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PUBLICATIONS

V

IMBSIR 80-2029

The Thermochemical Properties of

the Uranium-Halogen Containing

Compounds

I

Vivian B. Parker

Center for Thermodynamics and Molecular Science

Chemical Thermodynamics Division

National Bureau of Standards

U.S. Department of Commerce

Washington, D.C. 20234

July 1 980

interim Report

Prepared for

100

U56

80-2029

1980

International Atomic Energy Agency

i^nd

fice of Standard Reference Data, NBS

C.2

AnOMAL BURKALl

Pt nANDARUS

UBBART

SEP 1 9 1980

NBSIR 80-2029

» f « THE THERMOCHEMICAL PROPERTIES OF

THE URANIUM-HALOGEN CONTAINING

COMPOUNDS

Vivian B. Parker

Center for Thermodynannics and Molecular Science

Chemical Thermodynamics Division

National Bureau of Standards

U.S. Department of Commerce

Washington, D.C. 20234

July 1980

Interim Report

Prepared for

International Atomic Energy Agency

and

Office of Standard Reference Data, NBS

DATE DUE

NOV 0 6 ^ « Ta.^ ^!2.m

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GAYLORD PRINTED IN U.S.A.

U.S. DEPARTMENT OF COMMERCE, Philip M. Klutznick, Secretary

Luther H. Hodges, Jr., Deputy Secretary

Jordan J. Baruch, Assistant Secretary for Productivity, Technology, and Innovation

NATIONAL BUREAU OF STANDARDS, Ernest Ambler, Director
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The Thermochemica 1 Properties of the

Uranium-Halogen Containing Compounds

Vivian B. Parker

Table of Contents

I. General Description

1. Introduction 2

2. Explanation of the Contents of the Text and

Tables 3

2.1 Conventions Regarding Pure Substai.ves 3

2.2 Convention Regarding Solutions for the Tabulated

Values 4

2.3 Convention Regarding Solutions for Reactions in

the Text 6

2.4 Chemical Formulae and Physical States for the

Tabulated Values 6

2.5 Definition of S3nnbols for Thermochemical Properties

for the Tabulated Values 7

3. Unit of Energy and Fundamental Constants 7

4. Internal Consistency of the Tables 8

4.1 The Use of Auxiliary Data 8

4.2 Uncertainties 10

4.3 Relationship to Other Tables of Thermodynamic

Properties 10

5. Arrangement of the Tables 12

II. The Evaluation and Analysis of the Thermochemistry of

the Uranium-Halogen Containing Compounds

1. The Evaluation Procedure 18

2. The Analysis of the Data 20

3. The Evaluations of the Uranium-Halogen Containing

Compounds 21
21

3.100

U-X (X=F^Cl,Br or I) Compounds with/without 0 . .

3.101 U-F Compounds 21

3.102 U-F-0 Compounds 59

3.103 U-Cl Compounds 69

3.104 U-Cl-0 Compounds s6

3.105 U-Cl-F Compounds 105

3.106 U-Br Compounds 106

3.107 U-Br-0 Compounds 113

3.108 U-Br-Cl Compounds 119

3.109 U-I Compounds 121

3.110 U-I-Cl-Br Compounds 126

3.200 U-X Compounds with N 128

3.201 U-NH^-F-0 Compounds 128

3.300 U-X Compounds and the Alkali Metals^ M

(M = Li, Na, K, Rb, or Cs) 131

3.301 U-F-M Compounds 131

3.302 U-F-O-M Compounds 133

3.303 U-Cl-M Compounds 135

3.304 U-Br-M Compounds 138

II. Tables of Values for 4Gf®, S®,

298.15 K and 4Hf® at 0 K for the Uranium- Halogen Containing

Compounds 139

r/. Bibliography 149

V. Appendix 160

ii
The Thermochemical Properties of the

Uranium-Halo gen Containing Compounds

Vivian B. Parker

Abstract

A detailed analysis and evaluation of the thermochemistry of

142 uranium- halo gen containing compounds is presented, and a tabular summary of the thermochemical properties is given. The properties given, where data are available, are the enthalpy of formation,

<iHf®, Gibbs energy of formation, ^f°, entropy, S°, heat capacity,

Cp”, and the enthalpy difference, (H-Hq), all at temperature

298.15 K, and the enthalpy of formation at 0 K, The values

are consistent with the CODATA Key Values for Thermodynamics.

The analysis of the uranium-halogen containing compounds

includes some vapor pressure equations and Cp equations. Some

thermal functions which are not readily available in the accessible

open literature are presented in the Appendix for compounds that

reqtiired their use.

The tabular summary of zlHf”, ^f®, etc. includes the values for

those uranium compounds which were necessary for this evaluation.

Keywords: Data evaluation; enthalpy; entropy; Gibbs energy; heat

capacity; thermochemical tables; uranium-halogen containing compounds.

-1-
I, General Description

1 . Introduction

This report presents a detailed analysis and evaluation of the

thermochemistry of 142 uranium-halogen containing compounds and a

tabular summary of AHf °

' s, z2fif°'s^ S°'s^ Cp®'s^ (H-Hq)'s at 298 K,

AHf°Q's for these compounds. It is part of an independent ongoing

evaluation of the thermochemistry of the actinide compounds for inclusion in the NBS Technical Note 270 Series. The interest of the

IAEA in sponsoring a series on the thermochemistry of the actinides

by compound class ^ both at low and high temperatures has led to our k

participation in this effort .

Since this IAEA review uses the CODATA (1975) recommended values

for the auxiliary key compounds, all compounds considered here have

been evaluated on this basis and represent part of our contribution

to the Halide Chapter of the IAEA sponsored series, "The Chemical

Thermodynamics of Actinide Elements and Compounds".

The analysis of the uranium “halogen containing compounds includes

some vapor pressure equations and C^ equations. Some thermal functions

which are not readily available in the accessible open literature

are presented in the Appendix for compounds that required their use.

The tabular summary of AHf”, zXf^, etc. includes the values for

those uranium compounds which were necessary for this evaluation.

The sources for these values are given in 4.1.

Users are invited to comment on the selections, correct errors

and bring new measurements to our attention.

See Wagman et al. (1977). References are listed by author and year

in the bibliography (Section. IV).

- 2 -
2. Explanation of the Contents of the Text and Tables

The following material provides definitions and conventions used

in the tables.

2.1 Conventions Regarding Pure Substances

The values of the thermodynamic properties of the pure substances

given in these tables are for the substances in their standard states.

These standard states are defined as follows:

For a pure solid or liquid, the standard state at any temperature

is the substance in the condensed phase under a pressure of one

atmosphere

For a gas the standard state at any temperature is the hypothetical

ideal gas at unit fugacity, in which state the enthalpy is that of the

real gas at the same temperature and at zero pressure.

The phase of a substance is indicated in parentheses at the end

of the chemical formula (see Table A).

The values of ZlHf° and AGf° given in the tables represent the

change in the appropriate thermodynamic quantity when one mole of

the substance in its standard state is formed, isothermally at the

indicated temperature, from the elements, each in its appropriate

standard reference state. The standard reference state at 298 K for

each element except phosphorus has been chosen to be the standard

state that is thermodynamically stable at 298 K and at one atmosphere

pressure. For phosphorus the standard reference state is the crystalline

white form; the more stable forms have not been well characterized

-3 -
thermochemical ly. The same reference states have been maintained for

the elements at 0 K except for the liquid elements bromine and mercury

for which the reference states have been chosen as the stable

crystalline forms. The standard reference states for the elements

are indicated in the tables by the fact that the values of AHf® and

^f® are exactly zero.

The values of S° represent the virtual or "thermal” entropy of

the substance in the standard state at 298,15 K, omitting contributions from nuclear spins and isotopic mixing. Where data have been

available only for a particular isotope^ they have been corrected

when possible to the normal isotopic composition.

2.2 Convention Regarding Solutions for the Tabulated Values

For all dissolved substances the composition of the solvent is

indicated in parentheses following the chemical formula. Except in

special cases^ discussed below^ the number of moles of the solvent

associated with one mole of solute is stated explicitly. See

Table A for the conventions used.

In some cases the concentration of the solute can not be specified.

These are indicated as "AU" (aqueous, unspecified) for water solutions

and by ''ll" for non-aqueous and mixed media. In all these cases the

solution may be assumed to be "dilute".

The standard state for a non-dissociated solute in aqueous solution

is taken as the hypothetical ideal solution of unit molality, which

has been designated as "std. state, m = 1". For strong electrolytes

in aqueous solution the conventional standard state is the ideal

-4-
solution of unit activity (unit mean molality). The designation ”A"

is used for strong electrolytes in the standard state and "AO” for

undissociated species in water solution. In non-aqueous media two

standard states are commonly used. For the mole fraction scale,

"std. state, y:^ = 1”, x is added to the formula of the solvent. For

the molal scale, "std. state, m = 1", either s or M is appended to

the formula.

The value of ZlHf® for a solute in its standard state is equal

to the apparent molal enthalpy of formation of the substance in the

infinitely dilute real solution, since the enthalpy of dilution of an

ideal solution is zero. At this dilution the partial molal enthalpy

is equal to the apparent molal quantity. At concentrations other

than the standard state, the value of represents the apparent

enthalpy of the reaction of formation of the solution from the

elements comprising the solute, each in its standard reference state,

and the appropriate total number of moles of solvent. In this representation the value of for the solvent is not required. The

experimental value for a heat of dilution is obtained directly as the

difference between the two values of zlHf® at the corresponding concentrations. At finite concentrations the partial molal enthalpy

of formation differs from the apparent enthalpy.

The values of the thermodynamic properties tabulated for the

individual ions in aqueous solution are based on the usual convention

-5 -
that the values of AHf®, S® and Cp® for (aq, std. state,

m = 1) are zero. The properties of a neutral electrolyte in aqueous

solution in the standard state are equal to the algebraic sum of

these values for the appropriate kinds and numbers of individual ions

assumed to constitute the molecule of the given electrolyte. For an

ionic aqueous species e.g., HSO^ , the properties tabulated refer to

that undissociated ion, i.e., they are not equal to the sum of those

for its constituent ions. By adopting the above convention with

respect to aqueous H^, it follows that the thermodynamic relation

z2Gf® = AHf®-T(ZiSf® + n*0.5S®(H2)) holds for individual ionic species,

with n equal to the algebraic value of the charge. For neutral

electrolytes the normal consistency relationship applies. See

section 4.

2.3 Convention in the Text (Section II) for Reactions Involving Solutions

In the text reactions involving solutions are given as the authors

presented them in the literature. The discussion accompanying the

reactions gives the interpretation.

2.4 Chemical Formulae and Physical States for the Tabulated Values (Section

These tables were reproduced from computer printout in which only

capital (upper case) letters are available. Normal one-line chemical

! formulae are used, with the following modifications:

Subscripts (counts of atoms) : UCL'4 = UCl^ I The apostrophe ' appears after each letter in a chemical ||

symbol that would normally be written in lower case:

^

UCL'BR'3 = UClBr^ i,

-6 -
- The centered dot^ used in hydrates and minerals^ is shown

as a colon: UF4:2.5H20 = UF ‘Z.SH.O

4 2

- The physical state of the substance is appended to the

chemical formula in parentheses: UCL'4(C) = UCl^^ crystalline

Conventions with respect to physical state are given in Table A,

2.5 Definition of Symbols for Thermochemical Properties for the

Tabulated Values (Section III)

The headings used in the tabulated Thermochemical Values have the

following meanings;

DH298 = ZlHf ”, standard enthalpy of formation at 298.15 K;

DG298 = standard Gibbs energy of formation at

298.15 K;

The others are self explanatory. All H and G relationships are

in kcal/mol; all S and Cp® relationships ^e in cal/mol*K.

All values refer to one mole of substance for the formula given.

3 . Unit of Energy and Fundamental Constants

All of the energy values given in these tables are expressed in

terms of the thermochemical calorie. This unit, defined as equal to

4.184 joules exactly, is used throughout the IAEA review series.

Values reported in other units have been converted to calories by

means of the conversion factors for molecular energy given in

Table B.

Values in this report are consistent with the CODATA fundamental

constants (1973).

The formula weights in the tables have been calculated from the

molecular formula using the 1969 Table of Atomic Weights as given by

lUPAC (1970).

- 7 -
4. Internal Consistency of the Tables

The processes given in the text have been obtained from the

original articles, using consistent values for all subsidiary and

auxiliary quantities. The original data were corrected where possible

for differences in energy units, molecular weights, temperature scales,

etc. Thus we have sought to maintain a uniform scale of energies for

all processes in the text and tables. In addition the final tabulated

values of the properties of a substance satisfy all the known physical

and thermodynamic relationships among these properties. The quantities

zHif®, zlGf®, and S“ at 298.15 K satisfy the relation:

^f® = - T^f“

to the precision given. The special case of solutions is discussed

in section 2.2. Furthermore the calculated value of any thermodynamic

quantity for a reaction is independent of the path chosen for the

evalua tion.

4.1 The Use of Auxiliary Data

As indicated this evaluation uses as its basis the CODATA (1975)

recommended values for the auxiliary key compounds. Parker et al. (1976)

have incorporated these values into an extensive consistent set of

CODATA compatible values for auxiliary values needed in this evaluation.

Unless otherwise stated all values for the non-uranium containing

auxiliary compounds are taken from the latter report or are consistent with

it. Values for the thermal functions as a function of temperature for

the auxiliary data are from JANAF (1971)

-8 -
The tables in Section III also include values for those key non-halo gen

containing uranium compounds that were necessary for the present evaluation. These values also maintain consistency with other IAEA chapters.

In addition^ the values listed for the oxides are in agreement with

the independent evaluations of Parker (1975^ 1976a), GQDATA (1978),

Rand et al. (Part XI: The Actinide Oxides) and Glushk? (1978).

These compounds and the sources for their values are listed below.

U(cs) and U(g)

The tabulated values are from Getting et al. (1976).

3+ 4+ 2+

U (aq, std. state), U (aq, std. state) and UO^ (aq, std. state)

The tabulated values are from Fuger and Getting (1976).

U02(c)

S®, G

® and H-H_ are obtained from the G measurements of Westrum

P 0 p

and Huntzicker (1971); AHf® is from the combustion measurements of

Huber and Holley (1969).

U03(c,7)

The S®, Gp® and H-H^ are from the G^ measurements of Westrum (1966).

The AHf® is obtained from the solution calorimetry of Fitzgibbons et al.

(1967) on the uranium oxides and the decomposition measurements on

UG3 (c, 7) of Gordfunke and Aling (1965).

The S®, Gp® and H-Hq are obtained from Westrum and Gr^nvold (1959^

and Girdhar and Westrum (1968), the £iiHf® from the combustion measurements of Huber and Holley (1969).

-9 -
4.2 Uncertainties

Each selected value in the tabular summary, Section III ^contains

the final assigned uncertainty, i.e. the estimate by the evaluator of

the accuracy of the selected value. In some cases these uncertainties

are calculated from (1) the uncertainties assigned to the individual

reactions (usually given as 2^) and the auxiliary data and (2) how well

the calculated values for a property (obtained from the various reaction

paths) agree within their assigned uncertainties.

Where possible or warranted the various uncertainties and the

uncertainty on the selected value are also given in the individual

evaluation (Section II).

In many cases, however, the final uncertainty is based on (2) and/or

the evaluator's judgment, since a strictly mathematical calculation is

not warranted (i.e. (1) is not applicable since no measure of precision

is available, and a quantitative estimate of the uncertainty on the

individual processes can not be made. In these cases no uncertainties

are given in the individual evaluations (Section II)).

In both approaches, however, the final uncertainty is partly

subjective

.

The uncertainties for the Cp°'s and (H-Hq)'s appear in square

brackets and are arbitrarily assigned as "10 in the last place".

4.3 Relationship to Other Tables of Thermodynamic Properties

The chemical thermodynamic properties in the present tables of

halogen containing compounds may be combined with those published by

-10-
GODATA (1975, 1978) and Parker et al. (1976) in order to calculate

the change in a property for a process. However, we recommend against

these values being combined with those in any other tabulation or with

a property reported in an original research paper. In particular, we

warn against indiscriminate combination with the NBS Technical Note

270 Series. Values consistent with auxiliary data in the NBS Technical

Note 270 Series are available and will be included in that series.

There are several reasons for avoiding the combination of thermochemical data from more than one table. The most important is that

different large-scale tables use different thermochemical properties

of formation for substances that are ubiquitous in thermochemical

measurements. Outstanding examples are in the common inorganic acids

and their ions. Another reason is that the groups preparing different

tables may have relied on different measurements as the basis for

selecting property values.

It is difficult to predict a priori how a change in one selected

formation property would affect values assigned to other substances

because of the way these are linked by complex networks. In general,

it may be expected that the advantage of internal consistency of a

table will be lost if values from several sources are combined and

the experimental measurements may be reproduced poorly.

No general, simple algorithm can be suggested for overcoming

this problem. If it becomes necessary to extend a table of data to

substances other than those tabulated, the user is advised to consult

the group that prepared the table about the procedure that is

contemplated.

-11-
5. Arrangement of the Tables

The compounds in the tables are entered according to the Standard

Order of Arrangement, (see Figure 1), by the principle of latest

position. In this scheme, a compound is listed under the element

occurring latest in the list; water of hydration is neglected. Within

a given element-table will be found all of the canpounds of that element

with elements occurring earlier in the order; the arrangement within

a table follows the same ordering.

However, for the alkali metal halogen-containing compounds the

arrangement is by compound class, i.e., ^^4

^

in combination with the alkali metals.

-12-
TABLE A: Physical State Conventions for Section III

The following conventions are used to designate the physical

state of a substance. These apply to the tables in Section III.

This information appears in a parenthetical expression appended to

the molecule'^ formula. Some of the explanations i~ply a thermochemical value, particularly those for solutions. These normally

are used in describing enthalpy measurements.

Basic Symbols Explanation

(G) Gaseous, e.g., HCL'(G) for HCi(g)

(GS) Gaseous reference standard state for an element,

e.g. , 02(GS) for O^Cg)

(C) Crystalline, e.g., NH4CL'(C) forNH^C£(c)

(CS) Crystalline reference standard state for an element, e.g., RB'(CS) for Rb(c)

(L) Liquid, e.g., H20(L) for H^OCi)

(LS) Liquid reference standard state fo an element,

e.g. , BR'2(LS) for Br^C^-)

(AM) Amorphous

(GL) Glassy

(A) Hypothetical standard state of the ideal aqueous

solution at unit activity. For a neutral

electrolyte the value of a property is equal to

the algebraic sum of the values for the ions

assumed to constitute the molecule of the

electrolyte, e.g. HCL'(A) = H+(A) + CL'-(A).

For an ionic species this notation is commonly

used to refer to the undissociated ion as written

e.g., HS04-(A)

-13 -
Symbol Explanation

(AO) Hypothetical standard state of the ideal aqueous

solution at unit activity of the undissociated

(non-ionized) species, e.g., HF(AO) , HF2-(A0)

.

May also be used whenever the designation (A)

could be ambiguous. Note that the descriptions

HS04-(A) and HS04-(A0) are equivalent, but that

HF(A) and HF(AO) are not.

(AU) Aqueous solution of undefined, but usually

dilute, concentration, e.g., XE'03(AU).

The symbols used above occasionally are modified by numbers to

distinguish two substances in the same state that have the same

molecular weight, as for isomers: (AU2) , (C3) . They are also used

in combination with descriptive material, e.g. (C:HE’), (C:AL') etc.

to mean "crystalline, hexagonal , "crystalline, alpha form", etc.

Special notations for substances in solutions

The notations for the "state" of a substance in solution may

combine a definition of the system, e.g. HCi in 220 moles of water,

and a specification of the thermochemical property associated with

it. Usually the thermochemical property is the apparent integral

enthalpy or free energy of formation or an absolute entropy, i.e. the

formation properties of the solvent are not included. If a partial

molal property is tabulated the notation D: ("D" for "differential")

occurs as the first term in the state bracket. The notations given

below illustrate the differences for integral and differential

(partial molal) properties, and extrapolated values. Examples are

given for aqueous, mixed, and non-aqueous solvents.

- 14 -
Symbol Explanation

HCL' (200H20) An aqueous solution of specified composition,

e.g. one mole of HCi in 200 moles H

2 O. The

value of AHf represents the apparent integral

enthalpy of formation.

HCL’ (D:200H20) and

H20(D:HCL'+200H20)

These represent the partial molal (enthalpy of

formation of the substance in a solution of

specified concentration, e.g. the partial molal

enthalpy of formation of HC5. and H2O respectively,

in a solution consisting of 1 mole HC£ and 200

moles H

2

O.

UCL'4(HCL'04+50H20) This describes a solute dissolved in a mixed

solvent, e.g. one mole of UCJI 4 in a mixture of

1 mole of HCiO^ and 50 moles H^O. The value of

AHf represents the apparent integral enthalpy of

formation of the substance, UCi4 , in the medium.

UCL’4(HCL'04+

50H20:AU)

This represents a solute at an unspecified but

usually dilute concentration in a solvent mixture

of fixed composition.

- 15 -
CONVERSION

FACTORS

FOR

UNITS

OF

MOLECULAR

ENERGY

<u

m

m '^0m 1 I 1 1 3 m o O vO O o CN o cn ^ vO ^ LO <—t m

0) <r

X <rXX

X CO vO vO o O <7\ o cn cn m cn cn a o cn O • 'd' CN • . • • > f“^ cn CN

0)

o;3 CN f"4 cn

CJ 1 1 1 3 LO O lO o o cn r-( in 1—1 O o\ X X o X a m ON p" cn -O’ 'd' • • • 1 GO cn 00 ao

^4 CN o cn NO 1 a CN ON ON o 1 <Ti a\ 00 >o X

CN ONm

vl" cn 3 • • » O

(23

NO 00 1 CN 1 I 00 O CN O oo o O >—

CN f-l in

a

X CN

XX

r- NO 1^ 00 :s • • • CN CN

-O' 1 ON

00 o rs. cn cn 1^ St

ON St

X f*4

IS <r in o ON NO o • O cn 1—1 00

1—4 CN

CN ON

•1 o *

psi cn o

<r St 00

in sr « 1-^ 1—t m

ps NOt CN1 1 r-4 O 1 NO- O cn o

1—4 1—4 ON >—<

1^

—t X CN

X o X

NO CN 00 't cn NO • • • NO cn CN o rs. CN cn CN o 00 CN o

CN CO ON 1—4

m NO cn

a ON CN CN

X in o CN

u NO ON in ON • •v ca 00 CN m • oo CM • • • CN —4 LH

cn ON 1—4 cn CN a\ a >0

CJ

CN1 rs. 'o 1 O

O o m 1-4 cn |S. t-H St a O X X CM CN X 1—4

m o f—4

St Ctl ON o

ON CN s in in NO

ca cn St o m oo o

CJ • NO « • cn

CN CN OO

m CN CN

r—4

in NO

CN o O NO in

cn o o NO • o 1—4 o NO CN -O’ a 1-^ 00 o CN NO 00

1—4 o cn ON St

•n • . o • • NO o NO CN 1-4 ON — 1-4 cn

II

II

II CU II 1—4

1-4 3 3 o r—4 o r-4

II

a II o 3 3 a 1-4 3

II

a 1-4 1 o 3 g U o JO a 1-4

ca

a ^4 o o a a c^ J3 3 I "s. \ 03 a 3 u a > o .:ic CQ o 3

r—1 r]t

16

The

underlined

numbers

represent

the

fundamental

values

used

in

deriving

this

table.

The

remaining

factors

were

obtained

by

applying

the

relationships:
STANDARD ORDER OF ARRANGEMENT

- 17 -

Standard

Order

of

Arrangement

of

the

Elements

and

Compounds

based

on

the

Periodic

Classification

of

the

Elements
II. The Evaluation and Analysis of the Thermochemistry

of the Uranium-Ha logen Containing Compounds

1.

The Evaluation Procedure

Although the compounds evaluated are arranged systematically^

it is recommended that the user first read the evaluations for the

key compounds in the following order:

1. UCl^(c)

2. UO^Cl^Cc)

3. UF,(c) and UO.F^ (c)

O A A 4. UF^Cc) and UF^(c)

This arrangement follows the stepwise procedure of evaluating the

thermochemical properties by the sequential method described by

Garvin et al. (1976). In this method^ the values for the key

compounds are determined first, so that a framework or network

of values is developed for those compounds that are essential

for the determination of the properties of many other compounds.

Initially, this procedure involves the compounds whose properties can be determined independently, i.e., they involve no other

compounds of the same element (in this case, uranium) and/or they

depend only on known auxiliary data. Ideally, there should be

definitive measurement paths between these key compounds as

indications of mutual confirmation. Thereafter, the properties

of other compounds dependent on these first selections are chosen.

- 18 -
The halogen compounds listed above are only part of the

uranium key network, which includes all of the following

compounds

:

1. U(c), U^O^Cc), UO^Cc), and U0^(c,7)

2. UCl, (c), UO Cl (c), and UO (NO ) -6H 0(c)

4+^ ^ ^ ^ ^ ^ ^ 3. U (aq, std. state) and UO

2

(aq, std. state)

4. U(g), UN(c), US(c), and UC(c)

5. UFg(c) and UO^F^ (c) , UF^(c) and UF^(c)

The non-halogen compounds listed above in 1, 1 , and 3,

as U(g) from 4 are considered fixed auxiliary data for this

tion (see Section I 4.1).

as well

evalua-

- 19 -
2 . The Analysis of the Data

The purpose of thermochemical data evaluation is the selection

of reliable values^ i.e., ones that can be used with confidence.

When properly documented^ an evaluation not only presents the results

but also explains how they were obtained. This explanation is

analogous to the detailed experimental section of a thermochemical

measurement paper. An explanation of our evaluation is provided

here to assist the user and to reduce the need for re investigation

in the future.

