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Compounds of Alkali Metal Anions.

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Compounds of Alkali Metal Anions
By James L. Dye"'
Anions of sodium. potassium, rubidium, and cesium are stable both in suitable solvents and in
crystalline solids. The latter can be prepared either by cooling a saturated solution or by rapid
solvent evaporation. Thermodynamic arguments show that alkali metal anions can probably
exist in saturated solutions of the alkali metals in any compatible solvent. but that below saturation, dissociation into the cation and solvated electrons is favored in highly polar solvents
such as ammonia. The key to solvent-free salts of the alkali metal anions is stabilization of the
cation by incorporation into a suitable crown or cryptand complex. By using such complexes it
also appears possible to produce "electride" salts in which the charge of the complexed cation
is balanced by a trapped electron. The chemical, electrical, and optical properties of salts of the
alkali metal anions and "electrides" could provide useful applications.
1. Introduction
Alkali metals are exceptionally good electron donors,
tending to form monopositive ions by donation of one electron to an acceptor atom or molecule. The resultant cation
has a stable inert gas electron configuration with a filled outer p-shell. Because of this, the history of compounds of the
alkali metals has been that of the cations and it was long
thought that, except in metals and alloys, only the + 1 oxidation state could exist. Yet the existence of alkali metal anions
in the gas phase has been recognized for 30 years['.21.These
anions were proposed as major species in metal-ammonia solutions in 19651' 41 and in metal-amine solutions in 1969"'.
The isolation and characterization of a crystalline N a salt["-'' in 1974 firmly established that the - 1 oxidation state
of alkali metals could exist in condensed phases. This has
opened a new field of research on the properties of these
highly-reducing species both in solutions and in solids. It is
the purpose of this article to describe the current status of research in this field and to speculate on possible future developments.
2. Properties of Alkali Metals in Solution
The preparation of salts of the alkali metal anions has its
origins in the study o f metal-ammonia solutions and solvated
electrons[x ' I 1 . However, the discovqy of alkali metal anions
may have actually been delayed by the fact that ammonia is
"too good" a solvent so that the dissociation reaction
lies far to the right in dilute solutions and most properties are
only mildly influenced by the cation. This does not mean
that metal-ammonia solutions are simple: quite the contrary;
ion-pair formation, electron spin-pairing and the onset of
metallic behavior occur as the solutions are concentratedlxl.
Because of these complications, alkali metal anions could be
present in concentrated metal-ammonia solutions, but there
is no specific evidence of such species in this solvent.
2.1. Solutions in Amines and Ethers
When any alkali metal except lithium is dissolved in an aliphatic amiae, poly-ether or hexamethyl phosphoric triamide
(HMPA), two prominent optical absorption bands are observed. One, in the infrared region at 1200-2000 nm, is the
relatively metal-independent band of the solvated electron.
The other band, whose position depends upon the metal, solvent and temperature, can be assigned to the alkali metal anion"'. The proof of this assignment has been discussed in detail elsewhere[".7.'2 and will be included here only as
needed for the description of specific systems or properties.
Progress in identifying the species in metal-amine and metal-ether solutions was hampered by low solubility, rapid decomposition. and the ease of extraction of sodium from borosilicate glassl'xl.The problem of low solubility was overcome
by the use of crown ether'"' and cryptand["'' complexing
agents. Extraction of sodium from glass presents a major
problem only at low concentrations and can be completely
eliminated by using a fused silica apparatus. Unfortunately,
problems arising from solution decomposition are still with
us but can be minimized by working with rigorously clean
vessels[*'' and at reduced temperatures.
2.2. Optical Spectra
Each alkali metal anion is characterized by a broad, intense, structureless absorption band which depends upon
metal, solvent and temperature and is strongly asymmetric
on the high-energy side. These features are characteristid5' of
the so-called charge-transfer-to-solvent (ctts) bands[22.231of
other anions such as the halides, hydroxide, amide. etc. The
spectra are also similar in shape and in solvent-'"' and temperature-dependence to those of the solvated electron in various solvents["] but the absorption maxima occur at higher
energies. Figure 1 shows the spectra of Na ~,K ~,Rb , Cs
and e,,, in ethylenediamine'24'. The spectrum of e,;,l> obtained by pulse-radiolysis"'1 (solid circles) compares well
with that obtained by dissolving lithium (solid line) in ethy~
[*I
Prof. Dr. J . L. Dye
Department of Chemistry. Michigan State University
Eaht Lansing. Michigan 48x24 (USA)
16 14 12
-
I
8
10
v
07
06
I
I
I
I
I
I
12
14
I
6
d [pml
10 09 08
CcmC 10.~1
I
1
I
76
18
20
-
Excited states of the anion might lie above the ionization
limit in the gas phase. yet be observable in solution or in
crystals, either because of a strong perturbation of the continuum level or because of low probability for direct excitation to the continuum level of the solvent. In fact. auto-ionization apparently occurs from excited states of the alkali metal anions and other anions in solution[2h"I suggesting that
the continuum level lies at lower energies than does the first
excited state. It is probably appropriate to consider the intense absorption band of M in solution and in crystals as
originating from an s to p transition, modified by the presence of solvent molecules or neighboring ions and including
a contribution from direct bound-continuum transitions["I.
05
2.3. Conductivity
Fig. 1 . Optical \pectra 01. Li. N d . K. Rh. and C s in cthylrnedianiine 1241 w t h
identification of absorption peaka 0 1 N a . K . R h . Cs and e-<,,,.Solid clrcles are for e ~ , , ,produced
\
by pul\e radiolysir 1251.
lenediamine. The infrared shoulders obtained for solutions
of K, Rb and Cs are attributed to the solvated electron and
the ratio of this absorbance to that of the anions depends
strongly on concentration (see Eq. 1 ). The solvent-dependence of the peak positions is illustrated by Figure 2, which
compares the peak positions of Na , K - and e,,,,, in various solvents. The solvent-dependence of the absorption
maximum of e,,,, correlates in turn with that of I I2'I.
