non-aqueous solvents in inorganic chemistry

Non-aqueous Solvents in Inorganic Chemistry

by

A. K. HOLLIDAY, Ph.D., D.Sc, F.R.I.C. Reader in Inorganic Chemistry, The University of Liverpool

and

A. G. MASSEY, B.Sc, Ph.D., A.R.I.C. Lecturer in Inorganic Chemistry, Queen Mary College,

University of London

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PREFACE

BY THE end of the nineteenth century, two non-aqueous inorganic solvents — liquid ammonia and liquid sulphur dioxide — had already been investigated. Today, a large number and great variety of such solvents are known, but the early work has had a great influence on the way in which our knowledge of these newer solvents has developed. The investigations on ammonia and sulphur dioxide were notable in being concerned experimentally with the manipulation of low-boiling, reactive liquids and theoretically with the extension of the simple acid-base concepts of the original Ionic Theory to non-aqueous systems. These trends can be seen very clearly in Franklin's classical monograph on liquid ammonia chemistry — The Nitrogen System of Compounds (1935). The influence of this early viewpoint can be realized when one recognizes how many of the newer non-aqueous solvents are low-boiling reactive liquids and how often acid-base concepts have been invoked to explain the observations made. Much of this later work has been very fruitful, but until recently attention has not been directed to the use of substances with a higher-temperature liquid range, or to the use of non-aqueous solvents for (e.g.) oxidation or synthetic reactions. Only in the last decade has work on solvents such as pure sulphuric acid and fused salts begun to show valuable results, and this has been due in great measure to the development of newer methods — and especially spectroscopic methods — of studying solute-solvent systems.

In writing this book, our aim has been to give a concise treatment of the important inorganic non-aqueous solvents, emphasizing why they do in fact exhibit solvent power, how they are prepared and handled experimentally, how they can be used as

vii

viii PREFACE

media for the synthesis or analysis of inorganic and organo-metallic compounds, and how far the various acid-base concepts can be useful in accounting for many (but not all) of the reactions observed. It has not been easy to achieve this aim in the final chapter, on high temperature solvents, because the latter (mainly fused salts) have as yet been little used for synthesis or analysis in the laboratory, although their use on a large scale has been known for many years. However, a considerable amount of physico-chemical information is now available for fused salt systems, and we have endeavoured to show the relevance of this to the investigation of reactions in these systems. This field of study is likely to be of great importance in inorganic chemistry in the future.

This book is intended primarily for the undergraduate reader — both for the intending Chemistry Honours or R.I.C. graduate and the non-specialist student of chemistry. We have therefore tried to present the subject-matter in a simple and readable form, without the inclusion of elaborate tables of properties and with the minimum of detail necessary for comprehension. Therefore, those working for the A- and S-level chemistry examinations for the G.C.E. could read much of the book with profit; and the research student who aspires to work in the field of non-aqueous solvents will, it is hoped, find this book a useful introduction to a fascinating branch of inorganic chemistry.

A. K. H. A. G. M.

CHAPTER I

THE NATURE AND SCOPE OF INORGANIC NON-AQUEOUS SOLVENTS

WATER is such a common and therefore readily obtainable substance that it was an obvious choice as a solvent by the very early chemists. The extraordinary versatility of water as a solvent was soon recognized and the solubilities of many substances were determined over the range 0-100°C. It was not surprising that other solvents were almost completely neglected until the development of organic chemistry produced, simultaneously, organic substances which were often insoluble in water and organic liquids which could be used as solvents instead of water.

The process by which a simple organic molecule, e.g. a paraffin hydrocarbon, dissolves in (say) benzene is comparatively simple. The relatively weak intermolecular forces between the solute hydrocarbon molecules permit dissolution and are replaced by solvent-solute interactions which, again, are weak; the "driving force" leading to solution is here the change to a state of higher entropy which the solute molecules attain by solution and hence the solubility usually increases markedly with temperature.

By contrast, the process of solution in water is always complicated, and even now it is not in general possible to make quantitative predictions about solubilities. The complications arise because in water there are already relatively strong, specific and directed intermolecular forces — hydrogen bonds — which give to the liquid water some semblance of an ordered crystalline structure. The mechanism by which, for example, an ionic solid

1

2 NON-AQUEOUS SOLVENTS IN INORGANIC CHEMISTRY

such as sodium chloride dissolves, is not then simply a matter of the reduction of the strong inter-ionic attractions in the crystals by a continuous medium of high dielectric constant. The reorientation of the solvent consequent upon solvation of both cation and anion plays an important role in the energetics of solution. However, covalent solids can sometimes dissolve in water and here hydrogen bonding between solute and solvent is generally the factor favouring solubility; solutes containing —OH or —NH2 groups, such as alcohols, carbohydrates, amines and amino-compounds are typical examples of this type of behaviour.

