James L. Dye Michigan State University East Lansing, 48824
Alkali Metal Anions An unusual oxidation state
...there can no longer be any doubt that the -I
oxidation state o f the alkali metals exists under a variety o f conditions.
Accustomed as we are to the 1+ oxidation state of the alkali metals, it seems incredible that cesium, the most electropositive element in nature, can form the ceside anion with two electrons in the outer s orbital, and that a crystalline salt of the sodium anion has been isolated and characterized. Yet there can no longer he any doubt that the 1- oxidation state of the alkali metals exists under a variety of conditions. The evidence for distinct species in metal-amine solutions with the stoichiometry M- has been accumulatingfor over five vears. The isolation of a crvstalline salt of Na- and determination of its crystal structure (1,2) show that a solvent is not required for the existence of such anions. The nature of the 21Na nmr spectrum (3) of solutions which contain Na- (and similar results (4) . . for Rb- and Cs-) Droves that this anion is a "genuine" sphrrirallg symmetric aniun, consisrenr with a structure in which thr 2p electr~msare well-shielded by the ~s. although our "chemiral inspin-paired 3; e l t ~ c t n ~ tFinally, tuition"and uptringingat the periodic table may make alkali anions seem improhable, thermudgnamic arguments shou that their iormation is entirely reasunahlr r.51. I.rt us consirit,r ihe evidence uuhich requires that we add the 1- uxidation state ro the ramiliar 1 +oxidation state o i t h e alkali metals. Background
Solutions of alkali metals in liquid ammonia have been studied for over 110 vears (6-12). C. A. Kraus discussed the importance uf electronic conductivity in these solutions as rarlv as 1807. His classic studies (13 -21i over a time s m n of neaily 30 years yielded much of the basic data upon which todav's models of the solvated electron are based. he idea of an electron existing in solution as a solvated species (anion) has fascinated theoreticians for vears and a number of models fur rhr structurr ofthis speciri have Iwen rand cuntinue to he) rmvosed (221. Although the detailsdiffer somewhat, most reeeni models considerthe electron to he trapped in a potential well formed by orientation of polar solvent molecules in the vicinity as originally suggested by Ogg (23) and developed extensively by Jortner and coworkers (22, 24-27). Mutual repulsion of solvent lone pairs and of the added electron for these "closed shell" electrons (28) can lead to the formation of a cavity within which much of the solvated electron density resides. Solvated electrons are extremely reactive towards most reducible compounds. Even in solvents such as ammonia, amines. and ~olv-etherswhich form "stable" metal solutions. the stabilityis limited by the kinetics of solvent reduction since the solutions are thermodynamically unstable. Until the development of pulse-radiolysis techniques, this reactivity olaced severe restrictions on the solvents which could he used
overlap between the study of metal solutions and radiation chemistry (references (11) and (12)). 332 / Journal of Chemical Education
to study solvated electrons. When Hart and Boag, in 1962,first observed the spectrum of the hydrated electron (291, produced by a pulse of high-energy electrons, the door was opened for the study of solvated electrons in dozens of solvents (30, 31). Ionizine radiation (and even hiph-energy photons) can eject electrons from solvent moleculis. The ejected electrons become thermalized and solvated in just a few picoseconds (32-34). Althoueh thev usuallv react with the solvent or with bther radiation-produced speiies within a few microseconds, this period of time is long enough to permit the study of many properties of solvated electrons, such as optical (30,31) and esr (35) spectra and conductivities (36). Because the con-
