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ORNL-3391

Contract No. W-7405-eng-26

CHEMISTRY DIVISION

MIXTURES OF METALS WITH MOLTEN SALTS

M. A. Bredig

DATE ISSUED

BUG 15 1363

 

OAK RIDGE NATIONAI, LABORATORY
Oak Ridge, Tennessee
operated by
UNION CARBIDE CORPORATION
for the
U.S. ATOMIC ENERGY COMMISSION

 
 

 

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CONTENTS

Abstractl‘.l...I.........OQ-.....I'C....l‘......l..C.'I.

IntroduCtion......‘O.....I.......I..l....'..."C'.'.-l.....O‘.I...

"Metallic" Metal SolutionS.eec... T eesetesessessetsnessetseebense e

A]kali‘Metal_Alkali-Halide SystemS. ® 5 4 0 00BN LR ORI COIBETEONEPIETSIO RS

Phase Diagrams..........flfl.l

Physical Properties.....

S 9 8 48 48 %9 S e OS2

LI I I B BN TN B B B K N RE BN BN BN

Alkaline-Farth-Metal—Halide Systems.eceeceiecsnccsverercncnee
Phase Diagrams...‘...l....-..I..I...........'I..l...‘.l.

EleCtriCall Conductalnce....‘......‘.l...l.........l..l..l

Rare-Farth-Metal—-Trihalide Systems...

Hlase Diagrams....-.............'I..................

Electrical Conductivityeeseesccoacsnss

o 6 8 08 s e 80P

llNOnmetallic“ Metaal SolutionSO..iI.......l‘..CG.“.O'G...'I....-..
TranSition Metals..'......l.....ll....I.l‘..l..l..ll...ll'...

Posttransition MetalSeeeecesosesesesosasocccscsossssacossncacsssssae

SUNMAryYeeoesse

® o 0% 60 P e &t e e Pe e R

ReferenCGS.....‘...-........‘..............

& " @ o8 s et S e e b

Page

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oo oWWw M~ e

30
32
37
41
43
bds
51
52

 
 

 

 

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MIXTURES OF METALS WITH MOLTEN SALTS

 

M. A. Bredig

ABSTRACT

A review is presented of various types of solutions of
metals in molten salts, especially in their own molten ha-
lides. With relatively little reference to the older liter-
ature, the progress made in the last 20 years is discussed.
Roughly, the solutions are classified into two groups: The
metal may retain, to some degree, its metallic properties in
the solution, or it may lose them through strong interaction
with the salt solvent. The alkali-metal systems are typical
examples of the former type, while solutions of cadmium or
bismuth represent the second. Equilibrium phase diagram
data are presented in detail for many metal-salt systems.
These include critical solution temperatures, that is, tem-
peratures above which metal and salt are miscible in all
proportions. Electrical conductivity is singled out as a
most significant physical property from which conclusions on
the state of the electron in the solution may be drawn. In
the electronically conducting solutions, notably of the al-
kali metals, the electrons may be thought to resemble F
centers in color-center colored crystals. In solutions where
electronic conductance is absent, monomeric, dimeric, and
even more highly polymerized species of the solute metal in a
low valence state must be assumed to occur.

INTRODUCTION

The history of the subject of interactions between metals and their
salts in the molten state, though not well known to most chemists, is at
least 150 years old.* It includes observations made by Davy in 1807 on
the production of deeply colored melts and on difficulties in recovering
metal at the cathode upon electrdlysis of molten alkali-metal hydroxides.
Similar phenomena were encountered by chemists who, either in small-scale
laboratory research or in larger-scalé industrial production, have been
or are concerned with the recovery of certain metals by molten~salt elec-

trolysis.?

Colored metal "fogs" were recognized by various investigators?
as one important cause of low current yield in the electrolysis of molten

salts. Various theories ranging from colloidal suspensions to "subhalides"

1

 
 

 

 

were propoéed to explain the observations. One also finds in the liter- kfi;
ature the statement by Nernst:4 "For not a single metal do we know a

nonmetallic solvent in which the metal would dissolve without chemical

interaction and from which the pure metal could be recovered simply by
recrystallization ... ." We know today that many such molten-salt sol-

vents do exist. A point of view which appears to be exactly opposite

that Of Nernst, toward which one group of authors”~7 seems to be inclined,

stresses the notion that metal atoms may indeed be dissolved as such in

molten salts, especially when the interaction with the molten solvent is

not a strong one. Neither view, if taken in the extreme, appears to be

‘:T):

consistent with experimental observation. _

The present discussion may profitably be restricted to cases in
which a metal dissolves in one of its own molten halides. Salts contain-
ing anions such as nitrates, sulfates, or phosphates are likely to de-
compose in reactions with the metsl. Little of a precise nature is as
yet known about solutions of one metal in the molten salt of another,
but a few specific cases will be briefly mentioned. Facts will be
presented which demonstrate that the variety of true solutions of metals
in fused salts is considerable. This includes, at one extreme, dissolu-
tion with relatively little solvent interaction, which might be described
by the concept of electrons substituted for anions in cavities, or that
of solvated metallic electrons. At the other extreme, dissolution with
chemical reaction between solvent and solute occurs, as in the formation =
of "subhalides,” where unusually low ionic valence states of the metallic |
element are formed. Contrary to the older literature® as well as to more .
recent statements,® no evidence whatsoever seems to exist for the occur-
rence of the colloidal state of metals in fused salts (with the exception
of highly viscous systems such as silicates).

It is convenient to divide solutions of metals in their molten ha-
lides into two main categories: |

1. The metal imparts partially metallic character to its solution
in the salt. The properties of the solution reflect the presence of
mobile electrons. These are the true metal solutions. They are of par;

ticular interest because they represent a novel state of metallic matter. ‘ 0

 
 

 

 

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i»

 

They are somewhat similar to solutions of metals in liquid ammonia and
similar solvents, but are distinguished from these largely by the ionic
nature of the solvent (the salt), which, it is to be especially noted,
usually contains one of the constituents of the metal which it dissolves,
namely the cation. Semiconducting solutions‘(i.e., those having positive
temperature coefficients of the electronic part of the conductance) are
included and may, in fact, represent the majority of examples in this
class.

2. Strong interaction (i.e., chemical reaction, oxidation-reduction)
occurs between metal and salt. The metal assumes a valence higher than
zero but lower than normal. This class might be designated "subhalide"
solutions. One might distinguish two subgroups, depending on the nature
of the second, that is, the metal-rich phase (solid or liquid) that, at
saturation, is in equilibrium with the molten salt-rich phase:

(a) The reaction goes so far as to lead to the crystallization of a
solid "subhalide." ‘

(b) The metallic element itself (with a small amount of salt dissolved)
forms the second phase. In this group the lower valence state of
the metal is stable only in the molten solution.

The division between 1 and 2 is not clear cut, and both mechanisms
of dissolution may describe a single system. The distinction is between
relatively mobile electrons and electrons which attach themselves to, and
become part of, ions to produce a lower valence state. There are likely
to be intermediate cases not clearly defined, and the attachment of elec-
trons (or "subhalide" formation) may or may not be a matter of degree
only, rather than a matter of a statistical equilibrium between two
distinetly different states of the electron, attached and unattached.
Also, this may vary with temperature and composition. The variety in the
nature of metal solutions will be illustrated in the following by a
systematic consideration of a number of examples. As we are dealing with
mixtures of metals with salts in the liquid state, we shall also have
occasion to consider metal-rich phases and the solubility of salts in
liquid metals, which in many cases is greater than that of the metal in
the molten salt. The discussion will be based largely on measurements of

the phase behavior and the electrical behavior of these systems.

 
 

 

 

 

 

 

"METALIIC" METAL SOLUTIONS

The metal-~metal halide systems which might be termed "metallic"
metal solutions are mainly those of the metals of the first two main
groups of the periodic system and of some of the rare-earth metals.
Farlier literature, that of the last century and the beginning of the
present one, has been reviewed by Cu.bicciotti,1 and a rather complete
review has been published by Ukshe and Bukun.? Many of the earlier re-
sults must be considered as of rather limited value, except for the fact
that they indicated the existence of stable mixtures of metals with
salts, both solid and molten. With a few exceptions, we shall not deal
in any detail with these older data, many of which were relatively in-
accurate, or were given erroneous interpretations. We shall confine our-
selves to the more recent work, which is beginning to lead to a far more
satisfactory explanation of these solutions than was available 20 years

ago.

Alkali-Metal—Alkali-Halide Systems

Phase Diagrams

A study of alkaline-earth-metal—alkaline-earth-halide systems under-
taken under the auspices of the Atomic Fnergy Research Program in the
United States during World War II (the "Manhattan Project')® and the
subsequent investigations of the alkali-metal—alkali~hslide systems begun
in the early 1950's at the Oak Ridge National ILaboratory seem to have
initiated a new period of considerable interest in metal-salt solutions.
Except for rather demanding experimental conditions, the situation was
believedll to be particularly simple with the elements of the first main
group of the periodic system, where complications related to formation
of a lower valence state of the metal were not anticipated. Indeed, the
phase diagrams, in which the alkali metal is one component and one of
its halides the other, are relatively simple (Figs. 1-6). The prinéipal
data of these diagrams are given in Table 1. The temperature range of

C
 

 

 

UNCLASSIFIED
ORNL-LR-DWG, 21894 A

UNCLASSIFIED 69 OLE FRACTION MX
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1200 CRNL-LR-DWG. 38872 1200 |8 7 6 5 4 3 2 .1 1200
| i I [ i [ | | | © COOLING CURVES, THIS WORK
X EQUILIBRATION, MAB JWIWTS
1480° HIGH TEMPERATURE RANGE & EQUILIBRATION, HRE, MB
OF 1100 1080° —1100
- lggfii-lfa'-%&%- 214618 SODIUM METAL-SODIUM HALIDE
aAQ4° 0,0, . COOLING CURVES SYSTEMS
900 EQUILIBRATION AND SAMPLING _| 14150 — — 1000 TWO LIQUIDS —11000
G °C
NaCl - Na
900 900
ool | — 800 800
!
I
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,l NaCl- Na 00 700
'l CsF-Cs
1050 [—/ — 600 600
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10001 | — 1000 |- —1000
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TWO LIQUIDS [/ Two Liauibs
900} 4 —{900
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950~ x — 800 Nal - Na 800
N 740° \
¥ &
700+ {700
, | Ll g SOLID SALT + LIQUID METAL { Lx 656° |
10 20 30 40 50 60 70 80 90 l S0LID ¢ SALT+ LIQUID METAL i
K 900 L L1 b 1 1 1 | sool L1 1 1 11 i Tty Y eoo
mole % 00 20 30 40 50 60 70 B8O 909 4 23456.78.9 1.2.3 45 e 7 s 9
MOLE % SODIUM METAL MOLE FRAGTION OF METAL—> MOLE FRACTION OF METAL —>

Fig. 1. Potassium Metal-Potassium
Halide Systems [J. W. Johnson and M. Fig. 2. Sodium Metal-Sodium Halide Systems, Fig. 3. Sodium Metal—~Sodium Halide Systems [M.

A. Bredig, J. Pbys. Chem. 62, 606 (1958) High-Temperature Range [adapted from M. A. A. Bredig and H. R. Bronstein, . Phys. Chem. 64,
(reprinted by permission of the copyright Bredig and H. R. Bronstein, ]. Phys. Chem. 65 {1960) (reprinted by permission of the copyright
owner, the American Chemical Society)l. 64, 64 (1960} ]. owner, the American Chemical Society)].

 

 
 

UNCLASSIFIED
ORNL-LR-DWG, 310678

I ! I
1300 f— —

 

LiF - Li

12C0

2

HOO

 

 

 

 

 

 

1000 |
E
c |
900 {i— : \,
’ \
800i— -
RbF - Rb =
700 _
CsF - Cs
l I |
8005 25 50 75 100

MOLE % METAL

Fig. 4. Alkali Metal-Alkali Metal
Fluoride Systems [A. S. Dworkin, H
R. Bronstein, and M. A. Bredig, J.
Phys. Chem. 66, 572 (1962) (reprinted
by permission of the copyright owner,
the American Chemical Society)].

UNCLASS FIED
ORNL-LR-DWG. 37011

T 1 T

 

700

€50

600

°C

550

500

 

CESIUM METAL - HALIDE SYSTEMS

 

 

450 ' ' L
0 25 50 75 100

MOLE % METAL

 

Fig. 5.  Cesium Metal-Cesium
Halide Systems [adopted from M. A.
Bredig, H. R. Bronstein, and W. T.
Smith, Jr., J. Am. Chem. Soc. 77,
1454 (1955)].

