 -21.001 9.98 x lO"22 700 - 13.587 2.59 x 10-14 -16.010 9.77 x 10-17 800 -10.773 1.69 x 10- 11 -12.281 5.24 x lO" 13 900 -8.600 2.51 x 10-9 -9.399 3.99 x 10- 10

To PbO (yellow)

T so2 o2 S03 Total

°K

298.15 2.36 x 10"38 1.18 x 10"38 5.48 x 10-45 3.55 x 10"38 400 2.38 x lO"28 1.19 x lO-26 2.25 x 10-31 3.57 x lO-26 500 2.51 x 10-19 1.26 X10-19 2.03 x lO-23 3.77 x lO" 19 600 1.19 x 10-14 5.92 x 10- 15 4.04 x 10-18 1.78 x 10-14 700 2.57 x 10- 11 1.28 xlO-11 2.43 x 10- 14 3.85 x 10- 11 800 7.95 x 10-9 3.97 x lO"9 1.61 x 10-n 1.19 XlO"8 900 6.73 x 10-7 3.36 xlO-7 2.46 x 10-9 1.01 x lO-6 1000 2.26 x lO"5 1.13 x 10-5 1.30 x 10-7 3.41 x 10-5 1100 3.89 X lO"4 1.95 x 10-4 3.23 x lO-6 5.87 x 10-4 (1200) 4.40 x 10-3 2.20 x lO"3 5.08 x 10-5 6.64 x lO-3 (1300) 3.68 x lO-2 1.84 x lO-2 5.88 x 10-4 5.59 x lO-2 (1400) 2.27 x lO"1 1.13 xlO-1 4.76 x lO-3 3.45 x lO"1 (1500) 1.097 0.549 2.94 x lO-2 1.675

To PbO (red) 298.15 2.90 X lO"38 1.45 x 10-38 7.45 x 10-45 4.35 x 10-38 400 2.67 X lO"28 1.33 x lO-26 2.67 x lO"31 4.00x10-28

500 2.72 X 10-19 1.36 x 10- 19 2.28 x lO-23 4.07 x 10- 19 600 1.26 x lO" 14 6.29 x 10-15 4.42 x 10-18 1.89 xlO-14 700 2.67 x 10-“ 1.34 x lO" 11 2.59 x 10-14 4.01 x 10- 11 800 8.19 x 10-9 4.10 x 10-9 1.69 x 10- 11 1.23 x 10-8 900 6.83 x 10-7 3.41 x lO"7 2.51 x 10-9 1.03 x lO"6

Lithium

Decomposition of Li2 SO.i begins to be noticeable

not far above its melting point [59, 125], the Li2 0

product apparently dissolving in the Li2 SO.i, while

some volatilization of the Li2 S0.4 occurs. No

measurements of decomposition pressures have

been made, nor can the pressures be calculated

thermodynamically. High-temperature thermal

data exist for Li2 0 [2] but not for Li2 S04 .

Density o/Li2 SO.}

Phase T d References

c, II °K298 2.221 7

c, I 883 2.07 g/cm3 45

Liquid 1133

(crystallo- graphic) 2.004 59 1147 1.999 59 1170 1.989 59 1196 1.978 59

1236 1.962 59

1250 1.956 59

1274.4 1.947 59

1311 1.932 59

1330 1.924 59 1347 1.917 59

1363 1.911 59 1385 1.901 59 1430 1.884 59

1441 1.879 59

1457 1.873 59

1465.4 1.869 59

1487 1.860 59

19
Transitions o/Li2 S04

Phase changes T AH AS Refer- ences

°K kj mol~l J mol~1 deg~ l U2SO4, c II—* c I 859 27.2 31.8 5

c, I— 1. 1132 7.5 6.7 5

Thermodynamic data (298.15 °K)

S°

Li2 S04 (II) Li2 0

J mol 1 deg 1 cal mol 1 deg 1 References

37.89 9.056 10

AHf°

kj mol-1 kcal mol 1 References

Li2 S04(II) - 1434.4 -342.83 1, 8

Li2 0 -595.8 -142.4 1, 8 -598.7 -143.1 ±0.5 10

Lutetium

See “Rare Earths”

Magnesium

MgSC>4 begins to decompose below the melting

point, directly to MgO without an oxysulfate intermediate. Literature melting points therefore vary

widely.

Equilibrium decomposition pressures, by the

static method, of Marchal [75] (1223-1463 °K) are thought to be too high because of attack on the boat.

Knopf and Staude [64] made measurements (1144-

1313 °K) by a dynamic method. Their data ex- pressed as Pso3 are somewhat lower than Marchal’s,

the difference increasing with temperature. It is not yet possible to calculate these pressures thermodynamically for lack of high-temperature thermal

data for MgSC>4 , though such data are available

for its decomposition product, MgO [2, 134]. For

a thermodynamic treatment of Marchal’s data, see [60].

Kinetically the decomposition of MgS04 has been

studied in a flow system in the range 1223 to 1323

°K [98]. When N2 flows over the salt at 3 1/hr the

decomposition is linear in time. Under these conditions the product gases are continuously removed

and the equilibrium SO3 = SO2 + 1/2 O2 is not

established.

If these gases flow over the salt at 3 1/hr the re- action rate is nearly independent of oxygen content, but the proportion of S03 relative to SO2 decreases with decreasing 02 content. The acti- vation energy in the range 1293 to 1333 °K is 661

kj (158.0 kcal) [103].

Density o/MgS04

Phase T d References

°K

c, II c, II 288 2.71 130 298 2.66 7

Transitions of Mg MgS04

Phase change T AH AS Refer- ences °K kj mol~ l J mol~ l deg~ l c, II -* c, I 1283 No data 5

c, I —> 1. 1400 14.6 10.4 5, cf. 114

Thermodynamic data (298.15 °K)

S

O

J mol-1 deg-1 calmol~l deg_1 References

MgS04 91.63 21.9 1, 3, 8, 10 MgO (powder) 27.87 6.66 1,3

MgO (crystal) 26. 8 6.4 1,3

A///0

kj mol-1 kcal mol-1 References

MgS04 - 1278. - 305.5 1, 8, 10 MgO (powder) -598.10 - 142.95 1 MgO (crystal) -601.83 - 143.84 1 -601.2 - 143.7 8, 10 Manganese

Mn (il). The melting point of MnSCL, 700 °C,

quoted in handbooks, is erroneous. Decomposition begins before melting. The product usually

observed on pyrolysis in air is MnsCL, as in TGA

[28, 90], though Mn2 03 may appear mixed with it in

the lower temperature range, complicating the

kinetic picture [97]. Marchal [76] measured decomposition pressures in the range 1093 to 1383 °K.

We have calculated them thermodynamically, by

the method described below, using high-temperature thermal data from [2] with an extrapolation of

the free energy function of MnSCL to 1200 °K. We

also calculated the pressures for decomposition to Mn2 03. A comparison of these two sets of pres- sures with the calculated oxygen pressures for

pairs of the four oxides of manganese show that

Mn3 04 is the product expected thermodynamically

20
in an initially evacuated system containing only

MnS04 originally, in agreement with Mah [71].

Marchal’s measured decomposition pressures

are substantially lower than those calculated here.

Kelley [60] used them to calculate the heat and

free energy of formation of MnS04 , but the thermodynamic constants available then (1936) have been

mostly superseded, so that the near agreement of

this heat of formation with those based on the work

of Thomsen and of Berthelot is fortuitous. Mah

[71] calculated partial pressures of SO2 and 02 expected for the decomposition of MnS04 to Mn3 04 , using some constants different from those

used here, and a simplified method that ignores

Pso3 and assumes Pso2 = 3P02 , obtaining pressures

slightly lower than those tabulated here.