...pulse radiolysis techniques have also been used to verify the existence o f species o f stoichiometry Mand M - in certain solvents. centration of solvated electrons produced by radiolysis or photolysis is low and the concentration of other solute species can be controlled, radiation chemistry provides the most reliable identification of the spectrum of the isolated solvated electron. As we shall see, pulse-radiolysis techniques have also been used to verify the existence of species of stoichiometry M and M- in certain solvents. Metal-ammonia solutions are more complex than the foregoing discussion might suggest (37). Solvated electrons interact with each other and with cations to alter the solution properties as the concentration is increased. Ion-pair formation (expected for anv electrolvte in ammonia) aDDears to account for the effect ojcation-eiectron interactions.'~owever, the nature of the electron-electron interactions remains a mystery. Electron spin-pairing becomes substantial a t concentrations greater than 0.01 M a s verified by both static (38, 39) and esr (40) susceptibility studies. This spin-pairing has been variously attributed to the formation of such species as: the dielectron, ez2-, in which two electrons occupy a single cavity (23, 41); an ion triple, e-, M+, e- (42); and an alkali metal anion. M- (43). The optical, magnetic, volumetric, electrochemical, and thermodynamic properties of metal-ammonia solutions show only gradual changes with concentration and give no specific evidence for the existence of alkali metal anions. This lends support to a "weak interaction model" in which the cations and solvated electrons retain their basic identities hut form weakly interacting pairs and clusters as the concentration is increased (38). Although this picture is reasonable, the issue is by no means settled a t this time. In s h a r ~contrast to the behavior of metal-ammonia solutions, metal solutions in amines and ethers show distinct metal-dependent optical absorption hands (44-52) and esr spectra (53-56). The latter require the presence of distinct,
relatively long-lived species of stoichiometry M ("monomers"). Hyperfine splitting by the metal nucleus detected by electron spin resonance (esr) techniques, can he used to measure the "percent atomic character" of the monomer. The percent atomic character varies strongly with solvent and temperature and ranges from just a few percent to nearly 5096. Monomers comprise only a small fraction of the total dissolved metal in contrast t o the species M- which is responsible for the metal-dependent optical absorption hand. Evidence for Stoichiornetry M-
I t was early presumed by Symons (57, 58) and later by others (59) that the species responsible for the metal-dependent optical bands (V-hands) were diamagnetic since a t most only weak esr spectra could be observed in systems which showed onlv. V-hands. This contrasts with the strong esr singlet of the solvated electron in solvents which show only a metal-independent infrared optical absorption hand (irband). T h e "monomer" species, M, which give characteristic esr spectra (53-56), are present a t concentrations too low to give the intense optical V-hands. At this writing, all of the available esr evidence points to the diamagnetic nature of the species responsihle for the V-bands although this has not been directlv confirmed hv static susceptibilitv measurements. confusion abounded for years ahout the assignment (and even the number) of optical absorption hands in metal solutions. There seemed t d h e two metal-independent hands, one in their from solvated electronsand another a t about 700 nm. These hands appeared in addition to hands a t ahaut 850,9W, and 1020 nm for solutions of K, Rh and Cs, respectively. Various models for the species responsible for these ahsorptions were proposed and experiments in one lahoratory were often contradicted by similar experiments in another. The confusion was abruptly cleared up in 1968 when Hurley, Tuttle. and Golden (60) showed that the 700-nm hand resulted from contamination by sodium from the borosilicate glass used to prepare the solutions. Apparently other alkali cations, especially Li+ and K+, can rapidly exchange with Na' in the elass to eive rather laree concentrations (un to a t least G)of sodyum in a short ierieriod of time. M& df the earlier misassignments and disagreements can, in retrospect, he explained by the presence of the species Na- when, in fact, no sodium should have been present. Metal solutions in am&es, poly-ethers, and hexamethyl phosphoric triamide (HMPA) have optical spectra which can he characterized by two types of absorption hands. One of these (the ir-hand) is metal independent and can be assigned to the solvated electron. pulse-iadiolysis studies of thesolvated electron in ethylenediamine (EDA) (61) and in other solvents (62) confirm this assignment. The other band (the V-band) has a peak position which depends upon the metal used. Typical spectra are shown in Figure 1.All of the metals
I
I
I
I
4
6
8 10 12 WAVENUMBER
Figure 1 . Spectra at 25'C of e-,.l. drofuran (5i).
