Preparation of Compounds of Alkall Metal Anions
3065
Strategies for the Preparation of Compounds of Alkali Metal Anions J. L. Dye,’ C. W. Andrews, and S. E. Mathews Department of Chemlstry, Mlchigan State Unlverslty, East Lanslng, Mlchlgan 48824 (Received July 28, lQ75) Publication costs asslsted by the U.S. Energy Research and Development Admlnlstratlon
By using alkali cation complexing agents of the crown and cryptand classes (C) and appropriate solvents, solutions can be prepared which contain essentially only M+C and M- in equal concentrations. From estimates of the metal solubility in the absence of the complexing agent and of the cation complexation constant, the solubility in the presence of the complexing agent can be predicted. If, in the metal solution with C absent, the esolv-/M- ratio is appreciable, then when C is used and the solution is allowed to contact excess metal, this ratio will remain nearly the same. Therefore, in addition to MC+ and M-, the solution may contain relatively high concentrations of esolv-. Examples of the application of these equilibrium considerations to various metal-solvent combinations are described. Only Na+C.Na-, in which C i s 2,2,2 cryptand has been prepared as a crystalline solid. Gold-colored solid films have been obtained by rapid evaporation of solutions of Na, K, Rb, and Cs in one or more of the solvents methylamine (MA), ethylamine (EA), and tetrahydrofuran (THF) when the complexing agent 2,2,2 cryptand is present. In addition, gold-colored films are obtained with Na and K when 18-crown-6 is used. The former metal gives the film when MA is the solvent while the latter gives it with THF. In these cases, the gold film apparently yields the metal plus 18-crown-6 when the solvent vapors are removed. The criteria for thermodynamic vs. kinetic stability of salts of M- are considered.
Introduction Part of the fascination of metal-ammonia solutions stems from their complexity. We are accustomed to describing them as solutions which contain metal cations and solvated electrons. However, it has been evident since the days of Krausl that these solutions are more complex. In particular, electron spin-pairing occurs as the concentration increases. For example, a t -33’C only about 30% of the electron spins are unpaired at 0.02 A large number of models have been proposed3 for this spin-pairing phenomenon. Species of stoichiometry M- were proposed by Binge14 as early as 1953. Arnold and Patterson5 used electrochemical and magnetic data to argue that species of stoichiometry M- were to be preferred over those of stoichiometry M2. However, Golden, Guttman, and Tuttle617 first suggested the presence of “genuine” spherically symmetric alkali metal anions in metal-ammonia solutions. Their arguments were based largely on the stability of M- in the gas phase, which made its existence in solutions plausible. However, even today there is no specific evidence for alkali anions in metal-ammonia solutions although species of stoichiometry M- probably are present. The situation is very different in metal-amine and metal-ether solutions. In these solvents, metal-dependent optical absorption bands e x i ~ t ,which ~ , ~ demand that the alkali metal play more than a mere “bystander” role, The variation of the optical spectrum with metal, solvent, and temperature prompted Matalon, Golden, and OttolenghilO to suggest that this species in metal-amine and metalether solutions is a centrosymmetric anion of large radius. This assumption is reasonable and has been confirmed by alkali metal NMR studies11J2 and by the isolation of a solvent-free solid salt of Na-.13J4 The evidence for stoichiometry M- is overwhelming. It includes the diamagnetic nature of the species,15J6its conductivity,l7J8 the Faraday effect,lg oscillator strength,20 behavior upon photolysis,21-24 and the rate of its formation from M+ and
es01v-.25-27We can safely conclude that, a t least in amines and ethers, the formation of alkali metal anions for all members of the family except lithium is a common occurrence. Diamagnetic species of stoichiometry M- undoubtedly also exist in metal-ammonia solutions, but the present evidence indicates that they are probably best described as ion clusters or pairs with two solvated electrons or with the dielectron, e P . Since infrared bands attributable to “the solvated electron” also exist in metal-amine and metalether solutions, it is likely that ion clusters of stoichiometry M- as well as ion pairs of stoichiometry M are also present whenever the solvated electron concentration becomes appreciable. Such ion clusters, which in metal-amine solutions absorb in the infrared region, should not be confused with the alkali metal anions, M-. The latter species, which absorb at shorter wavelengths, are the subject of this paper. Equilibria Involving M- in Solution The species which can exist in metal solutions may be described by the following e q ~ i l i b r i in a ~which ~ ~ ~ C~ represents a cation-complexing agent of either the crown30 or ~ r y p t a n d ~class: l,~~,~~ M(s) i=i M+
+ esolv-
+ esolv- M M+ + 2esolv- s MM+ + C ~i M+C M+
E;!