Although many non-ionic substances undergo hydrolysis in water by an essentially bimolecular process in which a water molecule "attacks" the solute molecule, apparent hydrolysis — appearing as a decrease in pH — is not uncommon in solutions of salts of multivalent ions. Here, however, the change in pH is due to an increased dissociation of water molecules co-ordinated around highly charged cations, i.e. the process is essentially unimolecular.

The frequent occurrence of hydrolysis, real or apparent, does limit the usefulness of water as a solvent; the other limitation is of course the rather narrow liquid range which does not permit the study in solution of substances of low thermal stability or of species stable only at high temperatures. The use of organic solvents to overcome either of these limitations is itself subject to the severe limitation imposed by the very low solubility of many inorganic substances in such solvents; for this reason, organic liquids have not been widely investigated as solvents for inorganic systems, although solvents such as dimethyl formamide, dimethyl sulphoxide and the "glymes" (e.g. "diglyme", 1,2-dimethoxyethane) are now used to an increasing extent.

The approach to the problem of finding suitable non-aqueous solvents for inorganic chemistry has not in general been made systematically because of the difficulties, already mentioned, which attend any quantitative theoretical approach to the problem of solubility. In practice, two considerations have often influenced the choice of suitable solvent systems ; the possibility of dissociating

NATURE AND SCOPE OF INORGANIC NON-AQUEOUS SOLVENTS 3

the solute into ions and the possibility of setting up acid-base systems in which the solvent might participate. Hence solvents possessing a dielectric constant which is greater than about 10 and an electrical conductance, possibly due to some degree of self-ionization, have often been chosen for study. Even a low degree of self-ionization does not, however, preclude the use of a solvent for ionic substances, thus, for example, dinitrogen tetroxide, N204, whilst having a low degree of self-ionization can still be used as an effective solvent for many salts.

It is the application of the acid-base concept which has been the dominant theme in the development of non-aqueous solvents in inorganic chemistry. Franklin, who did much pioneering work with liquid ammonia, postulated the following self-ionization equilibrium for this solvent:

2NH3 ^ NH4+ + NH2-

compare 2H20 ^ H30+ + OH-

He observed acidic behaviour by the ammonium salts (for example NH4C1) and the basic behaviour of alkali amides (such as potassium amide KNH2) in liquid ammonia and was led to formulate new acid and base definitions; viz. an acid was a substance giving a cation characteristic of the solvent and a base was a substance giving an anion characteristic of the solvent. Otherwise, acids and bases had their characteristic properties, for example, an acid plus a base gave a salt and solvent; acids dissolved metals to produce salts, and so on.

It is important to note that Franklin's definitions are not restricted to hydrogen-containing substances; thus acid-base behaviour arising from the following possible self-ionization equilibria are covered by them :

Acid Base liquid N 2 0 4 ^ NO+ + i l liquid 2BrF3 ^ BrF2+ + BrF4~ liquid 2S02 ^ S02+ + S03

2-


Whilst there is good evidence for the first two equilibria, the third is very doubtful, but the important point is that investigations in all three liquids as non-aqueous solvents have been influenced and often aided by Franklin's definitions as have other investigations of hydrogen-containing solvents, for example liquid hydrogen fluoride (Chapter IV). The weakness of the Franklin definitions is that they are so wide as to be incapable of quantitative application — we cannot use them to compare one solvent with another in any quantitative sense.

If we consider those hydrogen-containing solvents which are protonic, then the familiar Lowry-Bronsted acid-base theory can be applied. Consider again the equilibria

2NH3 ^ NH4+ + NH2-2H20 ^ H3O+ + OH-

Here the Lowry-Bronsted definition concerns the processes

NH3 + H+ ^ NH4+ H20 + H+ ^ H3O+

and the position of equilibrium in each process is determined by the proton affinity of ammonia and water respectively. Since the proton affinity of ammonia is greater than that of water, addition of ammonia to an aqueous solution of a strong acid HA (present as H30+ and A" ions) produces the ions NH4

+ and A~

NH3 + H30+ ^ NH4+ + H20

and we say that ammonia is a more basic solvent than water. But addition of water to a solvent such as glacial acetic acid gives immediately the reaction

HOAc + H20 ^ H30+ + OAc-

Hence here the water is the stronger base or, putting it another way, pure acetic acid is a more acidic solvent than water. So far as protonic non-aqueous solvents are concerned, these differentiations from water — more basic or more acidic — have proved

NATURE AND SCOPE OF INORGANIC NON-AQUEOUS SOLVENTS 5

very useful guides to behaviour with inorganic solutes (see further Chapters III and IV).

