Solutions in liquid ammonia and other solvents

I Actually, the first observation was probably made by Sir Humphry Davy some 55 years earlier an unpublished observation in his Notebook for November 1807 reads When 8 grains of potassium were heated in ammoniacal gas it assumed a beautiful metallic appearance and gradually became of a pure blue colour . 

The interpretation of these remarkable properties has excited considerable interest whilst there is still some uncertainty as to detail, it is now generally agreed that in dilute solution the alkali metals ionize to give a cation M and a quasi-free electron which is distributed over a cavity in the solvent of radius 300-340 pm formed by displacement of 2-3 NHy molecules. This species has a broad absorption band extending into the infrared with a maximum at 1500nm and it is the short wavelength tail of this band which gives rise to the deep-blue colour of the solutions. The cavity model also interprets the fact that dissolution occurs with considerable expansion of volume so that the solutions have densities that are appreciably lower than that of liquid ammonia itself. The variation of properties with concentration can best be explained in terms of three equilibria between five solute species M, M2, M+, M and e  

Lithium, Sodium, Potassium, Rubidium, Caesium and Francium 

The lower solubility of Li on a wt/wt basis reflects its lower atomic weight and, when compared on a molar basis, it is nearly 50% more soluble than Na (15.66 mol/kg NH3 compared to 10.93 mol/kg NH3). Note that it requires only 2.34 mol NH3 (39.8 g) to dissolve 1 mol Cs (132.9 g). 

Solutions of alkali metals in liquid ammonia are valuable as powerful and selective reducing agents. The solutions are themselves unstable with respect to amide formation  

However, under anhydrous conditions and in the absence of catalytic impurities such as transition metal ions, solutions can be stored for several days with only a few per cent decomposition. Some reductions occur without bond cleavage as in the formation of alkali metal superoxides and peroxide (p. 84). 

---

Such a marked difference in the properties of solvated electron solutions in liquid ammonia and other solvents should be attributed to the extremely high donor number of ammonia (Table 3) and this shifts the equilibrium of Eq. (8) towards the cation. 

Alkali metals dissolve in liquid ammonia and other donor solvents, sueh as aliphatic amines (NR3, in whieh R = alkyl) and OP(NMe2)3 (hexamethylphospho-ramide), to give blue solutions beheved to eontain solvated eleetrons ... 

Although coordination of a ligand about a metal ion may either suppress or enhance reactivity, the present discussion is limited to the latter. Evidence is cited for enhancement of reactivity resulting solely from coordination, including cases where the evidence is indirect and involves species that exist only in solution. The principal emphasis is placed upon species that can he isolated as pure compounds, fully characterized, and which participate in reactions not exhibited by the uncoordinated ligand. Most of these species result from deprotonation of ethylenediamine and a variety of other ligands by means of amide ion in liquid ammonia and other bases in other solvents. In addition to a review of the pertinent background literature, numerous unpublished results of current studies are included. 

The alkali metals can be dissolved in liquid ammonia, and also in other solvents such as ethers and organic amines. Solutions of the alkali metals (except Li) contain solvated M- anions as well as solvated M+ cations ... 

Just as the cation produced by dissociation of water (H30+) is the acidic species in aqueous solutions, the NH4+ ion is the acidic species in liquid ammonia. Similarly, the amide ion, NH2, is the base in liquid ammonia just as OH- is the basic species in water. Generalization to other nonaqueous solvents leads to the solvent concept of acid-base behavior. It can be stated simply as follows A substance that increases the concentration of the cation characteristic of the solvent is an acid, and a substance that increases the concentration of the anion characteristic of the solvent is a base. Consequently, NH4C1 is an acid in liquid ammonia, and NaNH2 is a base in that solvent. Neutralization becomes the reaction of the cation and anion characteristic of the particular solvent to produce unionized solvent. For example, in liquid ammonia the following is a neutralization ... 

The metals, and to a lesser extent Ca, Sr, Ba, Eu, and Yb, are soluble in liquid ammonia and certain other solvents, giving solutions that are blue when dilute. These solutions conduct electricity electrolytically and measurements of transport numbers suggest that the main current carrier, which has an extraordinarily high mobility, is the solvated electron. Solvated electrons are also formed in aqueous or other polar media by photolysis, radiolysis with ionizing radiations such as X rays, electrolysis, and probably some chemical reactions. The high reactivity of the electron and its short lifetime (in 0.75 M HC104, 6 x 10"11 s in neutral water, tm ca. 10-4 s) make detection of such low concentrations difficult. Electrons can also be trapped in ionic lattices or in frozen water or alcohol when irradiated and again blue colors are observed. In very pure liquid ammonia, the lifetime of the... 

Table 5 compares the standard potential of the electron electrode in hexamethylphosphotriamide (5 °C) with the standard potentials of alkali metals (25 °C). Data for liquid ammonia are also given. In both solvents the rubidium electrode potential serves as a reference point since it depends very little on the solvent. It is seen from the Table that in both solvents the standard equilibrium potential of the electron electrode is more positive than that of a lithium electrode and is close to the potentials of other alkali metals. In the course of experiment, cathodic production of dilute solutions (10 — 10 mol/1) of solvated electrons takes place and this makes the electron electrode equilibrium potential more positive compared to the standard value. In case of hexamethylphosphotriamide the same happens when electrons are bound in strong non-paramagnetic associates by the cations of all alkali metals except lithium (see Sect. 4). This enables one to assume that under the conditions of the experiments the electron-electrode equilibrium potential in liquid ammonia and hexamethylphosphotriamide is more positive than the equilibrium potential of all alkali metals. This makes thermodynamically possible primary cathodic generation of solvated electrons in solutions of all alkali metal salts in the two solvents. 