The evaluator usually works with some incomplete records. There

are not enough data to solve all of the problems that arise. Therefore,

it is necessary to ’’squeeze” the maximum amount of information out

of the existing data. That will become apparent in the individual

evaluations that follow.

Complete thermochemical cycles, where possible, are given. The

same measurements are repeatedly considered from different points

of view. The work of a research group on related systems is considered in order to assess the likely accuracy for the system of

immediate interest and to modify or avoid suspect reactions and

paths in determining the "best” values.

- 20-
3. The Evaluations of the Uranium-Halogen Containing Compounds

3.100 U-X (X=F, Cl, Br, or I) with/vithout 0

3.101 U-F Compounds

UF^(g) (n = 1,2,3)

Gurvich al. (1977) have recently published the ideal gas

thermodynamic functions for UF(g), UF^Cg), and UF^Cg) based on estimates

for the molecular constants and frequencies. These estimates are:

Values for Molecular Constants of UF

Molecule

UJeUU X

e e B

e Dg-lO^ Number of

Electronic

States

-1

cm

UF 580 1.72 0.230 0.0014 1.5 82 with statistical weight

and Tg^51000 cm

Values for Molecular Constants of UF

2 and UF^ in the Ground State

Symme try

CTr(U-F), A

<F-U-F, degrees

,117

^1' cm

-1

CM cm

-1

y

cm

-1

cm

g *cm

2.00±0.05

.+50

-20

2920±2000

110

575±100

140±50

525 ±100

Other details are to be found in Gurvich et al.

2.00±0.05

13690±2300

600±100

100±50

550(2)±100

140(2)±50

(1977). The functions

are re tabula ted in the Appendix. They are to be considered approximate.

- 21-
Zmbov (1969) has nade a nass spectrometric study of high temperature equilibria (1201-1313 K) of the following:

UF^(g) -h Ca(g) - CaF(g) -h UF^(g) (1)

UF^(g) + 2Ca(g) - 2CaF(g) + UF

2 (g) (2)

From his reported ilH's at the mean temperature and his tabulated

K's, the following are obtained at 298 K:

kcal/mol j. 2nd law 3rd law

(1) 23.5±4.0 25.6

(2) 21.2±5.0 36.0

The 2nd law value for appears unreasonable compared to

the values for that = 23.5±4.0 kcal/mol is accepted,

resulting in _lHf° UF

2 (g) = -248±5 kcal/mol. 2Bf® UF

2 (g), then, may

be obtained from the 3rd law ~ ^

UF

2 (g) -h Ca(g) - -DF 2

(g) + CaF(g); = 10.4 kcal/mol (3)

The resultant 4Hf® UF

2 (g) = -129±7 kcal/mol.

From a comparison of the ' s for:

UF^(g) ^ U(g) + nF(g)

a reasonable estimate for when n = 1 is 148±5 kcal/mol resulting

in i_Hf® bT (g) = -2±5 kcal/mol.

Obtained using the thermal functions of UF^(g) without the addition

of 3.2 cal/mol*K to S® bF^(g), since a similar contribution may be

necessary on S® UF^^S^ Most of the contribution

would cancel resulting in third law values close to those given.

-22 -
UF^(c), UF^(c) and LT^*2 .SH^OCc)

The entropy of UF^(c) at 298 K has been estimated as 30.0rl.0

r\ cal/mol*K in comparison with iS [ PuF^ (c)-PuF^ (c) ] and S” UF^(c). Krestov

(1972) has estimated the Cp (298-1000 K) as:

Cp = 21.2 + .0073T cal/mol-K.

Thermal functions have been generated from this equation and the

estimated S°

.

Burns et al. (1960) measured the low temperature heat capacities

(1,3-20 K) of UF^(c). These measurements are in agreement (in the

overlapping range) with those of Osborne et al. 's (1955) measurements,

5-300 K. The smoothed functions tabulated by Burns et al. are accepted

here

.

Dworkin (1972) measured the heat content of UF, (c,l) (298-1400 K) .

His smoothed tabulated functions extrapolated to 1600 K are

accepted and are represented by :

~ -9650 + 29.53T + 1.15x10 T + 2.21:cl0^T ~ cal/mol:

(298-1309 K)

^fusion " ^1^230 cal/mol; = 8.6 cal/mol-K; (1309 K)

^‘^298 " 29.57T cal/mol (1309-1400 K)

The earlier results of King and Christensen (1961) are in fair

agreement

.

The measurements of Settle et al. (1963) and Ha\Tnan (1967) are

in good agreement on the direct fluorination of U(c) to UF-(c.g);

o this lent suonort to the _lrif®'s of l!F.(c) and UF, (c) obtained from

3 -iOsborne et al. (1974, 1975) reported S* = 30.14n0.10 and 35.lSn0.10

cal/mol-K for ?uF

2 (c) and ?uF^(c), respectively.
the f luor ina tiona> of UF^Cc) and UF^Cc) to UF^ (g) by Hayman (1967)

resulting in AHf® UF^Cc) ~ -357 and AHf® UF^(c) ~ -454 kcal/mol.

The reappraisal made here of the Ziflf^’s of UFg(c) and (g) necessitate

a complete reevaluation of the 4iHf®'s UF^^c) and UF^(c). A tabular

summary is also given (Table 1).

Ao The Measurements on the Enthalpy of Formation of UF^Cc)

1. Room Temperature Calorimetry

Khanaev and Khripin (1970) measured the following AH's (kcal/mol) at 323

(UO^Cl^ + 3FeCl^ + HCl + 3/4HBF, + 1/4 H^O) , ^

2 2 2 4 2 soln. I

UF^(c) + (3FeCl2 + 3/4H^B02) solv. ; AH^ = +34.43±.12 (1)

00201^(0) + (3FeCl2 + 3/4H2B02)solv. - (UO^Cl^ + ^FeCl^ +

ir^ = -8.79±.03 (2)

OFeCl, + UO.Cl, + HCl + 3/4HBF, + 1/4H.0) , ^ - 3FeCl.(c) +

Z 2 2 42 soln. I 2 (UO.Cl^ + HCl + 3/4HBF, + 1/4H_0) . AHo = 3x (2. 642+.035

)

^ ^ ^ Z, soi.li*

(3)

(UO-Cl, + HCl + 3/4HBF, + 1/4H_0) , + 3FeCl.(c) -» 2 2 4 2 soln, IX 3

(3FeCl„322 + U0_C1. + HCl + 3/4HBF,42+ 1/4 H„0) , soln. V

= 3(-7.198±.05) (4)

(U0_C1_ + 3FeCl. + 9/4H.0 + HCl) , ^ 2H.0 , +

2 2 3 2 soln. IV 2 soln.

(UO-Cl, + 3/4HBF, + 3FeCl- + 1/4H-0 + HCl) . 2 2 4 3 2 soln. V

AHg = +2.124±.01 (5)

Use was made of the following Khanaev (1968) measurements which they

confirmed

(UO.Cl. + 3FeCl- + 3/4H.B0J . __ + 3HF(1.33H,0) 22 3 33 soln. IX 2

(UO-Cl. + 3FeCl. + 3/4HBF, + 9/4H.0)^^, ___ + 3.99H 0(1) 22 3 4 2 soln. Ill 2

AH, = -15.18±,12 (6)

6

(UO-Cl, + 3FeCl. + 3/4HBF, + 9/4H_0) , ___ + HC1(3.91H_0) 22 3 4 2 soln. III 2

(UO^Cl^ + 3FeCl, + 3/4HBF, + 9/4H^0 + HCl) , + 3.91H_0(1) 22 3 4 2 soln. IV 2

^ = +4.04±0.004 (7)

- 24 -
The summation, AH^+ “ ^3 " ^4 + ^5 +

^ = ^8 ^ +30.292±0.25

corrected to 298 K using a ^iCp = 30 cal/mol'K results in:

UO^Cl^Cc) + SFeCl^Cc) + 3HF(aq) + HCl(aq) - 2 H20 ( 1 ) + 3 FeCl

2

(c)

+ UF^(c); = 29.54±0.32 kcal/mol ( 8 )

The integral values for AHf HC1(3.92H20) and AEIf HF( 1 , 33H20)

are appropriate here. The resultant AHf** UFg(c) is -354.9±0.8 kcal/mol.

Khanaev and Khripin also measured the ZXH , of UCl, (c) as part soln 4

of their series of measurements. The reaction follows:

UCl, (c) + (3FeCl_ + 3/4H-B0- + 2H.0) , - (U0.C1_ + 2FeCl- +

4 3 334 soiv. 44 4

FeCl. + 3/4H-B0. + 4HC1) , AH^ = -18.44±.02 (9)

3 33 soln. VI 9

The summation from a rearrangement of the reactions, AH^ + AHg

+ AH^ - 3xAH^ - l/3AHg - 1/3AH^ = AH^^ = -6 . 754 ±0.0 18 , corrected to

298 K with a ACp = -48 cal/mol*K is:

FeCl

2

(c) + UCl^(c) + 3HF(aq) - FeClg(c) + UFg(c) + 3HCl(aq);

AH^q = -5.564±0.20 kcal/mol (10)

so that AHf° UFg(c) = -354.67±0.8 kcal/mol. The two paths are in

good agreement.

As shown in the section on U02Cl2(c), the AHf® U02Cl2(c) =

-297.0±0.5 kcal/mol obtained from the above Khanaev and Khripin

reactions (relating U0^Cl2(c) to UCl^(c)) is in excellent agreement

with the selected value for U 02Cl

2

(c), and this lends support to the

above values for UFg(c).

- 25 -
It is important to note that the value for UF^(c) = -452.3±0.9

kcal/mol derived from Khanaev (1968) [see section on UF^(c) ] and the

above values for UF^Cc) lead to:

UF^(c) + l/ZF^Cg) - UF^(c); = -452.3 + 354.8 = -97.50±0.8 kcal/mol

( 11 ) Hayman (1967) has fluorinated UF

2

(c) as well as UF^(c) and U(c)

to obtain UF^ (g)

UF^(c) + 3/2F2(g) -*UF^(g); = "153.3910.45 (12)

U(c) + 3F2(g) ^UF^(g); = -510.4710.64 (13)

UF4(c) + F^(g) -*UFg(g); = -56.8210.17 (14)

It appears that 1^ 1^ error as shown in the discussion on

UF^(c) and UFg(g). If it is assumed that the error is systematic in

all the measurements, then AHf ° UF

2

(c) = = -357.1 kcal/mol.

If only 1^ 1^ error, and ^^^4 reasonably correct,

then using the selected AHf'* UFg(g), AHf** UF^(c) = -513.25 + 153.39 =

-359.8610.6 kcal/mol and AHf^ UF^(c) = -513.25 + 56.82 = -456.410.5

kcal/mol.

With either assumption one obtains = -96.610.5 kcal/mol,

as contrasted with -97.5 from the Khanaev measurements.

There are other possibilities in interpreting the Hayman results:

the errors in ^^^^4 proportional to the error in ’

the errors are random and the uncertainties on AH, _ , AH, _ , and AH,

,

12-^ 13-' 14

are too small and the <AHf**'s of UF

2 (c) and UF^(c) could be more

positive.

-26-
2. High Temperature EMF Measurements Involving UF^Cc)

The values cited here depend on the estimated thermal functions

for UF^ (c)

.

Heus and Egan (1966) and Markin et al. (1967) have measured the

e.m.f. of cells involving CaF^ as a solid electrolyte.

For

Al(c) + UF

2

(c) - AlF^(c) + U(c) (15)

the values for at 873 K are 4.84±1.0 and 5.40±1.0 kcal/mol,

respectively^ leading to = -3.73±1.5 and -3.16±1.5 kcal/mol at 298 K

and for UF

2

(c) = -357.3±2,0 and -357,8±2.0 kcal/mol, respectively.

The above relationship (15) was combined with other experimental

reactions of Heus and Egan, and Markin et al. to obtain the relationship of UF

2

(c) to MgF

2

(c), which on the NBS Technical Note Scale,

was less dependent upon their F (aq, std.). The following relationship is obtained which can be used with the CODATA ZlHf ° MgF

2

(c):

4/3U(c) + 2MgF2(c) - 2Mg(c) + 4/3UF^(c) (16)

and from the Heus and Egan values ZbGat 873 K = 50.73, so that

= 62.09±1.5 and z^Hf® = -356.5±2.0 kcal/mol. From Markin et al.

Z^ at

873 K = 50.27, = 61.63±1.5 and Z^® for UF^(c) = -356.8±2.0 kcal/mol.

These two relationships indicate that the measured ZiG's are

good within their experimental uncertainty of 1.0 kcal/mol. Because

of the uncertainty in -(G-H«q_)/T for UF (c) and the uncertainties

o j

in ilHf® of AlF^(c) and MgF

2

(c), the assigned uncertainties on the UF^

values are ±2.0 kcal/mol.

- 27 -
Markin et al. have also determined the relationship between UF^(c)

and UP, (c) .

4UF^(c) 3UF^(c) + U(c) (17)

^873K ~ 62.20 and AH® = 61.86±4.0 kcal/mol (3rd law)

This relationship will be used with the results of Heus and Egan^

and Markin et al. to calculate the AHf°'s of UF^(c). The calculated

then is ~ -98 kcal/mol.

B. The Measurements on the Enthalpy of Formation of UF^(c)

1. Room Temperature Calorimetry

Khanaev (1968) has made a series of measurements at 323 K on the

AH . 's, UF (c- monoclinic) = -3.11±.02 kcal and UCl, (c) = -17.91±.05

soln 4 ^ 4

kcal in various HC1-H_B0 aqueous solutions, and on the various AH 's 3 3 mix.

to obtain the relationship between UCl, (c) and UF, (c) . 4 4

The summation of the measured reactions leads to:

UF^(c) + 4HCl(aq) UCl^(c) + 4HF(aq); AH^ = 49.77±0.15 kcal/mol (1)

where ZlH^ includes the correction from AH at 323 K = 51.20±0.13 kcal/mol

using a ACp = 58 cal/mol*K.

The concentrations for HCl(aq) and HF(aq) are the same as for

the similar cycles for UF^(c) by Khanaev and Khripin (1970)^ so that

AHf® UF^(c) = -452.3±0.9 kcal/mol.

As indicated in the discussion on the AHf® UF^(c), the results

of Hayman (1967) on the fluorination of UF^(c) lead to AHf ® UF^

values of -453.7±0.7 (if a systematic error is assumed in all their

measurements) or -456.4±0.5 kcal/mol (if only AH^^ is in error).

The available information on UF^ • 2 .5H20(c) and the various reported

hydrates^ as well as on the other forms of UF^(c)^ must be considered.

-28-
The UF^‘2.5H20 (c) may be obtained from the Maltsev et al .

(1960) calorimetric measurements involving the precipitation of

UF^'Z.SH^OCc) at 293 K. Their measured reactions (kcal/mol) are:

UCl^(c) + 250HCl(aq) - [UCl^ + 250HCl]aq; AH^ = -43. 47 ±0.10 (2)

[UCl^ + 250HCl]aq + 18HF(aq) + 2.5H20(aq) ^ UF^ • 2 .SH^OCc) +

[254HC1 + 14HF]aq; = -20.9610.33 (3)

[250HC1 + 18HF]aq -250HCl(aq) + 18HF(aq); AH^ = 5.3010.12 (4)

The summation,^5AH_ = AH. + AH_ + AH, = -59.1510.4 is for:

2 3 4

UCl^(c) + [250HC1 + 18«F]aq + 2.5H^O(aq) - [254HC1 + 14HF]aq

+ UF^*2.5H20(c) (5)

The authors have calculated

4HF(g) + [254HC1 + 14HF]aq ^ 4HCl(g) + [250HC1 + 18HF]aq;

AH. = 4x(5 .6810 .50) (6)

6

using unevaluated literature data for the AH , HF(g) “* HF(aq) and soin

AHsoin HCl(g) — HCl(aq) to obtain:

4HF(g) + 4HCl(aq) -4HCl(g) + 4HF(aq); AH^ = 5.68x4 (7)

This AH was assumed to be equal to their •

This approach leads to

UCl^(c) + 4HF(g) + 2.5H20(li - UF^.2. 5H20(c) + 4HCl(g);

AHg = -36.4312.1 (8)

which corrected to 298 K = -24 cal/mol*K) is -36.5512.1 kcal

and AHf® UF^«2.5H20(c) = -623.913.0 kcal/mol.

The AH, , , , which will be discussed and evaluated separately

hydration-'

for

UF^(c) + 2.5H20(1) UF^.2.5H20 (c) (9)

- 29 -
is ~ -10 kcal/mol UF^(c). This results in a AHf ° UF^(c) = -443 kcal/mol

which is not reasonable.

A better approach would be to use the experimental summation,

(4) which contains a mixture of HF and HCl on both sides of the equation differing only by the removal of 4HF(aq) and the addition of

4HCl(aq) and assume that the differentials, /AHf for the HCl, HF, and

H^O at the appropriate concentrations are to be used; the assumption,

of course, is that the ZlHf of HCl(aq) and HF(aq) would be the same as

in pure solutions.

Schematically, then, and corrected to 298 K;

UCl^(c) + 4HF(aq) + Z.SH^OCaq) UF^ • 2 .SH^OCc) + 4HCl(aq) ; = -59. 68 ±0.5 kcal/mol (10)

With the differential AHf^’s for HCl (10.04H^O) , H^O, and HFdeOH^O),

the resultant AHf® UF, *2.5H.,0(c) = -633.211.0 kcal/mol which means

4 2

that AHf° UF^(c) is ~ -452.4 kcal/mol.

2. High Temperature Equilibria

Briggs (1960) has reported the K's in the range 622 to 955 K

for the equilibrium:

00^(0) + 4HF(g) ^ UF^(c) + 2H20(g). (11)

The AH = -50.6 kcal/mol UF^(c) at the mean T results

in AH'* = -50,84 kcal/mol and AHf® = -455.88 kcal/mol. The third law

value for AH° = -51.9010.92 results in AHf° = -456.99 kcal.

-30-
The earlier measurements were evaluated by Rand and Kubaschewski (1963)

who cite = -50,000 + 57. 7T cal mol ^ for the range 700 to 900 K.

This results in a 3rd law AH® = -50.12 and UF^(c) = -455.1

kcal/mol. The other measurements cited by Rand and Kubaschewski are:

(1) Johns and Walsh (1945) (temperature range 873-1073 K)

2nd law = -42.83; AHf ® = -447.82 kcal/mol

3rd law zlH = -56.46±2.8; 4Hf® = -461,45

(2) Domange and Wohlhuter (1949) (temperature range 374-773 K)

2nd law zlH = -27.4

3rd law AH = -49.1±9.0

These two sets are given no further consideration.

3. Room Temperature Aqueous Equilibria

Vdovenko et al. (1963) reported for the following equilibria:

U'^(aq) + F‘(aq) - UF^''‘(aq) ; log K = 7.15 (12)

U'^(aq) + 2F'(aq) UF

2 ^‘^(aq) ; log K = 12.4 (13)

U''^(aq) + 3F"(aq) - UF

2

'^(ai); log K = 17.7 (14)

u"^ (aq) + 4F (aq) UF^(aq) ; log K = 23.1 (extrap, value) (15)

Savage and Browne (1960) reported:

UF^-2.5H20(c) ^UF^(aq); K = 1.09xl0‘^ molal (16)

UF^(aq) - UF

2

‘'"(aq) + F"(aq) ; k = 5 .4x10'^ mo la 1 (17)

UF

2

'’’(aq) UF

2

'^^(aq) + F"(aq) ; k = 5.9x 10"^ molal (18)

From the combination of (16) and (15) we obtain:

UF^-2.5H20(c) U'^(aq) + 4P"(aq) + 2.5H20(1); AG® = +36.9 kcal/mol

(19)

which results in £i3f® UF^*2.5H20(c) = -574.9 kcal/mol.

-31-
If it is assumed that S° UF^*2.5H20(c) = 60.0±2.0 cal/mol*K,

then the calculated for UF^*2.5H20(c) = -631.0 kcal/mol and

ZlHf° UF^(c) ~ -450.2 kcal/mol.

Other combinations for the equilibria (19) are possible from

various combinations of the Vdovenko et al. data with the Savage and

Brown data^ such as = +35.4 from (16) + (17) - (14) which would

result in -629.5 kcal/mol for AHf° UF^*2.5H20(c) and 4Hf® =

-448.7 kcal/mol.

4. High Temperature EMF Measurements

As indicated under UF^(c), Markin et al. (1967) have determined

the relationship between UF^(c) and UF^(c)^ as:

^873 K

~ ^2.20 kcal/mol

for

4UF^(c) - 3UF^(c) + U(c) (20)

Their results and Heus and Egan's (1966) EMF measurements on

UF

2 (c) can then be converted to relationships involving UF^(c) and

MgF

2

(c), and UF^(c) and AlF^(c) negating the added uncertainty of

the estimated thermal functions of UF^(c) on the reactions involving

UF3(c).

Then for:

Al(c) + 3/4UF^(c) -*AlF^(c) + 3/4U(c) (21)

from Heus and Egan^ 4H° = -19.2 kcal/mol and ZHIf° = -455.7±1.0 kcal/mol^

and from Markin et al.^ 4H = -18.6 and zlHf’* = -456.5±1.0 kcal/mol.

-32 -
Similarly for:

U(c) + ZMgF^Cc) * 2Mg(c) + UF^(c) (22)

from Heus and Egan, AH® = 82.7 kcal and AHf® = -454,7±1.0 kcal/mol; from

Markin at al AH® = 82.3 and AHf® = -455.1±1.0 kcal/mol.

C. The Selection of Values for the Enthalpies of Formation of

UF^Cc) and UF^(c)

As can be seen from the range of values for AHf®'s for UF^(c)

and UF^(c), Table 1, the selected values can not be considered definitive;

however a reasonable presentation of the relationships can be made

with AHf® UF^(c) = -455.5±1.0 kcal/mol.

A(AHf® UF^(c) - AHf® UF^(c)) = 98.2±1.0 kcal/mol

AHf® UF^(c) = -357. 3±1.5kca 1/mol

The data derived from the aqueous equilibria measurements are

suspect; the hydrate system is a complex one; the K's given are not

true equilibrium constants.

The more disturbing values are those from the Khanaev measurements; many reactions are involved in their paths to UF^(c) and UF^(c)

from U02Cl2(c) and UCl^(c) ; it does not appear reasonable to assume

that all the measurements are in error and that these errors are

additive, particularly since the relationship between UCl^(c) and

is in excellent agreement with that from the selected

values for zlHf®'s UCl^(c) and UO^Gl^Cc). One of the reactions in

common for both compounds is the AH . of HF(1.33H„0) with the

appropriate solution; if the AH mix. should be 0.7 kcal/mol HF more

negative, the A(AHf® UF^-AHf® UE^) would become 98.2 in agreement

with the Markin et al (1967) EMF measurements and would result in

- 33-
AHf ** of UF^Cc) and UF^(c) -357 and -455 kcal/mol, respectively. The

UF^ *2. 511^0 value may have a similar error in AH mix. Obviously all

of this is conjecture.

D. The Enthalpy of Hydration of UF, (c, monoclinic) to

UF^*2,5H20(c, orthorhombic)

As indica+'ed in C, from the Maltsev et al, (1960) reactions

involving the precipitation of UF^»2.5H20 (c) from UCl^-HCl solutions^

AHf® UF^*2.5H20 (c) = -633.2±1.0 kcal/mol.

Gagarinskii and Khanaev (1967) have measured the AH's of solution

of UF^(c, monoclinic) and UF^*2.5H20(c) at 323 K and 298 K to the

same final solution.

Their measurements are:

UF^(c) + [166.5HC1 + 3 H^B02 ] ^ [UCl^ + 3H

2

O + HBF^ +

162.5HC1 + 2H2BO^]aq; AH

23 (23)

[UCl^ + 3H

2

O + HBF^ + 162.5HC1 + 2H2B02]aq + 2.5H20(liq) - [UCl^ + 5 . 5H

2

O + HBF^ + 162.5HC1 + 2H3B02]aq; AH^^ (24)

UF^-2.5H20(c) + [166.5HC1 + 3H2BO^]aq “* [UCl^ + 5 . 5 H

2

O

+ HBF^ + 162.5HC1 + 2H

2

BO

2

] ; AH^^ (25)

then for:

UF^(c) + 2.5H20(liq) - UF^-2.5H20(c) ; AH^^

^26 = ^^23 + ^24 - ^25

At 323 K: AH

AH

AH

23

24

25

-2.74±0.02 kcal

-2,629±0.001

+5.13±0,01

(26)

AH_, = -10.50 ± .025 kcal

26

-34-
= -22 cal/moL»K) results in

At 298 K: ^^23 = -4.73±0.05 kcal

= -2. 405 ±0.003

24

^25 = +3.21±0.03

zlH., = -10.34±.06

26

Correcting 4^

2 ^ ftoni 323 K to 293 K (4fi

AH

25 = -9.94 kcal/mol, A weighted average for then, is

- 10 . 2 ±0.2 kcal/mol.

From an inspection of the values calculated from AHf*^ UF^(c) and

zlHf* UF, •2.5H_0(c) from the Khanaev measurements and the Maltsev et al.

4 2

measurements, one can calculate:

= -633.2 + 2.5x68.315 + 452.3 = -10.11 kcal

26

which indicates a consistent cycle.

Popov et al. (1957) reported Che of UF^(c) at 293 K

to be -8.21±0.02 kcal/mol UF^. The samples of UF^ were obtained by

dehydrating the crystalline UF^* 2 . 5 H20 (c) in .vacuo, gradually heating

to 498 K.

In addition to this Popov and Gagarinskii (1957) reported from

tensimetric data

UF^-2.5H20(c) ^UF^(c) + 2.5H20(g) (293-318 K) ; (27)

log P , = 7,154 - 2,920/T atm '

so that for the process given:

M = 13.36x2.5 = 33.4 kcal

^ = 3.60x2.5 = 9.0 kcal

^ = 32.7x2.5 = 81.75 cal

This would result in = -33.4 + 2.5x10.52 =-7.1 kcal/mol of

26

UF^-2.5H20(c)

.

zlH H^0(lf = 10.52 kcal/mol at 298 K.

vap 2 - 35 -
This is clearly in error. In addition = 32,7 cal/mol H^O *k

is low; one would expect = 35-36 cal/mol H20*K. There is a

possibility of a lower hydrate forming.

