Elements
Volume
September 1972 and
September 1973
S e n i o r R e p o r t e r
C
D B Sowerby
All of Depor tment of Chemis t ry Unive is i ty of Not t inghan i
C o p y r i g h t 1974
The Chemical Society
Burlington House London
Printed in Northern Ireland
Preface
The framework used in Volume for reporting the Chemistry of the Main-
group Elements appears to have been generally acceptable, and has been
continued in Volume 2. The present volume therefore comprises eight
chapters, each concerned with one of the Main Groups as defined in the
abbreviated form of the Periodic Table given in the Preface to Volume and
it has now been agreed that the chemistry of zinc, cadmium, and mercury will
be included in the Specialist Periodical Reports concerned with the Transition
Elements
The relative sizes of the chapters are much the same as in Volume and this
again reflects the amount of published research in each Group. In Chapter 1
greater coverage is given to those properties of the metals which are relevant
to their use in the generation of electrical energy from batteries, or from
nuclear fission and fusion reactors, and both Chapters and include more
illustrative material. Chapter 3 reflects a steady increase in effort throughout
the Group, but an especially large number o papers have been published on
carbaborane r-complexes. Chapter s large, consistent with the considerable
amount of research which continues to be published on each of these elements.
Chapter 5 now includes a short section on ‘nitrogen oxides and atmospheric
chemistry,’ but the bulk of published material is again concerned with the
chemistry of phosphorus; there are some
500 references to phosphorus,
whereas arsenic, antimony, and bismuth together are covered by 240
references. Careful selection has been necessary in Chapter to avoid overlap
with other chapters or volumes. Thus, this chapter contains the chemistry of
sulphides of Main-group elements, but not sulphides of transition metals.
Again, S N compounds are dealt with in this chapter, whereas S B com-
pounds are in Chapter
5
The
halides of the elements are treated as they arise in Chapters 1-6 and Chapter
7 is restricted to interesting recent developments in halogen chemistry, such
as the superacids. Noble-gas chemistry is represented by a small number of
highly interesting papers, which are discussed in Chapter 8.
We have continued the policy of referring to physical properties and
particularly spectroscopic data) of compounds only where this is essential to
demonstrate some important chemical property, Similarly, we refer only to
those aspects of organo-derivatives which illustrate significant features in the
chemistry of the Main-group element involved. On the other hand, more
structures are becoming available often highly refined) now that X-ray
diffraction methods are becoming computerized; the chemistry becomes more
meaningful, and is more readily explained, once the structure is known, and
other physical measurements become less significant. We have therefore taken
every opportunity to include structures of key compounds.
 
iv reface
consultation has been possible. In spite of this, and in spite of the fact that the
period of coverage of Volume 2 is from September 1972 to September 1973
i.e. 12 months as against 15 months for Volume l), Volume 2 is appreciably
longer. This is not due entirely to the enthusiasm of the authors; with ex-
perience, it has become easier to identify developing themes, and to discuss
them meaningfully, and we have the impression that the amount of research
effort devoted to the Main-group elements is increasing.
C C. Addison
Alkali-metal Oxides
Alkali-metal Halides
Lithium Compounds
Sodium Compounds
Potassium Compounds
Rubidium Compounds
Caesium Compounds
Beryllium
Magnesium
Calcium
Strontium
Barium
1
1
1
11
13
17
25
34
38
42
49
53
56
57
58
6
61
64
73
73
77
86
92
95
99Analysis
V
B y G
Carbaboranes and their Metallo-derivatives
Boron Halides
Bonds
103
04
110
117
139
144
147
151
158
160
170
171
-N
Compounds containing Al-N Bonds 179
Compounds containing A1-0 or AI-S Bonds 183
Aluminium Halides 193
Compounds containing Ga-0 or Ga-S Bonds 201
Gallium Halides 204
Indium Halides 212
Other Indium Compounds
Group IV 225
By P G H a r r i s o n a n d P H u b b e r s t e y
1 Carbon 225
Oxidation Studies 232
Methane 242
Halogenomethanes 245
265
28 1
2
Graphite 228
Derivatives
Hydrides of Silicon, Germanium, Tin, and Lead
Halides
of
Synthesis
Physical Studies of Quadrivalent Silicon, Germa-
halides and Related Compounds
nium, and Tin Halides
(iii) N.m.r. studies
(iv) Mossbauer studies
(i) Halide donors 307
(ii) Oxygen donors 308
(iii) Sulphur donors 3
Molecular Oxides 354
Sulphur, Selenium, and Tellurium
Related Systems 369
-Te M Si, Ge, Sn, or Pb) Bonds 370
Compounds containing Silicon-, Germanium-, Tin-,
and Lead-Nitrogen Bonds 372
Germanium, and Tin. 382
Tin. 382
Tin-Main-group Metal Bonds 383
383
386
Halide Complexes 4 5
Transition-metal Derivatives
 
Lead(1x)
Miscellaneous Physical Measurements
By A Morris and D 6 owerby
1
Nitrogen
NOz-NZO4
2
Element
Phosphides
Hydrides
Bonds to Halogens
Bonds to Nitrogen
Bonds to Oxygen
4 Antimony
Contents
47
472
472
475
479
48
48
482
485
489
493
493
493
495
Chapter 6 Elements of Group VI
B y
Group IV Metal Sulphides
Group V Metal Sulphides
Fluorosulphates
61
3
Other Sulphur-containing Compounds
Tellurides
By M
3 Xenon 1v)
676
676
676
678
679
683
684
In this chapter individual references which are inter-related are grouped
together to make a section and, therefore, reference to several alkali metals
may feature in a single section. Each reference, however, appears once only
within this chapter sothat, if described in one section, it will
not
in any other. Single references to topics are presented systematically
in
the
section on the appropriate metal.
The elements of Groups I and I1 are so closely linked in some instances
that
a
section describing them jointly is presented to avoid duplication in
Chapter 2. Such a case is the section on ‘Molten Salts’, which covers the
chemistry of the molten salts of both Groups I and I1 but is presented only
in this chapter.
sn 3P -6,
-7 (Li 11) has been investigated by the beam-foil technique. Zero-field
quantum beats were observed in the intensity decays of transitions from the
1 sn
3P
terms n = 2, 3, or 4) in 6*7Li1 and the magnetic hyperfine coupling
constant was determined for each isotope for the 2p 3P
terms. Preliminary
=
cm-l
2.73
GHz), A ( l s , 2 p 3P, Li 11) = 0.239 f 0.002 cm-l (7.17 GHz). The
measured fine structures agree within a few percent with recent ca1culations.l
The Auger electron spectrum of freshly filed lithium contains an emission
peak at 51.7 eV which
is
attributed to the KL,L, Auger transition, and un-
identified peaks at 27.5 and 8 eV. On exposure to oxygen the peaks at 51.7
and 27.5 eV disappeared but the 8 eV peak intensified. The low-energy
spectrum was characterized by emissions at 13.3, 24.0, 33.0, and 40.0 eV due
to Auger transitions of lithium, oxygen, and lithium monoxide.2
A
value of
3.05 eV is reported for the work function of freshly prepared lithium films.3
H. G . Berry, J.
a R . E. Clausing,
D.
S. Easton, and G . L. Powell, Surface Sci., 1973, 36, 377.
1
Elements
Lithium, the ‘not so rare metal’, is reviewed. The discussion covers the
occurrence, production, and LISGSof the irietal and its compo~nds .~he metal
has considerable potential in the future generation of electrical energy from
the fusion reaction:
The supply of tritium for this process is derived from:
Within this context, the chemical, physical, and thermal properties of
lithium that are related to its use in fusion reactors have been reviewed.
These include natural abundance, thermodynamic and transport properties,
characterization, analysis purity control, and corrosion
of materials by the
molten liquid5 Problems associated with tritium in the metal are also covered,,
as is the separation of tritium from lithium by crystallization or diffusion.’
The pressures of hydrogen isotopes in equilibrium with their solutions in
liquid lithium have been measured. The square root of the hydrogen pressure
is proportional to the hydrogen concentration in accordance with Sievert’s
Law. Graphical data are presented for 2H.8The metal is also chemically very
reactive. The effect of temperature and pressure on the reaction of static
molten lithium with oxygen, nitrogen, and the compounds CCl,F,, C,F,,
and SF, has been studied, The metal is hsated inductively under vacuum and a
small known volume of gas exposed to the surface. Pressure and temperature
changes are followed by rapid-response instr~mentation.~he liquid metal
is a versatile solvent for both non-metals and metals. Non-metals when
dissolved in the metal may not have the same deleterious corrosive effect on
containment materials as they do with sodium. Chemical processes are
affected by the different thermodynamic stability of lithium compounds. This
is illustrated by the effect of oxygen on the chemical corrosion of niobium
and tantalum by static liquid lithium at 600 “C in capsules. An increase in
the oxygen concentration of lithium from 100 to
2000
is
centrations in liquid sodium or potassium. The free energy
of
formation of
lithium oxide is so great that the liquid metal getters niobium and tantalum to
an oxygen level 120p.p .m. regardless of the oxygen concentration in the
lithium. When the transition metals contain more than a threshold level of
oxygen (400 and 100 p.p.m. for N b and Ta respectively), chemical attack by
lithium occurs at the grain boundaries with the formation of ternary com-
pounds containing lithium, oxygen, and transition metal.1° Methods of
R. Feather, Philippine G eogr. J . , 1973, 17, 16.
R. G.
H . Weichselgartner, Reaktortagung, 1972, 751.
T . E. Little, U . S. Na t. Tec h. Inform. Serv. AD Rept, 1972, No. 759378.
t
V.
A. Cafasso, ANL-800 1 Rept. 1973.
* D. H.
J. Goodall and G. M . McCracken, P r o c . Symp. Fusion Technol., 7 t h , 1972, p. 151.
lo R. L. Klueh, ORNL -TM-4069 R ept. , 1973.
 
analysing the liquid are obviously important. Photon activation appears
applicable for the analysis of nitrogen and oxygen. These elements are deter-
mined by photon activation with a microtron as a pray source. The 13Nis
separated by distillation as ammonia and collected in sulphuric acid for
activity measurements. Oxygen is rapidly separated by distillation as water
for coincidence counting. Sensitivity is 2 x
In addition to its role in the fusion reactor, liquid lithium features promin-
ently in solid-state batteries,13an area which has been reviewed.14 Lithium and
another element are usually separated by a solid or liquid electrolyte perme-
able to lithium ions, which migrate to form a compound with anions of the
second element, thus driving electrons through the external circuit. The
second element has been halogen, though this may be replaced by a compound,
e.g.
vanadium pen t0~ ide.l~resent interest is in the chalcogens, and several
lithium-chalcogen systems have been investigated with this use in mind.