I
I
I 14
I
I
I
16
18
20
excent lithium show the metal-dependent hand in a t least some solvents. The ratio of the intensity of this V-band to that of the ir-band depends upon both solvent and metal. The less polar the solvent (for this purpose "polarity" decreases in the progression NHs, HMPA, MeNH2, EDA, EtNH2, PrNH2, polyethers) the lower is the solubility and the larger the ratio of the V-band intensity to that of the ir-band. The same progression of solubility and optical properties is found for a given solvent as one moves from Cs to Na. No metal-dependent hand has been positively identified in lithium solutions. A maior breakthrough in the identification of the species presentin metal s o l u & m was made in 1969 by Mitalon, Goldenand Ottolenghi (63) who compared the solvent, temperature, and metal dependence of the V-hand with the predictions of charge-transfer-to-solvent (ctts) theory. A comparison with the ctts hand of I- strongly suggested that the V-hand was a ctts hand of the alkali anion. This assignment was also in accord with the diamagnetic nature of the species, since M- would he expected to exist in a spin-paired ground state. -
[Welproduced Na- by the reaction o f a known amount o f a sodium salt with an excess o f solvated electrons. . . Characteristic of solvated electron absorption hands are oscillator strengths (64) between 0.7 and 0.9. (The oscillator strength is a measure of the intensity of the absorption band-a sinale weaklv bound electron would have an oscillator strength of [o.) since the species M- has two electrons which are relatively weakly hound, we expect the oscillator strength to be about twice that of e-.,,I,. DeBacker and Dye (65) produced Na- by the reaction of a known amount of a sodium salt with an excess of solvated electrons (produced by dissolving cesium) in EDA and thus determined the extinction coefficient and oscillator strength of Na-. The oscillator strength so obtained, 1.9 f 0.2, gave further evidence for the stoichiometry Na-. When solutions of sodium bromide or iodide in EDA were subjected to an electron pulse from a linear accelerator (66) the solvated electrons which were produced reacted with Na+ to form Na-. The rate of reaction was strictly second-order in ec,,,l, and the hand intensities were completely in accord with the assienment Na-. The same t ". w e of reaction has since " been confirmed by studies with methylamine, ethylamine, tetrahydrofuran (THF), and diethyl ether (62). Similar conclusions may be drawn from flash photolysis studies (67-72) in which M is subject to photoionization and in which the disappearance of M-, appearance of e - (and M), and recombination of M and e- to form M can be followed. Tuttle (73) has studied the Faraday effect with metal solutions and concludes that the V-hand arises from a diamagnetic species. The conductivities of solutions of sodium in ethylenediamine (74) and in methylamine (75) have been studied. In both solvents only the optical band of Na- is observedso that the major contributors to the conductivity should be Na+ and Na-. Solutions of simple salts in these solvents show a pronounced drop in equivalent conductance with increasing concentration as expected for solutes which are suhject to ion-pairing. However, the decrease in conductivity of sodium solutions with concentration indicates much less ion-pairing for Na+.Nac than for normal salts. This would only he expected for ions with a large value for the closest-approach distance. If we use the equation of Fuoss (76) for the ion-pair association constant
I~m-~xlo-~)
(1).Cs-(2). KK(3). and NaK(4) inTetrahy-
the maximum value of K A , 1.8 X 103, obtained from the conVolume 54, Number 6,June 1977 / 333
The lattice enthalpy o f the hypothetical ionic crystal Na + Na-(s) may be estimated by assuming that its structure would be the same as Na +I-(s) but with a different interionic distance.
.
-
I 'N'
ION TRIPLE
', ,H
1.
EXPANDED ORBITAL
H.
e%-
I I
'N' i
ION PAIR WITH DIELEC~RON M*C
Figure 2. Three "catim-based" models (4) far a species of stoichiometry M-. Note that ail three models permit interaction of the solvent with the outer p electmns of the cation. Ammonia is used to represent any amine or ether solvent mlsulle
durtivity of sodium in methylamine, requires a clusest-apnroarh distanre greater than G.0I[. This indicates that Nais a large ion. Finally, a puzzling solubility phenomenon is readily explained on the basis of M-. Cafasso and Sundheim (47) found that while sodium is insoluble in 1,2 dimethoxyethane, a sodium-potassium alloy will dissolve to give equimolar Na and K. This could be caused by the formation of K+, Na- in a solvent in which pure sodium will not dissolve. The results described in this section provide convincing evidence that species of stoichiometry M- exist in metal solutions. Although the simplest description of this species is that it is a solvated alkali anion with two electrons in the outer s orbital, a t least three other models, each of which has the stoichiometry M-,can be envisioned. These are illustrated in Figure 2. The interaction of the electron-pair with the cation might he strong enough in these solvents to give metal-dependent optical absorption bands and the species would he diamaenetic. A characteristic which is common to all of the speciesshown in Figure 2 is the presence of a soluated cation. This is imnortant to our later discussion of the nmr sDectrum of M-. Thermodynamics of Anion Formation
If anions of the alkali metals can be formed, isn't it possible tomake anions of nearly any element? The answer is no! It is the combination of the low ionization potential, relatively weak interatomic forces in the solid, and positiue electron affinity of the alkali metal atom which permits formation of both M+ and M- according to
assisted by the solvation energy of both M+ and M-. The role of the positive electron affinity may be illustrated by the approximate calculation of the energy of formation of the hypothetical solid salt, Na+.Na- from sodium metal (2, 5). In the thermodynamic cycle shown below, if the lattice enthalpy, AH3, of the hypothetical salt could he calculated, we would be able to evaluate the energy of formation of Na+. Na-.