Two other acid-base concepts require mention. The protonic solvents water and ammonia, as we have seen, can behave as bases by uniting with a proton. This, in terms of electron-pair bonding, requires donation of an electron pair from base to proton. G. N. Lewis proposed an extension of this idea from protonic acids and bases to the definition of any base as an electron donor, and any acid as an electron acceptor. Hence, for example, boron trifluoride is a "strong" Lewis acid because the position of equilibrium in the reaction

BF3 + NMe3 ^ Me3N.BF3

is well to the right. The Lewis definitions are clearly applicable to a wide range of substances, but they are not capable of quantitative application in many cases. Further reference to the Lewis definitions is made in Chapter VI.

In the Lux-Flood acid-base concept, oxide ions are regarded as the transferable species corresponding to protons in the Lowry-Bronsted scheme, e.g.

SO«, + O2- = S032~

acid base

Note that the base gives up oxide ions, and the acid gains oxide ions. This idea is discussed further in Chapter VI.

Experimentally, studies in a non-aqueous solvent require, first, preparation and purification of the solvent, then means of manipulating it, and finally methods for the observation of phenomena in it. Purity is of prime importance, and methods of purification (and of preparation where appropriate) will be mentioned briefly as each solvent is discussed. Means of manipulation depend upon physical properties, particularly the melting point to boiling point temperature range, and therefore short lists of the relevant properties of each solvent will be given. Observations of phenomena in the "older" solvents (e.g. liquid


ammonia, sulphur dioxide) were made by classical methods; in[the "newer" solvents (i.e. pure sulphuric acid, fused salts) such methods are insufficient, and a wide variety of physical methods must be used. It is worth noting, however, that the application of the newer methods to a classical problem — the nature of metal-ammonia solutions — has not been completely successful in elucidating the structure of these systems, which are, nevertheless, extraordinarily useful in preparative [inorganic and organic chemistry.

CHAPTER II

LIQUID AMMONIA

ALMOST 100 years ago the English chemist Gore investigated the solvent properties of liquid ammonia for some 250 substances. His work was greatly extended by the American Franklin who, over the period from 1900 to 1935, studied many chemical reactions in the solvent, formulated the definition of acid and base discussed in Chapter I and systematized liquid ammonia chemistry by relating solutes and their reactions to the characteristic element of ammonia, nitrogen, in a similar way that water-soluble substances so often relate to oxygen. Thus, for example, nitrides were considered to correspond to oxides and amides to hydroxides.

Since the time of Franklin, and notably since World War II, liquid ammonia chemistry has continued to expand until it is now a familiar solvent in organic chemistry, where solutions of the alkali metals are versatile and widely used reducing agents. Additional uses in inorganic chemistry are in the production of unusual valency states, unusual metathetic reactions and in reactions where hydrolysis would be a restriction.

This extensive usage is readily explained since, of all non-aqueous solvents, liquid ammonia approaches most closely to water in its physical properties as Table 1 shows. Like water its heats of fusion and vaporization (and melting point and boiling point) are higher than those of other hydrides in the same group of the Periodic Classification, such as phosphine PH3 or arsine AsH3. This indicates some considerable degree of structural association in the liquid and solid states attributable to hydrogen bond formation but, as the dipole moments indicate, the extent of

7


TABLE 1

Physical properties of ammonia and water

Property

Boiling point, °C Freezing point, °C Density, g/ml Heat of vaporization, kcal/mole Heat of fusion, kcal/mole Viscosity of liquid at 25°C, cP Dielectric constant Dipole moment, Debye units Polarizability, cm8 (x 1024) Specific conductance, ohm-1 cm-1

Ammonia

-33-4 -11Ί

0-68 (-33°C) 5-58 1-35 0135

22(-33°C) 1-46 2-25

4 x 10-10(-15°C)

Water

100 0

0-96 (100°C) 9-72 20

0-891 80 (0°C)

1-84 1-49

4 x IO"8

association is likely to be less with ammonia than water. This lower intermolecular attraction means lower viscosity and hence a high mobility of ions in liquid ammonia f. Furthermore, ammonia and water are non-corrosive liquids of limited reactivity and in both cases the liquid temperature range is reasonably wide. This means that manipulations in liquid ammonia are not restricted by attack on the containing vessels and a large range of substances are soluble without rapid solvolysis.

