J. J. Lagowski
The Chemistry of Liquid Ammonia
The University of Texas Austin, 78712
-
.. . one might well consider whether there are principles which govern chemistry. on a broader basis than those considered for aqueous media.
..
Solvents are ubiquitous substances to most chemists. Many reactions of interest are conducted in a solvent, and very often little attention is directed toward the influence of the major suhstance in the system-the solvent-when considering the detailed chemistry of such processes. Our perceptions of chemistrv are stronelv influenced bv ohservations on aaueous systems.?hus, the &solubility of sifver chloride, the solihility of potassium sulfide and the immiscibility of (C2Hd20 in water suggest certain implications with respect to the chemistrv and svnthetic usefulness of these suhstances. These imbressionb, gained from aqueous chemistry are, however, quite incorrect for liquid ammonia in which silver chloride is soluble, K&3 is insoluble, and (C2H5)20is completely miscible. Thus, it is as easy to precipitate K+ from liquid ammonia with SZ- as it is to precipitate Ag+ from water with C1-; the Ag/ AgCl electrode is not viable in liquid ammonia. Ether cannot he used to extract liquid ammonia solutions, hut it forms a good co-solvent with ammonia in which to run reactions of covalent substances. Faced with such observations. one mieht well consider whether there are principles which govern the chemistrv of. for examole. silver on a broader basis than those considered fbr aqueois reaction media. A consideration of the properties of ammonia compared with those of water (Table 1) leads to a superficial insight into the solvent properties of this suhstance. As early as 1899, Franklin and Krause ( I ) described ammonia as water-like in attempting to rationalize its solvent characteristics. In fact, the supposed similarity of ammonia and water led Franklin to propose (2) the concept of a nitrogen system of compounds, analogous to the water system of compounds in an attempt to systematize the chemistry of liquid ammonia solutions. Ammonia. like water. is a hiehlv associated liauid as is suegested hy'the relatively hi& value of ~ r o u t h ' sconstani. X-ray scattering experiments on liquid ammonia indicate that each ammonia molecule is surrounded by an average of 11 other ammonia molecules (3).This hieh decree of association is also reflected in the structure of solih ammonia. Crystalline Table 1. Some Physical Properties of Ammonla and Water NH.
Normal Bailing Point Normal Freezing Point Trouton's Cm~tant.callmaie Specific conductivity. ohm-' cm-' Dielectric constant Viscosity. centipoise Density, glmi. Dipole moment, Debye Poiarirabiiity (X loZ'), crn3/moie 752 / Journal of Chemical Education
H.0
ammonia is structurally more complex than ice (4); each nitrogen atom is surrounded by 6 nearest-neighbor nitrogen atoms, the unit being hound together by asymmetric and non-linear hydrogen bonds. The hydrogen atoms formally bound to the central nitrogen form hydrogen bonds to three ammonia molecules. The nitrogen atom of the central ammonia molecule is bonded. throueh bonds. to three " hvdroeen " more ammonia molecules. In effect, each nitrogen atom is associated with six hvdroeen atoms. three formine normal covalent bonds and tfnee at hydrogen-bond distances. A more meaningful understanding of the chemistry of liquid ammonia solutions can he obtained by considering in detail the nature of the interactions which can occur between solvent and solute species. Solution Phenomena
-
We shall d e v e l o ~first eeneral orincioles of solution ohenomena, discussing various modes of solvent-solute interaction for both molecular and ionic species. Chemical phenomena which arise from or are enhanced by the properties of ammonia are then discussed. Solvents intervene in chemical processes by producing species from solutes that are more reactive than if the solvent were not present. If the solute is This is the seventh in a series of Resource Papers, intended primarily fur college and university teachers. The publicatian of'the series is supported in part by a grant fmm the Research Corporation. J.J. Lagowski was graduated from the University of Illinois in 1952. He CWWNIan M i degree in lti:>i t r m the I'n~wriityoi . \ l n h ~ g a nm d a Phi)
7 -
the next two years he studied with Professor H. d. Emeleus at Cambridge University, Englana, as a Marshall Scholar and a memher of Sidney Sussex College, receiving the degree PhD (Cantab.) in 1959 for his research on ~erfluomalkvlmercurv comnaunds.
with solution phenomena in "on-aqueous solvents, particularly liquid ammonia; with metallo-organic derivatives of nonbenzenoid compwnds such as borazine; and synthetic methods involving the reactions of metal vapors. He also has an active research gnmp investigating the importance of computer-based techniques an chemical education. He is Editor-designate of this Journal.