Equilibrium phases in the Li-S system on the sulphur side of Li2S (the only
compound observed) are determined by using an unusual vapour-transport
technique. By this means equilibrium compositions of the melts at various
temperatures can be obtained by utilizing the transport of sulphur vapour
from one melt to another. The Li-S phase diagram exhibitsa large miscibility
gap which extends from the monotectic composition, 65.5 mol S , to
almost pure sulphur
63
to
98.8
>600 C?'
An e.m.f. method using cells of the type
LilLi halide eutectic mixturelLi in selenium is used to determine thermo-
dynamic quantities in the lithium-selenium system. From the cell data the
standard free energy of formation of Li,Se at 360 "C is calculated as -94.0
kcal mol-l.ls In the Li-Te phase diagram, eutectics occur at 179.9 C near
the lithium axis at
35.7
atom
Li,
and
423.1
"C at 10.5atom Li. Two intermediate compounds are present, Li,Te and
LiTe,, melting at 1204.5 and 459.9 C, respe~tive1y.l~
The spectrum of doubly ionized sodium
I11
2500-
1300A and the analysis has been revised and extended as regards the 2p44s,
wt for each element.11*12
l1 B.
E.
la B.
A.
Chapyzhnikov, Kh. N. Evzhanov, E. D. Malikova, L. L. Kunin, and V. N.
l3 M. Eisenberg, Intersoc. Energy C onvers. Eng. Conf., Con Proc. 7 t h , 1972, 7 5 .
lo B. Scrosati, J . Ap pl. Electrochem., 1972,2,231.
l6
1972, 119, 1439.
l7 P. T. Cunningham, S. A. Johnson, and E. J. Cairns,
J .
J . Electrochem. SOC.,1973,120,
328.
A. N.
May
1973.
2
classified lines is now 177, out of
which 110 are newly observed and Classified. Some of the older classifications
are altered and ca.
levels are new: (T)4s4P,
(lS)3p2P; 3P)3d4F; 3P)3d4P, F5/2;
3p, 3d4s
configurations are now complete.20 n the range 380-18OA about 90 lines
are measured, of which 50 are reported for the first time.21 The third, fourth,
and fifth
analyses.22New lines in
V, and Na VI are observed for the
first time and ~l as si fi ed .~ ~he K X-ray spectra of sodium excited by protons,
helium, and oxygen ions of
0.8, 3.2,
and 30 MeV, respectively, have been
measured. The strongest lines are the normal K, satellite spectra produced by
multiple electron vacancies in single ion-atom collisions. In the H- and He-
ion-induced spectra,
the spectrum of atomic sodium between
30
and 150 eV shows lines which can
be attributed to the excitation of a 2s or 2p electron. Considerably broad and
asymmetric absorptions above the lP1 eries limit are due to the simultaneous
excitation of a
coaxial-cylinder, two-electrode system at
of a model for the vapour consisting
of Na,, Na,,
and Na, moieties.26The
temperatures can be semi-empirically derived. In the case of unsaturated
vapours, the number of associates tends to zero with increasing size. The
symmetry of the associates, binding energies, and mobility for sodium and
potassium are given. The effect
of
temperature is calculated on the equilib-
rium between the concentration of free sodium atoms and those combined
in the cluster. Agreement with experimental data is sati~factory.~' he best
available thermodynamic data on liquid metals are tabulated and include
m.p., entropies of fusion, heats of fusion, and heat capacities. Graphical
correlations are presented between heats
of
between entropies
anomalies are discussed in terms of the electron configuration
of
the metal.2a
The surface tensions of molten alkali metals from their melting temperatures
L.
T. Lundstrom and L. Minnhagen, Physica Scripta, 1972, 5 , 243.
22 T. Goto,
23
T.
J .
Z. hysik, 1972, 257, 288.
25 H. W. Wolff, K. Radler, B. Sonntag, and R. Haensel, Z. Physik, 1972, 257 353.
26 R. Morrow and J. D. Craggs,
J .
Issled. Striikt. Mol. Krist. Krist.,
28 J. L. Margrave, CoIlog. Int. Cent. Nat . Rech. Sci., 972, No. 205,
p.
71.
up to
C have been determined in a special high-temperature, high-
pressure apparatus. The surface tensions/dyn cm-l as a function of tempera-
ture (t/"C) under an atmosphere of their own vapours are given by:
7Na = 193.6 - 0.094(t - 98)
YI = 107.1 - 0.069(t - 64)
68.8 - .045(t - 28)
The values correlate well with those previously pu b l i~hed .~~
As with lithium, the majority of the literature on the commercial uses of
metallic sodium is devoted to aspects of the generation
of
electrical energy
either where the metal is used as a coolant in fast nuclear reactors or used as
an electrode in high-power batteries. An indication of the extent of the
nuclear use of liquid sodium is provided in a review of the principal program-
mes involving fast reactors in the Technological aspects are also
repre~ented?l-~* hese applications steadily reveal new chemical properties
of sodium and its compounds, This is illustrated in the proceedings
of
a
conference on the Liquid Alkali Metals which covers fundamental chemistry,
physics, analytical and instrumentation techniques
,
carbon and fission-product behaviour in sodium, physical processes, corro-
sion, and mass transfer.35Also, chemical reactions in liquid alkali metals are
discussed, with particular emphasis on solvation aspects.
A
comparison is
made of the nature and properties of liquid metals, representing continuous
reaction media, with other non-aqueous solvents,
e g .
molecular liquids,
the purification, analysis, and corrosion areas. The non-metals oxygen
,
hydrogen, nitrogen, and carbon, when dissolved in the liquid metal, have a
deleterious effect on transition metals, which are invariably employed as
containment rnaterial~.~' urification and analytical techniques, therefore,
are primarily designed to remove38and m o n i t ~ r ~ ~ * * ~hese elements, in many
cases
in
sit^.^^
To
liquid sodium, ca.
29 A .
Fiz. Khim. Poverkh.
31 F. Chaminade, FRNC-CONF-38 Rept. 1972.
32 Sodium Technology, 1948-1961 [TID-3334 (Pt. l)].
1972.
Wykoff,
34 M . E. Durham, RD/B/M-2479 Rept. 1972.
35 Proceedings of the International Conference of the BNES, London, on Liquid Alkali
Metals at Nottingham University 4-6 April, 1973.
36 C . C. Addison, Sci . Progr. (London ), 1972, 60, 3 8 5 .
37
38
W.
39 K.
40 L.
F. Lust, F. A. Scott, and J. F. Jarosch, HEDL-TME-71-17 Rept. 1971.
1971, 108.
the
Maimgroup
Elements
liquid. These metals, with their strong chemical affinity for nitrogen, effec-
tively isolate the steel from nitrogenjl To analyse for hydrogen in liquid
sodium, a nickel thimble is immersed in the liquid and evacuated to
Torr on the inside. The process relies on the equilibrium between
dissolved and gaseous
Hydrogen leaves the liquid, diffuses through the
nickel, and establishes an equilibrium pressure, the magnitude of which is
dependent on its concentration in the liquid. As little as 0.02 f
0.01 p.p.m.
of hydrogen can be detected.42Hydrogen may exist in a sample of sodium in
several forms, i.e. dissolved sodium hydride, solid sodium hydride, or sodium
hydroxide. To distinguish between these requires several processes. All the
hydrogen is released as gas by vacuum fusion in
a
Subsequent heating to
ing hydroxide hydrogen may be determined by difference.43Alternatively, the
remaining sodium amalgam
400 C.
these conditions sodium reacts with sodium hydroxide to give hydrogen,
NaH, and Na,O. Hydride decomposes to give hydrogen, which is determined
by gas ~hromatography?~ydrogen is also soluble in liquid potassium. Over
the temperature range
340440 O C
log(C/p.p.m. by wt.) = 6.8 - 2930/(T/K)
The pressure
of
log(P/Torr)
= 11.3 -
5860/(T/K)
These pressures are the dissociation pressures of potassium hydride according
to:
KH
pressures is -13.7 kcal mol-l. The equilibrium pressures
of
P112 = C x 104/14.2
where C s in weight . Thus Sieverts' Law is obeyed (I'll2 z C), which
indicates that the species
hydrogen in the metal is m~ na to mi c.* ~
Most interest has centred on solutions of oxygen in liquid sodium since
this element, more than any other, renders the liquid metal corrosive.
41
A. K. Fischer, U.S. . 3 745 068 (CI. 176-38, B O l j G21c),
10 Jul 1973.
(Cl.
73/19;
8
May
1973.
43 Kh . Evzhanov, E. D. Malikova, and L. L. Kunin, Z h r . nnalit. Khim.,
1973, 28, 235.
44 M . Takahashi, J. Nuclear Sci Technol., 1973, 10, 54.
45 M .
N.
Ivanovskii,V. A. Morozov, S. S. Pletenets, and V. V. Sitn iko v,
Izvest .
 
I 7
Vanadium, niobium, and tantalum, and their alloys, have a low intrinsic
soIubility in liquid sodium and suffer but slight corrosion. The presence of
oxygen in the liquid, however, leads to penetration by non-metals into the
transition metal, internal oxidation, oxide scale formation, spallation or
dissolution of oxides, and, in some cases, penetration by the sodium.4G
Whether the transition-metal surface oxidizes or whether sodium extracts the
oxygen contained in or on the metal depends largely on the relative free
energies of formation of the transition-metal oxide and sodium oxide,
respectively. The situation is more complicated, however, since the energy
balance is affected by the activity (or concentration) of oxygen in the sodium
or in the solid metal, i.e. a dilute solution of oxygen in liquid sodium may be
reducing whereas a more concentrated solution will oxidize a particular
transition-metal surface. Further complications arise when ternary com-
pounds form which are stable in sodium. Most transition metals form at least
one ternary oxide with sodium. These points are illustrated below. Vanadium,
exposed at
to static sodium solutions containing oxygen up to 4000
p.p.m.,getters all oxygen from solutions which contain less than 2000 p.p.m.