UNCLASSIFIED
ORNL-LR-DWG, 32909A

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-X-.

i\
N\
/ RbF- Rb ~
0 T N
a
773 \\

750 — —1

706°
700 }— E‘ -Rb —
°C X 696°

—.

 

 

 

\\'\x Rb8Sr-Rb \1
——
650’( T —]
-—x—-x
,\.,,’N /—;bl Rb \
. /x‘
X SIS'
600 — —
550 I | |

 

25 50 75
MOLE % Rb METAL

Fig. 6. Rubidium Metal—Rubidium Halide
Systems [M. A. Bredig and J. W. Johnson,
J. Pbys. Chem. 64, 1900 (1960) (reprinted
by permission of the copyright owner, the

American Chemical Society)].

 

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Al )

 

 

 

 

 

 

Table 1. Principal Data for Alkali-Metal—-Alkali-Haelide Phase Diagrams
Monotectic

Salt- Consolute

Metal Salt Tesp Phase Comp. (mole % M) Eutectic Metal

System  mp ey Tem Com Comp™ mp
MX-M (°x) Salt-Rich Metal- ° § 1 i y  (mole % MX) (°k)
(MX) Rich X (mo e M

Solid  ILiquid  ILiquid

LiF 1121 1120 1b 3 1603 40 452
Licl 883 ggob 0.5b 452
LiBr 823 822b 452
LiT 742 74P 1b 452
NaF 1268 1263 3b 83 1453 28 10™9 370
NaCl 1073 1068 0.15 2.1 97.7 1353 50 370
NaBr 1020 1013 0.15 2.9 96 .6 1299 52 370
Nal 933 930 1.6 98.6 1306 59 370
KF 1131 1122 %49 51L.7 1177 20 3 x 104 337
KC1 1043 1024 0.04¢ 10.5 75,0 1063 39 109 337
KBr 1007 o981 0.03¢ 19.0 69.2 1001 4ty 10™° 337
KI 954, 931 13.5 82.5 990 50 10™° 337
ROF 1068 1046 9b 40b 1063 21 312
RbCl 995 969 18 57 979 37 312
RbBr 965 a a a d 910b,e 44b ye 312
RbI 920 888 22 73 907 51 312
CsF 976 d d d a a d 103 302
CsCl 918 a d a d el a 108 302
CsBr 909 d,a d,a d,a d,a d,a d,a 10™7 302
CsI 899 a d d d a a 107 302

 

b

cEstimated.
By extrapolation from ref 13.

e

SEstimate by extrapolation.

0 miscibility gap.
Unstable.

 

 
 

 

 

special interest is that near the melting points of the salts and above.
The melting points of the alkali metals are very much lower than those

of their salts and are not far above room temperature. The solubility

of the salts in the liquid metals, which is considersble in the higher
temperature range, decreases rapidly with decreasing temperature, so

that the composition of the eutectic liquids in these systems is that of
almost pure alkali metal. The mole fraction of the salt in these eutec-
tics is, with few exceptions, less than 1077, as estimated by extrapole~
fion of the liquidus curves. Although a eutectic of this nature has
sometimes been known under the term "monotectic," this usage will not be
followed here. We shall use this term? to designate the equilibrium
between two liquids and one solid of a composition not intermediate be-
tween those of the two liquids. Thus, in most alkali-metal—alkali-halide
systems the melting point of the salt is lowered by the addition of metal
until the monotectic temperature and composition are reached. Above the
"monotectic horizontal" we £ind the region of coexistence of one solid
and one liquid phase, and the region of the coexistence of two liquids,
one richer in salt, the other richer in metal. In the alkall-metal sys-
tems the compositions of the two liquid phases approach each other
monotonically with increasing temperature until, at the critical solution
or consolute temperature, they equal each other. At and above this tem-
perature, only one liquid phase exists under equilibrium conditions.
Below the monotectic temperature, solid sslt, containing in solid solu-
tion very small amounts of metal (which decrease rapidly with decreasing
temperature),’? is in equilibrium with a solution of the salt in liquid
metal.

The general déscription of the phase diagram given here fits the
1ithium,*% sodium,'+1% and potassium'® metal systems with each of their
halides, all of which, with the exception of the chloride, bromide, and
iodide systems of lithium, are known in detail. Among these three alkali
metals the potassium systems (Fig. 1) exhibit the highest degree of misci-
bility of the salt with the metal in the liquid state. The sodium sys- -
tems (Figs. 2 and 3) are intermediate, and lithium metal (Fig. 4) shows
the least tendency of all alkali-metal systems to mix with its halides

 
 

*

w?

'y

 

in the liquid phase. The critical solution temperature in these systems
represents a qualitative measure of the relative miscibility and of the
relative deviations from "ideal" solution behavior. This consolute tem-
perature is 1330°C in the lithium fluoride system, ranges from 1028 to
1180°C in the sodium systems, and ranges from 717 to 904°C in the potas-
sium systems (Teble 1). The temperature range in which two liquids
coexist is about 530°C in the lithium fluoride system, about 200 to
400°C for sodium systems, and only 20 to 60°C for the potassium systems.
No precise data appear as yet to be available for ternary alkali-metal
systems embodying two or more cations with one anion (besides the elec-
tron), or several anions with one cation.

It is evident that the cesium—cesium-halide systems!? are qualita-
tively different from the majority of the alkali-metal—alkali-halide sys-
tems, in that the liquids are miscible in all proportions (i.e., the
consolute temperature lies below the liquidus line) (Fig. 5). The liqui-
dus curve, depicting the temperatures for the solidification of the ce-
sium halide from the liquid mixture, descends without discontinuity
(except that there ought to be a slight kink at the transformation tem-
perature of CsCl) from the melting point of the pure salt to the eutectic
point, which is almost identical in melting temperature (30°C) and compo-
sition with the pure metal.” The four rubidium phase diasgrams!® show, as
might be expected, a behavior intermediate between the potassium and the
cesium systems (Fig. 6). The temperature range of only partial miscibil-
ity is less than 20°C in the fluoride, chloride, and iodide systems, and
the miscibility gap is absent in the bromide system. It is interesting
to note (Fig. 7) that the trend, with increasing atomic number, size, or
polarizability of the halide ion, toward greater miscibility as expressed
by the value of the critical solution temperature [T, = 1180, 1080, and
1026°C for Na(F, Cl, Br); 904, 790, and 728°C for K(F, Cl, Br); and 790,
706, and < 655°C for Rb(F, Cl, Br) respectively] is, in the iodide systems

 

¥Delimarskii and Markov (ref 18), in Table 29, page 211, "Solubility
of Metals in Fused Salts," misstate the case of the cesium systems, by
erroneously taking the metal concentration of the solution in equilibrium
with solid salt for a metal concentration of a salt-rich phase in equilib-
rium with liquid cesium metal. Many other metal solubilities in this
table are much outdated and quite in error.

 
b 4 e b

 

 

10

Fig. 7. Critical Solution Temperature in Alkali
Metal—Alkali Metal Halide Systems vs Molar Refrac-

tion of the Gaseous Halide lon.

Tes
cs

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Rx (cm3 mole™)

 

 

 

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01 02 03 0.4 05
2
*u

0.6

0.7 08 09 ({0

Fig. 8. Excess Free Energy of Mixing as Function

of Volume Fraction of Metal.

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11

of these three alkali metals, either reversed [Na(I):1033°] or consider-
ably diminished [K(I):708° and Rb(I):634°], with data lacking for the
lithium systems and for Cs-CsBr. Although it seems reasonable that the
solution behavior and T, might be related to the size and polarizability
of the anions, no simple quantitative relationship to these anion proper-
ties has been deduced.

Besides the value of the minimm temperature of complete miscibility
(Te), the ("eritical") composition at the consolute temperature is of
interest. The metal content of the "eritical" solution is seen; in
Table 1, to increase from the fluoride to the iodide for the sodium, po-
tassium, and rubidium systems; the same is true for the compositions of
the inflection on the solidus curve of the cesium systems, which is some-
what representative of the "submerged" (i.e., unstable) critical solu-
tions (28, 47, and 65 mole % cesium metal for CsF, CsCl, and CsI solution
with cesium respectively). It has been pointed out that the correspond-
ing "eritical" volume fractions approach the value 0.50 for all systems,
indicating the quite plausible significance of the molar volume in the
solution behavior.t®

The activity coefficients of the salts, calculated from the liquidus
curves for the precipitation of solid salts from the melts, increase from
unity with increasing metal concentration, corresponding to positive devi-
ations of the salt component from Raoult's law. This is usual in systems
exhibiting (or having a tendency to exhibit) liquid-liquid miscibility
gaps, when no strong ("chemical') interactions such as in the cadmium—
cadmium halide systems (cf. below) take place. In Fig. 8 are plotted
values of RT Iny_ ., Vs ¢§etal for the systems K-KBr, Rb-RbBr, and
Cs-CsI, where ¢metal is the volume fraction of the metal.

The temperature dependence of the solubility of the solid salts in
the liquid metals is another feature of interest. An unusually low value
of the partial molar heat of solution is derived from the temperature
depehdence of the solubility of the solid salts in the liquid metals for
the fluorides of cesium (10 keal/mole) and potaessium (13 kecal/mole) in
comparison with the other halides of these metals.

 
 

 

12

No explanation of these phenomena which is free of arbitrary assump-
tions has been proposed, although rather interesting speculationsils20
lead to a crude but instructive picture of the solution behavibr of these
systems.

Pitzer?l suggests that the excess free energy of mixing is associated
with the conversion of the metallic state of binding of the metal elec-
trons to an ionic type of binding and that the mixing of electrons with
halide ions occurs with little excess free energy, if any. He obtains
fair agreement between experimental excess free energies derived from
the phase diagrams and a theoretical calculation of the difference in
energy of a metallic and ionic model of an alkali metal.

An instructive view msy be obtained according to Blander?? by dividing

the process for the dissolution of metal into the two steps:

M(liquid) — M(gas) ,

M(gas) — M(in dilute solution) .

The first step is simply the vaporization of the metal for which data are
available.?? The thermodynamics related to the second step, which is the
dissolution of the gas molecules, is discussed elsewhere.?* The influ-
ence of the first step may be illustrated by the relative values of RT
1n p given in Table 2, where p is the partial pressure of the metal mono-
mer in the vapor in atmospheres. Aside from differences in the second
process, the influence of the vaporization step would tend toward an in-
creasing relative solubility of metal in salts in the order of increasing
vapor pressure, Li < Na < K < Rb < Cs.

Where Henry's law applies, the second process is best described in
terms of the Henry's law constant, K,,
Ke = csol/cgas ’
where Cool is the concentration of metal in solution in moles/cm® and
cgas is the concentration of metal in the gas phase in the same units.
The free energy of solution (M, = —RT 1n K,) is the sum of the free
energy of forming a hole the size of the metal atom (sbout 10 * 5 kcal/
mole) and the free energy of interaction of the metal atom with the salt.

9

 
 

 

i1

ag

u‘l)

 

13

Table 2. Miscibility and Metal Vaporization Data for the
Alkali-Metal—-Alksli-Halide Systems

 

 

RT In p (monamer) Metal bp Range of Values of Tg
(kcal mole) (OK) _ (oK)
(at 1000°K)
L - 13.8 1604 1603
Na -3.6 1163 1303-1453
K 0.8 1039 981~1177
Rb +0.3 97 9071063
s +0.6 958 (<850 to <950 "submerged")

 

As an example, one might consider the dissolution of sodium in a hypo-
thetical sodium salt having a molar volume of about 50 cm? at the boiling
point (1163°K) to form a 1 mole % solution. These parameters are not

very different from those for real salts. The free-energy change, M.,
for this process, if Henry's law holds even approximately, is ~7 kcal/mole,
and the interaction energy of a metal atom with the solution would have

to be about =17 * 5 kcal/mole. Since the solubilities of sodium are
higher than 1 mole %, the postulate of the dissolution of metal atoms with
weak interaction is seen to be improbable. That the dissolution of alkali
metals in molten salts is Influenced greatly by the same factors influ-
encing the vaporization is shown by the comparison in Table 2 of the boil-
ing points of the metals and the consolute temperatures. This is, of

course, only suggestive, and much more data and a more complete quantita-

- tive analysis are necessary for any definite conclusions.

