LIQUID AMMONIA AS A SOLVENT. V METALLIC SOLUTIONS This

LIQUID AMMONIA AS A SOLVENT. V METALLICSOLUTIONS JOSEPH F. CHITTUM A N D HERSCHEL HUNT Department oj Chemistry and the Purdue Research Foundation, Purdue University, Lajayette, Indiana Received December $8, 1986

This field has been ably reviewed by Johnson and Meyer (7), Johnson and Fernelius (6), Kraus (8), and Franklin (4),so that references to the literature need not be cited here. The experimental data of liquid ammonia solutions of metals is explained by a theory-namely, the dissociated metal theory-which we consider open to question. In liquid ammonia some metals, lithium, sodium, and potassium in particular, are claimed to dissociate to give a metallic cation and an electron which is solvated in dilute solutions. In more concentrated solutions even the existence of free electrons is claimed to account for the enormously low resistance of the solutions. The theory may be summarized into two equations: M=M++ewhere M represents Na, K, etc., and e-

(“I),

= e-

+ nNHI

The “solvated electron” is used to explain the anomalous conductance data, vapor pressure data, photoelectric properties, transference values, the blue color of the metallic solutions, electrolysis, and absorption spectra. In this paper we shall report a number of experiments which directly conflict with the “dissociated metal” theory, and also give our interpretation of the data in the literature. EXPERIMENTAL

Electrolysis A dilute solution of ammonium chloride was electrolyzed at -60’ to -8O’C., using small platinum electrodes. A blue color was detectable on the cathode surface. This blue color does not exist at the boiling point of liquid ammonia. Furthermore the evolution of hydrogen at the cathode shows that very little metallic ammonium is being formed. The small amount of blue color, however, must be due to the presence 581

582

JOSEPH F. CHITTUM AND HERSCHEL HUNT

of metallic ammonium. Other data lead us to believe that the ammonium metal exists in the solution as NH4(NH2-),. When a pure mercury cathode was used, with only a trace of ammonium chloride present as an electrolyte, no gas was liberated a t the cathode, but there was a large amount of gas evolved a t the anode. This we interpret to mean that ammonium metal is plated out and forms ammonium amalgam. A part of the cathode surface was a deep blue color during the plating process. When the amalgam was removed, it decomposed and hydrogen was liberated. KHe was first prepared in (l)NH3 solutions by Moissan (9) and Rich and Travers (11). The ammonium amalgam resembles the ammonium amalgam prepared in aqueous solutions. We then used sodium amalgam (unsaturated amalgam) as a cathode, and the only electrolyte present in the liquid ammonia was the small amount of sodium amide formed by reaction of the sodium in the amalgam with the ammonia. This solution was a very poor conductor, and it was necessary to use 220 volts across the system with the electrodes only a few centimeters apart. As soon as the current was applied blue streamers shot out from the cathode toward the anode. The speed of these particles is so high, a centimeter per second, that most of their energy must be supplied mechanically by surface tension effects and gas formation. Polarization causes the sodium and ammonium to plate out locally. The metal forms a colloidal particle and then moves away. Before reaching the anode the blue became fainter, and bubbles of hydrogen appeared which rose to the top of the liquid. There is no gas evolution a t the cathode surface. In our discussion we will explain how the blue color is due to a colloid, that is, free metal with adsorbed amide ions. This negative particle is pulled away from the cathode by the applied potential. Out in the solution it becomes electrically an unstable sol. The hydrogen is produced by the decomposition of the ammonium metal, two molecules of ammonium metal giving two molecules of ammonia and one of hydrogen. According to the dissociated metal theory the blue color would be due to EL solvated electron, but the decomposition of this ion could not give gaseous hydrogen. The liquid ammonia solution containing the sodium amalgam was perfectly colorless before the current was applied.

Reactions Sodium reacted very rapidly with ammonium chloride a t room temperature. At low temperature, about - 7OoC.,there is practically no reaction. The rate of the reaction was determined by the volume of hydrogen given off per unit of time. At room temperature, roughly 5 cc. of hydrogen was produced per minute. At the low temperature ten hours was required to produce 1 cc. The reaction rate is easy to follow qualitatively by the disappearance of the blue color, Ionic reactions should reach equilibrium

LIQUID AMMONIA SOLUTIONS OF METALS

583

very rapidly, but a reaction with a metal with adsorption interfering would be relatively slow. The reaction