Mn (ill). Mn2 (S 04)3 is said to be stable at 473 °K,

decomposition beginning about 100° higher, to MnS04 , 02 , and S03 . Mn (IV). Mn(S04)2 is somewhat unstable at room

temperature.

first step for mathematical convenience. K3 and

K4 now become:

Equating ( 1) and (2) to eliminate x: 6n8

-\- K4n3 — 2KaK4 = 0.

( 3) K3 and K4 are first evaluated for each temperature

of interest from thermodynamic data of the com- pounds involved. Equation (3) is then solved for n at each temperature.

Then:

Calculation of Partial Pressures

There are four equilibrium equations and corre- sponding equilibrium constants:

(a) S03 ^S02 + l/2 02

Pso3 ~ n

so2 K3 (PSO3)

2

(Psq2)(Pq2 ) 1/2 Densities

(PS03) Phase T d Refer-

(b) 3 MnS04 Mn3 04 + 3 S02 + 02 ences

K-2 = (PS02)3 (P02 ) MnS04 (c) °K

298 3.181 (25/4) 81

(c) 3 MnS04 ^ Mn3 04 -f 2 S03 + S02 Mn2 (S04)3(c) 288 3.24 (g/cm3

) 16

K3 — {Pso3)2 (Pso2) Thermodynamic data (298.15 °K)

(d) 6 MnS04 + 02 2 Mn3 04 + 6 S03 K4 =

(Pso3)6 (Pq2) Three relations are needed to solve for the three

pressures, but only two of these four equilibriumconstant expressions are independent; any two can

be derived algebraically from the other two. For ex- ample: K2 = K\K3 . The required third relation is obtained from:

S02 o2 SO3

X

X

3

0 moles formed in reaction b — n — n ~Y + n moles formed in reaction a

(x —n) (H) n moles present at final

equilibrium.

Reaction (b), rather than (c) or (d), is chosen for the

S°

J mol-1 deg-1 cal mol-1 deg-1 Refer- ences

MnS04 112. 26.8 1, 3, 8 MnO 59.71 14.27 3, 8 Mn304 (a) 154. 36.8 91 Mn2 03 IIO.5 26.4 3, 8 Mn02 (I) 53.05 12.68 1, 3, 8

AHf°

kj mol-1 kcal mol-1 Refer- ences

MnS04 — 1063.

4 -254.24 1, 4, 8 MnO -385.1 -92.05 1, 4, 8 Mn3 04 (a) - 1386.6 -331.4 4, 8 Mn2 03 -956.9 -228.7 91 Mn02 (I) -520.9 - 124.5 1, 8

21
Decomposition o/MnS04 Mercury

3 MnS04 (c) MnaO.Ca) + 3 S02 (g) + 02 (g) A//^98 = 914.1 Id, 218.4, kcal.

A. Free energies

T Mn3 04 MnS04 Afef AG°t

fef fef

°K J J J kj

298.15 -153.97 -112.13 -766.92 685.42 400 - 159.93 -116.32 - 766.32 607.55 500 -171.46 -124.77 -764.16 532.00 600 - 184.64 - 134.68 -761.10 457.42

700 -198.11 -145.10 -757.24 384.01 800 -211.40 - 155.44 - 753.28 311.45 900 -224.22 - 165.56 - 749.09 239.90 1000 -236.44 -175.35 -744.80 169.28 1100 -248.11 -184.77 - 740.56 99.46

(1200) -259.30 -194.18 -735.30 + 31.72

Hg (i). Pascal [1958], using very old references,

describes Hg2 S04 as unstable to light, disproportionating to Hg° and HgS04 , and melting with

decomposition to Hg°, O2 , and SO2 . Hg ( 11). According to a recent study by Hoschek

[57], HgS04 is unchanged after 2 hr at 673 °K, but

by 773 °K decomposition occurs to HgO • HgS04 and then 2HgO • HgS04 , identified through their

x-ray spectra. All three compounds are volatile

at elevated temperatures. The supposed next

product HgO was not observed because it is even more volatile. If HgS04 is heated rapidly to 973

°K, some Hg2S04 appears along with the oxysulfates

before volatilization is complete. No thermal

data above 298 °K exist for any of these compounds.

Densities

B. Equilibrium constants

T log K K x

°K M 298.15 - 120.076 8.40 x 10- 121 400 -79.331 4.67 x lO"80 500 -55.573 2.67 x IO

-56 600 -39.819 1.52 x 10-40 700 -28.653 2.22 x 10-29 800 - 20.334 4.63 X lO"21 900 -13.922 1.20 X 10- 14 1000 -8.841 1.44 X 10-9 1100 -4.722 1.90 x 10-5 (1200) -1.381 4.16 x 10-2

C. Equilibrium Pressures (atm)

T S02 o2 S03 Total

°K

298.15 1.26 X IO

-30 4.20 x IO-31 1.74 x IO-33 1.68 X lO”30 400 1.94 x 10"20 6.43 x lO"21 1.35 x 10-22 2.59 X 10-20 500 1.69 x 10- 14 5.57 x 10- 15 2.87 x 10" 16 2.27 X 10- 14 600 1.47 x 10" 10 4.81 x 10" n 4.50 x 10- 12 1.99 X 10-'° 700 9.09 x 10"8 2.96 X 10-8 4.14 x lO"9 1.25 X 10-7 800 1.09 X10-5 3.54 x 10-6 6.63 X 10-7 1.51 X lO"5 900 4.39 x lO"4 1.41 x 10-4 3.29 x lO"5 6.13 X 10-4 1000 8.20 x 10-3 2.61 x 10-3 7.19 x 10-4 1.15 x lO"2 1100 8.80 x 10-2 2.79 x 10-2 8.73 x 10-3 1.25 x 10- 1 (1200) 6.03 x 10- 1 1.90 x 10- 1 6.48 x lO"2 8.57 x 10- 1

Phase T d Refer- ences

HgS04 (c) °K

298 6.49 (pycnometric) 13 298 6.50 (crystallographic) 13

Thermodynamic data (298.15 °K)

S°

J mol-1 cal mol 1 Refer- de^1 deg-1 ences

HgS04 Hg2SU4 201. 48.0 1, 3, 9, 8 Hg2 0

HgO (red) 70.3 16.8 3, 8, 10 HgO (yellow) 72.4 17.3 3

AHf°

kj mol 1 kcal mol 1 Refer- ences HgS04 -704.2 - 168.3 1, 8 Hg2S04 -743.12 -177.61 9 Hg2 0 -91.2 -21.8 1 HgO (red) -90.7 -21.68 1, 8 HgO (yellow) -90.21 -21.56 1

22
Neodymium

For high-temperature properties, see “Rare

Earths.”

thermal data on NiS04 . Such data are available

for NiO [2]. The necessary thermodynamic constants at 298 °K are known, as tabulated below,

except the entropy of NiS04 . [1] gives a value for

it that seems much too low; [3] does not list one.

Density 0/Nd2(SO4 ) 3

Phase T d Refer- ences

c °K

294 3.961 17

Thermodynamic data (298.15 °K)

S°

J mol-1 cal mol-1 Refer- deg"1 deg-1 ences

Nd2 (S04 )3 Nd2 03

kj mol 1 kcal mol 1 References

Nd2(S04)3 . Nd2 03 - 1782.4 -426.0 8

Nickel

NiS04 decomposes, without melting, to NiO.

TGA curves have not shown evidence of an oxysulfate intermediate [22, 28, 73, 90]; some decomposition of the NiO may occur [22], especially if the

thermobalance is continuously evacuated [28].

Equilibrium decomposition pressures have been

measured by Marchal [76] (973-1218 °K) and by

Wohler and Flick [141] (1048—1119 °K) whose values

are about twice as high as Marchal’s. Neither

found an oxysulfate intermediate. For a thermodynamic treatment of Marchal’s data, see [60].

Both also worked with CoS04 and agree that NiS04 is the less stable in the sense that its decomposition

pressures are higher in these temperature ranges.