The distribution coefficient for oxygen between vanadium and sodium is
greater than lo4 at
By alloying chromium or molybdenum with
vanadium, the activity coefficient of oxygen in the solid alloy is increased
and hence the solubility is reduced.47 In sodium containing
2000
p.p.m.
internal precipitates of oxide during the gettering, and the concentration
of
oxygen dissolved in the alloy approaches that of the same alloy without
titanium or zirconium.48 When titanium and zirconium are immersed in
liquid sodium containing dissolved sodium oxide at 600 "C, the surfaces are
covered with the ternary oxides Na,Ti04 and Na,ZrO,, respectively. These
compounds were identified in situ by their X-ray diffraction patterns. The
compound Na4Ti04 was detected when the sodium contained from 100 to
12000
p.p.m. oxygen. At the end of long contact times the oxide Ti0 formed
below the ternary oxide, which suggests that the ternary oxide is formed
first and is followed by diffusion of oxygen into the substrate metal to form
TiO. With zirconium, a rapid formation of the oxide ZrO, is postulated which
is followed by a slow reaction with dissolved sodium monoxide to give Na,Q,
Zr0,.49 Liquid potassium, like sodium, also becomes more corrosive towards
transition metals when it contains dissolved oxygen. Analysis
of
potassium
after immersion of tantalum at 600, 800, and 1000"C shows that the amount
of tantalum finding its way into the alkali metal increases with the amount of
oxygen originally dissolved in the liquid metal. Again, a ternary oxide phase
is formed. Oxygen held in the tantalum also promotes corrosion when the
transition metal contains more than a threshold concentration of oxygen in
4 6 H.
U. Borgstedt and
47 R.
L. Klueh and J. H . DeVan, J. Less-Common Metals, 1 9 7 3 , 3 0 , 9 .
48 R.
25.
49 M. G. Barker and D. J. Wood, J.C.S. Dalton,
1972 2451.
solid solution; potassium penetrates the solid metal intergranularly and trans-
granularly
via
ternary oxide formation. The threshold levels of oxygen for
this type of attack at 400, 800, and 1000°C are 500, 700, and lOOOp.p.m.,
respe~tively.~~istribution of radioactive corrosion products is obviously
important in flowing sodium. Particulate material deposits according to flow
rate and geometry of circuit, size of particulate, and whether the species is
soluble in the sodium or reacts preferentially with metallic parts of the circuit.
Initial experiments have investigated the transport and deposition character-
istics of 59Fe,54Mn, nd 6oCo. The 59Febehaviour is characterized by its
appearance as a firmly adherent layer on pipework downstream of the test
section. 6oCo s similar to iron but the deposit is less strongly attached. The
behaviour of 54Mn is characterized by its rapid and highly preferential
migration to the coldest part of the circuit.51 Adsorption of caesium, a
product of the fission process, also occurs from solution in sodium at trans-
ition-metal surfaces. Between 100 and 200°C, caesium is adsorbed on to
nickel and steel (EN-58B) surfaces but at 800 O C the adsorption is eliminated
The mechanism of adsorption is not clear.52
Determinations of the solubility of oxygen in liquid sodium are numerous
and the values vary.
therefore, to derive the mean solubility relationship:
log(S/p.p.m.) = 6.1587 2386.4/(T/K)
from T = 387 to 828 K , using the least-squares method. This equation is
recommended for fast-reactor Methods of determining these small
concentrations differ widely. Thus at 350-530 O C , the solubility ca.
10-850 p.p.m. as determined by an e.m.f. method using the cell:
where the rare-earth oxides comprise the solid electrolyte which separates
the reference electrode Cu,OICu(or air]Au) from the second electrode, a
mixture of sodium with sodium monoxide.54Alternatively, a vanadium wire
is immersed in the molten sodium to allow oxygen to partition between the
two metals. The wire is subsequently removed and analysed for oxygen
content. The method relies on a knowledge of the equilibrium distribution
coefficient of oxygen between sodium and vanadium. These values (as
W. oxygen in V) are given at 750 OC over the range 0.003-16 p.p.m.
oxygen in sodium.55
Carbon dissolves in liquid sodium but to a lesser extent than do hydrogen
or oxygen, and methods for determining the carbon content of liquid sodium
continuously are generally less advanced than those for oxygen and hydrogen.
50 R. L. Klueh, Corrosion (Houstom), 1972, 28, 360.
61
J . Brit. Nuclear
Energy SOC., 973, 12, 63.
52 H. E. Evans and W. R. Watson, RD/B/N-2094 Rept. 1971.
53 J. D. Noden, J .
Brit. Nuclear Energy SOC.,
R .
H.
A technique, reminiscent of the electrochemical oxygen meter, is described,
however, which equilibrates the carbon dissolved in liquid sodium with a
membrane of a-iron at 500-700 "C. This membrane forms part of an electro-
chemical cell and is separated from a reference source of carbon by a fused
electrolyte of 1 :
Li,CO,-Na,CO,, which is able to transfer carbon in ionic
form. The voltage between the membrane and reference electrode gives a
measure of carbon activity in the membrane and hence in the sodium.56
Protection and security measures against accidents with liquid sodium are
reviewed.57A fire-extinguishing powder that is especially effective against
alkali-metal fires consists of 45.4 NH4H,P04 (fluidized with up to 6%
of its weight by Si02 and silicone resin), 45.4%urea, 9.1 polystyrene
micro-balls (d
covers the metal, whose temperature falls very rapidly.58
High-temperature (300
"C) storage b a t t e r i e ~ ~ ~ - ~ ~nvolving liquid sodium
utilize the chemical reactions between sodium and liquid sulphur. The
Na-S system is complex, however, and contains several polysulphides with
the general formula M,S, containing
S ,-
ions. Of about 15 polysulphides of
the alkali metals described in the literature, the crystal structures have been
determined for only about three, which reflects the difficulty of preparing single
crystals of the polysulphides. The Na,S-Na,S,-S region has been investigated
mainly by high-temperature microscopy but some complementary experi-
ments involve d.t.a., t.a., and quenching techniques, The components
S and
Na,S melt at 118 f 1 "C and 1168 f 10 'C, respectively. The intermediate
phases Na2S2,Na2S4,and Na,S5 which are formed melt at
478 f 5,294 f 2,
and
270 f 5 "C, respectively. Na,S, melts incongruently. The shapes of
polysulphide crystals appearing just below the melting points are detected
by high-temperature micro~copy.~~urther d.t.a. work reveals that when
Na,S-Na,S, or Na4S4-S, mixtures are heated, a reaction occurs near the
m.p. of sulphur with formation of Na,S5 as the initial step, Unless the S:Na
ratio is
>5:2 then further reaction between the sulphides occurs, until at
equilibrium only those species are observed corresponding to the given Na:S
ratio. The highest sulphide is Na,S,, and Na,S, does not exist at the m.p.;
this stoicheiometry is really a
1 :
Na,S,-Na,S, eutecticBgThe sodium poly-
sulphides Na,S, and Na2S5,however, can be prepared from the reaction of
56 M . R. Hobdell and D.
M.
5 7 M . D e la Torre Cabezas,
Energ.
Niicl.
1972, 16, 439.
5 8 E. Chahvekilian, R. Peteri, and A . Hennequart, Fr. Demande
2 102 424 (Cl.
A 6 2 4 ,
5 9
J. Fally, C. Lasne, and Y. Lazennec, Fr. Demande 2 142 695 (CI. H O l m ) , 9 Mar 1973.
6o S. Gratch , J. V. Petrocelli,
R.
61 T.
Nakaba yashi, Ger. Offen. 2 240 278 (CI. H Olm), 2 Apr 1973, Japan.
62 J.
Fally and
63
C.
Levine,
64 S . P. Mitoff,
U.S.
P. 3 672 994 (CI. 136-6, H O l r n ,
27 Jun 1972.
Chemica Scripta,
12 May 1972.
1972, p. 38.
1973, 12,43 5.
Main-group
Elements
sodium chloride with K2S, and K,S, respectively in liquid ammonia. Using
KZSG, however, the only polysulphide obtained is Na,S,. The physical
properties
of
of
of
structure of the polysulphide linkage is confirmed by photoelectron spectro-
~ c o p y . ~ ~odium tetrasulphide, Na,S,, is in fact tetragonal, with space group
1 4 2 d and cell dimensionsa = 9.5965, c = 11.7885 A and 2 = 8. The struc-
ture is built up of unbranched and separated Si- ions surrounded by Na+ ions.
Adjacent SZ- ions are
S-C
is
of
two types: a distorted tetrahedral
arrangement of sulphur atoms with two pairs at 2.826 and 2 . 8 4 2 A from a
central sodium atom; a sodium atom at the centre of a distorted octahedron
with three pairs of sulphur atoms at distances
2.887, 3.043,
respectively, from the sodium.G8The monosulphide, Na2S, forms several
hydrates, the stabilities of which depend on the temperature and partial
pressure of water vapour above the compounds.
By
thermally decomposing
N a 2 S , 9 H 2 0 he compound Na2S,M,0 was found to be stable over the largest
pressure and temperature range, with an enthalpy of hydration of 16.62 kcal
mol-l. Heats
6.76
and
1.96
kcal mol-l, re sp e~ ti ve ly .~ ~
The spectrum of doubly ionized rubidium (Rb 111) over the range
370-
3500 A has been re-analysed. The existing analysis is revised and extended.
Most levels of the 5s6s5p4d and 5d configurations are now known.70A second
analysis yields most
eV.'l The vapour pressure of
liquid rubidium from 129 to 278 C has been determined by thermogravi-
metric and mass-spectrometric techniques. Calculation of the latent heat of
vaporization from the vapour-pressure data yields a value
of
19.0 f 0.5
kcal (g atom)-l at 298 K. The dissociation energy of the Rbz molecule is
10.0
f
0.5 kca lm01- l.~~ oth rubidium and caesium have relatively high
vapour pressures which pose experimental problems in the handling of these
elements. The saturated vapour pressures, of Rb at
683-1649OC
and
0.97-101.5
775-1600 O@ and
nd
"C
Chem. ,
ge R. C. Kerby and M . R. Hughson,
Canad,,
Mines
Amer. ,
73 L. I. Cherneeva and
V.
The crystal structure of Li,CdPb is f.c.c., with
a = 6.837 A , d(expt) = 6.79
at
20 'C , and d(X-ray) = 6.93 for 2 = 4. The observed and calculated X-ray
intensities are tabulated. The most probable space group is T$ -
F 4 3 m 7 ,
The Li-Ga phase diagram, when investigated by d.t.a., reveals four new
intermetallic compounds. Of these, Li,Ga, LiGa,, and LiGa, are identified by
X-ray diffra~tion.~~he crystal structure of Li,Ge, formed in the Li-Ge
system has been determined by X-ray diffraction. The compound crystallizes
with the orthorhombic space group
Cmmm,
having
a
= 13.21,
c = 4.63A, d(expt) = 2.25, d(ca1c) = 2.28 for 2 = 4. Li,Ge, is not iso-
structural with Li,Si, but there are many similarities between the s t r u c t ~ r e s . ~ ~
For a series of molten Li-Sn alloys, direct measurements
of
e.m.f. from the
electrolytic cell (-) LilLiC1, LiFILi, Sn (+), carried out at 550 'C, indicate
substantial negative deviation from ideal beha~iour.~'hase equilibria in the
ternary system Wa-K-Rb have been investigated by thermal methods. The
ternary eutectic
lower than the Na-K binary
eutectic temperat~re. '~he gaseous equilibria
NaAg(g) = W g ) +
NaAg(g) + Ag(g) = Na<@
nd
have been determined by using a double-oven technique in combination with
mass spectrometric analysis of the composition of the vapour. The Third Law
enthalpies
AH:
f
respectively. From these enthalpies the dissociation energy Dg(NaAg) is
31.8
=t3.0 kcal mol-l. The corresponding 0gg8s 32.4 f 3.0 kcal mol-l.