NH4+

+ e-("&

+

= (n s")1

+H

would certainly proceed very rapidly and

is known t o go very rapidly with the production of much heat. Schlubach and Ballauf explain the deficit in the hydrogen evolution by the formation of ammonium metal. They observed that only 35 per cent of the theoretical hydrogen was given off when ammonium chloride acted upon potassium solutions in liquid ammonia at - 70°C. (12). Pure mercury was added to blue, unsaturated solutions of sodium in liquid ammonia. The mercury was shaken with these solutions for several hours. The solutions were, in sealed glass tubes; some of them were kept cold and some were allowed to react at room temperature. As soon as the formation of amalgam stopped, as judged by hardness and swelling, the amalgam was removed to dry tubes and the solvent ammonia allowed to evaporate. From these solid amalgams ammonia and hydrogen were liberated. When this latter decomposition took place the amalgams softened and decreased in volume. Tyndall cone

A blue solution of sodium in liquid ammonia in a closed tube was allowed to settle carefully so that the undissolved metal and the sodium amide precipitated fell t o the bottom. Dilute solutions prepared in this manner always showed a Tyndall cone. Dialysis

Several membranes were found which were not attacked rapidly by liquid ammonia. Sodium and amide ions would diffuse through these membranes. A commercial viscose product was very satisfactory. Metallic sodium was placed in such a bag and suspended in a glass cell between platinum electrodes. Ammonia was condensed in the cell and bag. The sodium dissolved, forming a deep blue solution in the bag and some sodium amide. The amide diffused throughout the cell, but there was no blue color outside the bag. Electrodialysis was then carried out, using 220 volts. With this large potential no blue color could be pulled outside the bag. Sufficient sodium and amide ions were dialyzed to make the solution a poor conductor. This we consider as good evidence that the metal exists as a colloid.

584


Photoelectric properties Kraus' experiment on the photoelectric properties of a lithium solution is confusing, since he reports a positive as well as a negative charge of the electrometer. We tested the photoelectric properties of sodium solutions in a closed cell a t a low temperature, so that the ammonia vapor pressure was only 1 to 5 mm. The electrodes were of platinum and the anode was about 6 mm. above the solution. No current was detectable in a sensitive galvanometer, even when a 220-volt driving force was applied. The experiment was performed under different light intensities with the same negative result. Under the same experimental conditions pure metallic sodium showed photoelectric properties. We found that unless extreme care is used to prevent it, the ammonia will condense on the wall of the cell and thereby produce an electrical leak between the solution and the anode. Such a leak may explain the data of Kraus. Densities of solutions We hoped to determine the densities of solutions of alkali metals and of salts as well as of their mixtures a t 25°C. However, we found that the dilute blue solutions of sodium became colorless with gas evolution when sodium bromide was added. This cannot be interpreted as common-ion effect. Surely it is due to a breaking down of the colloidal system. If concentrated solutions of sodium were used we obtained two layers, a dense blue solution which settled to the bottom like an oil and a less dense bronze-blue layer. These latter solutions are both stable for days. The so-called solutions of alkali metals when saturated, as a t the break in the vapor pressure curve, are not homogeneous in appearance, but are sludge-like. These data we cannot explain by the theory of Franklin and Kraus, and therefore we are offering a colloidal metal theory. The colloidal theory was first suggested by Ostwald (10). He did not propose the existence of ammonium metal in the alkali solution and therefore he tried to explain only the blue color of the solutions. THEORETICAL INTERPRETATION O F THE DATA I N THE LITERATURE

The authors think that the alkali metals are dispersed into colloidal particles when they come in contact with liquid ammonia. They are not dispersed to atomic or ionic dimensions, At the time of peptization 'the metal displaces ammonium metal from the liquid ammonia according to the reaction M + "4' e Mf + NHI (1) This makes the colloidal particle a colloid of an alkali metal and ammonium metal. The blue solutions can be obtained only in alkaline solution. It


585

is impossible to prepare ammonia so pure that it will not destroy a part of the blue color when a dilute blue solution is added to it. We do not think that the disappearance of the blue color is due to reaction with water adsorbed on the glass container. The extent to which reaction 1 takes place depends upon the activity of the two metals, the temperature, and the presence of impurities. The ion which is adsorbed in the peptizing process is the amide ion. Consequently the colloidal micelle is negatively charged. The ammonium and amide ions come from the reaction 2NHa rit NH4+

+

"2-

(2)

K , the dissociation constant, of which is reported to be 10-88 in pure NHs, but the reaction rate can be increased by removing the products as they are formed. The value of K increases very little with increase in temperature, but may increase markedly in the presence of alkali metals. The alkali ions from equation 1 are removed by the amide ions that are left from reaction 2 to the extent of the solubility of sodium amide. The reaction