It is not yet possible to calculate those pressures

thermodynamically for lack of high-temperature

Density of NiS04

Phase T d References

°K

c 288 3.543(25/4) 81

Transitions of NiS04

Phase

change

c —* c No information c—> 1. Decomposition

Thermodynamic data (298.15 °K)

S°

J mol 1 cal mol-1 References

deg-1 deg- 1

NiS04 NiO 38.0 9.08 3, 8

AHf°

kj mol-1 kcal mol-1 References

NiS04 -870.7 -208.1 ±0.2 11 NiO -239. 7 — 57.3±0.1 19

Palladium

Anhydrous PdS04 has not been prepared free

of H2 0 or H2 S04 . The hydrate PdS04

• 2H2 0

on heating loses H2 0, then S02 , yielding ultimately

the metal.

Density of PdS04

Phase T d References

°K

c 287 4.012 17

23
Plutonium Density of K2 SO<

For the oxidation states Pu (ill) and Pu (vi)

(PuO^"1") only sulfates in solution or double salts are known. Pu (iv) forms a tetrahydrate but,

contrary to earber reports, attempts [32] to prepare

a stoichiometric Pu(S04)2 by heating materials

initially containing H2 0 and H2 S04 let to a product

of composition varying with conditions. TGA

showed a Pu(S04)2 plateau that was not quite level,

followed by a drop to the PuO2 . 09 level.

Polonium

Like tellurium, Po forms a basic sulfate,

2Po02

• S03, but in contrast to Te it also forms a normal sulfate Po(S04)2 anhydrous and hydrated.

The anhydrous salt is stable to above 673 °K,

decomposing to Po02 near 823 °K [14].

Phase T d References

°K

c, II (a) 291 2.663 55

c, II 293.8s 2.662 12

c, II 298 2.662 12

c,I(/3) 298 2.621 ±0.004 120

c, II (a) 313.65 2.655 12 Liquid 1343.4 1.888 59

1376 1.870 59

1418 1.848 59

1472 1.818 59

1520 1.792 59

1579 1.760 59

1670 1.737 59

1645 1.724 59

1673 1.709 59

1713 1.687 59

1736 1.674 59

1763.6 1.660 59

1803.5 1.637 59

1859 1.607 59

1929 1.569 59

Transitions o/K2 S04 Potassium

In the short time of TGA runs [53, 73], K2 S04 appears stable to 1173 °K. A slight loss in weight

by 1273 °K was attributed to sublimation [73].

Spitsyn and Shostak [125], in maintaining K2 S04 for 2 hr at each of various temperatures, found no

loss in weight at 1073 and 1173 °K, a barely measurable loss at 1273 °K, and 3.6 percent loss at 1473 °K.

Because the residue showed no alkaline reaction

or change in percentage composition, even on

repeating the heating in a stream of steam, they

concluded that all the weight loss was due to volatilization and none to decomposition. They

believed they had ruled out the possibility that

K2 0, known to be volatile at high temperature,

might have formed by decomposition but escaped

as fast as SO3. On the other hand, Kroger and

Stratmann [67], by an application of the Knudsen

effusion method, reported both decomposition

pressures (1114—1272 °K) and sublimation pressures

(1085-1272 °K), the former being between one and

two orders of magnitude higher than the latter. An isolated value [21] for the equilibrium decomposition pressure of K2 S04 at 1223 °K by the static

method is two orders of magnitude higher than

that of Kroger and Stratmann.

These discrepancies could be resolved if it were

possible to calculate the decomposition pressures

thermodynamically, but the necessary high-temperature thermal data are available only for K2 S04 [2], not K2 0. Further, the 298 °K entropy of K2 0

is unknown.

Phase change T AH AS References

c, II -^c, I c, 1-^ 1.

°K

856

1342

kj mol~x

8.12

36.65

J mol~x deg~x 9.50

27.

3

2

1, 2, 5

1, 2, 5

Thermodynamic data (298.15 °K)

s°

J mol-1 cal mol-1 References

deg-1 deg"1

K2 S04(II) 175.7 42.0 1, 3,8

k2 o

ko2 116.7 27.9 3, 8

AHf°

kj mol-1 kcal mol-1 References

K2 S04 (II) - 1436.9 -342.66 1, 8 k2 o -361.5 -86.4 1, 8 ko2 -280 . -67. 1 -283. -67.7 8

24
Praseodymium

For high-temperature properties, see “Rare

Earths.”

Density of Pr2(S04 )3

Phase T d References

°K

c 297 3.797 17

Thermodynamic *data (298.15 °K)

s°

kj mol-1 kcal mol-1 References

Pr2 (S04)3 Pr2 03 -1827.6 -436.8 8 PrOi.72 -937.2 -224.0 8

PrOi.833 -952.3 -227.6 8 Pr02 -1004.2 -240.0 8

Radium

Thermodynamic data (298.15 °K)

S°

J mol-1 cal mol-1 References

deg"1 deg-1

RaS04 142. 34 1 RaO

AHf°

kj mol-1 kcal mol-1 References

RaS04 - 1473 . -352 1 RaO -523. -125 1

Rare Earths

On heating, all of the anhydrous sulfates of the

trivalent rare-earth elements and yttrium, type

formula R2 (S 04)3 , decompose, without first melting,

to basic salts (oxysulfates) of the type R2 03 -SC>3 , then to an oxide. The oxide final product is R2O3

for all these elements except cerium, praseodymium,

and terbium, which, in air, at 1 atm pressure, form

Ce02, PreOn, and the somewhat variable oxide

usually written Tb407 .

R2(S04)3 ^R2 03 -S 03 + 2S03 (a) R2O3 • SO3 —* R2O3 + SO3. (b)

There have been three comprehensive studies

which, taken together, reveal the trend in stability

of the sulfates and oxysulfates within the rare-earth

series. Wohler and Griinzweig [142, Ce] [143,

others] measured equilibrium decomposition pres- sures for reaction (a) for scandium, yttrium, and 8

of the rare-earth elements, in the range 973 to 1293 °K. Nathans and Wenlandt [83] performed

TGA and DTA, covering reactions (a) and (b) for

yttrium and all the rare earths. Ivakin [58], by a weight-loss technique in a stream of dry N2 , working with yttrium and 7 of the rare earths, investigated the kinetics of reaction (a) at 1223 and 1323

°K and of reaction (b) at 1573 °K. From these

three sources one may conclude that with the

exception of Ce the stability of the sulfates in

reaction* (a) decreases from La to somewhere in

the Sm Eu Gd group, and then increases to Lu,

while the stability of the oxysulfates in reaction

(b) decreases all the way from La to Lu. These

conclusions hold for either kinetic or equilibrium

definitions of “stability,” and they ignore a few

reversals of order which are probably attributable

to experimental errors in view of the close similarity of adjacent members of the rare-earth series.

Among the more stable sulfates the order of decreasing stability in reaction (a) is La-Y-Lu (or Yb)

according to kinetics [58], equilibrium pressures

[143], and DTA [83], but not by TGA [83]. In re- action (b) Y falls near Ho and Er [58,83] in accord- ance with the crystal radii of R+3 ions and the

lanthanide contraction.

By the equilibrium decomposition pressure cri- terion, Sc2 (S04)3 is less stable than any of the

rare-earth sulfates except cerium. However, unlike the rare earths, it seems to decompose directly

to Sc2 03 [143] rather than through reactions (a) and (b). Ce2 (S04)3 is much less stable than the sulfates

of the other rare earths [83, 142, 143], and less

stable than Sc2 (S04)3 [143]. The product is Ce02 rather than Ce2 03, but the mechanism is unknown.