If it is assumed that the is reasonable but that the slope

is in error, then = -(9.0 + 35x2, 5x. 29815) + 2.5x10.52 =

-8.7 kcal/mol UF^» 2. 5H20(c) , in somewhat better agreement with the

Popov et al (1957)

40

We have selected ZlH hydration from the results of Gagarins kii

and Khanaev (1967), = -10.2±0.2 kcal/mol, since these direct solution

measurements are on well characterized samples. This results in AHf”

IJF^*2.5H20(c) = -636.5±0.2 kcal/mol (relative to AHf° UF^(c)).

The previously mentioned = -633.2 kcal/mol from Pbltsev et al.

may suffer from incomplete precipitation to UF^»2.5H20(c) ; more probable

is an error in the AH , involving HF(aq); this would cancel in the

mxx

derived ZlH. , (obtained using the AHf UF, (c) from Khanaev hydration 4

(1968)).

The estimated S® UF^»2,5H20(c) = 60±2.0 cal/mol*K is obtained from

the measured S® UF^(c) plus an estimate for the contribution of the

2.5H20 's (2.5 X 9.5).

The calculated and accepted Z33f® = -580.4 kcal/mol is obtained

from the selected ZlHf® and the z2Sf® = -188.2±2.0 cal/mol*K.

The Z^f® obtained from the aqueous K's (see Section C) are in

poor agreement.

-36 -
E. Lower Hydrates of UF^ and the Various Metastable Forms of UF^

Gagarinskii and Khanaev (1967) have also prepared and measured

the AH soln. of various lower hydrates and metastable forms. These

lower hydrates and their dehydration products are ill-defined. The

values derived for the AH, . UF (c, monoclinic) to the hydration 4

lower hydrates are inconsistent^ i.e.^

UF^(c) + l.SH^OClj -* UF^-l.SH^OCc) ; AH = -5.1 kcal/mol

but

UF^(c) + l.aSH^OClj ^ UF^-1.33H20 (c)

;

AH = -7.44 kcal/mol.

The metastable forms of UF^(c) are formed from the lower hydrates

and are irreversibly converted to the monoclinic. Again^ they are

not well defined.

An earlier paper by Gagarinskii and Mashirev (1959) also reports

on the lower hydrates.

-37 -
Till)

le I

Obtained

from

Markin

et

al

(1967)

4UT„(c)

3111’,

(c) +

D(c),

AH“ 61. 86lA. 0

kca LIU^
UF^(g)

Hildenbrand (1976) has calculated the thermal functions for

UF^(g) using the estimates of Tumanov (1968) where a tetrahedral

structure was assumed and the vibrational frequencies were estimated

on that basis. These are the assumed parameters;

= 555 cm"^ (1) = 147 " (2) = 566 " (3)

V, = 177 ” (3)

T^ models tetrahedron

cr = 12

Singlet ground state

r U-F = 2.08A

However, an analysis of the vapor pressure of UF^(c) and (1) indicates

a serious discrepancy in calculated ill's of sublimation from 2nd and

3rd. law analyses (which also shows a trend with temperature). The

trend can be removed and agreement with the 2nd law values can be

made by increasing S° and -(G-H^gg) /T by 8.2 cal/mol*K (see Table 2

AH®subi

^

4

^^) (g))» One can account for about 3.7 cal/mol*K of this

difference if UF^(g) has the less symmetrical structure; it had

been assumed that the rest was due to a large electronic partition

function. However, Gurvich et al. (1977) have indicated that this

would not resolve the remaining 4 cal/mol*K discrepancy. For convenience the thermal functions are tabulated but with S° and -(Gi-H

2 gg)/T

increased by 8.2 ca 1/mol »K.

-39-
Table 2 shows the available vapor pressure measurements on

UF^(c) and (1)^ the

^and at the mean temperature, the 2nd

law calculated from the A(H-H_q_) for the gas and the conzyo zyo T

densed phases, and the increment in S° needed to resolve the discrepancy between the 2nd law and 3rd law (calculated

zyo zyo

using S® at 298 K = 80,18 cal/mol«K from the unadjusted thermal

functions for Uf^(g)). From these measurements the selections

at 298 K are :

AH® = 75,4±0,5 kcal/mol

subl

zlS® = 80.2 - 36.25 + 8.2 = 52.1±1.5 cal/mol-K

subl

leading to AHf® UF^(g) = -380.1±1.2 kcal/mol. The ASf® = -373.9±2.0

kcal/mol is calculated from the AHf® and the tabulated S® = 88.4

' cal/mol*K,
K

Oo CN LO a^ cn

X ^ • • • . • • 00 00 cO o cn

COyX

CN in , 0 m cn 00 O ft

lE-ie • • • • • • CN cn m CN cn o

<j CO <r

COy3

00 CO

O

ON'—' s cn 00 cn o^

CN • • . • • • 5 <r in cO 'O’ in Q U CO vO vO vO vO cO N3 cn y 3 ft

00 cO o m £ r*^ CN m in CN • • • . • • 5

"3 <r in vO o^ ^ 3 CO pN. CN y

CN

ca H

00 o o o o in

O cn m o ^4 00 in <r o o CN 1^4 o

tx* lE-t f-t 1^ ft 3T ^4o CN in CN

/f> s CO 00 <r o CO cn y • . • • • • 1 ^ ^ in C3^ o in <r ^ CO cn n- CO cO 3 y 33

CO

T3O T3 (U 3 3

CO O*

•T^

U 3

•> fi . O O

(U CO a 3 3

COO CM • 3 y o y ft

a. 01

01 3

CM

COu

y 3

•r^ a. oe u y Cll CO y CO CM ^4 -u y 3 y 5

CO <u CO !.t CO

y

b CO ft CM

Ml -M iS

O i3

U 3 CM 00 • cn S u S CO U in

00

H

y

X

00 oc CO o cn

m o II IM c o <r r>- r>. 3 o

ca in o CN CN CN 3

OS rM ^4 ^4 pM ft

O cn 1 1 \ 1 I 1 •rt y • <r o o o o o CO 3

O. ct CNm CN cn 3 e cn o o^ ft CO (Ml 03 y ^4 ?» H <] 0 . .. cn

a. > OcS ft s o f > 3 u as ft

CO

y as o o 3 3 cO y in O. cn y r>. 3 CJ\ o O » CTt o T-^ xf. f M f 'w' o y y 5 'O' > 3 w f4 • o^ 3 o y o » 3 3 -M ft 3 Q 3 M C 3 3 Cf' 3 3 3 3 3 y 3 3

00

C c M

M 3 > ^ 3

•M

y

O 3 3 -I

y

O

O CO

y

cO 3 cO CO 3 cO 3 JO

y o^

y 3 Oc > CO ft oc cr

y 00 ft 3 CO ft o 3 3 ft o > 3 ft ft 'f' Ou 3 3 'f' tJ E-t 3 cn ft 0 O M 3 3 < Ot 3 O i:

41

(g)

to

obtain

2nd

and

3rd

law

agreement.
UF (c) (n = 4.25, 4.5, 5) n

' ^

Katz and Rabinowitch (1951) and Agron (1958) reported on the

properties of UF^ 25^^^^ ^4 5^^^^ ^^'^(^,0^) and UF^(c,3). Although

the or-form of UF_ is the stable form above 398 K (under 1.76 nnn UF,(g)),

j o the '^-form has not been converted to the 8-form at lower temperature.

To affect complete conversion of the 3-form to the a-form requires

12 hrs heating at 458 K» Both a and 3 forms have tetragonal crystal

symmetry.

Brickwedde (1951) measured the of UF^(c,a) which contained

only 83 weight %

^

5

* reported S®UF^ at 298 K was 45±3 cal/mol*K.

This value lies between the values for UF^(c,a) and UF^(c,?) which

are derived from the disproportionation reactions of Agron (1958)

on UF (a), UF (3), UF (c) and UF ,.(<=). The values for S® UF

and S° UF^ estimated by Rand and Kubaschewski (1963).

Rand and Kubaschewski (1963) evaluated the disproportionation

reactions of Agron (1958) modifying the experimental data (382-621 K)

by using their estimated S“'s for UF^ ^(c) and UF^ 25^^^ ^

= -11 cal/mol*K . The following relationships are accepted:

8UF^^25<‘=> ^ (1)

= 33,400 - 116. 5T + 25. 3T log T cal/mol UF^(g)

(502-621 K)

at 298 K = 30.1 kcal/mol UF, (g)

0

AG** = 17.4 kcal/mol UF^ (g)

o

45** = 42.7 cal/mol-K UF, (g)

o

A slightly better for the disproportionations, -5 cal/mol*K,

would result in negligible differences in the 4Hf®'s.

-42-
(2)

= 31,100 - 114. 4T + 25.3 T log T cal/mol UF^(g)

(533 - 590 K)

at 298 K = 27.8 kcal/mol UF, (g)

D AG® = 15.7 kcal/mol UF, (g)

o 4S® =40.7 cal/mol'K UF, (g)

o For the disproportionations of cc-UF^Cc) and P-UF^Cc), however, Rand

and Kubaschewskii did not distinguish between the two forms, so that the

original measurements of Agron then were refit (with a 4Cp = -11

cal/mol*K) to obtain:

3UF^(c,a) - 2UF^^^(c) + UF^(g) (3)

AG® = 18,300 - 99. 8T + 25. 3T log T cal/mol

(382 - 469 K)

at 298 K AH® = 15.0 kcal/mol UF^ (g)

b

AG® = 7.3 kcal/mol UF^(g)

AS® = 26.1 ca 1/mol • K UF,(g)

o 3UFg(c,3) ^ 2UF^.5 (c) + UF^(g) (4)

z2G® = 24,000 - 114. OT + 25. 3T log T cal/mol

- (384 - 423 K)

at 298 K AH® = 20.7 kcal/mol UF, (g)

0

4G® = 8.7 kcal/mol UF^ (g) = 40.4 cal/mol-K UF, (g)

o These equations result in the tabulated values (Section III).

Then, from (3) and (4) for:

UF^(c,3) -* UF^(c,CL)

at 298 K AH = 1.9 kcal/mol

AG = 0.5 kcal/mol

^ = 4,8 cal/mol*K

(5)

-43-
Krohn et al. (1976) estimated the vibrational assignment for

the basis of the pyramidal structure (O’ =4). Hildenbrand (1976) used these with the U-F bond length = 1.995 K and a

doublet electronic ground state to calculate the thermal functions

for UF^(g), These functions are accepted as modified below.

Hildenbrand (1976) using high temperature mass spectrometry

studied the gaseous equilibrium (range 1012-1158 K) :

Ag(g) + UF^(g) - AgF(g) + UF^(g) ( 1 )

131.1 cal/mol*K, at 1100 K, ~ 3 cal/mol*K higher than the calculated S® = 128

cal/mol*K, The tabulated S and -(G-H-qq)/T have been increased by

these 3 cal/mol*K. These are to be considered approximate functions.

The 2nd law AH = 15.6±0.6 kcal/mol corrects to = 16.9±1.5

kcal/mol. Using;

AgF(g) - Ag(g) + F(g); AH®2 = 84.8±4.0 kcal/mol (2)

from Clements and Barrow (1968), and Hildenbrand ' s determination:

UF^(c) -»UF^(g); AH®2 = 76.2 kcal/mol (3)

the following is obtained:

UF^(g) - UF^(c) + F(g); AH®^ = 25.5±4.5 kcal/mol (4)

The resultant AHf® UF^(g) = -462. ±5.0 kcal/mol.

From electron impact threshold measurements Hildenbrand also

reports :

UFg(g) UF^(g) + F(g) ( 6 )

-44-
D°q = 2,955 e.v. or 68.0±0.2 kcal/mol

so that ~ 69 kcal/mol and ZlHf* UF_ (g) = -463.3±3.3 kcal/mol.

Z9o j

Wolf et al. (1965) determined the vapor pressure of the form

of UF^(c) from 555 K to the experimentally determined melting pointy,

621° K as well as that of the liquid to 685 °K by measurements of the

transpiration rates, using UF^ as the carrier gas to prevent disproportionation of the UF^. The following vapor pressure equations

were derived:

UF^Cc^a)

;

log P(atm) = -(8001±664)/T + (11. 113±1. 119) , (555-621 K)

UF3(liq):

log P(atm) = -(5388±803)/T + (6 . 938±1.236) , (621-685 K)

The calculated mean - and 4S° , t at 580 K = 36. 6 ±3,0 subl subl

kcal/mol and 50.8±4,5 cal/mol*K. Correcting to 298 using an estimated = -7 cal/mol*K results in = 38.6±3.2 and 4S° = 55.6±5.0,

P

leading to AHf° UF^ (g) = -456±3.6 kcal/mol and S °= 104±5.4 cal/mol*K.

The entropy appears particularly high. One would not expect S° UF

3 (g),

even including an electronic contribution, to be greater than 94

cal/mol*K. In addition, the calculated 4S_ . = 19 cal/mol*K

^ fusion

appears high; one would expect . = 12-14 cal/mol*K. fusion

A third law analysis of the vapor pressure measurements on the

crystal, however, results in = 32.6 kcal/mol. At 600 K,

subl

-A(Q-H

2

gg)/T = 44,0 cal/mol’K, based on the adjusted -(G-H

2

gg)/T for

UF

3 (g), S° UF^(c,<^) = 47,7 and an estimate for-(G-H

2

gg/T)gQQ ^

for

UF (c,A). This , results in 4Hf° ~ -462 kcal/mol for UF (g)

5 subl j

in agreement with the Hildenbrand measurements.

The Gibbs energy of formation has been calculated from AHf° and 4Sf°

-45-
UFg(c)

The S° , Cp* and values tabulated by Brickwedde et al. (1948)

from their G measurements on the crystal and liquid (14-370 K) are

accepted. These measurements also define the triple point (337.20 K)

and . = 4.588±0,045 kcal/mol. fusion

In addition, there are measurements by Llewellyn (1953) in the

range 213-373 K.

The UF- (c) has recently been redetermined by Johnson (1977)

o by direct fluorination of U(c) to be -525.13±0.44 kcal/mol. Although

still to be considered preliminary, this new determination confirms the

value = -525.1±Q.5 kcal/mol recommended by Parker (1976^^from the key

network analyses and the U0^(c,7) - UF^(c) - HF(aq) - U02F2(c)

cycles. These values supercede the value -522.6±0.4 kcal/mol that

/ had been considered to be definitive.

The analysis which indicated that the value -522.6 kcal/mol for

UF- (c) was questionable and that a direct redetermination was needed

b

to substantiate the indirect value, -525.1 kcal/mol, is presented

historically in order to indicate how values considered reliable

become suspect through a detailed analysis of the network. The

analysis includes much of the preliminary evaluation of AHf® U02F

2

(c)

and /lBf° '[]0^F^(aq, in HF) which brought the problem into focus.

A. The Initially Accepted AEIf® UF-(c) and UF (g)

b b

There were two determinations of the enthalpy of fluorination

of U(c) to UF^(c) and UF^ (g) that were in good agreement with one

another and were considered definitive for ZlHf° UFg(c) and /2JIf“ UF^ (g) .

-46 -
The first of these is by Settle et al. (1963) using bomb

calorimetry

U(c) + 3F^(g) -»UF,(c); zm, = -522.64±0.43 kcal/mol (1)

Hayman (1967) in a flow calorimeter obtained:

U(c) + 3F2(g) ^UFg(g); AH^ = -510.47±0.64 kcal/mol (2)

From the analysis of the vapor pressure measurements (see UF^(g)):

UF^(c) -UFg(g); = 11.85±0.10 kcal/mol (3)

so that AH® = t522.32±0,65 kcal/mol from Hayman 's data.

These measurements led to the accepted AHf® UF (c) = -522.6±0.4

o and AHf® UF^(g) = -522.6 + 11.85 = -510,75±0.4 kcal/mol.

B. The Inconsistency in Paths and How it Affected AHf® UO.F (c) and

AHf® U02F2(in HF)

The analysis of the data on the AHf® U02F2(c) involves the AHf

of the aqueous species. The measurements on the reaction of U0^(c,7)

in HF(aq) to form a species for convenience called U02F2(aq, in HF)

will be cited first. The measured reactions are of the form:

1102 (0 , 7 ) + 2HF(aq, in excess) "* U02F2(aq, in HF) + H20(aq, in HF)

Since all the measurements on the AH of U0_(c,7) were made in

excess HF of varying concentration and in varying UO^/HF ratios we

have corrected for the dilution of the excess HF in the final solution, either by using the appropriate cp^ values for the initial and

final concentrations, or by using the differentials L^'and if

there is no appreciable concentration change for HF(aq).

* All values listed by Parker et al. (1976) for the integral and

differential AHf ' s HF(nH20) should be made more negative by 0.07

kcal/mol. The value listed for AHf HF(aq, std. state) = -80.15

kcal/mol is correct.

-47-

i
The

Reaction

of

UO

(c,7)

in

HF(aq)

<u> LO

o* it

eg -u

m • • • » w a- • O o p o CNhH <u o +1 +1 fl +1

C£« !-i

u +1 o\ CM 00

CvJO CO 00

m M3

O CO

u • • • • o s <u 00 on CM cn cn cn

T3

O o O o o a ^

-it «ct

(U T3 wu X-N o y—s o

<u (U CO o • o CM £ u CM X CM X •H 0 s CO X M3 CN X 4J o o ni • CU (U <4-1 • o CM + CM

•H -U M CO CM

T3

•U ? + CM+ U -u fa (0 fa fa

00 ca S cn

X (U

<U

O • vO • V-/

u

O P cr in CO

cO 1 CM r—4 u

cn• -o

4J

N in CO o CM o o CM • o •s • _r^ X in X CM CO* c X mT f—4 X CO

CO M3 1-M o 1-1 o O S M3 CM + CM

•U CO CM CM + fa +

fa CM

1-1 + fa

X CM

C fa 8 S X

X o M S v—' cn X 00 • cn

M3 t

fi- CM o'a o it 00 CM O cn

CM

O un cn T3 cn cn o 1—4 CM 3 ^ r-4 • • • • • CO

CO o\ 00 a^ o^ CT\

CJ r—4

t-M 1—4 t-4 1 1 1 1 1 0) u

CO CO a^ 00 u o o o CO o • • • • 1/m +1 -H -H +1 • M3 1-M 00 cn TO to on cn M3 U

CO • • • • CO o in in in CM CM CM CM •s 1 1 1 1 o*

CO \ M3 1—4 1 r— fa

(ON cn

r—4

T—4 s V-/ M3 CO

CO pi. ;>

m (U cn o cu M3 ••—I f—4

-4-1

X

C3^

u s—/

1-4

u

r—4

1—4 1—4 v—/

(U s o s 3 o cu a • <5

CO fa

u

i-H

a 1-4 <u

CO

X

T3 o TO TO

u

C *n cu 4« <0 CO CO u

TO T3 CO

<u fM

C •H 1-4

u 1-4 cO 1-4 3

CO c X

O

1^ CO 3 (U CO 1—4 3 > M-4 P» CO O cO 'X3 iS 0> O T3 )-i X TO iM O — •r4 a > o o > it

-48
Table 3 shows the measured AH and the corrected

m

zlH for the reaction,

r

’

U0^(c,7) + 2HF(aq, std. state) "* H^OCl) + U02F

2 (aq, in HF),

the initial composition of the HF solution, the descriptive state

for the in the final solution and the calculated AHf

for that state.

As can be seen^ the results are in good agreement

with one another, although the Z^Jif calculated from Vidavskii et al.

(1965) is 0.75 to 1.0 kcal/mol more positive than the others. The

solution is also more concentrated with respect to the UO^F^.

Similarly, information on AHf UO^F^Caq, in HF) from the reaction

of UF^(c) in H^OCl) or in a solution of HF(aq) may be obtained. The

measured reactions are of the form:

UF- (c) + 2H_0(1) or (in HF soln.) U0_F_(aq, in HF soln.)

o Z Z Z + 4HF (soln.

)

The reactions have been corrected where necessary to 298 K and to the

form

:

UFg(c) + 2H20 ( 1 ) — U02F

2

(aq, in HF) + 4HF(aa, std. state)

by correcting as for the reaction of U02 (c, 7 ) in a solution of HF.

Table 4

The Reaction of UF^ (c) in H^O

6 Z

kcal/mol

AH^ Initial

Soln.

Descriptive

state of

UO

2

F

2

formed

AHf ^^2^2

descriptive

state

Popov et al.

(1957)

-49.73'"

±0,30

-61.874 I 6 O 2 H

2

O 4(HF+400H20) -400.50

±0.8

O' Hare and

Johnson(1976)

-49.918

±0.07

-62.35 21.33 (HF+

I47 . 8 OH

2

O)

25.33(HF +

I 24 . 4 H

2 O)

-400.98

±0.75

Corrected from AHm= -50, 22±0.30 at 305 K assuming AC^ = - 70 cal/mol * K

-49-
The two values are in good agreement with one another, and show

the same effect as noted earlier, but the values calculated from the

two tables differ by ~ 2.5 kcal/mol.

UO F (aq, in HF) Descriptive Difference

from UF^ from 7U0_(c)

o 3

State

in Paths

OTg-UOs

-400.50 -402.79 4(HF+400H20) 2.29

-400.98 -403.62 25.33(HF+124.4H20) 2.64

The following measurements have been made on the ^soln. 1102^2^“=^

in HF and in H^OCl) corrected to 298 where necessary.

^soln

kcal/mol

Solution

Suponitskii et al. (1971) -7.12 57 IH

2

O

Popov et al. (1957) -7.89±0.2* 4(HF4400H20)

Cordfunke and Ouweltjes (1976) -8.10±0.05 47.6(HF+264H^O)

These values are also in agreement with one another. Obviously

the values calculated for zlHf° UO^F^Cc) from the two available paths

will reflect the 2.5 kcal discrepancy, i.e. UO^F^Cc) = -402.79

+ 7.89 = -394.90±0.50 kcal/mol and -403.54 + 8.10 = -395.44±0.40

kcal/mol from the U0^(c,7) path and the appropriate

UO^F^Cc), and UO^F^Cc) = -400.50 + 7.89 = -392.61±0.85 kcal/mol

from the UF^(c) path and the appropriate AH , U0_F„(c).

o soin 2 2

At this point it is essential to explore the possible causes of

the difference.

1, There is a major error in either the experimental measurements

involving the reaction of UFg(c) or those of UO^Cc,/). This is

HeCorrected from = -8.10 at 305 K assuming ACp = -30 cal/mol*K.

- 50-
highly unlikely since the agreement within each set

(differing in their origin) is ^ood.

2 . The UO^Cc,/) used for the measurements is not the same form as that for which = - 292 . 5 ±0 . 3 , However_, O'Hare

confirmed that the sample of UO^Cc^/) used by O' Hare and

Johnson ( 1976 ) was the same as that used by Fitzgibbon

et al. ( 1967 ) in their determination of Z^® 1102(0,7)

from solution calorimetry involving 1102(0,7), U02(c) and

U20g(c) since they confirmed these solution measurements.

3 . The value used for 1102(0,7) = - 292 . 5 ±0.3 kcal/mole is too

negative by ~2,5 kcal/mol. This does not seem possible

unless a corresponding change is made in 1102(0) and an

even larger change is made in U^O (c) . If UO. (c) and

1102(0,7) were changed significantly the good agreement

obtained for the z2Hf® values for ( 1 ) UCl^(c) [only partly

dependent on zlHf 1102(0)], (see section on UCl^(c));

( 2 ) 1102012(0) [partly dependent on UCl^(c) and partly dependent on U0-(c, 7)l, (see seotion on UO_Cl„(c)); and

2+

^ ^

( 3 ) UO2 (aq, std. state) [partly dependent on U 02Cl

2

(o)

and partly dependent on UO2 (^02)2 ’ 6H20(c) which is completely

dependent on 1102(0,7)], (see Fuger and Oetting ( 1976 ))

would disappear.

4 . The value used for UF,(c) should be made more negative than

o - 522.6 by this 2.5 kcal/mol. As shown in A on the ziHf®'s for

UF^ (c) and UF^ (g), there is confirmation for - 522.6 kcal/mol.

o b

However, a more negative value for UF^bo(c) and UF^ (g) would

lower the value for zlH® U(c) derived from UN(c) and

US(c) and would bring them into better agreement with the

accepted ^^subl ~ kcal/mol (Oetting et al. ( 1976 )).

- 51-
5 . Neither the UO^Cc,/) nor the UF^Cc) are seriously

in error, but the auxiliary values for AEIf HF(aq) are. Since the UF^ (c)-U02 (c,7) relation involves six HF(aq), this

shift of 0,4 kca 1/mol HF would account for the discrepancy

in paths. However, the AHf HF(aq) values listed by Parker

et al, (1976) from the Johnson et al. (1973) determinations

are consistent with the CODATA (1975) AHf® F (aq, std, state) = -80,15±0«16 kcal/mol which was fixed from various cycles,

five of which were in agreement (within ±0.1 kcal),

6. None of the above, but an erroneous assumption is being made.

The first three possibilities are at present rejected. It is more

difficult to assess the remaining three; however, these 3 possibilities

all involve the reaction of UFg(c), either because the accepted AHf **

UFg (c) is erroneous, or because of the greater dependency on the

value for AHf ® F (aq, std. state), or because the assumption on final

solutions is not warranted.

The better path, then, to UO^F^ (c) and in HF) is from

the relationship to UO^(c,7), which results in -394.90±0.5 and

-395.44±0,4 kcal/mol for AHf® UO_F^(c) rather than from the UF (c)

Z Z o path which results in -392. 6±0, 85 kcal/mol.