6+
6-
-
(1)
D: + X-Y D-XC + YElectrophilic attack a t the potential anion site will also encourage the formation of ions in such systems (eqn. (2)). 6+
~~~~
-
Fioure 1. In the cubic structure tor solid ammonia, each nitraqen atom is formally bound to three hydrogen atoms, and hydrogen-bondedto threeother hydrogen atoms on adjacent ammoma malecuies.
ionic, the energy for the dissolution process is supplied almost entirely from the solvation of ions; this is primarily an electrostatic nrocess. However, two processes may occur when a covalent solute is dissolved; the solvation of molecules and the formation of ions. The solvation of molecular species usually involves (a) dipolar interactions, (h) specific interactions such as the formation of hydrogen honds, and (c) the formation of covalently bonded species via coordinate covalent bond formation. In many instances it is possible for such intermediate species to undergo ionization. The formation of ions from covalent molecules is attributed to the "ionizing power" of the solvent. Attempts have been made to relate ionizing power to the dielectric constant of the solvent, but there are a sufficient number of embarrassing contradictions to make such relationships unreliable. Two examples will suffice to make the point. Walden showed that triphenylmethylchloride, which is certainly a covalent substance in the pure state, is extensively dissociated in liquid SO2. (c 12 at 2Z°C), but is too weak an electrolyte in nitrobenzene (c = 34.5 at 25%) t o permit an estimation of the concentration of free ions in that solvent. The second examnle involves nerchloric acid which -is completely dissociated in water ( 6 = 58.5 a t 25"C), hut undissociated in anhvdrous sulfuric acid (c = 101 at 25'C). The heterolytic cleavage of a covalent bond (6+)X-Y(6-) to form the corresnondina ions can be imagined to occur by one of two processks. ~ucreophilicattack b; a donor atom at the nositive site can assist in the displacement of the negative as an anion (eqn. (1)).
-
~~
~
~
Table 2.
Doniclty (D.) and Dielectric Constants ( f ) of Some Common Solvents Solvent
1.2 Dichloroethane Nitromethane Nitrobenzene Selenium oxychloride Acetonitrile Methyl acetate Acetone Water Diethyl ether Dimethyl formamide Dimethylsulloxide Pyridine Hexamethylphosphoramide
0" 0 2.7 4.4 12.2 14.1 16.5 17.0 18.0 19.2 26.6 29.6 33.1 38.8
101 35.9 34.8 46.0 36.0 6.7 20.7 81.0 4.3 36.1 45.0 12.3 30.0
6-
X-Y+A-Xf+YA(2) In essence, this view of ionization involves a consideration of the ability of a solvent to form coordinate covalent bonds. Solvents, the molecules of which are electron donors (like ammonia) will tend to stabilize cations (eqn. (1)) whereas solvents that are electron acceptors stahilize anions (eqn. (2)). There are. of course. other factors. vide infra, which encourage anion formation. The extent to which a substance is ionized by neutral donor molecules should increase with increasing stability of the cation resulting from nucleophilic attack (eqn. (1)) compared with that of the unionized solute. We should expect that the strength of the cwrdinate covalent bond formed in the species DX+ would he related to the donor ability of D, the acceptor ahility of X in the species X-Y, steric effects, and the magnitude of s~ecificsolvent solute interactions such as Hbonding. ~ i donor e strengths of solvent molecules have been defined relative to a reference accentor viz., ShCk (5.6).The enthalpy for the reaction between Sbcls and a series of donors in an inert medium (ean. (3)) was found to he proportional to the log of equilihrium'constant for this reac$on, suggesting the entropy contribution to such processes are constant. ~
~~~
~
SbCk + D r D .SbCk (3) Thus, AH for reaction (3) was taken as a measure of the donor properties of D and the donicity (D,)for a solvent was defined by eqn. (4).
Dm= - ~ o . s b c l S
(4)
The donicity of a solvent has been interpreted as a measure of its donor intensitv. .. nucleo~hilicitv,or Lewis base strenah, of and i t can be a useful guide in assesding the ionizing a solvent. The donicities of some common solvents appear in Table 2. Inspection of these data indicated that there is no necessary correlation between the donicity of a solvent and its dielectric constant. As might he expected on intuitive grounds, nitrogen-containing solvents have a relatively high donicity. Interactions between Solution Species Solubility of a solute in a solvent involves consideration of the attractive forces which exist between the species in the mixture, i.e., solute-solvent, solvent-solvent, and solute-solute. Fundamentally three factors must he considered in a discussion of the interaction between molecular substances: (1) the polarities of the molecules involved, (2) the magnitude of the dispersion forces, and (3) the ability of the molecules to interact in a specific manner, e.g., via hydrogen bonding. Dipolar Interactions. The interaction between two molecules possessing permanent dipole moments depends upon their relative orientation. The most favorable orientation, i.e., both dipoles aligned anti-parallel, is opposed by thermal motion; at a given temperature a statistically preferred orientation of the dipoles occurs, the potential energy of which is given by eqn. (5) E
=- - 2 ~ 1 ~
3l"T
(5)
where pl and flz are the magnitudes of the two interacting dipoles at a distance r from each other and the remaining svmhols have their usual meanines. In addition, polar (solvent1 molecules mn induce a dipole in another (solute) rnolecule which leads to secondary dipolar interactions. Thus, it might he vxpected that ammonia w ~ t ha dipole moment shout 80% that of water (Tahle :31 could dissolve other polar molecules as well as ionic species. Dispersion Forces. London (7) has shown that two molecules which do not react chemically will attract each other Volume 55, Number 12, December 1978 1 753
. ..ammonia is a better solvent for substances with a low internal
pressure than is water.
because they possess fluctuating dipole moments, although neither mav- nossess a nermanent dinole moment. The fluc. tuating dipoles in a given molecule induce dipoles in neighboring molecules leading to dipole-dipole interactions; the forces which arise by this process are called dispersion forces. The magnitude of the dispersion energy is given by eqn. ( 6 ) .