Physical Properties

Optical Observations. — Aside from the early qualitative findings

 

about the deep coloration of molten alkali halides containing excess
metal, there appears to be, thus far,'only one brief study of a semiquan-
titative nature. Mollwo2? measured the visible and near-infrared absorp-

tion spectra of melts of alkali-halide crystals colored by exposure to

 
 

 

 

 

14

alkali-metal vapor. He found broad sbsorption maxima at 790 mp for all
three sodium salts studied (NaCl, NeBr, and NaIl) and at 980 mu for the
corresponding potassium salts. In contrast to this, the position of the
sbsorption maximum in the solid salts varies with the anion, or the lat-
tice parameter. Mollwo believed that the absorption bands of the melts
could be considered as resonance lines of the metal atoms, displaced and
broadened by intermolecular fields. However, it seems entirely possible
to ascribe the absorption spectrum to an electron whose wave function is
not confined to one metal core (M) but spread out over several M' ions
[e“(Nfi)X] as in an F center, but without the decisive influence, upon its
energy, of a rigid crystal lattice.

Electrical Properties. — In a very few rather crude experiments,

 

Mollwo?? found'no directional transport of the colored cloud from a
melting, F-center colored crystal. He considered this as a confirmation
of the optical finding, which indicated the presence of uncharged (i.e.,
nondissociated) metal atoms.2® This conclusion does not seem to be con-
firmed by later observations.

Bronstein and Bredig?7»28 measured the conductance of alkali-metal—
alkali-metal-salt mixtures and demonstrated very considerable transport
of electricity by electrons in solution. The stainless steel apparatus
for the determination of the specific electrical conductivity, which
they describe in detail, included a synthetic-sapphire capillary dip cell
which was found to be entirely inert to most alkali-metal-halide mixtures.
A1l of the seven sodium and potassium systems investigated (the tempera-
tures of the Na~NaF system were not accessible) had in common a rise in
the specific conductivity of the salt melt on addition of metal.27»28
This rise, however, varied greatly in extent and nature from system to
system. TFigures 9, 10, and 11 are typical examples. In the sodium sys-

tems the rapid initial rise in conductivity is characteristically slowed,

~ and only in cases of sufficient solubility (Na-NaBr and Na-Nal at higher

temperatures) is it finally accelerated as the solubility limit in con-
centration is approached. In the potassium solutions, on the other hand,
the rise is monotonically accelerated with increasing metal concentra=-
tion. In solutions of potassium in KI, a specific conductivity of 600

ohm™% em™t, that is, 400 times that of the pure molten salt, was measured

n

-

 
A

 

15

at 42 mole % potassium, indicating an electronic contribution to the con-
ductivity of this solution of more than 99.5%. The potassium fluoride
solutions are distinguished from those in the iodide by a far slower
initial rise in conductivity, with the chloride and bromide melts being

intermediate. Bronstéin and Bredig define a characteristic or apparent

 

W

&)

equivalent conductance Ny of the dissolved metal by the equation

 

 

 

A - (L —-DN_JA
_ “soln ( M)sfli
My = ’
Ny
ORNL-CR- DWa. 230484
10 L ‘I—fil T I T I T T ! T { T T T - I T 30
9 .
8 28
L 890° .
5
B I'_- 7 —26
™~ -
a 845 24
L Na - NaCl 1
2 122
) —{20
. [760" 118
KB 5700

—10

SPEGIFIC CONDUCTANCE (chmi! cm™)
]
>

SPECIFIC CONDUCTANGE (ohmtem™)

 

 

 

 

 

Q° ]
82 ]K-KCI 8
860° ]
9 16
— 4
5 Na-NaBr 4,
1 4L 1 L 1 | .l I 1 1 I 1 J L I 1 l i O
0 2 4 6 8 0 2 4 6 8 10

SODIUM METAL (mole %) POTASSIUM METAL (mole %)

Fig. 9. Specific Conductivity of Alkali Metal
Solutions in Alkali Halides [H. R. Bronstein and M.
A. Bredig, J. Am. Chem. Soc. 80, 2080 (1958) (re-
printed by permission of the copyright owner, the
American Chemical Society)].

 
 

 

 

 

x, SPECIFIC CONDUCTIVITY, ohm™! cm™!

UNCLASSIFIED

ORNL-LR - DWG. 54649

 

 

 

 

 

16

Fig. 10. Specific Conductivity of Solutions of
Sodium Metal in Sodium lodide [H. R. Bronsteif and
M. A. Bredig, J. Pbys. Chem. 65, 1221 (1961) (re-
printed by permission of the copyright owner, the

16 a T T 1 T T
14( * .
12| -
900°C
10}
8_
800°C
. Ameri ica i .

NoI- Na merican Chemical Society)]

2} ~
| | | | | 1

0 2 4 6 8 10 12 14

MOLE PER CENT SODIUM METAL

20

UNCLASSIFIED
ORNL-LR-DWG. 54650

 

18

- -~ -
N H o

x , SPECIFIC CONDUCTIVITY, ohm !cm!
o

 

_ 1

 

 

l [ [ {

 

4 6 8 10 12

MOLE PER CENT POTASSIUM METAL

Fig. 11. Specific Conductivity of Solutions of
Potassium Metal in Potassium lodide or Fluoride
[H. R. Bronstein and M. A. Bredig, J. Pbys. Chem.
65, 1221 (1961) (reprinted by permission of the copy-
right owner, the American Chemical Society)].

 
 

 

)

ok

%)

 

17

where AL oin and Ajppy BT the equivalent conductances of the solution
and salt, and Ny is the equivalent fraction of metal. The term Ay repre-
sents the change of the conductance caused by the addition of one equiva~
lent of metal to an amount of salt conteined in the solution of the given
composition which is between two electrodé plates 1 cm apart.

Figures 12 and 13 show, in a more striking form than Figs. 9, 10, and
11, the typical difference in the behavior of sodium and potassium. While
in the potassium systems the apparent equivalent conductance of the metal
solute, fys Tises monotonically and with increased rate with inéreasing
metal content, it drops rapidly in the sodium solutions toward a minimum,
which obviously would be followed by a rise if greater solUbility would
permit higher concentrations of metal to be dissolved.

These characteristic changes of Ay, with metal concentration mey be
interpreted as follows: At infinite dilution of the metal, in both the
sodium and potassium solutions, electrons are in a state in which they
can contribute to the carrying of current. Rice,?? assuming a "random
walk" of the electrons, has performed theoretical calculations for this
state at infinite dilution which are correct within an order of magnitude
of the observed conductivity. For the potassium solutions, Rice's theory
predicts an increase of Ay with metal concentration. Bronstein and Bredig
suggest that electron orbital overlap plays a role in this increase, and
doubtlessly this becomes important at higher metal concentrations. How-

ever, in the sodium solutions, another important factor appears to be

- present that actually removes electrons from the conduction process with

.dincreasing metal concentration.';The trapping of electrons in pairs to

form diatomic molecules of Na, waS'éuggested27 to rationalize this dif-
ference. The fact that Nap is relatively more steble in the vapor than

Kz (the heats of dissociation are 17.5 and 11.8 keal/mole respectively)
suggests that this is reasonable. Measurements at higherrtemperature in
the sodiumppontaining systems, 1if this hypothesis were true, would lead

to behavior more like therpotassiumpcontaining systems because of the
greater dissociation of Nas. In the lithium systems a still greater de-
gree of pairing should occur. This,'rather than a short relaxation time,3°
might be the reason for the absence of electron spin resonance in solutions

of lithium in molten lithium iodide.

 
 

EQUIVALENT CONDUCTANCE (ohm™! em?2 equivix 1073)

UNGLASSIFIED

18

ORNL-LR-DWG. 23047A

 

 

 

 

 

 

METAL (mole %)

Fig. 13. Equivalent Conductance of
Alkali Metal Dissolved in Alkali Halides
[H. R. Bronstein and M. A. Bredig, J.
Phys. Chem. 65, 1223 (1961) (reprinted
by permission of the copyright owner,

* the American Chemical Society)l.

 

A, , METAL EQUIVALENT CONDUCTANCE, ohmicm? equivix 1073

n
o

@

o

>

i
n

o

0

L)

o

N

Fig. 12. Equivalent Conductance of
Alkali Metal Dissolved in Alkali Halides
[H. R. Bronstein and M. A. Bredig, J.
Am. Chem. Soc. 80, 2080 (1958) (re-
printed by permission of the copyright
owner, the American Chemical Society)].

UNCLASSIFIED
ORNL-LR-DWG. 54651A

 

 

 

 

i |

1 | ] ] -
2 4 6 8 {0 12 14
- MOLE PER CENT METAL

 

(:)

 

 
 

o

9

N

 

19

Another interesting aspect is the variation of Afl (i.e., Ay extrapo-
lated to infinite dilution) with variation of either the anion or the
metal ion. Table 3 gives a summary of the equivalent molar conductances,
AE, of sodium and potassium metals in infinitely dilute solution in most
of their molten halides. These values were obtained by a short extrapo-
lation, to zero metal concentration, of the curves for fy vs metal con-
centration.

. Two definite trends may be seen in these data: First, the contribu-
tion to the conductivity by the metal solute increases greatly in going
from the fluoride to the iodide systems, that is, with atomic number, or
size, of the halide ion. Second, it decreases in going from sodium to
potassium (i.e., apparently, with increasing atomic numbef, or size, of
the metal ion). The former trend has been attributed?7+28 to the increase
in the polarizability of the halide ion with increasing size, which is
thought to facilitate transmigsion of an electron from one cation, or
group of cations, to another. This is reminiscent of the findings by
Taube and others that the rate constant for exchange of an electron be-
tween a complexed cation such as [Cr(NH3)sX]?t or [Cr(H,0)sX]?* and hy-
drated Ccret ions is very highly dependent on the nature of the ligand
halide ion, X~, presumably on its polarizability. However, while the ef-
fect in aqueous solution shows a very high power dependence on the polari-
zebility of X~ (the bimolecular rate constant for electron exchange be-
tween [Cr(NH;)s5X]2t and Cret changes from 2.7 x 10™% to 5.1 X 10™2 to 0.32
to 5.5 for X = F, Cl, Br, and I respectively), in the metal-salt melts

the apparent dependence of Afi'on polarizability of X is of a lower power.
Quite significantly, the iodide systems exhibit a weakening of the trend
established by the fluoride, chloride, and bromide systems. This is remi-
niscent of a similar change in the trend_for the critical solution temper-
atures mentioned sbove, & similarity which‘may not be merely coincidental.

Of considerable interest would be a discussion of the temperature
dependence of the electronic part of the conduction in infinitely dilute
solution. Unfortunately, the data are, by their very nature (i.e., small
differences of two large numbers divided by a small concentration value),
not nearly accurate enough for this purpose, except to show that the tem-

perature dependence at infinite dilution is very small. At higher metal

 
20

Teble 3. Molar Conductance A;Z of Sodium and Potassium in Infinitely
Dilute Solution in Their Molten Halides at 900°C (ohm™ cm? mole™')

 

Crystal Radius

 

 

 

 

 

 

 

 

 

 

 

of Cation (A) F ¢l Br 1
Ne, 0.95 6000 12,000  (16,000)®
X 1.33 800 2800 6,000 8,100
(820°c)  (870°c)
Molar refraction 2.5 9 12.7 19
of gaseous
anion
(em® mole™?)
Cube of crystal 2452 5.94 7.43 9.95
radius (A?)
a‘Es’t:.:i.mzad:,eél.
UNCLASSIFIED
50 103 ORNL-LR-DWG, 70744
gls aofznec e
o oNaBr o r o B
% o :NuI eKI .RbIr
S
3 20 f
3 /
= 10 ?
V4
x| -3t e
35 4.0 45 4 5 6 4 5 6 7

 

Fig. 14. Electron Mobility in Alkali Halide Crystals
[A. Smakula, Nachr. Ges. Wiss. Géttingen, U 1(4), 55

(1934)].

 

 

 

 

 

 

 

 

 

 

REDUCED TEMPERATURE, 7/

 
 

 

 

o}

»

"

13}

2

 

21

concentrations the temperature coefficient (positive) of the metal solute
conductivity appears to depend both on metal concentration and temperature
(cf. Fig. 13). This might be attributed??+28 to the equilibrium between
associated species such as Nap, and Ko and electrons, 2Kt + 2™ = K,. At
still higher metal contents, where the liquid is essentially a solution
of salt in metal, the temperature coefficient appears to become negative,
as expected for metals.