The conductivities of sodium and cesium solutions in ethylenediamine""' and methylamine""' have been
measured. The optical spectra show that sodium solutions in
these two solvents consist primarily of Na ' and N a - while
dilute cesium solutions contain mainly Cs and e,,,,,. The
differences in conductivity of these solutions are striking-as
shown in Figure 3. The rapid drop in conductivity with in+
I
3h
~
154
,
I
150
,
I
I
-6
OEACD ME
O,
I
y
I:"-//
1 30 - o HMPA o
I
4kDdA
~
.0
7
u
,1 111,2L6-
I
UDME U D E A
THF
DEY'
HMPA
DIPE
Fig. 3. Product of equivalent conductance ( A ) and solvent viscosity (TI) for solutions of Na in methylamine I
3 [30], and ethylenediamine A [29], and of Cs in
methylamine 0 [30]. ethylenediamine x [29]. and ammonia 0 [83].
creasing concentration for solutions containing Cs and
e,,,,, in the amines resembles the behavior in ammonia and
can be readily accounted for in terms of ion-pairing between
Cs+ and e - . The high limiting conductance of e,,,,,, in both
ethylenediamine and methylamine is also similar to the behavior in ammonia especially when the appropriate viscosity
corrections are made as in Figure 3. The conductivities and
the similarity in optical spectra thus provide strong evidence
that metal solutions in methylamine and in ethylenediamine
which d o not contain alkali metal anions are similar to metal-ammonia solutions. By contrast, the conductivities of sodium solutions in methylamine and ethylenediamine show
that the mobility of Na-- is "normal" and that the extent of
+
Maialun, Golden and Ottolenghi[51 showed that the
wavenumber of maximum absorbance of the alkali metal anions varies inversely with the estimated anion radius, in accord with the predictions of ctts theory. The observed temperature-dependence of the peak position is also predicted
We conclude that the absorption spectra
by this
of M in solution are similar to those of other anions.
An examination of ctts theory might lead one to conclude
that the excited state bears no relation to excited states of the
isolated anion. This does not necessarily follow. however.
~
588
Anxew. C l e m . In!. Ed. Engl 18. 587-.798 (1979)
ion pairing between Na' and Na
large interionic contact distance.
is small, as expected for a
spectra obtained in three solvents[331.The calculated diamagnetic shift of 6 = - 2.6 for Na compared with Na(g)13'.3hlis
-
I
2.4. The Effect of Crown Ethers and Cryptands
One of the major barriers to experimental studies of alkali
metal solutions in amines and ethers is the generally low soiubility. Combined with the ever-present tendency towards
decomposition, this low solubility has made reproducible
measurements difficult. By using crown-ether and cryptand
complexes such as those of [18]crown-6 ( l ) ,and C222, we
H-
1-1.
C r y p t and s
C222, m = n = o = 1
[18] crown-6
I
80
were able to dramatically increase metal solubilities in a
number of solvents['2.'y~20.3il.
For a cation complexing agent
(Cry = cryptand) the equilibrium (2) shifts the equilibrium
(3) to the right, thus increasing the concentration of M - in
M
+
+ Cry 5 M + C r y
2M,,,
e M++M
~
,
I
LO
20
-s
0
I
1
-20
-10
Fig. 4. "Na-NMR spectra of Na ' C222. Na solutions ( ~ I M0) in three m vents [33]. All shifts are referred lo Na,,,. Positive shifts are paramagnetic. The
peaks of Na ' C222 at the left are at a = t49.5.T h e peaks of Na at the right dre
at a = -2.
very close to the measured shifts[331
of 6 = - 1.4, - 1.6, and
- 2.3 for Na
in methylamine, ethylamine and tetrahydrofuran respectively. This is to be compared with paramagnetic
shifts of 60.5, 72.2 and 50.4 for N a + in water, methylamine
and the cryptate complex respectively. The absence of an appreciable paramagnetic shift for N a - in solution proves that
the solvent is prevented from interacting appreciably with
the 2p electrons by the presence of the outer 2s electrons.
This, in fact, provides the strongest evidence that Na in solution is truly a n alkali metal anion rather than, say, a solvated cation with two solvated electrons in the vicinity or a
solvated cation attached to a solvated electron pair.
The NMR spectra[331
of Rb in ethylamine and in tetrahydrofuran (THF) show single lines at 6 = +26.2 and +14.4
ppm respectively relative to Rb(g) while the peak for Cs- in
T H F occurs at 6 = + 52.3 ppm relative to Cs(g). These are to
be compared with values of 6 = +213 and + 348 for aqueous
R b + and Cs' respectively. Exchange effects1"' sometimes
prevent the observation of the N MR signal of M and may
be the reason that we have not yet been able to observe K by N MR even in solutions whose optical spectra show the
existence of high concentrations of this species.
-
the saturated solution. (This also applies for crown ethers as
complexing agent.) When excess metal is present, equilibrium (4) determines the relative concentrations of M - and
M
I
60
* M,,,+e,,,
(4)
e,:,,,. In the absence of excess metal, equilibrium (1) governs
the relative concentrations of M - and e,,,. The addition of
a cation-complexing agent shifts reaction (1) to the right by
reducing the concentration of M +. Equilibrium (3) requires
that the product of the activities of M - andfree cation cannot exceed the limit imposed by the solubility of metal. These
reactions permit a measure of control over the composition
tion of the solutions. The maximum value of the ratio [M -]/
[e,,,,] for a given metal is determined by the solvent and the
temperature according to reaction (4). If this ratio is large,
solutions consisting of primarily complexed M ' and M can
be prepared by maintaining an excess of metal in the preparation vessel. Alternatively, if K,[*Iis not too small and K2 is
large, the addition of equimolar amounts of complexing
agent and metal can yield solutions whose composition is
mainly complexed M + and e,,,. ,Section 3 considers the
thermodynamics of these reactions In more detail.
-
-
-
-
2.6. Other Solution Properties of Alkali Metal Anions
The rate of reaction of Na with water in ethylenediamine
is much slower than that of eSirvand the reaction order in
The ability to prepare concentrated metal-amine solutions
water is also different(2'.'71. The rate of formation of Na-.
(up to 0.2 M and higher) by using crown ethers and cryptands
from solvated electrons and Na' has been studied by pulse
opened the door to NMR studies of alkali metal a n i ~ n s [ ~ ' . ~ ~ l . r a d i o l y ~ i s ~ ~the
~ . 'reaction
~~;
is second order w.r.t. el,,,,. The
It was k n ~ w n ~ from
~ ~ . NMR
~ " studies with salts that the reldependence of the rate on the concentration of Na is comease of Na' from a Na'C222 complex is slow o n the N MR
plex and suggests['41that electron-pairing precedes the fortime scale, so that separate "Na-NMR signals were expected
mation of N a - . Photolysis of M - in the wavelength region
and found for Na Cry and for Na- . Figure 4 shows the
of the absorption maximum leads to the production of solvated electron^'^' 471, sometimes in a spin-polarized
['I KI IS the equilibrium constant for reaction ( I )
state[40.46.471 (CIDEP). Photoelectron emission has also been
2.5. Nuclear Magnetic Resonance
-
+
+
Angen G e m . Inl Ed. Engl. I d 587-598 (1979)
589
from e,;,,, and from Na in HMPA. These observations generally support the assumption that alkali metal
anions are new species and not just a combination of the solvated cation and two solvated electrons.
2.7. Ion pairs and "Monomers" in Solution
The electron spin resonance (ESR) spectrum of the solvated electron consists of a single narrow line with a g-value
near that of the free electron. When the electron interacts
strongly with an alkali cation so that appreciable electron
density exists at a metal nucleus with spin I. and when this
electron-cation pair has a long enough lifetime (generally
microseconds or longer) then a hyperfine pattern results with
21+ 1 nearly equally spaced lines[4x"I.
