Rational Bureau 9f SttfflM

Library, E-01 Mmin- W*

dec 1 q196-

A 1 1 1 2 14 Si

I

t

NBS

PUBLICATIONS

NSRDS-NBS 30

SRDS-NB^ 11021 45929

100 -US73 V30.1969 C.1 NBS-PUB-C 1964 ligh Temperature Properties

and Decomposition of

Inorganic Salts

Part 2. Carbonates

U.S. DEPARTMENT OF COMMERCE

NATIONAL BUREAU OF STANDARDS

N SRD S
NATIONAL BUREAU OF STANDARDS

The National Bureau of Standards 1 was established by an act of Congress March 3, 1901. Today,

in addition to serving as the Nation’s central measurement laboratory, the Bureau is a principal

focal point in the Federal Government for assuring maximum application of the physical and

engineering sciences to the advancement of technology in industry and commerce. To this end

the Bureau conducts research and provides central national services in four broad program

areas. These are: (1) basic measurements and standards, (2) materials measurements and

standards, (3) technological measurements and standards, and (4) transfer of technology.

The Bureau comprises the Institute for Basic Standards, the Institute for Materials Research, the

Institute for Applied Technology, the Center for Radiation Research, the Center for Computer

Sciences and Technology, and the Office for Information Programs.

THE INSTITUTE FOR BASIC STANDARDS provides the central basis within the United

States of a complete and consistent system of physical measurement; coordinates that system with

measurement systems of other nations; and furnishes essential services leading to accurate and

uniform physical measurements throughout the Nation’s scientific community, industry, and commerce. The Institute consists of an Office of Measurement Services and the following technical

divisions:Applied Mathematics—Electricity—Metrology—Mechanics—Heat—Atomic and Molecular Physics—Radio Physics -—Radio Engineering -—Time and Frequency -—Astrophysics -—Cryogenics.2 THE INSTITUTE FOR MATERIALS RESEARCH conducts materials research leading to improved methods of measurement standards, and data on the properties of well-characterized

materials needed by industry, commerce, educational institutions, and Government; develops,

produces, and distributes standard reference materials; relates the physical and chemical properties of materials to their behavior and their irtteraction with their environments; and provides

advisory and research services to other Government agencies. The Institute consists of an Office of Standard Reference Materials and the following divisions: Analytical Chemistry—Polymers—Metallurgy—Inorganic Materials—Physical Chemistry.

THE INSTITUTE FOR APPLIED TECHNOLOGY provides technical services to promote

the use of available technology and to facilitate technological innovation in industry and Government; cooperates with public and private organizations in the development of technological

standards, and test methodologies; and provides advisory and research services for Federal, state, and local government agencies. The Institute consists of the following technical divisions and

offices: Engineering Standards—Weights and Measures — Invention and Innovation — Vehicle

Systems Research—Product Evaluation—Building Research—Instrument Shops—Measurement Engineering—Electronic Technology—Technical Analysis.

THE CENTER FOR RADIATION RESEARCH engages in research, measurement, and application of radiation tb the solution of Bureau mission problems and the problems of other agencies and institutions. The Center consists of the following divisions: Reactor Radiation—Linac Radiation—Nuclear Radiation—Applied Radiation.

THE CENTER FOR COMPUTER SCIENCES AND TECHNOLOGY conducts research and

provides technical services designed to aid Government agencies in the selection, acquisition,

and effective use of automatic data processing equipment; and serves as the principal focus

for the development of Federal standards for automatic data processing equipment, techniques,

and computer languages. The Center consists of the following offices and divisions: Information Processing Standards—Computer Information — Computer Services — Sys- tems Development—Information Processing Technology.

THE OFFICE FOR INFORMATION PROGRAMS promotes optimum dissemination and

accessibility of scientific information generated within NBS and other agencies of the Federal

government; promotes the development of the National Standard Reference Data System and a system of information analysis centers dealing with the broader aspects of the National Measurement System, and provides appropriate services to ensure that the NBS staff has optimum accessibility to the scientific information of the world. The Office consists of the following

organizational units: Office of Standard Reference Data—Clearinghouse for Federal Scientific and Technical

Information —Office of Technical Information and Publications—Library—Office of Public Information—Office of International Relations.

1 Headquarters and Laboratories at Gaithersburg:, Maryland, unless otherwise noted: mailing address Washington, D.C. 20234.

- Located at Boulder, Colorado 80302.

0 Located at 5285 Port Royal Road. Springfield, Virginia 22151.
UNITED STATES DEPARTMENT OF COMMERCE • Maurice H. Stans, Secretary

NATIONAL BUREAU OF STANDARDS • Lewis M. Branscomb, Director

High Temperature Properties and

Decomposition of Inorganic Salts

Part 2. Carbonates

K. H. Stern

Electrochemistry Branch, Naval Research Laboratory

Washington, D.C. 20390

and

E. L. Weise

U.S. Patent Office, Washington, D.C. 20231

NSRDS

NSRDS-NBS 30

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

CODEN: NSRDA

Issued November 1969

For sale by the Superintendent of Documents, U.S. Government Printing Office Washington, D.C. 20402 (Order by SD Catalog No. C 13.48:30), - Price 45 cents
NATIONAL BUREAU OF STANDARDS

JUN 9 1370

Qd /DO

, 0 45 73

C-efi 2>

Library of Congress Catalog Card Number: 66-60081
Foreword

The National Standard Reference Data System provides effective access to the quantitative data of physical science,

critically evaluated and compiled for convenience, and readily accessible through a variety of distribution channels. The

System was established in 1963 by action of the President’s Office of Science and Technology and the Federal Council

for Science and Technology, with responsibility to administer it assigned to the National Bureau of Standards. The System now comprises a complex of data centers and other activities, carried on in academic institutions and other

laboratories both in and out of government. The independent operational status of existing critical data projects is maintained and encouraged. Data centers that are components of the NSRDS produce compilations of critically evaluated data,

critical reviews of the state of quantitative knowledge in specialized areas, and computations of useful functions derived

from standard reference data. In addition, the centers and projects establish criteria for evaluation and compilation of data and make recommendations on needed improvements in experimental techniques. They are normally closely associated with active research in the relevant field. The technical scope of the NSRDS is indicated by the principal categories of data compilation projects now active or being planned: nuclear properties, atomic and molecular properties, solid state properties, thermodynamic and transport

properties, chemical kinetics, and colloid and surface properties.

The NSRDS receives advice and planning assistance from the National Research Council of the National Academy of Sciences-National Academy of Engineering. An overall Review Committee considers the program as a whole and makes

recommendations on policy, long-term planning, and international collaboration. Advisory Panels, each concerned with

a single technical area, meet regularly to examine major portions of the program, assign relative priorities, and identify

specific key problems in need of further attention. For selected specific topics, the Advisory Panels sponsor subpanels

which make detailed studies of users’ needs, the present state of knowledge, and existing data resources as a basis for recommending one or more data compilation activities. This assembly of advisory services contributes greatly to the guidance

of NSRDS activities. The NSRDS-NBS series of publications is intended primarily to include evaluated reference data and critical reviews

of long-term interest to the scientific and technical community.

Lewis M. Branscomb, Director.
Preface

During the past decade or two, inorganic salts, particularly in the liquid state, have assumed increasing importance

in a variety of applications. They are useful as reaction media, in metallurgical processes, and in electrochemical power

sources such as fuel cells and thermal batteries. Experimentalists and theoreticians have found molten salts an interesting subject for study since these ionic fluids offer an unusual opportunity for the study of short-range ionic interactions in the liquid state. So far, attention in this field has largely been focused on the alkali and alkaline earth halides since their stability at high temperatures is well known. The only decomposition which they can undergo is dissociation to the elements. The

extent of this dissociation can be easily calculated from existing thermodynamic compilations. For most inorganic salts the situation is more complicated. In many cases the decomposition reactions are not well defined and high-temperature thermodynamic and kinetic data are either lacking or scattered through the literature. Thus,

although the study of many salts would undoubtedly prove interesting and useful, they have received little attention because

in many cases not even the range of thermal stability is known.

The purpose of the present publication is to alleviate this situation by publishing in concise form 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, carbonates in the present volume.

Thermodynamic information has been heavily stressed because it is most important for dealing with the stability

problem, and because it constitutes the bulk of what is available. Whenever possible we have used in our calculations data already critically evaluated by others — references are given in the appropriate places — but we have also included

qualitative and semiquantitative information on the assumption that some knowledge is better than none. However, in these cases we have tried to warn the reader by posting appropriate “Beware!” signs.

Thermodynamic variables which have been included are:

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

enthalpies and entropies.

(b) Equilibrium constants and decomposition pressures, as well as relevant free-energy functions from 298.15 K

to as high a temperature as data exist. A/// and S° values of reactants and products at 298.15 K from which

the above functions are calculated are also given.

(c) Densities at 298.15 K and above. Since the kinetics of carbonate decomposition have been more extensively studied than those of any other class of inorganic salts, an examination of the relevant literature has revegled many regularities as well as delineated more sharply our areas of ignorance. For example, although it is not yet possible to control all the factors determining rate constants,

the conditions for which activation energies are well-defined, reproducible quantities can now be stated rather precisely.

Rather than cite the bulk of the extensive literature in this field, much of which is applicable to very special conditions, we have contented ourselves with a brief review of carbonate decomposition kinetics in which the common features rather than the peculiarities of individual compounds are stressed. If such peculiarities are of special interest, they are described

in the appropriate section. However, our bibliography is not intended to be exhaustive. In the preparation of this monograph we have consulted many papers not referenced here. Our aim has been to describe for the nonspecialist those features

of the decompositions which are generally agreed on, and to illustrate these with representative references of good

quality work.

IV
Contents

Page Foreword iii Preface iv The carbonates 1 A. Introduction 1

1.

The structure of the CQf ion 1 2. General features of carbonate decomposition I

3. Phase transitions 2

4. Density 2

5. Decomposition equilibria 2

B. Kinetics of thermal decomposition 4

1. General 4

2. Mechanism of endothermic solid state decomposition 4

(a) Kinetic equations 4

(b) Surface area changes during decomposition 6

3. The effect of pressure on rate 7

4. Activation energy 8

5. The effect of method of preparation on rate 8

6. The effect of heat transfer on rate 9

7. The influence of crystal defect struc- ture on the rate of decomposition... 9

Page C. Suggestions for experimental work 10

Units, symbols, and abbreviations 11

C02 , CO, and 02 11

Barium 12

Beryllium 13

Cadmium 13

Calcium 14

Cesium 15

Cobalt 15

Copper 16

Iron 16

Lead 16

Lithium 17

Magnesium 18

Manganese 19

Nickel 20

Potassium 20

Rare Earths 21

Rubidium 21

Silver 21

Sodium 22

Strontium 23

Thallium 24

Zinc 24

References 26

v
'
High Temperature Properties and Decomposition of Inorganic Salts

Part 2. Carbonates

K. H. Stern 1 and E. L. Weise 2

The literature dealing with the high-temperature behavior of inorganic carbonates has been

critically reviewed. Free energy functions, of reactants and products of the decomposition reactions were calculated and have been tabulated from 298 K up to as high a temperature as possible. Free

energy functions and equilibrium constants of reactions were tabulated. Auxiliary data on phase

transitions, densities, and kinetics of thermal decomposition have also been included. The literature of the endothermic decomposition kinetics of solids, as it applies to carbonates, has been reviewed.

Key words: Carbonates; thermal decomposition; thermodynamic functions.

The Carbonates

A. Introduction

1. The Structure of the CO3 Ion

The carbonate ion is a well-defined entity in which

the three oxygens are arranged in the same plane as the carbon atom at a distance of 1.30 ±0.01 A and

with bond angles of 120°. This value of the C—

O

distance, calculated by Pauling [103] on the basis

of the resonance structures

O

C

/ \

O: : O

: 0 :-

I C

/ \

: O : : O • +

: 0 :~

I C

/ \ + 0 : 0 :

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

per mole.

2. General Features of Carbonate

Decomposition

All carbonates decompose with the evolution of

CO2 , the decomposition product being either the

corresponding oxide or a basic carbonate, the

latter decomposing with further evolution of CO2 as the temperature is raised. The oxides of some metals

are unstable in the same temperature range as the

corresponding carbonates so that the decomposition occurs as the simultaneous reactions

MCO3 = MO+ C02(g)

MO= M(g)+ 1/2 02 (g).

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

2 U.S. Patent Office, Washington, D.C. 20231. Formerly of the NBS Institute for Basic Standards.

Consequently the vapor phase contains C02 , M, and

02 . We have not calculated the oxide equilibria but

have indicated under the particular element whether

it is important, such as for the higher atomic weight

alkali metals. Thus the melting points of pure

carbonates are generally not measurable directly,

though they may be estimated by extrapolation of

phase diagrams or by rapid heating, particularly

under high CO2 pressures, which minimizes decomposition. The dissociation, CO2 — CO + 1/2 O2 , is too slight at temperatures relevant for carbonate

decomposition to need taking into account (see

below).

As pointed out in out previous paper [121],3 the

thermal stability of a polyatomic anion decreases

with increasing polarizing power of the associated

cation, since polarization of the anion structure

leads to the distortion and consequent weakening

of the intra-ionic bonds.

This argument applies to the carbonates. There

is, however, some question as to how the polarizing

power of the cations is to be measured. For example,

Alekseenko [13] noted that the decomposition

temperatures (Pc02 =1 atm) of carbonates with 18

electrons in their cation outer shell were lower than

those of carbonates with cations of comparable

size containing 8 electrons. Ostroff and Sanderson

[98] suggested that an anion is most stable when it has the greatest control over its valence electrons

and is in a nonpolarizing environment. Since the

polarizing power of the cations is inversely proportional to their size and since the electronegativity measures the power of an atom or ion to attract electrons, Ostroff and Sanderson suggested that

anion stability should increase as some power of

(r/S), where S is the Sanderson stability ratio, a quantity linearly related to other electronegativity

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

1
scales. As a measure of stability they chose the

“lowest temperature at which decomposition could

be noted,” and found that for several sulfates this

temperature varied as (r/S) 1/2

. In order to define stability in more quantitative

terms, it would seem that a thermodynamic measure

would be less ambiguous. If one examines the

thermodynamic functions for the thermal decomposition of the carbonates, it is evident that the rea- son for differences in stability are due to AH° since

AS° values are nearly the same for all the carbonates

and arise primarily from the formation of a gaseous

product. Since AH° values at 298 K and those at higher temperatures increase in a parallel manner,

the 298 K values are an adequate measure of stability. It has been shown by a semi-empirical method

[121a] that A/Z^g is a linear function of r

1/2 /Z*,

where r is the cation radius and Z* is the effective

nuclear charge, as calculated from Slater's rules

Figure 1. Decomposition enthalpies as a function ofcation r

1/2/Z*.

[119a]. The plot is shown in figure 1. Although this

correlation is probably not unique, it does illustrate

the kind of connections to be expected between

atomic and thermodynamic properties.

3. Phase Transitions

a. Solid Transitions. The literature on solid

transitions is confusing. Several systems of naming

them exist and differences in the reported values of

the transition temperatures frequently are so great

that it is not clear whether the same or a different

transition is being described.

A number of well-defined transitions have been

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

and others have been collected and are listed.

b. Melting Points. Melting points have been

reported for a few of the more stable carbonates.

We have selected what appear to be reliable values.

It should be noted, however, that the decomposition

which occurs below the melting point makes a

measurement of the pure compound difficult and

that the solid phase may contain some oxide. Melting

points are usually obtained by rapid heating of the

salt to minimize decomposition, but it cannot always

be avoided. Another technique is to measure the

melting point under a high C02 pressure. The listed

values should therefore be regarded as lower limits.

c. Boiling Points. All carbonates decompose so extensively in the solid or liquid range that boiling

points cannot be measured.