Least stable by far is Ce(S04)2 which also yields

Ce02 ultimately. Wohler, Pliiddemann, and

25
Wohler [144] measured equilibrium decomposition pressures in the range 634 to 755 °K in

what they believed to be the system 2Ce(S04 )9 ^Ce2(S04)3 + S03 + l/2 02 . Duval [34] by TGA

found an intermediate 3Ce02 -4S03 . Ivakin’s kinetic investigation [58] showed that

the oxysulfates in reaction (b) decompose in ac- cordance with a simple zero-order rate law, whereas

reaction (a) is more complex kinetically. Thus

a plot of the equation

1 — (1 — a)113 = kt

is a line whose slope changes at one point from

k\ to k>. This would suggest a compound intermediate between R2 (S04 )3 and R2 03 *S03 , but x-ray

diffraction patterns were interpreted as showing

that the only possible intermediate was a solid

solution of variable composition.

The relative stabilities of the sulfates of Sc, Y,

and the rare earths may be represented in the

following order of decreasing stability:

La-Y-Lu Gd-Sc-Ce(III)-Ce(IV).

This refers to decomposition of the normal anhydrous sulfate to its first decomposition product,

oxysulfate or oxide as the case may be for each

element. The dashed line preceding Gd represents the other rare-earth elements. With increasing atomic number the stability first decreases in

the order Pr-Nd-Sm-Eu and then increases in the

order Tb-Dy-Ho-Er-Tm-Yb.

The high-temperature behavior of the bivalent

sulfates SmS04 , EuS04 , and YbS04 has not been

investigated. They are expected to undergo self- oxidation, like FeS04 , to yield ultimately R2 03 . Rhodium

The compound Rh2 (S04 )3 is difficult to prepare

in pure form. The reported density needs confirmation.

Density o/Rf^lSCLL

Phase T d Refer- ence

c °K

291 3.44 17

Rubidium

In the short time of TGA runs [41, 73], Rb2 S04 appears stable to 1173 °K. A slight loss in weight

by 1273 °K was attributed to sublimation [73].

Spitsyn and Shostak [125], in maintaining Rb2 S04 for 2 hr at each of various temperatures, found no

loss in weight at 1073 °K, a very slight loss at 1173°K,

0.3 percent loss at 1273 °K, and 6.3 percent loss at 1473 °K (22.4% in 8 hr). They believed this was

due entirely to volatilization, none to decomposition. But see the discussion under Potassium.

It is not yet possible to calculate decomposition

pressures thermodynamically. Neither entropies nor high-temperature thermal data are available

for Rb2 S04 or Rb2 0.

Density o/Rb^SCL

Phase T d Refer- ences

c, II °K293 3.614 132

c, II 298 3.613 7

c, II 333 3.594 133 Liquid 1359 2.538 59

1385 2.521 59

1418 2.499 59

1468 2.466 59

1508 2.440 59

1562 2.403 59

1617 2.367 59

1670 2.331 59

1688 2.319 59

1755 2.275 59

1818 2.233 59

Transitions o/Rb-iSO-i

Phase change T Refer- ences °K

c, II^c, I 923 1 c, I—> 1. 1347 1

Thermodynamic data ( 298.15 °K)

S°

J mol 1 cal mol-1 Refer- deg”1 deg” 1 ences

Rb>S04 Rb2 0

&Hf°

kj mol 1 kcal mol 1 Refer- ences

Rb,S04 -1424.7 -340.5 1, 8 Rb.O -330. -78.9 1, 8

26
Ruthenium

No simple sulfates have been isolated. The

compound RuOS04 has not been prepared in welldefined form.

Samarium

For high-temperature properties, see “Rare

Earths.”

Density

Phase T d References

°K

SmS04(c) Sm2 (S04 )3(c) 297 4.059 17

Thermodynamic data (298.15 °K

S°

J mol 1 cal mol 1 References

deg-1 deg"1

SmS04 Sm2(S04 )3 Sm2 03

melts within this range. For a rough thermodynamic treatment of these data see [60]. These

pressures cannot be calculated thermodynamically

for lack of high-temperature thermal data for

Ag2 S04 , though such data are available [2] for Ag

(and for Ag2 0). The other data that would be re- quired are tabulated below.

For the transition temperature, a higher value

703 ±3 °K was obtained by [52].

Density of Ag2S04

Phase T d Refer- ence °K

c, II 298 5.363 145

Transitions o/AgjSO-i

Phase change T AH AS Refer- ences

°K kj moD1 J mol~l deg-1 c, II -*c. I 685 8.0 11.7 1, 5; see text

c. I->1. 933 1 7 . 1 7 . 1, 5

Thermodynamic data (298.15 °K)

AHf°

kj mol 1 kcal mol-1 References

SmS04 Sm2 (S04 )3 -3884.8 -928.5 82 Sm2 03 -1815.4 -433.9 8,82

Scandium

s°

J mol 1 deg- 1 cal mol-1 deg- 1 Refer- ences AgoSCMH) 200.4 47.9 3, 8 AgoO 121.3 29.0 3. 9. 8 Ag 42.70 10.20 3. 9. 8

ahp

For high-temperature properties.

Earths.”

Density of Sc2 (S04 )3

see “Rare

Phase T d References

°K

c 297 2.605 17 c 298 2.579 7

kj mol 1 kcal mol- 1 Refer- ences

AgoSCMH) -713.37 -170.50 1, 8 Ag.0 -31.05 -7.42 9 Ag 0.000 0.000

Sodium

Silver

Ag2 S04 decomposes above its melting point to Ag metal. In TGA [53] it has appeared stable to 1063 °K, but not beyond. Marchal [76] measured

decomposition pressures by the static method in the range 1090 to 1493 °K. The Ag metal product

In the short time of TGA runs, Na2 S04 appears

stable to beyond 1173 °K [39, 53]. Spitsyn and

Shostak [125], in maintaining Na2S04 for 2 hr at each of various temperatures, observed no weight

loss at 1073 and 1173 °K, 0.04 percent at 1273 °K

and 1.05 percent at 1473 °K. The composition of

the residue showed that both decomposition and

27
volatilization had occurred. Kroger and Strat- mann [67], by the Knudsen effusion method, measured the low-vapor pressure of Na2 SC>4 in the range

1025 to 1124 °K, but not the decomposition pressure. For comparison with their own measure- ments on K2 SO4 , they present a curve of the decomposition pressure of Na2S04 based on fragmentary data of Marchal [1929] and Terres [1954] in

the range 1473 to 1623 °K with extrapolation down

to about 1140 °K. But these data must be in error; near 1100 °K, the decomposition pressure would be

higher by a factor of 107 than that calculated thermodynamically.

The decomposition pressures in the tables below

are based on high-temperature thermal data from

[2] and thermodynamic data for 298.15 °K as tabulated. Na2 S04(V) is the mineral thenardite.

The nomenclature for the crystal forms is that of

Kracek [66]. The value for the heat of formation

of Na2 S04 , —330.90 kcal mol-1

, in [1] was calculated there from data of Pitzer and Coulter [113]

who used Na2 S04(V). According to Coughlin [26]

there is a constant A// = 0.74 kcal mol-1 between

types V and III from 298 to 450 °K, whence the

heat of formation of Na2 S04(III) at 298.15 °K is taken to be

Thermodynamic data (298.15 °K)

S°

J mol 1 cal mol 1 References

deg" 1 deg" 1

Na2 S04 (V) 149.5 35.73 1, 3, 8 Na2 S04 (III) 156.5 37.4 3 Na2 0 72.8 17.4 1 76.1 ±4 18.2 ± 1.0 3 Na2 02 94.6 22.6 3, 8 Na02 116. 27.7 3, 8

AHf°

kj mol 1 kcal mol-1 References

Na2 S04 (V) - 1384.5 -330.90 1 Na2 S04 (III) -1381.4 -330.16 see text Na2 0 -416. -99.4 1 , 10 Na2 02 -504.6 - 120.6 1 Na02 -259. -61.9 1 — 330.90 + 0.74 = — 330.16 kcal mol 1 = —1381.4 kj mol-1

. The entropy of Na20 is uncertain. [3] has an esti- mated value:

18.2 ±1.0 cal mol-1 deg-1 (=72.8 J mol-1 deg-1

).