There are several ways in which Uquid ammonia is used experimentally as a solvent. Much of the earlier work was carried out at ordinary temperatures under pressure and therefore required only limited use of the rarely available cooling baths. It is worth noting that many of the solubility values in liquid ammonia were determined at room temperatures in this way and solubility data over the liquid range at atmospheric pressure, —78 to — 33°C, are remarkably sparse.

Franklin was able to devise simple pieces of apparatus in which experiments normally carried out in test-tubes with water could be imitated with liquid ammonia under pressure. In later forms of

t Unlike the HsO+ ion in water, the NH4+ ion has no abnormally high mobility

in liquid ammonia.

LIQUID AMMONIA 9 apparatus the ammonia was cooled below — 33°C and transferred from one vessel to another by means of nitrogen pressure. In this way quite elaborate manipulations — titrations, filtrations and extractions — could be carried out without much difficulty.

FIG. 1. Manipulations in liquid ammonia; ammonia from the storage bulb F may be transferred to the cooled reaction vessel A, to which appropriate reagents are added through E. Warming A and cooling C permits filtration through the sinter B, and the filtrate can be recycled from C through the glass wool trap at D to wash the solid in A. H9 I

and / are connected to mercury manometers.

The convenient melting point of ammonia means that baths cooled with solid carbon dioxide (—78°C) can maintain the liquid at a low vapour pressure but just above the melting point.

B


In reactions where gaseous or volatile products are probable it is now more convenient to manipulate liquid ammonia solutions in a vacuum system by transferring the solvent and other volatile materials by means of traps and cooling baths — the latter being used to fractionate (i.e. separate) the volatile products from the solventf. A part of such an apparatus is shown in Fig. 1.

For larger scale work, particularly in organic preparations where very precise conditions are not essential, it is often convenient to use liquid ammonia in open Dewar vessels under a well-ventilated hood, the liquid being transferred by pouring as with water. The rather high heat of vaporization implies that solutions in liquid ammonia open to the atmosphere do not evaporate at a rapid rate.

Solubilities in Liquid Ammonia An inspection of Table 1 suggests that in general the solubilities

of various substances in liquid ammonia would follow the same pattern as in water — many salts soluble, organic substances soluble only if capable of hydrogen bonding with the solvent— because liquid ammonia is, in ways already mentioned, a similar liquid to water. This idea is borne out in practice provided that it is not pressed too far. The main difference between the two solvents is one of degree: most substances are less soluble in liquid ammonia than in water. This would be expected for ionic substances since the dipole moment and dielectric constant of ammonia are both smaller than for water; also the smaller dipole moment is, in part, an indication of weaker hydrogen bonding power and therefore of lower solubility for substances such as alcohols.

There is, however, one other significant difference between the two solvents in that the ammonia molecule has a higher polariz-ability than that of water. Hence, although ion-dipole attraction is less in the case of ammonia, ion-induced-dipo\e interaction

t Liquid ammonia may be dried over sodium; any hydrogen formed is pumped away and the dry ammonia distilled off.

LIQUID AMMONIA 11 may make a more significant contribution to the solvation energy especially when the ion concerned has high "polarizing power", i.e. is small and highly charged. The effect of this is seen in the lithium salts which are generally less soluble in water than those of the other alkali metals because of their higher lattice energy. In ammonia the ion-induced-dipole energy contribution outweighs the lattice energy effect in many cases, giving a higher solubility for the lithium salts than for those of the other alkali metals. For salts of the latter, and for salts in general, the normal rule is lower solubility in liquid ammonia than in water and lower solubility for higher lattice energy.

In water, salts possessing relatively high lattice energies may still be moderately soluble. This is notable even when the ions carry charges > 1 as in the case of some aluminium salts, and for sulphates and phosphates. In liquid ammonia very few salts containing polyvalent ions of this kind are soluble because the lattice energy is too high and is not "outweighed" by ion-solvent interactions; hence appreciable solubility is virtually restricted to 1 : 1 electrolytes.