4. Density

The total literature of high-temperature density

measurements on inorganic carbonates is very smalL

There exists only one set of reliable values for the

liquid salts, those of Janz and Lorenz [75] on the

alkali metal carbonates. We have reproduced their

original values and also their best fit to a linear

equation. On reason for the lack of data is that most

carbonates decompose considerably below the

melting point so that a measurement of any physical

property of the pure liquid is inherently impossible.

Except for a few limited studies below 100 °C, there

appear to be no studies of the temperature dependence over a wide range of temperature in the

solid state. 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 frequently

higher than experimental values. Since most

crystallographers do not indicate the temperature

for which the density is calculated, we have preferentially listed experimental values when these were available. An exception is the careful x-ray

work of Swanson and co-workers [123, 124] which

has been listed. Units for x-ray densities are g/cm-3

. In some cases experimental values seem to depend on the method of preparation, probably

because of differences in the defect structure. Thus,

although the density of a given sample can certainly

be determined to four or five significant figures, we

would caution against attributing too much significance to figures beyond the third. This is particularly true because the units of the reported

density values are frequently in doubt. In many

cases they are not stated at all; in others it is not clear whether the unity is g/cm3 of 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 difference between milliliters and

cubic centimeters is so small (the ratio is 1.000028)

as to be negligible. However, the difference in the

density of water between 4 and 20 °C is in the third

significant figure (1.00000 and 0.99823 g/ml, re- spectively) and no longer negligible for the better

measurements.

We have indicated the units as reported by the

original authors. When no units are given, it can

generally be assumed that g/ml is meant. In a few

9
instances we have listed the thermal expansion

coefficients, where these were listed in Vol. Ill of

[9]. The quantities given are:

, 10 6 dl 0 10«A/

a ——r~ at t ; ora= —tt

over the range At . / dt /At

energy functions” (fef) as discussed in reference

[6], pp. 166—9. Where the high-temperature thermal

data are available in the form of enthalpy and

entropy increments (H° — 8 and S° — Sms),

these are combined with 298.15 K entropies to get

the free energy functions:

A = 109 da

dt

^resp. 109 Aa \

At )’

i.e., 109 X the rate of

change of a with

fef

G

O TJO TJO TJC T n298 n T 112298 (St S298 ) 529!

t, at t° (resp. over At 0 ).

5. Decomposition Equilibria

As a measure of the stability of each carbonate,

we have calculated the equilibrium constants and

pressures of CO2 for each decomposition reaction.

Equilibrium pressures may either be obtained from

direct measurement in which the carbonate and

oxide are allowed to equilibrate with C02 at various

temperatures or they may be calculated by standard

thermodynamic procedures.

Direct measurements are frequently unreliable

since in many cases “pseudo-equilibria” appear, i.e., different constant pressures are obtained, depending

on whether the equilibrium is approached from

above or below [131]. If the same pressure is obtained by both approaches, this possibility is minimized and we have listed a few such values in cases where the necessary thermodynamic values were

not available. Whenever possible we have calculated the equilibrium properties by thermodynamic

methods.

For these calculations, four items of information

are needed:

(1) Knowledge of the course of the decomposition.

For the carbonates this is generally quite simple,

particularly if the carbonate and oxide coexist as pure solids. In that case the course of the reaction

is described by the overall stoichiometry, e.g., MC03 = MO + C02 , and the equilibrium pressure

of C02 is a function of temperature only.

Several carbonates, e.g., lead carbonate, do not

decompose directly to the oxide but rather to a basic

carbonate such as PbO • PbCOa. These cases are described in the individual sections, but equilibrium

pressures can usually not be calculated for the basic

carbonates because the necessary thermodynamic

data are largely lacking.

(2) High-temperature thermal data such as those

tabulated in reference [4] in the form of enthalpy

and entropy increments above 298.15 K, for each

compound.

(3) Heats of formation at 298.15 K.

(4) Entropies at 298.15 K.

These data are required for the decomposition

products as well as for the carbonates themselves.

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

Data” under each carbonate to show the data used

and to point out the gaps in existing data.

It has usually been most convenient to perform

the calculations through the medium of “free

where T is the temperature of interest in K. Then for a given T K, the fef’s for decomposition products and

sulfate are added and subtracted, in the same

manner as heats of formation in a thermochemical

equation, to yield a Afef for the decomposition

reaction. The heat of reaction at 298.15 K is obtained

from the heats of formation. Then the Gibbs energy

change attending the reaction is

AG°t = TAM+ A//2 98

from which the equilibrium constant for the reaction

,is obtained through the relation:

AG°T =-RT\n K.

For the case of simplest stoichiometry:

MC03 = MO + C02 K= PC02 .

In some cases the decomposing carbonate and the

oxide product may form a solid or liquid solution. In

that case the equilibrium constant

where aM0 and aMC0;i are the oxide and carbonate

activities in the condensed phase and the pressure

is now dependent on the solution composition. Such

information is not always provided. We would therefore caution against uncritically accepting measured

decomposition pressures as meaningful thermodynamic data unless the condensed phases have

been examined, for example, by x-ray diffraction

of the solids, or direct chemical. analysis.

Many of the older decomposition pressure

measurements were used by Kelley and Anderson

[3] in the calculation of thermodynamic properties

of reactions and compounds. When no newer data were available, we incorporated their results into

this monograph. The thermodynamic functions

in the tables were calculated in 1966 and the litera- ture survey is reasonably complete up to that date. In some cases newer 298 K values have been in- troduced when these represented major changes

from existing values or when this was necessary to be consistent with the ongoing revision of [2], The

values at higher temperatures were then revised

accordingly.

3
B. Kinetics of Thermal Decomposition

1. General

Although the literature on the thermal decomposition kinetics of carbonates is probably more extensive than that of any other class of compounds,

none of the data are of “reference data” quality, as

that term is commonly used. One reason for this

situation is that it does not yet seem to be possible

to prepare duplicate samples of any inorganic solid

salt which are identical in all the properties that

may determine the rate of decomposition, e.g., the

density of dislocations.

Nevertheless, the kinetics of decomposition are

of sufficient interest, both for theoretical and for

practical reasons, that it seems worthwhile to include the results of at least the more carefully

done studies. In this section we summarize those

features common to carbonate decompositions and

list factors affecting the rate of decomposition. The

discussion is essentially limited to solids, since most

experimental and theoretical work has been in this area. Kinetics of the individual compounds are treated in their respective sections.

2. Mechanism of Endothermic Solid State

Decomposition

Many authors have derived equations for the rate

of decomposition, A(s) —> B(s) +C(g), from theoretical models. The aim of all these studies is to find

functional relations between some measurable

quantity of the system, such as weight and time,

which can be tested against experimental results.

Although the usual caveat against taking agreement

between the predictions of the model and experimental results as evidence for correctness of the

mechanism of the model applies, nevertheless such

agreement, together with other experimental evidence, e.g., microscopic examination, is frequently

very helpful in at least narrowing the possibilities. In this section some of the models which have

been proposed are briefly summarized. Greater

detail can be found in books by Garner [54] and

Young [129] and in the original references.

Most models try to account for the shape of the

experimental a (= fraction decomposed) versus time t curves. These are generally S-shaped with a

relatively slow induction period, a steep portion

during which the bulk of the material decomposes,

and a slow final section. Experimentally determined

curves differ primarily in the relative duration of the

induction period. This induction period is commonly

identified with the formation of nuclei of the new

phase B imbedded in A; these nuclei grow, both by

spreading over the surface and into the bulk of the

decomposing particle. After the surface is covered

by B, this phase grows into the bulk of the material.

Decomposition thus occurs only at the reactantproduct phase boundary, as has been argued by

Langmuir from phase rule considerations [87]. It seems to be generally agreed that the evolution of

gas C during endothermic solid decomposition is

not a rate-limiting step in the process, i.e., the product B is sufficiently porous to permit the gas to escape readily. However, if the gas is generated in the body of the crystal, e.g., as a result of ionizing

radiation, and must diffuse out along grain boundaries, the rate of decomposition may depend markedly on the diffusion constant of the gas. This case has been discussed by Gafner [53].

Although the measured decomposition rate of carbonates is not rate-limited by diffusion

of the gas out of the crystal, it is clear that in the case of readily reversible decompositions,

such as those of the carbonates, the measured

rate will depend on the difference between the

equilibrium pressure pe and the actual pressure p

at the reaction interface. Thus, although the functional dependence of a on t may be independent

of pe —p, the actual rates are not. It is largely for

this reason that it is not possible to tabulate “standard reference data” rate constants, since the pres- sure p is not usually known. Instead one finds

decompositions carried out in streams of various

inert “sweep gases” which may or mav not remove

product gas from the reaction interface, or in vacuum. In the latter case, the question as to whether the manometrically measured pressure

equals the interface pressure needs to be answered.

Small sample size and a porous oxide favo.r this

condition.

(a) Kinetic Equations

The most complete classification of solid decompositions has been worked out by Jacobs and

Tompkins [73]. When the decomposition begins at some points on the A lattice where the local energy

is most favorable, small fragments of B are imbedded in A. These may at first retain the lattice

parameters of A and will therefore possess some

strain energy as a result of the deformation. Jacobs

and Tompkins show that the Gibbs energy accompanying the formation of a fragment of B containing

m molecules is AG = am213 — brn (1)

where a is proportional to the strain energy and b is the negative of the bulk Gibbs energy per molecule.

AG passes through a maximum at m — m* when a fragment has the critical size to be in equilibrium

with its surroundings. Smaller fragments are unstable and revert to A; larger ones are stable nuclei

and grow. Whether decomposition proceeds from

the growth of a few nuclei or whether many small

nuclei are formed depends on the relative magnitudes of the activation energies for nucleus formation (AGf) and the growth stage (AGjf). If AGf < AGf,

growth of existing nuclei predominates over formation of new ones. If AG|" — AGf , many small nuclei are formed, none of which grow to visible size. In

that case the induction period is much reduced.

The whole surface then nucleates rapidly and the

4
kinetics of the remainder of the reaction depends on

the rate at which the interface progresses into the

crystal.

We next consider various mechanisms for each

of the stages in the decomposition.

(1) Nucleus Formation. Jacob and Tompkins

consider only structure-sensitive nucleation, which

occurs at definite sites in the lattice where the

activation energy is least, such as lattice defects

and dislocations. The rate of nucleus formation

thus depends both on the defect density and on the

activation energy.

Jacobs and Tompkins distinguish two cases; (1)

the decomposition of a single molecule leads to the

formation of a nucleus. The probability of this

unimolecular decomposition is h = y exp (AGf/RT) (2)

where y is the frequency of lattice vibrations and

AGf is the activational Gibbs energy for nucleus

formation. The rate of nucleus formation at time t then is

dN/dt= kiN0 exp ( — kit)

or (3)

line of contact. Both of these effects result in rates

different from those predicted from models without

these refinements.

The simplest decomposition kinetics, commonly

observed for carbonates, result from a model in

which nucleation and subsequent surface growth

are extremely rapid. The rate of decomposition is then determined by the progression of the reaction

interface into the crystal. Rate laws have been

derived for various geometries, but the one for

spherical particles is most useful for carbonate

decomposition and has been given in several equivalent forms by various authors. In this derivation

it is assumed that the reaction interface proceeds

at constant velocity into the crystal, i.e., for a constant interfacial area, the rate of decomposition is a constant. For a spherical particle the reaction

proceeds along a constantly diminishing interface.

For such particles of initial radius, Ro, the fraction

decomposed at time t is

4/37r/?f) — 4/37r (/?(> — k-it)

z _ / k2 t\

3 “ =

4/377R* = 1

_3kot 3kit2 k$t3 r Ro Rg Rl

{ ]

and = A^o [

( 1 — exp ( — kit)]

where No is the total number of potential nucleus

forming sites. The possible loss of nucleus-forming

sites through their ingestion by growing nuclei is neglected. In the early stages of the reaction and

especially for large AGf and consequently small Ay

N — kiNot (4)

so that the number of nuclei increases linearly with

time.

(2) A power law results from two possible mechanisms: (a) a stable nucleus is formed in a bimolecular combination of two active intermediates, and

(b) several decompositions are required to form a stable nucleus. The first of these possibilities may

require the surface migration of several B molecules until they coalesce into a stable nucleus.

Some aspects of this phenomenon have been discussed by MacDonald [89],

(2) Growth of Nuclei. As each nucleus is formed,

it grows generally hemispherically, i.e., it spreads

across the surface of the crystal and into its interior.

Since the laws for the rate of formation of nuclei

are different from those for their growth, the overall

decomposition rate is likely to be complicated if the

two rates are comparable in magnitude. This is particularly true if the complications resulting from

overlapping nuclei are taken into account. As nuclei

grow, they ingest or overlap sites which would

otherwise yield nuclei. Also, as nuclei grow, they

impinge on one another and cease to grow along the

da 3k2 6kit 3k2 t

2

These equations were first derived by Hume and

Colvin [65].

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

( 7) where A is the total area of the reacting interface.

For a spherical particle A = k2 (l — af13

. where k'2 depends on Ro, and hence

( 8)

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

TrKmm < 9>

where k2 again depends on Ro. Equation (8) is usually

given in integrated form

k2 t= 1 — (1 —a) 1 /3

. (10)

The dependence of rate on particle size can he

included by the substitution k2 = k2/Ro.

5
The second limiting ease arises when the rate of linear propagation of the reacting interface

is so great that each particle decomposes as soon

as it has a nucleus. For this case, Hume and Colvin

give

Rate = koNt

i.e.,

dNjdt — koNt (11)

where Nt is the number of particles at time t. Integration from No to Nt gives

ln(NtlNo) — kot (12)

a first-order reaction. This behavior is most likely

to be approached by very small particles.

In general the kinetics of decomposition are more complex than in the two limiting cases just

outlined. Thus, if nucleation is not extremely

rapid, the number of nuclei will increase in time.

Each of these nuclei will then begin to grow at a

different time. Some of them will spread over

sites at which nucleation would have occurred at some future time. These sites are referred to as “phantom nuclei” by Jacobs and Tompkins. The

problem has also been studied in great detail by

Mampel [92]. Both Jacobs and Tompkins and

Mampel assume the rate of nuclei formation to be first order,

dN — k\Noe~kit (13)

where No is the total number of possible sites. For the induction period Mampel finds a t 4 law

for a, a relation derived by Jacob and Tompkins

for random nucleation and short times, and by

Fischbeck and Spingler [51] for the case of no

overlap, a constant rate of formation of nuclei, and

a large total number of nuclei.

For large radii (R>k2 t), Mampel finds

an eq (10) given previously for the contracting

sphere case. For small radii the fraction decomposed

is X=\ — ce~kt

where c is independent of t, but does depend on R in a rather complicated way. The rates are dependent on particle size and both small and large

radii give lower rates than intermediate radii.

A generalized equation has been derived by

Erofeyev [47] in terms of probability theory. In

his treatment

where p is the probability that a molecule will

react in the time interval dt.

The general solution of (14) for thermal decomposition is a=l — exp (— ktn

) (15)

where the constant n depends on the shape of the

nucleus and the number of electrons necessary

for the formation of a stable nucleus. For example,

cylindrical nuclei (the centers of formation are edges or surface cracks) give n = 3, flat nuclei

give n = 4. Thus the evaluation of n, most easily

carried out by transforming (15) into

In [ — In ( 1 — a)] = In k + n In t (16)

and plotting the left side of (16) versus In t , gives

some information about the nucleation mechanism.

(b) Surface Area Changes During Decomposition

Since the molar volume of the product oxides

is less than that of the corresponding carbonates,

decomposition results in a porous structure through

which the CO2 readily escapes. Although not

directly related to the decomposition kinetics,

the resulting increase in surface area during decomposition is of some interest in the production of

“active solids” and has been treated by several

authors. Accordingly, we briefly review it here.