.
where a, ~d a2 represent the polarizahifities of the molecules in auestiou. r is the distance between the molecules. and I their ionization potentials. Dispersion forces are short range effects. A comparison of the London forces for ammonia and water appears in Tahle 3. The orientation effect, which denends upon the dipole moment of the molecule, for water is-more than twice as large as that for ammonia, hut the reverse is true when dispersion forces are considered. Thus, ammonia may not he as good a solvent for polar molecules as is water, hut it is a hetter solvent for the less nolar suhstances than water. liildehrand and Scott (A, haveshuwn that thesquare root of the internal nressure ( 1 E , \ ' I ' ?. where 1E..is the enerrv of evaporation-per mole and is the molal vkume, can ;f used as an index of soluhilitv; suhstances that have similar internal pressures will dissolve in each rjther to the grenttst extmt;at O°C the internal pressure iurammmia isahout 55% thnt crt'wntpr (Tnhle 3). Although t:xtmsivt: quantitntivedata are nut available for the soluhiltty ofcuvalenr suhstsnn:~in liquid ammonia, there are indications (Table 4) that ammonia is a better solvent for suhstances with a low internal pressure than is water. Aliphatic hydrocarbons [(AE,)IV)"~- 71 are generally found to he "insoluble" in ammonia. On the other hand, aromatic hvdrocarhons and the haloeenated alinhatic hydrkarbons with larger internal pressures are morr soluble than the a l i ~ h a t i rhvdrocarhons: rhe suhititutiarn of a nitro group for hidtogen increases the solubility of an organic radical. The introduction of highly polarizable atoms into a covalent molecule can lead to strong interactions with Table 3.
Solvent Parametersfor Water and Ammonla
p. Debye
Orientation Effect.* erg cmB
Dispersion erg cme
AEJ Vw2
a caiculated from eqn. (6). "calculated from eqn. (5).
Table 4. Internal Pressure and the OualRatlve Solublllty of Selected Covalent Substances In Llquld Ammonia Substance
(AEv/ V)'I2
Saiubility
CH&H2CHs CHdCH&CH3 %He CsHGH3 CeHsCH=CH2 m(CHs)26& CHk3 C*Hd C&@ Br(CHd2Br ClzC=CH2 CIHC=CHCI CH3N02 CeHsNO*
6.0 7.3 9.2 8.9 9.3 8.8 10.5 9.4 8.9 10.4 9.1 9.8 12.6 10.0
Insoluble Insoluble Mcderately Soluble Slightly Soluble Soluble Slightly Soluble Miscible Soluble Soluble Soluble Miscible Sduble Miscible Solubility. 24%
754 / Journal of Chemical Education
ammonia dipoles, incrvusing the sduhility of these suhstances. For examnle, iodofrmn is renorted to he "ven easilv soluble" in liquid ammonia and a white crystalline compoukd, CHIT NH?, has been isolated from this svstem. ~ i d r o g e nBonding. Specific intkractions between molecular species, in contrast to the non-specific dipolar, coulombic, or didpersi6n interactions, are a ~ s bimportant in solvation processes. Hydrogen-hond formation is one of the most important of the specific interactions. Species which contain hydrogen atoms covalently hound to a highly electrouegative atom-usuallv N or O-can form hvdroeen . " honds to other atoms carrying unshared electron pairs, i.e., X-H . . . Y. Hvdroeen hondine- is.. of course. an imnortant consideration for solvents which undergo auto-protolysis, since molecules in such svstems can hvdroeen - - bond to themselves as well as to solutemolecules carrying unshared electron pairs. Molecules which contain highly polar functional groups such as the carhonyl moiety (aldehydes, ketones, acid amides, and esters) that can interact strongly with solvent dipoles might he expected to interact strongly with ammonia (Table 5). In addition, an enhancement of solubility would he expected in the presence of functional groups which can form hydrogen honds to ammonia (X-H . . . NH3, e.g., alcohols, primary, and secondary amines) or he hydrogen-bonded by ammonia (X . . . HNH2, e.g., ethers, tertiary amines, and nvridine). Some t -~. i c aaualitative l soluhilitv results which .. illustrate these arguments are given in Tahle 3. Hnwever, with tht, inlurmation available at nresent. it is difficult t o s e ~ a r a t e the dipole-dipole interactions from hydrogen-bonding interactions in many of these systems, e.g., compounds containing the carhonyl group. Finally, molecules possessing acidic hydrogen atoms such as carhoxylic acids, phenols, or imides react with liquid ammonia to form soluble ammonium salts. However, highly oxygenated compounds such as poly-acids are less soluble in Table 5.