C. The decision to change the AHf®'s UF^ (c) and UF^ (g)

As is indicated in B, either the AHf® UF^ is suspect or to a lesser extent

4Hf HF(aq, std. state). However, if one goes through the uranium key

network the more positive AHf ® HF(aq, std. state) (= -79.75 kcal/mol)

required here would not resolve other inconsistencies. (See UF^(c)

and UF^(c)) where a more negative AHf ® HF(aq, std. state) would be

- 52 -
required to resolve the inconsistencies. In addition, on a practical

note, many more new cycles would have to be considered before one

could say with certainty that the GODATA. HF(aq, std. state) is

too negative by 0.4 kcal/mol (the UF^ (c) -UO^ (c, 7) cycle from the

combination of Popov et al. (1957) and Vidavskii et al. (1965) had

been considered by the CODATA task force).

The most likely values then, by the process of elimination, are

£^®*s UF-(c) and UF^ (g) . In order to substantiate this the following

0 o was done :

It was assumed that the f luorinations of U(c) to UF, (c) and (g)

o were in error but that the zlH's of fluorination of UN(c), US(c) to

UF^ were correct.

6

The high temperature decomposition reactions of US(c) and UN(c)

were considered and the most reasonable reactions used. The imf**

+2

U(g) = 127

^ kcal/mol from Getting et al.(1976) was used in each

reaction to obtain values for ZlHf*’ of US(c) and UN(c) which were then

used in the fluorination reactions to obtain values for 4Hf'’ UF^ (c)

0

and ^Hf* UF^ (g) .

The reactions are

:

UN(c) ^U(g) + l/2N2(g); = 200.7±1.0 kcaiymol (1)

(avg. of 2nd and 3rd law)

from Hoenig (1971), so that UN(c) = -73.7 ±1.8 kcal/mol and 4Hf°

-53 -
UNq

gg^y

= -73.6±1.8 kcal/mol.

TOq. 9957

(c) + BF^Cg) - UFg(c) + 0.9957/2N2(g);

= -451.86±0.3 kcal/tnol (2)

from O'Hare et al, (1967) (corrected for composition)^ so that Zlflf®

UFg(c) = -525.5±1.9 kcal/mol.

US(c) -* U(g) + S(g); = 269.6±2.0 kcal/mol from Nater (1969) (3)

so that zmf° US(c) = -76.4±2.5 and US^

Oll^'^^ ^ -76.8±2.5 kcal/mol.

USi^Oii(c) 6.033F2(g) -UFg(g) + l.OllSFg(g);

= -732.59±2.0 kcal/mol from O'Hare et al. (1967) (4)

with zlHf“ UF.(g) = -514.4±3.2 kcal/mol and AHT* UF, (c) = -526.3±3.2

o o kcal/mol. These two paths lend support to more negative values for

AHf ° UF-(c) and AHf° UF, (g) . These are obtained from the UF^-UO„

cycles, so that AHf° UF^(c) = -522,6 - 2.29 = -524.89±0.6 kcal/mol

from the Popov-Vidavskii cycle and -522.6 - 2.64 = -525. 24 ±0.5 kcal/mol

from the O' Hare and Johnson cycle, resulting in a selected AHf® UFg(c)

= -525.1±0.5 kcal/mol.

D. Confirmation From the Direct Enthalpy of Fluorination

Obviously confirmation for this new value was needed. The

new measurements of Johnson (1977) on the direct fluorination to

* UF, (c) cited previously = -525.13±0.44 kcal/mol lends support to

o the indirect value

A o T A recalculation by Johnson (1979) using the values for the AHf s for UF3 (c) and UF^(c) and the UFg(c) selected in this evaluation results in AHf'’UF^(c) = -525.28±0.44 kcal/mol.

AA

One should also note that the new direct determination lends

support to the CODATA AHf° HF(aq, std. state) and AHf° F (aq, std.

state) = -80.15 kcal/mol.

-54-
We can now with certainty assign AHf® UFg(c) = -525.1±0.4 kcal/mol

and AHf® UF, (g) = -525.1 + 11.85 = -513.25±0.4 kcal/mol.

o There is at present no explanation for the two earlier direct

fluorination reactions, in agreement with one another, being in error.

The Gibbs energy of formation has been calculated from the z^If°

and the z^f®.

- 55 -
The , , subl

= 11.85±0.10 kcal/mol and the =1.14

subl

kcal/mol result in Z^® , , = 35.95±0,3 cal/mol*K which leads to S®

subl.

UFg(g) = 90.3±0.45 cal/mol-K.

The thermal functions for the ideal gas may also be calculated

from spectroscopic data. The frequency assignment by Glaasen (1959)

measurements of Gaunt (1953)^ Glaasen et al. (1956)^ Burke et al. (1952)^

and Bigeleisen et al. (1948)^ for the octahedron (symmetry = 24), with

r-TJ-F 1.999A from Seip (1965) results in S® = 90.23 cal/mol*K in excellent

agreement with the experimentally derived S®.

The more recent assignment of IfcDowell et al. (1974) 672(1),

cal/mol*K.

A table of thermal functions based on Glaasen et al.'s (1956)

assignment is presented here (Section V),

The ZlHf® of UF^ (g) is derived from the selected ZIH® ,, =

6 subl

11. 85 ±0.10 kcal/mol for:

This is obtained from both the calorimetric determinations of AH®

(277-363 K) of Masi (1949) and those calculated from

the various vapor pressure measurements, all in excellent agreement.

This results in zlHf® UF^ (g) = -513.25±0.4 kcal/mol.

667(1), 535(2), 623(3), 181(3), 202(3), 140(3) cm"^ based on the

540(2), 634(3), 186(3), 200(3), 143(3) cm"“ results in S® = 89.93

UF^(c) - UF^(g) ( 1 )

vap

-56 -
As indicated under UF^ (c) , Hayman (1967) obtained = -510.5±0.7

kcal/mol for:

U(c) + SF^Cg) UF^(g) (2)

This has been shown to be in error.

The ZiG®

^ = 1.14 kcal/mol^ result's in UF^(g) = -493.26

kcal/mol.

The various vapor pressure measurements on the solid and liquid

are in good agreement and can be represented by the equations of

Oliver et al.(1953):

uFeCc)

log P(atm) = 3.50282 + 0.0075377t - 942.76/(t+183.416)

(273-337 K)

UFed)

log P(atm) = 4.11383 - 1126.288 /(t+221. 96 3)

(337-390 K)

UFed)

log P(atm) = 4.80988 - 1683 . 165/(t+302 . 148)

(390-500 K)

The triple point calculated here is 337.20 K with P(iran) = 1139.6.

This calculated triple point is in excellent agreement with that

directly determined by Brickwedde et al. (1948) as 337.202 K with

^fusion “ ^*588 kcal/mol, = 13.61 cal/mol*K,

For UF^d) - UFg(g) :

at T = 337.202 K ^"“vap = kcal/mol

The sublimation temperature (1 atm.) is at 329.69 K.
For convenience Masi's (1949) smoothed zm** for the condensed phases

vap

(273-370 K) are tabulated. These values are within 50 cal/mol of the

recommended values given here.

TABLE 5 Enthalpies of Vaporization of UF^

T, K

crysta

1

ca 1/mole

273.15 12,023

280 11,988

290 11,929

298.15 11,872

300 11,858

310 11,772

320 11,666

330 11,537

337.20

liquid

11,429

337.20 6859

340 6817

350 6671

360 6533

370 6404

See Rand and Kubaschewski (1963) for citations to the individual

investigations for the vapor pressure measurements.

- 58 -
3.102 U-F-0 Compounds

UQF^Cc) and UOF^-H^OCc)

S® TJCF^Cc) has been estimated as 28.5±1.0 cal/mol*K based on

a comparison of the S®’s of UGl^(c), UF^(c), U0Cl

2

(c), UO^Cl^Cc)^

1102^

2

( 0), 1102 ( 0) and UO^Cc). The S® UOF^’H^OCe) is derived from the

estimated discuttad in the following section.

Vdovenko et al. (1967) isolated a black crystal hydrate which

corresponded to UQF

2

*H20 (c). In (1969) they measured its enthalpy

of reaction in HCl(aq) at 293 K. Their measured reactions are:

U0F

2

*H20 (c) + 254HCl(aq) - [UF^Cl^ + 252HC1 + 2H20 ]aq; = -5.42±0.04 kcal/mol (1)

[UCl^ + 250HCl]aq + 2HF(aq) - [UF^Cl^ + 252HCl]aq;

= -7.90±0.05 kcal/mol (2)

The concentrations used correspond to those used by Maltsev et al.

(1960) in their determination of 4flf® UF^« 2 . 5H20 (c) and use is made

of the Maltsev et al. reaction:

UCl^(c) + 250HCl(aq) ^ [UCl^ + 250HCl]aq; AH^ = -43.47±0.10 (3)

to complete the cycle. From these reactions the following summation

is obtained:

UCl^(c) + 2H^0(aq) + 2HF(407. soln.) UOF^-H 20 (c) + 4HC1(5N HCl) ; = -46.68±0.2 kcal/mol UCl^(c) (4)

From the individual reactions in the two papers involved it is

clear that the AHf of the HF 40% soln. HF* I. 667H

2 O) to be used is an

Corrected to 298 K, assuming ACp = -145 cal/mol‘K.

-59-
integral quantity = kcal/mol but the AHf's of HCl and H^O in

5N HCl solution , HC1(10H20)^ are differential quantities.

The resultant AHf® UOF^-H^OCc) = -430.7±0.8 kcal/mol.

In addition, the H

2

O vapor pressure (283 to 363 K) over UQF2*H20

was determined by Vdovenko et al. (1969). For:

U0F2-H20(c) U0F2(c) + H20(g) ; (5)

at 298 K they report a calculated AH^ = 11.1 kcal/mol H20(g)

= 2.82 kcal/mol H20 (g)

AS^ = 27.8 cal/mol-K

Use of the calculated AH^ = 11.1 kcal/mol results in AHf* U0F2(c)

= -361.8 kcal/mol. This would mean that for the reaction:

2UOF2(c) -* U02(c) + UF^(c) (6)

* AH > 8 kcal/mol UF, and AG > 9 kcal/mol UF, 4 4

This would indicate that the stability of UCF

2

(c) is greater

than that of U0Cl2(c) (AH and AG for the comparable chloride reaction

are 7.2 and 7.4 kcal/mol UGl^(c), respectively). This is questionable.

In addition the calculated AS and AH for reaction (5) are low. One

would expect AS^ to be ~ 35 e.u. If the measured pressure at 298 K

is used (-A3°^ = 2,82 kcal/mol) with the estimate AS^ = 35 cal/mol*K,

AH^ = 13.3 kcal/mol H20(g). This results in AHf° U0F2(c) = -359.6±1.5

* kcal and results in, for (6), AH 6 kcal/mol UF^ .

The Gibbs energies of formation have been calculated from the

AHf ** 's and ASf “ 's.

AH > 8 and AH 6 are used because of the uncertainty in the AHf°

UF^(c).

-60-
U02F2(c)

Wacker and Cheney (1947) measured the heat capacity from 13 K

to 418 K. The tabulated S°, C

” and are obtained from their

P 0

smoothed values. Kelley and King (1961) cite the above measurements

in their tabulation; however, they report S°

2gg

= 33.40 cal/mol’K.

This appears to be an error in transcription.

Cordfunke et al. (1978) determined the high-temperature enthalpy

from 375 to 811 K relative to 298 K. The results are expressed as:

^t“^298

" 29.53T + 0 . 5833xlO"^T^ + 5.193x10^t"^ - 10,598 cal/mol

(298-811 K)

This equation results in values for ^'^”^298 good agreement

(lower by v .03 kcal/mol at 400 K) with those tabulated by Wacker and

Chenev (1947) but leads to a calculated C

“ at 298.15 K vO.6 cal/mol *K

P lower than that obtained from the direct C

® determinations which

P

are, at present, preferred for the low-temperature properties.

Purity of the Wacker and Cheney sample, however, may be the

problem.

As indicated in the discussion of AHf° UF^(c), from

U0^(c,y) + 2HF(aq, std. state) U02F2(c) + H20(l)

AHf® U02F2 (c) = -394.90±0.50 kcal/mol from the combination of the

Popov et al. (1957) AH , UO.F.(c) in 4(HF + 400H.0) and the Vidavskii soin 22 2

et al. (1965) AHf° U02F2(aq) in the same final solution. Also AHf®

U0^F2 (c) = -395.44±0.35 from the Cordfunke and Ouweltjes (1976) values.

- 61 -
From the Popov et al. (1957) UF^-UO^F^ cycle AHf*’ = -403.0 +

7.89 = -395.11±0.6 kcal/mol.

The selected value of AHf° U0^F2(c) = -395.2±0.3 kcal/mol.

In addition to the measurements described above, Knacke et al.

(1969) reported a AH = 90.7 kcal/mol UFg(g) and AS° = 62.3 cnl/mol

UF^(g) in the range 1033 K to 1073 K for the decomposition;

3U02F2(c) -V 2/3U30g(c) + UFg(g) + l/302(g)

With a ACp = -3.7 cal/mol*K, AH° = 94.0 kcal/mol and AS° = 66 cal/mol*K.

This AH results in AHf° U02F2(c) = -392.3 kcal/mol and S® U02F3(c)

= 28.6 cal/mol*K, in sharp disagreement with the measured low temperature value. In addition, if one uses the measured low temperature

S°'s and the AG° at 1040 K, the resultant AH^^^g = 80.2 kcal/mol and

AHf^ = -387.7 kcal/mol, even further off. No weight can be given to

these measurements. There is a possibility that the products of the

decomposition are not as written.

The Gibbs energy of formation has been calculated from the

accepted AHf and the ASf ” .

i.e., -400.5 - 2.5 where -2.5 = -525.1 - 522.6.

-62 -
in HF)

The various values for the 4iHf® U02F

2

(aq^ in HF) cited in the

discussion under UF (c)

D from the 411° 's of reaction

f of UF. , 7UO- and 0^ 3

U02F

2

(c) in HF can now be consolidated. The major

be the concentration of the HF solution.

Final smoothed values are:

effect appears to

UO F (c) soln Z 2

4Hf UO^F^ (solution) Descriptive

kcal/mol kca 1 /mol State

-8.6 -403.8 aq. std. state

-7.12 -402.32 in 57 OH

2 O

-7.85 -403.05 in 4(HF + 4OOH

2

O)

- 8 . 20±.10 -403.40 in 47 (HF + 276 H

2

O)

-8.35 -403.55 in 25 (HF + I24 . 4 H

2

O)

-8.58

U02F2-nH20(c)

-403.78 in 17 (HF + 22 . 4H

2

O)

Suponitskii et al (1971) report the following zlH's soln. in STIH^O

UO^F^ (c)

uo^F^-i.en^oCc)

UO^F^OH^OCc)

U02F2*4H20(c)

zliH, kcal /mol

-7.12±.ll

-2.05±.07

-1.46±.06

-1.35±.02

Tsvetkov et al (1972) report from tensimetric measurements:

Reaction

zyo

kcal/mol H20 (g) kcal/mol H20 (g)

1 12.8±0.8 2.32±0,08

2 12.6±0.8 2.54±0.07

3 10.7±1.2 2.54±0.08

4 13.3±1.4 3.25±0.10

5 11.7±1.9 3. 25 ±0.12

6 13.3±2.2 3.93±0.13

7 11.3±2.4

-63-

3.93±0.15
for the following:

U02F2*4H20(c) - 1OO^F^OH^OCc) + H^OCg) (1)

UO^F^OH^OCc) U02F2*2. ISH^OCc) + 0.82H20(g) (2)

U02F2-2 .ISH^OCc) = UO^F^ •1.75H20(c) + 0.43H20(g) (3)

UO^F^-l .ySH^OCc) •1.18H20(c) + O.fyH^OCg) (4)

UO^F^-l .ISH^OCc) = tlOjF^ •0.85H20(c) + O.OOH^OCg) (5)

UO^F^-O .85H20(c) = TO^F^ .0.25H20(c) + 0.60H20(g) (6^

UO^F^-O .25H20(c) = UOjFj (c) + 0.25H20(g) (7)

Then for :

UO

2

F

2

•nH20(c) - UO^F^Cc) + tiH^Od)

we obtain:

n S°

kca 1/mol kcal/mol kcal/mol cal/mol*K

Suponitskii Tsvetkov et al

4 5.77 7.54 3.45 13.7

3 5.66 5.46 3.14 7.8

2 5.36 3.97 2.73 4.2

1.6 5.07

1.0 2.09 1.7 1.3

ZIH = 5.6±1.0 kcal/mol is accepted for :

UO

2

F 2*3H20(c ) - UO^F^Cc) + OH^Od)

leading to AHf® uo^;F2*3H20(c) = -605.8+1. 0 kcal/mol. No values are

given for the other:s becaus e of the disagreement and the imprecise

states of the hydrates.

-64 -
UOF(OH)(c) and UOF(OH) -O.SH^OCc)

Vdovenko et al. (1970a) prepared a black finely crystalline

precipitate (by the addition of alkali to an aqueous solution con3+

taining UF ) which has the composition UOF(OH)

0 . 5 H20 (c) . They also

measured the thermal properties at 293 K, i.e. in 5N HCl (1970b).

U0F( 0H)* 0 . 5 H20 (c) + 254HCl(aq) [UFCl^ + 251HC1 + 2 . 5 H20]aq; = -13.84±0.15 kcal (1)

and

[UCl^ + 250HCl]aq + HF(aq) ^ [UFCl^ + 251HCl]aq;

= -4 .

6

±0.1 kca 1/mol (2)

The concentrations appear to correspond^ as in their similar

solution measurements on UQF^ *

H

20 (c) , to those used by Maltsev et al.

(1960) in their determination of _lHf® UF^* 2

.

5 H20 (c) and use is made

of the Maltsev et al. reaction;

UCl, (c) + 250HCl(aq) [UCl, + 250HCl]aq;

= -43.47±0.1 kcal (3)

so that

UCl^(c) + HF(407o soln.) + 2 . 5 H20(aq) ^

UQF( 0H)* 0 . 5 H20 (c) + 4HCl(aq); (4)

= -34.8±0.6*kcal/mol

4

With the integral ilHf® HF(407o soln. , HF (I. 667 H

2

O)) = -76.42 kcal/mol

and the differential ±Jif 's for HCl and H^O in a 5N HCl solution, i.e.

HCKIOH^O) = -376.7±1.5 kcal/mol for the hemihydrate.

Corrected to 298 K,- assuminga ±C„ = -120 cal/mol*K.

p

- 65 -
Vdovenko et al. (1970b) also measured the vapor pressure (283-363 K)

and reported for:

U0F(0H)-0.5H20(c) UQF(0H)(c) + 0.5H20(g) ; (5)

= 10,4 kcal/mol H20(g) and = 25.6 cal/mol H20(g)*K

If one assumes (5) to be the process, one \</ould expect Z:S to be 35-36

cal/mol*K. If this ZlS is used and one assumes the is reasonable ^ 2.8 kcal/mol, a ~ 13.2 kcal/mol H20(g) is obtained, and ZlHf°

U0F(0H)(c) ~ -341±3.0 kcal/mol.

I

-66 -
U 02 ( 0H)F‘H20(c) and UO^ ( 0H)F- 2 H20 (c)

Tsvetkov et al. (1973) observed the existence^ in the UO^-HF-H^O

system, of two hydrates of a basic salt of variable composition;

they are UO^(OH) F •2H„0 and UO (OH) F„ where x = 0,3- 1.2.

The thermodynamic characteristics have been determined by DTA,

They report the following for the case when x = 1.

U02(0H)F.H20(c) - 1/2UC2F2(c) + l/2UO^ .H20(c) + E^O(g) (1)

log P(atm) = (7.60±0.84) - (3100±246)/T (303-408 K)

U02(0H)F-2H20(c) - UO^ (OH)F-H^O(c) + H20(g) (2)

log P(atm) = (6.23±0.52) - (2360±141)/T (303-379 K)

The calculated values at 298 K are

:

For reaction (1) = 14.2 kcal/mol H20(g)

= 3.82 kcal/mol H20(g)

= 34.8 cal/mol H20(g)

for reaction (2) = 10.8 kcal/mol H20 (g)

=2.3 kcal/mol H20(g)

= 28.5 cal/mol H20 (g)

The formation properties of UO^ (0H)F*H20(c) are obtained from

the values for reaction (1) since the formation properties of

U02F

2

(c) and U0^*H20 (c) are known.

For reaction (2), the Z!;iH and are unreasonable. The normal

zlS® , , - ~ 35-36 cal/mol H_0(g) . A more reasonable = 13.0

dehyd 2

kcal/mol results from using the measured 4G®2 = 1.11 kcal/mol at

T = 340 K and the estimated zlS®= 35 cal/mol.

zlHf® UO •H20 (c,'^) = -366.5±0.2 kcal/mol from Gordfunke and O'Hare

(1978) rs used. Parker (1976) recommends -366.6±0,3 kcal/mol.

-67 -
Then for the dehydration

U02(0H)F*2H20(c) - UO^ (0H)F.H20(c)

= 13^000 - 35T (cal/mol)

H20(g)

(298-380 K)

log P(atm) = 7.65 - 2,840/T.
3.103 U-Cl Compounds

UCl^Cc)

The 298 K values for S° and H-Hq are taken from Katz and

Rabinowitch ' s (1951) tabulation of the smoothed thermal functions

obtained from the low temperature specific heat measurements of

Ferguson and Prather (1944) (15-380 K) made on samples of unknown

purity. They are considered accurate to 0.57o. The measurements

extrapolated to 0 K assuming no abnormal behavior results in S° =

38.0 cal/mol*K; however^ as pointed out by MacWood (1958)^ the UCl^

curve appears unusual between 15-20 K, appearing to approach a minimum

at ~ 10 K. If this is so^ then a better value for S® UCl

2

(c) would

be between 38.0 and 40.75 cal/mol*K where R In 4 = 2.75 is the total

magnetic entropy present in UCl

2 (c). However^ until more definitive

measurements are made^ S“ = 38.0-0.3 cai/mol*K is accepted.

Ginnings and Gorruccini (1947) measured the heat content relative

to 273 K in the range 273-998 K using a higher purity sample (~ 99.8%)

than Ferguson and Prather. In the overlapping range, 273-373 K, the

derived Cp's are 0.5% lower.

Kelley (1960) tabulated the , ^^“^298^1 ^^~^298^T

the Ginnings and Gorruccini results which, with the low temperature S®,

results in the thermal functions to 1000 K. The results can be

expressed as:

^^298 K

" 20.98T + 3.72xl0"\^ - 1.16x10^t"^ - 6,197 cal/mol

(298-1,000 K)

-69-
Fontana (1947) and (1958) measured the following:

UCl^(c) + [1/2U02^'^ + 2H^ ] in 0.5m HCIO^ -

[3/2u'^'‘’ + 3Cl' + H^O] in 0.5m HCIO^; = -69.07±0.5 kcal/mol

( 1 ) With the values for the AHf ' s for (0.5m HCIO^) and

(0.5m HCIO^) from Fuger and Getting (1976) and the assumption that the

enthalpy of formation of the Cl in 0.5m HCIO^ is the same as in 0.5m

HCl, AHf” UCl2(c) = -207.3±1.3 kcal/mol is calculated.

However^ since Fontana made a whole series of calorimetric

4+ 2+ oxidation reduction measurements on 1101^(0)^ UCl^(c)^ and U UO2

(see Fuger and Getting (1976))^ another way of treating the above

measurements would be to rearrange his reactions and obtain the

^soln. UCl3(c):

UCl3(c) + 3Cl"] in 0.5m HCIO^; = -28.97±0.5 kcal/mol

( 2 )

Assuming z^Uf® (in 0.5m HCIO^) ~ AHf ® U^"^(aq, std. state)^ one

obtains AHf® 0013(0) = -206.5±1.3 kcal/mol.

Another approach, which minimizes the assumptions regarding

the value to be used for the formation of Cl in 0.5m HCIO, and ties

4

AHf® 11013(0) to AHf® UCl^(c) is through the rearrangement of Fontana's'

reactions to:

0013(0) + HCl(in 0.5m HCIO^) 001^ (c) + l/2H2(g);

AH3 = +3.01±0.6 kcal/mol (3)

Although, again the partial AHf of HCl is to be used, the solution

now involves only one mole of HCl, an introduction of only ~ 0.1 kca 1

uncertainty. The resultant AHf® 0013(c) = -207.0±0.9 kcal/mol.

- 70-
The experimental calorimetric results of Barkelew given by

MacWood (1958) result in:

UCl^Cc) + Cl^Cg) ^UCl^(c); = -38.1 kcal/mol (4)

and UCl^Cc) = -205.4±2.0 kcal/mol.

MacWood (1958) also tabulates the K’ s from Altman (1944) and

Gregory (1945) for the UCl^ hydrogen reduction equilibrium (673 to

823 K) . The third law AH*’ values for the reaction:

UCl^(c) + l/2H2(g) ^ UCl3(c) + HCl(g) (5)

are :

1 Altman = 14.36±0.1 kcal/mol UCl^(c)

Gregory = 14.38n0.3 kcal/mol UCl, (c)

The value obtained for AHf® 11013(0) from the data of Altman and

Gregory is -207.0±0.7 in excellent agreement with the calorimetric

value from Fontana's results. However, as MacWood (1958) points

out, if a value for S“ = 40.5 cal/mol*K is used for 11013(0), then

the 3rd law AH'* = 16.21 kcal/mol and AHf* 11013(c) from Altman and

Gregory is -205.2 kcal/mol in better agreement with the calorimetric

value of -205.4 from Barkelew and MacWood. At present, the value

-207.0±1,0 is accepted.