The oxide is first formed at sites on the carbonate

lattice and is therefore in a strained state. Gregg

[61] describes these as small crystallites or micelles.

This strained pseudo-lattice will tend to recrystallize

into a stable oxide lattice with a resulting decrease

in surface area. Three mechanisms for this re- crystallization are generally recognized. At temperatures T/Tm < 0.2, where Tm is the melting point

of the oxide, adhesion predominates. This mechanism results from contact between parts of the

micelles which are atomically contiguous. For

0.2 < T/Tm < 0.35, surface diffusion predominates.

Here, mobile atoms or ions in the surface layer

leave their position on the lattice to migrate to sites of lower energy, i.e., to stable oxide lattice

positions. In DTA studies of carbonate decomposition [109]

small exothermic peaks (200-1000 cal/mol) were

observed at temperatures 40 to 80 °C above the

main endothermic decomposition peak, consistent

with such an oxide lattice rearrangement. The

formation of the surface mobile species from the

active species on the original lattice is considered

by Ilschner [71] to be slower than the prior formation

of the active species and the subsequent incorporation of the mobile species into the stable oxide

lattice. Above T/Tm = 0.5 sintering, i.e., lattice or bulk diffusion, is the most significant process

leading to decreased surface area. Thus as decomposition proceeds, there are basically two processes

leading to changes in surface area: the formation

of the metastable oxide produces an increase in

surface, and the formation of stable oxide decreases

6
the area. Consequently the “surface area” versus

“fraction decomposed” plot exhibits a maximum

before decomposition is complete. Nicholson [97]

has developed a mathematical model for this

piocess which agrees well with experimental

observations. If the purpose of the decomposition

is the preparation of an “active solid,” the reaction

should not be carried to completion in order to achieve the desired maximum surface area.

3. The Effect of Pressure on Rate

When the pressure of CO2 in a carbonate-oxide

system is equal to the equilibrium pressure, pe , no

net reaction occurs. When p < pe , the thermodynamic driving force favors oxide formation; conversely, when p > pe , carbonate formation is favored.

In actual systems the favored reaction may not

occur, however, because kinetic factors prevent it. Particularly when p is not too far from pe , the reaction may not proceed because some rate-limiting

process, such as nucleus formation, is proceeding

too slowly. The resulting spurious equilibria [131]

give rise to hysteresis effects, i.e., decomposition

stops for some p < pe , recombination stops for

p > pe . It is for this reason that we have largely

relied on thermodynamic methods for the calculation of equilibrium pressures. When the carbonate

and oxide are present as pure phases, the value of

pe is completely unambiguous, i.e., at a given temperature pe = Kp, the equilibrium constant for the

reaction MC03 =M0+C02. However, when the

condensed phase consists of a solid or liquid solution, pe no longer has a unique value since

mixture of C02 and some inert gas. Independence

of decomposition rate from flow rate is generally

taken as evidence that the reaction is not diffusioncontrolled, e.g., [69]. If the surface area is constant, the decomposition

rate is proportional to pe — p, i.e.,

v = k(pe — p).

Such a relation was also found to hold for both

cylindrical pellets of calcite and powdered CaC03 [72].

Cremer and Nitsch [39], in studying the decomposition of CaC03 , found that for samples which

followed a 2/3 rate law the pressure dependence of

the rate (in m1/3£ _1

) was given by

(17> A similar equation has been derived [69] from a two-step model in which the CaO is assumed to be

formed first in some “active” state CaO* which

occupies a constant fraction of the surface. The

reactions are then

kt CaC03 ^ CaO*+ C02

k<i 1-0 e p

k$

CaO* ^ CaO

/14

both of which are reversible. The rate K

, aMO

aMCO:, and

pe now depends on the solution composition. Since

the rate of decomposition is generally some function

of pe — p, we consider only systems without solution

formation. We have also omitted discussion of the

effect of “foreign” gases, i.e., gases other than

CO2 , on the rate. Such gases, e.g., 02 , H2 O, fre- quently function as catalysts or inhibitors, but

these effects are difficult to relate to well-defined

properties of the system [25, 66, 68]. In order to study the dependence of reaction rate on CO2 pressure, it is necessary that this pressure remains

fixed throughout the course of the decomposition

while more CO2 is evolving. Since the significant

value of the pressure is that at the reaction interface

it is possible that in a closed system where one depends on diffusion to remove the excess CO2 , the

pressure at the interface will exceed the nominal

value, particularly if the sample size is large and

the reaction is rapid. The effect of a pressure

gradient on the reaction rate has been investigated

[20] by the application of Fick’s law. Both for slow

and for fast reactions, the rate is proportional to pe ~ p<>, where p» is the pressure at the reaction

site but the proportionality constants differ in the

two cases. One way to avoid the problem of pressure

gradients is to use sweep gases made from a known

R = — dw/dt =

hobi k3 PA'oA'2 /1'4

ki + A3 + A'4 + k%P (18)

where the constant A0 is introduced to convert the

surface area 6 to surface concentration in molecules

per square centimeter, to account for the roughness

factor, and for the fact that only a portion of the

total surface sites may be potentially reactive.

Equation (18) can be written as

r =L^p.

VP + D

WhenP=0, R — Ro—l/D.

When P = Pe , R = 0 and C=l/Pe . Hence

P- 1 ~ PlPe BP+l/Ro

When 1/Ro > BP

1 1 R ~ w~K

(19)

( 20)

( 21 )

which is essentially the equation given by Cremer

and Nitsch [39],

7
4. Activation Energy

The activation energy E as defined by the usual

Arrhenius equation

k = Ae~E*lRT

is usually at least as great as the thermodynamic

enthalpy for the reaction, AH, though it is fre- quently greater. When E* = AH, the reaction is often described as occurring “without activation.”

The question as to whether E* = AH for the endothermic carbonate decompositions was first con- sidered by Fischbeck and Schnaidt [50]. In

comparing experimental values of E* and AH, they

found approximate agreement, but the scatter of

the then available data prevented rigorous comparison.

Shortly thereafter Zawadzki and Bretsznajder

[130] showed that for a reversible reaction 4 of the

type A(s)— B(S )-i- C(g), the experimental value of E

for the forward reaction, as determined from the

temperature dependence of k, must necessarily in- crease with rising pressure and will have its minimum value in the absence of the backward reaction,

i.e., in vacuum. This dependence was verified for

the decomposition of CaC03 . CdC03 , and Ag^CO;}.

In vacuum the temperature dependence of k was

the same as that of the equilibrium pressure pe , i.e., E

9L = AH. The dependence of E

on the pres- sure has recently been reexamined by Pavlyuchenko and Prodan [104], who distinguish two

cases: (a) if E* is determined from the temperature

dependence of the rate, with the pressure p held

constant at the various temperatures, then E* in- creases as p —» pe . (b) If, however, the ratio p\pe is held constant as the temperature is changed, then

E* will be independent of the actual numerical value

of the ratio and will have the same value as in vacuum. This conclusion was reached by a theoretical argument and verified by data on the decomposition of CdC03 . Since condition (a) is the

more usual experimentally, particularly in the older

literature, it is not surprising that widely varying

values of E* for the same reaction have been re- ported [121] by authors who failed to recognize

the dependence of E* on pressure.

The atomic mechanism responsible for monomolecular reactions, including thermal decompositions, was first discussed by Polanyi and Wigner

[105]. Their model assumes that decomposition

occurs when, due to energy fluctuations in the bonds

of the molecule, the bond strength is exceeded; or more precisely, that the bond energy “resides in

harmonic vibrations and that decomposition occurs when their amplitude is exceeded.” The resulting

4 For this type of reaction reversibility implies that the pressure of the gas C is the equilibrium pressure Pe and that infinitesimal changes in this pressure will produce the appropriate shifts in the direction of the reaction, i.e., to the left for p > Pe , and to the

right for p < Pe. Moreover, the solid phases must be the most stable forms, rather than metastable phases.

expression for the first-order Polanyi-Wigner rate constant is k—v exp (— ElRT) (22)

where v is the atomic frequency of vibration and E

is the activation energy. Thus a molecule having

energy E would leave the surface.

Shannon [118] has analyzed in some detail the

theory as it applies to the thermal decomposition

of solids. He finds that of the 31 reactions for which

he compared experimental rate constants with

those calculated from the Polyani-Wigner equation,

only a third show order-of-magnitude agreement.

In Shannon’s view, this lack of agreement stems

from neglect of rotational and other vibrational

degrees of freedom.

Shannon’s is probably the most ambitious attempt

to date to calculate a rate constant for the thermal

decomposition of a solid from absolute reaction rate theory. The resulting equation is of the same form

as eq (22), but v is replaced by a partition function

ratio,

kTQ*

k —

-fr-Q

exP

(— E/kT

)

(23)

where Q* is the complete partition function for the

activated complex excluding that for the reaction

coordinate and Q is the complete partition function

for the reactant. Detailed calculations of the partition functions were carried out for the decomposition of CaC03 and MgC03 , using known values of

the rotational and vibrational partition functions

and several models for the activated complex. The

best agreement with experiment was obtained from

a model in which the carbonate ion in the lattice is free to rotate and the C02 molecule leaves directly

from the surface of the solid, rather than from a mobile surface layer.

5. The Effect of Method of Preparation on

Rate

Various authors have 'recognized that the rate of

thermal decomposition may depend on the method

of preparation and the thermal “history” of the

sample, but systematic studies in this area are in- herently difficult and few have been carried out. For example, even if the effect of surface structure on the rate of nucleation is recognized, the characterization of the surface in quantitative terms

and the reproducible preparation and systematic

variation of surface and bulk properties is as yet a largely unsolved problem. Additional problems

arise if impurities present in the samples at very

low concentrations, such as H2 0, catalyze the decomposition [108]. In that case, the resulting

change in mechanism will affect not only the rate but also the activation energy.

If the carbonate is a single crystal, the reaction

will propagate from the surface into the interior. In polycrystalline materials the reaction may ad8
ditionally propagate along grain boundaries [63].

A systematic study of the effect of sample constitution on decomposition rate was carried out

by Cremer and Nitsch [40J. They decomposed four

calcite crystal samples at 850 °C and CO2 pressures

ranging from 120 to 220 torr. The CaO product

exhibited increasing grain size with increasing pressure. The four CaO samples were then reacted

with CO2 at 850 °C and 420 torr to produce four

CaCOs samples with similarly increasing grain size.

When these four samples were then decomposed

under identical conditions (850 °C, 200 torr C02 ), the rate of decomposition increased in inverse

order of crystallite size. The disappearance of the

induction period with decreasing crystal size was

particularly noticeable. These results are explained

in terms of more rapid nucleus formation at lower

pressures, leading to smaller CaO crystals and

consequently a larger number of active sites in the

CaC03 subsequently formed.

It is evident from these results that the course

of the decomposition is affected by the microstructure of the sample and that the reproducible preparation and characterization of nominally identical

materials is an important prerequisite for the understanding and systematization of decomposition

kinetics.

6. The Effect of Heat Transfer on Rate

An endothermic solid-state decomposition proceeds as the produce-reactant interface advances

into the interior of the sample. The interface can,

however, advance only if the necessary heat of re- action is supplied to it. Since the source of heat is outside the sample, the rate of heat transfer to the

interface may become rate determining if the in- herent rate of reaction is greater than the rate of

heat transport. The rate at which heat is transported

depends not only on the properties of the product

through which heat must be transported, but also on the general experimental arrangement. It is for

this reason that one expects and finds the literature

to be conflicting.

For example, Narsimhan [96] has derived an equation for the rate of thermal decomposition based

on the idea that the entire particle rapidly reaches

the decomposition temperature so that there are no

internal temperature gradients. All of the heat

reaching the reaction interface is used up in the

decomposition so that the reaction rate depends on

the rate of heat transport. For both spherical and

cylindrical particles, good agreement with some

literature values is found. Under these conditions

the chemical steps are no longer rate-determining.

It can thus be seen that in any solid-state reaction

whose enthalphy is appreciable, the establishment,

maintenance, and measurement of the temperature

at the reaction interface is a complicated problem.

Thus, careful studies of the decomposition of

PbCOa [76] and several other inorganic salts [133]

have shown that the induction period may disappear

if the entire sample is rapidly brought to the reaction

temperature. There is therefore now some question

whether the commonly reported induction periods

in the decomposition of so many substances are

artifacts of the experimental arrangement. Considerable care in the experimental design is certainly

required to ensure a constant and known temperature at the reaction interface throughout the course

of the decomposition.

7. The Influence of Crystal Defect Structure

on the Rate of Decomposition

The idea that the defect structure of a solid

reactant affects its rate of decomposition seems

to be generally accepted; but of all the factors

influencing the kinetics, this one is the most diffi- cult to characterize quantitatively. The literature on the subject has been reviewed by Boldyrev

[26], who then classified decompositions according,

to which aspects of the defect structure needed to be considered in particular cases. Because of its general interest, his classification is reproduced

here.

Effect of crystal defects on the rate of thermal decomposition

of solids (26)

Decomposition proceeds through

Crystal defects

affecting the rate

Breaking of bonds within

cationic and anionic

lattice constituents Electron

transfer Reversible

decomposition

Irreversible decomposition

from anion

to cation

Change of habit Growth figures, vicinals, macro-fissures | Effect

1 Effect Dislocations, and

groups of them Effect Impurity inclusions

in lattice No

1 effect Ionic defects Electronic defects No

effect

According to this scheme, carbonate decomposition constitutes the simplest type, since it is a re- versible reaction in which only intra-ionic bonds

are broken. Therefore, only the first two types of

defects affect the rate of decomposition.

The main distinction between reversible and ir- reversible decompositions is the more complex

mechanism of the latter, whereas reversible decomposition proceeds through only one step, such

as the breaking of a C—O bond in the CO" ion.

Since these reactions start on the surface, a change

in the crystal habit which affects the surface area

ratio of more/less reactive faces will necessarily

affect the rate, particularly since such a change of

crystal habit also changes the ratio of surface/bulk

ions. Growth figures and macrofissures have their

greatest effect on the initial rate since nucleation

commonly begins at the most reactive surface sites

and these are usually defects in the structure.

9
If Boldyrev’s arguments are correct, it would

seem that the study of reversible decomposition

offers the best chance for a quantitative study of the

relation between the various effects discussed above

and the rate of decomposition. This is particularly

true for the main course of the reactions which is much less affected by the defect structure than is the initial nucleation.

Boldyrev and Medvinskii [27] have attempted

a quantum-mechanical interpretation of the above

decomposition scheme by considering changes in

the electron energy levels which correspond to the

various types of decomposition, primarily in the

language of band theory. The theory is not yet suffi- ciently well developed to permit quantitative predictions, although it provides a rational framework

for the classification scheme. In terms of this

theory, intra-ionic decompositions involve very

local changes in the electron levels, i.e,, “excitons

of small radius.” The authors predict that these

reactions will in the future be describable by energy

level diagrams.

C. Suggestions for Experimental

Work

In view of the preceding discussion, it is clear

that considerable care in experimental design and

execution is required if meaningful thermodynamic

and kinetic data on the decomposition of inorganic

salts are to be obtained. For the measurement of

equilibrium pressures, the requirements are rather

simple: (a) the measured temperature must be the

temperature of the sample; this requires placing

the measuring device close to the sample, (b) The

measured pressure must be the equilibrium pres- sure; precautions against pseudo-equilibria require

that the same pressure is obtained when approached

from above and below, (c) The condensed phases

must be well characterized in order for the measured

pressure to be thermodynamically meaningful. In

the case of a solid decomposition, both the reactant

and product should be examined for evidence of

solid solution formation. If a liquid phase forms,

pressure measurements are probably of very limited

usefulness since not only the composition of this

phase, but the component activities must be known

in order to calculate the equilibrium constant of

the reaction. The characterization of the solid

phases is also important for another reason. In cases where the reactant can exist in more than one crystalline modification, it is possible that a form other than the most stable one becomes

“frozen” in a metastable state at the temperature

of the decomposition. The decomposition pressure

and the reaction Gibbs energy calculated from it will be different for the two forms and hence the

reaction to which AG° corresponds will not be properly identified. Hence AG° will have its correct

equilibrium value only if the solid phases are the

most stable ones, (d) Just as the vapor pressure of

small drops is greater than that of large ones, so

the decomposition pressure of small crystals is

larger than that of large ones. Experimental evidence for this phenomenon exists [36A], For equilibrium measurements, relatively large ( > 10-2 cm)

crystals should probably be used, or the pressure

studied for several particle sizes large enough for

the pressure to be independent of particle size.