Qualitative Solubllltles of Some Selected Organic Compounds
Compound
Soiubility
Compound
Solubility
CH&HO CH&H2CH0 CHs(CH&HO CHZ(CH~)~CHO CH3CONH2 CH3CONHCH3
miscible soluble easily soluble soluble very easily soluble soluble
miscible miscible miscible miscible ml~cible miscible
CH3CH2CONH2 CH3CSNH2 INH&CS HC02C2Hs (CHd2CO CH3CHzC02CH3 CH3CH2COr (CHhCH,
very easily soluble
CsHsCH2C0&H3 CH30H CH3CH20H CH3(CH2),0H CsHsCH20H HOCH2CH(OH)CH20H CHSNHZ ICHM CH,(CH,),NH, CeHsNHe C5W (C4WJ [CHdCH2)&0
miscible miscible miscible miscible miscible miscible Soluble
C6HsOCH3
miscible
easily soluble solubility 37% miscible miscible miscible very easily soluble
Table 6. Solublllties (molesll) of Selected Sodium and PolassIum~Salbin Ammonia and i n Water Ammonia (-0.1') Na+ K+ CIBr1NOS-
2.20 6.21 8.80 15.00
0.0177 2.26 11.09 1.04
Water (0') Na+ 6.10 7.71 10.72 8.62
K+ 3.76 4.49 7.72 1.30
ammonia, because they form ammonium salts of multiplycharged anions which are less soluble than uni-univalent electrolytes (vide infra). Electrolytes
The solubilities of ionic substances depend not only on the dielectric constant and the polarity of the solvent but also upon the lattice energies of the solutes. Thus, in comparing the solubilities of ionic substances in ammonia with their solubilities in water, it is apparent that the differences in solvent properties are the predominating factors. Since the dielectric constant and dipole moment of ammonia are less than the corresponding values for water (Table I), the iondipole interactions might he expected to be diminished, as would the solvation energy of an ion. However, dispersion effects for ammonia are significantly more important than for water; although dispersion forces generally can be neglected in ionic solvation, these forces may become important for large ions such as I- that have relatively high polarizabilities. The maximum potential energy arising from the interaction between an ion of charge Ze and a molecule with a dipole moment fi a t a distance r apart is given by eqn. (7). E = - - Zerr -2
such as SOaZ-, P043-, C032-, 0 0 2 - , As04-, BOZ-, S02-, 02-, OH- and dicarboxylate anions are virtually insoluble in liquid ammonia; perchlorates and nitrates show the greatest solubilities of the oxygenated anions. Compounds containing cations with a relatively high charge density (e.g., Li+, Mg2+, and Ca2+) dissolve to form ammoniates as do salts containing easily polarizable anions (e.g., NaI and RbI). In general, the ammonium salt of a given anion is the most soluble salt; indeed, a saturated liquid ammonia solution of NHaN03 (Divers solutions) has a vapor pressure of less than one atmosphere a t 0%. lonic Equilibria
Solutions of ionic species in liquid ammonia are appreciably associated. because of the moderate dielectric constant of the solvent. I& pairs predominate a t relatively low analytical concentrations of ionic solutes. and higher ion agmeaates are formed with increasing concentration-(eqn. (9)j.' ~~~~~~
~
~
(7)
Thus, it is apparent that the higher dipole moment of water should lead to a stronger interaction with a given ion. The solvation energy, AE, of a gaseous ion of charge Ze and radius r in a solvent with a dielectric constant D is given by the Born equation (eqn. (8)).
-.
The hvdration enerev -.of a "easeous ion should be -greater than its ammoniation energy. Although extensive data are not available. a comoarison of the solubilities of a few ionic suhstnncri in liquid ammunia and in water at the same temperature r'l'ahle 6 ) indicntrs that water is aeneralls a better solvent f i r electrolytes than is ammonia inaccord&ce with eqn. (8). The trends in solubilities generally reflect trends in the lattice energies of a metal salt. Thus, at a given temperature the solubilities of metal halides increase in the order F- < C1< Br- < I- which is also the order of decreasing lattice energies. The majority of salts containing highly oxygenated anions
1.f
1s f
9 X
W
Conductivity data for solutions of substances that might be exoected to be comoletelv ionized. e.a.. the alkali metal halides, indicated thaiion p & ~are the &;dominant aggregates M ; the existence of more complex (charged) to about aggregates a t higher analytical concentrations is also indicated. Thus, to a first approximation all completely ionized substances behave as weak electrolytes in liquid ammonia because of ion association; in the electrolytic sense "strong electrolytes" do not exist in liquid ammonia solutions. Thus data from conductivity experiments on liquid ammonia solutions cannot be used to establish the ionization constant for weak electrolytes as it is often used for aqueous solutions. In liquid ammonia some fraction of the ions which might arise from the ionization of, e.g., a weak acid will form non-conductine ion-oairs which are indistineuishable from the molecula;spec& Other methods mustLe used to establish the ionization constants of weaklv ionized substances in liquid ammonia. Although the existence of ion-pairs in liquid ammonia solutions of electrolytes can be inferred from conductivity data, the results of spectroscopic investigations indicate that the formation of ion pairs does not affect the spectra of the species involved. An insight into the nature of solvated ionic species in liquid ammonia is obtained from a spectroscopic study of solutions containing the iodide ion (9). The alkali and alkaline earth iodides exhibit an ahsorotion band (Fie. 2) in the ultraviolet region of the spectrumihich has beenidentified with a charee-transfer-to-solventorocess. In effect the iodide occupies a solwnr cnvity and the excited stntc corresponds to an electron in an orbital drfinrd hv the orientated solvent molecules that form the cavity. ~ h u sthe , energy of transition is dependent upon the cavity size which can vary with temperature. Although the absorption maximum in liquid ammonia moves to lower energies with an increase in temperature (Fig. l), a t a given temperature the band is unaffected by the presence of relatively large amounts of inert salt. I t appears that the ion-paired species consists of associated solvated ions
O.! Table 7. ion
Figure 2.