The very small temperature dependence of the metal solute conductivity
at infinite dilution, A;, significantly distinguishes the melts from the
F-center colored crystals studied by Smakula.3! In the crystals the rela-
tively large activation energy of conductance by F-center electrons is
directly related to the thermal motions of the ions (i.e., the character-
istic temperature of the crystal as derived from the specific heat). Here,
the large differences in the electron mobility for different halide crys-
tals of the same alkali metal are wiped out by the use of a reduced tem-
perature scale, Tabs/e (Fig. 14). A similar relationship using, for ex-
ample, the melting point of the salt obviously (cf. Fig. 13) does not
exist for A&,‘which seems nearly constant with temperature and corresponds
to an electron mobility several orders of magnitude greater than that in
the crystal. It appears, then, that when the thermal motion of the ioms
and the electron mobility are as high as they are in the melt, the re
maining individual differences in Afi may be expected to be due to other
aspects of the conduction mechanism, such as the polarizability of the -
anions which act as negative potential barriers, as discussed above.,

In both the crystals and the melts, the electrons in the sodium sys-
tems exhibit higher mobility than in the potassium systems. This does
not seem to be well understood at present, although a connection with the
smaller interionic distances in the sodium case may exist.

'~ We have seen that it is possiblé to prepare binary mixtures of many
alkali metals with their molten halides in all proportions of the two
components, salt and metal. Compared with the study'of the con&ucfivity
of the salt-rich solutions, only a start has been made in measuring the
conductivity of metal-rich solutions. Figure 15 shows the whole concen-
tration range for several potassium—potassium-halide systems and demon-

strates the large deviation from additive behavior. Bronstein, Dworkin,

 
 

 

 

 

22

UNCLASSIFIED

 

17 — — ORNL-LR-DWG, 56263
I T 1T 1

6 .

15 |-

14
% sl
";x 12~
§ .
e 1 Fig. 15. Specific Conductivity of Solutions of
°~1ol— - Potassium Halides in Potassium Metal [H. R.
ool i Bronstein, A. S. Dworkin, and M. A. Bredig, J.
é Chem. Phys. 37, 677 (1962) (reprinted by permission

|- -

i of the copyright owner, the American Institute of
g 7 | . Physies)].
2 of 1
s 1
n

4% . .

/
3'— ,I —
/
| _
i ”’,/ -
"

 

 

 

1 1 | 1 1 1
0 40 20 30 40 50 & 7O 80 90 100
MOLE PER CENT POTASSIUM

and Bredig3? found that up to 20 mole % KI the resistivity of potassium
metal increased linearly with increasing salt concentration.

Subsequent work>? showed that the proportionality factor & in the
equation for the specific resistivity, p (uohm-cm) = N, + b, increases
with the size of the anion, from 420 for KF (700°), to 740 for KBr (740°),
to 920 pohm-cm for KI (700°). An appropriate explenation appears to be
the increase in the cross section of the larger anions for scattering
moving electrons. An alternative explanation, that of the exclusion of
conducting volume by the introduction of insulator (salt) molecules, is
entirely unsatisfactory: It is true that the equivalent conductivities,
that is, the conductivities of 1 mole of solution of salt in metal, of
KI, KBr, or KF in potassium, fall essentially on the same curve when
plotted against volume fraction. However, the conductivity drops far
more rapidly, initially by a factor of 5 to 10, than the exclusidn of
conducting volume would require, namely, at 700°, from approximately
990,000 ohm™ cm? mole™! for pure potassium metal to 620,000 at a volume

 
 

- (.& I

 

v

i

1

 

23

fraction of potassium of 0.95, that is, (990,000 — 620,000)/0.05 x
990,000 = 370,000/49,500 or 7.5 times faster. It is not unexpected that
the macroscopic picture of exclusion of conducting volume does not work
in this situation where the mean free path of the electrons is large com-
pared with the excluding particle size, which makes the scattering process

the all-important feature.

Alkaline-Earth-Metal—-Halide Systems

Phase Diagrams

The older literature's> contains a rather bewildering variety of
claims for the existence of extensive solubility of alkaline-earth metals
in their molten halides and the existence of subhalides of the formula
MX or MpX, (similar to HgpX,). The assumption of solid alkaline-earth
subhalides, for example, the claim by WShler and Rodewald’?“ regarding

the existence of a s0lid CaCl, based on chemical analysis of deep-red
crystals prepared by fusing together calcium and CaCls, has not been borne

out by later investigations of Eastman, Cubicciotti, and Thurmond©s»33 and
of Bredig and Johnson.2% (Considerable solubility of the metals in their
salts and some solubility of the salts in the liquid metal were found, but
no indication of solid reaction products, subhalides (MK).' A more recent
suggestion®7+38 of a layer~type crystal structure for a preparation pre-
sumed to be "CaCl" turned out soon thereafter>? to be erroneous; the crys-
tals investigated were those of the ternary compound CaCl,*CaH,, or CaClH.
Only & start has asctually been made to obtain truly accurate phase
diagrams of the alkaline~earth-metal-~halide systems. Schifer and Niklas4©
in 1952 very briefly reported a study of the system Ba-BaCl,, which is
characterized by & solubility of 15 mole % barium in molten BaCl, and of
5 mole % BaCl, in liquid barium metal at the monotectic temperature of
878°C, and by a relatively low consolute temperature of 1010°C, only 50°
above the melting_point of the salt. These results, confirmed by Peterson
and Hinkebein*l»4? (Fig. 16), differed greatly from those obtained ear-
lier by Eastman, Cubicciotti, and Thurmond}®»35 which gave little, if any,

 
 

 

 

2

UNCLASSIFIED
ORNL-LR-DWG. 70760

1050 T ! T 1T 1 | 1 T T -
o THERMAL ANALYSIS L
X QUENCHED SAMPLE

 

   

1000

 

 

 

 

 

950 3
a\; J
—~— 1
o 900§ X
2 |
w °
|
E gs50f D
= 1
< 1
s i
o BOOR | -
= |
w {
- 1
750]- atl, :~
¥
| n.o D o O /
700 - a+p a+ly
] 1 i 1 1 ] ! 1 1
0 10 20 30 40 50 60 T0 80 90 100
BaCl,p MOLE PER CENT Ba Ba

Fig. 16. Barium Metal-Barium Chloride System (J. A.
Hinkebein, Ph.D. Thesis, lowa State College, 1958).

indication of complete miscibility at higher temperatures. The difference
is due to faults of the earlier technique in determining the temperature-
concentration area of the coexistence of two liquids by analysis of
quenched melts. As emphasized by Bredig and co-workersil»15:43 gng py
others,4® it is in general impossible to preserve, by quenching, the
equilibrium concentrastion of a liquid phase saturated with respect to
another liquid phase, because of rapid precipitation and segregation of
the second liquid phase. What one may obtain, instead, is a phase re-
sembling the monotectic liquid in composition. Another occasional source
of error is the contraction of the salt phase on freezing, which produces
a sump hole and/or cracks which fill with excess liquid metal. To avoid
these errors it is necessary to resort to entirely different techniques
such as thermal or differential thermal analysis, that is, the taking of
cooling curves, as Schi8fer and Niklas did, or to & process of separating
a sample of each liquid phase at the equilibration temperature as de-
scribed by Bronstein and Bredig.?7? A still different method, which relies
on diffusion of the metal into the salt contained in a separate compart-

ment of a capsule,%* is subject to some doubt, as the wetting character-

- istics of the salt and metal and the creeping of liquid metal along the

 
 

 

i
1
§
E

g

b

Y

b

 

 

25

inside walls of the metal containers may also lead to distorted results,
aside from the inconvenience of long periods of time for equilibration.

Besides the measurements on the Ba-BaCl, system, the results are
believed to be reasonably accurate for the Ca~CaF2 system reported by
Rogers, Tomlinson, and Richardson;45 for the Ca-CaCl, system reported by
Hinkebein and Peterson“l:42 and by others;%4»%% and for the higher tem-
perature range of the Ca-CaBr, and Ca-Cal, reported by Staffansson,
Tomlinson, and Richardson.#

A low solubility of magnesium in molten MgCl, was observed, 0.90
(Zhurin*7) and 0.30 mole % (Rogers, Tomlinson, and Richardson%’), both
at 800°C. A comparison (cf. Tables 1 and 4) of the solubility of the
various alkaline-ecarth metals in their molten chlorides near the melting
points of the salts reveals a trend somewhat similar to that found in
the alkali-metal systems, that is, a rapid increase in solubility with
atomic number of the metal; Mg-MgClp, < 0.5; Ca-CaClpy, 2.7; Sr-SrCl,,
5.5; Ba=BaCls, 20 mole %.

In the Ca-CaF,; system a rather high miscibility, that is, a rather
small temperature range (30°) for coexistence of two liquids, was found,45
with a critical solution temperature of 1322°C, almost 100° below the
melting point of the salt (Fig. 17). The critical solution temperatures
in the other three calcium-halide systems are very similar, yet the tem-
perature range of only partial miscibility in these systems covers 500°.
This, of course, is caused by the relatively much lower melting points
of CaCl,, CaBr,, and Cal, (772, 742, and 780°C respectively), as compared
with 1418° for CaF, with an entirely different crystal structure. From
a comparison of the melting-point depréssion of CaF, produced by calcium
with that produced by NaF, Rogers, Tomlinson, and Richardson“® concluded
that the metal dissolved either as calcium atoms or as Ca22+ molecule
ions, nemely, according to either Ca — Ca® or Ca + Ca?+ — (Cap?*, both
giving a cryoscoplc number, n, of unlty, or approx1mately unity, in the
Raoult—van?t Hoff formula AT = (T, - 1T) = (Rmz/m%)-nng , where N, 1s the
mole fraction of calcium metal. In_both cases, solid solution was as-
sumed not to oceur, and it was concluded that the reaction Ca + Cat —
2Cat did not take place. However, in this case, the AH in the formula

 
Teble 4. Principal Data for Some Alkaline-Earth-Metal—Halide Systems

 

 

 

 

 

 

 

 

, Monotectic
Salt- -
g;::im Sz;t Phase Comp (mole % M) Eutectic Consolute Mzgal
}&2{ j (°K) ?3%) Salt-Rich M;ii-,z%l_ ?f;l)’ (moggm% ) 'J(?g%a (mogzmi . (°x)
Solid Liquid ILiquid
MgCl, 987 ~987 ? 0.2 ~923  ~100 923
CaFa 1691 1563 & 25.5 67 1094 98.6 1595 % 5 45 1110
caCl, 1045 1093 2.70 99,5 1033 2.0 1610 + 5 62 1110 o
CeBr, 1015 1100 2.3 99.6 1000 3.02 1610+ 5 64 1110
Cals 1053 1104 3.8P 99,7 1033 2.0 7 1650 £ 5 74 1110
srcl, 1145 1112 8 5.5 ? ? ? ? ? 1044
BaCl, 1235 1163 c 15.0 95 985 ~99 1290 50 1002
a

yeonsiderable metal solubility in (fluorite type of) solid is likely.

See ref 46,

Cconsidereble metal solubility in high-temperature crystal form of BaCl, (fluorite type ?) is
likely.

 
 

 

 

¥

b

-

N

1

 

27

UNCLASSIFIED
ORNL- LR-DWG. 70746
T ¥

T

 

     
 

—

1300

 

 

 

 

 

S 1200
Fig. 17. Calcium Metal-Calcium Fluoride System §
(P. S. Rogers, J. W. Tomlinson, and F. D. Richard- E 1100
son, p 919 in Physical Chemistry of Process Metal- &
. = o DIFFERENTIAL
lurgy, Interscience, New York, 1961). M 000\~  THERMAL ANALYSIS
0 SOLUBILITY
200 |-
gool
o 100

COMPOSITION (mole % Ca)

would become approximately 10 kcal, quite different from the calorimetric
value of 7.1 (Naylor48), which is hardly in doubt. It seems, then, more
reasonable to assume solid solubility actually to occur both with NaF and
Ce (as CaF??). Mollwo,4° in measurements on F centers, found the solu-
bility of caleium in solid CaF, at 1300°C to be approximately 30 mole %,
which might well be too large by as much as a factor of 10. With the
value of 7.1 kecal for A , the modified formula

A7 = B2 1[N, (11q) — N, (solid)]
ANy - |

would leave the possibility open for a cryoscopic n larger than 1,
possibly 2; at low metal concentration; This cryoscopic n = 2 would be

in agreement with the findings in the éorresponding chloride system,

CaCl,-Ca, where Dworkin, Bronstein, and Bredig3® found n also to be only
slightly less than 2. Here,'howevér, no solid solution was assumed, the
crystal structure being'quite different, with sixfold rather than eight-
fold coordination of snions around cations. ~In both cases the reaction
would then seem to be Ca — Ca?¥ + 2¢~, as also suggested by the elec-

trical conductivity (discussed below), with a possible, though much less

 
 

 

 

28

plausible, additional dissolution mechanism, Ca + Ca?t — 2¢a*, not ex-
cluded at this stage, and with increasing formation of (Cap)?* at higher
metal concentration.