The standard enthalpy of formation
AH;,298
61.4 f 3.4 kcal mol-l. The experimental dissociation energy
of
NaAg(g)
is
considerably lower than the value of 51 kcal mol-l calculated using the
Pauling model of a polar bond.79 n an investigation of Na-Cd alloys from
20 to 86 atom Cd by d.t.a., single-crystal and polycrystalline forms of the
compounds Na,Cd,, and NaCd, have been prepared. NaCd, is cubic,
a = 8.04
A , d(expt) = 7.1.
At 20% Cd, prismatic crystals are obtained which probably contain less Cd
than NaCd,.80 The phase diagram of Na-Hg has been re-investigated, and
changes in enthalpy, free energy, and entropy at 648
K
74 K . Kuriyama, M . Yahagi, and KIwamura, Jap. J. Appl.
Phys. ,
1973,12, 743.
75 S. P. Yatsenko, K. A. Chutonov, S . I. Alyamovskii, and
E.
N.
Dieva, Izuest. Aknd.
7 6 V. Hopf, W. Mu eller, and H . Schaefer, Z. Naturforsch., 1972,
27b,
1157.
N. Gerasimenko, A.
Elektrokhimiya,
1972, 8, 1622.
B . J. Ott, J. R . Goa tes, and C. C. Hsu,
J. Chem. Thermodynamics, 1973, 5 , 143.
Nauk S.S .S .R . , Metal . , 1973, 185.
7 9 V. Piacente and K.
A.
8o V . N. Kornienko and N. N. Zhuravlev, Kristallografiya,
1972, 17, 863.
Maingroup
Elements
well known, the system forms a series of intermetallic compounds, amongst
which the congruently melting NaHg, is the most stable.*l
The velocity of sound in liquid metals and alloys, its variation with tempera-
ture, and its relation to the adiabatic and isothermal compressibility of the
metals has been reviewed. Experimental results are compared with theoretical
considerations.*2The velocity of sound at 7 .5 MHz
and the density have been
measured
as
functions of temperature for a number of solutions of metals in
liquid sodium. The results cover the whole concentration range for mercury
and the range 0-5 atom solute for Au, Cd, In, Pb, and Sn. Changes in the
adiabatic compressibility and the mean molar volumes provide no convincing
evidence for compound formation in the Zipid phase of Na-Hg. There is,
however, a substantial volume contraction on mixing which apparently
causes the compressibility, when plotted against concentration, to fall well
below the straight line linking pure sodium with pure mercury. The con-
traction and the associated reduction in the adiabatic compressibility are not
peculiar to Na-Hg since the effects for all the different solutes used are
remarkably similar.83Evidence that compound formation does occur in the
liquid, however, comes from potentiometric measurements.
The
potentials
of homogeneous liquid sodium and potassium amalgams containing 0.65-
7.19 (g atom Na)(l Hg)-l and 0.053-2.62 (g atom K)(l Hg)-l have been
measured at 25 and
theoretical values calculated from the Turin-Nernst equation, are explained
by chemical interaction between Na (or K ) with Hg. The number of atoms of
mercury co-ordinated with Na (or K) in liquid amalgams is calculated by the
Hildebrand equation, and, near saturation, is 5-6 and 15-16 for Na and
K,
respe~tively.~~nthalpies of mixing sodium and gallium and the enthalpies of
formation of the two compounds NaGa, and Na,Ga, have been determined
calorimetrically at 723 K. Referred to the pure constituents in the liquid
state, the following values are found: AH, (NaGa,) 17.5 f
1.0,
AH,(Na,Ga,)
of
was determined as 13.5 f 0.5 kJ
m01-l.~~he density and surfxe tension of
five solutions of indium containing 0.5-7.0 atom% in liquid sodium have
been measured at 170-400 O C
by means of the maximum-bubble-pressure
technique. The gram atomic volumes of these solutions, calculated from the
density, indicate a substantial contraction on mixing which is about twice
that for analogous Na-Cd solutions. The surface tension of liquid sodium
increases slightly on adding indium, indicating
a
bu lk ,
active behaviour of cadmium.86Two compounds are reported in the Na-Bi
a1 M. A. Bykova and
A . 6
a 4 L.
M.
Ruban,
A. I. Zebreva, and V. P. Gladyshev, Eiektrokhimiya, 1972, 8, 1021.
1 3 ~M . Gambino and
J. P.
G. M . B. Webber, Ultrason. Ind. Con Pap., 1970, 22.
H. A .
 
Elements of Group I 13
system with very high melting points: Na,Bi, m.p. 775OC; NaBi (incon-
gruently melting at 446 C).
These compounds are so stable that it is advan-
tageous to consider them as independent components and to use for the liquid
alloys of this system a quasi-ideal-solution Activities of potassium,
deduced from e.m.f. data, show that concentrations below 1 atom in
liquid indium at 577 O C
obey Henry's law, with a linear dependence of the
activity of K on its concentration in the alloy. The study also reveals a limited
solubility of potassium in indium.88 Thermodynamic data on liquid alloys
of potassium with thallium over the complete concentration range have been
obtained from 350 to
K-glasslK, Tl(1). Partial molar enthalpies and entropies of mixing are
calculated from the results. Some interesting maxima and minima are
observed which indicate
potassium and thallium.sgThe density of caesium vapour above solutions of
caesium
in
liquid rubidium has been determined from 293 to 313 K and mole
fraction Cs, x
with increasing caesium concentration. For a given concentration, the
density increases with increasing temperature. For all values of x
at
K ,
and for x 2 0.35 at 2 3 1 3 K , the solutions showed negative deviations from
Raoult's law. If the solutions are sprayed on the vessel walls, effects due to
surface adsorption are observed.g0
Aqueous Solration.-A review, covering the 1968-1 972 publications, deals
with physical properties, thermodynamics, and structures
of
limits.g1 Thermodynamic aspects of ionic hydration also reviewed include
the thermodynamic theory of solvation; the molecular interpretation of ionic
hydration; hydration of gaseous ions (AG's, AH'S , and A S S ) ; thermody-
namic properties of ions at infinite dilution in water, solvent isotope effect in
hydration; reference solvents; and ionic hydration and excess proper tie^.^^
A
third review on the hydration of ions emphasizes the structure of water in
the gaseous, liquid, and solid states; the size of ions; and the hydration
numbers of ions and the structure of the hydrated shell from measurements
of mobility, compressibility, activity, and from n.m.r. spectra.93Pure water
and aqueous LiCl t concentrations up to saturation have been examined by
neutron and X-ray diffraction. For the neutron studies 'LiCl and D20 are
employed. The data are consistent with a simple model involving only
87 A. G.
Rasplav. Solei, 1971, 37.
M . A. Bykova and A. G. Morachevskii, Izvest. V. U.
Z., Tsuet. M et . ,
( A ) , 1973,
89 S. Aronson and S. Lemont, J . Ch em. Thermodynamics,
1973, 5, 155.
1972,
46,
2987.
Water: Compr. Treatise 1973, 3 , 1.
93 H. Ohtaki,
 
Inorganic Chemistry of the Main-gvoup Elements
nearest-neighbour interactions. In the pure liquid and in dilute solutions the
structure of the water molecule resembles that in the vapour phase. Each
water molecule in the liquid is tetrahedrally co-ordinated by four others. This
basic water structure gradually diminishes with increasing LiCl concentration
and is not seen at mole ratios of ten or below. The co-ordination around Li+
appears to be tetrahedral, with co-ordination through oxygen. The C1- ions
are octahedrally co-ordinated through
H.94A theoretical study concludes that
energetically the most favourable aquo-complexes of the ions Li+, Na+, and
K+ contain six, eight, and eight water molecules, respectively. The energy
E of [M(H,O),] t complexes as a function of the distance between the metal
ion and the atom is calculated according to the extended Huckel method,
assuming a donor-acceptor bond between H,O molecules and alkali-metal
ions. The bond energy of the aquo-complexes, AE, is obtained from
AE
=
E - n E ~ ~ 0 ,here E H ~ Os the energy of a free water molecule. The potential-
interatomic distance curves show a sharp minimum (AE,) for Li+ and Na+
ions but are shallow for K+. The values of AE,/n can be used as a measure
of the respective aquo-complex stability at higher temperatures. The flat
potential curve for K+ suggests both a larger contribution from the second
hydration shell to the total hydration shell and a greater mobility of water
molecules in the hydration layer of K+ than in that for Li+ and Na+.g5The
equilibrium constant K has been determined for the Li-isotope exchange
reaction
6Li(s)+ LiCI(aq)
using an electrochemical quadruple cell without liquid junction. The value of
K at 296.6 K is 1.046 f
0.13 and, when compared with values of K calculated
from statistical thermodynamic theory for various model reactions, is
consistent with a tetrahedrally co-ordinated structure for the aquated lithium
Hydration numbers, however, when determined by acoustic-inter-
ferometer measurement of the adiabatic compressibility of the solution at
infinite dilution at 35 C,have values
Lif, 3 ; Rbf, 2; Ca2+,
7,
8.97
From literature data on the temperature dependence of the conductivity of
electrolytes, the radii rs of hydrated M+,
M2+, and anions have been calcu-
lated using the equation Y, = O.82O/q,RO, where 7 is the viscosity/centipoise
and A, is the limiting equivalent conductance/cm2 equiv-l Selected
values of r , /A at 25
O C
Table 1 Valuesof rJA andPaulingcrysta1 radii
r c / A or ions in aqueous solution
Ion
rc
0.60
1973,
A. Zhogolev, Yu.
A. Kruzlyak, B . Kh. Bunyatyan, and I . V. Matyash, Teor. i
O 6 G. Singh and
P. A. Rock, J . Chem. Plzys., 1972, 57, 5556.
97 A.
eksp. Khim. , 1972, 8 ,745.