Pulse radiolysis[iX.52.531 and p h o t o l y ~ i s [ ~studies
'~
in the presence of alkali
cations have shown that the optical spectrum of e,:,,, is
blue-shifted by formation of the ion-pair (also referred to as
an alkali metal "monomer"). The magnitude of this blueshift increases as the polarity of the solvent decreases and
correlates well with the electron contact density at the alkali
nucleus as determined by ESR measurements[5". Changes in
the cation-electron contact density and in the spectrum with
solvent and temperature have been attributed to the formation of solvent-shared and contact ion-pairs"". Although the
sensitivity of the ESR technique permits easy detection of
the monomer species, the equilibrium concentration of monomers, M, in amines and ethers is generally small compared
with M - and e,o,v. Crown ethers or cryptands shift the
equilibrium
M+M+ t e -
(5)
for salts of the alkali metal anions (which we refer to as "alkalides") was also based on thermodynamic predictions of
their stabilities and we continue to be guided in the search
for appropriate metals, solvents and complexing agents by
approximate calculations of stability constants. Questions
have also arisen about the possible existence of anions of
other metals such as the alkaline earths. I t is a simple matter
to show thermodynamically that the existence of an ion such
as Ca in solution is thermodynamically impossible while
the gold anion, Au -, may well be thermodynamically stable.
In this section some of the pertinent thermodynamic arguments will be given. Some of the predictions can be made
with reasonable accuracy; others are far less certain, either
because key measurements have not yet been made or because of limitations of the theory of solvation free energies or
lattice energies. In either case, it is often possible to make
reasonably accurate relative calculations which at least permit comparisons among the alkali metals.
3.1. Stability of Metal Anions in Ammonia as Solvent
It is possible to calculate accurate values of AH" and AG"
for the process
from electron a f f i n i t i e ~ [ ~properties
.~~],
of the solid and gaseous
and statistical thermodynamlcs calculations of entropies['"]. The values of AH',' and AG',' for the alkali metals and some other metallic elements at 25 " C are
given in Table lI*].
Table I . Thermodynamic estimates [a] used to Judge the stability of.anions in ammonia to dissociation into metal and solvated electrons
Metal
Estimated
anion radius
r
Li
Na
K
Rh
cs
Au
AL:
cu
Ba
PI
Te
Ph
61
TI
Sh
Sn
Electron
Affinity
A ~!
A 6::
AH!
A G<.,
(A)
59.8
52.9
48.4
46.9
45.5
222.7
125.7
118.3
- 52.1
205.3
f83.3
101.3
101.3
4X.2
101.3
120.6
2.35
2.72
3.27
3.39
3.55
2.00
2.00
I .93
3.5
3.55
2.3
2.2
2.2
2.1
2.2
2.2
[a] All values given in kJ mol
I:
96.7
53.4
39 I
36.4
30.1
140.9
156.4
217.5
225.4
357.5
I1 0
91.2
103.3
131.5
15X.6
179.0
67 7
30 6
18.2
14.4
10.9
109.0
125.4
1x5 7
202 3
320.6
- 20.x
66.0
12.3
104.6
126.3
152.2
-
65.8
- 18 I
298 K
23X K
1.x
5.6
13.5
15.2
I x.5
5. I
21 4
72.5
207.9
32X 2
-91.4
- 14 1
- 8.4
12.8
45.6
7I .4
- 1I.X
- 20.2
- 1.1
-
59.3
56.3
54.9
- 61.5
- 46.0
5.5
137.7
27 I .9
- 156.5
- X6.X
- 14.1
5x I
- 195
-
I.o
0.X
38
- x3
1.9
59 0
193 7
3 16.9
- 104.5
- 29.2
- 21 8
14
32 5
51.3
AH:= 247-535(1/r); Ad:= 334- 51 I ( I / r )evaluated by fitting data for the halide ions (see text).
to the right in the absence of excess metal. Although the concentration of monomers is low, they may play important
roles in the exchange kinetics of complexed M i and M and
in the rate of decomposition of metal solutions.
If the solvation free energy of M,, is more negative than
that of e;, by an amount equal to or greater than that given
in Table 1 then the reaction
e , , , + M,.,
-
M,,,
(7)
3. Thermodynamic Considerations
metal anions in ammonThe proposed existence Of
ia[3.41was based on thermodynamic arguments. Our search
590
['I Many o l t h e thermodynaniic ealimates described in this article have not been
previously puhliahed. Details of the calculations arc dvallahle f'rom the author on
requebl
Angew.
Chem. Inr. Ed. Engl. I H . SH7-398 (1979)
is feasible and we can expect M - to be stable to dissociation
into the (solid) metal and solvated electrons. According to
the Born e q ~ a t i o n l " both
~
the Gibbs free energy of solvation
and the enthalpy of solvation of an ion are inversely proportional to the ionic radius. Since the Born equation is based
on a continuum model and fails to account for specific solvation of the ions, it is not appropriate to use it directly. A better procedure is to use the known free energies and enthalpies of formation of CI-, Br- and 1- in ammonia[5x1together with the thermodynamics of formation of the corresponding gaseous anions to calibrate the dependence of the solvation enthalpy and Gibbs free energy upon radius. Specifically, AC$ and AH'; for
are linear functions of the reciprocal of the anion radii of
CI-, Br . and I -. These thermodynamic functions were obtained by combining the reactions
For ions which do not have the added electron in an s-orbital, the calculations predict stability for T e - , Pb , Bi .
and TI- . The first three of these elements do form anions in
ammonia'h'.h21which have also recently been isolated in crystalline salts[" "I, but in these cases the anions are polymeric
rather than monomeric. Salts which contain polymeric an0ur
ions of antimony and tin have also been i ~ o l a t e d ~. ~
'~~'~~~~
calculations indicate instability of monomeric Sb and Sn
in ammonia but the addition of metal-metal bonding would
stabilize polymeric anions.
The existence of detectable concentrations of such ions as
Ba- and Pt- in ammonia is clearly impossible from a thermodynamic viewpoint. In the former case, the negative electron affinity of barium shows that Ba- is not even stable in
the gas phase. On the other hand, while platinum has a large
electron affinity the lattice energy is so great that conversion
of Pt,,, to Pt- by the action of solvated electrons is not thermodynamically permitted.