-
(- -).
Theabsorption spectrum of iodide ion at -77'C
(-)andat -35°C
Na+ Kt N H ~ Ht OHNH,-
Equivalent lonic Conductances of Some Selected Species in Water and in Ammonia H20(25OC)
50.00 73.52 72.4 349.82 198.0
...
NHs (-33.5%) 130 168 131
... ...
133
Volume 55, Number 12, December 1978 / 755
The markedly basic nature o f liquid ammonia leads to a greater degree o f ioniza tion o f potential acids in this solvent.
rather than solvent shared ion-pairs he., pairs of ions sharing a single oriented solvent molecule) or contact ion-pairs. The conclusion is consistent with the previous discussion in which were developed arguments for strong ion-solvent interactions based upon the properties of ammonia. Chemical Phenomena
The high basicity of liquid ammonia leads to chemical phenomena that are unique to this solvent system and are of special interest to chemists. Acid-Base Behavior As is the case in water, acid-base reactions occur in liquid ammonia by the participation of the solvent. The fact that pure liquid ammonia conducts electricity, though to a lesser extent than does water (Tahle I), indicates that the solvent undergoes some degree of autoionization (eqn. (10)). 2NH3 e NHn+ + NH2-
(10) Comparison of eqn. (10) with the related process that is known to occur in water (eqn. (11)) 2H20 e H3Ot
+ OH-
(11) suggests that the NH4+ ion and the NH2- ion in ammonia correspond to the hydronium and hydroxide ions, respectively, in aqueous systems. The ion-product for NH3 (lOV%t -50°C) is considerably lower than that of water (10-l4 a t 25"C), which could he a reflection of the higher basicity of NH3 or the lower temperature of the measurement, or a combination of hoth factors; it is most likely, however, that the greater basicity of NH3 is the more important factor in the extent of autoionization of NH3 compared to that of H20. Neither the solvated proton in ammonia (NH4+) nor the amide ion (NH2-) exhibits an unusual conductance when compared with other ionic species in that solvent, which is in contrast to the marked differences in the conductivity of H3O+ and OH- and other ionic species in water (Tahle 7). These observations suggest that weaker hydrogen bonds are formed by the nitrogen atom of an ammonia molecule or that the proton affinity of ammonia is greater than that of water; both factors would suppress the Grotthuss mechanism (eqn. (12)) which has been suggested as the source of the abnormally high conductivity of protons and hydroxide ions in water.
water. For example, solutions of ammonium salts dissolve many metals to form hydrogen (eqn. (13)) Mg + 2NH4+
756 1 Journal of Chemical Education
+ + NH3
Mg2+ Hz
KNH2 + NH&l e KC1
(13)
+ 2NHs
(14)
the latter process can he followed to a stoichiometric end point using indicators such as ~henolnhthalein. ~ r a n k l i nrecognized (2) that some formal analogies could he drawn between the reactions of certain nitrogen-containina compounds in ammonia and those of the corresponding o x y gen-containing compounds in water. The analogy rests on a comparison of the ions generated in the two solvents, NH3 and HzO, and their relationship to the parent solvents (eqns. (10) and (11)). In the water system of compounds, many oxygencontaining compounds can he considered to arise from replacing one or more hydrogen atoms of a water molecule by other radicals. Thus, alcohols (ROH) and carhoxylic acids RCOOH are conceptually generated by the replacement of a hydrogen atom in water by an organic radical (R-) and an acyl radical RC=O, respectively. An ether (ROR) can be derived by replacing hoth hydrogen atoms by organic radicals. Franklin (2) suggested that just as these oxygen-based compounds behave in a characteristic way in aqueous solutions, a corresponding nitrogen-based system of compounds exists in liquid ammonia with similar characteristic reactions. A partial list of some aquo compounds and their corresponding ammono derivatives appears in Table 9. The markedly basic nature of liquid ammonia leads to a greater degree of ionization of potential acids in this solvent than in water. Some substances that are characterized as weak acids in water are more ionized (but not necessarily more dissociated) in liquid ammonia (e.g., compare acidic acid in water and ammonium acetate in ammonia), while other suhstances which show little or no acidic behavior in water give liquid ammonia solutions that exhibit acidic properties. Thus, a variety of substances such as acid amides (eqns. (14) and (16)), sulfonamides (eqn. (17))
Table 8.