In the Sr-SrCl, system, Staffanssont* found n closely equal to 1,
assuming a heat of fusion of 4.1 + 0.6 kcal. Dworkin, Bronstein, and
Bredig®? questioned this latter heat value as possibly being too low.
However, a recent calorimetric determination (Dworkin and Bredig®*) essen-

tially confirmed the low value, which actually gives fair agreement in the

rather low entropies of fusion of CaF, and SrCl, (4.2 and 3.4 e.u.) of
similar, fluorite-type crystal structure. Mollwo found the solubility of
strontium in solid SrCl, to be high, similar to that of calcium in CaF,.4°
As in the case of CaFjy-Ca, the existence of solid solution in the fluorite
type of structure could lead, then, to an interpretation of the melting-
point depression in terms of n ® 2, corresponding to Sr — Srét — 2e”
(and/or, less likely, Sr + Sr3t — 2sr¥),

Solid solution has also been reported*®s4* for barium in the high-
temperature form of BaCl,, most probably also possessing a fluorite type

of structure. For

$m-— T AT 43
= ~ = 1100°C (observed experimentally4l) 3
No(£) — Na(e) AN, 0.04

 

JA'S) AT 4 x 1100
m ~
n= X

~ N1.8% 2.
RT AN, 2 x 1233

 

Again, a solution mechanism corresponding either to Ba — Balt + 2e™
and/or (less likely) Ba + Ba?t — 2Bat seems indicated by these experi-
mental data. On the basis of activity measurements carried out by equi-
libration of MgCl, with MgAl alloys of known activity, Rogers, Tomlinson,
and Richardson*’ suggest for Mg-MgCl, the solution mechanism Mg + Mgt —
(Mg2)2+. The precision attained does not exclude the alternate mecha-
nisms Mg — Mg?* + 2¢~, or Mg + Mgt — 2Mg*, and the data appear to fit
best a mixed mechanism, or in other words, an equilibrium 2~ + 2Mg*t+ =
(Mgs )%+, or 2Mg* = (Mgy)?t, discussed more generally below.

 
 

b2l

4

 

29

Electrical Conductance

Dworkin, Bronstein, and Bredig®?:52:53 found it impossible to apply
the method used for measuring conductivity in the alkali-metal systems to
the alkaline-earth systems. While the components, molten salt or liquid
metal, were found singly not to react extensively, if at all, with the
synthetic sapphire cell, their mixtures did attack the latter readily.
Recourse was taken to an all-molybdenum-metal apparatus embodying an as-
sembly of two parallel rods as electrodes which were immersed into the
melt from gbove to a varying, accurately determined depth. With pure
molten salts the low resistance of this cell led to large effects due to
electrochemical polarization of the electrodes and resulted in a large
frequency dependence of the measured resistance which could not be ex-
trapolated to infinitely high frequency with any accuracy at all. How-
ever, the addition of even small amounts of metal, acting as a depolar-
izer, removed the frequency dependence. Thus, it was possible to use the
apparatus with the metal solutions and to calibrate it with solutions of
cadmium in cadmium chloride. The specific conductivity of these solu-
tions is known from the early measurements of Aten in 1910 (ref 54) and
was confirmed by new measurements with the sapphire cell,3?

Only results for solutions of calcium and strontium in their diha-
lides, excepting the fluoride, have been reported thus far.46:30 It was
concluded, as far as the limited data permit, that calcium behaves some-
what like sodium in that the rate of increase in specific conductivity
with metal concentration decreases (Fig. 18). The same is true for
strontium in SrBr, at the relatively low temperature of 700°C, but in the
dichloride, at 910°C, the behavior of strontium, giving a linear increase
of conductance with metal concentration, lies somewhere between that of
sodium and potaessium, with barium expected to resemble the latter. The
decreasing equivalent conductance of the calcium-metal solute was attri-
buted to the formation, not of diatomic molecules, Ca,, as with sodium,
but of hypothetical molecule ions, (Ca2)2+, in which the trap for elec-
trons is the single two-electron bond similar to that in Nap, or in Cdy?*
and Hg,?t.

 
 

A1 AR R et A el et e e bttt B Y b 1t

 

 

30

UNCLASSIFIED
ORNL-LR-DWG.70122

 

3.4 ]

3.2

3.0

   
 
  
   
      
  
 

o
™

o
>

Fig. 18. Specific Conductivity of Alka-
line-Earth Metal Solutions in Alkaline-
Earth Halides (partially unpublished data).

o
b

n
M

D
©

 
 

”
// Sr-SrBr, 700°

SPECIFIC CONDUCTIVITY, ohm™cm’
©

7’
7%
+
/ +/D/
/ /O)/Cu-cuz 830° _

o4

1.6

 

 

) l l
0 1 2 3
MOLE PER CENT METAL

From both the melting-point depression, above, and the decrease in
metal solute conductance from 800 to 450 ohm™' cm? equiv™', it would seem
to follow that at saturation (approximately 3 mole %) somewhat less than
half of the dissolved calcium metal is associated to Cay?t.

More information is certainly required to put these ideas about the
constitution of the solutions of the alkaline-earth metals in their di-
halides on a firm basis, and especially to exclude with certainty the

existence of Mt alkaline-earth ions in limited concentrations.
Rare-EFarth-Metal-Trihalide Systems

The systems comprising rare-earth metals with their trihalides (see
Table 5) represent an especially interesting and more complex group of
metal-metal-halide systems, as they occupy an intermediate position.
They are farther than the alkaline-earth systems along the line leading
from the extreme of the alkali-metal solutions, with metallic properties

and with no solid subhalide formation, to the systems of the transition

L7

 
 

 

Table 5. Principsal Data for Some Rare-Earth-Metal—Halide Systems

 

 

 

= MX3; mp Depression ‘Butectic "Subhalides” mp (°K) SOlubIEjty 10 Metal mp
o
MX3 Melting Asm(a) . Ja) () ) (Congr.:C; incongr.:i) (°k)
Point (dT dn 0 n Conp Temp 1 1 1t " " n n n Mole % Temp
(°x) (eu) w (wole % M) (%K) Fo.us)' Mpaz" MKy 5" MK poin" (ex) Pwre  HEK
LaCls = 1131 1.5 470 2.4 2.0 1099 None None None None 12.0 1293 1193 1187
LaBrs(d) 2061 (11.8)(e) 375 (2.4) 14.5 1001 None None None None 15.5 1173 1193
Lals 1052 (12.0) 490 (2.8) 8.2 1007 10231 None None 1103¢ 33 1173 1193 1173
CeCls 1090 1.7 480 2.6 9.0 1050 None None None None 9 .73 1068
ceBrp'd) 1005  (11.8) 345 (2.0) 12.0 960  None None None  Nome 14 1173 1068
Cels 1034 12.0 460 2.7 8.8 o288 10044 None None 10811 32 1173 1068 1064
PrCls 1059 11l.4 565 3.0 17.0 919 None 9321 None None 19 1073 1208
prar, (3 966 (12.0) 545 3.4 16.0 852 8744 None None  Nome 18 1023 1208
Pri, 1011 12.6 500 3.1 11.9 939 946C None None 10311 29 1073 1208 1180
N4AC1, 1032 - 11.6 590 3.3 14.0 913 None 9531 9751 11141 31 1173 1297
NdBrs 955 (12.0) 1297
NAI; 1060 9.2 740 3.2 26.5 764 None None None 835¢ 37 1073 1297
car, () 1204 (12 7) 555 (2.8 2) 14.0 1098 None None None 11041 14 1173 1585
YI, 1270 (10 ?) 390 (1.5 22) 12.0 1221 None None None None 15 1423 1782

 

(a)A. S. Dworkin and M. A, Bredig, J. Phys. Chem., March 1963.

(v)

{e) _

n = (ar,/aN,), (45 /RT ).
(gee ree 113.
e)Figures in parentheses are estimstes.

Liquidus slope at infinite dilution, mostly from work by Corbett et al.

¢

 

4]

 

 
 

32

and posttransition elements, where metal-salt miscibility in genersl be- kfifi
comes more limited, where electrical conductivity of the solutions gives
no indication of metallic character, and where solid "subhalides" do

occur.

Phase Diagrams

An early phase diagram by Cubicciotti®? of the system Ce-CeCl; in~
dicating a solubility in CeClas as high as 33 mole % cerium metal was
later’® found to be in error: A solubility value of 9 mole % cerium,
apparently rather constent between 780 and 880°C, was obtained by Mellors
and Senderoff,’® and confirmed by measurements of Bronstein, Dworkin, and
Bredig53 at 855°, On the metal side, the phase diagram proposed by
Mellors and Senderoff seems to require further investigation, as both the
true melting point of cerium metal (817°C) and & transition occurring in
the solid state (730°C) were not taken into account.

The phase diagram of the corresponding lanthanum system, as deter-
mined by Keneshea and Cubicciotti®? (Fig. 19), gives a solubility of the
metal in the salt very similar to that of cerium. No intermediate solid
phases, that is, no solid compounds containing cerium or lanthanum in a
valence state lower than 3, were observed., In this respect, then, these
two systems resemble the alkaline-earth systems. Also, in both cases, it
appears from the melting-point depressions of the trichlorides that n,
the apparent number of new particles produced on dissolving one atom of -
metal, is approximately 3 or slightly less, perhaps 2.5. This signifi-
cantly rules out the smaller values, n = 1 for the dissolution of these
metals as atoms and n = 1.5 for Mt ions®® formed according to 2Ce +
CeCl; — 3CeCl or 2Ce + 4CeCls — 3Ce(CeCl,), unless no or few (CeCl,)™

- complex ions were present in pure CeCls, so that (CeCl,)™ ions also would

act as new particles, in which case n would be approximately 3. No com-

pounds of monovalent rare-earth elements seem ever to have been reported,

even though some of these elements have long been known in the divalent

state, notably Yb2+, Eu?*, and sm?+, with 14, 7, and 6 electrons in the

4f shell, that is, with the 4f shell full, half-filled, or nearly half- )
filled respectively. We shall, however, see below that an equation Q-J

 
 

 

 

¥

 

33

UNCLASSIFIED
ORNL-LR-DWG. 70745

 

 

 

 

 

 

 

1000 T : T T I T I T
L, Lo+ L, Ly
S147 — fsigopo-aatk]
o
© 9001 e % ° 4
2z L+7-Lo(sS) |
w ° i
860
E st
o o
<t L+ -La
« ~ *R 826°
; -; . = d%: ° ° o o o ]
L + LaCl °
i 800 ! 3 LaCl, + 8 - Lo o oo
e o
o
° THERMAL ANALYSIS %
e SALT- PHASE SAMPLED
o QUENCHED
700 [ | | ! L1 | 1 |
o 10 20 30 40 50 60 70 80 90 100

MCLE PER CENT Lo

Fig. 19. Lanthanum Metal-Lanthanum Trichloride System [F. J.
Keneshea and D. Cubicciotti, J. Chem. Eng. Data 6, 507 (1961) (copy-
right 1961 by the American Chemical Society and reprinted by permission
of the copyright owner)].

Ce + 2¢e?t — 3ce?t (cryoscopic n = 3) would not seem to do justice to
the results of the measurements of electrical conductance, which suggest
Ce — Ce3t + 3¢, .