Density o/Na2 S04

Phase

' T d Refer- ences

°K

c, V 298 2.664 ±0.001 66

c. III 298 2.697 ±0.001 66

c, I 993 2.64 g/cm3 (crystallographic) 123 Liquid 1173 2.061 59

1218 2.039 59

1263 2.017 59

1305 1.997 59

1350 1.971 59

Transitions. of Na2 SC>4

Phase change T AH AS Refer- ences

°K kj mol~x J mol~x deg~ l c, V—> c, III 450 3.1 0 6.9 5

c, III-+C, I 515 7.49 14.56 5

c, I-> 1 . 1157 23.85 20.63 5

Decomposition of Na2 SC>4

(1) Na,S04(c) ^ Na2 0(c) + S03(g) from Na2 S04 (III): A//|98 15 = 570.32 kj, 136.3, kcal from Na2 S04 (V): AH°29e ,5 = 573.4, kj, 137.05 kcal

(2) Na,S04 (c)?± Na,0(c) + SO,(g)+ 1/2 02 (g) from Na2 S04 (III): Atf 2°98.,5 = 668.73 kj, 159.83 kcal from Na2 S04 (V): AA/^g.,

5

= 671.8, kj, 160.5? kcal.

A. Free energies

T Na2 0 Na2 SC>4 Reaction (1)

fef fef Afef AG°t

°K J J J kj

298.15 -72.80 - 156.48 - 172.99 + 518.74

(HI) 400 -75.65 -161.77 -172.62 501.27 500 -81.04 -172.51 -171.43 484.61 514 -174.22 - 170.90 482.48 (HI-* I) 600 (I) -87.28 -187.36 - 167.68 469.71 700 -93.68 -202.34 -164.18 455.39 800 -99.87 -216.81 - 160.97 441.55 900 - 105.90 -230.62 -158.15 427.99 1000 -111.63 - 243.80 - 155.49 414.83 1100 -117.19 - 256.44 - 153.04 401.98

298.15 - 149.49 -179.97 + 519.76

(V) 350 -151.13 -179.96 510.43 400 - 154.77 - 179.63 501.57 450 - 159.70 - 179.29 492.74

28
B. Equilibrium constants

T log Ki log K2 Ko

°K

298.15 -90.877 1.33 X10-91 -103.206 6.22 XlO”104 (HD

400 — 65.454 3.52 X10-66 - 73.292 4.06 XlO”74 500 -50.622. 2.39X10”51 -55.980 1.05 XlO”56 514 -49.026 9.42 xlO"50 -54.106 7.83 xlO-55 600 -40.889 1.29 xlO"41 — 44.535 2.92 xlO-45 700 -33.979 1.05x10”34 - 36.402 3. 96 xlO”37 800 -28.828 1.49 X10”29 -30.336 4.61 xlO-31 900 -24.838 1.45 XlO”25 -25.636 2.31 xlO”26 1000 -21.667 2. 15 XlO-22 -21.901 1.26 xlO-22 1100 -19.085 8.22 XlO”20 -18.859 1.38X10- 19

298.15

(V) 350

-91.054 8.83 XlO-92 - 103.383 4.14X10-104

-76.171 6.75 xlO-77 -85.946 1.13 XlO-86 400 -65.492 3.22 XlO”66 -73.430 3.72 xlO-74 450 -57.191 6.44X10-58 -63.697 2.01 xlO-64

C. Equilibrium pressures (atm)

T S02 o2 so3 Total

°K

298.15 1.98 x 10-69 9.89 x 10-70 1.33 x 10-91 2.97 x 10-69 am

400 1.49 x 10-49 7.44 XlO”50 3.52 x 10-66 2.23 x lO”49 500 6.03 x 10-38 3.02 x 10-38 2.39 x 10-51 9.05 X lO”38 514 1.07 x 10-36 5.35 x 10-37 9.42 x 10-50 1.61 X 10-36 600 2.57 x 10-30 1.29 x lO"30 1.29 x 10-41 3.86 x !0-30 700 6.80 x 10-25 3.40 x 10“25 1.05 x 10-34 1.02 xlO-24 800 7.52 x 10-21

' 3.76 x 10-21 1.49 x 10-29 1.13 xlO”20 900 1.02 x 10- 17 5.11 x 10- 18 1.45 x 10-25 1.53 x lO" 17 1000 3.16 xlO-15 1.58 xlO-15 2.15 x lO"22 4.74 x lO" 15 1100 3.37 x 10- 13 1.69 x 10- 13 8.22 x lO"20 5.06 x 10-13

298.15 1.51 x 10-69 7.54 x 10-70 8.83 xlO-92 2.26 x lO"69 (V) 350 6.35 x 10-58 3.18 x 10-58 6.75 XlO"77 9.53 X 10"58 400 1.40 x 10-49 7.01 x 10-50 3.22 x lO”66 2.10 x 10-49 450 4.32 x 10-43 2.16 x 10-43 6.44 x 10“58 6.48 x 10-43

Strontium

SrS04(II) is the mineral celestite (celestine).

There has been some disagreement on the thermal

stability of SrS04 . Gmelin quotes Grahmann

[1913] as finding a 4 percent loss in weight at 1273

°K, while [106] shows a TGA curve level only to 573 °K, then falling off gradually for a total weight

loss of 1 percent by 1157 °K. On the other hand,

[53] found no weight loss by TGA to beyond 1273

°K, and [90] by a special TGA technique found a "decomposition temperature” of 1647 °K (0.6%

loss in 15 min after rapid heating to just below this

temperature). Finally, two DTA curves [22, 50] on celestite to 1473 and 1573 °K respectively show

notable deflections only for the transition II —» I, but no sure sign of a thermal effect attending

decomposition.

It is not yet possible to calculate decomposition

pressures thermodynamically because high-temperature thermal data are lacking for SrS04 , though available for the SrO product [2]. The

other necessary data are hsted below. The en- tropy of SrS04 is uncertain. [3] does not accept

the value hsted by [1], but has an estimated 28.2

± 1.0 cal mol-1 deg-1 (118.0 ±4 J mol-1 deg-1

) taken

from the earlier [60]. We would estimate 28.7 cal

mol-1 deg-1 (120.1 J mol-1 deg-1

).

Density o/SrSO-t

Phase T d Refer- ence

c, II °K

ambient 3.91 129

Transitions of SrS04

Phase changes T Refer- ences

°K

c. II -> c, I 1425 1. 5 6 7 1878 1, 5

Thermodynamic data (298.15 °K)

S

o

J mol-1 cal moR1 Refer- deg-1 deg-1 ences

SrS04 121.8 29.1 1 118.0 28. 1±1 .0(2.0) 3, 8 SrO 54.4 13.0 1, 3. 8

AHf°

kj mol 1 kcal mol 1 Refer- ences

SrS04 -1445. -345.3 1. 8 SrO -590.4 -141.1 1. 8

Terbium

For high-temperature properties, see “Rare

Earths.”

Density o/Tb2(S04 )3

Phase T(°K) d Refer- ences

29
Thermodynamic data (298 °K) Transitions

S°

J mol 1 cal mol 1 - Refer- deg deg ences Tb2 (S 04 ) ;i Tb.O.i TbOi.71 TbOi.s

AHf°

kj mol 1 kcal mol 1 Refer- ences

Tb2 (S04)3 Tb,03 - 1827.6 -436.8 8 TbOi 71 -934.3 -223.3 8 TbOi.8 -947.3 -226.4 8

Thallium

Tl(l). According to most of the older literature, as collected in the handbooks of Gmelin, etc.,

TI2 SO4 may be heated as far as the melting point

without noticeable decomposition or volatilization,

after which it is volatile without decomposition,

although one author reported formation of TI2 O3 . More recently two TGA curves have shown contradictory results. One [108] has a plateau as far as 628 °K, then loss of weight interpreted as for- mation of TI2 O3 , not quite complete by 1153 °K.