 
for comparison, With increasing temperature two phenomena affect hydration
numbers and hence radii; destruction of the free water structure, which leads
to higher hydration numbers since more water molecules become available,
and thermal motion, which acts in opposition. The decrease inY, with increas-
ing temperature is observed with Li+ and Mg2+ ions only. With larger ions
such as K+, Rbf, and Cs+ the negative hydration gradually changes to
positive hydration and Y, increases. In N-methylacetaniide, values for Y, of
Lif, Na+, and K+ ions decreased and those for R4N+ and halides increased
with increasing temperature. In structureless solvents such as nitrobenzene,
phenylacetonitrile, and dimethoxyethane, values of Y of Nas and Css ions
are independent of temperature.9s The characteristic frequencies of water
molecules in hydration complexes have been investigated by neutron inelastic
scattering. For concentrated solutions of lithium and magnesium salts and of
calcium nitrate, the relaxation time of primary water of hydration increases
and exceeds interaction time. The self-diffusion coefficients for such complexes
decrease rapidly with decreasing temperature; with increasing anion basicity;
with both increasing mass of the cation and the number of water molecules in
its hydration sphere; and with increasing mass of anions and their bonding
to their primary water of hydration. In ternary solutions significant changes
occur in the bonding and ordering or water molecules in the hydrated spheres
of the component cations, in the diffusion kinetics of the water molecules,
and in the distribution of the anions relative to the cations.QgThe hydration
ability of concentrated hydrochloric acid
(5.8 moll-l) is stronger than that of
lithium chloride. From the heats of mixing of these species with 1-6 moll-1
aqueous solutions of bivalent chlorides of M g , Ca, and Sr it appears that
HCI removes more water molecules from the first co-ordination sphere of the
metal cation than LiCl. At lower salt concentrations the free water molecules
of the solution and not those bound to the inner co-ordination sphere of the
metal transfer to Li+ (or H+j.lo0 In the LiC10,-Ca(ClQ4j2-H,Q system at
25 C, however, Ca2+ does not have a dehydrating effect on Lif since these
ions possess similar enthalpies of hydration in saturated aqueous solutions.lo1
The nuclear magnetic relaxation rates and shifts of 'Lit and 133Cs+n aqueous
solutions containing Fe3 - and various counter-anions are interpreted in terms
of a
dipolar attraction between Lis and the unpaired electrons on the Fe3+
ion, and the formation of an ion pair betweenCs+
and ferric halide complex.lo2
bar results in an enhancement
in the hydration of the ions Naf and K+ in their aqueous chloride solutions.
The enhancement is more pronounced at 20 than at 45 "C. These conclusions
B.
S .
Krumgal'z,
O 9 G. I. Saford and P. S. Leung,
US. ffice Saline W a t e r , Res. D ev el op . P r o p . R e p t . ,
looZ. G . Szabo, M. Palfalvi-Rozsahegyi, and K . Burger,
Proc. Symp. Co-ordination
Pedagog. Inst., 1971, No. 95, p. 11. (Ref.
 
Inorganic Chemistryof the Maingroup Elements
are drawn from an analysis of the literature data on the pressure dependence
of
the adiabatic cornpre~sibility.~~~o-ordination of water molecules by
alkali-metal ions induces a distortion which shows up in the i.r. spectrum
of
the molecule.l@This is more far-reaching, however, and the water structure
is destroyed by electrolytes in the order KI
>
>
LiOH. Obviously the effect of any
cation on the structure of water depends greatly on its anion and vice versa.lo5
Calculations based on electrochemical data indicated that the primary
hydration number of caesium iodide is zero. Thus, the standard free energies
of transfer of Cs+
and I- from water to aqueous organic solvents should be
very low.lo6This is compatible with the order of extraction of alkali-metal
halides from aqueous solution byp-alkylphenols. Cs, Rb, and K are extracted
from chloride solutions at pH 9.5-13.8 by
0.5-4
mol
1-1
of
ROM,-
nROH
in the organic phase, n being 1.2, 1.2, and 2.2 for Cs, Rb, and K ,
respectively. The distribution ratios between the organic and aqueous phases
from the aqueous solutions containing <0.15 mol 1-1 chlorides are D,, >
D,, >
DK. On extraction, however, from solutions containing >0.15
mol 1-1 chlorides the sequence is reversed. The extraction from low metal
concentrations is governed mainly by the dehydration energy (low for
caesium) of the cations. In concentrated solutions, the dehydration energy of
the ions is levelledoff to some extent, and the formation energy of the phenol-
ate becomes the prominent factor influencing the extraction.lo7
As
the dielectric constant of the solvent in which the alkali-metal ion is
dissolved decreases, then the degree of pairing with its counter-ion increases.
Thus the equivalent conductance of 2 x 10-4-1.5 x 10-2mol 1-1 lithium
chloride solutions in water-sulpholane mixtures
(0-100
D
from
43
steadily decreases
as sulpholane is added, with noticeable ion-pair association for mixtures with
D
62.1°s Similarly with caesium bromide solutions.10g The diffusion
coefficients of the ions Cs+ and Br- in water-dioxan mixtures decrease with
increasing salt concentration. When interpreted on a model based on ap-
plicable ionic interaction theory, this is evidence for progressive ion-pair
formation. The vapour pressures above salt solutions (50.05moll-l) in
water-tetrahydrofuran mixtures are used to obtain the rate of change of the
standard chemical potential
composition and to obtain solvation numbers. The compounds NaNO,,
lo30 .Ya. Sam oilov, A . L. Seifer, and N . A . N evolina, Zltur. s trukt . Khim., 1973, 14, 360.
lo*L. I. Gudim, Dopooidi Akad. Nauk Ukrain. R . S . R . , Ser . A , 1972, 34, 904.
lo5A.
A.
Lapshin, Zhur. strukt. Khim., 1 9 7 3 , 1 4 , 2 1 .
Io6 T . Mussini and P. Longhi, Chimica e
Industria, 1972, 54, 1093.
L.
I. Pokrovskaya, A . M . Reznik, and V. E. Plyushchev, Izuest. Akad.
lo*
G. Petrella, M. Castagnolo, A. Sacco, and L. Lasalandra, Z . Naturforsch., 1972,
27a,
lo9H. Latrous, P. Turq, and M . Chemla, J . Chim. phys . , 1972, 69, 1644.
Nauk Kazakh. S.S.R. , Ser . khim., 1972, 2 2 ,
2 8 .
<
0).l1O
In aqueous solutions containing sulpholane (1.01-95.33 %) at 25 'C, the
limiting equivalent conductances]cm2e q u i r l L2-l for KClO,, the association
constantslequiv 1 1 for ion pairing (K+
+ ClO, = K+ClQT), and the ion-ion
contact distances/A in the solvated ion pair, respectively are 136.17, 0.69, 3.9
(1.01 ); 129.35, 0.60, 4.2 (5.62%); 70.14, 1.27, 4.6 (49.87%); 33.843, 5.23,
3.2 (80.78 ); 17.894, 18.25, 2.1 (95.33 sulpholane).lll I n aqueous 10.15
dioxan solutions at 25 O C , the preferential solvation of NaCl and NaOH by
water molecules and of NaBPH, and Ph,AsCl by dioxan molecules is indi-
cated by changes in the standard chemical potentials (Ape) for transfer of the
electrolytes from pure water into the mixed solvent. The A p o values are
obtained from measurements, by gas-phase chromatography, of the com-
position of the vapour above the solutions. The selectivities (dpo]dx,; x1 =
mole fraction water) for dioxan solvation of the ions decrease in the order
BPh, > AsPhz >> Na+
> OH-, where Na+ is approximately equally
solvated by water and dioxan molecules. With decreasing xl, he differences
between the d,uo/dx, values decrease markeclly.l12 The limiting electrical
conductivities of 10-4-2 x
DMF mixtures show minima near 25 and 4
DMF. These could be
attributed to either an exchange of the smaller water molecules for larger
DMF molecules in the solvent sheath of the ions, or a change in the solvent
structure, i .e. due to the appearance of DMF,3H,0.113 Sodium ion pref-
erentially forms long-lived complexes with DMSO and DMF in mixtures of
either of these solvents with ethanenitrile or propanone. This is deduced from
the quadrupolar broadening of the 23Nan.m.r. in these solutions. In mixtures
of water or methanol with ethanenitrile or propanone, however, no long-
lived complexes are 0b ~erved. l~~
Some aqueous systems that have been investigated are listed in Table 2.115-160
Non-aqueousSolvation.-The absorption spectra of solutions of lithium and
of potassium in liquid ammonia at -7OOC have been measured using a
spectrophotometric cell incorporating sapphire windows. The envelope of
strongly overlapping bands in the 3200cm-l region of the pure solvent is
resolved into components which are assigned to 2v4, vl and v in increasing
energy. The
with increasing concentration of both lithium and potassium metal. Solvating
molecules have stretching vibrations of lower energy than bulk solvent
molecules. The solvated molecules, therefore, give an additional band
displaced from that for the solvent. The shift is independent of the nature of
the metal over the concentration range 5 x 10-,-5
x
10-2moll-1. This
C. Treiner, J. F. Bocquet, and M . Chemla, J. Chim. phys. , 1973, 70, 472.
A. D'Aprano and I. D. Donato, G a z z e f f a , 1972, 102 ,92 3.
112 S . Villermaux, V. Baudot, and J. J. Delpuech,
Bull.
SOC.
Sci . Hung., 1972, 17, 49 .
 
Zhur. neorg. Khim.,
1 1 19.
Zhu r. neorg. Khim., 1973,18,246.
I N .
Lepeshkov and
R. Vilcu and F. Irinei,
Bull. SOC. him. belges, 1972, 81, 479.
L.
Zhu r. priklad. Khim ., 1973, 46, 648.
M. G. Manvelyan, V. D. Galstyan, E. A. Sayamyan, and A. G. Alakhanyan,
Armyan.
khim. Zhur., 1 9 7 2 , 2 5 , 8 4 0 .
Ya.
S.
Shenkin,
S.
Zhur. neorg. Khim., 1973,
Inst., 1971,
N o .
N o . 7B905.)
P.
F.
Zssled. 061. Neorg. Fiz. Khim., 1971, 143. (Ref.
Zhur., Khim.,
I. E
A. S.
Zhur. neorg. Khim.,
1 9 7 2 , 1 7 , 2 5 2 6 .
M . A. Durymanova and A. E. Telepneva,
Zhu r. priklad. Khim., 197 2,45 , 1610.
V . . Mikheeva.
S. M.
Ser.
Nippon K aisui Gakkai-Shi,
1972,
No.
266,
p.
78.
S.
Poletaev,
985.