-
3.2. Dissociation of M
Solution
XI,] +el,,
+
XI,,
(11)
The values of AG?; and AH': were obtained from the
tabulation of Latimer and Jolly[5x1and are based upon
AG: = 0 and AH'; = 0 for H in ammonia. The values of AGO
for reactions (10) and (1 1) were obtained from known atomization free energies and enthalpies["I of the halogens and
their electron affinities[s41.By assuming that the solvation
free energy and enthalpy of other anions depend upon radius
in the same way as those of C1-, Br- and I - , we can calculate AGA and AH': for any anion if we know its radius. Following the suggestion of Malalon, Golden and Ottolenghi'51
we can estimate the radius of an alkali metal anion by subtracting its ionic radius from the interatomic distance in the
metal. The radius of Au- (2.0 A) was obtained by subtracting the radius of Cs' from the cesium-gold distance in the
ionic compound CsAulSY1.The values'sxl AG= - 182 kJ
mol ' . A M , l = - l 5 9 k J m o l - I f o r N a ' i n N H 3 a t 2 5 " C a n d
AH','= - 182 kJ mol-' for eYNH,,combined with A G : 2 ~ 0 . 0
at 25 oC16cr1
for the process
+
-
to M
+
-
and Solvated Electrons in
The calculations described above were made for ammonia
solutions because we have sufficient data to calculate Gibbs
free energies and enthalpies of formation of ions in ammonia, but not in amines and ethers. Similar calculations could
be made for amine and ether solvents if sufficient solubility
and/or e.m.f. data were available. Yet ironically, we have no
XpecQ'ic evidence for the existence of alkali metal anions in
ammonia. Let us examine the predicted concentrations for
the most favorable case, that of Na -. From the data in Table
1 we can calculate that for the reaction
at - 35 "C, K I 4= 3 x lo4. However, in a dilute solution away
from metal we must also consider the equilibrium
for which K I 5 = 4x lo-' M *. This predicts that at a sodium
M (neglecting the effect of activity
concentration of
coefficients) the ratio [Na-]/[e-] should be about 0.15. If the
optical spectrum of N a - were appreciably different from
that of e in ammonia as it is in amine and ether solvents,
then the presence of this concentration of Na would be easily detectable. Certainly, upon addition of a sodium salt the
presence of N a - should be observable. However. no prominent optical band of N a - has been observed either in the
presence or absence of sodium salts. We must conclude that
either the calculation given for N a - in Table 1 is in error by
15 kJ mol-' or more, or the optical band of Na is shifted
into the infrared in ammonia as originally suggested by
Golden et al.[31.
We conclude that, while such calculations as
these cannot be used to determine concentrations, they do
demonstrate that the existence of alkali metal anions in ammonia, amines, and ethers is not prohibited on thermodynamic grounds.
~
-
gives AH':3= AGy3 = - 182 kJ mol- I for
at 25 "C based upon the convention that the enthalpy and
free energy of formation of HG,,, are zero. Addition of
reactions ( 6 ) , (8) and (13) gives reaction (7). The values of
AH'; and Ad: at 25 " C and Ad: at -35 "C for a number of
metallic elements are given in Table I. The results show that
the reaction of solvated electrons with metal to produce ammoniated metal anions with filled outer s-orbitals may be
thermodynamically permitted for all of the alkali metals and
for gold1*Iand silver. It appears that Cu would not be stable
to dissociation into Cu,,, and e,;,,.
Angen: ('hpiit Int. Ed. Engl I X . 5x7-598 llY7Y)
['I W. J. Peer and J. J. Lagowki [J. Am. Chem. Soc. 100. 6260 (1978)) have recently shown that Au can be formed in liquid ammonia.
591
as good a solvent as water, means that the gas phase complexation reaction
3.3. Stability of Metal Anions in Amine and Ether Solvents
Alkali metal solubilitiesl2'I and s p e ~ t r a l * ~ . have
~ ' ] been
measured for ethylenediamine solutions at 25 "C. The extinction coefficients of N a - and e5& in ethylenediamine are
'~~.
8.2 x lo4 and 2.0 x lo4 M ' c m - r e s p e ~ t i v e l y " ~ -These
measurements permit estimation of the equilibrium constants K , , K 3 , and K, for ethylenediamine as solvent (see Table 2).
Although (excluding Li-) N a - is most stable and K - is
least stable relative to e,;,, in the presence of excess metal,
Table 2. Estimation of equilibrium constants K , , K , and
Metal
Path lenght
icm)
Solubility
(M)
Li
Na
0.29
2 . 3 9 ~in '
1.04~10
K
Rb
1.31~10
cs
5.4x 10
'
'M.
+ M,,,, + M,,,,
M :Cry Mi,,,
2
0.09
<0.05 [bj
0.Y I
0.46 [c]
0.6s
[c] In the presence of' I
(M
(241
1691
(17)
+
3.4. Stability of Metal Anions in Crystals
Attempts to grow crystals containing alkali metal anions
were based upon their stability in solution, and the realization that stabilization of the cation by its inclusion in a cryptand cavity might prevent the spontaneous recombination of
M' and M - to form the metal. The fact that alkali cations
are strongly complexed by the appropriate cryptand, even in
K ; [aj
'1
(241
1691
1241
1241
K7
(hl I)
> 4 x 10
3x10
I.hX10 6
2.3 x 10 ''
1 . 7 ~ 1 0"
1 O X 10 '
4x10
(691
0.40
0.6K
0.9
x 10 ' M
K,
Ref.
'
t2x10
1 . 4 ~ 1 0"
2.7 x 10 '
2.7x 10 '
43x10 '
4.3 x i n
7x10
-
c7
220
41
3.4
50
20
5
added K ' .
position into its components Cry, M, and M,. By adding all
of the AH" and AG" values for the reactions listed in Table 3
one obtains AHY7 and AG?,. While the absolute values
of AH:', and AG:7 given in Table 4 may be in error by as
much as t 5 0 kJ because of errors of measurement and the
various approximations made, their relative values should be
much more reliable. Since we can prepare the crystalline salt
N a + C 2 2 2 N a - , even in the presence of an excess of sodium
and cryptand, it appears to be thermodynamically stable.
The calculated value of AG?, for this salt is +32 kJ mol I
which is within the estimated absolute error of the
value A G ~ , s O required for thermodynamic stability. It appears safe to assume that any proposed salt with a value
of AGO, equal to or more negative than that of the reference
salt Na C222 Na is likely to be thermodynamically stable.
Of course, these calculations do not preclude decomposition
via destruction of the cryptand moiety, a process which we
have found to be very common.
A critical quantity in these calculations is the interionic
distance in the crystalline salt. The sum of the calculated raand rNa is larger than the measured distance in
dii rN1
the salt Na C222 Na by 14%.This could be caused by lack
of sphericity of the cryptated cation, an error in the estimate
of its size or that of N a - , and/or compression of these relatively "soft" ions in the crystal. The cation-anion distances in
Na+C2221-[761and K'C222Iare within a few percent
of those calculated on the basis of the ionic radius of I - and
the radii of Na C222 and K C222 calculated from the ligand thi~knessl'~'.Therefore, the most likely source of error
in the estimate of the interionic distance in NaiC222Na- is
an error in the estimated radius of N a - or neglect of the
compression of this ion and of the cryptated cation in the
crystal. In order to compensate for this reduction in interionic distance, all of the calculated anion radii were reduced
by 14% in computing the lattice energy and entropy in step
(h) of Table 3.