A value of -29.5±0.6 kcal/mol (Fuger (1976)) for soln

11013(0) at infinite dilution can be estimated from the experimental

AHsoij. Pu0l3(c) and Am0l3(c). This value with the AHf* U^*^(aq,

std. state) results in AHf® 11013(c) = -207.2±1.1 in support of

-207.0±1.0 kcal/mol.

- 71 -
There are other measurements pertaining to UCl^Cc).

Hardy-Grens (1964) reported pressure measurements for the

decomposition of UCl^Cc). The assumed reaction is:

4 /3UCl

2

(c) UCl^(g) + l/3U(c) ( 6 )

The reported (1030 to 1173 K) is 52,900 - 27. 9T cal/mol UCl^(g)

with = 57.7 kcal/mol (assuming = - 6 cal/mol*K) and ^^*293 = 35.8

cal/mol*K, resulting in -188.3 kcal/mol for 4Hf® UCl

2

(c) and S® UCl

2 (c)

=51 cal/mol*K.

Similarly Shchukarev et al. (1956) reported, in the temperature

range 863 to 1063 K:

^ = 51,000 - 27. 5T cal/mol

This results in 4H

2 gg = 55.0 kcal/mol UCl^(g) and = -186.3

kcal/mol for UCl

2

(c).

These two sets of measurements are in reasonable agreement with

one another but differ with both the calorimetric values and the

UC 1^-H

2 reduction equilibria values. It may be that the measurements

made do not pertain to the disproportionation process.

The tabulated z^f® is obtained from the accepted z^fflf® and ^£^f®.

- 72 -
UCl^(c)

The low temperature heat capacities (15-355 K) have been measured

by Ferguson et al. (1944a). The Gp®, (H-Hq)^ and S* at 298 K are

taken from Katz and Rabinowit ch ' s (1951) tabulated values. It is

appropriate at this point to point out that the S® UCl^ at 298.15 K (47.1^)

has been listed in various places as 47.4 cal /mol *K [for example Rand

and Kubaschewski (1963) and Rossini et al. (1952)], There is no

evidence that a reevaluation of the low temperature Cp and S° was

made. It would appear to be an error in transcription.

Ginnings and Corruccini (1947) measured the heat content of

UCl^(c) in the range 700 K to 273 K. Popov et al. (1959a) measured

the Cp frcm 450 K to 920 K (above the melting point). A reasonable

presentation of the high temperature measurements has been made by

Rand and Kubaschewski (1963) :

UCl, (c) G^ = 27.2 + 8.57xl0"\ - 0.79 x10^t"^ cal/mol*K

(400-800 K)

UCl^(l) Cp = 25.8 + 14.4xl0“\ ca 1/mol *K

(890-920 K)

In combination with the vapor pressure measurements (see UCl^(g))

the following is accepted :

at the melting point, 863 K (Mueller (1948)):

UCl^(c) -UCl^d)

= 10.9 kcal/mol

^ = 12.6 cal/mol-K

UCl^(c) is a key compound in the evaluation of a consistent set

of thermodynamic data for the uranium compounds, and until 1971 the

- 73 -
had been considered to be firmly established as -251 kcal/mol.

[See Rand and Rubaschewski (1963), Rossini et al. (1952), Brewer et al.

(1945)]. However, an inconsistency noted by Rand and Rubaschewski

2+

existed in the paths to zlHf® tJO^ (aq, std, state) and the various forms

of UO^(c), These inconsistencies can now be traced primarily to the

AHf^'s of UCl^(c) and U02Cl2(c).

These early measurements and subsequent ones in support of the

above mentioned value for AHf° have been recalculated and are listed below.

They are given no further consideration for the selection of

UCl^(c) .

These values were obtained from measurements of U(c) and

UCl, (c) in excess HCl so that:

4

U(c) + 4HCl(aq) ^ UCl^(aq,in excess HCl) + 2H2(g); (1)

UCl^(c) -* UCl^(aq.in HCl) ; (2)

and by difference

U(c) + 4HCl(aq) ^ UCl^(c) + 2H2(g); (3)

" ^2 " UCl^(c) - 4zlHf HCl(aq)

•k The measurements of Barkelew were for oxidation in HCl-FeCl^ soln.,

but by difference ZHI^ is for the same reaction.

- 74 -
Investigator T K of Cone. HCl, dHf'* UCl^(c)

measurement HCl*nH20 kcal/mol

Smith et al. (1969) 298 8.13 -251.0

Argue et al. (1961) 298 8.13 -250.5±0.6

Barkelew [MacWood (1946)] 273 * -249.8*

Biltz and Fendius (1928) 273 8 . -248.2**

Fitzgibbon et al. (1971) redatermined the dUf® UCl^(c) by two

independent paths^ and obtained significantly different results from

those tabulated above. Their measurements "were confirmed by

Cordfunke et al. (1976); the resultant newly established AHf® UCl^(c) =

-243.5±0.6 kcal/mol has resolved many of the earlier inconsistencies in the

uranium network of key values. Measurements by Fitzgibbon et al. also

offer a reasonable explanation for the more negative values arrived

at earlier.

One of Fitzgibbon’s paths is a repeat of the above; i.e., reactions

(1) and (2) for which = -137.5±0,5 kcal/mol when = 4.35, or

= -136,2±1.0 kcal/mol when = 6.83

HCl :8. 13 H

2 O) and = -45.7±0.5 and -39.3±0,5 kcal/mol, respectively.

Using the differentials 4Hf HCl (I 2 . 76 H

2

O) = -37.856 kcal/mol and

^

HCl ( 8 .I3H

2

O) = -36.671 kcal/mol, values of -243.22±0.70 and

-243.58±1,1 kcal/mol are obtained for zlHf® UCl^(c).

HCI/I2 . 76 H

2

O) and

Barkelew used a 12M HCl-FeCl^ solution and experimentally determined

the 4Hf HCl in this medium at 273 K £S -33.0 kcal/mol; certain auxiliary

components needed revising so that AHf HCl for use in the Barkelew

HCl-FeCl

2 medium is -32.7 kcal/mol.

Earlier evaluations had 4Hf° = -251, in exact agreement with Barkelew

(MacWood) ; an erroneous interpretation of 4Hf° HCl and the experimental

4H mix had been made.

- 75 -
The s for the individual reactions, particularly (1), are

significantly less negative than the earlier measurements. It would

appear, as explained by Fitzgibbon et al., that the earlier measurements for the reaction of U(c) with HCl(aq) were not for process (1).

Probably not all dissolved oxygen had been eliminated; nor had a

2 - scavenger been used (SiF^ in the acid) to prevent the formation of

insoluble residues from the reaction of U(c) in HCl solutions.

The other path, which confirms the above values is dependent

upon UO^Cc). Fitzgibbon's experimental measurements, and those of

Cordfunke et al. (1976) who repeated the measurements are cited in

the following Table 6 .

- 76-
•K «

vO on

m o o K o m O CN o CN

0) • • • • • • js: c O o o o o c -H +1 -^l -rl -rl -rl 3 CVJ o CO a ON

<4-1 ON

m C7N <r

T3 • • • • • • • ;4 on m a O o o m 'O 1 in I 1 o o 1 1 1 • o^4M

CN •f4 y-N NO

>0" c • • CN o o o • • • o o • p—4

•i4 CO a o O o +1 -rl o • 4J a •f; -fl -rl NO a -l-l o o

0) •H NO <r CN

m >ct -rl o eo on 'Cf ON a CN N • • r*4 s on m o a ON o o CJ o •H m -O’ 1 1 -O’ 1 m 1 1 I 1 CO o O CN • o o +

<u <• XJ

O

a c

=1 o 33 • ^4

a

CO • o c s <^J c a o 0 o • 0 CO CNJ 2

<u s CO « o m

CN • CN on < O o on O S • in o LO y-^ on CN c 3 • JS o ^4 <t o CN 0

CO . O <r

a ji o o w 1 •H a 4J CN -o + <

vO OJ S3 • 3

•4^ ai + < 0

(U 0 i-i c • <r ^4 •t4 <U + a y" 0 44 c <4-1 0 • + a o

CO o OJ S3 CO o • i-i H • • a < y-s 44 4J c o o • • o ic <- c

CDs

(a ^ ON CO o s . Csl y-N ^4 _CN <• 0)o • CO ONJ on o NO c c • o • o a ^4 c o

-<j“ + 'O

a • • o O CN o O o < s-y CN a a HM

<4-4 CO CO

a CO • y-s • • 'O’ o CN J c T o O ai O fi CN o u CN a a <u >4 •<f a (U o a • CN-4 CU o T OT o CO < a o

CO • • a rcr E

CO CN • + CN

T • o 0) s S3 CN o ^4 1 <r a a -o- • a on o • 4-» O 2 • X a < < r o • [2 04 ^ o CN • CN CO O O O CO "4 CJ c 4-1 '0-

CU in CO o a c < CN ^4 • o . (U CO <!• a o o • o w H •H CN O • T ON a a i cn <u <U ^ CN ^•<r • ^4 <u P> a o <t , o + o a i-l -4

4J CO o + CO

a + Oa • 0

(U CO • a y—s c <u > m W /-S —o c y^ CN o + t-4 a c u qj CJ a o s CN o 3 *f4 5J O o 'W CN ss a 3

<4-1 CN <T CO <r <4-1 /—V

3J y—\ •

-fa N-y + a o Ui + CN CN • o !-l O o o y—N

cn CN y-s

a < m on on CJ • IZS p^ o CJ a US 'w' <• c • N—' CN CN ONJ >3- o <r c (U o c CN O o O • CN •3 ^ C N—

a 'W/ CJ >w' 4J 44 o a a a a <r ^4 ONI CJ yo U

y-^ y^ y-N y—\

<r un NO

a ON o ^4 'W Nw' v.^ r-4 -x

77
The summation results in:

UO^Cc) + 4HCl(aq) - UCl^(c) + aH^Od) ; ^12)

and AH^2 ~ 23.37±0.9 and 33.55±0,60 kcal/mol, respectively. With

zmf® UO^Cc) = -259.3±0.2 based on the combustion work of Huber and

Holley (1969) and the integral values for AHf® HCl(aq), AHf® =

-243.95±1.1 (Fitzgibbon) and -243.27±0,65 kcal/mol (Gordfunke).

A weighted average from these four values results in the selected

AHf° UCl^(c) = -243.5±0.6 kcal/mol.

These two paths to ZiiHf® UCl^(c), one independent of other uranium

compounds and one dependent upon AHf® UO^Cc)^ are mutually consistent;

i.e._, the AHf® UGl^(c)^ independent of other uranium compounds supports

a U02(c) = -259.310.2 kcal/mol.

The Gibbs energy of formation has been calculated from the AHf°

and ^f'*.

UCl^ in Various Aqueous Media

The concern here is with the bulk ZlH . , that is, the AH . soln-’ ’ soln

of UCl^(c) in various aqueous media to form the real solution, which

in HCl may be highly complexed and in dilute HCl and HCIO^^ a partly

hydrolyzed one, correcting only for the presence of the foreign

silicate ion when present, temperature (if not at 298.15 K) and

possible impurities if known.

In all measurements cited here the concentration of the UCl,4

in solution is below .01m. There is at present insufficient

- 78-
experimental information to correlate the ZMi , UCl, soln 4 as a function

“uci^*

1. Measurements in HCl

In Fuger and Getting

\ (1976) the measurements were corrected and

extrapolated from HCl and HCIO, solutions in order to

4

obtain the

^f®, and S® for U^'*’ (aq^ std. state). For convenience they are

repeated here (uncorrected for hydrolysis) along with additional

measurements

.

Cone, m. Investigator moles HCl /kg H^O AH 1 kcal/mol

soln

Biltz and Fendius (1928) 7.2

* -39.4

Kilner (see Brewer et al

.

(1958)) 7.2 -36.2

Smith et al. (1969) 6.90 -40.0

Fitzgibbon et al. (1971) 6.82 -38.9

Argue et al. (1961) 6.82 -42.7

Maltsev et al. (1960) 5.53

** -43.47

Fitzgibbon et al. (1971) 4.35 -45.3

Hearne and White (1957) 2.08 -51.53

Fuger and Brown (1970) 1.02 -54.46

Argue et al. (1961) 1.02 -57.0

Measurements at 273 K, uncorrected.

Measurements at 293 K, uncorrected.

- 79 -
2. Measurements in HCIO,

4

Investigator Cone, m, ^soln

moles HClO^/kg H

2

O kcal/mol

Argue et al. (1961) **

.01 -54.6*

.05 -57.3

0.135 -55.6

0.200 -59.1

1.02 -61.0

Fontana (1947) 0 -44.7

0.1 -52.7

0.5 -55.7

2.0 -53.6

3. Measurements in Ac id -Salt Mixtures

Investigator Solvent '^soln

kcal/mol

Fontana (1947)

^ 0.1m HCIO,,

0.4m LiClo,4 [concentration UCl,~5xl0 m] -54.22

Hearne and White (1957) ^ HCl-LiCl with

[concentration TJCl, ~5xl0 ] constant |i=2.0

[H ] = 2.0 -51.53±0.7

1.6 -54.40

1.2 -56.10

1.0 -56.89

0.6 -57. 15

0.4 -57.96

0.2 -58.31

0.1 -58.43

0.04 -58.50

0.02 -58.46

The final smoothed values recommended for . soln of UCl^(c) as

a function of m of HCl and HCIO^ are given in Tables 7 and 8

.

Corrected for pol3rmerization of U(IV).

Measurements by Argue are for molar concentrations.

-80-
Table 7 Recommended Values of ZIH , UCl, (c) in HCl soin 4

Cone, m. AH , soln

AHf UCl^Caq,

moles HCl /kg H

2

O UCl^(c)

kcal/mol

1.0 -54.6 -298.1

2.0 -51.5 -295.0

3.0 1 00

•

00 -292.3

4.0 -46.2 -289.7

5.0 -43.7 -287.2

6.0 -41.3 -284.8

7.0 -39.25 -282.7

Table 8

Recommended Values of AH , UCl soln 4

(c) in HCIO^

Cone, m AH , soln AHf UCl^Caq

moles HOI 0, /kg H_0 UCl (c) in HCIO^)

‘r kcal/mol

0 -44.7 -288.5*

0.1 -52.7 -296.2

0.5 -55.7 -299.2

1.0 -55.1 -298.6

CM • 0 -53.6 -297.1

in HCl)

** This represents the 'real' in H2 O, i.e. the hydrolyzed solution^

not to be confused with the AH°gQ2.n “ (-141.3 - 4x39.933) + 243.5 = -57.5

kcal/mol for the ideal unhydrolyzed solution + 4C1“], The difference

between the two ZiH's soln. is a measure of ^yjrolysls'

-81 -
UCl^(g)

The vapor pressure measurements of Choporov and Chudinov (1968)

(650-750 K; Knudsen effusion) on UCl^(c) are in good agreement with

the evaluation of Rand and Kubaschewski (1963); a readjustment,

however, of the parameters of the

^

expressions for UCl^(l) (which

retains the fit of the experimental measurements of Young and Grady

(1958), Gregory (1948), and Jenkins and Anderson (1948)) has been

made in order to obtain slightly better agreement with the measurements

of Popov et al (1959a) who, from Cp measurements, calculated the AH =

11.9 kcal/mol of all transformations between 834 and 882 K (including

fusion)

,

The adjusted expressions are;

UCl^(c) UCl^(g) (1)

AG^= 51,900 - 93. OT + 13.8 T log T cal/mol (298-863 K)

at 298 K; AG^^ = 34.3 kcal/mol

AH®j^ = 50.1 kcal/mol

= 52.9 cal/mol-K

UCl^(l) UCl^(g) (2)

^7,000 - 134. 8T + 29.9 T log T cal/mol (863-1063 K)

at 1063 K: AG = 0

AH = 33.2 kcal/mol

= 31.2 cal/mol*K

The calculated tabulated values for ZlHf®, ZlGf° and S° for UC1^(£;)

are obtained from the sublimation process.

The earlier measurements are cited by Rand and Kubaschewski,

- 82 -
UCl-(c) and UCl,(c)

J 0

The low temperature thermal functions of UClg(c), 0-350 K, from

the specific heat measurements of Ferguson and Rand (1945)^ have been

tabulated by Katz and Rabinowitch (1951),

S® UCl^(c) at 298.15 K has been estimated from the measured values

for UCl^(c) and UClg(c).

The enthalpies of formation of UCl_(c) and UCl, (c) are discussed

D O together.

MacWood (1958), reporting on the experimental results of Barkelew,

cites the AH's at 273 K for the oxidation of UCl^(c) (zl^ = -24.1±0.1

kca 1/mol), UCl_(c) (AH = -33.9±0.1 kcal/mol) and UCl, (c) (AH = -44.6

5 o ±0.1 kcal/mol) in 12N HCl, 107o FeCl^ solutions, so that:

UCl^(c) + FeCl

2

(aq) ^UCl^(c) + FeCl

2

(aq) ; AH^ = 9.8±0.15 kcal/mol (1)

UCl_(c) + FeCl,(soln) ^ UCl.(c) + FeCl-(soln); AH„ = 10.7±0.15 kcal/mol

J J o Z Z

( 2 ) Using the relationship:

FeCl

2

(soln) + FeCl^ (soln) ; AH^ = -20.2±0.5 kcal/mol (3)

which incorporates the Barkelew experimental components and pertains

to the specific medium of a 12N HCl, 107 FeCl^ soln. at 273 K, the

following relationships are obtained:

UCl^(c) + l/2Cl2(g) ^UCl^(c); AH^^ = -10.4±0.6 kcal/mol (la)

and

UCl^(c) + l/2Cl2(g) ^UClg(c); AH^^ = -9.4±0.6 kcal/mol (2a)

-83 -
More recently Gross et al. (1971) reported the direct chlorination of U(c) with 012 ( 1 ), so that:

U(c) + 5/2Cl2(g) -*UCl^(c); = -247.8±1.0 kcal/mol (4)

which results in = -4.3±1,2 kcal/mol.

Obviously there is a major disagreement in the relationships

for UCl^lc; DOand UCl. (c) . If, from Gross et al., ZXEIf® UCl_(c) =

J

-247.8 kcal/mol one would expect ~ -4 and UCl, (c) ~ -252

za 0

kcal/mol, as opposed to 4iHf° UCl^(c) = -253.9 and AHf'* UClg(c) =

-263.3 kcal/mol from the MacWood-Barkelew measurements. Although

the Gross et al. measurements appear to be definitive for UCl^(c),

the overall picture they present appears unreasonable, as compared to

the other halides.

The Barkelew measurements for ZlH of oxidation of U(c) in this medium

are known to be in error; however, some of the other measurements

appear reasonable and lead to calculated values for z:lHf'’ that are in

agreement (within 2 kcal/mol) with the presently accepted values,

such as 11012 (0), UBr^(c), U(X]l

2

(c) and U0Br2(c). It may be that

is erroneous, and that a better value is ~ -19 kcal/mol for these

conditions. See Rand and Kubaschewski (1963) for the variation of

with HCl concentration. Use of -19 kcal/mol for results

in: = -9.2 kcal/mol and AEIf° UCl_(c) = -252.7 kcal/mol; and

la 5

ZIH- = -8,2 kcal/mol and UCl, (c) = -260.9 kcal/mol. These

43 O

values have been rounded. Further work is necessary on the

UCl^-UCl^-UClg relationships.

Use of ~ kcal/mol would result in ZMlf° UCl^ (c) = -206.6

kcal/mol from the MacWood-Barkelew investigation, in better agreement with the selected ZlHf° = -207.0.

-84 -
The Gibbs energies of formation have been calculated from the

selected £!Hf®'s and the calculated ^f®'s.

UClg(g)

Rand and Kubaschewski (1963) reviewed the vapor pressure data

through 1961 and reported 298.15 K to be 18.8 kcai/mol and

= 35.1 cal/mol*K.

The 4lCf“ UCl-(g) is derived from the

sub 1*

Gruen and McBeth (1969) studied the reaction of UCl^(c) with

Cl

2 (g) (450-650 K) spectrophotometrically ; the uranium pentachloride

was found to vaporize as a dimer molecule.

For the reaction:

2UCl^(c) + 01^(3) - U

2 Cl^Q(g);

= 15,130 - 15. 4T cal/mol (450-650 K)

The Uranium Pentachloride-Aluminum Chloride Vapor Complex

Gruen and McBeth (1969) studied the reaction of UCl^(c) with

AlCl

2

(c) and Cl

2 (g) spectrophotometrically; they report:

UCl^(c) + l/2Al2Clg(g) + l/2Cl2(g) UCl^ -A 1C 13 (g)

^ = 8,910 - 10. 7T cal/mol (440-630 K)

and

UCl^(c) +Al

2

Clg(g) UCl^*Al

3 Clg(g)

:1G = 15,780 - 15. 3T cal/mol (600-800 K)

-85 -
3.104 U-Cl-0 Compounds

UOCl(c) and U03r(c)

Katz and Rabinowitch (1951) tabulate the K's of Gregory (1945)

for the supposed equilibrium of the reduction of UOX

2

, where X =

Cl or Br. It was assumed that the products were UCK(c) and HX(g)

,

No analyses of the products were made.

The K's lead to the following for:

00X

2

( 0 ) + l/2H2(g) UQX(c) + HX(g)

X 4St

kcal/mol HX kcal/mol HX cal/mol-K

Cl 10.6 6.0 6.9

Br 10.1 6.1 6.0

If it is assumed that the 4Cp's are negligible, the resultant Z^“'s

for UOCl(c) and UOBr(c) are -222 and -213 kcal/mol, respectively;

however, the resultant calculated S^'s are 11 cal/mol*K for UOCl (c)

and 12 cal/mol*K for UOBr(c), obviously impossible values. Reasonable

values for S® UOCl(c) and S® UOBr(c) are ~ 24 and ~ 27 cal/mol*K,

respectively (based on a comparison of the measured S®'s for U(c),

U02(c), UCl

2

(c), UCl^(c), UOCl^, UO^Cl^ and U0Br2(c)).

If it is assumed that the measured K's are reasonable, but the

T dependency is in error, then, with the estimated S®'s for UOCl

and UOBr (and assuming ~ 0) one obtains ^ = 20 cal/mol*K and

4H = 19.5 kcal/mol for the UOCl(c) equilibrium and ZlisS = 21 cal/mol*K

and 4H = 20.2 kcal/mol for the UQBr(c) equilibrium. These AH's result in

AHf® UOCl(c) = -213 kcal/mol and AHf® UOBr(c) = -203 kcal/mol.

Hovever, these zlHf® values still appear unreasonable in comparison

-86 -
with the relationships derived for the PuOCl(c)-PuCl

2

(c) , PuOBr(c)-

PuBr^Cc), and AmOCl(c)-AmCl^ (c) systems from Fuger (1976). No

recommendations are made at this time.

U0Cl2(c)

Greenberg and Westrum (1956b) measured the heat capacity from

11.7 to 348 K (also some measurements from 5 to 10 K of lower precision). Thermal anomalies were not observed. The extrapolation

below 10°K was made assuming the absence of magnetic transformation

below 5 K. Their tabulated values are accepted.

Rand and Kubaschewski (1963) evaluated the earlier measurements

(700 to 800 K) on the equilibria:

2UOCl2(c) U02(c) + UCl^(g) (1)

and arrived at =57.1 kcal/mol UCl^^ with perfect agreement on

AS® from the experimental 4S^ at 750 K (46.4 cal/mol*K) and that

obtained from the calorimetric S°'s for UOCl^(c) and U02(c), and the

calculated S® UCl^(g) obtained from the UCl^(c) and a

= -6 cal/mol*K, These earlier measurements will not be cited here,

or reinterpreted. The resultant AHf® U0Cl2(c) = -254.9 kcal/mol.

Recently Knacke et al. (1972) reported measurements on the

same equilibria (in the range 700 to 1023 K), which result (with a

zXp = -6 cal/mol*K) in AH®^ = 57.3 kcal/mol, and AS®^ = 52.4

cal/mol'K, also in perfect agreement with the independently calculated

AS®. This AH® results in AHf® U0Cl2(c) = -255.0 kcal/mol.

-87 -
Barkelew, as cited in MacWood (1958), calorimetrically determined

the AHf°'s of many of the uranium-halogen containing compounds;

however, his determination of AHf° UCl^(c) is seriously in error

because of the reaction of U(c) in HCl; the AHf®’s of UCl-Cc) and UCl,(c)

j b

agreement with better determinations. Where possible, little weight

is given to the determination, even if it is in agreement with better

determinations, but calculations will continue to be made of the

dBf^'s from the Barkelew data, possibly giving alternative ways of

handling his data, in order to apply this knowledge to compounds

where other data are scanty or non-existent.

Barkelew 's measuren^nts on UOCl^ are viewed in this regard.

Since the measurement of the oxidation of U(c) in a 12N HCl,

10% FeCl^ is seriously in error, Barkelew 's zlH = -16.7 kcal/mol for

the oxidation of U0Cl

2

(c) in this medium is combined with the oxidation

of UCl^(c) in the same medium, AH = -24.1 kcal/mol. Since the final

solutions are the same, UO^Gl^ and FeCl^ forming in the medium:

UCl^(c) + H^OCaq) ^ UOCl^Cc) + 2HCl(aq); = -7.4 kcal/mol (2)

Barkelew 's values for ZAHf 's of HCl and H20(aq) in the 12N HCl, 10%

FeCl^ medium (which are partly experimental) have been corrected for

the presently accepted AHf® HCl ( 3 . 49 H

2

O) = -36.05 kcal/mol and AHf°

H20(1) = -68.315 kcal/mol and recorrected to 273 K, incorporating

the experimental determinations of ZAH . and AH,,, to obtain for mix diln

- 88 -
's HCl and H

2

O in the medium at 273 K, -32.7 and i69.45 kcal/mol,

respectively. These AHf * s and -243.6 for AHf UCl^(c) (at 273 K)

result in AHf® UOCl^Cc) = -255.0 at 273 K and -254.9 kcal/mol at 298 K.

