GEORGE W. WATT The University of Texas, Austin, Texas
IN
ANY comprehensive discussion of the chemistry of inorganic nitrogen compounds, it is entirely appropriate that liquid ammonia chemistry should be emphasized. From several different points of view, ammonia is certainly one of the most important compounds of nitrogen. Equally significant is the fact that for more than fifty years noteworthy contributions to our knowledge of the chemistry of inorganic nitrogen compounds have resulted from the use of liquid ammonia as a medium for the conduct of a wide variety of interesting reactions. The selection of studies for inclusion in this discussion is designed to provide scope a t the expense of detail; this is in no sense intended to be a comprehensive review of even a small segment of the field. Quite arbitrarily, the studies discussed in the paragraphs that follow are limited to work published during or since 1950, together with some unpublished observations. Comprehensive reviews of much of the earlier work are available in monographs (1, 2) and review papers; typical titles in the latter catego*. include "Ammonolysis in Liquid Ammonia" (S), "Reactions and Reagents in Liquid Ammonia" ( ) "Reactions of Inorganic Substances with Solutions of Metals in Liquid Ammonia" (5), and "The Chemistry of the Alkali Amides" (6). For the purposes of the present discussion, reactions in liquid ammonia may be classified into five categories; these together with typical examples are enumerated below.
REACTIONS IN WHICH PARTICIPATION BY THE SOLVENT IS EITHER ABSENT OR INCIDENTAL
This category includes cases that are often concerned with synthesis and that take advantage of favorable solubility relationships, the stability of reaction products at low temperatures, and/or the lesser tendency (as compared with water, for example) toward complications arising from solvolytic reactions. These and other characteristics of liquid ammonia as a reaction medium permit both the synthesis and study of a considerable variety of species that may be observed less readily or not at all in other solvents. The synthesis of pure crystalline ammonium nitrite (7) is a typical example; NaNO*
+ NHICI
-
NH,NOr
+ NsCl
the sodium chloride was separated owing to its limited solubility in a m m ~ n i a the , ~ nitrite was recovered by 1 Presented as a part of the Symposium on Recent Advances in the Chemistry of Inorganic Nitrogen Compounds before the Division of Chemical Education at the 131st Meeting of the American Chemical Society, Miami, April, 1957. A more favorable separation would result from the use of potassium nitrite since potassium chloride is less soluble in ammonia than is sodium chloride.
533
evaporation of the solvent and purified by sublimation below 80'. The work of Fischer and co-workers (8)has led to an apparently general route t o the synthesis of dicyclopentadienyls of transitional metals by the general reaction
where MT'=Mn, Fe, Co, Xi, and M1=I,i, Na, K. The corresponding compound of C++ has also been prepared and studied. Upon ~ublimat~ion of the hexammines in uacuo, complete deammoniation occurs, [MI' (NHsM (CJTdz
-
MI' (C&h
+ 6NHx
Liquid ammonia has also been shown to be a highly advantageous medium for the synthesis and study of a considerable variety of carbonylcyano complexes of transitional metals. Thus, Nast and Roos (9)have demonstrated reactions of the type,
+
--
K,INi2(CN),] 2CO 2K,[Ni(CN),(CO)I X,[Ni(CN),] 2CO Kn[Ni(CN)a(CO)sI 2KCN
+
+
Although these products are extremely hygroscopic and unstable upon exposure to the atmosphere, they are formed as stable crystalline solids in liquid ammonia a t -40". AMMONIATION
The formation and characterization of ammoniates or apparent ammoniates usually arise as byproducts of studies having other objectives. These solvates, that in many instances are strictly analogous to the more familiar salt hydrates, are commonly encountered when salts are used as starting materials in other reactions, or they may be of interest in relation to structural prob1err.s. Examples of the former include the ammoniates of VC13.xNH3(10) where x=12, 7, 6, 5, 3, and 2; the ammonia-insoluble ammoniates of BiC13.xNH3 (11) where x = > 2 (at -SO0), 2 (25"), 1 (150°), and 0 (250"); and SiFe.2NHa (IS). The lesser tendency toward ammonolysis as compared with hydrolysis is well illustrated by the 2-ammoniate of silicon(1V) fluoride. Although this solvate is entirely stable toward ammonolysis up t o 300°, it undergoes hydrolysis rapidly, e.g., 3SiF,.2NHI
+ 2Hz0
-
2(NH,)&Fs
+ Si02 + 2NH1