(1) (2)
(3) (4)
It is convenient for some cases to combine equilibria 1 and 3 to give 2M(s) i=i M+
+ M-
(5)
with Kb = K12K3. Very little quantitative information is available for these reactions. However, by considering solubilities, conductances, and optical spectra, some qualitative trends can be esThe Journal of Physical Ch8mktry, Vol. 79, No. 26, 1975
3066
J. L. Dye, C. W.Andrews, and
S.E. Mathews
TABLE I: Qualitative Trends in Metal Solution Equilibria Effect of solventa on equilibrium constant
Reaction K
(1)M(S) d=M+ + e( 2 ) M+ + e-
(3)M+ + 2e-
(4)M++C
K
& ~b K
EM-
K4
Order of decreasing equilibrium constant
Decrease
Li, Cs, Rb, K , Na
Increase
Cs, Rb, K , Na (Li)C
Increase
Na, K , Rb, Cs, (Li)d
M+C
Unknown but expected Depends upon nature of C to increase aSolvent order: NH,, HMPA, MA EDA, EA, THF, DEE. b M refers to the monomer species which gives hyperfine splitting in the ESR spectrum. c No Li hyperfine splitting has been observed. d Li- has not been observed.
-
TABLE 11: Estimated Concentrations of Various Species in Equilibrium with Pure Sodium Metal in Two Solvents Computed concentrations Values used for Complexing Solvent agent (0.1 M ) K , K, K, MeM+ M+ C C MA None 1011 10-8 10-4 3x 10-4 10.' 0.10 10-4 0.097 3.1 x 1 0 - 3 10" 10'0 MA CZ22 8.2 x 1 0 - 3 MA 18-C-6 10" 108 10-8 0.089 2.8 x 10-3 1.1x 10.' 0.092 10" 0.050 0.016 2 x lo-" 0.066 0.34 THF CZ,, 1013 THF 1842-6 1013 109 0.008 0.003 1 x 10-l0 0.01 0.09 a As discussed in the text, these constants have not been measured but are estimated from solubilities, optical spectra, and conductivities where these data are available. tablished. The various equilibria are dependent upon the metal, solvent, and (for reaction 4) the complexing agent which is used. The general trends are summarized in Table I. Solvents which are referred to are ammonia, hexamethylphosphoric triamide (HMPA), methylamine (MA), ethylenediamine (EDA), ethylamine (EA), tetrahydrofuran (THF), and diethyl ether (DEE). Other solvents such as the longer chain primary amines and diamines, dimethoxyethane, and other linear polyethers could be added to the list. The value of K4 will generally be larger when the appropriate cryptand is used rather than a crown ether. Since equilibria 1-3 will still be applicable even when the complexing agent is present, the relative values of K1, K2, and K3 will determine the (e-)/(M-) ratio even when C is added to enhance the solubility, provided equilibrium with the metal is retained. Estimates of the solubility in the absence of C and of the complexation constant, K4, permit one to predict whether the metal solubility will be large or small when C is used. In the examples considered in this paper, a liberal amount of guesswork has been used to arrive at equilibrium constants. These examples should be considered illustrative only since they are not based upon firmly established data. Furthermore, such data are not easily obtained. Solubilities in the absence of C are very low, the saturated solutions are only slowly formed, and decomposition problems are severe. The values of Kq are so large in these nonaqueous solvents that they cannot be readily measured. To determine K3 would require quantitative measurements of absorption spectra or conductances, or combined ESR and optical spectra. Finally, the determination of the extent of reaction 2 to form the monomer species which shows hyperfine splitting by the metal nucleus in the ESR spectra would require quantitative ESR and optical spectra. We do not include the formation of loose ion pairs, M+.e-, in considering reaction 2. In the present treatment the formation of monomers is considered negligible. In spite of these difficulties, we have found estimates of The Journal of Physical Chemistry, Vol. 79, No. 26, 1975