+
iCr,
~
+
+
592
+
in which M, and M, may be the same or different. If
AG';, is negative then the salt M: Cry M; is stable to decom-
A hl
the concentration of M in the saturated solution increases
regularly from N a - to C S C . The dissociation of M to
M' + 2es;,, is least likely for N a - and most likely for K - .
This apparently anomalous position of potassium solutions
seems to reflect extra stabilization of the solvated electrons in
the presence of K +,perhaps oia ion pairing. Such an effect
has actually been observed both by the addition of a potassium salt to a solution of sodium in ethylenediaminel"] and
by pulse radiolysis
The appearance of the spectral
peak of the solvated electrons occurs in this case but not
when Rb or Cs are added.
Solubility and spectral data in other solvents are too sparse
to permit evaluation of equilibrium constants. However, it is
clear that the values of K 2 increase while K , and K 3 both decrease as the solvent polarity decreases. Thus, all metals except lithium when dissolved in ethylamine show primarily
the optical band of M - with little contribution from e , , , .
Based upon the effect of C222 on the optical spectrumi701we
estimate that K , < iO-*O for sodium in tetrahydrofuran.
Therefore, solutions of sodium in this solvent, even in the
presence of excess cryptand consist essentially of
Na'C222Na- with no detectable solvated electron concentration. By contrast, solutions of potassium in tetrahydrofuran with cryptand contain largely K'C222K- when excess
potassium is present, but mainly K' C222e,;,, when an
equivalent of cryptand is present[701.
+
Cry,,,
A,
2.0
10
(16)
M + Cry,,,
for reactions of alkali metal anions in ethylenediamine at 25 "C from optical spectra and solubiltties.
0.2
0.84
0.36
0.6
x
+
is strongly exothermic and highly favored. There are no direct data for this reaction, but indirect methods can be used
to evaluate AH:, and AG:, for the reaction
Absorbance Data Used
01
10
I .0
0.01
1 .0
01
0.01
[a] K,= K , KS. [bj Estimated conc. of L1=3
K7
MfLi+ Cry,,,
+
Angew. Chem. In(. Ed. Engl. I X . 587-598 (1979)
Table 3. Steps in thermodynamic cycle used to evaluate AH';, and AS:',.
-
Redction Ytep
Origin of ddta or estimate
Typical
data
source
Enthalpies of complexation and equilibrium constants
Solubility of cryptand a s a function of temperature. enthalpies of solution of pure cryptand in water
At€' and AG" of atomization plus ionization potentials of M,,,. Use Sackur-Tetrode equation for entropies
of gaseous speciea.
Tabulated hydration enthalpies and free energies
Born equation plus estimate of size of M ' Cry from ligdnd thickness
Sum of a-e
AH" and AG" of atomization plus electron affinity of M,,,, and calculation of entropy as in (c) above
Kapustinskii equation far lattice energy; AS" from si7e and comparison with alkali m e t d halides [a]
Sum off. g, and h
1711
Ibl
[55. 561
[721
[57. 73. 741
154. 551
(74. 751
[a] A plot ol.AS(:for formation of alkali halide lattices from the ga7eoua ions L'S ( r M t + r x - j ' is linear, following the equation A$:= - 47.8 - 29.2 ( r M b + r x ) I . while for
tlie alkaline earth5 tlie equation is AS:= -72.2-53.6 ( r M ?I + r x ) ' ( S i n J K ' mol '.r in A).
[b] Data not available so AHA=AC,'=O.O was used in preparing Tables 4 and 5 . Recent calorimetric measurements give AH;= - 24.5 kJ mol ' 1 M. H . Ahrohum. University 01. Surrey (England). private communication. (1979).
Calculated values of A H o and AG" for steps (0,(h) and (i)
of Table 3 are given in Table 4 for a few selected salts. Calculations involving the cryptated alkaline earth cations included the necessary valence corrections but are probably
less certain than those for the alkali metals. Nevertheless,
modynamic stabilities of salts of the alkaline earth cryptates
are particularly striking. In the following sections we will describe our progress in the preparation of "alkalide" and
"electride" salts and speculate about possible future developments.
Table 4 I:stimated thermodynamic data [a] for assessment of the stability of some crystalline salts of the type M: Cry M ,
( M y =M,. M,#M,) and M + C r y e - at 25 "C. The subscripts f, h. and i refer to the corresponding reactions in Table 3.
M:CryM.
M'Crpe
Na ' C222 N a
LI ' C21 I LI
K ' C222 K
K C222 Na
Ba' C222(Na
K ' C222e
Li ' C21 I Au
Ba' C?22(e
259
245
23 1
23 1
715
23 I
)?
245
715
[a] All values in kJ mol
I.
AM,' and
AQl
AG,'
A q:
255
226
227
227
72 I
227
226
72 I
- 323
- 356
- 303
- 322
- 966
- 294
- 373
-
- 258
-291
- 238
- 257
- 866
- 229
- 307
- 782
880
+ M,.,
+
10
-
14
- 33
- 38
- 197
63
13
- 165
-
28
2
6
0
-114
- 3
-
27
61
not included (see however footnote [b] of Table 3 )
they indicate that salts such as Ba2+C222(Na-)* should be
thermodynamically stable.
These calculations also provide estimates of the thermodynamic stability of " e l e ~ t r i d e s " ~ ~T~ol . calculate
and
AGVx for the process
Cry,,,
-
M Crye,,
+
(1 8)
we assume that the electron occupies the octahedral holes in
a closest-packed lattice of cryptated cations. Therefore the
effective "radius" of the trapped electron is 0.414rM+Cry and
the lattice energy is assumed to be the same as that of a salt
in which a hard monovalent anion just fits into the octahedral hole. Compression of the packed cations or delocalization of the electrons would be expected to predict even
greater stabilities. The conservative estimates given in Table
4 suggest that "electrides" should be stable. The values
of AH:',, AG:)6, AH?8, and AG:, for a number of other possible compounds are given in Table 5 .
At the very least, the results given in Tables 4 and 5 for a
variety of crystalline salts of "alkalides" and "electrides"
suggest that this field is potentially a very rich one for the
synthesis of a variety of new compounds. The predicted therAngeu. Chem. Inr. Ed. Engl. 18. 5N7-5YH (1979)
Table 5. Selected estimated values of the enthalpy and Gibbs function (in kJ
mol ' j at 25 'C for the formation of crystalline salts of type M : C r y M , or
MCrye from the metal and the cryptand [a].
:
M Cry M;
M'Crye
A B ' [b]
Li C211 Na
Li ' CZI 1 K
Li ' C211 Cs
Li ' C211 e
Li ' C211 H
Li C211 Ag
Li C211 Cu
N a C222 Li
N a ' C222e
Na C221 Na
Na C221 e
Ba2+C222(Li.j 2
Ba' ' C222(Cs j?
Ba' ' C222(Au )?
Ba' ' C222(Ag )2
Ba' ' C222(Cu )>
C a l i C221 (Na j z
Ca' C221 (e j.