Zhur. priklad . Khim., 1972,
136
137
13
A. M . Babenko, Zhur. neorg. Khim., 1972,17, 3059.
A.
G.
Components
Compounds
Ref.
;
K HaP04--NH4H,P04
K 2 ° ~ B 2 0 3 ~ 2 H 2 0
K,SO,-CsNH,
142
K,SO,-CsNH,-MnSO,
143
K,S04-(NH4)PS04-ZnS04 (N H 4)z S0 4,Z nS 84 ,6H 20 ; 144
K,SO4-CdSO, 3CdSQ4,8H20;K2S0,,3CdS04,5M20; 145
- 149
K,SO,,ZnSO, ,6H 2 0
-
159
CsI-CsBr
- 160
138 V. A. Tatarinov, Uch. Zap., Yaroslav. Gos. Pedagog. Inst., 1971,
N o .
3368.
A . P. Solov’ev, Uch. Zap., Mord. Univ., 1971, N o . 81, p. 39. Ref. Zhur., Khim., 1972
Abstr.
0. V.
Zhur. neorg. Khim., 197 2,17, 3379.
143 N. N. Runov, Uch. Zap., Yaroslav. Gos. Pedagog. Inst., 1971, No. 95,
p.
neorg. Khim., 1972, 17, 3361.
lP6
V.
1 8 , 1682.
146 V. K. Bishimbaev, I. N. Shokin, and A. G. Kuznetsova, Khim. Khim. Tekhnol.,
Alma-Afa),1971, N o . 12, p. 203. Ref. Zhur., Khim., 1972,
Abstr. N o . 18B766.)
14’ J. Balej. Coll. Czech. Chem. Comm., 1972, 37, 3855.
148
149
K . G. Myakishev and A . P. Kostin, Zhur. neorg. Khim., 1973, 18, 271.
160 Yu. M. Timoshenko, Zhur. neorg. Khim., 1973, 1 8 , 854.
151 T. V.
Mozharova, V. A. Borovaya, and E. N. Pavlyuchenko, Zhur. priklad. Khim.,
1972,45, 1872.
lSzV. I . Vereshchagina, V. N. Derkacheva, L. F. Shulyak, and L. V. Zolotareva, Zhur.
neorg.
Khim., 1973, 1 8, 507.
163 M. Nagatani and R. Kubo, Kagoshima Daigaku Kogakubu Kenkyu Hokoku, 1972,79.
L. V . Savel’eva, S. B. Stepina,
V. E. Plyushchev, and A. P. Rysev, Izoest.
V.
U.
Khim. i khim. Tekhnol., 1972,15, 1605.
155 N. Nishimura, T. Higashiyama, S. Yamamoto, and S. Hasegawa, Nippon Kagaku
Kaishi, 1973, 1059.
156
A. A. Maksimenko and V. Shevchuk, Zhur. neorg. Khim., 197 3,18, 1401.
15’ G. Gode and
G.
160
Ref. Zhur., Khim., 1972, Abstr. No. 7B927.)
3
Inorganic Chemistry of the Main-group Elements
contrasts with the cation dependence of this band in solutions of salts in
ammonia, where the absorption is attributed to Mf NH, dipole inter-
actions. Also the magnitude of the
v 3
greater for
than for Mf (400 times greater for K than for
KI).
These
results are attributed to the formation, with increasing metal concentration,
of a ncw solvated species of cation which incorporates the solvated electron,
i.e. a form of solvated cation-electron pair.161 The magnetic susceptibility of
solutions of alkali metals in liquid ammonia has been measured over the
concentration range where the solutions show a progressive transition towards
the metallic state. The general transport properties of metal-NH, solutions
have been analysed and a model is proposed for the mechanism of the transi-
tion to the metallic state.162This transition occurs at ca. 1-9 mole per cent
metal. Over this region metal-ammonia solutions exhibit rapid variations
with concentration of their static and transport properties. Concentrated
(>10mole per mole) metal-ammonia solutions are more like liquid metals
and show electronic transport. It is proposed that the intermediate 1-9
mol (mol)-l region is microscopically inhomogeneous, so that the solution
consists of a mixture of metallic clusters of a mean concentration > 9 mol
(mol)-l together with small cation-electron diamagnetic c0mp1exes.l~~
Photoelectron emission by potassium (0.002-3 moll-l) and sodium
(0.013-
0.21 and 7.2mol1-l) in liquid ammonia at -6OOC has been investigated
from 1.55 to 5.4 eV. The spectrum is composed of the previously reported
band at
ca.
4.6 eV. The highly asym-
metric low-energy band, 3.2 eV, is assigned to photoelectron emission by
solvated electrons. The 4.6 eV symmetrical photoelectron emission band
results from bound-bound transitions followed by a u t o i ~ n i z a t i o n . ~ ~ ~lkali
metals dissolved in liquid ammonia react with zirconium disulphide
ZrSz in
sealed tubes at 20°C to form several phases. The alkali metals partially or
wholly occupy the layers of empty sites in the host structure to form NaZrS,
or KZrSz when the ammonia solutions contain high concentrations of
alkali metal. At lower concentrations, only one layer out of two is occupied.
No reduction of ZrS,, occurs. At very high concentrations, sodium occupies
octahedral sites whereas potassium occupies either octahedral or prismatic
sites. A comparison of the X-ray diffraction data of NaZrS, with those of ZrS,
indicates an increase in the lattice parameter
c
of
K)
0.56-1, 0.08-0.10, and 0.03-0.04. The lattice
parameters of these compounds are compared with those of the sodium and
P.
F. Rusch and J. J. Lagowski, J. Phys. Chern., 1973,77,210.
J P. Lelieur, Report 1972, CEA-R-4333.
163 J. Jortner and
G 4
H. Aulich, B. Baron, P. Delahay, and R. Lugo, U.S. Nat. Tech. Inform. Service, A D
165 J. Cousseau, L. Trichet, and
J.
x
monotonically as the Li atoms progressively fill the octahedral sites
of
the
of
the different types of phases are classified
on the basis of the co-ordination number of the intercalated alkali-metal
atom, the periodicity of the occupation of the available sites between
successive (STiS) layers of the host structure, and the mode of succession of
the sulphide layers.166 Single crystals of MoS, can also be intercalated with
alkali metals (Li, Na, K , Rb, and Cs) by the liquid ammonia technique.
X-Ray results show considerable expansion of the c-axis after intercalation
(Ac; 6.745, 2.704, 4.286, 4.899, and 7.312 A for Li,.,MoS,, Na,.,MoS,,
KO,MoS,, Rb,, ,MoS,, and Cs,.
,MoS2, respectively). All intercalated crys aIs
are superconducting, which is attributed to electron transfer from the alkali
metal to an empty band of M 0 S p 7Potassium naphthalenide solutions also
react with MS, (M
W) to give K,MS2 x = 0.49-0.76). With
sodium naphthalenide, the first step in the reaction gives Na,MS,, but
lithium naphthalenide and subsequently sodium naphthalenide reduce the
sulphides to the respective metals.16*
The absorption spectra of solutions of alkali metals in various amines
display up to two bands; one at 1400 nm, which is independent of the metal
and which is attributed to a solvated electron, and a second band which peaks
between 660 and 1OOOnm which is independent of the metal and which is
attributed to a metal-containing species. Evidence has been accumulating
which suggests that this is a metal anion, M-,
(J . Php. Chem.,
1971, 75,
3092). The 660-1000 nm band obeys the criterion for a charge-transfer-to-
solvent transition that the position of the peak, Y, be a linear function of the
temperature coefficient, dv/dT. Thus the existence of the anion, M-, is
proposed in solutions
(Me,N),PO, Et,O, PrizO, THF, MeOCH,CH,OMe, and diglyme, in which
the solubility of the metal has been enhanced by employing cyclic polyethers
of the crown and cryptate classes to complex the alkali-metal cations.170
For a spherically symmetrical Na species with J =
1
a magnetic moment of
1.0 B.M. is to be expected. Experimentally, values ranging from 0.5 to 1.3
B.M. are found for Na- in EtNH,.171
Modern advances in solvation theory are reviewed.17, A second critical
review of the thermodynamic functions and crystallographic data of some solid
solvates (e .g . NaI,3MeOH; LiC1,py; and CoC1,,6NH3) shows that the relative
acid-base properties of the constituent cation, anion, and solvate molecule
can well be described by the HSAB (hard and soft acids and bases) concept.173
lG6 J. Bichon, M. Danot, and J. Rouxel, Compt. rend., 1973, 276, C, 283.
lG7 R. B. Somoano, V. Hadek, and A. Rembaum, J . Chem.
Phys., 1973, 58 , 697.
lG8 E. Bayer and W . Ruedorff, Z. Naturforsch, 1972, 27b, 1336.
lGS
K.
170
M . T. Lok, F. J. Tehan, and J. L. Dye, J . Phys.
Chem., 1972, 76, 2975.
Letters, 1973, 20, 371.
300.
173
C.
 
22 Inorganic Chemistry of the Main-group Elements
The enthalpies of ammoniation of gaseous cations relative to that of Na+, and
of gaseous anions and the electron relative to that of I-, have been determined
from an analysis of enthalpies of solution of salts and metals in liquid
ammonia. Selected experimental enthalpies of formation, AH,/kcal m o F , of
ion pairs in liquid ammonia at -33 "C are shown in Table 3. The AHf of a
Table3 Experimental values of AH,/kcal mol-l for ion pa irs in liquid ammonia
at
Li
Naf
K +
Rbf
Csf
Ca2+
Sr2+
Ba2+
9.7
-20.7
-19.0
certain ion pair, e.g. Na+ C1-, is found by combining the standard enthalpy of
formation of NaCl (-98.2 kcal mol-l) with its heat of solution in liquid
ammonia (-1.5 kcal mol-l). AH, of an ion-electron pair is merely the heat
of solution of the metal in liquid ammonia. Values in parenthesis are extrap-
olated. It is demonstrated that the relative ammoniation enthalpies of the
individual ions can be obtained from the Born equation if an effective radius
which is 0.61 13 greater than the crystalline radius is assumed for each ion. In
addition, the absolute ammoniation enthalpies for the gaseous ions are
evaluated at -33OC and compared with those deduced from the Born
equation. Values of AH/kcal mol-l for selected ions are shown in Table 4.174
Table 4
Ion Li+ Na+ K + Rb+ Cs+
Ca2+
Born
Ion Sr2+
Vibrational spectra, mostly from Raman scattering, are reported for LiNO,
and NH,NO, solutions in liquid ammonia ranging from dilute to the com-
positions LiNO,,l .6NH3 and NH,NO,,l .ONH, at ambient temperatures,
Large changes in the N-H stretching region are correlated with effects of the
electrolytes on the solvent structure. Low-frequency bands associated with
Li-NH, modes are observed. Some special features are indicative of both
114 N.M . Senozan, J .
of Group I 23
solvent-separated and contact ion pairs involving Li+ and NO; ions.175 The
enthalpiesof solution at 25 OC for the salts LiCl, LiNO,, NaBr, Nal, NaNO,,
KBr, KI, KNO,, RbNO,, and CaNO, in DMF have been measured calori-
metrically and the enthalpies of solvation
of the component ions ca lc~1at ed .l~~
Conductance measurements have been made at 25 "C for solutions
of
the
potassium picrate, and Pr,NBr in hexamethylphosphortriamide (HMPT).