There are, however, two further effects which increase solubility in liquid ammonia for specific groups of substances, both effects stemming from the higher polarizability of the ammonia molecule. The first of these concerns transition metal salts. When one of these is placed in a solvent, the transition metal cation-solvent interaction will be greater than can be accounted for by purely electrostatic ion-dipole or ion-induced-dipole interaction to an extent determined by the ligand field stabilization energy associated with the grouping of solvent molecules in a definite configuration as ligands around the metal ion. For many transition metal cations the stabilization energy is greater for ammonia than for water and this implies heats of solvation, greater for ammonia than for water, often by several kcal/mole, and hence higher solubility.

The effect of this is to increase solubility in ammonia as against water and again this increase (as with the lithium salts) is sufficient to outweigh adverse effects of high lattice energy — thus, for example, anhydrous cupric sulphate even though a 2 : 2 electrolyte


is nevertheless quite soluble in liquid ammonia. (It must be noted in any discussion of the solubilities of transition metal salts, very few of the latter are known in the non-sol vated state; most "solubilities" in water aie those of hydrated salts.)

The second effect associated with the higher polarizability of the ammonia molecule becomes important with solutes which are themselves highly polarizable or which contain highly polarizable ions. Dispersion forces, which depend upon the polarizabilities, then become important between solvent and solute and may make an appreciable contribution to the solute-solvent interaction and so favour solubility. Thus, aliphatic hydrocarbons of low molecular weight are virtually insoluble in both water and liquid ammonia, but the aromatic hydrocarbons where the conjugated ring systems give rise to considerable polarizability, are generally more soluble in ammonia than in water. The same principles apply when organo-metallic compounds are examined (for example cyclopentadienyl metal compounds are often soluble) and when a salt contains large, highly polarizable anions such as iodide or thiocyanate; hence iodides and thiocyanates (which in any case have low lattice energies) often exhibit high solubilities in liquid ammonia.

Ammonolysis Hydrolysis and ammonolysis are, of course, examples of the

general phenomenon of solvolysis. Reference has already been made (p. 2) to the two kinds of "hydrolysis", unimolecular and bimolecular, e.g.

[Μ(Η 2 0)^ ^ = ^ [M(H20)n_1(OH)]<^)+ + H3O HaO

BX3 + H20 = = * BX2(OH) + H3O+ + X-Similar equations can be written for ammonolytic reactions in liquid ammonia:

NH8

[ΜίΝΗ.) JH- ^ = ^ [ M i N H ^ ^ N H ^ ] ^ 1 ^ + NH4+ NH 8

BX3 + NH3 = * BXaiNHa) + NH4+ + X~

LIQUID AMMONIA 13 It should be noted that the first equation of each of these pairs really represents only an acid-base reaction (in the Lowry-Bronsted sense), the solvated ion (acid) losing a proton to the solvent (base). In the second equation, the BX3 molecule loses an X and gains an OH or NH2 group from the solvent, by what may be termed a metathetic reaction involving the solvent. However, it is convenient to use solvolysis to describe both kinds of reaction. Two general considerations can be given before ammonolysis is discussed in detail. First, N—H bonds of ammonia are less easily broken than O—H bonds of water and hence ammonolysis is, in general, less likely to occur than hydrolysis. This has important consequences since it means that liquid ammonia can often be used in the preparation of substances which would undergo hydrolysis if the synthesis were attempted in aqueous solution; examples of this will be given in the next chapter. The metal-ammonia solutions to be discussed later in this chapter also owe their stability to the fact that the reaction (M = alkali metal):

M + NH3 -> MNH2 + ±H2

occurs much less readily than the hydrolysis

M + H20 -> MOH + £H2

Secondly, even if ammonolysis does occur, the rate at which it does so will depend upon temperatuie. In many cases, since the reactions in liquid ammonia can be carried out at temperatures as low as — 78°C, the rate of ammonolysis can be made slow enough not to matter especially where, for example, metathetic ionic reactions (which are still rapid at low temperatures) are carried out.

As with hydrolytic reactions, ammonolytic reactions will be favoured if an ammono-base is present such as potassium amide (cf. KOH); but here we shall assume that only solvent is involved.