Since the measurement of meaningful decomposition pressures is difficult, work on the determination of high-temperature heat capacities would

be appropriate. A perusal of the tables in this monograph shows that only about half of the carbonate

decomposition pressures can be calculated from

existing thermodynamic data. Most of the remainder

(Cu, Fe, Pb, Zn) could be calculated if these Cp values were available. Co and Ni also lack reliable

S?298 values.

The proper execution of kinetic studies is more

difficult since more factors determine the rate of

decomposition than determine the equilibrium.

Since these factors have already been discussed

in the previous section, we only emphasize here

their consequences for the proper experimental approach. We assume that in a chemical study the

main interest lies in elucidating the chemical kinetic

steps, i.e., the reaction mechanism. Therefore it is desirable to ensure that purely physical steps, such

as diffusion of the gas and heat transfer to the re- action site, not be rate-limiting. In planning a ki- netic study, the following factors should then be

considered:

(a) The significant temperature is that at the re- action site, i.e., at the phase boundary between the

reacted and unreacted material. Placement of the

temperature-measuring device is thus of critical

importance. Particularly in very endothermic or exothermic reactions, temperature gradients within

the furnace and even within the sample may be

large.

(b). In order to minimize temperature gradients

within the sample, to ensure that the rate is not controlled by the diffusion of gaseous products and

by the rate of heat transport between sample and

surroundings, the sample should be small and consist of a thin layer of finely divided material unless

the study is primarily concerned with nucleus formation on the surface, in which case a large sample of

known geometry and perhaps surface structure,

e.g., a single crystal, may be preferable. If the same

rate is attained with samples of different thickness,

temperature gradients are probably absent.

(c) Depending on the mechanism of the reaction,

the particle size of the sample may or may not affect

the rate. Although the theoretical models discussed

in section B2 predict such dependence, experimental evidence for both dependence and non- dependence exist. It is therefore advisable to use as narrow a size distribution as possible in each experiment and to measure the rate of several such

distributions with the total sample size held fixed.

(d) Particularly for readily reversible reactions,

the observed decomposition rate depends on the

difference between the equilibrium pressure, p(

-

,

10
and the actual pressure, p. Therefore, only limiting

rates (p = 0) are meaningful for the forward reaction.

Measurement should therefore be carried out at a

sufficiently low pressure for the mean free path

in the gas to be of the order of the pore size. If an inert sweep gas is used to remove gaseous

products, it must be carefully controlled so that the

product is really removed from the reaction

interface.

(e) Since heat transfer to .the sample is more

difficult in a vacuum, special care must be taken to provide good thermal contact between sample and

heat source. The sample should be kept from

decomposing until the desired temperature is reached by maintaining the pressure above the

equilibrium value and then rapidly decreasing it to zero at the beginning of the reaction.

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

(1) The crystal habit. Samples may be crystalline or amorphous. In some cases the crystal structure

may depend on the method of preparation, e.g., high-temperature modifications may be frozen on by quenching the molten salt.

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

(3) The nature and concentration of substances

adsorbed on the crystal surface. These may function

as catalysts or inhibitors for the decomposition.

(4) The defect structure. If decomposition starts

at dislocations on the lattice, the rate of decompositions may be anomalously high if the defect density

is high.

It is primarily the difficulty of reproducibly preparing well-characterized samples which accounts for

the present lack of meaningful rate constants. The

measurement of rate constants for solid decompositions which are as significant and reproducible as those for gas reactions will require great care, both

in the preparation of materials and in the execution

of experiments.

CO2, CO, and 02 The CO2 produced by the decomposition of a carbonate dissociates partly to CO and O2 . However, at temperatures for which most of the car- bonate decomposition equilibrium pressures have

been calculated, the decomposition of C02 is very

slight. We have therefore given the decomposition

pressures in terms of CO2 only. However, for those

wishing to refine the calculations still further, we

have calculated the free energy functions and equilibrium constant for the reaction CC^g)^ CO(g>+ 1/2

02(g)-

Entropy and enthalpy values at 298.15 K are taken

from [2]. Free energy functions above 298.15 K are based on [10] for C02 and O2 and on [8] for CO.

Thermodynamic data (298.15 K)

S°

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

co2 213.64 51.06 2 CO 197.56 47.219 2 02 205.03 49.003 2 AHP

kj mol-1 kcal mol -1 References

C02 393.51 94.051 2 CO 110.52 26.416 2 02 0 0 2

Dissociation of C02 C02(B, = CO(K ,+ 1/2 02( B ) Af/Jw ,s = 282.98 kj, 67.635 kcal

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

AHf°— standard enthalpy of formation

S° = standard entropy

G° = standard Gibbs (free) energy

fef = free energy function

r^o I/O T

TGA = thermogravimetric analysis

DTA = differential thermal analysis.

A. Free Energies T C02 CO o2 Reaction

fef fef fef Afef HG°t K J deg ' J deg~' J deg~' J deg ' kj

298.15 -213.63 -197.56 - 205.04 -86.45 257.21 400 -215.15 -198.71 - 206.20 -86.66 248.32 500 -218.15 -200.87 -208.41 -86.95 239.51 600 -221.62 -203.32 -210.93 -87.18 230.60 700 -225.14 -205.80 -213.50 -87.32 221.86 800 -228.81 -208.21 -216.02 -87.40 213.07 900 -232.36 -210.53 -218.44 -87.42 204.31 1000 -235.75 -212.76 -220.77 -87.40 195.59 1100 - 239.03 -214.87 -222.99 -87.35 186.90 1200 -242.18 -216.89 -225.10 -87.28 178.24 1300 -245.19 -218.82 -227.12 -87.19 169.64 1400 -248.10 -220.67 -229.05 -87.09 161.06 1500 - 250.90 -222.43 - 230.89 -86.98 152.52 1600 - 253.58 -224.12 - 232.66 -86.86 144.00 1700 -256.20 -225.75 -234.36 -86.75 135.52 1800 -258.70 -227.30 -235.98 -86.62 127.07 1900 -261.11 -228.81 -237.55 -86.49 118.65 2000 -263.43 -230.25 -239.06 -86.36 110.26 2100 - 265.68 -231.64 -240.51 -86.24 101.89 2200 -267.86 - 232.99 -241.91 -86.11 93.55 2300 - 269.96 -234.29 - 243.22 -85.98 85.23 2400 -272.01 -235.55 - 244.58 -85.86 76.93

11

357-056 0-69—2
B. Equilibrium constants Barium

T(K) log K K

298.15 - 45.062 8.67 x 10~ 46 400 - 32.427 3.74 x IO- 33 500 - 25.021 8.53 X 10- 26 600 - 20.082 8.28 X 10“21 700 - 16.555 2.79 X 10~ 17 800 - 13.912 1.22 X 10- 14 900 - 11.858 1.39 X 10- 12 1000 - 10.216 6.08 X 10 _n

1100 - 8.875 1.33 X 10- 9 1200 - 7.759 1.74 x 10~ 8 1300 - 6.816 1.53 x 10- 7 1400 - 6.009 9.79 X 10- 7 1500 - 5.311 4.89 X 10 -6 1600 - 4.701 1.99 X 10- 5 1700 - 4.164 6.85 X 10- 5 1800 - 3.687 2.06 X 10-4 1900 - 3.262 5 . 47 X 10- 4 2000 - 2.880 1.32 X 10- 3 2100 - 2.534 2.92 X 10-3 2200 - 2.221 6.01 X 10- 3 2300 - 1.936 1.16 X 10 -2 2400 - 1.674 2.12 X 10- 2

The effect of the BaCC>3 transitions at 1079 and

1241 K on the calculated equilibrium pressure is virtually nil. Calculations above 1600 K were made

by extrapolating the fef values of the individual

compounds.

The solid-solid transitions of BaCOs exhibit

pronounced hysteresis, i.e., different temperatures

are obtained, depending on whether the transition

is approached from above or below [86]. The melting

point can only be measured under high CO2 pressure. AH and AS for the transitions are taken

from (1).

Density oj BaCC>3

Phase T( K) d References

c 293 4.287 15

c (witherite) 299 4.308 (x-ray) 123a

c (cubic) 1348 3.889 (x-ray) 123c

Transitions o/BaC03

Phase change T( K) AH AS References

W mol -1 J mol -1 deg-1

c,orthorhombic (y)—» c,hexagonal (/3

)

1076]

14.9 13.7 1. 86 c,orthorhombic (y) <— c, hexagonal (/3 ) 1023J c,hexagonal (/3)—»c,cubic (a) 1249'

3. 4. 86

c, hexagonal (/3) <— c,cubic (a) 1228

.

I c(a)—»

1653 1

Thermodynamic date (298.15 K)

S°

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

BaC03(c,II, witherite) 112.1 26.8 1, 10 BaO 70.3 16.8 1, 5, 6

AHf°

kj mol -1 kcal mol -1 References

BaC03 -244.7 -297.5 12 BaO -582.0 -139.1 91

12
Decomposition o/BaC03

Bat-O.Tty)— BaO(c)+ C02(u)

A//2V,5 = 269.2 IcJ. 64.35 kcal

T BaO BaC03 Reaction

K fef fef Afef ac;° log K K

Jdeg~ l Jdeg- 1 Jdeg~ l kj

298.15 -70.3 -112.1 -171.8 218.0 -38.19s 0.36 x 10-39 400 -72.1 -115.7 -171.5 200.6 -26.20„ 6.32 X 10-27 500 -75.8 -122.9 -171.0 183.7 — 19.194 6.39 X 10-20 600 -79.9 -131.1 -170.4 167.0 - 14.53h 2.89 X 10“ 15 700 -84.2 -139.7 -169.7 150.5 -11.22a 5.91 X 10- 12 800 -88.4 -148.3 - 168.9 134.1 — 8.7567

i 75 X 10-9 900 -92.4 -156.7 - 168.0 118.0 — 6.8493 1.42 X 10- 7 1000 -96.2 - 164.8 -167.1 102.1 — 5.334 4.63 X 10-6 1079* -99.1 -171.10(a) - 166.39 89.70 — 4.342 4.55 X 10-5 -171.08(0) -166.37 89.72 — 4.343 4.54 X 10“5 1100 -99.9 -173.1 - 165.8 86.8! -4.12s 7.54 X 10~5 1200 - 103.4 - 182.2s - 163.3 73.2k — 3.190 6.46 X 10-4 1241* - 104.7 - 185.8-,(/3, y) -162.3 67.8s -2.85s 1.39 X 10-3 1300 -106.7 -191.1 - 160.8 60.2n -2.41a 3.81 X 10-3 1400 - 109.9 -199.6 - 158.4 47. 4o -1.77* 1.69 X 10-2 1500 -112.9 -207.7 -156.0 35.1t -1.22s 5.96 X 10-2 1600 -115.8 -215.5 -153.9 23.

0

3 — 0.752 1.77 X 10- 1 (1700) -118.6 -222.9 -151.8 11-1.3 — 0.342 4.55 x 10-' (1800) -121.3 -230.0 - 150.0 -0.791 + 0.23o 1.05

(1900) -123.9 -236.6 -148.4 - 12.6s + 0.348 2.23

Beryllium

Normal anhydrous BeC03 probably does not

exist at ambient temperatures. Dehydration of the

hydrate at elevated temperatures probably forms

basic carbonates [83]. The value given for AHf° of

BeC03 is calculated from a measured value of AH

for the reaction BeO + C02 = BeC03 [83] and

the current value of BeO to be published in a continuation of [2].

Thermodynamic data (298.15 K)

S°

BeC03

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

BeO 4.15 3.38 2 aHr

kj mol 1 kcal mol 1 References

BeC03 c5 To -245 2 BeO - 145.0 2

Cadmium

The rate of decomposition has been measured by

Centnerszwer and Bruzs [34] in a C02 atmosphere

and by Prodan and Pavlyuchenko [107] as a function of C02 pressure. The decomposition yields

CdO without the formation of intermediates. In vacuo the activation energy is 36 kcal. The rate of

decomposition depends on the method of preparation. The kinetics follow the contracting sphere

model [50, 65]. Experimental dissociation pressure

measurements have been reviewed by Kelley and

Anderson [3], who selected the measurements of

Andrussow [17] for their analysis. For the reaction

CdC03 = CdO + C02 they obtain

AG°(cal) = 23,200 -10.46 T log T+8.71

X 10-3 r

2 - 13.80 T.

The density must be regarded as very uncertain

since there are no 20th-century measurements and

DeSchulten reports a value of 4.960 [42].

Density ofCdC03

Phase r(K) Density References

c 293 4.25k g ml-1

9, Vol. I

Thermodynamic data (298.15 K)

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

CdC03 92.5 22.1 2 CdO 54.8 13.1 2 aHr

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

CdC03 -750.6 -179.4 2 CdO -258.1 -61.7 2

13
Calcium Thermal expansion coefficient: More kinetic studies of the thermal decomposition

of CaC03 have been reported than for any other

carbonate, and perhaps for any other salt. Many

of the general features of carbonate decomposition

described in the Introduction have been elucidated

from CaCC>3 decomposition and will not be repeated

here.

The reaction CaC03 = CaO + CO2 is reversible.

Decomposition occurs without formation of intermediate products. For compacted powders the

contracting sphere model applies [72]. A comparison of experiments done with various particle sizes

and with single crystals indicates that even the

reaction order may vary, depending on which mechanism is rate-determining at each phase of the re- action. For example, in a careful analysis of

composition versus distance in a rectangular block

of powder, Kappel and Huttig [82] found different rate laws applying at varying distances from the

surface. Such apparent changes in mechanism can

easily occur if different processes, such as heat

transport to the reaction site or escape of CO2 , become rate-determining at various stages of the

decomposition. In decomposing a cube of CaC03 with thermocouples imbedded in it, it was found

that the center of the cube was colder by from 15

to 85° than the nominal furnace temperature [18].

The authors suggest from this and other evidence

that the decomposition rate is proportional to the

heat absorbed. Such a mechanism would most

likely not be applicable to thin layers of fine-grained

powders.

Many values of the activation energy have been

reported, nearly all near 167 kj (40 kcal)/mol,

a value approximately equal to the standard

enthalpy of decomposition.

The aragonite—* calcite transition has been

studied as a function of temperature and pres- sure [37, 90]. At ambient pressure, the transition occurs near 730 K [62, 122]. In addition, a transition from vaterite to calcite occurs in the range

620 to 670 K, with small impurity concentrations

producing large changes in the transition temperature [122].

Density of CaC03

Phase T( K) d References

c, calcite 273 2.771 gml"1 9, Vol. Ill c, calcite 298 2.713 15

c, calcite 299 2.711 (x-ray) 123a

c, aragonite 297 2.930 15

a (2° to 81°) =25.1353 + 0.0118*, || opt. axis;

CaC03 : Calcite . .

and- 5.5782 + 0.00138*,

_L opt. axis a

40 = 26.21 + 0.0160U - 40),

|| opt. axis; and = — 5.40

• 0.0087(f — 40), _L opt. axis a(50° to 60°) = 1.447

9, Vol. Ill

CaC03: Aragonite.