Potassium salts are more conducting than the corresponding sodium salts
and the ions migrate independently in HMPT. Alkali-metal cations are sol-
vated by ca.2HMPT moleculeswhereas the anions can be considered as naked.
All of the salts, except KNO, and Pr,NBr, are completely dissociated in this
s01vent.l~~n solutions of perchlorates in hexamethylphosphoramide HMPA),
the ionic equivalent conductivities of cations obtained on the assumption
that the conductivity of the perchlorate ion is 15.5 are Li+ 5.2, Na+
5.8 ,
K+6.1, Rb+6.1, (36.4, Ca2+8.6, Sr2+8.6, Ba2+8.4, Me,N+7.9, Et,N+9.3,
Pr,N+ 6.8, Bu,N+ 5.9, and Hex,N+ 4.5 cm2 equ i r l L2-l. The Stokes' Law
radius of a univalent cation is smallest for Et,N+. The solvation number
obtained from the effective ionic radius of Robinson and Stokes (1965) is
1
for alkaline-earth ions.178The association
of the ions Lif, Na+, and Mg2+with perchlorate in methyl cyanide has been
studied spectroscopically. On increasing the salt concentration, additional
bands appear in the i.r. spectrum of the C 1 0 ~ons owing
to the formation of
M+
log ion pairs: for solutions of Na+ these are at 1119, for Li+ at 1070
and 1132, and for Mg2+ at 1040 and 1154 cm-1.179 The enthalpies of solution
of the salts LiClO, and LiNO, in isoamyl alcohol decrease with increasing
concentration, especially 50.1 moll-l. At
25 C,
he values at infinite
dilution are -9.90 kcal mol-l for LiClO, and -2.75 kcal mol-l for LiNO,.
The enthalpies of solution of the salts in the alcohol are more exothermic
than those in water,ls0A quantitative study of n.m.r. chemical-shift variation
as a function of salt (LiCl, LiBr, LiI, NaI , and KI), molality, and temperature
gives values for effective solvation numbers of 4 for lithium salts and 6 for
sodium and potassium iodides.lS1Conductivity data for solutions of sodium
iodide in acetone from -50 to +50
OC
provide ionic association constants,
which with increasing temperature show a rise that is explained by the
endothermic effect of ion de-solvation. Positive values
of
indicate that disorder increases during the ionic association of sodium
iodide.182The solvation numbers of Na+, K+, and I- ions at 25 "C, calculated
176 D.
V.
A . Zverev and G. A. Krestov, Teor . Rastuorov, 1971, 148. ( R e f .Zhur., Khim., 1972,
Abstr. No. 6B1205.)
P. Bruno, M. R . Della Monica, and Righetti, J . Phys. Chem., 1973, 77, 1258.
178 T. Fujinaga,
K .
Izutsu, and S . Sakura, N i p p o n Kagaku Kaishi, 1973, 191.
179 I. S. P erelygin and M . A. Klimchuk, Ural. Konf. Spektrosk., 7th, 1971, No. 2, p.
125.
180
P.
I. Klement'eva, Zhur . f i z . Khim., 1973, 47, 692.
lE1 A. Lindheimer and B. Brun. J . chim. Phys.,
1972,
69, 1454.
B. S. Krumgal'z and Y u. I. Gerzhberg, Zhur. obshchei Khim., 1973, 43,
462.
from conductance data for solutions of NaI,
KI,
and KSCN in acetone,
when compared with solvation numbers for the same ions in ethanol and
water, illustrate that the intrinsic structure of the solvent plays a major role
in the solvation process.ls3 The i.r. spectra of solutions of Li, Na, K, Rb,
and Cs salts in sulpholane show absorptions which are independent of the
anion. Solvation appears to take place in a manner similar to that observed
for D M S O solutions but sulpholane, a dipolar solvent with a high (44)
dielectric constant, is a far weaker donor. I t is proposed that solvent-separated
ion pairs are formed in sulpholane solutions.1s4 The electrical conductivities
of solutions of the Li, Na, K , and Rb salts of 2,4-dinitrophenol in THF at
25 O C have been measured with and without added triphenylphosphine oxide.
Ion-pair dissociation constants of the salts are derived from variation of the
conductivities with salt concentration. These constants increase with increas-
ing cation size, as expected for contact ion pairs. Cation-ligand association
constants,
KL,
are derived from the increases in the conductivities due to
added
Ph,PO.
Values for KLare for Li 3500, Na 250, K 87,
and Rb
I mol-l.
This is just the order of association expected from the cationic radii. An ion-
dipole model, modified to take into account the influence of surrounding
polarizable solvent, is able to explain the re~u1ts .l~~on-solvent interactions
are detected in cyclohexane solutions of NaAlBu, by means of differential
vapour pressure analysis. The observed departure from ideality is attributed
to an aggregation process. The apparent degree of aggregation varies from
pairs of ions at ca. 0.0025 niol
1 1
ca.
2 moll-l. Addition of THF to the salt solutions lowers the apparent molecular
weight of the solute. The magnitude of this effect increases with both an
increase in the ratio of THF:NaAlBu, and with an increase in NaAlBu,
concentration. This is attributed to a loss
of aggregate stability in both
instances.ls6 Thermochemical measurements on solutions
of
of
solution.
With solutions of NaI, an increase in salt concentration causes an increase in
the exothermic effect attributed to the ability of NaI to form solvates with
Me,SQ. The energetics of solvation of Na+ ions are appreciably greater than
those for
0.032 mol 1-1 metal acetate at 25
O C
M+OAc-
PKB
f 0.10 for M =
K, Rb, and Cs,
respectively. For M = Li, Na, K, Rb, and Cs the linear plot of pKBus.
log a3
of
+1, where a is the radius of M+.ls8 The optical absorption
spectra have been measured for solutions of triphenylene reduced with alkali
lS3 Yu. I. Gerzhberg,
B. S . Krumgal'z, and D. G. Traber, Teor. Rastvoroo, 1971, 115
lE4 T. L. Buxton and J. A.
Caruso,
R. Gilkerson, J .
1973, 77, 1421.
lS6 J. H . Muller and M . C. Day, J . Phys. Chem.,
1972,
76,
3472.
lS7 0 . I . Kyabchenko, M. L. Klyueva, K . P. Mishchenko, and N. P.
Novoselov,
Zhur.
lg8
1972,
11,
2556.
5B1485.)
Group I 25
metals in 2-methyltetrahydrofuran from -155 to +75 'C. Several ion pairs
are observed; a linear correlation between v and reciprocal cation radius is
observed for both solvated and contact ion pairs. The low-temperature spectra
show a clear equilibrium between two types of ion pairs at high triphenyl-
ene anion concentrations. The high-temperature spectra show a gradual
increase in as the temperature is increased, which is interpreted on the
basis of an equilibrium between solvated and contact ion pairs.lE9
5 Compounds
Molecules or Complex Ions
Certain naturally occurring antibiotics show a high degree of cation specificity
for the alkali-metal ions in biological systems. Selectivity is in the order
Li+
<<
Naf, Csf <Rb+, K+. Significantly, this is the order of the magnitude
of the stability constants for their complexes with alkali-metal ions. Simula-
tion of the selective behaviour in simpler systems than those of naturally
occurring antibiotics can be achieved with cyclic polyethers. The structures of
some polyether-alkali-metal salt complexes are already known. The first
compounds to be isolated have the stoicheiometry 1:1,
i.e.
one mole of
cyclic ether to one mole of metal salt. The stability constant is greatest when
the metal cation fits the centre of the ring of oxygen atoms in the organic
molecule. Given the optimum match between size of hole and size of cation,
K + forms the strongest complexes. This can be explained on an ionic model;
the cation is not small enough for its hydration energy to be
so
large that
complexing with the cyclic polyether is precluded and, moreover, the cation
is not large enough to prevent strong ion-dipole interaction with the oxygen-
ring in the polyether. Recently, a meta1:ligand ratio of 2: 1 has been found
in (KCNS),L,, and complexes of ratios 1 2 and
2:3
if
the
cation is too large to fit the hole in the ring. The most recent study is on the
crystal structure of the 1 : 2 complex KI(benzo-15-crown), formed between KI
and
2,3,5,6,8,9,11,12-octahydro-l,4,7,10,13-benzopentaoxacyclo-pentadecin,
a, b
17.84, c = 9.750 A, d(X-ray) 1.51 for 2 = 4, d(expt) = 1.51, space group
P4/n.
Bond lengths/A are shown in (l), the standard deviations in the K-0
bonds being 0.007-0.008
A. K+ is sandwiched between two organic molecules
and is co-ordinated to ten oxygen atoms in an irregular pentagonal prism
as shown in Figure
The complex cations are separated by iodide ions in an
arrangement resembling that in CsCl. There is no interaction between K+
and I-.190 The equivalent conductance for NaCl, KCl, CsCl, NaBPh,,
KBPh,, and CsBPh, in methanol and acetonitrile solutions containing the
cyclic polyethers dibenzo-1%crown
have been
measured at 10 and 25 "C and in salt concentrations up t o 0.004 molld1. The
189 R . M. Arick, J. A . M . Van Broekhoven, F. W . Pijpers, and E. DeBoer,
J .
P. R . Mallinson and M . R. Truter, J . C. S .
Perkiiz
A3
decrease in equivalent conductance caused by the addition of a cyclic poly-
ether is analysed to give association constants and equivalent conductances
of the cation-polyether complexes. The selectivity observed for complexing
is a function not only of ionic size but also
of
2; R1= R2 = H, X =
0 or CH,;
benzo, R1= H, X = 0 ;RIR1
R2R2 = benzo, X = 0) and the macrocycle (3) with Naf, Kf, and Ba2f
ions have been determined in 95:
MeOH-H20 mixtures and compared with
lS1 D. F.
Evans, S . L. Wellington, J. A. Nadis, and E.L. Cussler,
J .