The other [53] shows a plateau to 1000 °K, then

rapid volatilization.

Favorskaya and Ponomaryeva [44] measured by

a dynamic method vapor pressure of TI2SO4 in

the range 973 to 1233 °K, finding no evidence of

decomposition. Decomposition pressures cannot

be calculated thermodynamically for lack of any

high-temperature thermal data. Even the data at 298.15 °K are quite fragmentary. The entropy

of T12 0 listed by [ 1 ] is ignored by [3]. Tl(in). Little is known about TblSCbL- It has

been reported (see Gmelin, etc.) to decompose

readily to TI2SO4.

Densities

Phase T d Refer- ences

TbSO.|(c)

°K

298 6.77 ±0.02 46

Tl2 S04 (c) 303 6.768 ±0.003 24

T12(S04) 3(c) 288 5.22 16

Phase change T AH AS Refer- ences

Tl.SO-i, II -> I T12 S04 . c * 1.

°K

765 905

kj mol~ l

23

J mol~ x deg- 1

25. s 461

Thermodynamic data (298.15 °K)

s°

T1,S04 Tl2 (S04 ) 3 T1,0

T12 03

J mol 1 deg" 1 cal mol 1 deg” 1 Refer- ences

99.0 23.8 1

AHfO

kj mol- 1 kcal mol-1 Refer- ences

ti,so4 -927.6 -221.7 1 ti,(S04 )3 T1,0 — 175.3 -41.9 1 T1,03

Thorium

Th(S04)2 is stable to at least 673 °K, the temperature commonly used in preparing the anhydrous salt from the hydrates. At higher temperatures it decomposes before a melting point is reached, to ThCL without passing through an

oxysulfate intermediate [144, 78]. The equilibrium decomposition pressure measurements

(848-1058 °K) of Wohler, Pliiddemann, and

Wohler [144] have been superseded by the better

values of Mayer et al. [78], (907-1057 °K). The

latter combined their pressure data with their

dropping-calorimeter heat-capacity measurements

(623-897 °K) to obtain analytic expressions for AG°,

A5°, A//°, and log K for the decomposition reaction.

It is from these that we have calculated the pres- sures tabulated below. [ 2] has high-temperature

thermal data for TI1 O2 but not for ThlSCLV

Density o/Th(S04 )2

Phase T d References

°K

c 291 4.37 16

30
Transitions <?/Th(S04 )2 C. Equilibrium pressures (atm)

Phase change References

c —> c None, 297-1057 °K 78

c * 1

Decomposes

Thermodynamic data (298.15 °K)

S°

J mol 1 cal mol-1 References

deg" 1 deg" 1

Th(S04 )2 148. 35.4 78 Th02 65.3 15.6 3, 8

AHf°

kj mol 1 kcal mol-1 References

deg-1 deg-1

Th(S04 )2 -2541.4 -607.4 78 Th02 -1226.7 -293.2 8

T S02 O2 so:i Total

°K

298.15 400 500

600

700 800

900

1000

1100

1200

5.47 X IO

-32 2.90 x 10-21 5.38 x 10- 15 7.84 x 10- 11 7.09 x 10-8 1.12 x 10-5 5.50 x 10-4 1.19 x 10-2 1.40 x 10- 1 1.05

2.74 x IO

-32 1.45 x 10"21 2.69 x 10- 15 3.92 x 10- 11 3.55 x 10-8 5.60 x 10-6 2.75 x 10-4 5.92 x 10-3 7.00 x 10-2 5.26 x 10- 1

1.93 x IO

-35 9.55 x 10-24 6.35 x 10- 17 2.17 x 10-' 2 3.54 x 10-9 8.54 x 10-7 5.73 x 10"5 1.56 x 10-3 2.20 x 10-2 1.88 x 10->

8.21 x IO

- 32 4.35 x 10-21 8.13 x 10- 15 1.20 x 10- 10 1.10 x 10- 7 1.77 x 10-5 8.82 x 10-4 1.93 x 10"2 2.32 x 10- 1 1.77

Thulium

For high-temperature properties, see “Rare

Earths.”

Density o/Tm2 (S04 ) 3

Phase T d References

°K

c 297 3.906 17

Decomposition of Th(S04 )2

(1) ThfS04 )2(c)^Th02 (c) + 2S03 (g) A// 2,8 15 = 524.2 kj, 125.3 kcal

(2) Th(S04 >»(c) Th02(c) + 2S02 (g) + 02 (g)

- A//L» 13 = 720.9 kj. 172.3 kcal.

B. Equilibrium constants

T log Ki Kx log K2 Kz

°K

298.15 -69.429 3.72 x IO- 70 -94.087 8.19 x IO

-95 400 -46.040 9.12 x IO

-47 -61.916 1.21 x lO"62 500 -32.394 4.04 x IO

-33 -43.110 7.76 x IO-44 600 -23.326 4.72 x 10-24 -30.618 2.41 x IO

-31 700 - 16.903 1.25 x 10- 17 -21.749 1.78 x 10-22 800 -12.137 7.30 x 10- 13 -15.153 7.03 x 10- 16 900 -8.484 3.28 x 10-9 - 10.080 8.32 x 10- 11 1000 -5.612 2.44 x 10-6 -6.080 8.32 x 10-7 1100 -3.314 4.85 x 10-4 -2.862 1.37 x 10-3 1200 -1.450 3.55 x 10-2 -0.234 5.83 x 10- 1

Tin

Sn (il). At room temperature S11SO4 is stable

for several months in air and much longer in the

dark in vacuo, but on heating it undergoes self- oxidation to Sn02 + SC>2 . Thus above 633 °K in

CVfree N2 complete conversion has been found at 651 °K, but little decomposition below this temperature [31]. Similarly, on heating in air for 2 hr at each of several temperatures, no change was observed at 573 °K, SnS04 + SnC>2 (but no oxysulfate)

at 623 °K, and only Sn02 remained at 673 °K; the

reaction became fast between 623 °K and 673 °K

[57]. In TGA with continuous evacuation, heating

at 100°/hr, decomposition began at 593 °K and was

complete to SnC>2 by about 870 °K [28].

Sn(iv). Anhydrous Sn(S04)2, very hygroscopic,

becomes thermally unstable not far above 373 °K

[51]. However, in vacuum TGA, heating at 90°/hr

to 573 °K, then at 150°/hr, decomposition first became noticeable near 573 °K with complete conversion to SnC>2 by about 860 °K [28].

No calculation of decomposition pressures is possible. High-temperature thermal data exist

only for Sn(>2 [2], and the 298.15 °K data are incomplete as shown below.

31
Densities Thermodynamic data (298.15 °K)

Phase T d References

SnS04(c) °K

ambient 4.185 g/cm3 (crystal- 118 Sn(S04 ),(c) 291

lographic)

3.45 16

Thermodynamic data (298.15 °K)

s°

J mol 1 cal mol 1 References

deg-' deg-1

SnSO-i

Sn(S04 ), SnO 56.5 13.5 1, 3, 8 SnO, 52.3 12.5 1, 3

\Hf°

kj mol 1 kcal mol 1 References

SnS04 Sn(S04 ), - 1646. -393.4 1, 8 SnO -286., -68.4 1, 8 SnO, -580.7 -138.8 1, 8

S°

Ti2 (S04 )3 Ti(S04 )2 TiOS04 Ti02 (rutile) Ti02 (anatase)

J mol-1 deg- 1 cal mol-1 deg- 1 References

50.38

49.92

12.04

11.93

3, 4

1, 3, 4, 10

AHj O

kj mol-1 kcal mol-1 References

TKSO.h

Ti(S04 ) 2 TiOS04 Ti02 (rutile) -943.5 -225.5 10 Ti02 (anatase) -912.5 -218.1 10

TiOi.!»_ 2 .o -943.5 -225.5 8

Titanium

Ti (II). TiSO-i, uncertain existence.