O af = 34.60 + 0.0337* 1 of =17.19 + 0.0368* a3° = 10.16 +0.0064* )

Transitions of CaC03

Phase Change T( K) References

c (aragonite.—» c (calcite) 728 62

c (vaterite)—* c (calcite) 623-673 122

Thermodynanuc data (298.15 K)

S°

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

CaC03 (aragonite) 88.7 21.2 1, 5 CaC03 (calcite) 92.9 22.2 5. 6 CaO 39.7 9.5 5. 6

AHr

kj mol 1 Kcal mol 1 References

CaC03 (aragonite) -1207.04 -288.49 1 CaC03 (calcite) - 1206.87 -288.45 1 CaO -635.09 — 151.7s 6

14
Decomposition of CaCOCaCO;j((.)= CaO,, ; • (

from calcite 15 = 178.28 kj, 42.61 kcal from aragonite AH°m , 5 = 178.45 kj. 42.65 kcal.

A. Calcite

T CaO CaC03 Reaction

fef fef Afef AC°T log A A

./ cleg~ l J deg-' J deg kj 298.15 -39.7 -92.9 -160.5 130.4 -22.850 1.41 x 10 - 23 400 -41.5 -96.4 - 160.3 114.2 -14.909 1.23 X 10 - 15 500 -44.9 - 103.3 -159.7 98.43 -10.283 5.21 X10-"

600 -48.7 -111.4 -158.9 82.

2 -7.219 6.04 X 10-8 700 -52.7 -119.8 -158.1 67.6! -5.045 9.01 x 10- e 800 -56.6 -128.2 -157.2 52.

4 -3.430 3.71 x 10- 4 900 -60.3 -136.4 -156.3 37.63 -2.184 6.55 X10- 3 1000 -63.8 5 -144.2 -155.4 22.90 -1.196 6.36 X 10 “ 2 1100 -67.2 -151.8 -154.5 8.360 -0.397 4.01 X 10 - 1 1200 -70.5 -159.0 - 153.6 - 6 .O63 + 0.264 1.84

B. Aragonite

T CaO CaC03 Reaction

fef fef Afef AG? log K A

298.15 -39.7 - 88.7 -164.7 129.3 -22.661 2.19 x 10- 23 400 -41.6 -92.0 -164.6 112.6 -14.707 1.97 x 10 _lS 500 -44.9 -98.7 -164.3 96.30 -10.060 8.71 x 10-"

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

Cesium

Very little information on CS2 CO3 is available.

The decomposition pressure has been measured

[ 88]. Decomposition becomes noticeable near 880 K (2.6 X 10~3 atm at 883 K) and reaches 0.207

atm at 1453 K.

Transitions of CS2CO3

Phase change T( K) References c—> 1 1065 111

Thermodynamic data (298.15 K)

S°

J mol 1 deg 1 cal mol 1 deg 1 References

Cs2 C03 Cs2 0

aHr

k Jmol 1 kcal mol 1 References

Cs2 C03 — 1118.

-267.4 1, 10 Cs2 0 -318. -75.9 10

Cobalt

Information on the thermal decomposition of

C0 CO3 is scarce. A number of basic oxides have

been reported [85], but their existence cannot be

regarded as proved. In a more recent study [30], no intermediates were reported, but the decomposition pressure exhibited an unexplained break

in the log p versus 1 \T plot.

in view of several density values in the 4.1—4.2

range, the value of 2.818 for spherocobaltite given

in [9, Vol. I] is probably in error.

Density of CoCO.t

Phase d nk) References

c 4.07 116

c 4.24 (x-ray) 48

c (spherocobaltite) 4.214 (x-ray) 299 123

15
Thermodynamic data (298.15 K) Thermal expansion coefficient:

S°

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

CoC03 88.5 21.2 80

CoO 52.93 12.65 10 C03O4 102.5 24.5 10

FeC03 : Siderite..

*

40 = 19.2 with A(20°-70°) = 25.5 | and a

40 = 6.05 with

A(20°-70°) =917 1 9, Vol. Ill

Thermodynamic data (298.15 K)

S°

AHf° J mol -1 deg-1 cal mol -1 deg-1 References

kj mol-1 kcal mol 1 References FeC03 92.9 22.2 1 , 10 CoC03 -722.6 -172.7 1 , 10 1 eO (1 e. 9470) 57.49 13.74 5 CoO — 238.H -57.1 10 1 e3 04 146.4 35.0 5 Co304 -905.0 -216.3 10 1 e2 03 87.4 20.9 5

Copper

No high-temperature thermodynamic data are

available and decomposition pressures of pure

CuC03 have not been measured. The temperature

at which decomposition of the minerals, azurite

and malachite, becomes noticeable varies, not sur- prisingly, with the rate of heating [119].

Thermodynamic data (298.15 K)

S°

J mol 1 deg 1 cal mol 1 deg 1 References

CuC03 87.9 21.0 10 CuO 42.63 10.19 5, 6 , 10 Cu20 92.38 ±0.42 22.08 ± 0.10 91a

aHr

kj mol 1 kcal mol-1 References

CuC03 -595.0 -142.2 10 CuO -155.77 ±0.63 — 37.23 ±0.15 91a

Cu2 0 -170.83 ±1.26 -40.83 ±0.30 91a

Iron

AHf°

kj mol-1 kcal mol-1 References

FeC03 -740.6 -177.0 270-3 unpubl.

FeO (Feo.9470 ) -266.5 -63.7 1 Fe304 -1117. -267.0 1 , 10

1 e2 03 -822.15 -196.5 1 , 10

Lead

In a careful study of PbC03 decomposition,

Centnerszwer, Falk, and Awerbuch [36] determined the following successive equilibria at 1 atm

C02 :

274° 286°

PbC03

<—> 3 PbO • 5 PbC03 «—

360° 412°

PbO • PbC03 «—> 2 PbO • PbC03

<—> PbO.

In a kinetic study of PbC03 decomposition, these

authors (31) found an induction period which was

accelerated by small amounts of H2 0, and a first- order decomposition for the main course of the

reaction independent of moisture content.

Glasner and Hodara [57] have reexamined the

basic carbonates and find 2 PbO • 3 PbC03 to be

the carbonate of lowest oxide content. They list the following pressures:

Anhydrous FeC03 can be prepared by heating

the 1/2 hydrate in C02 at high pressure (40 to 50

atm) near 300 °C. The thermal decomposition

always leads to a higher-valent oxide, since the

C02 liberated immediately oxidizes the FeO to Fe3 0.i [22]. The reaction may also be written as

3 FeC03 = Fe3 04 + 2 C02 + CO.

The oxidation occurs so fast that even in vacuum

no FeO is formed.

Density o/FeC03

Phase T{ K) d References

c, siderite 293.9 .851 16

°C

p (mm)

PbC03 2 PbO • PbC03 5 PbO • 6 PbC03

222 271

72 690

240 246 320

75 81 690

290 380

101 690

PbO • PbC03 3 PbO • 2 PbC03

°C

p (mm)

350 390

123 690

370 430

150 690

Kelley and Anderson [3] have used the enthalpy

data of Marshall and Bruzs [93] to calculate the

thermodynamic functions for the reactions

(1) 2 PbC03 = PbO • PbC03 + CO,

A//(cal) = 20,230 + 4.23 T- 8.21 X 10-3 r

2 AG°(cal) = 20,230 — 9.74 log I’d- 8.21

X 10-3 r

2 - 13.60 T

16
(2) PbOPbC03 = 2 Pb0 + C02 A//(cal) = 20,910 + 4.23 T-8.21 X 10-3 P

AG°(cal) = 20,910 — 9.74 T log T

+ 8.21 X 10-3 r

2 - 12.967\

i In studying the decomposition of PbC03 in vac- uum as a function of temperature and particle size,

Kadlets and Dubinin [81] found that, in contrast to the work of Centnerszwer and Awerbuch,

the induction period was absent if the sample was

brought to the reaction temperature very rapidly.

The rate (— dm/dt) was proportional to the sur- face area. The expected conformity with eq (10) was observed. Rate constants, expressed in g

s _1 cm-2

, were independent of particle size. Hence

the decomposition of PbC03 conforms to the contracting sphere model with rapid surface coverage

by nuclei. The kinetic results are summarized

by the relation

k(g cm_2s _1

) = (0.5 ± 0.5) exp (—41,500 kcal//?7)

in the range 236 to 335 °C.

These results were confirmed by a similar study

in nearly the same temperature range [115]

for which £ = 40.2 kcal was found. The reaction

rate with freshly prepared and with ground samples differed probably because of changes in the

surface structure.

Density o/PbC03

Phase nk) d References

c 275.3 6.533 16

c 299 6.582 (x-ray) 123a

Thermodynamic data (298.15 K)

S°

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

PbC03 131.0 31.3 2 PbO (yellow) 68.70 16.42 2 PbO (red) 66.5 15.9 2 PbO-PbCOs 204. 48.8 2 aHr

kj mol 1 kcal mol 1 References

PbCOa -699.1 -167.1 2 PbO (yellow) -217.3 -51.94 2 PbO (red) -219.0 -52.34 2 PbO-PbCOs -816.7 -195.2 2

Lithium

The melting point of Li2 C03 is still uncertain.

Two recent measurements [75, 111] are considerably lower than the value of 1008 K given in (1),

but differ from each other by 6°. In both of these

studies, the salt was protected from decomposition

by being kept under a pressure of C02 high enough

to prevent weight changes. AHw values given by

[77] and [112] differ by 7 kj.

Equilibrium pressures of C02 above liquid Li2 C03 were measured by Janz and Lorenz [75], using a dynamic method in which the temperature was

varied until weight changes (losses on heating, gains on cooling) were noted. The resulting pressures are

larger by several orders of magnitude than K values

determined from thermodynamic data. This apparent discrepancy can be accounted for by considering that K= (dhnolau2 CO:t )Pco2 and therefore

(auzolai ioco,) =KIPco.2

. For example, at 1000 K we

calculate K= 2 X 10-6

, whereas Pco 2 — 1.9 X 10-2

, i.e., aLi2o/«u2co:i ~ 10-4

. This is not unreasonable

since Li2 0 almost certainly dissolves in molten

Li2 C03 . It is for this reason that we have relied,

whenever possible, on thermodynamically calculated equilibrium constants.

Density o/Li2 C03

Phase T (K) d References

c 290.7 2.111 9, Vol. I

1 1012.2 1.8246 75

1027.4 1.8189

1029.7 1.8190

1052.1 1.8119

1069.7 1.8036

1082.7 1.7987

1105.1 1.7910

1120.1 1.7843

The liquid data are fitted by

d = 2.2026 — 0.3729 X 10_3!T (1010-1120 K).

Transitions of Li2 C03

Phase change T(K) AH kj mol 1 References

c —* 1 993 111 999 42. 75 1008 1

Thermodynamic data (298.15 K)

S°

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

Li2 C03 90.17 21.55 10b

Li2 0 37.89 9.056 10a AHf°

kj mol 1 kcal mol 1 References

LfiCOn -1216.04 -290.64 10b

Li2 0 -598.7 -143.1 10a

17
Decomposition of Li2 CO:

Li2 COs(c ) — Li2 Ol<

. ) + C02(e, A//?slil5 = 223.8okJ. 53.49 kcal

T Li2 0 Li2 C03 Reaction

fef fef Afef AG°t log K K

J deg~ x J deg-' J deg-' kj

298.15 - 37.89 - 90.17 - 161.35 175.69 - 30.780 1.66 x IO-31 400 - 40.15 - 94.16 - 161.14 159.34 - 20.808 1.56 x 10-21 500 - 44.71 — 102 . 1 a - 160.65 143.47 - 14.988 1.03 x 10- 15 600 - 50.08 — 111 . 8s - 159.82 127.90 - 11.135 7.33 x 10- 12 700 - 55.67 - 122 . 50 - 158.39 112.93 - 8.427 3.75 x 10-9 800 - 61.25 — 133 . 57 - 156.50 98.60 - 6.438 3.65 x 10-7 900 - 66.71 - 144.38 - 154.67 84.60 - 4.910 1.23 x 10-5 1000 - 72.00 - 192 . 9* - 114.76 109.04 - 5.695 2.02 x 10-* 1100 - 77.12 - 203 . 7 ! - 112.42 100.13 - 4.755 1.76 x 10-5 1200 - 82.06 -214.0b - 110.16 91.60 - 3.988 1.03 x 10-4 1300 - 86.82 - 224. 0 ! - 108.00 83.40 - 3.351 4.46 X 10-4 1400 - 91.41 -233.5b - 105.96 75.46 - 2.815 1.53 x 10-3 1500 - 95.85 - 242 . 7 , - 104.03 67.75 - 2.359 4.37 x 10-3 1600 - 100.13 — 251 . 5(i - 102.22 60.25 - 1.967 1.08 x 10-2 1700 - 104.3b - 259 . 94 - 100.50 52.94 - 1.627 2.36 x 10-2 1800 - 108.26 - 268.04 - 98.90 45.78 - 1.329 4.69 x 10-2 1900 - 137.00 — 275.83 - 122.25 - 8.473 + 0.233 1.71 2000 - 141.09 - 283 . 34 - 121.16 - 18.53 + 0 . 484 3.05

Magnesium

In vacuum, MgCOs decomposes to MgO at temperatures as low as 320 °C, the solids exhibiting

x-ray lines corresponding to MgC03 and MgO

only [60]. In earlier work, Centnerszwer and

Bruzs [32, 33] had reported the following series of

equilibria at 1 atm C02 (in °C).

Density of MgC03

Phase UK) d References

c 294.6 2.980 15

c 298 3.037 9. Vol. I c (magnesite) 298 3.009 (x-ray) 123b

Thermal expansion coefficient:

373° 442°

MgCO.3 <—» MgO-MgCO.i <—> 3

MgCOsrMagnesite a

40 = 21.3 with A(20°-70°)‘ = 33.9 1

1 ; and a

40 = 5.99 with A(20°- 70°)= 24.3 1 9, Vol. Ill

469°

MgO-MgCO.i «—» 4 MgO + C02

These equilibria are based entirely on pres- sure measurements and weight loss data. Until

structural evidence for the existence of these compounds is produced, their existence must be re- garded as doubtful.

The decomposition follows the “rapid nucleation and contracting sphere” kinetics with an

activation energy of 150 kj (36 kcal), somewhat

higher than the enthalpy of decomposition [29,

38, 95],

Thermodynamic data (298.15 K)

S°

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

MgC03 65.86 15.74 10b

MgO (periclase) 26.

4 6.44 10b aHr

kj mol-1 kcal mol-1 References

MgC03 - 1111.7 -265.7 10b

MgO (periclase) -601.2 -143.7 10b

18
Decomposition of MgC03 MgC03(c)=Mg0(C i+ C02( B ) ^298.15= 116.9 kj, 27.95 kcal

T MgO MgC03 Reaction

M fef Afef AG°t log K K

J deg"' J deg-' J deg ' kj

298.15 — 26. 94 -65.86 -174.72 64.85 -13.361 4.36 X 10-14 400 — 28.4g -69.03 -174.60 47.10 -6.151 7.07 X 10- 7 500 -31.5S -75.43 -174.23 29.82 -3.115 7.67 X 10-4 600 — 35.0q -83.46 -173.23 13.00 -1.132 7.39 x 10-2 700 — 38.75 -91.08 -172.90 -4.088 + 0.305 2.02 800 -42.