Electrochemical dissolution of electrons and electron thermoemission may be regarded as two parallel and independent processes. It is seen from Fig. 12 that electrochemical dissolution ig preferred from the thermodynamic viewpoint, because electrons are transferred to a lower level. The mechanism of the process is, however, dictated not only by the thermodynamic factor, but mostly by the activation energy for one or the other pathway of reaction. Electrochemical dissolution demands reorganization of the solvent, while thermoemission does not. From Table 1 and Fig. 11 it follows that electrochemical dissolution is observed in liquid ammonia and hexamethylphosphotriamide solutions (transfer coefficient a = 0.5-0.75). 

Of the non-substituted alkali amides, lithium and sodium amide are commercially available. In many cases, however, the bases are prepared from the alkali metals in liquid ammonia, which is often also used as the solvent for the subsequent reaction [6, 7]. A small amount of ferric nitrate or chloride is necessary for the conversion of the metals into the amides. The actual catalyst is zero-valent iron, which is formed as a black or grey colloidal solution. In contrast to the other alkali amides, potassium amide is soluble in liquid ammonia and therefore it is kinetically more active. Deprotonations of compounds having pK values near that of ammonia are more complete with potassium amide than with the insoluble sodium or lithium amide [26]. If desired, lithium or sodium amide can be isolated in a dry ... 

The formation of the above anions ("enolate type) depend on equilibria between the carbon compounds, the base, and the solvent. To ensure a substantial concentration of the anionic synthons in solution the pA" of both the conjugated acid of the base and of the solvent must be higher than the pAT -value of the carbon compound. Alkali hydroxides in water (p/T, 16), alkoxides in the corresponding alcohols (pAT, 20), sodium amide in liquid ammonia (pATj 35), dimsyl sodium in dimethyl sulfoxide (pAT, = 35), sodium hydride, lithium amides, or lithium alkyls in ether or hydrocarbon solvents (pAT, > 40) are common combinations used in synthesis. Sometimes the bases (e.g. methoxides, amides, lithium alkyls) react as nucleophiles, in other words they do not abstract a proton, but their anion undergoes addition and substitution reactions with the carbon compound. If such is the case, sterically hindered bases are employed. A few examples are given below (H.O. House, 1972 I. Kuwajima, 1976). 

The chemical resistance of PTFE is exceptional. There are no solvents and it is attacked at room temperature only by molten alkali metals and in some cases by fluorine. Treatment with a solution of sodium metal in liquid ammonia will sufficiently alter the surface of a PTFE sample to enable it to be cemented to other materials using epoxide resin adhesives. 

Solvated electrons were first produced in liquid ammonia when Weyl (1864) dissolved sodium and potassium in it the solution has an intense blue color. Cady (1897) found the solution conducts electricity, attributed by Kraus (1908) to an electron in a solvent atmosphere. Other workers discovered solvated electrons in such polar liquids as methylamine, alcohols, and ethers (Moissan, 1889 Scott et al, 1936). Finally, Freed and Sugarman (1943) showed that in a dilute metal—ammonia solution, the magnetic susceptibility corresponds to one unpaired spin per dissolved metal atom. 

Hz. Ebsworth and Sheldrick studied the dependence of the chemical shift and H—P coupling constant of phosphine on concentration, temperature and solvent. Two phases are formed in fairly concentrated solutions of phosphine in liquid ammonia below -30 °C, one of these is phosphine-rich and the other ammonia-rich. In the phosphine-rich phase, the coupling constant, /h-p, increases from 185.2 to 186.6 Hz on cooling from -32 to -79 °C, while in the ammonia-rich phase between the same temperatures it increases from 191.1 to 195.1 Hz. 

It was also observed, in 1973, that the fast reduction of Cu ions by solvated electrons in liquid ammonia did not yield the metal and that, instead, molecular hydrogen was evolved [11]. These results were explained by assigning to the quasi-atomic state of the nascent metal, specific thermodynamical properties distinct from those of the bulk metal, which is stable under the same conditions. This concept implied that, as soon as formed, atoms and small clusters of a metal, even a noble metal, may exhibit much stronger reducing properties than the bulk metal, and may be spontaneously corroded by the solvent with simultaneous hydrogen evolution. It also implied that for a given metal the thermodynamics depended on the particle nuclearity (number of atoms reduced per particle), and it therefore provided a rationalized interpretation of other previous data [7,9,10]. Furthermore, experiments on the photoionization of silver atoms in solution demonstrated that their ionization potential was much lower than that of the bulk metal [12]. Moreover, it was shown that the redox potential of isolated silver atoms in water must... 

Although the entire discussion of electrochemistry thus far has been in terms of aqueous solutions, the same principles apply equaly well to nonaqueous solvents. As a result of differences in solvation energies, electrode potentials may vary considerably from those found in aqueous solution. In addition the oxidation and reduction potentials characteristic of the solvent vary with the chemical behavior of the solvent. as a result of these two effects, it is often possible to carry out reactions in a nonaqueous solvent that would be impossible in water. For example, both sodium and beryllium are too reactive to be electroplated from aqueous solution, but beryllium can be electroplated from liquid ammonia and sodium from solutions in pyridine. 0 Unfortunately, the thermodynamic data necessary to construct complete tables of standard potential values are lacking for most solvents other than water. Jolly 1 has compiled such a table for liquid ammonia. The hydrogen electrode is used as the reference point to establish the scale as in water ... 

The symbol for the electron in tables of values of E° in liquid ammonia is thus equivalent to NH3(1) + /iH2(g /= 1) NH4+(liq NH3 a = 1). As for aqueous solutions, several secondary reference electrodes have proved more convenient for the actual measurement of E° in liquid ammonia, e.g. silver/silver chloride. This procedure has been applied to other inorganic solvents and numerous organic solvents, and tables of values are readily available.32... 

---