M e N N M e
those for nonactin. The stabilities of the cryptates are controlled by the thick-
ness, s, of the organic ligand layer separating the complexed cation from the
solvent. Increasing S decreases the attraction of the complexed cation for the
polar solvent,
e.g. H,O, and destabilizes the complex. Selectivity for Ba2+
Over K+ is introduced since, though of similar size, the attraction is stronger
for bivalent than univalent cations. The presence of fewer binding sites, i.e.
oxygen atoms, in the ligand also destabilizes the c0mplex.1~~otassium
cryptate, (Cl,H3,N,0,K)+I-, crystallizes in the monoclinic system with four
molecules in a unit cell
of
b = 8.334, c = 20.638 A,
=
and is located at the centre of the molecular cavity
of
the bicyclic ligand,
surrounded by all eight heteroatoms of the ligand as shown in Figure 2.
N
potassium
cryp tate,
The six oxygen atoms comprise a prism distorted towards an antiprism (the
torsion angle between the two triangular faces is 22'53'). The two nitrogen
atoms lie above and below, respectively, the centres of the two triangular
faces on a line through the central K atom. Using the given notation, the
K-N, K-Q(4), K--0(21), and K-0(7) distances are 2.874, 2.776, 2.790,
and 2.789A, respectively. The rest of the complex has been omitted for
clarity but each
is linked through-(CH,),-to its three adjacent oxygen
atoms. Pairs of oxygen atoms in top and bottom triangular faces are similarly
luaB.Dietrich, J. M. Lehn, and J. P.
Sauvage, J . C. S . Chem. Comm . , 1973, 15.
 
linked, e.g. o(7)-(CH2)2-o(4). The K-I distances are all greater than
6.5A lg3Analogous complex cations are found in rubidium and caesium
cryptates, C18H,,N206,MSCN,H20. This indicates the flexibility
of
the
organic bicyclic ligand, which adapts its molecular cavity to the ionic radius
of the metal. These isomorphous monoclinic compounds with space group
P2/c and four formula units per cell have dimensions
a =
NaI, the structure
to
a,
b =
8.630,
c =
18.570&
7 =
of
ionic
I-
of 50 OC, separate 23Nan.m.r. signals can
be obtained for the solvated Naf and for the Na+ ion complexed in a cryptate.
The doublet coalesces and the exchange narrows with increasing temperature.
The first-order rate constant for complex cation dissociation in ethylene-
diamine varies from 300s-l at 34OC to 2560s-I at 75
O C .
An activation
of
dissociation is similar
to that found for aqueous solutions.196 The final compound in this series,
lithium cryptate, C14H2,N204,LiI, s tetragonal with space group p4,2,2,,
with 2 =
= 24.361 A. The Li+
cation is again enclosed within the molecular cage but the co-ordinating
ligand in (4) contains two fewer
0
(4) M =
K, Rb, and Cs. The small Li+ ion occupies a
severely distorted octahedral environment comprising two nitrogen atoms at
2.288, two
ion.lS7The crystal structures of complexes of alkali-metal thiocyanates with
the antibiotic tetranactin have been determined. The biological property of
nonactins is to transport alkali-metal ions through biological membranes in
the form
complexes. The uncomplexed tetranactin molecule is fairly flat,
elongated, and twisted; the molecular outline resembles that of a propeller.
lS3 D. Moras, B. Metz, and R. Weiss,
Acta Cryst., 1973, 293, 383.
lS4 D. B. Metz, and
Acta Cryst . , 1973, 29B, 388.
lg5
lg6
Amer.
 
Group
I
29
On complexing with alkali-metal ions the molecule curls up to enclose the
cation in an eight-membered distorted cube of oxygen atoms (labelled 4, 54,
12, 62, 17,67 ,22, and 72, according with the original article) as shown in (6).
The complex compounds are all monoclinic, with the properties shown in
Table
5.1g8
Table 5 Properties of complexes of tetranactin with alkali-metal ions
Coiilplex
Na
K
Rb
Space
Form I I1
Other complexes involving thiocyanate are those of LiSCN with 1,5,9,13-
tetraoxacyclohexadecane and its
2
moles
of
LiSCN or
1 mole of LiBr or LiI, as shown by changes in the
lH
ligand.lgQThe cyclic quadridentate secondary amine 1,4,8,1 -tetra-azacyclo-
tetradecane (cyclam) also reacts with lithiumsalts to give complexes of the type
(LiX),(cyclam) when
X = Br, I, or ClO,. A 1:1 complex is also produced
with LiClO,. The i.r. spectra of the 2: 1 complexes are similar to each other
but differ from that of free cyclam. (LiClO,),(cyclam) and LiI(cyc1am) are
triclinic, with
2 = 2 ,
so that the N atoms are crystallographically equivalent and probably tetra-
hedrally co-ordinated to Li+.,O0 The interaction of sucrose with alkali-metal
l9* Y . itaka, T. Sakamaki, and Y . Nawata, Chem. Letters, 1972, 1225.
Ig9 J. Dale and J. Krane, J .
C.
1972
1012.
D. E. Fenton, C . Nave, and M . R. Truter, J. C.
S. Chem.
 
halides
L =
4-1,
and rn = 1-3. The complexing ability of the cations decreases in the order
Na+
NaCl has been prepared. The spectrum shows important differences between
the complex and the free ligand at 650 and 1200 cm-l. X-Ray data show octa-
hedral geometry round both Na+ and C1-ions.202 The following addition com-
pounds
of
,
8-10t~.~O~he crystal structure of the complex triphenylmethyl-lithium tetra-
methylenediamine, Ph3C- Li+ Me2N(CH,),NMe2, has been determined from
X-ray data. The compound is monoclinic, space group
P2, / c ,
1.07.
The structure can be described as a contact ion pair consisting of a lithium
cation co-ordinated to a tmeda group and a triphenylmethyl carbanion (7).
Each Li atom is co-ordinated to two N atoms of the bidentate chelate and to
one ?r-carbanion. Each Li-N distance is 2.10 A. The Li atom is not located
directly over any one carbon atom but has four close contacts to the carb-
anion; 2.49 and 2.51
to two C atoms on one phenyl group, one close
contact of 2.54A to one C atom of a second phenyl group, and one 2.23 A
distance to the central carbon atom. The twist angles of the phenyl rings
depend on their interaction with the Li atom, giving an overall propeller
geometry to the triphenylmethyl carbanion. I t is not possible to explain the
stereochemistry by two-centre 0 or ionic interactions alone, and a delocalized
bonding mechanism which involves all the 2s and 2p orbitals of the Li
201 A. J. Dangre,
J . Univ. Poon a, Sci. Techiiol., 1972, No. 42, p.
131.
ma F. Klema, Mitt. Che m. Forschungsinst. Wirt. Oeste rr. Oesterr. Kun ststojinst., 1972,26,
Io3
1961.
101.
atom is inv0ked.2~~here are distinct similarities between this compound
and [K(MeOCH2CH2),0] Ce(C&)2]. The structure of the latter consists of
an ether-co-ordinated potassium cation [K(MeOCH,CH,)20]f combined in a
contact ion pair with a discrete [Ce(C,H,)]- anion (8). One side of the cation
is
co-ordinated by the complexing ether diglyme, which has normal bond
distances and angles. The opposite side is co-ordinated by one of the two
cyclo-octatetraenyl dianion rings in the structure. The potassium lies on the
crystallographic mirror plane. The two independent K-0 distances from the
0
atoms in the ether are 2.792 and 2.741 A and the five independent K-C
distances between the K+ ion and the C,H;- ring average 3.16 A. The salt
crystallizes from diglyme as large green plates for which the crystal data are
orthorhombic, space group Pnma, a
=
d(obs) =
=
bis(salicylideneiminato)] have been prepared from the components in THF.
Crystals of the cobalt compound are monoclinic, space group P 2 , / c , with
a =
14.415,
b
= 14.329,
c =
cations and [BPhJ anions. The anions, held together by phenyl-phenyl
interactions, are arranged in planes with the complex cation between, as
shown in Figure 3. The sodium atom has approximate octahedral co-ordina-
tion, being surrounded by six 0 atoms, two from each Co(sa1en) group
(Na-0 distances ranging from 2.391 to 2.482 A) and one from each of the
two THF groups (Na-0 distances 2.35 and 2.37
A)
in the complex cation,
which has approximately C2symmetry. This is an example of an alkali-metal
ion complexed to a co-ordination compound. The following adducts are
2 0 4
J. J.
Brooks, W. Rhine, G . D. Stucky, J . Amer. Chem. SOC., 972, 94, 7346.
a05 K. 0. Hodgson
and K.
 
e o
bond distances.
similarly prepared:
X =
SCN); 3-MeOCo(salen),-
LiBr,THF; and ~-M~O.CO(S~I~~),L~SCN,TKF.~~he crystal structure of ru-
bidium ammonium hydrogen fluorocitrate dihydrate Rb(NH4)H(C*H40,F),-
2H,O has been determined. The crystals are triclinic, of space group PI,
2
c = 7.516 A, a = 102.80 , /3 = 115.90°,
y = 80.50°,
to
as shown in (9). The fluorocitrate ion
forms
0 of the carboxylic acid group adjacent to
F,
the hydroxy-group. The central carboxy-group is not
involved in this chelation, contrary to the case for citrates and isocitrates.
The fluorine, by virtue of its electronegativity, in
no
H
atom but is chelated to the metal. Overall there are six 0 and
two F
atoms
round any Rb+ ion but only five0 atoms round the NH$ cation; F is always
206 C. Floriani, F. Calderazzo, and L. Randaccio, J .
C. S.
co-ordinated to Rb+ rather than to
N H i
The compound is monoclinic, space group C2/c, a =
32.81, b = 5.47,
c = 9.95 A, , = 95"54', 2 = 8. The rubidium ion is tetrahedrally (distorted)
co-ordinated by four oxygen atoms at distances 2.86 and 2.88 A . Each 0
belongs to a different homophthalic acid molecule. The two acid molecules
linked with one Rb+ ion form a chelate complex which is repeated throughout
the structure. Linking between the acid molecules
is
hydrogen-bonds between oxygen atoms of
carboxy-groups.208n potassium boromalate monohydrate KBC8H8O1o,H@,
the alkali-metal ion is again co-ordinated by the oxygen atoms of the carboxy-
groups but in this case K+ exhibits five-co-ordination209The thermal de-
composition of alkali-metal halogenoacetates has been studied by gas
chromatography, d.t.a., and thermogravimetric, X-ray, and chemical
analyses. The decomposition of the salts (M = Li, Na, K , or NH,) proceeds
according to CH,ClCO,M -H,O + CO
+
the phase equilibria between alkali-metal chloroacetates (MA)-CH,ClC