The agreement here indicates that the Barkelew-MacWood data on

the reaction for U0Cl

2

(c) can be considered reasonable and can be

used in combination with the AH of U0Br«(c) to obtain reaction 2

AHf® U0Br2(c).

The Gibbs energy of formation is derived from the ASf® and the

AHf®.

-89 -
The low temperature heat capacities (6-350 K) have been measured

by Greenberg and Westrum (1956a). The smoothed thermodynamic functions

are accepted, Prins (1973) determined the high- temperature enthalpy

from 392 to 696 K relative to 298 K. The results are expressed as:

^“^298 " 22.93T + 5.562x10" V + 0.4643x 10^T"^ - 7,494 cal/mol

(298 - 700 K)

These functions join the low temperature results smoothly.

The recent measurements of Gordfunke at al. (1976) have clarified

the situation regarding the U02Cl2(c). Two different cycles

existed prior to this time for obtaining AHf®, both involving AHf°

UCl^(c). Rand and Kubaschewski (1963) obtained AHf® = -302.9 kcal/mol

(based on AHf° UCl^(c) = -251.3) from the Shchukarev et al. (1958a)

path which at that time was the only available one: a complete recalculation using presently accepted auxiliary /lHf°'s and /IHf ° UCl^(c)

= -243.5 kcal/mol results in AHf ° U02Cl2(c) = -292.0±2.0 kcal/mol.

The more recent cycle from Khanaev and Khripin (1970) results in

-297.0±1,0 kcal/mol.

k

More recently Gordfunke et al. (1978) reported measurements in the

range 366 to 649 K relative to 298 K:

Ht-H298 = 27.55T + 2. 178x10"V + 2.729x105t”^ - 9323 cal/mol

At 500 K and 700 K, H'j’-H293 are 1% and 0.2"% higher than those of

Prins (1973).

Another value for ZlHf*’ U02Gl2(c) can be derived, = -291.5 kcal/mol,

if one uses the Na 3U0£^(c) from the mass spectrome trie studies

of Battles et al. (1972) in the thermochemical cycle from O'Hare

et al. (1972) involving Na3U0^(c) and U02Gl2(c). Since this Z!\Hf®

Na 3 U04 (c) has a high uncertainty, a better approach is to obtain

zllif® U02Gl 2 (c) from the other thermochemical cycles and to obtain

zlHf® Na3U0^(c) from the O' Hare measurements as has been done by

Gordfunke and O' Hare (1977).

- 90-
Cordfunke et al. determined the relative to zlHf° UO^Cc^y)

and determined UO^Cl^Cc) = -297.1±0o3 and also repeated the

"Shchukarev" cycle and obtained -297.9±0,9 kcal/mol.

The thermochemical cycles studied are presented^ beginning with

the Cordfunke "yUO^" cycle^ which is the most straightforward of

the four available cycles. The measurements have been made in excess

5.545 molal HCl solutions.

7-U0^(s) + 2HCl(aq, excess) ""UO^Cl^Cin HCl) + H^OCin HCl);

= -15.53±0.07 kcal/mol (1)

UO^Cl^Cs) -* UO^Cl^Caq, HCl); = -17.06±0.05 kcal/mol (2)

so that

:

U0^(c,7) + 2HCl(aq, excess) “* UO^Cl^Cc) + H^OCin HCl);

4^2 = 1.53±0.09 kcal/mol (3)

With the differentials^ HCl (10.0 IH

2

O) and AHf H^O in HCl lO.OlH^O)

= -37.291±0.01 and -68 .440±0 .001 kcal/mol, respectively, UO^Cl^Cc)

= -297.12±0.3 kcal/mol.

The second of Cordfunke 's cycles, the "Shchukarev cycle”, and

the original Shchukarev results are tabulated.

TABLE 9

The "Shchukarev cycle" for the Enthalpy

of Formation of UO^Cl^Cc)

Reaction AH kcal mol

(Cordfunke) (Shchukarev)

(4)

(5)

( 6 )

(7)

UCl^(c) + 2FeCl^ (soln. ) + 2H20 (soln.) - UO^Cl^ (soln. ) + 2FeCl

2

(soln. ) +

4HCl(soln.)

U02Cl2(c) -.U02Cl2(soln.)

2FeCl^ (C) - 2FeCl

2

(soln.

)

2FeCl

2

(C) -* 2FeCl^(soln.)

-45.11±0.37

-24.56±0.17

-64. 64 ±0.15

-36.74±0.11

-45.50±0.10

-25. 44 ±0.07

-61.14±0.54

The solution refers to 1.98

containing 0.5 mass percent

mass percent of HCl^ i.e. HCl (IOO. 6 H

2

O)

of FeCl^, i.e. FeCl

3 (I8 OOH

2 O)

.

-91-
From the Cordfunke Ifeasurements

UCl^(c) + 2FeCl^(c) + ZH^OCsoln.) 2FeCl2(c) + 4HCl(soln.) +

UO^Cl^Cc); /^g = -48.45±0.45 kcal/mol (8)

Using auxiliary values for AHf® FeClg(c), FeCl^Cc) AHf

KCl (lOO.SH^O) and zlHf H^O in HCl lOO.dH^O) from Parker et al. (1976)^

/SSf° UO^Cl^Cc) = -297.89±0.9“ kcal/mol.

From the Shchukarev measurements

UCl^(c) + 2FeClg(c) + 2H20(soln.) ^ UO^Cl^Cc) + 2FeCl2 (soln.

+ 4HCl(soln. ) ; AHq = -81.20±0,55 kcal/mol (8a) oa The value for ZlHf® FeCl^Csoln.) = -101,2 kcal/mol from Wagman et al.

(1969) used to obtain the ZlHf*’ U02Cl2(c) = -292.0 cited previously

was probably not appropriate for this solution; a larger uncertainty

(±2.0 kcal/mol) was therefore assigned to the calculated 4flf°. Also

Shchukarev 's value for differs considerably from that reported

by Cordfunke, although the measurements for reactions (4) and (5)

are in reasonable accord. If the measurements of Shchukarev

for ZXH^ and are combined with Cordfunke 's values for AHg and

AHg =? -47.96±0.8 kcal/mol is obtained, which results in AHf® = -297.67

±1.0 kcal/mol, in good agreement with the two Cordfunke values and

that derived from the Khanaev and Khripin work cited earlier.

Cordfunke et al. (1976) at the same time measured the AHf® UCl, (c) = -243.27±0.65, relative to AHf® U02(c). This value was used to

be consistent with Cordfunke 's cycle.

-92 -
A detailed analysis of the Khanaev and Khripin measurement is

given under the discussion for UF^Cc) where the summation equations

are

:

UO^Cl^Cc) + SFeCl^Cc) + 3HF(aq) + HCl(aq) ^ ZH^OCliq) + 3FeCl^(c)

d-UF^Cc); AHg =+29. 54 ±0.32 kcal/mol (9)

and

FeCl^Cc) + UCl^(c) + 3HF(aq) FeCl^Cc) + UF^(c) + 3HCl(aq) ; = -5.564±0. 20 (10)

By difference , reaction (8) is obtained with AHg = -35.104±0.50

kcal/mol. However, here the H^O/HCl ratio is 3.91 and the integral

ZHlf® HCl(aq) = -36.449 kcal/mol is needed, resulting in the earlier

mentioned AHf® = -296.98±1.0 kcal/mol.

AHf° UO^Cl^Cc) = -297.2^±0.3 is the selected value.

A3f® has been calculated from AHf’ and ASf®.

- 93 -
U02Cl2(in HCl solutions)

There are scattered values in the literature for the AH

soln

HCl solutions. The following table sunnuarizes the

available measurements. In all cases the concentration of the

UO

2 CI

2

, although different in each case, is below

Table 10

1x10 ^ molal

The Enthalpy of Solution of UO

2

CI

2

in HCl

^C1 AH . soln

moKKgH^O)"^ kcal /mol

Fitzgibbon et al (1971) 6.82 -15. 34 ±0.5

Cordfunke et al (1976) 5.55 -17. 06 ±0.1

Prins (1973) 5.55 -17.30±0.23

Fitzgibbon et al (1971) 4.35 -18.75±0.5

O'Hare and Hoekstra (1974) 1.0 -22. 18 ±0.15

Shchukarev et al (1958a) 0.552

(0.57o FeCl3)

-25. 44 ±0.07

Cordfunke et al (1976) 0.552

(0.57o FeCl^)

-24.56±0,17

O'Hare et al (1972) 0.25

(containing NaCl)

-24.27±0.07

O'Hare and Hoekstra (1973) ft -24.19±0.09

Cordfunke (1975) 5.6x10'^ -26.03±0.03

Lipilina and Samoilov (1954) measured the the trihydrate

(using one mole of the salt in 50 moles H^O) as a function of the

HCl concentration. The experimental AH's have been corrected for

Revised from original -22.49±0.13 as explained by O'Hare et al (1976)

- 94 -
the dilution of 3H20(aq). These results are tabulated.

Table 11

The Enthalpy of Solution of UO^Cl^’SH^OCc) in HCl

®HC1

AH

exp

kcal/mol

AH corrected

kca 1/mol

0.80 -8.86 -8.86

0.97 -8.75 -8.74

1.55 -7.78 -7.75

1,98 -6.95 -6.91

4.44 -2.80 -2.56

In addition Prins measured AH , in 5.55 molal HCl (concen- soln

tration of UO

2 CI2 ~ .03 mol* kg = -2.41±0.01 kcal/mol, which corrected

for the 3H20(aq) formed = -2.04 ±. 01 kcal/mol.

As can be seen from the tabulated results the data on the AH , soln

U02Cl2(c) below one molal HCl are discordant; the measurements (m =

4.4 to 0.8) for UO^Cl^’OH^OCc) appear to be reasonable (the concentration of the UO

2

CI

2

, though, is much greater). From a smoothed plot

of the U0-C1„*3H«0 data we obtain at m = 4.40 AH , = -2.65 kcal/mol. 22 2 soln

From the U0«Cl_(c) data (m = 4,35 to 6.82) we obtain AH , = -18.65

2 2 soln

kcal/mol for m = 4,4 If the difference in the concentrations of

UO Cl (in HCl) is ignored, a AH , = -16.0 kcal/mol is obtained for

2 2 hydration

U02Cl2(c) + 3H20(1) - U02Cl2*3H20(c)

If, however, it is assumed that Ap^ (cp 1,11m - cp m < 0.001) = +0.5

Li 1j kcal/mol, estimated on the basis of the behavior of BaCl

2

(Parker et al

(1971)), AH = -15.5 kcal/mol. This confirms the AH, , . . = nydration

-15.3±0,23 from Prins (1973) measurements. In view of this confirmation,

- 95 -
all measurements on U02Cl

2

(c) and UO

2

GI

2

*3H20(c) have been merged to

obtain a set of smoothed values for AH , UO-Cl_(c) as a function soln 2 2

of the HCl concentration (Table (2.). The U02Cl2*3H20(c) measurements

were converted using the AH hydration = 15.3 and Aip^ =0.5 kcal/mol.

_3

The corresponding values for the AHf 's of U02Ci2 (in<10 and m = 1.11)

are also given in Table 1”^. This approach supports the Cordfunke

and Shchukarev measurements in the dilute region and is in agreement

with a AH® = 26.3 kcal/mol.

Table 12

The Recommended AH , U0-Cl«(c) in HCl soln 2 2

m AH soln AHf AHf

HCl U02Cl2(c) UO2CI2

^ m<10"^

kcal/mol

UO2CI2

, «-3

m UO2Cl2<10" I.-H •

Ir-I eII

10 -26.0 -323.2 -322.7

0.5 -25.2 -322.4 -321.9

1.0 -24.3 -321.5 -321.0

1.5 -23.4 -320.6 -320.1

2,0 -22.6 -319.8 -319.3

2.5 -21.7 -318.9 -318.4

3.0 -20.9 -318.1 -317.6

3.5 -20.1 -317.3 -316,8

4.0 -19.3 -316.5 -316.0

4.5 -18.6 -315.8 -315.3

5 . 0 -17.8 -315,0 -314.5

5.5 -17.1 -314.3 -313.8

6.0 -16.4 -313.6 -313.1

6.5 -15.7 -312.9 -312.4

7.0 -15.1 -312.3 -311.8

The Heat Capacity of Aqueous UO

2

CI

2

Solutions

Kapustinskii and Lipilina (1955) measured the specific heai

various aqueous solutions in the concentration range m =0.9 to

-9.6-
raols (kg H^O)

vs. m results

A graph of cp (the apparent molal heat capacity)

in the following tabulated values:

tn U0-C1„ 'Pp cal/mol*K

2 2 ^ 0.2

0.3

0.4

0.5

0.6

0.7

0.8

0.9

1.0

- 33.0

- 28.0

- 23.0

1 f“ -imKj m -) - 14.0

- 10.0

- 7.0

- 3.0

0

- 97 -
UO^Cl^-H^OCc) and UO^Cl^-SH^OCc)

The S®'s at 298 K have been estimated from the known S® U02Cl2(c), the

estimated contributions of the additional H^O and the Gibbs energies

of formation.

There are two self contained sets of measurements on the AH's

solution of the two hydrates and the anhydrous UO^Cl^Cc), those of

Prins (1973) [measurements in 5N HCl]^ and those of Shchukarev et al

(1959b) [measurements in H^O] . The following summarizes the

measurements and the calculated enthalpies of dehydration for:

U02Gl2-nH^0(c) -UO^Cl^(c) + nH^O(liq)

2 2

Prins Shchukarev et al

Hydrate AHgoin (1959b); AHgoin

(in 5N HCl) (in H^O)

kcal/mol

-2.41±0.01 -10.00i0.il

Prins Shchukarev

AHdehydration

kcal/mol

3

1 -10.19±0.3

0 -17.30±0.23

-13.32±0.23

-23.86±0.13

15.26±0.23

7.24±0.38'

13.86±0.17 (1)

10.54±0.26 (2)

As is obvious the results are not in agreement. Other information,

though, is available which supports the Prins measurements. A detailed

analysis is given elsewhere of the measurements of AH , of U0_Cl_(c)

in varying concentrations of HCl [from different laboratories], and

the Lipilina and Samoilov (1954) measurements of the AH , of the

trihydrate in HCl solutions over the range to 4.4, where

the molal concentration of UO

2

CI

2

is 1.1 (a much higher concentration

of UO^Cl^ than for the measurements on the anhydrous). Exclusive

* Corrected for the formation of H«0(aq) in 5N HCl, H O(liq) H 0(aq)

;

AH = -0.125 kcal/mol.

- 98 -
of the Prins neasurements they also result in a AH dehydration =

15.1 to 15.6 kcal/mol, confirming the Prins (1973) relationship for

the trihydrate complete dehydration. In addition, the individual

measurements from Prins are in agreement with other AH . measure- ° soln

ments

.

Now, AH^ can also be obtained:

UO^Cl^-SH^OCc) -* UO^Cl^-H^OCc) + 2H20(1) (3)

which from Prins' measurements is = 8 . 02 ±0.3 kcal/mol and from

Shchukarev's measurements is = 3.32±0.25 kcal/mol. Cordfunke (1965)

represents his vapor pressure measurements for:

U02Cl2*3H20(c) - UO

2

CI

2

-H20(c) + 2H20(g) (4)

as

log P ^ = -3,406±32/T + 8.497±0.10.

This converts to:

AG^*^ = 31,170 - 11 cal/mol

with AH**^ = 31.2±0.3 kcal/mol and AS”^ = 77.8 cal/mol*K. Converting to

equation (3) using for H20 (l) -* H^O(g), AH® = 10.52±0.01 kcal/mol

and AS® = +28.4 cal/mol*K, we obtain AH®^ = +10.6 kcal/mol and

^

3 ~ 21.0 cal/mol*K. Obviously this derived value for AH does

not confirm either the Prins measurements or Shchukarev's; the

AS® for the decomposition of a hydrate should be ~ 33-36 ca 1 /mol *H

2 0(g) *K

and ~ 5-8 cal/mol*H20 (l) *K. The derived AS®^ is 5-8 cal/mol *K too

high. If the assumption is made that the measured pressure is

- 99 -
reasonably correct, i.e., log = -2.93, then = 7.98±0. I kcal

and with the estimate = 71 cal/mol*K, AH, =29.1 kcal and AH_ = 4 ^4 3

8.1 kcal, in excellent agreement with the Prins determination.

The Prins' values for the dehydrations result in zlHf° UO^Cl^* SH^OCc)

= -517.4±0.4 kcal/mol and AHf* UO

2

CI

2

* H^OCc) = -372.8±0.5 kcal/mol.

With an estimate of 46±2 cal/mol *k for S* UO

2

GI

2

'H^OCc) , ASf®

UO^Cl^'H^OCc) = -124.0 cal/mol *K which in combination with AHf®

UO^Cl^’H^OCc)' = -372.8±0.5 kcal/mol results in AGf® = -335.8±0.8

kcal/mol.

The experimental AG°^ = 7 .98 ±0.1 kcal/mol, results in AGf°

UO^Cl^OH^OCc) = -453.0±0.8 kcal/mol. The calculated S'* UO^Cl^ -SH^OCc)

= 65.4 cal/mol*K is rounded to 65±2 cal/mol*K.

The experimental vapor pressure measurements of Cordfunke (1965),

with modifications, have been used to obtain the following equation:

UO^Cl^-SH^OCc) - UO^Cl^-H^OCc) + 2H20(g)

AG® = 29,040 - 70. 6T cal/mol

so that:

log P ^ = -3,170/T + 7.71

atm

- 100-
uo^ci^Cg)

Cordfunke and Prins (1974) chlorinated U_0 (c) (1140-1330 K)

.

J O

Volatile UO

2

CI

2

is formed. The transpiration method was used to

measure the pressure of U02 Cl

2 (g) formed.

The equilibrium:

U30g(c) + 3Cl2(g) - 3U02Gl2(g) + 02 (g)

can be expressed by:

= 142,300 ± 1600 - (67.8±1.3)T cal/mol O

2 (1140-1330 K)

A = -10 cal/mol*K is assumed so that at 298 K, ZlH = 151.5

kcal/mol and = 82 cal/mol*K, resulting in 3:Hf° = -234±3 kcal/mol

and S° = 87±4 cal/mol *K. The S appears low and is not tabulated

as a recommended value.

The results of Sietz cited by Kangro (1963) are not in agreement.

- 101-
UOCl^(c) and UOBr^Cc)

An estimate is given for UOCl^(c) S® = 41±2 based on the measured

S® for UCl^(c), UCl^(c), UOCl^Cc), and UO^Cl^Cc).

Similarly for UOBr^Cc) from the above pattern and the measured

UOBr^Cc), S® = 49±3 cal/mol is estimated.

Shchukarev et al, have determined the AH's of reaction of UCX^l^Cc)

= -28.55±0.13 (1958a) and UQBr^(c) = -45.42±0.21 kcal/mol (1958b) in

2% HCl, 0,5'%FeCl^ solution, as well as UCl^(c). [See discussion on

U02Cl2(c).]

The reactions of U0Cl2(c) and UQBr

2

(c) involve oxidation to the

(VI) state, e.g.,

UOCl^(c) + FeCl^(aq) + H20(aq) U02Cl2(aq) + 2HCl(aq) + FeCl

2

(aq)

as does the reaction of UCl^(c), so that'

U0Cl2(c) + FeCl

2

(c) + 2HCl(aq) UCl^(c) + FeCl

2

(c) + H20(aq)*; (1)

AH^ = 30.90±0.22 kcal/mol

and AHf® U0Cl2(c) = -277. 5±0. 8 kcal/mol.

The difference between Shchukarev ’s measurements on UOCl_(c) and UOBr^(c) yiei

3 3 I"' U0Cl2(c) + 3HBr(aq) - U0Br2(c) + 3HCl(aq); = 16.87±0.25kcal/mol (2)^

and AHf® UOBr^(c) = -228.2±0.84 kcal/mol

The value for UOCl^(c) appears to be reasonable; however, the

value for UOBr^(c) appears questionable; one would expect, if AHf®

UOCl^(c) is correct, AHf® U0Br2(c) ~ -238 kcal/mol. Since there is

As in the case of U02Cl2(c) the Cordfunke et al. (1976) measurements

on FeCl

2

(c)-FeCl

2 (c) are used.

-102 -
I

i

i

no other information at present to resolve this^ the uncertainties

have been enlarged to ±5 kcal/mol.

(U02)2Cl3(c)

Cordfunke et al. (1977) have prepared a new oxide-chloride by

two different methods, one the thermal decomposition of U02Cl

2

(o) in

vacuum at ~ 723 K, and the other by heating mixtures of UO^Cc) and

UO^Cl^Cc) (in a molar ratio of 1:3) ~ 773 K. They have characterized

this black-brown solid as (1102 ) 2 ^ 13 ( 0 ) with an orthorhombic unit

cell and have determined its

^

solution in 1.505M H

2

S 0^, (-46.025

±0.8 kcal/mol).

Previously Cordfunke et al. (1976) had determined the 4H , ^ soln

of UCl^(c) (-S0,94±0.20) and HC1\10.01H20U-0.245±0.01) and Cordfunke

and Ouweltjes (1977) had determined the of U02 (c, 7)

(-20.13±0.04 kcal/mol), all in 1.505M H^SO, .

A summation (after correction for dilution effects to the same

final solutions) results in:

3U02 (c, 7 ) + UCl^(c) + 2HCl(aq)* 2 (UO

2

) 2 CI

3

(c) + H20 (l) ; ^ = -19.64±0.3 kcal/mol UCl,4

The recommended value is derived from this reaction.

The integral AHf° HCl IO.OIH

2

O

-103 -
U02(0H)C1*2H20(c)

Prins (1973) measured the , of UO- (0H)Cl*2H_0(c) in excess soin 2 2

HCKIO.OIH^O) = -2.04±0.07 kcal/mol. With his UO^Cl^’SH^OCc)

= -2,4±0.01 kcal/mol, one obtains the relationship:

U02(0H)C1*2H20(c) + HCl(aq) -* UO^Cl^ -SH^OCc) ; ^ = 0.37±0.07 kcal/mol

from which AHf° = -480.5±0.4 kcal/mol.

- 104 -
3.105 U-Cl-F Compounds

UF^Cl(c), UF^Cl^Cc), and UFCl^(c)

The tabulated values for Cp® and are obtained from the

thermal functions estimated by Maslov (1964) on the basis of UF^(c)

and UCl^(c) (0-350 K)

.

No enthalpy or Gibbs energy of formation data are available.

- 105 -
3.106 U-Br Compounds

UBr^Cc)

Krestov (1972) has estimated the Gp as a function of the

temperature as:

Gp = 24.1 + 6.3xlO"^T cal/mol‘K (298-1000 K)

.

MacWood (1958) cites the measurements of Altman (1944) and Gregory

(1945) in the range 648-798 K on the reduction equilibrium:

UBr, (c) + l/2H.(g) - UBr-(c) + HBr(g) ( 1 )

From the K's as a function of T, a second law = 16.0 kcal/mol

(AHf° UBr

2

(c) = -167.0 kcal/mol) is obtained which is in agreement

with the calorimetric determinations cited below_, so that this AH® in

conjunction with the

^(^’^298^T^ ^^^”^298^T

used to

obtain S® UBr^(c)-S® UBr

2

(c) = 11.0 cal/mol, resulting in S® UBr

2

(c)

= 46±2 cal/mol*K,

MacWood (1958) reporting on the measurements of Barkelew cites

the AH's for reaction of UBr

2

(c) (= -49.5 kcal/mol) and UGl

2

(c) (= -42.0

kcal/mol) in solutions of 12N HCl, 10% FeGl^. In both of these solutions

the uranium is oxidized to the hexavalent state so that one can write:

UBr

2

(c) + 3HCl(aq) -» UGl^(c) + 3HBr(aq) ; AH^ = -7.5 kcal/mol (2)

with AHf® HCl(aq) = -32.7 and AHf HBr(aq) = -22.0 kcal/mol which are

consistent with Barkelew 's experimental components^ AHf® UBr

2

(c) =

-167.4 kcal/mol.

One can also use the measurement of UBr^(c) in the same media

(AH = -33.1 kcal/mol) to obtain the relationship:

UBr

2

(c) + HBr(aq) + FeGl^(aq) -* UBr^(c) + FeGl

2

(aq) + HGl(aq);

AH^ = -16.4 kcal/mol (3)

-106-
With the Barkelew relationship (ZXH = -20.2 kcal/mol) cited previously

(see UCl^(c) and UClg(c)) for the FeCl

2

-FeCl

2 relationship^

UBr^Cc) = -165.8 kcal/mol. This provides a crosscheck.

As cited previously, the high temperature equilibrium studies

result in zlHf® = -167.0 kcal/mol. The selected value is =

-167.0±1.0 kcal/mol.

The Gibbs energy of formation is derived from the selected Z^Hf*^

and z25f® = -20,6 cal/mol*K.

-107 -
UBr3(g)

Rand and Kubaschewski (1963), using the information given byKatz and Rabinowitch (1951) on the vapor pressure of UBr

3

(c,l)

(measurements by Altman and by Webster), obtained the following:

UBr

3

(c) -UBr^(g)

^ = 75,100 - 91. 8T + 13. 8T log T cal/mol

(298 to 1000 K)

at 298 K:

^ = 73.4 kcal/mol

= 51.6 cal/mol *K = 58.0 kcal/mol

UBr3(l) - UBr3(g)

zlG = 68,600 - 112. 8T + 23. OT log T cal/mol

(1000 to 1810 K)

^ = 0 at 1810 K

and

AW = 50.5 kcal/mol

=27.9 ca 1/mol *K

At 1000 K: UBr3(c) -*UBr3(l)

AH = 10.5 kcal/mol

AS = 10.5 cal/mol*K

These values are accepted.