These suggestions are borne out by experiment (Tahle 8).The hydrogen-bond energy between water molecules is about 14% greater than that between ammonia molecules, however, the proton affinity of ammonia is about 14% greater than the proton affinity of water. The high basicity of ammonia, as compared with water, is reflected in the greater stability of the solvated proton in ammonia. The ammonia solvates of strong acids such as HCI and H N O s e.g., NH4CI and NH4N03, can he readily isolated as stable solids, at room temperature, whereas the corresponding hydrates are markedly less stable. In other words, all the known soluble ammonium salts are acids in liquid ammonia because they can increase the concentration of the solvated nroton (NH*+) . . in this solvent system. The ammonium ion in liquid ammonia possesses many of the same chemical characteristics of the solvared p m t m in
-
and can be neutralized with ammonia solutions of alkali amides (eqn. (14));
A Comparison of the Hydrogen Bond Energy and Proton Affinity of Water and Ammonia H-bond energy. Kcallmole
Proton Affinity. Kcallrn~le
4.4 5.0
209 182
NH. Hz0
Table 9.
Some Aquo and Ammono Analogs
A ~ U Family O Hz0 KOH Z~(OH~ K2Zn(OHL SOAOHh ROH RC(0)OH ROR RC(0)OR'
Ammono Family NHs KNH2 Zn(NHd2 K2Zn(NH2), SOdNHh RNH2 RC(0)NH2 RN I RC(O)NR',
+
RSO2NHz KNHz
-
[RSOzNHIK + NH3
(17)
amines (eqn. (18)) and hydrorarhons (eqn. (1911 undergo reaction with metal amides to yield metal derivatives that bear a formal analogy to salts. R2NH + KNHz- RzNK + NH3 (18) (19) (C6H&CH + KNHz [(C6H&C]K + NH3 Acid-bare processre are also observed in liquid ammonia solutions containinr certain metal ions. The reactions of metal compounds in liquid ammonia can be broadly classified as ammoniation reactions, in which ammonia acts as a Lewis base (eqns. (20) and (21))
-
-
TiCL + 4NH3 TiC12(NH& (20) Mn(SCN)z + 6NHs [Mn(NH3)d(SCN)z (21) and ammonolysis reactions, which involves coordinated solvent molecules acting as protonic acids (eqns. (22) and (23)). Mn(NH3)c2++ NH3 a Mn(NHdsNHzt + NHI+ (22) AI(NH3)P NH3 a AI(NH&NHz2++ NHIC (23) In fact, compounds containing elements such as zinc are amphoteric in liquid ammonia as well as in water (eqns. (24) and (25)). (24) Zn(NHs),Z+ 2NH2- t Zn(NHz)z(NH& + 2NH3 (25) Zn(NHz)z(NH.&+ 2NHZ- a Zn(NHz)? + 2NHs The insoluble zinc amide precipitates from an ammonia solution of ZnCln upon the addition of KNHz (eqn. (24)) and redissolves in excess KNHz (eqn. (25)); these observations parallel the chemistry of zinc in aqueous solution. (22)-(25), Deorotonation of comolexed amines, e.g., .. ems. . is a characteristic process in ammonia; evvn ethylenediamine comnlrxes of transition metals are sufficiently acidic in ammoAa to undergodeprotonation by strong basis. For example, the complex Pt(e&Iz reacts stepwise with KNHz to form compounds which contain 1 to 3 less hydrogen ions (eqns. (26)-(28)). Pt(en)& + KNHl F. [Pt(en-H)en]I+ NH3 + KI (26) [Pt(en-H)en]I+ KNHz t [Pt(en-H)zl+ NHs + KI (27) [Pt(en-H)z]+ KNH2 a [Pt(en-2H)(en-H)]K+ KNH3 (28) These processes can he reversed by the addition of an ammonium salt. The ability of liquid ammonia, compared with water, to enhance the apparent acidity of substances is also manifested in the case of complexes; the aqueous pK, of [Co(en)#+ in water is -12.5 whereas this substance has a pK, of -7 in liquid ammonia (10).