7 -99.19 -172.00 -20.66 + 1.349 22.35 900 -45.87 -107.19 -171.02 -36.98 + 2.146 140. 1000 — 49.2,4 -115.01 -169.97 -53.03 + 2.770 539. Manganese

The thermal decomposition of MnC03 is complicated by the easy interconvertibility of the

various manganese oxides. When MnC03 is decomposed in 02 -eontaining atmospheres, the

oxide formed depends on the temperature. In

addition, a number of nonstoichiometric oxides

are also formed [24], the particular oxide formed

depending primarily on the temperature and

partial pressure of oxygen. In the absence of

oxygen, the decomposition product is MnO, at least up to 300 or 400 °C [43]. Above this temperature range, MnO may be oxidized by the CO2

formed according to [94]

MnO+ 1/3 C02 = 1/3 Mn3 04 + 1/3 CO.

The decomposition can thus also be thought of

3 MnCO.3 = Mn3 04 + 2 C02 + CO.

The actual temperature at which such trans- formations occur therefore depends on the efficiency

of C02 removal and on the partial pressure of 02 . The thermodynamics of many reactions involving

manganese compounds have been given by Mah

[90a]. The predominent mode of decomposition

of MnC03 is MnC03 = MnO + C02 and we have

therefore given equilibrium constants only for

this reaction. Our results agree very closely with

Mah’s. Thermodynamic functions for other oxides

are listed in [11 and 90a]. The kinetics follows

the contracting sphere model.

Density of M11CO3

Phase T( K) d References

c 293 3.125 gml"1 9, Vol. I c, rhodochrosite 294.4 3.633 16

Thermodynamic functions (298.15 K)

S°

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

MnC03 85.8 20.5 5, 10 MnO 59.71 14.27 5. 10 Mn3 04 (a) 154 36.8 99 Mn2 03 HO.., 26.4 5 Mn02 53.05 12.68 5

AHf

KJ mol" 1 kcal mol 1 References

MnC03 -94.96 -213.9 5 MnO -385.1 -92.05 6, 10 Mn3 04 (a) -1386.8 -331.4 6, 10 Mn2 03 -956.9 -228.7 99. 100 Mn02 -520.9 -124.5 1, 10

Decomposition o/MnC03 MnC03 ( c )= MnO(C ) -C C02(g> A//°298.i5= 116.3 2 KJ, 27.80 kcal

T MnO MnC03 Reaction

fef fef Afef A G° log K K

J deg~' J deg-' J deg-' kj

298.15 -59.71 — 85.

7 — 187.5s 60.39 -10.580 2.63 X10"14 400 -61.53 — 89.

0 — 187.

7 41.32 -5.396 4.02 XI 0"« 500 -64.98 -95.9s -187.1a 22.75 -2.377 4.20 X 10"3 600 -68.95 — 103.

7 -186.6, 4.34a -0.378 4.19 X 10" 1 700 -72.97 — 1 12.

0 — 185.0 — 13.81 5 + 1.031 10.74

19
Nickel

There is no evidence for the existence of an- hydrous NiC03 , although hydrates are well known.

These lose H2 O on heating up to 470 K, but some

CO2 is probably also lost before all the H2 O is re- moved [41, 52], In addition, the NiO product seems

to lose oxygen rather easily since metallic Ni has

been found in NiO formed by the decomposition

of NiC03 below 700 K [114]. S?9H.i 5 of NiC03 has

been calculated from low temperature Cp measure- ments to be 85.4 J mol-1 deg-1

[80].

Potassium

The free energy functions of K3 0 in [10a] are based on estimated Cp and S£98 values. In the 1966

revision [10b] fef values of K3 C03 are listed to 2500 K, but decomposition pressures cannot be

calculated above 1100 K because fef data for K2 O

are lacking above this temperature. K2 0 is suffi- ciently unstable for its dissociation to K<g) and O2 to be significant. Equilibrium constants for this

reaction are given in [10a] and can be combined

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

Density of K2 C03

Phase T (K) d

c 293 2.330 gml-' 9, Vol. Ill c 288 2.29 gml-1 9, Vol. I

1 1180.8 1.8922 gem -3 75

1185.3 1.8904

1193.9 1.8864

1195.9 1.8848 1206.6 1.8824

1213.5 1.8778

1220.0 1.8749

1223.6 1.8735

1234.1 1.8684

1236.3 1.8670

1245.3 1.8640

1257.4 1.8584 1270.0 1.8527

1283.3 1.8467

The liquid data are fitted by

d= 2.4141 -0.4421 x 10~3 T (1180-1280 K).

Transitions of K2 C03

Phase

change

nK) AH AS References

c, IV -»c. III c, IIWc, II C, II —» c, I c, I-» 1

523

701 895

1171 ±2

KJ

mol -1 J mol -1 deg -1

11

1 27 75, 77, 111, 112

Thermodynamic data (298.15 K)

S°

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

k2 co3 155.5 37.17 10b

k2 o 94.1 22.5 10b

AHP

KJ mol-1 kcal mol 1 References

K2 CO3 -1150.2 -274.9 10b

k2 o -363. 2 -86.8 10b

20
Decomposition of K2 CO;

K2C03(c)— KijOi, ! - CO2( g> A//',298.,5 = 393.5 KJ, 94.05 kcal

T K2 0 k2 co3 Reaction

fef fef Afef AG°t log K K

Jdeg ~' Jdeg-' Jdeg-' kj

298.15 - 94.1 - 155.5 - 152.3 348.1 - 60.987 1.03 x 10-61 400 - 97.5 - 160.2 - 152.5 332.5 - 43.420 3.80 X 10-44 500 - 104.2 - 169.4 - 152.9 317.1 - 33.123 7.53 X 10-34 600 - 111.8 - 180.3 - 153.1 301.6 - 26.260 5.49 X 10-27 700 - 119.5 - 191.6 - 153.1 286.3 - 21.365 4.31 X 10-22 800 - 127.1 - 203.0 - 153.0 271.1 - 17.703 1 . 98 X 10- 18 900 - 134.5 - 214.2 - 152.7 256.1 - 14.864 1 . 37 X 10- 15 1000 - 141.7 - 225.2 - 152.2 241.3 - 12.603 2.50 X 10- 13 1100 - 148.5 - 235.9 - 151.7 226.7 - 10.763 1.73 X 10 -”

Rare Earths

The existence of the anhydrous carbonates at ambient temperatures has not been proved, al- though hydrates are well known [126]. Such hydrates decompose on heating, losing water at first,

and then forming a series of basic carbonates, but

it is not certain whether the initial loss of water is also accompanied by some decomposition of the

carbonate ion [14]. However, anhydrous carbonates

of La, Ce, Pr, Nd, Sa, Eu, and Gd, have been re- ported by Head and Holley [64], who carefully decomposed the hydrates by TGA and removed the

theoretical amount of H2 O.

The formation of carbonates also results from

the thermal decomposition of oxalates, but these

seem to be formed as basic carbonates or in mixtures with oxides [58, 70, 101]. No thermodynamic

data are available.

Similarly, the existence of Pu(C03)2 as an intermediate in the decomposition of Pu(C2 04)2 has been

reported [78].

Rubidium

Very little information is available. The melting

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

Transitions of Rb2 C03

Phase change T (K) Reference

c —> 1 1146 111

Thermodynamic data (298.15 K)

S°

J mol 1 deg 1 cal mol”1 deg”1 References

Rb2 C03 174.9 41.8 2 (unpubl.)

Rb2 0 (124.) (29.7) 1

aHr

kj mol 1 kcal mol 1 References

Rb2 C03 -1128.0 -269.6 1 , 10 Rb2 0 -330. -78.9 1, 10

Silver

The thermal decomposition of Ag2 C03 is complicated by the fact that the decomposition product,

Ag2 0, decomposes further (Ag2 0-*2 Ag+1/2 02 , for thermodynamic data see [100a]). Therefore

weight loss experiments can only be carried out in

the relatively narrow temperature region, < 250 °C,

where Ag2 C03 decomposes but Ag2 0 does not. Alternatively, at higher temperatures the reaction can be followed by absorbing CO2 and O2 in separate

traps and measuring their respective weight changes

[81].

The kinetics are most consistent with a model of

rapid nucleation followed by contracting sphere

kinetics [120]. However, the actual rates are ex- tremely dependent on the method of preparation.

Kadlets and Dubinin [81] studied this effect by

preparing Ag2 C03 by two methods — one by precipitation from concentrated solutions which precipitated rapidly, the other by precipitation from

21
very dilute solution so that the salt first formed as Transitions of Ag^CO;

a supersaturated solution. Although both preparations exhibit the same x-ray diffraction pattern, the

second preparation decomposes much more rapidly

than the first, an effect which the authors attribute to a higher defect density. Direct evidence is lacking,

however.

Free energy functions of Ag2C03 and Ag^O are calculated from data in [4],

Phase change

c —* c No information

c —* 1 Decomposes

Thermodynamic data (298.15 K)

S°

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

Ag2 C03 167.4 40.0 2 Ag2 0 121.3 29 5

Density of Ag2C03 AHf°

Phase T (K) d References kj mol 1 kcal mol 1 References

c 293 6.077 gml-1 9, Vol. I Ag2C03 -505.8 -120.9 2

c 298 6.131 (x-ray) 124 Ag2 0 -31.0s -7.42 2

Decomposition of Ag2 C03

Ag2C03(C)— Ag20<c )+ COfew A//^8 15 = 81-291 kj, 19.429 kcal

T Ag2 0 Ag2 C03 Reaction

K fef fef Afef AC°T log K K

J deg- 1 ./ deg-1 J deg-1 kj

298.15 -121.3 -167.4 -167.6 31.32 -5.487 3.26 X 10-6 350 -122.2 -168.7 -167.5 22.66 -3.381 4.16 X 10-4 400 -124.0 -171.9 -167.3 14.38 -1.878 1.33 X 10-2 450 -126.5 -176.0 -167.0 61.4s — 0.7130 1.85X10-'

500 -129.2 -180.7 -166.6 — 2.087 + 0.211! 1.63

(550) - 132.0 -185.6 -166.2 -10.09 + 0.9587 9.09

Sodium

The literature on the phase transitions is some- what contradictory. The transitions at 629 and 759 K

but not the one at 891° were recently observed by

Reisman, as was the melting point. Two other transitions, at 593 and 723 K were observed by Ginzburg

[56] and Popov [106], and very similar values also

by Jaffrey and Martin [74].

AHm values of [77] and [112] differ by 5 kj.

We have listed the average value. According to Khlapova [84] , who studied the polymorphism of

Na2C03 by DTA, the actual transition temperature

will vary somewhat, depending on sample preparation and treatment, in some cases by 10 to 15°.

If this is so, the solid-solid transitions would re- quire careful reinvestigation.

Decomposition pressures were measured by

Janz and Lorenz [75], but the same comments apply to them as were made for Li2 C03 . The diffi- culties encountered in measuring equilibrium pres- sures above the liquid salt by the effusion method

have been described in great detail by Motzfeld

[94a]. He showed that above the melting point

C02 , Na(g ), and O2 vaporize from the melt, but that

the activation energy for evaporation of C02 is

quite high. Consequently both the observed pres- sure and the steady-state composition of the melt

depend on the size of the effusion orifice.

In this paper we have not calculated the partial

pressures of the above three gaseous species, but

they can easily be calculated from the equilibrium

constants for the dissociation of the oxide given in

[10a], and the table in this work.

Density of Na2 C03

Phase T d References

c 293 2.533 g ml- 1 9, Vol. I

1 1137.9 1.9685 gem-3 75

1146.2 1.9666 1165.5 1.9576 1178.8 1.9508

1184.7 1.9477 1223.3 1.9288 1233.2 1.9264

1244.8 1.9211 1260.3 1.9142 1277.0 1.9080

The liquid data are fitted by

d = 2.4797 — 0.4487 X lQ-3 r( 1140 —1280 K)

22
Transitions of Na2 CO.

Phase change T(K) AH References

c, IV -^c, III 629

K J mol~ 1

1 c, III-*c, II 759 1 C, II —> c, I 891 1 C, 1-^ 1 1129 ±2 31 75, 77, 112

Thermodynamic data (298.15 K)

S°

Na2 C03 Na2 0

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

138.7 75.27 33.17

17.99 10b

10b aHr

kj mol-1 kcal mol-1 References

Na2 C03 — 1130.

7 -270.26 10b

Na2 0 415.9 -99.4 10b

Decomposition of Na2 C03 Na2 C03( c) = Na2 0(c ) + C02( g ) AW?98. 15 = 321.4 kj, 76.8, kcal

T Na2 0 Na2 C03 Reaction

fef fef Afef AG°t log K K

J deg- 1 J deg- 1 J deg-1 kj

298.15 -75.27 — 138.

g -150.1, 276.62 -48.462 3.45 X 10-49 400 -78.19 — 143.

2 -150.0, 261.36 -34.131 7.40 X 10-35 500 -83.82 — 152.

3 — 149.

2 246.56 -25.758 1.75 X 10-28 600 -90.31 — 163.

0 -148.8, 232.08 -20.205 6.25 X 10-21 700 -96.98 — 174.

0 — 147.0 218.12 -16.276 5.29 X 10-17 800 -103.58 — 186.5s — 145.

2 204.71 -13.366 4.30 X 10- 14 900 -110.03 - 198.00 — 144.

7 191.44 -11.111 7.75 X 10- 12 1000 -116.26 — 208.

8 — 143.3 178.34 -9.316 4.83 X 10-1° 1100 -122.30 — 219.59 — 141.

2 165.47 -7.858 1.39 X io- 8 1200 - 147.25 — 249.09 - 140.32 152.98 -6.659 2.19 X 10- 7 1300 -154.13 -259.4s -139.85 139.57 -5.608 2.47 X 10-6 1400 - 160.65 — 269. 4.J — 139.

2 126.31 -4.713 1.94 X 10-5 1500 - 166.85 — 278.96 -138.7s 113.19 -3.942 1.14 X 10-4 1600 -172.74 — 288.09 — 138.

9 99.77 -3.271 5.36 X 10-4 1700 -178.35 -296.84 -137.69 87.

0 -2.682 2.08 X 10-3 1800 -183.70 — 305.24 — 137.

4 74.51 -2.162 6.88 X 10-3 1900 -188.82 — 313.30 -136.6, 61.81 -1.699 2.00 X 10-2 2000 -193.73 -321.

6 -136.08 49.21 -1.285 5.19 X 10-2 2100 - 198.44 — 328.54 — 135.

7 36.68 -0.912 1.22 X 10- 1 2200 -202.97 -335.74 — 135.

6 24.23 -0.575 2.66 X io-> 2300 -207.33 -342.6s -134.5s 11.84 -0.269 5.38 X 10- 1 2400 -211.52 -349.4! -134.09 -0.454 + 0.009s 1.02

Strontium

In contrast to the numerous studies of CaC03

and MgC03, the decomposition of SrC03 has re- ceived very little attention. Wanmaker and Radielovic [128] studied the decomposition in air thermogravimetrically and report that it commenced at

850 °C and was finished at 1175 °C. The rate for

fine and coarse powders was nearly the same. Up to 950 °C the decomposition rate was constant,

at higher temperatures a change from zero-order

to first-order kinetics occurred as the reaction

progressed beyond a = 0.5.

Lander [86] observed considerable hysteresis

in measuring the orthorhombic-hexagonal transi23
tion, but even the higher of his two temperatures Thermodynamic data (298.15 K)

is 13° lower than the value given below.