M

+ n(NHd+, NHz-) + nNH4+ + C(NHz-)*

(where C represents colloidal metal) leaves NH4+ ions in the solution, that is, this reaction produces ammonium ions in solution equivalent to the colloid in solution until the electrical characteristics of the sol are such that the system is stable. The theories of metals have not been developed to the point where they would predict the appearance of a completely dispersed metal. Therefore we cannot judge the color of the solutions. On the other hand almost all colloidal solutions are colored, the color being dependent upon the particle size. There is a point of apparent saturation of the metal in liquid ammonia. This fact can be explained using the colloidal metal theory if it is assumed that a certain definite concentration of amide ions is necessary to peptize the metal. This type of fact is known in the case of the peptization of iron in aqueous solutions of sodium hydroxide (2). The process of sol formation uses up amide ions, and when they are exhausted to a critical concentration then another type of colloid, such as the copper-colored colloid, is formed. The less active the metal the fewer amide ions required, and consequently the higher its solubility. There is a very small change in the solubility of the metal with the change in temperature. There is obviously, therefore, very little heat of solution. This fact is very hard to explain using the metal dissociation theory. If this theory should be completely vindicated, the solution of a largely ionized substance with an infinitesimal heat of solution such as lithium in liquid ammonia would remain a remarkable phenomenon, since the solvation of the electron would give off a large amount of heat energy. The small change in the solubility with change in temperature is much more

586


plausible in the light of the colloidal metal theory; however, little is known about the energy of peptization of metals. The heat of reaction 1 is also an unknown quantity. There is a decrease in the vapor pressures of the solutions with increased concentration of the metal, that is fairly reproducible. Very little is known experimentally about the effect of colloids upon the vapor pressure of solvents, but theoretically the particles should serve as ions in the solution. Effectively we should have ammonium colloid in solution. If this is the case, then the vapor pressure data at least do not disprove our theory. Concentrated lithium solutions have very low vapor pressures,much lower than Raoult’s law could possibly explain. If the limiting factor i s not the concentration of the lithium ion, but the adsorption of nTHs on the metal atom as NHa and NH2-, then one does not have to postulate a large ratio of adsorbed particles per lithium atom to have a large percentage of the solvent molecules tied up. The enormous increase in activity with a slight dilution cannot possibly be due to metallic dissociation; it must be due to a decrease in the per cent of solvent adsorbed. The molecular weight and activity of sodium have been calculated from the change in the vapor pressure with concentration of the metal. The molecular weight approaches less than 20 in dilute solutions. All of the calculations in the past have been carried out either to find out what value of molecular weight is approached a t dilute solutions or to determine what the activity of the metal is, assuming Kraus’ theory. The facts might be considered as some substantiation of the “dissolved and dissociated metal” theory if more were known about the extent to which salts dissolved in liquid ammonia obey Raoult’s law. One of the authors (5) has carried out the determination of the activity coefficient of salts in liquid ammonia a t room temperatures and has discovered marked deviations from Raoult’s law. The facts about the change in the vapor pressure of the solution with change in metal concentration can be accounted for qualitatively on the basis of the formation of ammonium colloid electrolyte in solution and the marked departure from Raoult’s bw which would be expected in solutions of such electrolytes. The charge on the colloid particle per atom of sodium determines whether or not Raoult’s law could be considered as being obeyed by these solutions. B y postulating the necessary charge on the colloid (not unreasonable charges) the data could be accounted for assuming Raoult’s law. The disappearance of the blue color around the anode when a dilute metallic solution is electrolyzed we explained by the oxidation of NH, to ”,+, which immediately reverses equilibrium 2 with an increase of WHa or solvent in the anode compartment. Amide ions are plentiful, since they were adsorbed on the colloidal metal. This, of course, removes the source of the blue color, and since the ions are removed the solution immediately


587

around the anode will become a very poor conductor. Furthermore, at the same time NH4+ is removed from the anode by attraction to the cathode. This explains the absence of visible products of electrolysis a t the anode. The NH4+ will be an excellent carrier of electricity in its parent solvent (1). The cathode compartment will become richer in the blue color because

NH4+ + e (furnished by cathode) % NH4 and secondly because of transference of the existing blue particle. The ionic atmosphere of the fast moving ammonium ion will contain the colloidal negative particle C(NH2-)z. If the sodium or potassium ion is present it will also act as a positive carrier. These ions are not liberated, however, since this would require a greater potential. By our theory or the accepted theory, the free metal simply would ionize again. Reaction 1 will go to completion only at the anode. If the concentration of the solution is increased the conduction should become more and more metallic, owing to the increase in concentration of the colloidal particles. This will also shift the fraction of the current carried by the positive and negative carriers. The N&+ will carry most of the current in the dilute solutions, since the larger colloid particle is less mobile, but when the concentration reaches a certain limit, the mobility of the negative carrier will no longer be an important factor, as the conduction is by electronic transfer (3). As a higher metal concentration is reached one would expect the ionic concentration to decrease and consequently the apparent molecular conductance. The two will not necessarily counterbalance, and we should expect a minimum conductance. This minimum has been reported by all investigators in the neighborhood of 0.05 normal. Our theory would predict a higher conductance than for a similar concentration of a salt because of the nature of the conductors. The specific conductance of mercury is approximately two and six times, respectively, the specific conductance of a saturated and a 2 N solution of sodium. Preliminary experiments in this laboratory showed that the resistance offered a D.C. current by a sodium solution is decreased roughly 50 per cent by the addition of sodium chloride. This must be explained by an orientatian of the conductors into a complicated “bridge-work” which favors electronic conduction. The addition of a common ion should decrease the concentration of the free electron, and therefore greatly increase the resistance of the solution. There is a large positive temperature coefficient of the electrical conductivity that does not decrease in value with increasing temperature. This fact cannot be explained using the dissociated metal theory. Ordinary electrolytes have their conductivities pass through a maximum. Metals have negative coefficients. The increasingly large positive tem-