Cs'C322Cs
Rb ' C222 R b
-
+
+
+
+
AG" [b]
41
- 19
35
33
- x
74
19
-
-
26
85
19
- 35
- 5
- 33
- 197
- 136
- 197
- 188
- 134
- 177
- 154
-
2
- 28
to
29
58
41
99
51
26
17
II
120
- 51
- 123
- I I3
-
- 59
-
I15
70
30
I2
[a] Values of AM2 and Ad: not included (however. see footnote [b] i n Table 3).
orCry,,,+M,,,-M ' C r y e , ~ , .
[bl Reaction: Cry,,,+M,,,,+M,,,,-M:CryM,
593
4. Solid Compounds Containing Alkali Metal Anions
By late 1973 the evidence for alkali metal anions in solution was very strong, and thermodynamic estimates indicated
potential stability of crystalline ionic salts of the type
M 'CryM -. Rapid evaporation of methylamine or ethylamine from concentrated solutions of sodium or potassium in
these solvents in the presence of cryptand C222 left deposits
of metallic-looking gold-colored films on the glass walls.
Upon cooling a solution of Na+C222, N a - in ethylamine
( = 0.1 M) from + 5 " C to dry-ice temperatures, a gold-colored
powder precipitated from solution. Slow cooling resulted in
the growth of hexagonal crystals of the salt Nat C 2 2 2N a -.
Analysis[61and a crystal structure determinationl7]confirmed
the existence of this first solid compound containing a n alkali metal anion. The isolation and characterization of other
members of this class of compounds has been hampered by
difficulties with stabilities and with finding the proper solvents for crystallization, but recently we have made considerable progress in this area.
on the metal at 0°C. After pouring the supernatant liquid
into E for the last time it is distilled into a waste flask and
either diethyl ether or n-pentane is condensed on the salt in
order to remove any unreacted cryptand or other organicsoluble impurities. Provision is usually made on vessel G for
transfer of the washed salt into a number of tubes which can
be individually sealed-off under vacuum. Crystals of
Na+C222Na- up to zz 1 mm in diameter can be grown by
slowly cooling a solution which is formed by redissolving the
salt in ethylamine.
4.1.2. Crystal Structure of N a C222 Na
+
In its gross features, the crystal str~cturel'~
of this salt may
be viewed as hexagonal closest-packing of the large cryptated cations with the smaller N a - ions occupying the octahedral holes. The cryptand forms a cage for Na with threefold symmetry and a n anti-prismatic arrangement of the
ether oxygens. Views down the three-fold axis and side-on
are shown in Figure 6. The shape of the cryptate moiety and
relative positions of N a ' and N a - are similar to those in
+
4.1. Preparation and Characterization of Na +C222Na
0
0
0 0
4.1.1. Preparation of Powdered and Crystalline Forms
The preparation of this salt is relatively simple provided
one pays close attention to solvent purity and cleanliness of
the glassware. The type of vessel used is shown in Figure 5.
0
0
0
0
0
0
Na
C
A /
bl
a1
I
Fig. 6. Views of the structure of Na 'C222Na j?]. (a) Looking down the threefold axis; filled-in circles represent the "top" Na ions: open circles the bottom
ones. Triangles represent 0 atoms, squares represent CH-. groups The two N
atoms and Na' are co-linear at the center. Solid lines connect atoms above the
center; dotted lines below the center. (b) Perspective side-view of the structure.
~
O
D
D
W
Fig. 5. Apparatus for the preparation of the crystalline salt Na ' C222Na
A lo G see text.
[74]
N a C222 I [761. The closest Na to Na distances are 8.83
A in the packing planes perpendicular to the three-fold axis
while the closest distance between N a - ions in different
planes is 11.O
This could result in considerable anisotropy
of those properties which depend upon the distance between
anions such as the electrical conductivity.
-
+
The sodium, previously transferred under vacuum into small
diameter glass tubing is inserted into glass tube A after scoring the metal-containing tube with a glass knife. Heatshrinkable Teflon tubing B is used to connect tube A to sidearm C. Cryptand is then introduced into vessel G and the
entire apparatus is evacuated. Then the tube containing sodium is shaken into the flexible connection B between A and
C, broken under vacuum and moved to the constriction. A
vacuum seal-off is made behind the metal sample and with
the vessel re-attached to the vacuum manifold, metal is successively distilled into vessel E through two or three constrictorr). (With volatile cryptands
tions while pumping (=.
vessel G must be cooled during this operation.) Solvent sufficient to produce about 0 . 0 5 ~cryptand is condensed in G,
the cryptand is dissolved and the solution is transferred into
E and allowed to contact the metal mirror at 0 ° C for a few
hours. The solution is then poured through the frit F into G
and cooled to - 78 "C, resulting in the precipitation of finely-divided crystals of Na C222 Na To improve the yield,
the supernatant blue solution is poured from the crystals at
- 78 " C back through the frit F into E for further agitation
+
594
A.
4.1.3. Properties of N a C222 Na
+
Crystals of this salt frequently grow as thin hexagonal
plates with shiny metallic-looking faces. Their color at liquid
nitrogen temperature is bright yellow-gold, which darkens
reversibly to a yellow-bronze appearance at room temperatures. The crystals melt with decomposition at 83 " C to yield
gray particles suspended in a clear fluid. Analysis after cooling shows that most of the cryptand remains intact and apparently sodium metal is released upon melting. By contrast,
crystals left in evacuated tubes for several days at room temperature and under ambient light conditions darken irreversibly, apparently with decomposition of the cryptand. When a
sample was kept in the dark at room temperature for three
weeks, no apparent decomposition occurred, but the apparent photosensitivity has not yet been investigated further.
AnKen.. Chrm Int. Ed. EnRl I K . 5X7-598 lIY701
Packed powders of Na C222 Na have the temperaturedependent electrical conductivity expected for a semiconductor. As shown in Figure 7, logp u s l / T gives a straight line
over six decades of resistance where p is the specific resistance. The magnitude of the resistance depends upon the
packing pressure and decreases with time as the sample anneals. Considering this salt as a semiconductor gives a band
+
n
-
lot
1_
I0 5
?
.
.
..
I
.
..
.
L
P
,... .__..
..’.
’.......” ’
1L
c
t
t
./
1
5
10
I
I
15
20
5 [crn-’ 10P I
/
25
30
Fig. 8. Absorption spectra of a thin, solid solvent-free film of‘ Na ’ C222Na
(solid line) 1781 and of a film of Na ’ 18C-6 Na containing some methylamine
(dotted line) [Sl].
4.1.5. Summary of the Properties of Na C222 Na +
Fig. 7 Semi-log plot of specific resistance of powdered samples of
Na ’ C222 Na wrsu.7 the reciprocal of the temperature. Sample A was relatively
loosely-p;icked. The data shown for the tightly-packed sample B were measured
at both increasing and decreasing temperatures [SZ].
gap of 2.4 eV. This salt is diamagnetic, as indicated both by
ESR studies and by static susceptibility measurements. No
magnetic anomalies have been found between 300 and 4.2 K.