+
+
~
.~ ~
~
~
Although there are numerous observations which suggest that certain substances behave as acids in liquid ammonia, there has been little success until recently (11) in establishing an acidity scale in liquid ammonia. A combination of nmr spectroscopy and ultraviolet-visible spectroscopy has led to the acidity scale shown in Table 10 for a series of carbon and nitrogen acids. These substances in Table 10 are either neutral or basic in aqueous solutions, but become more or less acidic in liquid ammonia in agreement with previous arguments. Metal Solutions
Perhaps one of the most unusual properties of liquid ammonia is its ability to dissolve active metals to form homogeneous solutions that are active reducing agents. Dilute metal-ammonia solutions are blue, but in more concentrated solutions a bronze phase separates which floats on the blue solutions. The electrical, optical, and magnetic properties of these solutions indicate that they contain solvated electrons and solvated cations. Although there is still a great deal of controversy concerning the detailed constitution of the solvated electron in liquid ammonia and the species which may be present a t higher metal concentrations, it is generally agreed that the dilute blue solutions contain electrons trapped in solvent cavities (Fig. 3). These charged species appear to be electrostatically associated with solvated cations in much the same way as are the ordinary ion-paired species in this solvent. Metal-ammonia solutions are perfectly stable in the absence of catalysts such as finely divided metals or metal oxides. When these solutions decompose, hydrogen is liberated (eqn. (29)). (29) e(NH& + NHa t %HZ+ NHz- + XNHB Metal amides also appear to be good catalysts for this reaction which means that reaction (29) is autocatalvtic. The reaction of solvated electrons with am&onia to f o r k hydrogen (eqn. (29)) . .. is reversible. i.e.. it is nossible to DreDare . . the characteristic blue electron sohtiois by introduriny hydrogen under oressurr into a solution conraining NHI-. The value for the kquilihrium constant for eqn. (29)has been estimated from electron spin resonance data and from optical spectroscopy t o he 5 X 104at 25°C (12). The utility of liquid ammonia as a medium in which to study redox processes using conventional electrochemical techniques has not been widely recognized even though metal-ammonia solutions have found extensive use as chemical reducing agents in the synthesis of a variety of organic
Table 10. Acidities of Selected Carbon and Nitrogen Acids in Liquid Ammonia Compaund
P Ka
Di-(4-methoxyphenyl)methane Diiptoly1)methane N,Ndimemyl-pphenylenediamine Triphenylmethana 4-Methylaniline Kthylaniline 2,5-Dimethylaniline 3.5-Dimethylaniline 4-Fluoroaniline Aniline CChioraaniiine 3-Trill~~r~methylaniline 3.4-Dichioroaniline 2.4-Dichloroaniline Uyanoaniline 3.5-Bistrifluoromethylaniline 2.5-Dichloroaniline
28.6 27.1 25.7 -23.6 22.1 22.0 21.8 21.7 21.4 21.2 19.9 18.7 18.0 16.7 15.9 15.6 15.5
Figure 3. The propetties of metal-ammonia solutions suggest that several species might be present, in concentrated solution, solvated electrons can be 8550Cialed wiih cations by sharing solvent molecules (A). In dilute solutions, me soivated electron is believed to occupy a cavity formed by solvent molecules (B).
Volume 55, Number 12, December 1978 / 757
Liquid ammonia has proved to he a useful medium for the stabilization o f species containing a toms in low oxidation states.
compounds. As discussed in the first section of this paper, ammonia is a very good solvent for organic substances; a wider variety of compounds can he dissolved if a co-solvent is present. Thus, it is possihle to obtain experimental conditions suitable for the homogeneous reduction of covalent molecules. Numerous kinetic studies involving reduction by metalammonia solutions have been made in an attempt to formulate general mechanisms by which such processes occur. In general, electron addition to saturated compounds lends m. to hond cleavage with the formation of either radicals (. e . (30)) andlor ionic species (eqn. (31)). X-Y+e--X'+YX-Y
+ 2e--X-+
(30) Y-
(31)
The ions formed in such reactions are usually conjugate bases of very weak acids and abstract protons from either the solvent (eqn. (32))
and anions which are conguate bases of weak acids can abstract protons (eqns. (49)-(51), from the solvent. 'X-Y-X-Y-YXXY-
-
+ NH8 + 2NH3-
'XYH + NH1-
(49)
+ 2NHzHYXXYH + 2 NHPHXYH
+ 2NH3-
(50) (51)
Again we have numerous examples of these possible pathways, hut only a representative sample is given in eqns. (52)-(58).
+
+
+ -
RCH=CH2 2NH3 2Na (CsHS)2C=C(C6H~)z + 4NH3 RCECR'
+ 2Na + NH3
RzC=O
-
+
RCHzCH3 2NaNHz
2Na
Z(CsH&CHz
+ 4NaNHZ 152) ~--,
+ 2NaNHs
+ 2Na + 2NH3
X-+RH-HX+R-
(33)
formed in process (30) can take up another electron (eqn. (34) X+e--X-
(34)
undergo dimerization with other radicals (eqn. (35)) 2X'- X-X
(35)
or they can abstract radicals from molecular species (eqn.
(36)). X'+X-Y-X-X+Y
(36)
The dimerized species can, of course, arise from the reaction of an anion with the molecular species (eqn. (37)). Although there are numerous examples of each of these pathways availahle for discussion only a representative member are presented in eqns (38)-(42).