Ionic compounds — salts of many of the metals — do not in general undergo ammonolysis, although the cations are


ammonated in solution. Relatively little is known of equilibria such as

[M(NH3)6]3+ ^ [M(NH,)5NHJ*- + H+ I

[M(NH3)4(NH2)2]+ + H+ etc., where M is a transition metal.

With wholly or partly covalent compounds, ammonolysis is more common, the extent to which it occurs being determined in many cases by the same factors as affect hydrolysis. Thus to take an example from Group III of the Periodic Classification, the covalent BX3 compounds ammonolyse and hydrolyse readily when, for example, X = halogen or some similar electronegative group but not when X = alkyl ; the same applies to the compounds of silicon (SiX4) in Group IV and phosphorus (PX3) in Group V. In these reactions, intermediate formation of an ammonia adduct precedes ammonolysis ; for this reason the carbon tetrahalides do not undergo ammonolysis at low temperatures since carbon has no vacant orbital of suitable energy to accept electrons from an ammonia molecule. This behaviour is parallel to that observed with water; carbon tetrahalides are not easily hydrolysed. The compound CI4.2NH3 formed when carbon tetraiodide reacts with ammonia at — 33°C is thought to be held together by purely van der Waals forces, the ammonia being lost when the solid is warmed up to room temperature. Rather similar behaviour to that just described is observed with many covalent halides of the transition metals such as titanium tetrachloride TiCl4, molybdenum pentachloride MoCl5 and the halides of tin(IV). In these ammonolyses, reaction may be complete

BC13 + NH3 -> [C13B-NH3] j N H 8

C1B(NH2)2 £^- C12BNH2 + NH4C1

| N H 3

B(NH2)3

LIQUID AMMONIA 15 or incomplete as in

ZrCl4 + 2NH3 -> ZrCl3NH2 + NH4C1 where no further ammonolysis takes place. In certain cases decomposition of the amide to the imide takes place as happens, for example, in the ammonolysis of germanium tetrachloride where the final product of the reaction is the imide

GeCl4 — ^ Ge(NH2)4 -> Ge(NH)2 + 2NH3

This step can be taken further since, on heating, several amides and imides can evolve ammonia until finally the nitride of the element remains.

In the presence of an excess of an ammono-base (i.e. an ionic amide) some of the ammonolytic products will dissolve to form soluble complexes. Thus in the presence of potassium amide, iridium(III) bromide in liquid ammonia gives first the amide Ir(NH2)3 and this then dissolves in excess of the potassium amide

Ir(NH2)3 + 3KNH2 -> K3Ir(NH2)6

Here again there is a parallel with the amphoteric behaviour of some metal hydroxides in presence of excess hydroxyl ion.

Metals in Liquid Ammonia It is probable that liquid ammonia, despite its resemblance to

water in solvent power and behaviour, would have been investigated only to a very limited extent had it not been for its remarkable power to dissolve certain metals. All the alkali and alkaline earth metals, except beryllium, dissolve to some extent in liquid ammonia to give stable blue solutions; certain of the lanthanide metals also do this. Although it is now known that other solvents, notably certain ethers and di-ethers, will dissolve alkali metals to give blue solutions, the solubility is much lower and the solutions are less useful. Solutions of alkali metals in some molten salts are also blue; these are discussed in Chapter VI. It is useful at this point to list some of the facts which are known about the alkali metal-ammonia solutions, taking sodium as an example:


(1) 23 g of sodium dissolve in 100 g of liquid ammonia at 0°C; the solubility rises slightly with a fall in temperature.

(2) The blue solution is stable if kept in the dark but in the presence of light (specifically of wavelength 2150-2500 Â) or of certain catalysts (e.g. iron, iron salts or platinum) decomposition with the loss of colour occurs

Na + NH3 ->· NaNH2 + £H2

and sodium amide, being almost insoluble, is precipitated; similar behaviour is observed with lithium but the other alkali metal amides are soluble.

(3) When the metal dissolves there is a great increase in volume which may be as much as 70 cm3/mole of metal for dilute solutions.

(4) Evaporation leaves a coppery mass probably consisting of metal and absorbed solvent; the latter can be pumped away slowly to leave the pure metal again. In the case of the alkaline earth metals the final product is not the metal but its ammine, M.6NH3; the co-ordinated ammonia can be removed by heating but not without some formation of the amide.