The ESR spectrum shows only an extremely weak signal
probably caused by trapped electrons.
The structure and properties of the salt NatC222Na- in
the crystalline state and in solution, confirm that N a- is a
“genuine” spherical anion with two electrons in the 3 s orbital. Were it not for the tendency of N a - to donate electrons
to any available acceptor, a variety of salts could probably be
prepared. However, stability requires that the counterion be
extremely reluctant to pick up an electron. Any bare metal
cation would probably be reduced by Na- and we will probably be forced to continue to use complexed metal cations or
perhaps such species as tetraalkylammonium cations. While
the latter will probably not pick up the electron by simple attachment, it is known that solvated electrons can react with
small tetraalkylammonium ions, breaking the C N bond.
Therefore, salts such as R4N Na may not be stable.
+
4.2. Other Salts of Alkali Metal Anions
4.1.4. The Optical
Na C222 Na
+
Spectrum
of
Thin
Films
of
-
Powdered or crystalline samples of this compound can be
dissolved in a number of amines and ethers to give deep blue
solutions. Upon cooling such solutions (except in methylamine and ammonia) the salt precipitates. When solutions in
methylamine, ethylamine or T H F are rapidly evaporated,
the walls are covered with a thin film which appears gold by
reflected light and blue by transmitted light. The absorption
spectrum of such a film is shown in Figure 8. In all cases with
this compound the three distinctive features shown in Figure
8 were observed[7x1.
The major peak at 15400 cm-‘ (650 nm)
is characteristic of N a - in solution[”] and probably arises
from a 3 s to 3 p transition, modified by crystal-field effects.
The origins of the pronounced shoulder at 18 900 cm - - I and
the small but distinct peak at 24500 cm-’ are unknown although the latter could arise from a bound-free transition
corresponding to the band-gap of 2.4 eV shown by conductivity. Figure 8 also shows that the gold-colored film formed
by using [ 1 81crown-6 in the presence of solvent vapor has the
major absorption band at 15400 cm-I but shows neither the
extra peak at 24500 c m - ‘ nor the shoulder at 18900 cm-I.
The spectrum shows clearly why crystals of Na’C222Naappear metallic by reflected light. The high absorption in the
near-infrared and sharp cutoff in the visible mimic the plasma absorption of metals. This can lead to a striking metallic
appearance even for nonconducting ~ r y s t a l s [ ~ ~ l .
The crystalline compound described in the previous section can be prepared either by crystallization from a solution
containing largely Na C222 and Na - or by rapid evaporation of solvent from such a solution. By controlling the complexing agent, metal, and solvent, one can prepare a wide variety of concentrated solutions containing M Cry (or M +
crown ether), MY and/or e,,,. The metals M, and M ymay
be the same or different. In this section we describe some (as
yet incompletely characterized) solids prepared either by
crystallization from such solutions or by rapid solvent evaporation.
+
+
4.2.1. Solids Prepared by Crystallization
The only crystalline solids which we have prepared to date
contain the cryptand C222 with various alkali metals-no
crystals have been made with other cryptands or with crown
ethers, nor have any “electrides” been crystallized. However,
there seems to be no fundamental restriction on the preparation of such compounds if one can find the appropriate solvent and conditions for crystallization.
In addition to the salt of Na-, we have prepared other
crystalline solids which appear to be K C222 K -,
KtC222Na-, Rb’C222Rband C s + C 2 2 2 C s - . The
methods of preparation are similar to those used for
Na C222 N a - (s) except that isopropylamine was used rather than ethylamine for salts of K -, R b and Cs . All of the
+
+
-
~
595
crystals have a metallic appearance which ranges in color
from light gold for Na C222 Na - to greenish-gold for
K C222 Na - to dark copper-bronze for Cs C222 Cs -. Solutions containing K - , Rb-, Cs- and e,;,, are much more
subject to irreversible decomposition than are those containing Na -. This makes it difficult to prepare crystalline salts of
these anions since all solution manipulations must be done at
low temperatures. Once the solvent has been removed, however, all of the crystalline solids prepared so far seem to be
reasonably stable and it is likely that we will be able to handle them and characterize them in much the same way as for
Na+C222Na-.
+
+
+
4.2.2. Preparation of Thin Films by Solvent Evaporation
By adjusting the solution stoichiometry and flash-evaporating the solvent (practical only with highly volatile solvents) it is relatively easy to prepare solvent-free transparent
films or powders whose spectral
indicate the
presence of M and/or e -. Such films have been prepared
from solutions of all of the alkali metals in methylamine in
the presence of both crown ethers and cryptands. Films
which show optical bands of M - appear blue by transmitted
light and vary in color from gold to dark-bronze by reflected
light, depending on the metal, while films which have only
an infrared band (presumably the absorption band of trapped electrons) appear blue by transmitted light and very
dark blue-black by reflected light.
Attempts to prepare crystals or powders by slow solvent
evaporation have been generally unsuccessful, with irreversible decomposition usually occurring instead of precipitation
of the desired compound. Solutions and solids that contain
[18]crown-6 tend to be more stable to irreversible decomposition than those that contain C222. However, in the former
solid films of N a + 18C-6Na- also require the presence of some solvent. Removal of solvent leads to the formation of solid sodium and [18]crown-6, without apparent decomposition of the crown ether. Evidently the axial position(s) in the N a + 18C-6Na- salt must be occupied by solvent in order for the salt to be stable. It is interesting that
films prepared by evaporating solutions containing
[18]crown-6 and either K, R b or Cs give spectra which do not
change when all of the solvent is removed.
and films by solvent evaporation from solutions which contained largely M + cryptand (or M + crown ether) and solvated electrons[’]. The ESR spectra of such powders consists
of a very intense narrow line at the free-electron g-value,
suggesting the presence of trapped electrons and extensive
motional narrowing of the ESR signal. Recently Harris and
LugowskiIXol observed narrow ESR lines from solids produced by freezing metal-ammonia solutions which contained
[18]crown-6. They attribute the signal to an ammoniated
electron in the vicinity of a complexed cation. In our case,
solvent-free blue solids are obtained so that the electrons can
no longer be solvated but may simply occupy the anion sites
in the crystal lattice. We refer to these solids as “electrides”
although the disposition of the electrons with respect to the
complexed cations is unknown.
We have measured the absorption spectra of “electride”
films with both cryptand C222 and [18]crown-6 as complexing agents (cf. Fig. 9)[R’1.Often the infrared band which we
-
5
10
15
3 [cm”
20
25
30
Fig. 9. Spectra of solvent-free thin “electride” films produced by rapid evaporation of methylamine from solutions of K C222e,,,
(solid line),
Cs’ 18C-6e,,, (dotted line) and K ‘tXC-6K (with e,,>,”?)(dashed line) [Xl].