Extensive, careful work on rare~earth-metal-halide phase diagrams
is being done by Corbett and associates.’®"®3 gtable, solid compounds
corresponding to, or at least approximating, the composition MX,; were
found to exist in the systems Nd-NdCl;, Nd-NdLz, Pr-PrIs;, Ce-Cels, and
La-Lalj, besides distinct solid compounds intermediate between MX; and
MX3; such as Nd012,37, Nd012,27, Pris,s, Cels,,, and La12,42 (Figs. 20-23).
The investigators appear to have had reasons for accepting the latter,
nonstoichiometric compositions in preference to simple ones such as
NACLl32NdCls, NACls<3NdCl,, and MI3+MI, (M = Pr, Ce, or La). Some of
these compounds melt congruently, others incongruently. Significantly,
the solid diiodides of lanthanum, cerium, and praseodymium are charac-
terized by melting points higher than those of the triiodides (Figs. 22
and 23) and, above all, by high metal-like electrical conductivity
(>1 x 103 ohm™?! cmfl), unigue among halides, except perhaps for Ag,F.
solid NdI,,gs, on the other hand, has a different crystal structure, is

 
 

 

 

 

 

   

 

 

 

 

34
UNCLASSIFIED
ORNL~LR-DWG. 56845
500 l | T 1 1
®
] = u.fi-t
ST
> !
800}~ / ¢ —
o
158' + L
T‘C. \‘ * » ,.’ @
L ]
] 102
7m_ Q’.Tu’—-_— p—
680%*1
S —pr—
e |
- * ' ” <+ oft‘ ’l “dClu
S i
&
€00~ Nd ! Nd Gl
o |
! ' &
| | i 14 | 1 )
Nd CI3 5 o 15 20 25 30 as

MOLE % Nd N NdCls

Fig. 20. Neodymium Metal~Neodymium Trichloride System
(thermal analysis: +, equilibration: @) [L. F. Druding and J. D.
Corbett, J. Am. Chem. Soc. 83, 2462 (1961) (reprinted by per-
mission of the copyright owner, the American Chemical Society)].

UNCLASSIFIED
ORNL- LR-DWG, 56816

 

 

 

 

 

 

 

 

T l v ‘ ¥ j’ -7 I v l
800 1
« 787
750 .
J Fig. 21. Neodymium Metal-Neo-
Too— dymium Triiodide System (thermal
4 analysis: +, equilibration: o) (L. F.
50— . - Druding and J. D. Corbett, J. Am.
| Chem. Soc. 83, 2462 (1961)(reprinted
TE.’gP' 600 " . by permission of the copyright owner,
1 (562 [Nd  the American Chemical Society)].
550|— "]
L 4 .
500~ 49 . Yy A -
1." T T¥ e+ T T¥Y ¥V Nl
450 Ndll; o5
o ey i
NdL, 8 16 24 3z 40

MOLE% Nd IN NdI,
 

 

1]

"

 

35

UNCLASSIFIED
ORNL-LR-DWG, 70759

 

 

 

 

 

 

I | [ I | | | I |
880 —
+ THERMAL ANALYSIS ;Tr
860~ o EQUILIBRATION 1 7
saol x LITERATURE oo *—*‘i
T 1

~ 820 + t o
9 / o
~— ++ 0

800}~ _
x J
-
2 780 y -
x £
Lk £
AN o
o 760 *, —
= +, -

\ X)Atq-‘_—‘_-lfi-i-_—_
740 |- + + .4 -]
—*—%*fl .
720 —
° I“1'12.42 L"Iz.oo
| | | | | | | | | | 2l
7000 4 8 12 16 20 24 28 32 36 ‘96 100
L0I3 MOLE PER CENT La IN LuI:5 La

Fig. 22. Lonthanum Metal-Lanthanum Triiodide System [J.
D. Corbett et al.,, Discussions Faraday Soc. 32, 81 (1961) (re-
printed by permission of the copyright owner, the Faraday
Society)].

UNCLASSIFIED
ORNL-LR-DWG. 70758

A
v

 

 

 

 

 

 

 

820 — | | | | ] [ I ! | ]
8001~ + THERMAL ANALYSIS
© EQUILIBRATION
780 x LITERATURE
760}
S 7404
Lt \
T 720 * + -
= * /
w 700t ' o —
s N\ F
i T,
- 680 \ g / . | .
- : :t:: + +
—-—_‘.'+'h#—+";+"-# #H‘T"Z—";T + ¥ +
6601 +* W —
6401 Pris s Priz.0, _
| | | II I 1 | | ] | | ,
o 4 8 12 16 20 24 28 32 36 40 100

PrI, MOLE PER GENT Pr IN Prl, Pr

Fig. 23. Praseodymi.um Metal ~Praseodymium Triiodide
System {J. D. Corbett et al., Discussions Faraday Soc. 32, 81
(1961) (reprinted by permission of the copyright owner, the
Faraday Society)].

 
 

 

 

36

seltlike, nonconducting, and melts 225° below NdIj (Fig. 21). The prop-
erties of the metal-like diiodides were interpreted as being due to the
presence of "metallic" lattice electrons plus tripositive metal cations,
M3*te~(I~),, while the saltlike NdCl, is simply NG**+(C17),. This inter-

- pretation is further supported by the relative weakness of the paramag-

netism of solid LaI, with a susceptibility of 220 x 10™° emu/mole, in-

‘dicating the absence of La?t ions.

The solubility of the metal in the triiodide at 830°, that is,
above the melting point of any diiodide, decreases from 37 mole %
in NdI; (Fig. 21) to 29 in Pris (Fig. 23) and 30 in Cels, but rises again
to 33 mole % in Lal; (Fig. 22). A similar deviation from the trend is
shown by lanthanum in LaCls; with a solubility of 10.3 mole % metal at
900° (Fig. 19), much higher then expected in the series: 30 mole % in
NaC1l; (Fig. 20), 20 in PrCls;,5t»5% and 9 in CeCl;.’>?»?® The reason for
this "anomaly" with lanthanum is not quite clear; it may, however, be
remembered that La?¥, like Sc3* and Ce*t, possesses, in contrast to the
other rare-earth ions, a noble-gas type of electronic shell.

Examination of the liquidus curves for the metal-like diiodides
(Figs. 22 and 23) confirms the high degree of dissociation of molten
Lal, into La?t ions: The depression of the melting point by the addition
of Lal; (Fig. 22) is only of the order of 0,50° per mole % MX3, corre-
sponding, with an estimated heat of fusion of 10 kcal/mole, to & cryo-
scopic n, that is, number of particles per triiodide molecule dissolved,
of only 0.2. For CeI,®% these numbers are larger, namely, 1.16° per
mole % Cel; and n = 0.5, and are still slightly larger for PrI, (Fig.
23): 1.4° per mole % PrI; and n = 0.6, indicating less dissociation ac-
cording to M** — M?* + €=, The apparent alignment of Cel, with PrI,
rather than with Lal, again is of interest and not fully explained.

The increase in the stability of the M2t rare-earth ion with in-

creasing atomic number leads to a sharp maximum with Eu2+, which contains

seven 4f electrons, that is, half of the full 4f shell, and is more

stable than Eu?¥ even in aqueous systems. This stability, which is very
low with gadolinium, perhaps gradually rises again in going toward the
well-established Yb?¥, which contains the complete set of 14 4f electrons.

1]}

 
 

e s e i

 

h

n

 

37

Electrical Conductivity

The first measurements of the electrical conductance of solutions
of cerium in cerium trichloride, by Mellors and Senderoff,58 indicating
a rapid rise in conductivity up to 0.65 mole % metal followed suddenly by
a very slow rise above that concentration, were shown by Bronstein,
Dworkin, and Bredig®?:?3 to be in error because of reaction of the solu-
tions with ceramic crucible material. (The results of a study of emf and
their interpretation in terms of a Ce¥ ion® must have been equally af=-
fected.) Actually, the conductivity in the Ce-CeCl; system, as in La-
LaCls,”® rises with a monotonically increasing rate to a value, in the
saturated metal solution, approximately five times that of the salt. In
the neodymium system this ratio amounts to only 1.7, and in the praseo-
dymium system it is intermediate, 2.5.%® A similar relationship holds

for the equivalent conductance of these solutions (Fig. 24) obtained with

UNCLASSIFIED
ORNL-LR -DWG. 70125

 

 

 

 

 

B | T T r
]
1 | ]
|
i
10} | _
fi,otos
. 9 |K'8 —
IE ,
S | B
e 8 i
s i L
Fig. 24. Specific Conductivity of Rare Earth - 7I- ; .
Metal-Rare Earth Metal Chloride Systems. S 'l Lo, 910°
5 e ! ]
=2
O
& —
O
Q
T .
2
o ————
m —r
Pr, 830° ,
Nd, 856°
] ] ] | ]
5 10 15 20 25 30

MOLE PER CENT METAL

 
 

 

38

estimates of the equivalent volumes on the basis of the density measure- \EJ
ments of Mellors and Senderoff in the CeCli-Ce system.’® The results

obtained with the corresponding iodide system367 are not very different;

however, the increases from the pure salts to solutions containing, for ex-

ample, 10 mole % metal, correspond to factors as high as 15, 19, 6, and 2

for the La, Ce, Pr, and Nd systems respectively. These findings may be ex-

plained in terms of the systematics of the stability of the divalent rare-

earth cation, M*¥, as observed in the phase diagram studies above. Where

no solid compound of the composition MXp; is found, or where such a solid

exhibits metallic character,?®»60 conductivity of the molten solution is ]
high, indicating the equilibrium M?**+ = M3* + e~ to be far to the right.
Praseodymium, forming no PrX, in the chloride system, but & solid phase -

PrClp, s of mixed valency which is stable in a very limited temperature
range, % and forming, in the iodide system, a diiodide, not saltlike
but metallic in character,®? also occupies an intermediate position be-

tween cerium and neodymium as far as conductivity is concerned.

UNGLASSIFIED
ORNL~LR-DWG. 70423

 

 

 

 

TT 1T ] I I T
2 | T —
|
;'
" — i -
}
10 ! -
] N
- N )
E I
L !
e s ~ Fig. 25. Specific Conductivity in Rare - |
2 A ,’ | Earth Metal—-Rare Earth Metal lodide Systems |
& ,’goo% [R. A. Sallach et al., J. Pbys. Chem., in press ‘
£l A (1963) (reprinted by permission of the copyright
3 ! owner, the American Chemical Society)].
Ss | ol /b .
o 840° 780°
= /
o 4 | -~
W /
) !
3l -
i
’ -
2;" ' Nd -
820°
= — . ]
——././ g
|1 Loy | -
0 5 10 15 20 25 30 35 (‘i')

MOLE PER CENT METAL

 
 

 

"

u

 

39

Of special interest is the maximum in the curve of conductivity vs
concentration for solutions of neodymium in NdIs at approximately 23
mole % neodymium, which corresponds to a maximum deviation from addi-
tivity of the conductivity of mixtures of NdIs and NdI, at 16.7 mole %
neodymium metal, corresponding to 50 mole % NAI, (Fig. 25). Bredig

et al.%7 interpret this maximum in terms of an electron exchange between

 

Nd** and Nd** for which the probability has a meximum at a concentration
ratio of NdI,:NdIs = l:l, since both these salts as pure liquids conduct
only ionically. The same effect is thought to cause the curvature of
the conductivity~concentration curve in the NACl3-NdCl, system,50 but,
probably because of the smaller polarizability of C1™ compared with I,
the rate of the electron exchange i1s lower and a maximum is not produced.
In any comparison of the conductance observed in the rare-earth-
metal—halide systems, with both the alkaline-earth and alkali-metal
systems, it is reasonable to meke the correlation for the equivalent
volume, that is, a volume of solution containing the same ratio of the
number of obstacles affecting electronic conduction, namely, the number
of anions to the number of metal electrons. However, the number of
cations in such a volume decreases in going from the alkali-metal halides
to the rare-earth-metal halides by a factor of 3. Now, according to
Bredig,68 the electronic conductivity, that is, the contribution of the
metal in the dilute state, may be thought of as being proportional both
to the total number of cations and to the number of valence electrons
introduced by the addition of metai, that is, to the concentration of
the two components making up a metal. The product of these two quanti-
ties 1s, then, thought to be a significant parameter. It equals the
product of the equivalent or mole fraction of the metal, NM’ representing
the concentration of electrons from M = M?* 4+ ze™, and the reciprocal of
the valency of M, 1/z, representing the concentration of MZt ions.
Figures 26 and 27 show the contribution of the metal solute per equivalent
of mixture, AM'NM,= Ayorn — (lf— NM)ABalt’ plotted against NM/Z, for sev-
eral chloride systems. It is clear that this simple correlation can be
expected to be significant only for those few MXZwM systems in which all
M atoms react according to M — M?* + ze™, or, in other words, where there

is no or little complication by the trapping of electrons in pairs such

 
 

 

 

 

 

 

 

 

 

 

 

 

 

UNCLASSIFIED
ORNL-LR-DWG, 66930
20 400 T T 1 T
18
16
300 -
14 T,
S
‘T -
> o
2>
z 12 *5
N .
5 £ KCI-K
T 10 ° 200
. $
T8 %
- '4:
s LoCI3-Lu
< g )
100 CeCIB-Ce i
4
SrCl,-Sr
2 LaCly-La . 2
e -
CeCiz-Ce —SrC12SY, | 1 L !
004 008 042 046 001 002 003 004 005
N, [e1x(M*] = N, /2

Fig. 26. Electronic Conductance in Some Metal-Metal
Chloride Systems. Left: Equivalent Conductance of Metal
vs Mole Fraction of Metal. Right: Electronic Contribution
to Equivalent Conductance of Solution vs Product of Elec-
tron and Cation Concentrations [M. A. Brédig, J. Chem.
Phys. 37, 914 (1962) (reprinted by permission of the copy-
right owner, the American Institute of Physics)].