The value of AH1 is 2 X A H s u b = 51.6 kcal while AHz is the ionization potential of Na(118.5 kcal) minus the electron affinity of Na (12.6 kcal following the convention that positive electron affinity refers to energy release upon electron attachment). The lattice enthalpy of the hypothetical ionic crystal Na+.Na-(s) may be estimated by assuming that its structure would he the same as that of Na+I-(s) but with a different interionic distance. The use of sodium iodide is suggested by the fact that Na- should be a t least as large as I-. Of course the Na+Na- distance in this hypothetical solid is unknown hut in order to make an estimate we set it equal to the interatomic distance in sodium metal. This implies a decrease in density because of the change in crystal structure from body-centered cubic in N a b ) to interpenetrating facecentered cubic in Na+.Na-. Other alternatives would he to retain ihe hcr structure and rhediitances in Y a w or to keep the densitv constant but change to the fcc structure. The first of these is-unrealistic in view of the expected radius ratio for Na+ and Na- and both would lead to even lareer lattice energies than that obtained by the first method. ~hkreforewe chose (5) the more conservative route to calculate the enerev . . of formation of Na+.Na-. The lattice energy of Na+Na- is computed by scaling that of Na+I- by the ratio of the interatomic distances since the lattice energy varies nearly as llr. The value of AH3 computed in this way is -143.8 kcal mole-'. This leads to AH, = +13.7 kcal mole-'. Similar calculations for the other alkali metals are given in Table 1.
The estimates described above provide a strategy for the preparation of a solid salt o f Na - from sodium metal. The small value of AH shows that stabilization of M+ could easily lead to a salt of M- for the alkali elements. Of course anything which increases the M+-M- distance will decrease the lattice energy and tend to counteract the effect of cation stabilization. The estimates described above provide a strategy for the preparation of a solid salt of Na- from sodium metal. The appropriate overall reaction is
in which C is a comolexine apent which interacts stronelv with the sodium cation and s., J. Pl?jr Chrm.. 79. 27%5 Fllr references to thenretical
.
(19751. (23) Ogg, Jr., R A , J Amer. Chem.Soc..68,155 119461;J Chem. P h w . 14,114,295 (19161: Phys. Re", 69.243.668 11946).
124) (251 (261 (271 (281 (29) (301
J0rtner.J.. J. Chem Phya., 30.839(19591. Cope1and.D. A.,Keslner,N. R.,endJortner, J.. J . Chem Phys., 53,1189 (1970). Keatner.N.R..ReL (9J.p. I. Kestner,N.R.,and Jortner, J., J Phyr. Chem., 77,1040(19731. 0'Roilly.D. E.. J . Chem. Phys., 41.3736 11964). Hart. E. J..andBoag, J . W.. J. A m r r Chem Soc., 84,409011962). Mslheson. M. S., and Dorfmsn, L. M.. "Pulse Radiolysia: M.I.T. Press, Cambridge. Mass 1960 . .. ,. ....