~
attribute to trapped electrons is accompanied by an M - absorption. In some cases the relative intensities of the two
bands shifts in favor of M - as the solution is allowed to contact excess metal[7x1(cf. Fig. 10). In addition to the “electride” band at 7400 cm-‘ and the band of M - , intermediate
absorption peaks of unknown origin are often observed (cf.
Fig. 9).
4.2.3. Absorption Spectra of Films
Films prepared by evaporation of methylamine from solutions with stoichiometric amounts of M + , C222 and M (M C222 M --) show strong absorption bands which are
characteristic of the anion. The shoulder and extra peak seen
with films of Na C222 Na are not observed for any of the
other metals. Indeed, when films are prepared from solutions
of K C222 Na the peak of Na is a single narrow band.
The origin of this change in spectrum and that obtained with
[18]crown-6 (see Fig. 8) remains a mystery although we expect a change in crystal structure from Na’C222Na- to
K C222 N a - or to Na 18C-6 Na -. Since we are now able
to prepare crystalline samples of K+C222Na-, it may be
possible to determine the origin of these spectral differences.
Even before the isolation and characterization of
Na C222 Na -, we had prepared dark blue solid powders
+
-
+
+
+
-
-
+
+
596
5. Speculation about Properties and Future
Potential
Since we have only limited information about the structure and properties of “alkalide” salts, and even less information about “electrides”, it is difficult to predict in what ways
these substances might be useful. Certainly, as examples of a
new oxidation state of a whole family of elements, alkali metal anions are fundamentally interesting. The strongly reducing nature of these compounds may prove useful in syntheses, and the solid-state properties may provide useful materials in such areas as semiconductor behavior, photoelectron emission, photoconductivity, etc. In this section we will
speculate about the nature of these compounds and indicate
some areas of potential usefulness.
Angew C‘hem. Inr. Ed Engl. I X . 5117-59H (1979)
tion of alkali metal anions-the complexity and cost of cryptands.
If the solids which we call “electrides” prove to be stoichiometric compounds with one trapped electron for each
cation, then it might be possible to prepare a variety of electride salts such as Bat +C222(e-)2. The increase in electron
density in such compounds compared with M C222 e
might be enough to give them the properties of “expanded
metals”, similar to the metallic compound Ba(NH,),[’’. The
optical and ESR properties of uni-univalent “electrides”
such as K C222 e suggest that the electrons are largely localized, perhaps in the traps formed by the octahedral holes
between closest-packed cations. As the electron density increases, enough electron-electron overlap may occur to give
metallic behavior.
+
+
10
5
i;
20
15
Ccm-’ 10.~1
25
Fig. 10. Spectra of solvent-free thin films produced by evaporation of methylamine from solutions containing cryptand C222 and various relative concentrations of potassium [78]. Spectra for progressively higher concentrations of potassium are in the order, dashed line (“electride”), solid line (mixture), and dotted
line K (“potasside”).
5.1. Solution Properties
The use of crown-ether and cryptand complexing agents
makes it possible to prepare concentrated solutions containing either M or e,;,, in a variety of amine and ether solvents. For example, a solution of 0.1 M sodium in T H F can
be prepared by using the cryptand C222. Since this solution
contains mostly Na+C222 and N a - and very little e,,, it
might be useful for two-electron reductions. If a one-electron
reductant in this aprotic solvent were desired, a solution of
potassium with an equivalent quantity of the cryptand C222
would produce K+C222 and e,iiV with little interference
from K .
~
5.2. Possible New Compounds of “Alkalides” and
“Elec trides”
The thermodynamic estimates described in Section 3 suggest that compounds such as Ba+ +C222(Na-)2 should be
thermodynamically stable. Similarly, we predict that solid
compounds containing Li- should be stable. The present
challenge is to find appropriate synthetic routes to such substances. In solution, Li- tends to dissociate into Li’ and
e,,,. Perhaps a direct vapor-phase reaction or reaction in an
inert matrix would avoid some of the problems associated
with the use of a polar solvent to dissolve the reacting species. One of the problems which we have encountered in
working with lithium and the alkaline earths is the inability
to purify these metals by distillation in glass or quartz. As a
result, solutions tend to be unstable. Distillation of these
metals directly into the apparatus from a metal container
might eliminate some of these difficulties.
It should be possible to substitute other cations for cryptated cations provided they are non-reducible. Long chain
tetraalkylammonium ions seem attractive at present since
they are stable in solution in the presence of Na-. We are attempting to find suitable conditions for metathesis reactions
in solution such as
M’CryM-+R,N+X-
+
M+CryX,J+R,N+M-
(19)
If tetraalkylammonium “alkalides” could be prepared, they
might eliminate one of the major barriers to practical utilizaAngew Chem In1
Ed Engl I8 587-598 ( I g 7 9 )
~~
-
5.3. Possible Uses for the Solid Compounds
Whether the semiconducting properties of Na ’ C222 Na
and similar salts of the other alkali metal anions will be useful depends primarily on their stability. It now appears that
the parent compound is stable at room temperature in uacuo
in the dark. Even its reactivity with air is not as great as we
first thought. For example, a compressed pellet of powdered
Na’C222 N a - remained gold-colored and shiny in air until
reaction with moisture in the air began on the surface. Further surface reaction, as observed under a microscope, occurred at water droplets with little or no general surface corrosion.
Light absorption by N a - in the solid is very efficient and
covers a broad spectral range extending throughout the visible and near infrared (cf. Fig. 8). If solids can be found in
which this absorption is followed by auto-ionization as in solution126**], and if the lifetime of the conduction electron so
produced is long enough, sodium anions might be efficient
for photoelectric energy conversion. Because the electron is
only weakly bound to such anions as Cs- and may be even
more weakly bound in “electride” salts the photoelectric
properties of these substances would be most interesting. At
least in principle, such compounds have the potential to extend the wavelength of photoelectron emission farther into
the red than is possible with currently available photoemissive surfaces.
Of course, such speculation about potential uses for “alkalides” and “electrides” is based upon very little fundamental information about their properties. Before more can be
done to utilize these compounds, a great deal of study is
needed to determine their properties. Measurements are not
easy with such reactive materials even when they can be prepared in pure form. However, we hope that the demonstration that the - 1 oxidation state of the alkali metals is stable
will stimulate others to investigate such systems and find uses
for them.
The author is grateful to the National Science Foundation
for current support of this research under Grant No. DMR7722 975 and to the U.S. Department of Energy f o r prior support
under Contract No. EY-76-S-02.0958. Thanks are due to the
many graduate students and post-doctoral research associates
who worked on this project, most of whom may be identified as
co-authors of papers from this laboratory. Previously unpub-
591
lished wofk described in this paper was done by M. DaGue, H.
Lewis, P. Smith, B. VanEck and M . Yemen.
Received. August 9. 1978 [A 282 IEf
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