588


perature coefficient is just what the colloidal metal theory would predict. There would be an increasing number of collisions between metal particles, and consequently an increasing percentage of metallic conductivity with the increase in temperature. The absorption spectra of the alkali metal solutions are the same. This interesting fact can be interpreted as indicating that the colored component of all these dilute solutions is the solvated electron. The colloidal metal theory postulates that the metal colloid, especially in the dilute sol, is largely ammonium metal with a core of alkali metal. In the formation of a dilute colloidal solution the procedure is quite uniform. Since the composition of the outer part of the colloidal particles (no matter what alkali metal is dispersed) is the same, since the methods of formation of the various colloids are the same, and since all of the colloids are quite stable, it follows from the principle of Hevesy that the rotational-vibrationalelectron-transition spectra should be quite similar for the sols of the various alkali metals. The densities of the solutions of the alkali metals in liquid ammonia pass through a maximum density with increasing concentration of the metal. The change in density of these solutions is credited to the existence of electrons associated with ammonia molecules. I t seems t o the authors that this explanation is especially weak, since metals are in general very dense and the more dense the metal, the greater the possibility of having more free electrons per atom. It is well known that colloids greatly change the internal energy of liquids under most circumstances. Colloids lower the surface tension of a liquid; thereby they decrease the force conipressing a liquid, and the density would be decreased. The charged forces in the colloidal solution will increase the free space between particles, so that the density will be decreased. The density of the free ammonium metal is an unknown factor, but we predict that it is less than 0.5 g. per cubic centimeter and that it will therefore decrease the density of the solution. The relative amount of free ammonium metal will change with $he concentration of the dissolved metal as well as the closeness of packing of the adsorbed particles, so that the magnitude of the final density cannot ,be estimated. The closeness of packing should be greater for potassium than for sodium; therefore, one would predict a smaller increase in volume for potassium than for sodium solutions. This prediction agrees with the experimental data. The metallic ionic radii of sodium and potassium are of the wrong relative order to explain it, and the “solvated electron” should have the same density in both solutions. Metals such as platinum will catalyze the decomposition of NH4 through the removal of Hz. 2KH4 = 2NH3

+ Hz


589

Therefore the rate of formation of sodium amide from liquid ammonia and metallic sodium will be increased. If the metal furnishes a colloidal metal surface under the experimental conditions, it will be a better catalyst for the reaction. SUMMARY

1. Methods for producing ammonium metal in liquid ammonia have been described. Its presence in alkali metal solutions has been proven. 2. The “dissociated metal’’ theory of metallic ammonia solutions is critically discussed, and a colloidal metal theory is outlined to explain the properties of these solutions. 3. The alkali solutions are evidently mixtures, and therefore the physical data are not capable of interpretation on the basis of the well-known laws of physical chemistry. REFERENCES (1) BERNAL AND FOWLER: J. Chem. Physics 1, 515 (1933).

(2) CHITTUM AND RITCHEY: Unpublished work. (3) FARKAS, L.: Z.physik. Chem. 161A,355 (1932). (4) FRANKLIN, E. C.: The Nitrogen System of Compounds, American Chemical Society Monograph. The Reinhold Publishing Corporation, New York (1935). (5) HUXTAND LARSEN:J. Phys. Chem. 38, 801 (1934); 39,877 (1935). (6) JOHNSON AND FERNELIUS: J. Chem. Education 6,20 (1929). (7) JOHNSON AND MEYER:Chem. Rev. 8, 273 (1931). (8)KRAUS,C. A. : Properties of Electrically Conducting Systems, American Chemical Society Monograph. The Reinhold Publishing Corporation, New York (1922). (9) MOISSAN:Compt. rend. 144, 790 (1907); 193, 803 (1901). (10) OSTWALD, WOLFGANG: Kolloidchem. Beihefte 2, 437 (1910-11). (11) RICHAND TRAVERS: Proc. Chem. SOC.22, 136 (1906); Trans. Chem. SOC.89, 872 (1906). (12) SCHLUBACH AND BALLAUF: Ber. 64B, 2825 (1921).

LIQUID AMMONIA AS A SOLVENT. V METALLIC SOLUTIONS This

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