Density of SrCOs

Phase T(K) d References

c,strontianite 298 3.736 15

c, precipitated salt 3.585 21

Transitions of SrC03

Phase change T(K) AHf References

c.orthorhombic—» c, hexagonal 1198

kj mol~' 17 1, 19, 86

c, hexagonal— c, cubic 1770 19

c —> 1 1770 1

S°

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

SrC03 97.1 23.2 1, 5 SrO 54.4 13.0 10

AHP

kj mol-1 kcal mol-1 References

SrC03 -1220.1 -291.6 12 SrO -592.0 -141.5 91

Decomposition of SrCOs

SrC03( c)= SrO(c )+ CCk(B) A//£98 ,5 =234.57 kj, 56.05 kcal

T SrO SrCOs Reaction

fef fef Afef AG°t log K K

J deg~ x J deg~' J deg-' kj

298.15 -54.4 -97.1 -171.0 183.7 -32.18 6.5 x lCF55 400 -56.3 -100.6 -170.8 166.4 -21.72 1.9 x 10~22 500 -59.8 -107.6 -170.4 149.5 -15.62 2.4 x 10- 16 600 -63.8 -115.7 -169.8 132.8 -11.56 2.7 x lO -12 700 -68.0 -124.0 -169.2 116.2 -8.67 2.1 x 10-9 800 -72.0 -132.2 -168.7 99.7s -6.51 3.1 x 10-7 900 -76.0 -HO., -168.1 83.

7 -4.84 1.4 x 10~5 1000 -79.7 -147.9 -167.5 67.1 4 -3.51 3.1 x 10-4 1100 -83.3 -155.4 -167.0 51.0, -2.42 3.8 x 10“3 1200 -86.7 -162.6 - 166.2 35.

0 -1.53 2.9 X 10-2 1300 -89.9 -170.8 -164.3 21.1* -0.85 1.4 x 10-' 1400 -93.0 -179.6 -162.5 7.24 -0.27 5.4 x 10-' 1500 -95.9 -186., -160.7 6.36 + 0.22 1.7

Thallium

There is disagreement about the course of the

decomposition. According to TGA studies [44]

,

TI2CO3 is stable well into the liquid range, up to 645 K, and then decomposes to a basic salt. Both

Gattow [55] and Rossa [110] find decomposition

directly to T12 0, without intermediates.

Density o/T12 C03

Phase T(K) d References

c 293 7.11 9, Vol. I c,y monoclinic 7.24 (X-ray) 125

Transitions o/T12 C03

Phase

Change

T (K) AH AS References

C, II^c, I 591 kj mol~' J mol~ ' dep~'

127

C, 1-^ I 546

536

18 36 1 125

Thermodynamic data (298.15 K)

S°

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

ti2 co3 -155.2 37.1 2 T1,0 -126. 30. 2 AHf°

KJ moP1 kcal mol-1 References

T12 COs - 700.0 -167.3 2 ti2 o -178.7 -42.7 2

Zinc

The decomposition of ZnCCL has been studied

by several authors. Huttig, Meller, and Lehman

[67] used natural ZnCCL, smithsonite. Bretsznaider

and Cibor [28] used synthetic material. In neither

24
case was the material pure Z Transitions o/ZnC03 11 CO3 . There is also

some indication that the decomposition is not entirely reversible [113, 132]. It is generally agreed

that over most of the concentration range, contracting sphere kinetics are followed, but the details

of the induction period are not certain. Various

values have been reported for the activation energy

and it is not yet certain whether the value at zero pressure is equal to the enthalpy of the reaction.

Density of ZnC03

Phase T (K) d References

c 293 4.44 9, Vol. I c, smithsonite 296 4.348 16

c (smithsonite) 298 4.333 (x-ray) 123c

Phase change T (K)

Thermodynamic data (298.15 K)

S°

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

ZnCO.3 82.4 19.7 2 ZnO 43.64 10.43 2 aHr

kj mol-1 kcal mol 1 References

ZnC03 -812.78 -194.26 2 ZnO — 348.2s -83.24 2

25
References

[1] Rossini, F. D., Wagman, D. D., Evans, W. H., Levine, S., and Jaffe, I., Selected values of chemical thermodynamic properties, NBS Circ. 500 (U.S. Government

Printing Office, Washington, D.C., 20402, 1952).

[2] Wagman, D. D., Evans, W. H., Halow, I., Parker, V. B., Bailey, S. M., and Schumm, R. H., Selected values of chemical thermodynamic properties, NBS Tech. Note 270—3 (1968) (An advance issue of a revision of the

tables of [1] containing the first 34 elements in the standard order of arrangement. Some selected, but not

yet published, values are also included.).

[3] Kelley, K. K., and Anderson, C. T., Contributions to the data on theoretical metallurgy. IV. Metal Carbonates —

Correlations and Applications of Thermodynamic

Properties. U.S. Bur. Mines Bull. 384, 73 pp. (1935). Reissued as part of Bull. 601 (1962). [4] Kelley, K. K., Contributions to the data on theoretical metallurgy. XIII. High temperature heat-content, heat

capacity, and entropy data for the elements and in- organic compounds, U.S. Bur. Mines Bull. 584, 232 pp.

(1960). [5] Kelley, K. K., and King, E. G., Contributions to data on

theoretical metallurgy. XIV. Entropies of the elements

and inorganic compounds, U.S. Bur. Mines Bull. 592,

149 pp. (1961). [6] Lewis, G. N., Randall, M., Pitzer, K. S., and Brewer, L., Thermodynamics, 2d ed. (McGraw-Hill Book Co., New

York, N.Y., 1961). [7] Wagman, D., Evans, W., Halow, I., and Dudley, R., Section 4, pp. 170—216 in American Institute of Physics

Handbook, 1963.

[8] McBride, B. J., Heimel, S., Ehlers, J. G., and Gordon, S., Thermodynamic properties to 6000 °K for 210 sub- stances involving the first 18 elements, NASA SP-3001,

(Washington, D.C., 1963). [9] International Critical Tables (McGraw Hill Book Co., New York, N.Y., used for some densities in Vol. I, III., 1928). [10] Kubaschewski, O., and Evans, E. LL., Metallurgical

Thermochemistry, 3d ed. (Plenum Press, New York and

London, 1958). [10a] JANAF Thermochemical Data compiled and calculated by the Dow Chemical Co., Thermal Laboratory, Midland, Mich. (1965).

(b) First Addendum 1966.

(c) Second Addendum 1967. [11] Coughlin, J. P., Contributions to the Data of Theoretical

Metallurgy. XII. Heats and Free Energies of Formation

of Inorganic Oxides. U.S. Bur. Mines Bull. 5 84 (1954). [12] Adami, L. H., and Conway, K. C., U.S. Bur. Mines, Rept.

Invest. No. 6822, 7 pp. (1966) CA 65, 9830 (1966). [13] Alekseenko, L. A., Sbornik Nauch Rabot. Akad. Nauk.

Beloruss. SSR, Inst. Khim. 36 (1956). [14] Ambrozhii, M. N., Luchnikora, E. F., and Sidorova, M. I. Zhur. Neorg. Khim. 5, 366-71 (1960), Engl. Transl.

p. 176. [15] Anderson, C. T., J. Am. Chem. Soc. 56, 340-2 (1934). [16] Anderson, C. T., J. Am. Chem. Soc. 56, 894 (1934). [17] Andrussov, L., Z. Phys. Chem. 115, 273-88 (1925). [18] Auffredic, J. P., and Vallet, P., C. R. Acad. Sei., Paris, Ser. C 263, No. 10, 652-5 (1966). [19] Baker, E. H., J. Chem. Soc. 2525 (1962). [20] Barret, P. and Bonnefair, L., Bull. Soc. Chim., France 576

(1961). [21J Batuecas, T., Rancano, A., Ibarz, J., An. Soc. Esp. Fis. Quim. 27, 290 (1929). [22] Berg, L. G., and Buzdov, K. A., Zhur. Neorg. Khim. 6,

2003-8(1961).

[23] Berg, L. G., and Buzdov, K. A., Zh. Neorgan. Khim. 7,

1773-8 (1962).

[24] Bergstein, A., and Vintera, J., Coll. Czech. Chem. Communs. 22, 884-95 (1957). [25] Bischoff, F., Zanorg. Allgem. Chem. 2 5 0, 10 (1942);

ibid. 262, 18 (1950). [26] Boldyrer, V. V., Kinetika i Katalio 1, 203 (1960). [27] Boldyrev, V. V.. and Medvinskii, A. A., Kinetika i Kataliz 4, 481 (1963). [28] Bretsznajder, S., and Cibor. E., Bull. Acad. Polon. Sci. Chim. 13, 709-12, 713-16, 765-9 (1965). [29] Britton, H. T. S., Gregg, S. J., and Winsor, G. W.. Trans. Faraday Soc. 48, 63—69 (1952). [30] Buzdov, K. A., Zhur. Neorg. Khim. 11, 1251 (1966). [31] Centnerszwer, M., and Awerbuch, A.. Z. Phys. Chem.

123, 127-33 (1926). [32] Centnerszwer, M., and Bruzs, B., Z. Phys. Chem. 114,

237-50(1924).

[33] Centnerszwer, M., and Bruzs, B., Z. Phys. Chem. 115,

365-76 (1925). [34] Centnerszwer. M., and Bruzs. B., Z. Phys. Chem. 119,

405-18(1926).

[35] Centnerszwer, M., and Bruzs. B., Z. Phys. Chem. 123,

111-26 (1926). [36] Centnerszwer, M., Falk. G., and Awerbuch. A.. Z. Phys.

Chem. 115, 29-53 (1925). [36a] Centnerszwer, M., and Krustinsons, J.. Z. Phys. Chem.

130, 187 (1927), 132, 185 (1928). [37] Clark, S. P., Jr., Am. Mineralogist 42, 564-6 (1957). [38] Cremer, E., Allgauer, K., and Aschenbrenner. W.. Radex

Rundschau 195 3 , 494—502.

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

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

(1962). [41] Deren, J., Haber, J., and Sloczynski, J., Bull. Acad.

Polon. Sci., Ser. Sci. Chim. 9, 245-52 (1961). [42] DeSchulten. Bull. Soc. Chim. (3) 19, 34 (1898); Bull. Soc. Min. 20, 195 (1897). [43] Dollimore, D.. and Tonge, K. H., Proc. 5th Symp. Reactivity of Solids, 1964, pp. 497—508 (Elsevier Publ. Co., Amsterdam, The Netherlands, 1965). [44] Duval, C., Wadier, C., and Servigne, C., Anal. Chim.

Acta 20 , 263-7 (1959). [45] Elliot, N., J. Am. Chem. Soc. 59, 1380 (1937). [46] Erdey, L., Liptay, G., and Gal. S., Talanta 12, 883-92

(1965). [47] Erofeer. B. V., Dokl. Akad. Nauk. SSSR 52 , 511 (1946). [48] Ferrari, A., and Colla, C., Atti Accad. Lincei 10, 594

(1929). [49] Fischbeck, K.. and Schnaidt, K.. Z. Elektrochem. 38,

199-205 (1932). [50] Fischbeck, K., and Schnaidt, K.. Z. Elektrochem. 38,

649 (1932). [51] Fischbeck, K., and Spingler, H.. Z. f. Anorg. u. allgem.

Chem. 241, 209 (1939). [52] Francois. J., Compt. Rend. 230, 2183-4 (1950). [53] Gafner, G.. Trans. Faraday Soc. 55, 981 (1959). [54] Garner, W. E., Ed., Chemistry of the Solid State (London:

Butterworths, 1955). [55] Gattow, G., Pure Appl. Chem. 2, 121-3 (1961). [56] Ginzburg, D. M„ J. Gen. Chem. USSR, 26, 1103-5 (1956). [57] Glasner, A., and Hodara, I.. Bull. Res. Council (Israel) 7A, 66 (1958). [58] Glasner, A., Levy, E., Steinberg. M.. and Badenheimer.

W„ Talanta 11, 405-14 (1964). [59] Goldschmidt, V., Ber. 60, 1263 (1927). [60] Gordeeva, V. A., Zhabrova, G. M., and Kushnerer, N. Y., Dokl. Akad. Nauk. SSSR. 129, 835-8 (1959). [61] Gregg, S. J., J. Chem. Soc. 3940 (1953). [62] Gruver. R. M., J. Am. Ceram. Soc. 33, 171-4 (1950). [63] Hashimoto, H., Nippon Kagaku Zasshi. 82, 1605 (1961). 26
[64] Head, E. L., and Holly, C. E., in Proc. 3d Conf. Rare

Earth Research (Gordon and Breach, New York, publ.

1964), pp. 51-63.

[65] Hume, J., and Colvin, J., Phil. Mag. (7), 8, 589 (1929). [66] Huttig, G. F., and Herrmann, E., Kolloid Z. 92, 9 (1940). [67] Huttig, G. F., Meller, A., and Lehmann, E., Z. Phys.

Chem. B19, 1 (1932). [68] Huttig, G. F., and Strial, K., Z. Anorg. Allgen. Chem.

2 09, 249 (1932). [69] Hyatt, E. P., Cutler, I. B., and Wadsworth, M. E., J. Am.

Ceram. Soc. 41, 70—4 (1958). [70] Illner, E., Z. Chem. 6, 190—1 (1966). [71] Ilschner, B., Z. Metallkunde 55, 153 (1964). [72] Ingraham, T. R., and Marier, P., Can. J. Chem. Eng. 41,

170-3 (1963). [73] Jacobs, P. W. M., and Tompkins, I. C., Ch. 7 in “Chemistry of the Solid State,” W. E. Garner, Ed. (Butter- worths. 1955). [74] Jaffray, J., and Martin, P., Compt. Rend. 2 36, 1755-6

(1953). [75] Janz, G. J., and Lorenz, M. R., J. Electrochem. Soc. 108,

1052 (1961). [76] Kadlets, O., and Dubinin, M. M., Izvest. Akad. Nauk.

SSSR, Otdel Khim. Nauk, 1031 (1958). [77] Janz, G. J., Neuenschwander, E., and Kelly, F. J., Trans. Faraday Soc. 59, 841-5 (1963). [78] Jenkins, I. L., and Waterman, M. J., J. Inorg. Nucl. Chem.

26, 131-7 (1964). [79] Kadlets, O., and Dubinin, M. M., Izvest. Akad. Nauk.

SSSR, Otdel Khim. Nauk. 1031—6 (1958). [80] Kostryukov, V. N., and Kalinkina, I. N., Zh. Fiz. Khim. 38,

780-1 (1964). [81] Kadlets, O., and Dubinin, M. M., Izvest. Akad. Nauk.

SSSR, Otd. Khim. Nauk. 390 (1961). [82] Kappel, H., and Huttig, G. F., Kolloid Z. 91, 117-34

(1940). [83] Kapustinskii, A. F., and Stakhanova, M. S., Dokl. Akad.

Nauk. SSSR 57, 575 (1947). [84] Khlapova, A. N., Doklady Akad. Nauk. SSSR 116,

979-82 (1957). [85] Krustinsons, J., Z. Elektrochem. 39, 193 (1933). [86] Lander, J. J., J. Chem. Phys. 17, 892-90 (1949). [87] Langmuir, I., J. Am. Chem. Soc. 38, 2263 (1916). [88] Lebeau, P., Compt. Rend. 136, 1256-7 (1903); 137,

1255 (1903). [89] MacDonald, J. Y., Trans. Faraday Soc. 47, 860 (1951). [90] MacDonald, G. J. F., Am. Mineralogist 41, 744 (1956). [90a] Mah, A. D., U.S. Bur. Mines, Rept. Invest. 5600 (1960). [91] Mah, A. D., U.S. Bur. Mines, Rept. Invest. 6171 (1963). [91a] Mah, A. D., Pankratz, L. B., Weller, W. W., and King,

E. C., U.S. Bur. Mines, Rept. Invest. 7026 (1967). [92] Mampel, K., Z. Phys. Chem. A187, 43,235 (1940). [93] Marshall, A. L., and Bruzs, B., J. Phys. Chem. 29, 1184

(1925). [94] Mohanty, S. R., and Ray, M. N., Indian J. Chem. 3, 37

(1965). [94a] Motzfeld, K., J. Phys. Chem. 59, 139 (1955). [95] Muraishi, K., Yamagata Daigaku Kiyo, Shizen Kagaku 6,

187-96 (1964). [96] Narsimhan, G. V., Chem. Eng. Sci. 16, 7 (1961). [97] Nicholson, D., Trans. Faraday Soc. 61, 990 (1965). [98] Ostroff, A. G., and Sanderson, R., J. Inorg. Nucl. Chem.