- 108 -
UBr^(c)

3° has been estimated as 57.0±2.0 cal/mol. Krestov (1972) has

estimated the Cp UBr^(c) as:

Cp = 28.5 + .0071T cal/mol -K (298-792 K)

From the discussion on UBr^(g) we have_, at the melting point, 792 K:

. =11.6 kcal/mol fuston

. = 14.6 cal/mol*K fusion

Cp(l) = 41 cal/mol*K

Fuger and Brown (1973) measured AH’s of soln. of UBr^(c) and UCl^(c)

in IM HCl solutions , i.e., HC1(54.4H20) of -63.24±0.09 and -54.46±0.23

kcal/mol, respectively. From these measurements the following is

obtained

:

UBr^(c) + 4HCl(aq) UCl^(c) + 4HBr(aq) ; AH^ = -8.78±0.25 kcal/mol

( 1 )

It is assumed that the formation of HBr in a IM HCl solution is

equivalent to the formation of an additional mole of HBr in an HBr

IM solution. Using differential zisHf's, then for HCl and HBr of

-39.284 and -28.53 kcal/mol, respectively, zlHf° of UBr^(c) = -191.70

±0.7 kcal/mol.

Shchukarev et al (1959a) measured the AH for the reaction of UBr^(c) wit

27o HCl, O.fToFeCl^ solution as -53.27iC.36 kcal/mol. In an earlier

report (1958a) they measured the zlH for reaction of UCl^(c) as -45.50±0,10

kcal/mol in a similar solution. These reactions involve oxidation

to UO

2

CI

2

(in HCl) as:

UBr^(c) + 2FeCl^(aq) + 2 H

2

O (1) - U02Cl2(aq) + 2FeCl2(aq)

+ 4HBr(aq)

-109-
By difference, though, one obtains = -7. 77 ±0.37 kcal/mol

and with the same assumptions as earlier regarding the use of the

differentials and with z^Sf HCKlOOH^O) and

^

HBrClOOH^O) =

-39.489 and -28.692 kcal/mol, respectively, zliHf® HBr^(c) = -192.5^±0.8.

Similarly, Vdovenko et al (1973) obtain ZlH for UBr, (c) to be

-54.32±0.1, and -45.34±0.1 for UCl^(c), from which zlH^ = -8. 98 ±0.14

and Z^® = -191.4±0,6 kcal/mol.

MacWood (1958), reporting on the measurements of Barkelew (1946)

cites the ZlH's at 273 K of the reaction of UCl^(c) and UBr^(c) in 12N

HCl, 10% FeCl^ solutions, from which ZlH^ = -9.0 kcal/mol. With their

values for z^Hf's of HCl and HBr = -32,7 and -22.0 kcal/mol, respectively

(these values for the differentials at 273 K incorporate the Barkelew

experimental components), AHf® = -191.8±2.0 kcal/mol.

A weighted average results in ZlHf® = -191,7±0,6 kcal/mol.

The z2Sf® = -27.77 cal/mol*K and the AHf® results in AGf® =

-183.4+0.9 kcal/mol.

-110-
UBr^(g)

Rand and Kubaschewski (1963) reviewed the vapor pressure data

on the crystal and liquid and using an estimated = -6 cal/mol*K obtained

for the sublimation process^ (Erom the excellent agreement in the

measurements of Thompson and Schelberg (573 K - 723 K; effusion)

and Nottorf and Powell (723-773; gas saturation) given by Katz and

Rabinowitch (1951)):

UBr^(c) - UBr^(g)

= 49,400 - 92. 7T + 13.8 T log T cal/mol

These results lead to, at 298.15 K:

=47.6 kca 1/mol

zlG® = 31.9 kcal/mol

4^° = 52.6 cal/mol-K

These values are accepted in the absence of newer data.

A similar equation, with 4Cp = -11 for the vaporization process

UBr^(l) - UBr^(g)

results in . at the melting point, 792 K, = 13.2 kcal/mol and fusion or} ,

a 4S. , = 16.5 cal/mol*K.

fus ion

In view of the large . a reexamination of the vapor pressure ° fusion

data was made using the measurements of Gregory (815-1033 K; boiling

point), and Nottorf and Powell (798-898; gas saturation) as given by

Mueller (1948). The measurements (although there is some scatter)

are in agreement, and with 4Cp = -13 cal/mol*K, the following is

obtained

:

AG** = 43,300 - 131. 9T + 29. 9T log T cal/mol

- 111-
= 0 at 1040 K, 4H =29.8 kcal/mol and = 28.6 cal/mol'K.

At 792 K (melting point):

4H = 11.6 kcal/mol

4S = 14.6 ca 1/mol *K

UBr^(c)

_3

Krestov's (1972) estimated Cp = 36.0 + 8.0x10 T cal/mol'K

is accepted. The S® was estimated as 70 ±3 cal/mol'K. These values

result in a reasonable fit of the Blair and Ihle (1973) thermal decomposi

tion measurements of UBr^(c) to a third law plot.

The thermal decomposition of UBr^(c) was studied in the range

298-400 K.

UBr^(c) ”* UBr^(c) + l/2Br2(g)

The decomposition is reversible below 353 K. Above this temperature

the compound decomposes irreversibly into UBr^(c) and Br

2 (l). From

the vapor pressure measurements below 383 K a third law =6.0

kcal/mol (4Hf® = -194.0 kcal/mol). A second law = 5.4 kcal/mol

(AHf° = -193.4 kcal/mol). The average ZXHf® UBr^(c) = -193.7±2.0

kcal/mol is selected. The 4Gf*’ is calculated from the ^Hf” and the

estimated 4Sf°.

- 112-
3.107 U-Br-Q Compounds

UOBr(c)

See UOCl(c)

UQBr^Cc)

The from Greenberg and Westrum's (1956b)

tabulation from their heat capacity measurements from 5 K to 339 K.

I^ieller (1948) reported the vapor pressure measurements of

UBr^(g) over U0Br2(c) from the results of Gregory. For the range

710 to 960 K,

log P(atm) = -10,870/T + 8.32

=49.7 kcal/mol UBr^ at the mean temperature for:

2UOBr2(c) - UBr^(g) + U02(c) (D

which, with = -6 cal/mol*K results in AH^= 52.9 kcal/mol and

ilHf® = -228.2 kcal/mol. However, the =38.1 cal /mol *K at the mean

temperature results in a at 298 K = 44 cal/mol*K, whereas the value

calculated from the S®'s of U0_(c), 2^2U0Br_(c) and UBr, (g) [obtained

4

from 4^® , , and an estimate for S® UBr, (c)] results in

subl 4 -1- = 53 cal/mol*Ko The difference in z^® is too great to be rationalized;

the S® UBr^(g) accepted here does not appear to be appreciably too

high since 110 cal/mol*K is reasonable with regard to S® UCl^(g)

obtained from zlS® , ,, UCl, (c) which is in excellent agreement with

subl 4

the value obtained from the comparable U 0Cl

2

(c) decomposition to

UCl^(g).

- 113 -
In addition, using the AHf® UOBr^Cc) = -228.2 kcal/mol for

2UOBr

2

(c) ^ UO^Cc) + UBr^(c), (2)

^

2 ~ kcal/mol UBr^(c) which indicates a lower stability for

UOBr^Cc) than for UOCl^Cc). Again this does not seem reasonable.

For these reasons the Qregory data has been rearranged by:

1. Assuming that his pressure measurements are reasonable,

with log P(atm) = -4.47 at 850 K (obtained from his

equation)

2. Assuming at 850 K ~ 47 cal/mol*K (using the at 298 K

53 cal/mol*K and 4C ® = -6 cal/mol*K).

P

3. Recalculating at 850 K to be 57.3 kcal/mol from

the above listed values for log P and 4S°

results in AH®

^ at 298 K = 60.6 kcal/mol.

The resultant ZlHf® U0Br2(c) = -232.0 kcal/mol.

The Barke lew (1946) measurements on the oxidation reactions

of (1) U0Br2(c) and (2) U0Cl2(c) in a 12N HCl, lO^FeCl^ aqueous

solution at 273 K are now considered. As indicated in the discussion

on U0Cl2(c) the values obtained for the AHf® U0Cl2(c) from the

Barkelew data are in reasonable agreement with the preferred data,

although the sample of U0Cl

2

(c) used was not of high purity.

The ZlH of oxidation of U0Br2(c) in a 12N HCl, 10% FeCl^ aqueous

solution = -16.3 kcal/mol. This reaction was paired with the comparable

reaction of U0Cl2(c), zlH = -16,7 kcal/mol^ tc obtain Al^at 273 K =

-0j4 kcal/mol for

U0Cl2(c) + 2HBr(aq) -» UOBr^(c) + 2HCl(aq). (3)

As for the differential AHf HCl(aq), the Barkelew value for the

differential AHf of HBr(aq) in a 12N HCl, FeCl^ aqueous solution has

-114 -
been recalculated equal to -22.03 kcal/mol at 273 K (this value

incorporated Barke lew's experimental components)^ so that at 273 K

= -234.2 for UOBr^Cc) and at 298 K = -234.1 kcal/mol.

None of these values by themselves can be considered good.

However^ the recalculations have brought them, from very different

paths, into reasonable agreement.

4Hf° UOBr^Cc) has been taken to be -232.7±2.0 kcal/mol. This

value results in = 14.4 kcal/mol UBr^(c).

The zlGf® is calculated from the z4Hf® and S®,

- 115-
UO^Br^-nH^OCc) (n = 0, 1, 3)

The entropy for UO^Br^Cc) at 298 K is estimated on the basis of

the values for U0Cl

2

(c)^ UOBr^Cc), and U02Cl

2

(c); those for the hydrates

using a contribution to S® of 9.4 cal/mol*K per mole of H^O.

Recently Cordfunke et al. (1978) determined the enthalpy of

UO^Br^Cc) in the range 345 to 454 K relative to 298 K. The results

are expressed as

:

= 24.92T + 4.534xlO"V - 7833 cal/mol

(298 - 460 K)

The values for the zlHf®'s are obtained from the recent measurements

of Prins et al. (1978) on the AH's of solution of lJ0^(c,7)

,

UO^Br^ *0 .024H20 (c)

UO^Br^* 1.05H20 (c) and U02Br2*2.92H20(c) in excess HBr (14 .SBH^O) . The

AH's are given, corrected to the appropriate hydrates.

Compound AH , soln

kcal/mol

U0g(c,7) -17.88±0.05 (1)

U02Br2(c) -25.38±0.10 (2)

U02Br2*H20(c) -17.61±0.15 (3)

U02Br2*3H20(c) -10.43±0.20 (4)

The AHf® U02Br2(c) = -271.84±0.32 is obtained from the summation:

00^(0,/) + 2HBr(aq) U02Br2(c) + H20(aq) ; AH^ = 7.50±0.11 kcal/mol (5)

using the differentials, AHf® HBr(aq) = -27.60 kcal/mol and AHf® H20(aq)

= -68.361 kcal/mol.

For the AH's of hydration:

UO-Br^(c) + H^O(l) - U0„Br„*H-0(c) ; AH, = -7.82±0.18 kcal/mol (6)

Z Z Z Z Z Z 0

U02Br2(c) + 3H20(1) - UO^Br^ • 3H20(c) ; AH^ = -15.09±0.22 kcal/mol (7)

- 116 -
from which AHf° UO^Br^ (c) = -347.97±0.34 kcal/mol and AHf°

U02Br2 • 3H20 (c) = -491.88±0.36 kcal/mol.

The AH's of hydration for UO^Br^Cc) and U02Cl2(c) appear to be

essentially the same.

The earlier measurements of Shchukarev et al. (1958a) on AH , soln

UO„Cl«(c) = -25.44±0.07 and (1958b) on AH ^ UO„Br.(c) = -31.23±0.2 22 soln 2 2 kcal/mol in 2% HCl, 0.5% FeCl^ aqueous solutions lead to:

U02Cl2(c) + 2HBr(aq) ^ U02Br2(c) + 2HCl(aq)

;

AHg = 5.79±0.5* kcal/mol (8)

which, with the differentials for the formation of HX(100H20) , AHf HCl

= -39.489 and AHf HBr = -28.689 kcal/mol, results in AHf° U02Br2 =

-269.8 kcal/mol, in poor agreement with the Prins et al. results.

From the Shchukarev et al. (1959b) AH , measurements of the soln

anhydrous U0^Br2(c) and the two hydrates in H

2 O, AH^ = -8.9 kcal/mol

and AHy = -11.8 kcal/mol, also in poor agreement with the Prins et al

.

results

.

The Shchukarev et al. (1958a, 1958b, 1959b) measurements on the

U02Br

2

'TiH 20 system are suspect and are rejected for the following

reasons

:

1. The AH , FeCl-(c) in a solution of 2% HCl by Shchukarev

soln 3

et al. (1958a) has been found to be in error.

2. The AH's of hydration of U02Cl^(c) to the two hydrates

from Shchukarev et al. (1959b) appear to be incorrect.

See U02Cl2-H20(c) and U0^Cl2 ' 3H20(c)

.

3. A comparison of the AH's of solution of U02Br2 in a solution of 2% HCl, 0.5% FeCl^ (1958b) and in H.O(l) (1959b)

See U02Cl2(c) discussion on the "Shchukarev" cycle.
with those for UO

2

CI

2 CC) indicates an inconsistency which

may be caused by erroneous measurements on U02Br

2 (c)

.

The Prins measurements, in contrast, offer a consistent

picture on the behavior of UO

2

CI

2

*nH20 (c) and U02Br

2

•nH20 (c)

.

The AGf°’s have been calculated from the AHf°'s and the ASf°’s.

U0Br2(c)

See UOCl^Cc)

U02(0H)Br*2H20(c)

Prins et al. (1978) measured the AH , = -6.67±0.01 kcal/mol soln

of UO

2

( 0H)Br* 2 H20 (c) in HBr (14

.

63 H

2

O) , Combining this with their

measurement for AH , of U0_(c,7) in HBr (14.63H„0) = -17.88±0.05

soln 3 2. one ob ta ins :

U0^(c,7) + HBr(aq) + 2H20(aq) ^ UO

2

( 0H)Br« 2 H20 (c) ; AH = -11.21±0.05

and AHf° = -468,0210.3 kcal/mol.

The differential AHf's to be used are AHf HBr (14.63H20) = -27.60

and AHf H20 (aq) = -68,361 kcal/mol.

- 118 -
3.108 U-Br-Cl Compounds

UCl. Br (c) (x=l,2) and UCl, Br (c) (x=l,2,3) 3-x X 4-x X

^ ^

MacWood (1958) reports the K’s for: (1) the reduction of

various UCl, Br (c) in the range 650-750 K from the results of ^“X X

Gregory (1945) and (2) the exchange equilibria with HBr(g) in the

range 600-773 K from the results of Altman (1944), Examples of

these equilibria are

:

1. UCl^_^Br^(c) + l/2H2(g) ^ UCl^_^Br^_^ (c) + HBr(g)

and

UCl^_^Br^(c) + l/2H2(g) - UCl3_^Br^(c) + HCl(g)

2. UCl^_^Br^(c) + HBr(g) UCl

3

_^Br^_^^ (g) + HCl(g)

These K's and the ZlH's calculated from the ZlHf®'s selected

here were used to obtain the S°'s for the various III and IV mixed

halides. They are to be considered approximate values.

MacWood also reports the measurements of Barkelew on the zlH's

of solution in a 12N HCl FeCl

3

solution of the mixed halides

(III and IV) as well as UCl

3

(c)j UCl^(c)^ UBr

3

(c), and UBr^(c) from

which the AHf®'s can be obtained by summation reactions. Examples

follow, as well as the measured AH's of solution.

- 119 -
and

UCl^BrCc) + HCl(aq) -* UCl„(c) + HBr(aq);

^ " ‘^soln - ^soln

UClBr^Cc) + HBr(aq) - UBr^(c) + HCl(aq);

^ = ^soln 'JClBrjCc) - UBr^(c)

^soln kcal/mol

UCl^Cc) -42.0

UCl^BrCc) -44.2

UClBr^Cc) -48.2

UBr^Cc) -49.5

UCl^(c) -24.1

UCl^BrCc) -25.7

UCl^Br^Cc) -29.2

UClBr^Cc) -31.8

UBr^(c) -33.1

-120-
3.109 U-I Compounds

Ul3(c)

Krestov (1972) has estimated the Cp as a function of temperature

Gp = 25.1 + 5.8xl0‘\ cal/mol-K (298-800 K)

MacWood (1958) tabulated the decomposition pressure measurements

(523-666 K) of MacWouu et al. (1944) for:

UI^(c) Ul3(c) + l/2l2(g) (1)

From the values for ZHif" 1113(c) listed below; zlHf® = -111,7±1.0

kcal is selected. The resultant = 19.7±0.5 kcal/mol is used in

combination with the above pressure measurements and A(H-H-_o) and

zy o T

A(S-S2gg)^ to obtain AS^ = 21.8±1.0 cal/mol‘K and A(S° UI^(c)-S° Ul3(c))

= 10 cal/mol*K, S° 1113(0) = 53±2 appears reasonable in comparison with

the measured S® 11013(0) and 4S® (Cl-I) . This value for S® 1113(0) in

conjunction with the estimated Gp equation completes the tabulated

thermal functions.

The second law value for AH®^ from the above decomposition

pressure measurements is 18.5±1.0 kcal/mol which results in AHf® =

-113.0±1.2 kcal/mol.

MacWood (1958) reporting on the measurements of Barkelew (1946)

cites the AH's reaction of 0X3(0) = -56.5 and UI^(c) = -44.2 kcal/mol,

respectively, with a solvent of 12N HCl saturated with 13(0). In

these final solutions the uranium is considered to be in (IV) state.

Then for the relationship:

UI^(c) - 013(c) + 1/213(0); AH3 = 12.3±0.6 kcal/mol (2)

and AHf® = -111.7±1.0 kcal/mol.

-121-
Tveekrem and Chandrasekharaiah (1968) from cell measurements

obtained - -98.8±0o4 kcal/mol for: 643 K

U(c) + S/ZI^Cg) -» UI^(c) (3)

The third law AH® = -131.5±1.0 and AHf® = -110.2±1.0 kcal/mol.

The ASf® = -0.6 cal/mol*K results in Z^f® = -111.5±1.2 kcal/mol.

The S® has been estimated as 63±2 cal/mol based on a comparison

of the differences in measured entropies of chlorides and iodides for

the polyvalent halides.

Popov et al. (1959a) have measured the Cp of UI^(c) and (1) from

373 K to 873 K.

Rand and Kubaschewski (1963) used these measurements to obtain

for UI^ (c)

Cp = 34.8 + 2.38xl0"\ - 4.72xlo\"^ cal/mol*K (380-720 K)

The results of Popov et al., for the phase changes have been

accepted. (See the discussion of the vapor pressure data for UI^(c)

and (1) under Ul^(g).) The total for the phase changes of UI^(c)

to UI^(l) (m.p. 779 K) is 9.2 kcal/mol. For UI^(l) : Cp = 39.6

cal/mol*K (820-870 K)

The selected Cp® at 298 K is obtained from the extrapolation of

the measurements below 473 K. Smoothed tabulated functions are given

here incorporating the estimated S® . zyo

The measurements of Fuger and Brown (1973) on the enthalpies of

solution of UI^(c) (-69.2910.3 kcal/mol) and UCl^(c) (-54.4610.23

-122 -
kcal/mol) in IM HCl, i.e., HCK 54 . 4H

2 O) lead to:

UI^(c) + 4HCl(aq) UCl^(c) + 4HI(aq); = -14.83±0.38 kcal/mol (1)

With the assumption that the differential AHf Hl(aq) in HCl is

the same as in an HI solution = 13.265 kcal/mol and using AHf HCl =

-39.284, AHf^ = -124.55±0.07 kcal/mol.

In addition, from their measurements in 6M HCl (i.e. HCl

(

8 . IdH^O)

)

( 8 .I3 H

2

O) of UI^(c) = -57.41±0.07 and UBr^(c) = -49.17±0.20, there is:

UI^(c) + 4HBr(aq) ^ UBr^(c) + 4HI(aq); AH^ = -8.24±0.21 (2)

With the same assumption, i.e., that the differentials, AHf ’ s of

HBr = -26.103 and HI = -11.202 in HCl are equivalent to the differentials in their own respective solutions, AHf° = -123.51±0.7 kcal/mol.

The average AHf° = -124.0±0.7 kcal/mol is accepted.

The calculated ASf° = -4.52±2.0 cal/mol*K and the AHf° = -124.0

±0.7 results in AGf° = -122.7±0.9 kcal/mol.

- 123 -
The vapor pressure of UI^(c) has been measured by Thompson and

Schelberg (as given by Katz and Rabinowitch (1951)); 573-683 K;

effusion method and that of UI^(l) by Gregory (1946) ; 823-923 K;

boiling point measurements against a known pressure of 12 (g) to

prevent decomposition of These have been reviewed by Rand and

Kubaschewski (1963); however^ the calculated values obtained for

. (16.9 kcal/mol) and . (21.6 cal/mol*K) at the melting fusion fusion

point, 779±3 K, (Gregory (1946)) are unreasonable.

Popov et al. 's (1959a) measurements on UI^ (c) and (1) indicate a

total = 9.2 kcal/mol and 4S = 12.1 cal/mol*K for the phase changes

(zlH = 3.53±0.1 kcal/mol at ~ 723 K, AH. . = 5.64±0.2 at the

trans ^ fusion

melting point) which are reasonable although possibly low. This

indicates that either one or both sets of vapor pressure measurements

are in error.

If the UI^(c) measurements are correct, S° Ul^(g) ~ 125 cal/mol*K

at 298 Ko If the UI^(l) measurements are correct, then Gregory’s

(1946) measurements can be described (with ACp = -15 cal/mol*K) by:

UI^(l) Ul^(g) (1)

AG^ = 47,770 - 150. 67T + 34.54T log T cal/mol; (792-1000 K)

and at 900 K:

AH^ = 34.3 kcal/mol

AG^ =4.00 kcal/mol

AS^ = 33.7 cal/mol*K

Using '^(^"^298^900 '^^^4 system = -21.3 cal/mol*K (Wagman

et al. (1977)), results in AS®^^ at 298 K = 55 cal/mol*K and S° Ul^(g)

= 118 cal/mol*K, which appears more reasonable, in a correlation of

- 124 -
the S®'s for the Th-U halides, than 125 cal/mol*K, For the vaporization, the expression for is accepted. The extrapolated boiling

point is 1020 K with diS = 31.9 cal/mol*K. For the sublimation,

vap

UI^(c) Ul^(g), (2)

with AH

2 = 34.3 kcal/mol at 900 K and the comparable A(H-H

2 gg)gQQ

=

-16,00 kcal/mol from the Thl^ system (Wagman et al. (1977)) one

obtains at 298 K:

kcal/mol

cal/mol*K

= 33.9 kcal/mol

The tabulated values are obtained from these values.

An approximate equation then for this process is

:

AS

2 = 52,400 - 101. 9T + 16. IT log T cal/mol; (550-700 K)

-125 -
3,110 U-I-Gl-Br Compounds

UCl^r(c) and UBr^I(c)

MacWood (1958), reporting on the measurements of Barkelew, cites

the reactions of UCl^I(c) (zlH = -28.0 kcal/mol), UCl^Cc) = -35.7

kcal/mol) and UBr^lCc) (AH = -36.7 kcal/mol) with a solvent of 12N

HCl saturated with 12 (c) at 273 K, Since the resultant final °''lutions

are considered to contain the U(IV) species, a summation results in:

001 ^( 0 ) + 17212 (c) - UCl^Kc); AH^* = -7.7 kcal/mol (1)

The calculated AHf® = -214.7±2,0 kcal/mol.

UCl^Kc) + 3HBr(aq) UBr^Kc) + 3HCl(aq); AH

2

'' =+8.7 kcal/mol (2)

With AHf HCl(aq) = -32.7 kcal/mol and AHf HBr(aq) = -22.0 kcal/mol

which pertain to the particular solutions used by Barkelew and incorporates the experimeptal components measured by Barkelew, AHf°

UBr

2 l(c) = -173.9±2.0 kcal/mol.

These values are accepted.

MacWood (1958) cites the measurements of Davidson et al. (1945)

for the I

2 pressure (effusion method) over UCl

2 l(c) and UBr

2 l(c).

UCl2r(c) -*UCl2(c) + 17212 (g) (3)

(472 - 575 K)

UBr^Kc) - UBr

3

(c) + l/2l2(g) (4)

(551 - 655 K)

With the accepted AHf® UCl^Kc), AH^ = 15.16 kcal/mol at 298 K and

with a 4Cp = -2 ca 1/mol *K the equilibrium K’s can be fit to:

= 15,760 + 4.61 Tlog T - 31. 5T cal/mol

it The 4Cp for this is negligible.

- 126-
This results in = 9.8 kcal/mol and = 18 cal/mol’K^ from

which the tabulated values for and S° for UCl^lCc) are obtained.

Similarly = 14.0 kcal/mol at 298 K. Here^ however^ a

similar treatment of the K' s results in a ^ 13 cal/mol which is

4

too low (i.e., the calculated S° UBr^lCc) = 64 cal/mol«K^ too high

a value); no improvement is made by adjusting the 4Cp. In the absence

of more information on this system^ no recommendation is made for S“

or ^f® for UBr^Kc).

-127 -
3.200 U-X Compounds with N

3.201 U-NH^-F-0 Compounds

(NH^)

3

U02F^(c) and NH^[ (U02 ) 2

F

5

] (c)

Mukhametshina et al. (1974b) from a series of reactions in 2M

HCl reported

3NH^F(c) + UO^F^Cc) - (NH^)^U02F^ (c) ; AH^ = -17.74+0.27 kcal/mol (1)

which, using ZlHf® NH^F(c)* = -111.7±0.2 kcal/mol, results in =

-748.0±0o7 kcal/mol.

Similarly, Suponitskii et al. (1974) obtained:

NH^(c) + 2U02F

2

(c) - NH^ (UO^) 2^5 (c) ; ^2 = -12.99±0,50 kcal/mol (2)

which results in AHf® = -915.1±0.8 kcal/mol.