Ti (III). Ti2 (S04 )3 Thermal properties unknown

except that it yields TiO, on ignition.

Ti (IV). Anhydrous Ti(S04 )2 has only recently

been prepared [51, 23]. Hayek and Engelbrecht

[51] found that it decomposed in 2 steps, beginning

at 423 °K, ultimately yielding Ti02 . Chuklantsev

[23], by a procedure resembling TGA, heating at 1.7°/min in dry argon, found decomposition beginning at 423 °K to produce TiOS04 ; then near 703 °K

this began decomposing to Ti02 , complete by about

873 °K. The decomposition pressures reported by

[144] for TiOS04 (741-853 °K) are not to be trusted.

No pressures can be calculated thermodynamically

because high-temperature thermal data are available

only for Ti02 [2].

Densities

Phase T d References

°K

Ti(S04 ),(c) 293 2.47 ±0.16 23 TiOS04 (c) 293 3.53 ±0.21 23

Uranium

U (ill). Difficult to isolate from solution; oxidized

readily by H2 0.

U (iv). It is not certain that a pure anhydrous

U(S04 )2 has been obtained, although hydrates are well known. The latter, on heating, are said to yield first the anhydrous salt (near 570 °K), then

U02 S04 even in the absence of air. U (vi). Anhydrous U02 S04 and several of its hydrates are known. The anhydrous salt is suffi- ciently stable at 773 °K to permit its preparation by

slowly heating the hydrates to that temperature.

At high temperatures it decomposes, without melting, to an oxide, the stoichiometry of which depends

on conditions, varying between UO3 and U02 . Formulas such as U3 O8 and U4 09 are approximated.

Because of this nonstoichiometry, Owens and

Mayer [92] chose not to attempt equilibrium decomposition-pressure measurements. They report

values of Ht—H2 9h for U02 S04 in the range 373 to 820 °K. These high-temperature thermal data

could be used along with those in [2] for U02 and

UO3 to calculate two sets of decomposition pressures

between which observed values would usually lie, but this hardly seems worth while in view of the

uncertainty in the U02 S04 entropy at 298.15 °K

upon which such calculations would strongly depend.

32
Densities Thermodynamic data (298.15 °K)

Phase T d References

U(S04 )2(c) °K

291 4.60 16 U02 S04(c)

Thermodynamic data (298.15 °K )

s

O

J mol-1 deg-1 cal mop1 deg-1 References

U(S04)* uo2 so4 155 ±25 37 ±6 92 u3 o8 282.4 67.5 3, 116a uo2 11.% 18.63 1, 3, 4, 116a, 8 uo3 98.66 23.58 1, 3, 4, 116a, 8 u4 o9 336 80.3 3, 116a, 8

AHf°

kj mol 1 kcal mol-1 References

U(S04 )2 -2259 -540 116a uo2 so4 -1888 ±6 — 451.2 ±1.5 92 u3 o8 -3574 -854.1 116a uo> -1084 -259.2 4, 116a, 8 uo3 -1226 -293 116a u4 o9 -4510 -1078 116a

Vanadium

s°

J mol-1 cal mol-1 References

deg-1 deg” 1

voso4 V2 05 131 31.3 1,3, 4, 8

\Hf°

kj mol-1 kcal mol-1 References

voso4 - 1307.5 -312.5 I v2 o5 -1561 -373 1,4

Ytterbium

For high-temperature properties, see “Rare

Earths.’'

Density

Phase T d References

°K

YbS04 (c) Yb^SO^c) 297 3.964 17 Yttrium

For high-temperature properties, see “Rare

Earths.”

V (II). Only the hydrate VS04 * 7H2 0 is known.

V (ill). Rivenq [119] described the preparation

and pyrolysis of anhydrous V2 (S 04)3 . On heating in

vacuo, the compound was stable 573 to 598 °K but

decomposed above 653 °K. According to the

author’s interpretation of weight-loss curves, the

decomposition is complicated, its course dependent

on temperature. The final product below 673 °K was V2 O5 , while above 681 °K a mixture of V2 O4 and

V2 O5 was formed. The results need confirmation

by other methods.

V (IV). Anhydrous VOSO4 decomposes, without

melting, above about 673 °K, with oxidation to V2 O5 , SO3 , and S02 . The reaction is reversible. Neumann and Sonntag [86] measured equilibrium decomposition pressures in the range 704 to 826 °K.

For a thermodynamic treatment (sigma function) of

these data see [60]. These pressures cannot yet

be calculated thermodynamically because hightemperature thermal data are available only for

V2 O5 [2], and the entropy of VOSO4 is unknown.

V (v). Anhydrous V2 05

• 2S03 , V2 O5

• 3S03 , and

V2 O5• 4S03 are known. The preparation of the

last of these has been described by Hayek and

Engelbrecht [51] who found it to be unstable 373

to 473 °K.

Density ofY2 ( 804)3

Phase T d References

°K

c 297 2.949 17

Thermodynamic data (298.15 °K)

S°

J mol 1 cal mol-1 References

deg deg

Y2 (S04 )3 y2 o3

AHf°

kj mol-1 kcal mol-1 References Y2 (S04 )3 y2 o3 -1905.4 -455.4 8

33
Zinc Density of ZnS04 Hoschek [57] quotes Shargorodski [1949] as finding a transition a (3 at 1008 °K, and Gmelin ref- erences Friedrich [1911] as reporting one at 1013 °K.

Gruver [50] published DTA curves of two samples

of ZnSO.i, one showing clearly a thermal effect

beginning at 1013 °K, the other none. [1] and [5]

do not list this transformation. Extrapolation of a binary phase diagram [62] yields a melting point

at 1003 °K, seemingly impossible. A melting point

has not been measured directly and probably cannot be because of decomposition.

Decomposition is through an oxysulfate intermediate, the formula of which has been variously

reported. A review is given by Hoschek [57]

whose own data yield the formula 2ZnO • 3ZnS04 . In TGA curves the oxysulfate either does not appear

definitely [22] or appears only as a nick, not a plateau [28, 53, 90, 127].

Terem and Akalan [127] report activation energies

for the two steps in the decomposition of ZnS04 , using a thermogravimetric technique. Between

1113 and 1173 °K, Ea for the reaction

5ZnS04 = 3ZnS04 • 2ZnO + 2S03

is 212 kj (50.6 kcal). Between 1233 and 1293 °K Ea for the reaction

3ZnS04

• 2ZnO = 5ZnO + 3S03

is 77.0 kj (18.4 kcal).

Pechkovsky and coworkers [103], on heating

ZnS04 under isothermal conditions between 1148

and 1163 °K, observed a change in slope for the

degree of decomposition versus time. From the

slopes the activation energy for the sulfate decomposition was found to be 271 kj (64.7 kcal), and for

the oxysulfate 385 kj (92.0 kcal). The latter value

differs markedly from that of Terem and Akalan

but it is not clear whether both studies are dealing

with the same oxysulfate.

Wohler, Pliiddemann, and Wohler [144] measured

decomposition pressures in the range 948 to 1073

°K, but these data are considered unreliable [60].

The pressures tabulated below are based on hightemperature thermal data in [2], and involve an

extrapolation of the free-energy function of ZnS04 above 1000 °K. Being calculated for the decomposition ZnS04 —»ZnO, they are intermediate between the somewhat higher pressures expected for

ZnS04 ^oxysulfate and the lower pressures for

Oxysulfate —> ZnO. The entropy of ZnS04 is uncertain. The value in [1] is too high. That of [3] used here is an estimate. The heats of formation

of ZnS04 and ZnO in [1] are subject to revision, but

had to be used because they are part of a self- consistent set.