+ + + -
+ 2Na [(CH3)3SnjNa + NaBr (38) + 2Na + NH3 [ArOjNa + Ar'H + NaNH2 (39) RC(0)OR' + 2Na + NH3 [RCNHjNa + [R'OINa + Hz (40) R4NC+ 2Na + NH3 R3N + RH + NaNHz + Na+ (41) (CH& SnBr
ArOAr'
-
2R2S 2Na (CsHs)&NH2 2RC1+ 2Na
2Na
NHB
2RSNa + R2
+ NaNHz RH + RNHP + 2NaCI
[(CsH&C]Na
(42) (43) (44)
The reaction of solvated electrons with unsaturated bonds generally follows similar pathways except that hond fission need not occur. Either one or two electrons can he added to unsaturated bonds to form a radical anion (eqn. (45)) or a dianion (eqns. (46) and (47)). Radicals can dimerze (eqn. (48))
758 1 Journal of ChemicalEducation
S,
(54)
H' H '
-
RzCHOH
+ 2NaNHz
02+e--0s-
or from other strong proton donors in the system (eqn. (33)). The radicals
(52)
+ 2e-
9Pb + 4e-
-
(55) (56)
SZ2-
(57)
Pb&
(58)
Liquid ammonia has proved to be a useful medium for the stabilization of species containing atoms in low oxidation states, undoubtedly because the solvent is resistant to highly reducing conditions. For example, polyhomoatomic anions are formed when the heavier members of Group IV, Group V, and Group VI elements are allowed to react with metal ammonia solutions. Little is known concerning the constitution of these highly colored species, but their stoichiometry, e.g., NaPbz, NaPh4, N%Phs, and Li4Pb,, suggests they might be similar to the so called intermetallic compounds. More of the Group V elements form homonuclear polyatomic species, Nap3, KP5, &As4, NasShs7, NaBis, NaBisri, than do those of Group IV. All the elements of Group VI yield polyatomic anions-M02, M202, M204,M2Xn (X = S, Se, Te; n = 1-7) -upon reaction with metal-ammonia solutions. Metal-ammonia solutions will react with substances that arc acidic in liquid ammonia to form metal salts and hydrogen gas. Many covalent hydrides and their derivatives react with metal-ammonia solutions to form the corresponding metal salts that are stable in liquid ammonia (eqns. (59)-(61)). AsHx + Na pH3 + Na
--
PHa + 2Na
NaAsHs
+ %Hz
NaPHz + 'hH2 NazPH
+ HZ
(59) (60) (61)
In fact, some organic compounds which are not sufficiently acidic t o react with amide ions react withmetal ammonia solutions to form stable metnl derivatives which are usrful as svnthetic intermediates. Fur exarn~le,alkali metal derivatives prepared in liquid ammonia can be used as alkylating agents (em. (62)). R-M+ + CH31 RCH1 + MI (62)
-
R = CHaCONH, ( C s H d G C6HsNH, CzHsO, C~HSS, (C6Hd3Ge. (CsHdzP
Finally, metal-ammonia solutions find considerable use in inorganic svnthesis in the formation of derivatives of transitionmetali in low oxidation states. In many instances the metallic elements are precipitated when their compounds are treated with metal ammonia solutions. However, a complex anion containing chromium in a zero oxidation state can he obtained from the reduction of potassium hexacyanochro-
mate(II1) with potassium in liquid ammonia (eqn. (63)). K3Cr(CNh + 3K KGCI(CN)~ (") This complex is a useful intermediate for the formation of a variety of other neutral compounds containing chromium in a zero oxidation state (eqns. (64) and (65)). +
+
-
K6[Cr(CN)s] 3D CrD3 + 6KCN (64) D = (CfiH5)2PCH3P(C6H5)2. bipyridyl, 1,lO-phenanthroline
solvent as being secondary only to water have been amply verified in the intervening years. In a certain sense ammonia is. uerhaus, a more versatile solvent than is water because of itsability to soluhilize, without reaction, highly negative or reducing species. Much interesting chemistry has been initially revealed in liquid ammonia solutions which could not have been easily achieved in other solvents. Literature Cited
The elements of the nickel-subgroup also form complex species in which the metal exhibits a zero oxidation state (eqns. (66)-(68))
-
K2M(CNl4+ 2K K&M(CNh M = Pt, Ni, Pd K2Ni(C=CHh + 2K KaNi(C=CH)&
+
K2Pd(CN)2(C=CR)2 2K
(66)
(67) KzPd(C=CR)z + 2KCN (68)
and these species also can be used as intermediates for other complexes (eqns. (69)-(70)).
+
K,Ni(CN)& 4L
-
NiLl
+ 4KCN
(69)
The prophetic comments of Franklin and Kraus in concerning the (then) perceived abilities of ammonia as a
and Harris, P. M.. J. Chem. Phys.. 35.1730(1961). (5) Gutmann. V.. Angeiu. Chemlnl~rnol Ed., 9,843 (1970). (61 (a) Gutmann,V..andWychera,E.,Rau.Chem. Min., 3,941 (19661.(b)Gutmann,V., and Wyehera.E..Nvel. Chem.Lat., 2,ZSI (1966). (7) London, F.. Trans. Forodoy Soe., 33.8 118371. (8) Hildehrand, J. H..and Scott, R. L.. "Solubilities of Non-Electrolyte&" Chapter V. Reinhold. New York. 1950 19) Nelson, J. R.,andLagourki,J. J.,J. Chem.Phys.,70.1492 (19661. (10) Moczygemha,G. A.,and Laxowski, J. J.,J Coord. Chem., 671 (1976). i,l l 1~Takemnh.1. H..and Lamwski. J. JLInore.Nucl Cham. Lett.. fi.315119701 ~ , (12) Kirschke,E. L a n d Jolly, W. L:,lno;g. C&rn.. 6.86511967) ~
General Bibliography
mania." Voi. 1, part 2, Interscience. New York, 1963. Jdl" W. L.. and Hailad'. C. J.. "Liquid Ammonia." in "Nonaqueoua Solvent Systems." (EditorT. C. Waddingtan)Academic Press, New Ymk. 1965. T ~ O ~ ~ c. O . ,~ . ,, EJ .I i~n ~ i~q u~i d~~ ~m m~o n i a . " ~ ~ s r e n d o n ~ r e ~ ,1976. ~~fod.
Volume 55. Number 12, December 1978 1 759