the type illustrated here to be useful in predicting which metal-solvent-complexing agent combinations can be expected to yield relatively high concentrations of M-. In all cases, activity coefficient effects are neglected. The first illustrative example is sodium in methylamine at temperatures of -30 to -80°C. The solubility of sodium M based upon the conin this solvent is a t least 1.2 X ductance data of Dewald and Br0wa11.~~ Their conductance data also indicate that the ratio (e-)/(Na-) is probably less and K3 2 lo1'. By than 0.04. These data yield K5 2 examining the variation of K4 from one solvent to another,35we estimate K4 2 lolo for 2,2,2 cryptand ( C ~ Z and Z ) K4 2 lo8 for 18-crown-6 (18-C-6). (See the companion paper12 for the structures of these complexing agents.) Suppose we prepare solutions of Na in MA by three methods: (1)a saturated solution of Na in MA; (2) a saturated solution of Na in MA which contains 0.1 A4 total (2222 concentration; (3) same as (2) but with 18-(2-6. By using the limits given above for K3, K4, and Kj, we obtain the results shown in Table 11. Note the prediction that both 18-c-6 and C222 yield high concentrations of M-, nearly equal to the concentration of added C and that the concentration of esolv- remains relatively low. Both of these predictions have been shown to be correct in this case. The solubility of sodium in MA in the presence of either complexing agent is a t least 0.1 M and the integrated intensities of the *3Na NMR absorptions of Na+C222 and Na- are equal within experimental error,11,12showing that the relative concentration of esolv- is low. The relative concentration of esolv- can be increased by increasing the concentration of C in a solution which does not contact the metal. For the constants given in Table I1 for Na in MA with (2222 as the complexing agent, removal of the solution from the metal and the addition of enough C222 to make its total concentration 0.2 M would yield (M-) = 0.068 M, (esolv-) = 0.061 M , (M+C) = 0.13 M, and (C) = 0.071 M . Even in this favorable solvent, the addition of excess cryptand away from the metal does not complete-
Preparation of Compounds of Alkali Metal Anions
3087
TABLE HI: Solvent and Complexing Agent Combinations Which Have Been Used to Enhance Metal Solubility Solvent
Comptexing agents
Metal
MA DCH 18-C-6,a 1 8 4 - 6 , C,,, Na, K, Rb, Cs EA DCH 18-C-6,18-C-6, C,,, Na, K, Rb, Cs Dig1yme DCH 18-C-6 K THF 18-C-6 Na, K Na; K, Rb, Cs THF c,,, THF DCH 18-C-6 _K. _ Cs DEE 18-C-6 K’ Na,b K DEE C,,, DEE DCH 18-C-6 K, Cs K DimethyloxyC,,, ethane 1,a-PropaneDCH 1842-6 K diamine Diethylamine C,,, K Diisopropyl C,,, K ether Di-n-propylC,,, K amine Di-n-propyl c,,, K ether a Dicyclohexyl-18-crown-6.b Light blue solution indicates low solubility. ly dissociate M-. The second example shows the effect of solvent on the solubility equilibria. Consider a saturated solution of Na in T H F in the presence of 0.1 M total (2222 or 18-C-6. Sodium does not dissolve in THF, but potassium does, to an extent M.36-3sLet us assume that Na in the saturatof about ed solution has a concentration which is less than M. Otherwise a blue color would be easily discernible. Since the solvation of cations by T H F is not as favored as their solvation by MA, we assume that the complexation constants, K4, for both Cz22 and 18-C-6 are an order of magnitude greater in T H F than in MA. Also in accord with the qualitative trends shown in Table I, K3 is assumed to be two orders of magnitude greater in T H F than in MA. As shown in Table 11, these assumptions predict a total sodium solubility of about 0.2 M with 0.1 M C222 and about 0.02 M with 18-C-6. We find that C222 actually gives high concentrations of sodium in T H F (C > 0.1 M ) while 18-C-6 produces a dark blue solution which is, however, easily transparent in thin cells and has an estimated concentraM . On the other hand, potassium distion of to solves readily to high concentrations in THF when either complexing agent is used. The narrow NMR line of Na- in T H F and the equal integrated intensities for the Na+C222 and the Na- NMR absorptionsl1J2 indicate that K3 is probably much larger than that used to prepare Table 11. Preparation of Solutions The combinations of metal, complexing agent, and solvent which have been used to obtain metal solutions are summarized in Table 111. These results show that the solubilization of metals by C222 and by crown ethers is very general. I t may be of particular interest to chemists who are engaged in synthesis, that concentrations of sodium as high as 0.4 M can be obtained in methylamine by using the relatively inexpensive complexing agent 18-crown-6. If a solvent such as T H F is to be used instead with this complexing agent, then a switch to potassium still permits the production of high concentrations of M-. The metals will even dissolve in benzene and t o l ~ e n with e ~ ~the ~ ~aid~ of complexing agents of this type, although in these cases the aromatic radical anion is formed rather than M-.