UNCLASSIFIED
ORNL-LR-DWG, 66931

 

 

 

 

 

   

 

 

 

 

 

400 T T 77
300 -
T.Z -
g 2
o 3
Q o
T £
£
5 7, 200 -
7 g
< ?: KI-K Loly-La
. =
100 Ce13—0e —
4l Cels-Ce
Lolx-La
2 P
1 o Lt o | 1
004 008 O{2 016 001 002 003 004
Ny leXM*1= m,, /2
Fig. 27. Electrenic Conductance in Ce, La, and K

Metal—Metal lodide Systems. Left: Equivalent Conductance
of Metal vs Mole Fraction of Metal. Right: Electronic Con-
fribution to Equivalent Conductance of Solution vs Product of
Electron and Cation Concentrations (M. A. Bredig).
 

 

 

"

»n

 

41

as in molecules (e.g., Na,), in molecule ions (Ca,?*), or singly in lower
valence states such as Pr?t, guitable systems appear to be KCl-K, SrCl,-
Sr, and La-LaCl; or Ce-CeClz. Figure 26 shows La-LaCls; and Ce-CeCls to
have indeed a dependence of AM-NM on NM/z rather similar to that of K-KCl.
It is perhaps possible to ascribe the fact that the curve for Sr-SrCl, is
considerably below the other ones to the trapping of electrons in (Sr2)2+
molecule ions. However, the attempt to interpret these finer details is
not really satisfactory, as is illustrated by Fig. 27, where AM'NM for
CeI3~Ce first deviates negatively and thereafter positively from AM'NM
for KI-K. An explanation does not seem available,

Similarly, the fact that the rise in fyy with the metal concentration
is much faster for cerium in CeIs; (ref 69) than in CeCl; remains unex-
plained, except that it possibly has something to do with the stability
of the metallically conducting solid Cel, and the nonexistence of a

similar solid CeCls.
"NONMETALLIC" METAL SOLUTIONS

The term used here to classify a second large group of solutions of

metals, essentially the heavy metals, in their molten halides is based on
the observation, made on a number of such systems, that the electrical

conductivity of the melt is not increased, if altered at all, on addition
of metal. In contrast to this viewpoint about the nommetallic character
of these solutions, Cubicciotti’ has pointed to the metallic luster which
deeply colored (black) solutions of bismuth in molten BiCl; reportedly
exhibit. The significance of this apparent discrepancy is not clear, and
we shall proceed with the descriptlon of this second class of solutions,
bearing in mind that further study may tend to blur the sharp line of
separation found to be a convenience at this time.

There is, however, one more criterion, which rather sharply sepa-
rates many of the systems to be discussed helow from the solutions of the
electropositive metals dealt with on the preceding pages: This is the
formation of electrically nonconducting, stoichiometric solid halides of

the metal in a low valence state. Corbett et al. have suggested that the

 

solubility of the posttransition metals in their molten normal halides
parallels or, rather, depends on their ability to form a cationic species

of lower than normal charge. We shall see that this tendency increases

 
 

 

 

 

Table 6. Solubility of Some Transition and Posttransition Metals in Their Molten Halides

 

 

Solubility Temp Solid _ Solubility  Temp Solid
M-mz (mOle % M) ( ° C ) "Subhali deS " Rers M wcz (mole % M) ( o c ) "S’llbhalides " RefS
Ni-NicCl, 9.1 977 None 70 In-InCl InCl, In;InClg 97
"InClz 1
(InAlCL,)
Ag-AgCl 0.03 490 None 71, T1~-T1CL 0.09 550 T1Cl 71
0.06 700 0.09 650 (T1A1C1,7)
Zn-7nCl, 0.18 500 None 80 Sn-SnCl, 0.0032 500 71
1.64 600
Zn-7Znlis 0.28 500 None 80 Sn-SnBro 0.068 500 71
1.65 670
0d-Cacl, 14..0 550 None 54,79 Pb=-PbCl, 0.020 600 71
21..0 800 0.052 700
30.0 1000 0.123 800 ~
Ca-CdBr, 14.0 550 None 79 Pb=FbI, 0,024 440 ®
20.0 700 0.15 600
28.0 900 0.41 700
Cd~-Cal, 2.5 400 None 79 Sb~SbC13 0.018 270 80
-';2-8 ggg Sb-SbIs 1.69 200 SboTy, 80, 102
’ 3.5 300
Heg~HeCl, 7.0 280 HgoClo 73 5.8 400
13‘8 ggg Bi-BiCl, 28.0 202 107
: ' 4640 320  "pic1"
Hg-HeTp 25.0 230 HeoIo 76 28.0 550 [Bis(AlC1,)5] 112
'35.0 280 100.0 >780
Ga=-Ga,Cl, 3.7 180 GaCl, GaGaCl, 71 Bi-BiBrj; 21.0 205 - 108
(GaAlcl,) 57.0 294  "BiBr"
45,0 440
Ga-GasBr, 14.0 180 GaBr, GaGaBr, 95 100.0 5538
(G&AJ.BI‘4 ) 80 *
100.0 >458  "BiI"

 

 

 

 

 
 

¥)

 

43

rapidly within each group with increasing atomic weight (see Table 6). It
also depends on the halide ion, and Corbett et al. attribute this latter

trend to the relative tendency toward stabilization of the higher oxida-

tion state by complexing with the halide ion, normally in the sequence

Cl > Br > I, but reversed for Cd and Hg.
Transition Metals

The occurrence in this group of elements of more than one stable
valence state is a well established general feature of their chemistry.
However, very little seems to be known about the interaction of these
metals with their molten salts. Solid salts of a monovalent state of the
metal seem to be almost entirely missing, the bivalent state being the
lowest in which the great majority of stable solid or liquid halides of
the transition elements are known to occur. There is a notable exception
[besides the compound KNi(CN)s] in the study of the nickel-nickel chloride
systems by Johnson, Cubicciotti, and Kelly,’® in which a solubility as
high as 9 mole % nickel metal was found at the eutectic temperature of
980°C. The authors consider the formation of Nit ions from Ni + Ni?t —
2Nit as being indicated by the freezing-point depression in conjunction
with the heat of fusion, but do not rule out the possibility of an alter-
nate, "physical," interpretation.

However, in the latter case an asymptotic approach of the activity
of NiCl, to the curve of ideal activity vs concentration near afiiClg =1
would be expected but is not found. A third possibility which was not
discussed by the authors, and the odds of which are hard to evaluate in
this case, is the splitting-off of two electrons, according to Ni — Nilt +
2e~, similar to the dissolution of electropositive metals in their molten
halides discussed in the preceding sections. Electrical conductance meas-
urements will have to be made to settle this question. Such measurements
are likely to show either very little change in conductance, indicating
Nit ions, or some small rise resembling that in the NACl;-NACl, system,
above, entailing a similar interprétatidn in terms of electron exchange
between metal ions of different valency, Ni%t and Ni?*, A large rise in

conductivity, corresponding to the introduction of truly mobile electrons,

 
 

 

 

4l

that is, of electrons in shallow, F-center-like traps, is considered the

third, least likely possibility.
Posttransition Metals

Tt appears that in the first B subgroup of the periodic system, only
the solubility of silver in silver chloride has been determined and found
by Corbett and Winbush?* to be as low as 0.06 mole % at 700°C. The only
solid subhalide known in this group is Ag,F, prepared in aqueous solution
and readily decomposed at temperatures above 90°C.7?

In the second B subgroup the products of interaction of mercury metal
with its dihalides, that is, the "subhalides" Hg,X, of mercury, have been
well known for a long time and are stable, even in contact with water.

The phase diagram of the (anhydrous) HgCl,-Hg system (Fig. 28) as deter-
mined by Yosim and Mayer’? shows solid HgoCl, to be stable up to the
"syntectic" temperature of 525°, where it decomposes into a salt-rich melt
of almost the same composition and a small amount of a liquid mercury
phase containing 6.8 mole % HgCl, in solution. With the use of a calori-
metric value of the heat of fusion of HgCl,,’% the depression of its

UNCLASSIFIED

ORNL-LR -DWG, 56700

T
sk
3 2

600

 

   

550

500}

4501 Hg,Cl,+L,

 
 
    

wezs RUFF 8 SCHNEDER
0O SMITH & MENZIES

 

  
 

TEMPERATURE °C

THIS WORK
© THERMAL DATA
A VISUAL DATA

      
 

 

 
 

 

ngmzagm_ o
10 20 30 40 50 60 70 80 90 100
MOLE PERCENT Hg

 

Fig. 28. Mercury—Mercuric Chloride System [S.
J. Yosim and S. W. Mayer, J. Phys. Chem. 64, 909
(1960} (reprinted by permission of the copyright
owner, the American Chemical Society)].

o+

 
 

 

"

¥

 

45

freezing point on addition of mercury metal was interpreted in terms of
the dissolution either as mercury atoms or as undissociated Hg,Cl, mole-
cules from Hg + HgCl, — HgpClp. According to a thermodynamic analysis
of the liquidus curve for HgpCl, by Yosim and Meyer,?2 Hg,Cl, is a more
likely solute species than the mercury atom in that concentration region.
Since HgoCl, does not impart electrical conductivity-to its solution in
molten HgCl,,”” the former must be undissociated, as is the latter. The
deepening in color of these solutions either with increasing temperature
or mercury-metal content merits further investigation.

A study of the HgI,-Hg system by PElaebon and Laude, 1929,7° seems
to be too incomplete to allow definite conclusions to be drawn.

The cadmium halide systems have been’” and still are under quite
extensive investigation. With one exception,78 most probably in error,
if for no other reason but that the reported phase diagram for CdCl,-Cd
violates the phase rule, no solid subhalides Cd,X, have been observed.
However, considerable solubility of metal in the molten halides exists,

and it rises with temperature.79 Like that in the alkali-metal halide
systems, the solubllity seems to be largest in the bromide and lowest in

the iodide system. A very low value of 1.5 mole % at 600° given for the
iodide system,go however, seems to be in error, éompared with 10 mole %,
according to Topol and Landis.”’® The latter determination is not com-
pletely satisfactory either, because of a probably too rapid rise in
solubility at low temperature which produces an entirely improbable
curvature in a logarithmic plot of solubility vs l/T. Consolute tempera-
tures may be estimated to be very high, in the general vicinity of 1500°,
and are not verifiable because of high vapor pressure.

A considerable amount of data exists for the cadmium—cadmium chloride
system.77»81=83  aten’4 proved the solutions to be not colloidel in nature
because of the large melting-point depression of CdCl; on addition of cad-
mium metal. He also showed that the electrical conductivity does not in-
crease, but rather decreases slightly. Farquarson and Heymann®% found
the solutions to be diamagnetic, and this was confirmed by Grjotheim
et al.8% and by Nachtrieb.8® It indicated that the metal dissolves

 

either in the form of atoms, molecules, or molecule ions such as (Cdp)3T,
analogous to (Hgy)?*, and not as ca* ions.87 Cryoscopic data83:79 to-

gether with a recent calorimetric heat of fusion”4 of 7.22 kcal/mole for

 
 

 

46

CdCl, lead to the suggestion of ideal behavior for concentrations of
CdgClg in CdCl, below approximately 2 mole % CdzCl,, but increasing nega-
tive deviation from Reoult's law for CdCl, with increasing metal (Cd,Cl,)
solute concentration. This negative deviation (7CdClz = 0.96 or 0.99
for a cadmium atom or a CdpClp solute, respectively) at the monotectic
(15 mole % Cd,Cl,) may be explsined by the substitution of large cations
(cd; )2t for smaller Cd?T ions. It is analogous to similsr deviations
found in systems with other large cations such as KCl-CdCl, and may or
may not be thought of in terms of the formation of complex anions such as
(cac1,)?~.88 perhaps to explain the asymptotic decrease in the negative
deviation with decreasing metal concentration, the formation of undis-
sociated species Cd,Clg end CdsX, has been suggested,’’ but does not ap-

pear plausible in a highly ionized and ionizing medium such as molten

- CdClp. An alternative explanation is that in small concentrations the

complex anion may act as a "common ion" with respect to some of the ions
of the solvent salt, CdCl,, or, in other words, that a certain concentra-
tion of complex anions such as (CdCl4)2‘ is present, even in pure molten
CdClp,. Solid solution formation which would invalidate any such inter-
vretations of cryoscopic data is deemed unlikely.