(31) Hsrt. E. J.. and Anbar. M.,"The Hydrated Electrnn;~Wiley.lntpraeience, New Yurt. 19711. ~~
~~
1321 Bronnk1ll.M. J.. Wolff, R. K.,snd Hunt, J. W.. J . Chrm. Phys.. 53,4201 (1970). 1331 Renteepia,P.M., Jones,R. P.,and Jortner, J.. J. Chom. Phys.. S9,766(1973). (34) Chase, W. J.,andHunt, J . W., J . Phys. Chem., 79.2835 (19751. (35) Auery,E.C..Remko, J . R.,andSmaller,B., J C h r m Phys., 49,951 (1968). 1361 Schmidt,K. H.,andBuck, W. L.,Scienee, lSI,7011966). 1371 See far example: Dyc,J. L., Ref. (81, p. 1. (381 Huster, E.. Ann. Phys.. 33.477 (1938). (39) Freed,S.,andSugarman, N.. J. C h a m P h y s . 11,354 (19431. (401 Huuhim,Jr.,C. A.,sndPasfur,R.C.,R~u. Mod. Phys.. 25,285(1953): J. Chsm. P h w 21.7959 (19531. 111) Kaplan, J..and Kittpl,C.,J. Cham. P h y r , 21.1429 (1953). (421 Co1d.M.. Jolly, W. L.,and Pitz8r.K.S.. J Amer Chem. Soc, 84,2264 (19621. (431 Golden, S.,Guttman,C., andTuttle, Jr.,T. R., J. Alner C h m . Soe, 87.185 119651; J. Chem. P h m . 44,3791 11966l. (44) GIbson,C.E.,and Phipps,T. E., J . Amrr. Chem. Soe., 48.312(1926). (461 ~ i s d e r ~. . . ~ ~ d ~ w ~ . .dc agn . iJ. ~cham., ~ , 33.411 ~ . (19551. (46) D0wn.J. L., Lewi3.J.. Mmre,B.,and Wilkin~on,G.,J.Chem. Soc., 3767 (1959). (471 Cafsssu. F..andSundheim.B. R., J . Chem Phys., 31.809(1959). 1481 Daintun, F. S., Wiles. D. M.,and Wrieht,A.N.,J. Chem. Soc., 4283(1960). 1491 Dewald, R. R., and Dyo,J. L . , J Phyx Chsm.. 68,128 (1964. 1601 Hurley, L..Tuttle,Jr..T. R.,andCc,lden.S.. Ref. (R1.p. 503. (511 l u k , M.T..Tohan,F. J.,sndDye, J. L.. J. P h w . C h m . , 76,2975(19721. (62) Fur additional ieferenees fuupticalnpectrain smines and ethen,see Dye, J. L., ref(9),
" 7" -...
(5:lJ Vm, K. D.,endDye, J. L., J. Chem. Phys., 38,2033 119631. 154) Bar Eli. K.,andTutBe. J r . T R.. J. Chem. Phyr.. 40,2508 (1964). (551 Catterall. R..and S m o n s , M. C. R.. J. C h m Sm.. 6656 11965). 1561 Fur other rrhrcnees t o ear studiesof metal~olutionss8eNicely, V. A , and Dye, J. L., J. Chrm. Phys., 53.119 (19701. (671 Fuwle~.G. W. A . McGregur. W. R.. and Symsns, M. C. R., J. Chem Sr)r., 3329 (19571. (58) Symons. M. C. R., J. Chem. Phya.. 13.99 (1959): Quan.Reu, a0.1628 19959). csa ~ d f u " . L. R., ~ ~ ~ bJ. D., ~ H~ ~ ~E. d M., t ~. ~ J. L.. ~J chpm. , phyx. 44,3969
M-.seeFletcher, J. W.,andSeddon, W.A.. J Phi., Chem.. 19,9055 (19751. I631 Matalsn, S.. Gulden. %,and Otvdenehi. M.. J. P h v ?hem.. 7'1. :I098 i19691. 164) Barrow. G.M.,"Molceulsr Spectroscopy: MeGraw~Hill,New York. 1962. p. 81. (65) DeBaekor, M.G.,andDye, J: L., J . P h u s Chrm.. 7S.1092 (19711. (66) Dye, J. L., DeBacker, M. G., E m ,J. A , and Dorfman, L. M., J Phys. Chom, 76,839
,.-,-,.