uum
Ive
I
CryDtQnd
Figure 1. Apparatus for the preparation of solutions of M+C,M-
Solutions with the maximum (M-)/(e-) ratio are easily prepared in an apparatus such as that shown in Figure 1. Provision is made for the separate addition of metal and complexing agent. Following evacuation, the metal is distilled from the side arm to form a mirror on the walls of vessel A. Then the solvent is distilled under vacuum into B. When the complexing agent has been dissolved, the solution is poured through the coarse frit onto the metal. Agitation at about O°C produces a saturated solution which can then be poured back through the frit to remove it from the metal. Solutions of sodium in MA, EA, and T H F in the presence of C222 prepared in this way were studied by 23Na NMR,12 and in each case the area under the Na+C peak was equal to the area under the Na- peak within a few percent. This shows that the stoichiometry of the dissolved metal was essentially Na+C,Na-.
Solids Which Contain MWhen a -0.2 M solution of sodium in EA with Cz22 is cooled from about +10 to -2OoC, shiny gold-colored hexagonal crystals form s p o n t a n e ~ u s l y . ~ Their ~ , ~ ~analysis and crystal structure show that the solid has the composition Na+C-Na-. The Na+C moiety has virtually the same structure as it does in salts such as Na+C-I-. Since the Na+C. Na- salt forms spontaneously and reversibly in the solution even in the presence of pure solid sodium, it appears that this compound is thermodynamically stable with respect to Na(s) and C(so1n). When the solid is rapidly heated in a sealed tube, it melts at 83OC (compared with 68OC for the pure cryptand) to yield C222 and free sodium.14 It is also possible, however, that other decomposition products form. The existence of salts which contain alkali metal anions is not altogether unexpected. Because of the relatively low lattice energies and ionization potentials of the alkali metals, and particularly because of the positive electron affinities of the gaseous alkali atoms,6~7~41~42 the energy required to convert the metallic solid into a hypothetical ionic solid according to 2M(s)
-+
M+-M-(s)
is less than about 15 kcal mol-l. The calculation of the enthalpy of this reaction is made by means of a Born-Haber cycle of the type shown in Figure 2 for sodium. Known thermodynamic quantities are used throughout except for the lattice energy of the hypothetical salt. This was estimated by assuming that the structure would be the same as that of sodium iodide but with a different interionic distance. The M+M- distance in the hypothetical salt was set equal to the interatomic distance in the corresponding metal. This implies a decrease in density because of the change in crystal structure from body-centered cubic The Journal of Physical Chemktry, Vol. 79, No. 26, 1975
J. L. Dye, C. W.
3068
Andrews, and S.E. Mathews
TABLE IV: Estimated Enthalpy Changes for Various Reaction Steps M-M Enthalpy change, kcal mol" (see Figure 2) distance Metal in metal, A AH,, AHio AHea AH1e AHF
Li Na
K
Rb
cs
3.03 3.72 4.5 4.87 5.42
-14.3 -12.5
124.3 118.5 100.1 96.3 89.7
33.17 24.03 19.35 18.64 16.82
-11.8 -11.1 -1 0.8
-175.0 -142.7 -117.9 -109.4 -97.9
Radius of (AP9(AH",cryptated AHqo, A f J O s h a x , AfJO