Similarly, the ammoniates of aluminum(I1I) iodide (29) exhibit both surprising thermal stability and resistance tosolvolysis by ammonia. At -7O0, the stable species is A118.20NH3; the 6-ammoniate is stable in the presence of liquid ammonia over the range -33.5 to 110°, but thermal decomposition at 300' yields the 5-ammoniate. An excellent example of the importance of apparent ammoniates in connection with structural problems is provided by the work of Parry and co-workers (14) JOURNAL OF CHEMICAL EDUCATION
on the so-called "diammoniate of diborane," B2H6. 2NH3. As a structural problem, this compound has long been controversial both per se and in relation to the structure of diborane. In a series of papers noteworthy for their thoroughness, Parry and his associates have brought together a wealth of both chemical evidence and physical data that point to the conclusion that this "diammoniate" is probably best formulated as [H2B(NHJ2](BH4); the evidence in support of this conclusion is too extensive for inclusion here. AMMONOLYSIS
Studies of solvolytic reactions in ammonia may have widely differing objectives. Many such investigations have been concerned with reaction kinetics, including acid and base catalyzed reactions. From the general standpoint of inorganic nitrogen compounds, however, ammonolysis is primarily of interest as a route to ammonobasic salts, amides, imides, nitrides, and related mixed types. This kind of reaction in ammonia is illustrated by the work of Goehring and co-workers (15) who have shown that the ammonolysis of sulfuryl chloride yields sulfamide and ammonium imidodisulfamide,
It is not uncommonly true that the solvolysis of a particular type of compound is strongly influenced by the nature of the more electronegative constituent. Thus, whileBClaandBBr, areammonolyzed to the corresponding amide, Keenan and McDowell (16) have shown that the imide results from the amrnonolysis of the iodide, 2BIs
+ 9NHa
-
B2(NHh
+ 6NHJ
As more or less typical of the kind of reactions that occur upon ammonolysis of halides of transitional metals, the work of Fowles and Pollard (17) on the ammonolysis of niobium pentachloride may be cited. The sequence of products observed in the presence of solvent (-63" to -23") and upon thermal decomposition of the low-temperature solvolysis product may be summarized schematically as follows, -63' + 18NHa -Nb(NHn),Clr.l2NHa
+
2NHJ21.3NHs 160' -63 to -23' 2NH41-------tNb(NHz)Cla.SNHr 200" 800' Nh(NH1)*Clr NbN Nb(NH2)C18 NHCI
NbCls
+
REACTIONS OF ALKALI AND ALKALINE EARTH METAL AMIDES
Reactions in this category are similar in some respects to those involving ammonolysis. Some ammonolytic reactions are base-catalyzed and the soluble alkali metal amides are the bases most commonly employed as catalysts. Also, the reaction products are sometimes similar, but differ in that the solvolytic reactions yield free bases whereas those involving the metal amides most commonly lead to salts of these and similar bases. I n order to maintain a distinction, the examples given here are restricted t o reactions of alkali and/or alkaline earth metal amides with species that are not susceptible to ammonolyeis at a2preciable rates. Here again one may select an example involving a potentially important practical process. I n 1945, a patent was issued with reference to a process for the production of sodium azide from nitrous oxide and sodium amide (22) ; N1O
+ 2NsNH2
-
NaN1
-
+ NaOH + NHz
more recently, this reaction as well as, 4Na
+ 3N,O + NH,
NaNs
+ 3NaOH + N2
have been the subject of extensive rate studies designed to establish conditions optimum for maximum conversion (23). Yields of 95% or greater were realized in both cases and both are essentially temperatureindependent over the range -40' to 30". The second of the above reactions proceeds a t the greater rate but is economically less advantageous. The interaction of sulfur nitride and potassium amide in ammonia has been shown t o lead to a mixture of yellow crystalline potassium salts (84), SINI
+ BKNH,
-
2S(NK)2
+ 2S(NK) + 4NHz
that are stable in an inert atmosphere and are of interest for their potential use in synthesis. Upon exposure to the atmosphere, these salts are converted to mixtures of sulfite, sulfate, and thiosulfate. The use of reactions of alkali am'des in the form* tion of amides, imides, and nitrides is well illustrated by the work of Schmitz-Dumont and Raabe (25). By varying the mole ratio of potassium amide to potassium hexanitratothorate(IV), these workers have reported a variety of imidoamides, i.e.,
+
Fowles and Pollard have also published the results of similar studies on the ammonolysis of TiC14 (fa), ZrCL and ThC14 (19), and from work as yet unpublished or still in progress, data on the ammonolysis of VC4, MoC16, SnCL, SnBr4, and Sn14 are to be anticipated
(W.