Phase T d References

c 288 3.74, 7 c 298 3.546 81

Transitions of ZnS04

Phase change T References

°K

a—* fi 1013 see text

Thermodynamic data (298.15 °K)

S

O

J mol-1 cal mol-1 References

deg"1 deg"1

ZnS04 113 27.0 3 ZnO 43.64 10.43 3, 4, 8

AHf°

kj mol 1 kcal mol 1 References

ZnS04 -975.3 -233.1 ±0.2 11 -978.55 -233.88 1, 8 ZnO -348.0 -83.17 1, 8

Decomposition of ZnS04

(1) ZnSO,(c) ZnO(c) + S03(g) A//I98.i 5 = 235.4 kj, 56.26 kcal

(2) ZnS04(c) ZnO(c) + S02(g) +1/2 02 (g) A//|98 . 15 = 333.8o kj, 79.78 kcal.

A. Free energies

T ZnO ZnS04 Reaction (1)

fef fef Afef < of~

°K J J J kj 298.15 -43.64 -112.97 -187.34 + 179.54 400 -45.33 -116.98 -187.10 160.55 500 -48.66 - 124.73 - 186.82 141.98 600 -52.47 - 133.65 - 186.58 123.44 700 -56.36 - 142.95 - 186.28 105.00 800 -60.19 - 152.23 - 185.87 86.70 900 -63.85 -161.49 -185.23 68.68 1000 -67.36 -170.58 - 184.45 50.94

(1100) -70.71 - 179.43 - 183.56 33.47

(1200) -73.93 - 187.92 -182.72 16.129

34
B. Equilibrium constants Transitions o/Zr(S04 )2

T log Ki Ki log K2 K2

°K

298.15 -31.452 3.53 x lO"32 -43.781 1.66 x 10-44 400 -20.964 1.09 x 10-21 -28.902 1.25 x 10-29 500 -14.831 1.48 x 10-15 -20.189 6.47 x 10"21 600 -10.746 1.80 x 10-11 - 14.392 4.06 x 10- 15 700 -7.834 '1.47 x 10-8 -10.257 5.53 x lO"11 800 -5.661 2.18 x 10-6 -7.169 6.78 X lO"8 900 -3.986 1.03 x 10-4 -4.784 1.64 x 10-5 1000 -2.661 2.18 x 10-3 -2.895 1.27 x 10-3 (1100) -1.589 2.58 x 10-2 -1.363 4.34 x lO"2 (1200) -0.702 1.99 x 10- 1 -0.094 8.05 x 10- 1

C. Equilibrium pressures (atm)

T S02 o2 S03 Total

°K

298.15 8.19 x IO

-30 4.09 x 10"30 3.53 x IO

-32 1.23 x 10-29 400 6.80 X 10-20 3.40 x 10-20 1.09 x 10-21 1.04 x 10- 19 500 4.38 x 10" 14 2.19x10-14 1.48 x 10- 15 6.71 x 10~ 14 600 3.20 x 10- 10 1.60 x 10-10 1.80 x 10-11 4.99 X lO" 10 700 1.82 x 10-7 9.15 x 10“8 1.47 x 10-8 2.89 x 10-7 800 2.09 x 10-5 1.05 x lO"5 2.18 x 10-6 3.36 x 10"5 900 8.15 x 10-4 4.07 x 10-4 1.03 x 10-4 1.33 x 10"3 1000 1.48 x 10-2 7.40 x 10-3 2.18 x 10-3 2.44 x lO"2 (1100) 1.56 x lO"1 7.77 x 10-2 2.58 x lO"2 2.59 x 10- 1 (1200) 1.09 5.45 x 10_1 1.99 x lO"1 1.84

Phase change Refer- ences

c —* c None observed. 298-1050 °K

c * 1 . Decomposes

Thermodynamic data ( 298.15 °K

S°

J mol-1 cal mol-1 Refer- deg" 1 deg- 1 ences

Zr(S04 )2 133.1 31.8 estimated

Zr02 (a) 50.71 12.12 3, 4, 8

AHf°

kj mol-1 kcal mol 1 Refer- ences

ZrfSO-tb -2500. -597.4 1 (see text) Zr02 (a) - IO8O .3 -258.2 1 - 1085.- — 259.5 8 -1094.1 -261.5 4. 10 Zirconium

Von Hevesy and Cremer [135] attempted to measure equilibrium decomposition pressures of Zr(S04)2 in the range 849 to 905 °K, but were not sure of

obtaining true equilibrium. Decomposition was

direct to Zr02 , not through an intermediate oxysulfate, and the sulfate appeared to be less stable than

Hf(S04) 2 . A TGA curve [68] also shows no sign

of an oxysulfate.

We calculated decomposition pressures using

high-temperature thermal data for Zr(S04)2 from

[124] and for Zr02 from [2], estimating the entropy

of Zr(S04)2; but the pressures so obtained were

too low by far more than could be attributed to the

uncertainty in the entropy estimate. The difficulty

lies rather in the reported [1] heat of formation of

Zr(S04)2 .

Density of Zr(S04 )2

Phase T d Refer- ence

c °K

289 3.22 16

Relative Thermal Stability of Some

Sulfates

In order of increasing stability; criterion: equilibrium partial pressure of SO3 at 1000 °K. Note

that the products are not always the first actually

formed.

Sulfate Product PSO:.

A12(S04 )3 A12 03 atm

3.74 X 10-' CuS04 CuO 1.36 X 10-2 BeS04 BeO 1.19 X 10-2 ZnS04 ZnO 2.18 x 10-3 Th(S04 )-2 Th02 1.56 x 10-3 MnS04 Mn3 04 7.19 x 10-4 PbS04 PbO (yellow) 1.30 x 10-7 CaS04 CaO 4.74 x 10- 12 BaS04 BaO 1.45 x 10- 17 Na2 S04 Na2 0 2.15 x 10-22

35
Relative Thermal Stability of Some

Sulfates

In order of increasing stability; criterion: equilibrium partial pressure of SO3 at 298.15 °K. In all instances in this table, the immediate product is believed to be the metal oxide with no change in

oxidation number.

Sulfate f*soa

Sn(S04 )2 atm2 x 10- 13 A12 (S04)3 2 x 10~24 NiSO.i 6 x IO

-32 C0SO4 1 x IO

-34 Th(S04 ) 2 2 x IO

-35 MgS04 6 x IO

-40 CaS04 3 x 10-61 Li2 S04 3 x 10-69 SrS04 5 x IO

- 71 BaS04 3 x 10-80 Na2 S04 1 x 10-91 k2 so4 7 x 10_n0

Rb2 S04 6 x 10- 114 Cs2 S04 3 x 10- 115

Melting Points in Increasing Order

Compound MP MP

T1,S04 °K

905

°C

632 Ag2 S04 933 660

ZnS04 1003 730

Li2 S04 1132 859

Na2 S04 1157 884

CdS04 1273 1000 Cs2 S04 1277 1004 k2 so4 1342 1069 Rb,S04 1347 1074 PbS04 1360 1087 MgS04 1400 1127 BaS04 1620 1347 SrS04 1878 1605

Boiling Points

None have been measured. The following are calculated from the Law of Corresponding States

by Lorenz and Hertz [70].

Arrhenius Activation Energies for the

Thermal Dissociation of Some Inorganic Sulfates

Reactant Product Ea

Fe2 (S04 )3 Fe2 03 kj

83.3, 154. FeS04 Fe2 03 253. A12 (S04 )3 A12 03 268. CoS04 CoO 319. PbS04 PbO 385. ZnS04 2Zn0.3ZnS04 212, 271. CuS04 CuO.CuS04 212, 238. 2Zn0.3ZnS04 ZnO 77, 385. CuO.CuS04 CuO 346, 269.

Compound BP

°K

Na2 S04 1702 k2 so4 1971 Rb2 S04 1981 Cs2 S04 1900

We thank Miss Susan E. Meador for helping with

the literature search and translation of Russian

papers; W. H. Evans for helpful discussions and

the use of his extensive files; and John E. Ricci for

considering the applicability of the phase rule to the decomposition of MnSCL.

36
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