The above values are accepted, in preference to those which can

be derived from the decomposition pressure measurements of Sudarikov

et al. (1970) on (NH^)^U02F^ (c), the ammonium dioxopentaf luorouranate

(VI) and the intermediate compound, NH^[ (1102 ) 2^^] (c), the ammonium

tetraoxopentaf luordiuranate (VI).

•This is compatible with the CODATA selections. It is derived from

the following:

(1) Higgins and Westrum (1961) measured the AH's soln NH-(g) = -42.53±0.06 and NH,F(c) = -14.31±0.03 kcal/mol in HF(1) so that - NH3 (g) + HF(1) - NH^F(c); AH = -28.22±0.07 kcal/mol

From Johnson et al. (1973) AHf° HF(1) = -72.55±0.06 kcal/mol, so that

AHf“ NH^F(c) = -111.75±0.06

(2) Armstrong et al. (1959) obtained:

NF

3 (g) + 4NH3 (g) 3NH^F (c) + N2 (g); AH = -259.5±1.0 kcal/mol

which using AHf’ NF

3 (g) = -31.60±0.1 kcal/mol (evaluated from sources

in JANAF (1971)) results in AHf® NH^F(c) = -111.67±0.4 kcal/mol.

-128 -
The decomposition pressures are described by:

(NH^)^U02F^(c) - 0.5NH^[(U02)2F3l(c) + + 2.5HF(g) (3)

log P(atm) = 8.01 - 4. 468x10^ /T (516-538 K)

= 102.3 - 176.4xlO"^T kcal/mol U

and

0.5NH^[(U02)2F^](c) - UO^F^Cc) + 0.5NH^(g) + 0.5HF(g) (4)

log P(atm) = 6.927 - 4.878x10^/1 (647-703 K)

= 22.3 - 30.3xl0"^T kcal/mol U

Assuming the 4Cp's for these decompositions are negligible and

with the use of:

NH^F(c) - NH

3 (g) + HF(g); = 35.40±0.20 kcal/mol (5)

we obtain + 3zlH^ = -18.4±6.0 kcal/mol and =

-2z^, + = -9.2±2.0 kcal/mol.

4 5

Since the temperature range for these decomposition measurements

is short, these results are not unreasonable. If the measured

pressures are used with the preferred values for AH:^, the following

relations result for the decompositions:

= 99,700 - 171. 6T cal/mol

log P(atm) = 7.80 - 4.36xlO^/T

and

AG^ = 24,200 - 33. IT cal/mol

log P(atm) = 7.53 - 5.29xl0^/T

Approximate S° values of 86r5 and 90±5 ca 1/mol *K are calculated

for NH^ ( 1102 ) 2^

3

( 0 ) and (NH^)

2

U02p^ (c) , respectively.

-129 -
NH^[(U02 ) 2

F

5 ^ 3H20 (c) and NH^ [ (UO2

) 2^

3

]

4 H20 (c)

Tsvetkov et al (1971) measured the vapor pressure of the hexagonal

tetrahydrate NH^[ (1102 ) 2^

3

]

* 4H20 (c) and the orthorhombic trihydrate.

They reported the following:

NH^[(U02)2F3]-3H20(c) - NH^ [ (UO^)

2

F

3

] (c ) + SH^OCg) (1)

log P(atm) = 7.02±0.54 - 2,684±181/T (303-372 K)

and

NH4[(U02)2F3]*4H20(c) NH^[ (U02)2F3] -3H20(c) + H20(g) (2)

log P(atm) = 6.84±1.10 - 2,533±350/T (303-338 K)

These vapor pressure equations lead to the following:

• ( 1 ) ( 2 )

= 12.3±0.8 kcal/mol H20 (g) 11.6±1.6 kcal/mol H20 (g) = 32.2±2.5 cal/mol H20 (g) = 2.7±1.0 kcal/mol H20 (g)

31.3±4.5 cal/mol H20 (g)

2.3±1.6 kcal/mol H20 (g)

In both cases appears low. Normally a Z:S of 35-36 cal/mol

H20 (g) would be expected. The ^®'s have been accepted and have been

used with the estimated = 35±2 cal/mol H20 (g) to obtain ZlH® ^ =

13.1±1.3 kcal/mol H20(g) and z^®2 = 12.7±1.7 kcal/mol H20(g)

.

These lead to the following approximate formation properties:

trihydrate tetrahydrate

-1128 -1198 kcal/mol

Z^f -1015 -1072 kcal/mol

S 116±10 12 6±10 ca 1/mol *K

-130-
3.300 U-X Compounds and the Alkali Metals, M (M=Li, Na, K, Rb , or Gs)

3.301 U-F-M Compounds

MDFg(c) (M = Na(3)^ K, Rb, and Gs)

Kudryashov et al, (1978a) measured the enthalpy of oxidation of

the alkali metal uranium (V) hexafluorides, MUF^ (c) in aqueous solutions containing 0.8% H

2

O

2

(i.e. (ZSOH^O)) , 1.72% K^CO^ + 2.5% KOH

(i.e. KOH( 120H20) ) , and the reaction of UCl^(c) in the same media

(AH

2 = -175.4±0,3 kcal/mol).

MUFg(c) + [ 3K

2

C 0

^ + 4K0H + I/2 H2

O

2

] (aq) -* [K^UO^CCO^)^ + 6 KF

+ MOH + 2H

2

O] (aq) ; (1)

UCl^(c) + [BK^CO^ + 2K0H + H^O^] (aq) [K^U02 (C 0^)^ + 4KC1

+ 2H20](aq); (2)

With their measured AH , KCl(c) (3.96±0.01 kcal/mol) and their AH's

soln

of solution MF(c) tabulated below, the following composite is obtained from

which the tabulated AHf®'s are derived:

UCl^(c) + 0 . 5 H202 (aq)* + MF(c) + 5KF(c) MUF^(c) + K0H(aq)'+ 4KCl(c) ; AH

3 (3)

MUF,

6

AH^

kcal

AH T MF(c) soln

kca 1 /mol

AH

3 kcal

Na(,3) -117.0 ± 0.3 +0.05 ± 0.01 -94.59 ± 0.6

K -114.1 ± 0.1 -4.08 ± 0.01 -101.62 ± 0.5

Rb - 111.8 ± 0.8 -6.34 ±0.10 -106.18 ± 0.7

Gs -109.6 ± 0.5 -8.77 ± 0.06 -110.81 ± 0.7

The values used for AHf KOH(120H_0) = -115.077 kcal/mol and zlHf

H^O (250H 0) = -45.69 kcal/mol are^^ad jus ted from Wagman et al. (1973, etc.)

to Be compatible with GODATA selections.

-131-
NaUF_(c) and Na.lIF_ (c)

/ Z Q

Katz (1964^ 1966) measured the vapor pressure of UF^ (g) (gas

o evolution) over UF^*NaF(c) and UF,*2NaF(c) and reported (1966):

o o

2(UFg*NaF(c) "* UF^(g) + UF^*2NaF(c) (1)

log P(atm) =8.18 - 3.4xlO^/T (340 - 400 K)

= 15,900 - 37. 4T cal /mol

and (1964)

UF,*2NaF(c) ^ UF^ (g) + 2NaF(c) (2)

b b

log P(atm) =6.37 - 4. 18x10^ /T (500 - 573 K)

^2 = 19,100 - 29. IT cal/mol

Gathers et al, (1958) reported the existence of UF^*3NaF(c) and

measured the UF, (g) vapor pressure over it by the transpiration method

o (626 to 873 K) as did Groves (1961).

However, Katz (1964) found no evidence for UFg*3NaF and indicated

that the measurements by Gather et al. (1958) were on UFg*2NaF(c).

Malm et al. (1966) prepared UF-»NaF and measured the UF^ vapor pressure

o o at 373 K ~ 35 mm, in fair agreement with Katz's measurements.

Tentatively the measurements of Katz for the systems specified

are accepted and, with an estimated ~ -11 cal/mol*K, = 16.6

and ZlH- = 21.5 kcal/mol UF^ from which the tabulated Zllfflf^'s are

2 6

obtained.

-132 -
3,302 U-F-O-M Compounds

M^UO^F^Cc) (M = Na, K, Rb, Cs)

Mukhametshina et al. (1974b) measured the AH's of solution of

MF(c)^ UO^F^Cc)^ and M^UO^F^Cc) in 2M HCl. These measurements lead

to the following:

3MF(c) + UO^F^Cc) M^UO^F^Cc); AH

M AH kca 1/mol

M^UO^F^Cc)

Na -7.58±0,27

K -28.82±0.26

Rb -32.95±0.66

Cs -34.86±1,38

The selected values are based on these relationships.

M(U02)2F5(c) (M = Na, K, Rb, Cs)

Suponitskii et al. (1974) from a series of zXH's of solution

these compounds in 2M HCl obtained the following enthalpies for

formation from MF(c) and U02F2(c):

MF(c) + 2U02F2 (c) ^ M(U02)2F5 (c) ; AH

AH kca 1 /mol

M M(U02)2F3

Na -8. 15 ±0.22

K -18.64±0.24

Rb -23.28±0.62

Cs -27 ,32±0.70

The selected values are based on these results.

-133 -
M^(U02)2Fg(c) (M = K, Rb, Cs)

liikhametshina et al (1974a), from a series of /^'s of solution

of these compounds in HCl reported, for the enthalpies of reaction

of alkali metal fluorides with U02F

2

(c) :

5MF(c) + 2U02F2 (c) - (UO

2

) 2

F^ (c) ; AH

M

K

Rb

Cs

AH

kcal/mol

Mg(U02 ) 2

Fg (c)

-52.76±1.2

-63.38±1.5

-68. 16 ±2. 25

The selected values are based on these tabulated numbers.

- 134-
3.303 U-Gl-M Compounds

MUCl^(c) (M = K, Rb, Cs)

The following summarizes the AH's for the formation of

MUCl^(c) from the components salts (MCl(c) and UCl^(c)) from measurements in 27o HCl, 0.57o FeCl^ solutions, they are accepted.

MCl(c) + UCl^(c) - MUCl^(c); AH

AH kcal/mol MUOl^

K Rb Cs

Martynova et al. (1968) -5.5±0.5

Vdovenko et al. (1974) -6.8±0.6 -10.5±0.8 -13.6±0.8

M2UC1^(c) ( M = Li, Na, NaK/2, K, Rb, Cs)

The following summarizes the AH's for the formation of

M

2

UClg(c) from the component salts (MCl(c) and UCl^))

:

2MCl(c) +UC1,

4

(c) -M-UCl,

Z o (c); m

AH kcal/mol M-UCl,

2 6

M = Li Na NaK/2 K Rb Cs

Martynova et al. (1968) -1.5±0.6 -6.0±0.6 -9.8±0.6

[measurements in 2%HC1,

0.57, FeCl^]

Vdovenko et al. (1974) 1.1±0.6 -1.7±0.7 -9.4±0.5 -16.1±0.6 -24.1±0.S

[measurements in 27,HC1,

0.57, FeCl^]

Fuger and Brown (1971)

[measurements in IM HCl] -26.7±0.3

[measurements in 6M HCl] -25.8r0.4

The selected values were obtained from the average AH.

-135 -
Rb^UClg (c)

The results of Vdovenko et al. (1974), from measurements in 27o

HCl, 0.5% FeClg solutions lead to:

4RbCl(c) + UCl, (c) ^ Rb,UCl„(c); AH = -16.5±0.8 kcal/mol Rb,UClo(c)

4 4 o 4 o This is the basis for the selected value,

GsU^ClgCc)

The results of Vdovenko et al, (1974) from measurements in 2% HCl,

0.5%FeClg solutions lead to:

CsCl(c) + 2UCl^(c) CsU

2Clg(c); AH = -13.1±1,4 kcal/mol CsU^Clg

This is the basis for the selected value.

- 136-
MUClg(c) (M = Na(a and 3)^ K, Rb, and Cs)

Kudryashov et al. (1978b) measured the enthalpy of oxidation of

the alkali metal uranium (V) hexachlorides^ MUCl^^ in TU HCl, 0.5%

FeCl^ aqueous solutions

:

MUClg(c) + [FeCl

3 + 2H

2

O] (aq) - [UO

2

CI

2

+ FeCl^ + 4HC1 + MCl](aq); (1)

With their measured AH's of solution for FeCl^(c) (-32.55±0.1 kcal/mol),

FeCl^Cc) (-18.7±0.2 kcal/mol) and MCl(c) cited below and the results

of Vdovenko et al. (1973) for the reaction of UCl^(c) in the same

media:

UCl^(c) + [2FeCl3 + 2 H

2

O] (aq) ^ [UO^Cl^ + 2FeCl2 + 4HC1] (aq) ; (2)

AH

2 = -45.3 ±0.1 kcal/mol

the following summary reaction is obtained:

UCl^(c) + FeCl

3

(c) + MCl(c) - FeCl

2

(c) + MUClg(c); AH

3 (3

MJCl, (c) AH. AH MCl(c) AH

0

Na(a)

1

-61.6 ± 0.5

kcal/mol

1.19 ± 0.02

j

3.6 ± 0.6

Na(3) -61.4 ± 0.4 1.19 ± 0.02 3.4 ± 0.5

K -51.7 ± 0.3 4.37 ± 0.02 -3.0 ± 0.5

Rb -45.0 ± 0.2 4.08 ± 0.02 - 10.1 ± 0.4

Cs -41.9 ± 0.2 4.01 ± 0.02 -13.2 ± 0.4

The AHf°'s are obtained from AH^,

-137 -
3,304 U-Br-M Compounds

M^UBr^Cc) (M = Rb, Cs)

The results of Vdovenko et al. (1973), from the 4fl's of solution of the

components in 0.5% FeCl^, 2% HCl lead to:

2RbBr(c) + UBr^(c) ^ Rb^UBr^Cc); AH^ = -14.86±0.15 kcal/mol Rb^UBr^

( 1 )

2CsBr(c) + UBr, (c) -» Cs^UBr, (c) ; AH_ = -23.02±0.15 kcal/mol Cs-UBr^

4 Z o Z Z 0

( 2 ) These are accepted.

-138 -
o III. Tables of Values for Z^f% S°

,

Cp®, and H-Hq

at 298.15 K and Z^Hf® at 0 K for the Uranium- Halogen

Containing Compounds

The values are consistent with the CODATA key values for thermodynamics and the IAEA series^ "The Chemical Thermodynamics of Actinide

Elements and Compounds",

Included in the tabular summary are values for the non-halogen

containing uranium compounds which were necessary for this evaluation.

The compounds covered are listed in the standard order of

arrangement except for those containing the alkali metals which are

arranged by compound class of the uranium- halogen compound with the

alkali metal.

See Section 1-2.5 and Section I - 2.4 for the conventions

used for the Table headings^ the chemical formulae and physical

states

.

-139-
Note Added in Proof

There are significant new results that necessitate modification

of the values selected for certain uranium fluorides. These adjusted

values are included in the tabulation.

The calorimetric measurements of Cordfunke, (private communication, June 1980) j.esult in:

AHf’ UF^Cc) = -360.6±1.0 kcal/mol

AHf° UF^(c) = -459.1±1.0 kcal/mol

These values are based on two different measurement paths.

The paths are:

1. The direct fluorinations of UF^Cc) and UF^(c) to UF^(c).

2. Solution measurements of UF^(c), UF^(c), and UO^CcjY) and

and TJ^Og(c) in a H^SO^, Ce(S0^)2> H^BOg aqueous medium.

The above values are accepted and result in the adjustment of

the values for UF(g) , UF

2 (g), UFg(g), UF^(g) and UF^ • 2 . SH^OCc)

as well, since they are all dependent upon the AHf° UF^(c).

The values for UF^ 25^^^’ ^^4 5^^^’ UF^(c,6) and

UF^(g) are also dependent upon UF^(c); however, there are indications from solution calorimetric measurements in progress by

O'Hare (private communication, June 1980) on UF^(c,a) and UF^(c,3)

that their AKf^'s should be more positive, rather than more negative as required by the above cited AHf’ UF^(c) and the Agron

(1958) data.

Until the situation regarding UF^(c,a) and (c,3) is clarified,

the values for these two compounds and for UF^ UF^ ^(c), and

UF^(g) are maintained as in Section II, but with larger uncertainties.

- 140-
In addition, Westrum (private communication, June 1980) measured

the low temperature heat capacity for UF^Cc) resulting in S® =

29.50±0.10 and = 22,73±0.10 cal/mol'K which replace the estimated

S® = 30.0±1.0 and C

® = 23.4±1.0 cal/mol-K.

P
tills

liltI

MU

PHl)»>

OF

U-M^LOGEN

CONTG

CMPOS

ON

COOATA

BASE

IN

KCAL/MOL

C

CAL/MQL.K

o o o o

CM • • — O 'C O

O W 'O W

• I • I O > lO

(S (Ni • !

\no

— O

nj • m

o in • I •

O o o o o

— •- — — o o o — o

• • • • • «• • • • o o rvj o N o o m « mm o

C\i w in 0* w O w o o a N a

• 1 • \ • 1 • 1 • 1 » 1 • 1 • 1 • 1 m o >0 O' o 4> eg N -f in eg " eg eg

m in m o n o O' o o

• • N • • • • o • • in • • o OJ o 0* o in o in o eg o

• 1 • 1 • 1 • 1 • 1 • 1 • 1 • 1 • 1 • 1 a (\J +- o + <¥ O' <¥ O' +• a « a

fV eg o N eg N n a

1 o o o o

•« O O Q a •« o

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Westrim, L. F . > and HuntzicUer, J. J . ; J. CheM. Tl'ierMod y nan . 3_-. hi

•1971)

l4olt» A. S. ; PoseVv J . C. j and Rapp, i< . t. . 5 Inorn, C1k?m . A. 73:1

( 1965 ) Younq, H. S.i and Gradv> H. r . ; Paper 70 in Katz, ,T . ,T . ; and Rabin

owitch , L . , v?d'.a . ; U . 0 . AFC T .1 D-3590 > ( 19 30 )

Za'ibo',; , K. F. i 'Heats of ForMation of Lower Actinide Fluorides froM

Hass Spectroeetr ic 'studies', Proc. 1st iiit. Coni'. 1. or iMetr •; and

1 herMod ynanics > Warsaw p. 40.3., (Auq. 35 -Sept. 4, 5.969)

-159-
rI
V. Appendix: Thermal Functions for Some

Uranium- Halogen Containing Compounds

-160-

T K Cp

cal /mol

100 7.01

200 7.99

298. 15 9.04

300 9.05

500 10.11

1000 10.76

1500 10.93

2000 11.01

100 9.48

200 11.93

298.15 16.00

300 16.06

500 18.33

1000 16.83

1500 17.35

2000 18.10

100 12.61

200 15.41

298.15 17.76

300 17.79

500 20.27

1000 21.49

1500 22.05

2000 22.58

UT(g)

H-Hq S®

ca 1/mol cal /mol*

696 51.83

1439 56.94

2278 60.34

2294 60.39

4229 65.31

9492 72.58

14919 76.98

20405 80.14

UF2(g)

858 58.48

1908 65.66

3279 71.18

3309 71.28

6887 80.37

15552 92.45

24060 99.34

32937 104.44

UF3(g)

1009 63.11

2411 72.72

4044 79.33

4077 79.44

7929 89.23

18457 103.78

29348 112.60

40504 119.02

-Al-

-<G-H

29g)/T

cal/mol*K

67.65

61.14

60.34

60.34

61.41

65.37

68.55

71.07

82.69

72.52

71. 18

71.18

73. 16

80. 18

85.49

89.62

93.46

80.89

79.33

79.33

81.46

89.36

95.74

100.78
UF^Cc)

T K Cp “-’^98 s" -(G-H29s)/T

cal/mol*K cal /mol cal /mol • K ca 1/mol • K

298.15 23.38 0 30 30

300 23.39 43. 30.145 30.

400 24.12 2419. 36.974 30.927

500 24.85 4867. 42.434 32.700

600 25.58 7389. 47.029 34.715

700 26.31 9983. 51.027 36.766

800 27.04 12651. 54.588 38.775

900 27.77 15391. 57.815 40.714

1000 28.5 18205. 60.779 42.574

-A2-
T K Cp H-Hq -(0-H

2 gs)/T"

cal/mol'K cal /mol cal/mol*K cal /mol •

K

298. 15 21.815 4764 80.176 + 8.2 80. 176 + 8.2

300 21.852 4804 80.311 + 8.2 80. 177 +8.2

400 23.304 7070 86.818 + 8.2 81.054 + 8.2

500 24.114 9434 92.113 + 8.2 82.753 + 8.2

600 24.598 11882 96.556 + 3.2 84.693 + 8.2

700 24.906 14358 100.373 + 8.2 86.667 + 8.2

800 25.114 16860 103.713 + 8.2 88.593 + 8.2

900 25.259 19379 106.680 + 8.2 90.441 + 8.2

1000 25.365 21910 109.347 + 8.2 92.201 + 8.2

1100 25.444 24451 111.768 + 8.2 93.871 + 8.2

1200 25.505 26998 113.985 + 8.2 95.456 + 8.2

1300 25.553 29551 116.028 + 8.2 96.961 + 8.2

1400 25.591 32109 117.923 + 8.2 98.392 + 8.2

1500 25.622 34669 119.690 + 8.2 99.753 + 8.2

298.15 26.222

UF

3 (g)

5597 90.055 + 3.0 90.055 + 3.0

300 26.270 5646 90.218 + 3.0 90.056 + 3.0

400 28.222 8380 98.069 + 3.0 91.113 + 3.0

500 29.342 11263 104.498 + 3.0 93. 167 + 3.0

600 30.022 14233 109.913 + 3.0 95.519 + 3.0

700 30.460 17259 114.576 + 3.0 97.916 + 3.0

800 30.756 20321 118.664 + 3.0 100.259 + 3.0

900 30.965 23407 122.298 + 3.0 102.510 + 3.0

1000 31.117 26512 125.570 + 3.0 104.655 + 3.0

1100 31.231 29629 128.541 + 3.0 106.693 + 3.0

1200 31.319 32757 131.262 + 3.0 108.629 + 3.0

1300 31.388 35892 133.772 + 3.0 110.468 + 3.0

1400 31.443 39034 136.100 + 3.0 112.217 + 3.0

1500 31.487 42181 138.271 + 3.0 113.882 + 3.0

S and -(G-H«q_)/T values for UF^(g) are increased by 8.2 cal/mol'K to be

consistent with equilibrium data; similarly 3.0 cal/mol*K are added for

UF^Cs).

-A3-
I

I

uFe(g)

T K Cp H-Hq S* -(°-«298

cal/mol* K ca 1/mol cal/mol*K cal/mol

298.15 30.985 6384.4 90.233 90.233

300 31.044 6441.7 90.425 90.233

4w0 33.436 9677.4 99.717 91.484

500 34.797 13095.1 107.337 93.916

600 35.620 16619.1 113.760 96.702

700 36.149 20209.4 119.293 99.543

800 36.506 23843.2 124.145 102.322

900 36.757 27507.0 128.460 104.991

1000 36.941 31192.4 132.343 107.535

1100 37.078 34893.7 135.871 109.953

1200 37 . 184 38607.0 139.102 112.249

1300 37.267 42329.8 142.081 114.431

1400 37.333 46059.9 144 . 846 116.506

1500 37.387 49796.0 147.423 118.482

1600 37.431 53537.0 149.838 120.367

1700 37.468 57282.0 152.108 122.168

1800 37.499 61030.3 154.251 123.892

1900 37.525 64781.5 156.279 125 . 543

2000 37.547 68535.1 158.204 127 . 129

)/T

-A4-
UBr^ (c)

T K Cp «-«298 S* -(G-H2,g)/T

cal/mol*K cal /mol ca 1 /mol •

K

cal/mol*K

298.15 25.98 0. 46.00 46.00

300 25.99 48. 46. 16 46.00

400 26.62 26’’?. 53.72 47.03

500 27.25 5372. 59.73 48.99

600 27.88 8129. 64.76 51.21

700 28.51 10948

.

69.10 53.46

800 29. 14 13831. 72.95 55.66

900 29.77 16776. 76.42 57.78

1000 30.40 19785. 79.59 59.80

UBr^(c,l)

T K Cp «-«298 S® -(G-H2,g)/T

cal/mol* K cal/mol cal/mol.

K

cal/mol*

K

298 .15 30.62 0. 57.00 57.00

300 30.63 57. 57.19 57.00

400 31.34 3155. 66.10 58.21

500 32.05 6325. 73.17 60.52

600 32.76 9565. 79.07 63.13

700 33.47 12877. 84.18 65.78

792 34.12 15986. 88.35 68.17

792 41.00 27586. 103.00 68. 17

800 41.00 27914. 103.41 68.52

900 41.00 32014. 108.24 72.67

1000 41.00 36114. 112.56 76.44

1040 41.00 37754. 114.17 77.86

-A5-
T K Cp S® -(G-H29g)/T

cal/mol

ca 1/mol cal/mol* K cal/mol*K

298.15 26.83 0. 53.00 53.00

300 26.84 50. 53.17 53.00

400 11 .M 2763. 60.97 54.06

500 28.00 5534. 67.15 56.08

600 28.58 8363. 72.30 58.37

700 29.16 11250. 76.75 60.68

800 29.74 14195. 80.68 62.94

UI^(c,l)

T K Cp «“^98 S® -(G-H29g)/T

cal/mol * K cal/mol ca 1/mol'

cal/mol *K

298.15 30.20 0. 63.00 63.00

300 30.27 56. 63.19 63.00

400 32.80 3226. 72.29 64.22

500 34.10 6577. 79.76 66.61

600 34.92 10031. 86.06 69.34

700 35.50 13553. 91.48 72.12

779 35.88 16373. 95.30 74.28

779 39.60 25573. 107.11 74.28

800 39.60 26404. 108 . 16 75.16

870 39.60 29176. 111.49 77.95
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