Solvolytic and other types of reaction in liquid ammonia are by no means lacking in potential practical significance. The ammonolytic reactions, for example,
+
--
+ +
C12 2NHa CINE* NHd2l NzH4 NH,Cl CINHn 2NHs
+
provide a possible route t o the production of anhydrous hydrazine (21). These and related reactions are the subject for more detailed discussion in another paper in this symposium. VOLUME 34, NO. 11, NOVEMBER, 1957
The thermal decomposition of these products has been studied in some detail; for example, the decomposition of the first of the above products is believed to proceed as follows,
This particular case is only illustrative of a broader program on the aetion of alkali amides on salts of transitional elements that is in progress in Professor SchmitzDumont's laboratory (26, 27,281. A somewhat different type of reaction of an alkali amide in ammonia has been studied recently in the writer's laboratory by McCarley and Dawes (99). The reaction between bisethylenediamioeplatinum(I1) iodide and potassium amide a t -33.5" has been shown to proceed as folloms,
--
+ KNH1 + 2KSH
[Pt(en)tlI, [Pt(en)sl12
+
+ +
[Pt(en-H) (en)jI XI XH3 [Pt(en-H)?]' 2KI 2NH8
+
where (en-H) denotes the abstraction of a proton from a nitrogen atom coordinately bonded to the central platinum ion. Both of these products are crystalline solids that are stable in an inert atmosphere; they are reconverted to [Pt(en)& upon reaction with one and two equivalents of HI, respectively. REACTIONS OF SOLUTIONS OF METALS
Probably no single feature of liquid ammonia chemistry has attracted more widespread interest than the fact that this liquid is a solvent for certain of the more active metals, notably the alkali and alkaline earth metals. Although these solutions have been the subject of a wide variety of both chemical and physicochemical investigations, interest in recent years has centered about the use of these solutions as reducing agents in the production of unusual oxidation states of the elements. For example, Hieber and Rart,enstein (SO) have shown that the reduction of potassium hexacyanocobaltate(111) with potassium in ammonia yields potassium tetracyanocohaltate(O), K8[Co(CN).I
+ 3Rt + 3
-
e
+
K,[CO(CX)~] 2KCX
This over-all reaction was subsequently sholvn to involve the intermediation of a cyano complex of the I + oxidation state of cobalt @ I ) , i.e.,
+
--
+
+
Kz[Co(CN)sI 2Kf 2eK3[Ca(CN)41 2KCN Ks[CO(CN)II K + eKA[CO(CN)~I
+
+
By similar studies, Christensen, Kleinberg, and Davidson (32) have shown that the reduction of potassium hexacyanomanganate(II1) with potassium is best interpreted as leading to a product involving both the 1 and zero oxidation state of manganese,
+
2K,[Mn(CN),l
+ 5IZ+ +
-
+
2NH, I Z J M ~ ( C N ).K.[Mn(CN)8] ~~ .ZNH*
56-
This interpretation is consistent with the analytical composition, reducing action, and magnetic properties of the product in question. Also by seduction with potassium in ammonia, ammines of platinum (55) and iridium (54) in the zero oxidation state have been produced by the reactions,
+ 2Kt + 2
[Pt(NH&IBr8 lIr(NH&BrlBrl
e
--
+ 3Kt + 3e-
[Pt(NH,),Io [Ir(NH,),I0
+ 2KBr + 3KBr
These products are nnusual in that their stability depends only upon a bonds between nitrogen and metal. Finally, there are two additional items of recent information on reactions in liquid ammonia that merit special consideration. The first of these is particularly novel in that it indicates that some reactions of solutions
of metals in ammonia may depend upon the nature of the metal. For many years the accumulation of evidence indicated that the reactions exhibited by these solutions are reactions of relatively free solvated electrons and are independent of the partirular alkali or alkaline earth metal employed. However, Keenan and McDowell (56) have shown that the ieactious of boron trifluoride l-ammoniate with lithium, sodium, and potassium in ammonia lead to distinctly different horon-nitrogen compounds, i.e., BF..NH.