9, 45 (1959). [99] Otto, E., J. Electrochem. Soc. 1 1 1, 88 (1964). ,[100] Otto, E., J. Electrochem. Soc. 112, 367 (1965). [100a] Otto, E., J. Electrochem. Soc. 113, 643 (1966).

| [101 1 Pannetier, G., Nataf, J., and Dereigne, A., Bull. Soc. Chim. France 1965 (2), 318-21.

[102]

Pauling, L., The Nature of the Chemical Bond, 3d Ed.

(Cornell Univ. Press, Ithaca, N.Y., 1960), p. 93.

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

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

1965), pp. 407-21.

[105] Polanyi, ivi., and Wigner, E., Z. Phys. Chem. A139, 439

(1928). [106] Popov, M. M., and Ginzburg, D. M., J. Gen. Chem. USSR

26, 1107-17 (1956),

[107] Prodan, E. A., and Pavlyuchenko, M. M., Geterogennye

Khim. Reaktsii 1961, 79-91.

[108] Prodan, E. A., and Pavlyuchenko, M. M., Geterogennye

Khim.Reaktsii, Inst. Obshch. i Neorgan. Khim. Akad.

Nauk. Belorussk, SSR 1965 , 64—79.

[109] Rao, C. N. R., Yoganarasimhan, S. R., and Lewis, M. P., Can. J. Chem. 38, 2359-62 (1960). [110] Rassa, M„ Compt. Rend. 2 36, 1673-5 (1953); Ann. Chim.

(12) 8, 755 (1953). [111] Reisman, A., J. Am. Chem. Soc. 80, 3558-61 (1958). [112] Rolin, M., and Recapet, J. M., Bull. Soc. Chim. France 1964 , 2504-10.

[113] Rose, A. J., Bull. Soc. Franc. Min. 71, 16-108 (1948). [114] Rufov, Yu. N., and Markova, Z. A., Kinetika i Kataliz 6,

731-2 (1965). [115] Samal, G. I., Geterogennye Khim. Reaktsii 1961, 92—8.

[116] Saslawski, I. I., Z. Anorg. Chem. 146, 319 (1925). [117] Sass, R. L., Vidale, R., and Donohue, J., Acta Cryst. 10,

567 (1957). [118] Shannon, R. D., Trans. Faraday Soc. 60, 1902 (1964). [119] Simpson, D. R., Fisher, R., and Libsch, K., Am. Mineralogist 49, 1111-14 (1964). [119a] Slater, J. C., Phys. Rev. 36, 57 (1930). [120] Spencer, W. F., and Topley, B., J. Chem. Soc. (London)

192 9, 2633-50.

[121] Stern, K. H., and Weise, E. L., High-Temperature Properties of Inorganic Salts, Part 1, Sulfates, NSRDS-NBS 7

(1966). [121a] Stern, K. H., J. Chem. Educ., in press. [122] Subba Rao, G. V., Natarajan, M., and Rao, C. N. R., J. AmCeram. Soc. 51, 179 (1968). [123] Swanson, H. E., et al., Standard x-Ray diffraction powder

patterns, Nat. Bur. Stds. Circular 539.

(a) with Fuyet, R. K., Part 2 (1953).

(b) with Gilfrich, N. T., and Cook, M. I., Part 7 (1957).

(c) with Gilfrich, N. T., Cook, M. I., Stinchfield, R., and

Parks, P. C„ Part 8 (1959).

(d) with Cook, M. I., Evans, E. A., and deGroot, J. H., Part 10 (1960). [124] Swanson. H. E., Morris, M. C., Stinchfield, R. P., and

Evans, E. H., Standard x-Ray diffraction powder

patterns, Nat. Bur. Stds. Monograph 25, Part 1 (1962). [125] Tranguard, A., Lacassagne, C., Boinon, M. J., and Capella.

L., C. R. Acad. Sci., Paris. Ser. C264, 1111—4 (19677. [126] Vickery, R. C., Chemistry of the Lanthanons (Academic

Press, New York, 1953), pp. 217-219.

[127] Vorlander, D., Hollatz, J., and Fischer, J., Ber. 65, 536

(1932). [128^ Wanmaker, W. L., and Radielovic, D., 5th Intern. Symp.

Reactivity Solids, Munich 1964, 529-39 (Pub. 1965). [129] Young, D. A., Decomposition of Solids (Pergamon Press, New York, N.Y., 1966). [130] Zawadzki, J., and Bretsnajder, S., Z. Elektrochem. 41,

721 (1935). [131] Zawadzki, J., and Bretsnajder, S., Trans. Faraday Soc. 34, 951 (1938). [132] Zawadzki, M. J., and Szamborska, W., Bull. Intern. Acad.

Polon. Sci., Classe sci. Math. Nat., Ser. A., 1940-46,

27-31 (1948). [133] Zhabrova, G. M., and Gordeeva, V. A., Izvest. Akad.

Nauk. SSSR, Otdel Khim. Nauk. 1959, 596-9.

27

U.S. GOVERNMENT PRINTING OFFICE : 1969 01-357-056
'

'
"

(cut

here)

Announcement of New Publications on

Standard Reference Data

Superintendent of Documents,

Government Printing Office,

Washington, D.C. 20402

Dear Sir:

Please add my name to the announcement list of new publications to be

issued in the series: National Standard Reference Data Series — National Bureau

of Standards.

Name

Company

Address

City State Zip Code

(Notification Key N337)

(cut

here)

Publications in the National Standard Reference Data Series

National Bureau of Standards

You may use this listing as your order form by

checking the proper box of the publication(s) you

desire or by providing the full identification of the

publication you wish to purchase. The full letter

symbols with each publications number and full

title of the publication and author must be given

in your order, e.g. NSRDS-NBS-17, Tables of

Molecular Vibrational Frequencies, Part 3, by

T. Shimanouchi.

Pay for publications by check, money order, or Superintendent of Documents coupons or deposit

account. Make checks and money orders payable

to Superintendent of Documents. Foreign remitNSRDS-NBS 1, National Standard Reference Data SystemPlan of Operation, by E. L. Brady and M. B. Wallenstein, 1964

(15 cents). NSRDS-NBS 2, Thermal Properties of Aqueous Uni-univalent

Electrolytes, by V. B. Parker, 1965 (45 cents). NSRDS-NBS 3, Sec. 1, Selected Tables of Atomic Spectra,

Atomic Energy Levels and Multiplet Tables, Si II, Si III, Si IV, by C. E. Moore, 1965 (35 cents). NSRDS-NBS 3, Sec. 2, Selected Tables of Atomic Spectra,

Atomic Energy Levels and Multiplet Tables, Si I, by C. E. Moore, 1967 (20 cents). NSRDS-NBS 4, Atomic Transition Probabilities, Volume 1, Hydrogen Through Neon, by W. L. Wiese, M. W. Smith and

B. M. Glennon, 1966 ($2.50).

NSRDS-NBS 5, The Band Spectrum of Carbon Monoxide, by

P. H. Krupenie, 1966 (70 cents). NSRDS-NBS 6, Tables of Molecular Vibrational Frequencies,

Part 1, by T. Shimanouchi, 1967 (40 cents). NSRDS-NBS 7, High Temperature Properties and Decomposition of Inorganic Salts, Part 1, Sulfates, by K. H. Stern and

E. L. Weise, 1966 (35 cents). NSRDS-NBS 8, Thermal Conductivity of Selected Materials, by

R. W. Powell, C. Y. Ho, and P. E. Liley, 1966 ($1).

NSRDS-NBS 9, Bimolecular Gas Phase Reactions (rate coefficients), by A. F. Trotman-Dickenson and G. S. Milne, 1967

($2). NSRDS-NBS 10, Selected Values of Electric Dipole Moments

for Molecules in the Gas Phase, by R. D. Nelson, Jr., D. R. Lide, Jr., and A. A. Maryott, 1967 (40 cents). NSRDS-NBS 11, Tables of Molecular Vibrational Frequencies,

Part 2, by T. Shimanouchi, 1967 (30 cents). NSRDS-NBS 12, Tables for the Rigid Asymmetric Roto: Trans- formation Coefficients from Symmetric to Asymmetric Bases and Expectation Values of P*, P*, and P«, by R. H. Schwendeman, 1968 (60 cents). NSRDS-NBS 13, Hydrogenation of Ethylene on Metallic Catalysts, by J. Horiuti and K. Miyahara, 1968 ($1).

NSRDS-NBS 14, X-Ray Wavelengths and X-Ray Atomic Energy

Levels, by J. A. Bearden, 1967 (40 cents). NSRDS-NBS 15, Molten Salts, Vol. 1, Electrical Conductance,

Density, and Viscosity Data, by G. Janz, F. W. Dampier, G. R. Lakshminarayanan, P. K. Lorenz, and R. P. T. Tomkins, 1968

($3).

tances should be made either by international

money order or draft on an American bank. Postage stamps are not acceptable.

No charge is made for postage to destinations

in the United States and possessions, Canada,

Mexico, and certain Central and South American

countries. To other countries, payments for documents must cover postage. Therefore, one-fourth

of the price of the publication should be added for

postage.

Send your order together with remittance to

Superintendent of Documents, Government Printing Office, Washington, D.C. 20402.

NSRDS-NBS 16, Thermal Conductivity of Selected Materials, Part 2, by C. Y. Ho, R. W. Powell, and P. E. Liley, 1968 ($2).

NSRDS-NBS 17, Tables of Molecular Vibration Frequencies,

Part 3, by T. Shimanouchi, 1968 (30 cents). NSRDS-NBS 18, Critical Analysis of the Heat-Capacity Data

of the Literature and Evaluation of Thermodynamic Properties of Copper, Silver, and Gold From 0 to 300 K, by G. T. Furukawa, W. G. Saba, and M. L. Reilly, 1968 (40 cents). NSRDS-NBS 19, Thermodynamic Properties of Ammonia as an Ideal Gas, by L. Haar, 1968 (20 cents). NSRDS-NBS 20, Gas Phase Reaction Kinetics of Neutral Oxygen Species, by H. S. Johnson, 1968 (45 cents). NSRDS-NBS 21, Kinetic Data on Gas Phase Unimolecular Reactions, by S. W. Benson and H. E. O’Neal (In press).

NSRDS-NBS 22, Atomic Transition Probabilities, Vol. II, Sodium Through Calcium, A Critical Data Compilation, by W. L. Wiese, M. W. Smith, and B. M. Miles (In press).

NSRDS-NBS 23, Partial Grotrian Diagrams of Astrophysical

Interest, by C. E. Moore and P. W. Merrill, 1968 (55 cents). NSRDS-NBS 24, Theoretical Mean Activity Coefficients of Strong Electrolytes in Aqueous Solutions from 0 to 100° C, by Walter J. Hamer, 1968 ($4.25).

NSRDS-NBS 25, Electron Impact Excitation of Atoms, by B. L. Moiseiwitsch and S. J. Smith, 1968 ($2).

NSRDS-NBS 26, Ionization Potentials, Appearance Potentials,

and Heats of Formation of Positive Ions, by J. L. Franklin,

J. G. Dillard, H. M. Rosenstock, J. T. Herron, K. Draxl, and

F. H. Field ($4).

NSRDS-NBS 27, Thermodynamic Properties of Argon from the

Triple Point to 300 K at Presures to 1000 Atmospheres, by

A. L. Gosman, R. D. McCarty, and J. G. Hust ($1.25).

NSRDS-NBS 28, Molten Salts, Vol. 2, Section 1, Electro- chemistry of Molten Salts: Gibbs Free Energies and Excess

Free Energies From Equilibrium-Type Cells, by G. J. Janz and

C. G. M. Dijkhuis. Section 2, Surface Tension Data, by G. J.

Janz, G. R. Lakshminarayanan, R. P. T. Tomkins, and J. Wong

($2.75).

NSRDS-NBS 29, Photon Cross Sections, Attenuation Coefficients and Energy Absorption Coefficients From 10 keV to 100 GeV, J. H. Hubbell (75 cents).

U.S. GOVERNMENT PRINTING OFFICE : 1 969 0-357-056



NBS TECHNICAL PUBLICATIONS

PERIODICALS

JOURNAL OF RESEARCH reports National

Bureau of Standards research and development in

physics, mathematics, chemistry, and engineering.

Comprehensive scientific papers give complete details

of the work, including laboratory data, experimental

procedures, and theoretical and mathematical analyses. Illustrated with photographs, drawings, and

charts.

Published in three sections, available separately: • Physics and Chemistry

Papers of interest primarily to scientists working in

these fields. This section covers a broad range of

physical and chemical research, with major emphasis on standards of physical measurement, fundamental

constants, and properties of matter. Issued six times

a year. Annual subscription: Domestic, $9.50; foreign, $11.75*. • Mathematical Sciences

Studies and compilations designed mainly for the

mathematician and theoretical physicist. Topics in

mathematical statistics, theory of experiment; design,

numerical analysis, theoretical physics and chemistry, logical design and programming of computers

and computer systems. Short numerical tables. Issued quarterly. Annual subscription: Domestic,

$5.00; foreign, $6.25*. • Engineering and Instrumentation

Reporting results of interest chiefly to the engineer

and the applied scientist. This section includes many

of the new developments in instrumentation resulting

from the Bureau’s work in physical measurement,

data processing, and development of test methods.

It will also cover some of the work in acoustics,

applied mechanics, building research, and cryogenic

engineering. Issued quarterly. Annual subscription:

Domestic, $5.00; foreign, $6.25*.

TECHNICAL NEWS BULLETIN

The best single source of information concerning the

Bureau’s research, developmental, cooperative and

publication activities, this monthly publication is designed for the industry-oriented individual whose

daily work involves intimate contact with science and

technology—for engineers, chemists, physicists, re- search managers, product-development managers, and

company executives. Annual subscription: Domestic,

$3.00; foreign, $4.00*.

• Difference in price is due to extra cost of foreign mailing.

Order NBS publications from:

NONPERIODICALS

Applied Mathematics Series. Mathematical tables,

manuals, and studies.

Building Science Series. Research results, test methods, and performance criteria of building materials, components, systems, and structures. Handbooks. Recommended codes of engineering

and industrial practice (including safety codes) developed in cooperation with interested industries,

professional organizations, and regulatory bodies.

Special Publications. Proceedings of NBS conferences, bibliographies, annual reports, wall charts,

pamphlets, etc. Monographs. Major contributions to the technical

literature on various subjects related to the Bureau’s

scientific and technical activities. National Standard Reference Data Series.

NSRDS provides quantitive data on the physical

and chemical properties of materials, compiled from

the world’s literature and critically evaluated.

Product Standards. Provide requirements for sizes,

types, quality and methods for testing various industrial products. These standards are developed cooperatively with interested Government and industry groups

and provide the basis for common understanding of

product characteristics for both buyers and sellers. Their use is voluntary.

Technical Notes. This series consists of communications and reports (covering both other agency and

NBS-sponsored work) of limited or transitory interest.

Federal Information Processing Standards Publications. This series is the official publication within

the Federal Government for information on standards

adopted and promulgated under the Public Law

89—306, and Bureau of the Budget Circular A—86

entitled, Standardization of Data Elements and Codes

in Data Systems.

CLEARINGHOUSE

The Clearinghouse for Federal Scientific and

Technical Information, operated by NBS, supplies

unclassified information related to Government-generated science and technology in defense, space,

atomic energy, and other national programs. For

further information on Clearinghouse services, write:

Clearinghouse

U.S. Department of Commerce

Springfield, Virginia 22151

Superintendent of Documents

Government Printing Office

Washington, D.C. 20402
U.S. DEPARTMENT OF COMMERCE

WASHINGTON, D.C. 20230