,>0""5
167) Gi1linz.L. J..Kloostorboer. J . G..Rottschnick. R.P. H..andvanV~xmt.d. D. W..Chem phi;. ti., R . ~ s ~ , ~ R z ~ I w I ~ . l68l Csbor, G.,and Bsr-Eli, K. H.. J. Phys. Cham., 75.286 (19711. 169) Fisher, M., Rsmme. G., Claosron. S., and Szware. M., C h s m Phyr. Lett.. 9. 30'3 ,,"7.>
,,>,>,.
(701 Gaathun,A.,andOttolcnghl, M.,IsioZ J . C h ~ m . 8, . 165 (19701. (711 Huppert. D..and Bar-Eli. K. H.. J. Phyr. Chrm., 74.1286 11970). 1721 G1arum.S. H..snd Marshall. J. H.. J. Chem Phyx.. 52.ShSS 11970). (731 Tuttla,Jr.,T. R.,Chem. P h y s Left.. 20,371 119731. (741 Dewald, R. R.. and Dye, J . L., J Phya Chem., 68,128 i1964). (751 Dews1d.R. b a n d Bruwal1,K. W., J . Phyr. Chem.. 74, 129 11970). (761 Fums,R. M . J . Amer Chpm Sor.. 80.5059 119681. 1771 Pedersen.C. J., J . Amer Chom. Soe. 89.7017 11967l,92,386(19701. (78) Diotrich,B.,Lehn,J. M..sndSauva~e,d.P., T ~ i r o h r d r o n ~ ~ i2885.288911969). r.. J. Amp,. Chem.Soc., 92.2916 (L9701. 179) Values fur AHIO, the ionization potentials and atomic and ionic radii were ahtsined fmm theC. R.C. HandboukufChemistryand Phyaics,Vul.S0. 1969 Valuer ihrthe elecVunaffinitips w r e t a k s n fromchen, E. C. M..and Wentwarth. W. E.. .I. CHEM. EDUC.52.486 (19751. (Wl Kspuslinskii.A. F.. Q. R w . Chem. SIC..10,283 119561. (811 Guldman,S.,and Baies,R.G., J . A m w Chem. Snc, 94.1476 (1972). (82) Chriatpnsen, J. J.,Eatough,D.J..and lzatt, R. M.,Chem. Rau., 74.351 11970. (831 Lewis. G. N., Randall. M., Pitrer. K. S.. and Brewer. L.. "Theimsdynamics: McGraw-Hill. New York, 1961. pp. 503. 528. (841 Lehn, J . M.,SLlurl. Rmdinp (Rerlin). 16. L(19781. (861 Dye, .I. L.. DeBacker. M. G.. and Nicely. V. A,, J . Am?, C h r m S s r . 92. ST26 11970). 186) Dye. J.L.. Luk, M.T.. Tohan,F. J.. Cm11en.R. B.,Papadakis, N., Cwani,.J. M.,and DeBacker. M..Ref liZi.o.659. 1871 Ceran,. J . ~ : , k dDye. J. L . J . Amer Chpm Sor.. 95.41432 119731 (881 Luk. M. T.. Ph.D Theaii. Michiean State Univerritv. 1971. (89) Moms. D.. and Weirs. R.. Act# Cryriollrr~r.,B29.:196 (1971l. (901 Cuhcn. M. H.. Jortner, J., and Thompson, J . C.. J. Cham. Phys.. 63.:1741 119751. (911 Beckmann.A..BBklen, K. D.,snd Elke. D . Z . Phys.. 27% I73 119741. (921 Malli, G..sndFraxa,S., Thvar Chim Auto. 5.275 il966J. (9:IJ Edlch, R. H.. Rnach.E.,and Popw.A. 1 . L Amer Chem Srii.. 92.4989 119701. 1941 Edieh.R. H.,and P0pov.A. I., J. Amer C h r m Svc. 93,.%20(197li. 195) Herlem. M.. and Popov. A. I.. J . Amrr. Chnm. i r 91. 14:11 119721. (%I Crrmbrrg,M. S.,Bodnel, R. L.,and P0pov.A. I., J. Phys. Chpm., 77,2449(197:31. (971 Cutmann,V.. and Wyrhera, E..lnorp. Nurl. C h r m L d . . 2.257 (19661. (981 Gutmann, V., "Cuurdinathm Chemipiry in Nonaque,,ur S,~lvenlr."Sprin#er-Ve,la& Vienna, 1968.
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