+ Li
--
+ Na +K
(NH&BNHBNH (NH&BNHB(F)NH2 BFINH.
In addition t,o the major products shown above, these reartions yield alkali fluoride and hydrogen, and the M/BF1 ratios are all different, i.e., 3 : 1 for Li, 5:2 for Ka, and 1 : 1 for K. Differencesof this kind are without precedent in liquid ammonia chemistry and it will be of interest to learn whether similar differences are exhibited in still other instances. The other case that is equally unusual involves a related but somewhat different reducing system. All the reactions of solutions of petals that are discussed above involve (only) the interaction of a reducible species and a solntion of a metal in ammonia. There is a considerable body of information on reduction reactions in ammonia that differ to the extent that they include also the presence of a source of active hydrogen such as ammonium ion, an alcohol, water, etc. Evers and co-workers (36) have published the first evidence that indicates that, under competitive conditions, a reaction of ammonium ion takes precedence over a reaction of a solution of a metal in ammonia. Thus, although tetrasodium diphosphide is not reduced directly by sodium in ammonia, reduction does occur following an initial reaction with ammonium ion. The over-all result is represented by the equation,
The reaction with the metal solution, however, does not occur until exactly four molar equivalents of ammonium bromide are added. This is interpreted as indicating that phosphorus-phosphorus bond rupture is preceded by the reactions,
The biphosphine thus formed is then believed to he reduced by the sodium solution as follows, H,PPH1
+2 e
-
2(PH.)-
A somewhat different interpretation of these reactions has heen suggested by Royen and Zschaage (37). I t is hoped that this necessarily brief account of types of reactions in liquid ammonia serves to illustrate both the versatility of this reaction medium and the manner in which reactions in liquid ammonia continue to provide new information in the general area of inorganic nitrogen compounds. LITERATURE CITED (1) FRANKLIN, E. C., "The Nitrogen System of Compounds," Reinhold Publishing Corporation, New Yark, 1935.
(2) AUURIETH,L. F., A N D J. KLEINRERG, "Non-aqueous S01vcnts," John Wiley & Sons, Inc., New York, 1953. JOURNAL OF CHEMICAL EDUCATION
FERNELIUS,W. C.,
AND
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3 (1940).
KRAUS,C. A,, Chem. Revs., 26,95 (1940). WATT,G. W., Chem. Revs., 46, 289 (1950). LEYINE,R., A N D W. C. FERNELIUS,Chem. Revs., 54, 449 119.541. (7) LARBOUILLET-LINEMANN, L., Cmnpl. rend., 238,
- -..
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(8) FISCHER,E. O., A N D R. JTRA, Z. Nalurfo~seh.,8b, 217, 327 E. O., AND W. HAFNER, (1953); 9b, 618 (1954). FISCHER, ibid., 8b, 1327 (1953). (9) N A ~R., , A N D H. ROOR,Z . anorg. u. allgem. C h m . , 272, 242 (1953). (10) JANTSCH, G., K. BERGMAN, A N D H. RUPP,Z. anorg.u.allgem. Chem., 262, 223 (1950). (11) GALLAIS, F., A N D J . FANILIADES, Bd1. SOC. ehim., 1954, 4fiQ
(12) SISLER,H. H., A N D n . U. MILLER,J . Am. Chem. Soc., 77, 4998 (1955). (13) WATT, G. W., A N D J. H. BRAUN,J . Am. Chem. Soe., 78, 5494 (1956). (14) PARRY,R. W., ET AL., J. Am. C h m . Sac., in press. (15) GOEHRING, M., J . HEINKE,H. MALE,AND G. ROOS,Z. anorg. u. allgem. Chem., 273, 200 (1953). (16) KEENAN, C. W., A N D W. J. MCDOWELL, J . Am. Chem. Soe., 78,2069 (1956). J . Chem. Soe., (17) FOWLES,G. W. A., A N D F. H. POLLARD, 1952,4938. (18) FOWLES,G. W. A,, A N D F. H. POLLARD. J. Chem. Soc., 1953,2588. (19) G. W. A.. A N D F. H. POLLARD, J . Chem. Soc., 1953, . . FOWLER. 4128. (20) FOWLEE, G. W. A,, private commrrnication
VOLUME 34, NO. 11, NOVEMBER, 1957
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