The observed negative deviation from Raoult's law would be more dif-
ficult to reconcile with the “"physical" model (cadmium atoms), although

Nachtrieb®® points out that the true state of affairs may be a subtle one,

perhaps corresponding to a cadmium atom "solvated" by a Cd?* ion82 or to
a solute species Cdp?t with an asymmetric charge distribution.
Observations by Grjotheim et 81.8° and by Hertzog and Klemm®? on the
electrical transport of Cd with respect to C1~ were interpreted in terms
of strong interaction between Cd and Cd?*. Other effects such as those
produced by the addition of a third component upon the solubility of
cadmium metal may be taken as an indication that a fairly stable {(Cd,)2+
ion exists. According to Cubicciotti®® the addition of XCl to CdCl, grad-
ually lowers the solubility of cadmium in the melt. Corbett et al.
explained this by assuming that the excess chloride ions stabilize the
higher oxidation state of cadmium by complexing.8°:88:91 However, by
adding A1Cl; (a strong Lewis acid) they prepared a solid, diamagnetic
compound of the formula Cd,(AlCl,), conteining all of the cadmium in oxi-

dation state I,9° In the melts the increased stability of the cadmium(I)

n

 
 

 

”

LAY

 

47

oxidation state when the halide ion X~ (C1™, Br~, I~) is replaced by the
larger AlX,~ was considered to result from the decrease in the interaction
of the Cd*¥ cation with the anion, and, in the solid, from the decrease in
the difference in lattice energies of the salts in the two oxidation
states. The contrast between the dark~brown color of melts containing no
AlClz and the light-green color of those in which ALCl,” is present was
attributed to the role, in the former case, of the (highly polarizable!)
uncomplexed halide ion as a bridge in a weak association between the ions
in the two different oxidation states, Cd,2* and Cd?*t. Other examples of
deep color are found even in aqueous systems containing Fe, Sn, Sb, or Cu
in two valence states.®3

In the zinc systems the solubility of the metal does not much exceed
1 mole %.80 Because of the strong halide complexing tendency of Zn?*, it
is doubtful whether addition of AlX; would raise the solubility of zinc
in ZnX, to any considerable extent by allowing the lower valence state
(zn,%%?) to form in greater concentrations.

In the third B subgroup of the periodic system the halides of mono-
valent indium and thallium have long been known, and some of those of gal-
lium have also been prepared recently.?498 "Dihalides" of these elements
have been recognized as ionic salts of the monovalent metal with anionic
complexes of the trivalent metal, for example, Ga(I)[Ga(III)cl,].?9=101
Again, A1Cl; was employed by Corbett et al.91+24:95:98 0 aid in the com-
pletion of the reaction 2Ga + Ga’*t — 3@at, that is, to convert all of the
gallium into the low oxidation state (I), as in Ga(I)ALCLl,.° The solu-
bility of aluminum metal in molten AlI; was reported’t to be as high as
0.3 mole % at 423°C.

In the fourth group of posttransition elements, dihalides of germa-
nium as well as those of tin and lead are well known stable compounds.

The solubilities of the latter metals in the molten dihalides are very
small, 7> 80 and it does not seem to have been shown as yet that addition
of ALCl,; incréases the solubility, although this would seem to be a test
for the existence of an oxidation state lower than II.

All three elements of the fifth poéttransition group, arsenic, anti=-
mony, and bismuth, possess a very stable oxidation state (III). The
solubility of the metal in the molten trihalides rises rapidly with in-

creasing atomic number, and in the bismuth systems even leads to the for-

 
 

 

 

48

mation of solid "subhalides" of the approximate stoichiometry BiX. Vapor
pressure and emf studies of the SbI3-Sb system by Corbett et al.}02,103
demonstrated the existence of a catenated I,Sb-~SbI, species. Older
claimsi®4 for the existence of dihalides BiX, of bismuth were not substan-
tiated later,1°5 especially not in more recent phase behavior and related
studies by Sokolova, Urazov, and Kuznetsov:©® and particularly by Yosim
et al.t%7 10 (rigs. 29-31). The latter study is distinguished by the ex-
perimental feat of determining the consolute temperature (778°) under a
BiCl; pressure of 80 atm. The latest suggestion for solid bismth "mono-
chloride," "BiCl," on the basis of x-ray diffraction measurements,l'! is
the stoichiometric formula BigCl, or (Big)(BiCls),(BisClg), allegedly con-
taining, besides anions BiCl3?*™ and Bi,Clg?™, the remarkable complex
cation BigӴ corresponding to the unique oxidation state of 5/9. On the
other hand, in the solid complex compound with A1Cls, BiAlCl, (first pre-
pared by Corbett and MbMullangl), Levy gz_g}.llz propose the trimeric

 

  

 

 

 

 

 

 

800 T T T I T T T T
700
m.—
S
'&J 500 L,
E Fig. 29. Bismuth Metal~Bismuth
sc( 400 - Trichloride System [S. J. Yosim et
& L o al.,, J. Pbys. Chem. 63, 230 (1959)
300 LBl (reprinted by permission of the copy-
= @ "z Yy
E L right owner, the American Chemical
4e ' LetBl Society)]
200 e : . ’
E (BicCl) + Bi
100 = BiCly+ (BiCl) E .
o Loy
0O 10 20 30 40 50 60 70 B8O 90 100
MOLE PERCENT B/ISMUTH
® VIsUAL

DECANTATION
& THERMAL ANALYSIS

 

 

 

©

"

 
 

"

 

 

 

49

form of oxidation state (I), namely, triangular (Bi;)3*, which they also
find in the molten state to represent the best interpretation of radial
distribution functions bbtained by x-ray diffraction.

Various other studies have been made on the molten solutions of
bismuth metal in its trihalides.tl3 Keneshea and Cubicciotti have meas-

ured the vapor pressures of BiX; over these solutions?4~16 and their

 

i I | [ | ' | | {
THERMAL ANALYSIS -0
500 |- DECANTATION— @
VISUAL—A

  
 
 
   
     

400 —

TEMPERATURE (°C)
O
s
I

    
   

 

Ly+Bi

OO0 c 0
p-BiBry + BiBr

BiBr +BI

200 [©
| O
a-BiBry + BI Br

ool—4— L L 1t 1 11 1 |
O 10 20 3 4 50 €60 7 80 9 00

MOLE % BISMUTH

Fig. 30. Bismuth Metal-Bismuth Tribromide
System [S. J. Yosim et al., J. Pbhys. Chem. 66,
28 (1961) (reprinted by permission of the copy-
right owner, the American Chemical Society)].

 

 

 

 

 

O THERMAL ANALYSIS
® DECANTATION

 

 

 

 

 

 

A VISUAL
500 500
M 400
)
.
w Bily+L, - : o 300
—-emaee. - - - - - - o oo &
l_:_ 300: ----- q-nu‘a’nfl---3—-:6'.-:3‘)-::8‘-&--3—0w- :
< Qo C o o o O :
& 1 ABitHL,
a '
5 3
- 200 Bily HBIT) i (BiD+B 4200
{
§
| ]
1
: :
|00 i 1 i 1 1 i al 1 1 |°0

 

 

0 0 20 30 40 850 60 70 B0 90 I0O
MOLE % BISMUTH

Fig. 31. Bismuth Metal=Bismuth Triiodide
System [S. J. Yosim et al., J. Phys. Chem. 66,
28 (1961) (reprinted by permission of the copy-
right owner, the American Chemical Society)].

 
 

 

 

 

 

50

densities.”*117,118 ymile the chloride and bromide systems show large
positive deviation from Raoult's law for the solvent, BiX3, when atoms

are assumed to be the solute species, Bil; seems to behave ideally up to
30 mole % bismuth. The reasons for this peculiar behavior of the iodide
system are not apparent. Bredigtl? interpreted the vapor pressure meas-
urements in the Bi-BiBr; system in terms of an equilibrium between Bijp
molecules, (Bi,)?t ions, and BiBrs, namely, 2Bi, + 2Bi’* — 3(Bi,)?%, but
an equilibrium Bi, + Bi?*t = (Bi;)?* might explain these data equally
well.112 At higher temperatures the equilibrium appears to shift to Bij
and its dissociation products, Bi3* and electrons. A multiple mechanism
for the dissolution of bismuth in BiCls was also proposed by Msyer, Yosim,
and Topol,L09»120 corpettl?l suggested an interpretation of the same
vapor pressure data in terms of a tetramer, (BiCl),, possibly with Bi-Bi
bonds to explain the diamagnetism of solid "BiCl."22 The diamagnetism

of the molten solutions was discussed by Nechtrieb.8® On the basis of emf
and polarographic measurements, Topol, Yosim, and Osteryoungl?3:124 pro-
pose the existence of an equilibrium nBiX ==1Bian in both the chloride and
bromide systems, with n = 4 apparently given preference. An equilibrium
of this nature is also indicated by the optical absorption measurements

of Boston and Smith,'2° in which the ratio of two optically distinct
species, probably the monomer and a polymer of Bit, was shown to depend
on the bismuth-metal concentration. Thus, the majority of authors are
found to prefer the assumption of polymeric subhalide formation in solu-
tions of bismuth in the trihalides to that of bismuth atoms,7»117,118

The early measurements of the electrical conductance by Atent?6:127 ghowed
a decrease in the equivalent conductivity of solutions with metal concen-
tration, This is much better understood in terms of subhalide and Bi,
molecule formation than of dissolution of bismuth atoms.

The solubility of salts of several posttransition metals in the
liquid metals was dealt with by Yosim and Luchsinger.128 At sufficiently
high temperatures bismuth halides are completely miscible with bismuth
metal; HgoCl, was found to be soluble in mercury metal to the extent of
7 mole % at 600°C. The solubility of PbCl, in lead at 1000° is given as
1 mole %, while the salts of the remaining metals are considered insoluble
in their metals. There appears to be no simple relationship governing

this behavior.
The question of specificity was also examined, that is, the question

- of whether metels dissolve salts other than their own halides. The solu-

i%

 

 
 

 

v

»d

 

51

bility of a foreign salt can be explained in terms of oxidation-reduction
reactions to form the halide of the solvent metal, which then dissolves
in the metal. |

SUMMARY

Mixtures of metals with their molten halides are not colloidal sus-
pensions, but true solutions. On the basis of both thermodynamic and
electrical conductance measurements, it is clear that many of the mixtures
are solutions in which the electrons introduced by the metal, especially
in the case of electropositive metals, are in shallow traps and conse-
quently mobile, as, for example, in potassium or in lanthanum systems.

The mixtures also include various other types of solutions in which the
electrons are partly or wholly in traps of greater depth. Such traps are
believed to be diatomic metal molecules such as Na,; diatomic molecule
ions such as Hgy?¥, Cdy?*, Ca,?t; more complex cations such as Bi,3*t,
Bi,*¥, or perhaps even Big®¥; and simple monomeric ions of the metal in a
lower than "normal' valence state such as Bi* and Nd?*. Electron exchange
between cations of different valence states of the metal seems to occur
and to contribute slight electron mobility. The electronic conduction
process and the state of the electron in solutions, such as the alkali-
metal solutions, in which the electrons are quite mobile, are not too well
understood as yet. However, there are a few first attempts to describe
theoretically the structure and electrical behavior of such solutions in
terms of F(color)-center-like electrons, but the degree of delocalization
of the mobile electron or, in other words, the number of metal cations
with respect to which each elebtron‘fiay be considered gquantized needs
further clarification. Solute metal atoms resembling gaseous atoms do
not deserve serious consideration, as their very high polarizability must
lead to strong interaction with the ions of the molten salt; that is, for-
mation of F-center-like configurations, if not with the molecule ions or
similarly deep traps Jjust mentioned above. Measurements of magnetic sus-
ceptibility, of paramagnetic and nuclear magnetic resonance, and of the
Hall effect may perhaps bring further enlightenment regarding this more
detailed aspect of the shallow traps in which the electrons contributing

considerable electronic conductivity are thought to be located.

 
 

 

8.
9.

10.

1

15.
16.
17,

18.
19.

20.

21.
22
23,

 

52

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49.
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51.
52

53.

54 .
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42.
43,
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72

 

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