(5) The solution shows electrical conductivity (Fig. 2). In the dilute solution this behaviour is characteristic of a salt, the equivalent conductivity reaching a limiting value at infinite dilution. In concentrated solution the conductivity rises to a very high value being of the same order as that of a metal. For example, the equivalent conductivity of sodium at — 33°C in a saturated solution is IO6 ohm"1 cm2 equiv-1 whilst that of liquid mercury is 0-15 x 10e ohm-1 cm2 equiv-1.

(6) The solution is paramagnetic but the static field molar susceptibility decreases with increasing concentration of the dissolved metal. The paramagnetic resonance absorption line is notably sharp having a width of only about 0-05 gauss between points of maximum slope; this technique also shows that at a metal concentration of 0·5 molar the strength of the resonance signal is only about 3 per cent ofthat expected if all the dissolved sodium dissociated to give sodium ions and unpaired electrons.

LIQUID AMMONIA 17

(7) The blue colour is due to a broad and intense absorption centred in the infrared at about 7000 cm-1 but with a tail extending into the visible part of the spectrum. In dilute solutions this absorption is identical for all the alkali and alkaline earth metals. For more concentrated solutions (and for solutions of metals in amines or amine-ammonia mixtures) a second absorption band may appear at 15,000 cm"1 in the visible region of the spectrum.

2000 l·

t ì 1000

3 er tu

Dilution, log10V

FIG. 2. Equivalent conductivity-dilution curve for sodium dissolved in liquid ammonia (—33°).

(8) The nuclear magnetic resonance spectra of sodium solutions in liquid ammonia show very large shifts to low fields for the 23Na and 14N nuclei, whereas the 1H signal remains completely unaltered from the standard. The large shifts are thought to be due to a paramagnetic effect — the presence of unpaired electrons can

c


influence the magnetic field at the nucleus being studied and cause an apparent increase in the field at that nucleus. It follows from the equality hv = gßH that, at the fixed frequency of the spectrometer, the resonance observed for the nucleus in question will occur at a lower field than is usual.

Of the theories put forward to explain these observations, one of the earliest, that the metal-ammonia "solutions" are really colloidal suspensions, now has little to support it. Pure, freshly prepared solutions do not display Tyndall scattering of light; moreover, the visible band in the absorption spectrum, which is observed in all the solutions, is unlikely to arise from scattering by colloidal particles, for if this were the case, its position and intensity would be expected to vary greatly from metal to metal. In considering any other theory, the notable similarity of properties of all the metals in both ammonia and other solvents must be an important consideration.

It is also important to notice that solubility of a metal in liquid ammonia seems to be associated with a large negative electrode potential in aqueous solution, i.e. the metals are good reducing agents, and the process:

M -> Mn+ + ne-metal solvated

ion

is energetically favourable. This in turn implies a low ionization energy and a low sublimation energy for the metal as well as a high solvation energy for the cation. Since, in liquid ammonia, the latter quantity will be of comparable magnitude to the value in water, then the process M -> Mn+ + ne~ will again be favoured.

In the more recent theories of alkali metal solutions, removal of an electron from the metal atomic orbitals is envisaged but there is some divergence of opinion as to the exact environment of this exiled electron. Two main postulates have been brought forward to describe the known phenomena. In the "expanded metal" theory the electron is supposed to move in an expanded orbital defined by the (innermost) layer of ammonia molecules

LIQUID AMMONIA 19

co-ordinated around the cation. The other theory considers the solution to contain solvated cations and electrons held in solvated cavities — in effect, solvated electrons — the latter behaving like the anions of a dissolved electrolyte.

Since both theories are not capable of explaining all the phenomena, a composite theory embodying the two has been described and fits, at least qualitatively, all the known facts. In concentrated solutions the high conductivity and lack of para-magnetism are explained by assuming that the expanded metal monomers, M(NH3)6, pair their outer, valence electrons and form dimers

o · o-oo Monomer Monomer Dimer (diamagnetic)

The close proximity of the dimers in concentrated solutions allows metallic conduction of electricity by means of electron tunnelling between neighbours.

As the concentration falls (to about 0Ό5 molar) the dimers are too far apart, on the average, for effective electron tunnelling to take place and so the conductivity decreases. Since there is still no paramagnetism exhibited by the solutions there can be little unpairing of electron spins by either of the two processes

Dimer -> 2 Monomers (1)

Dimer -> Monomer -> M^iv. + eloiv. (2) At lower concentrations (below about 0Ό5 molar) the solutions

begin to show paramagnetism and high conductivity which cannot be explained on the expanded metal theory, eqn. (1); however, the process shown in eqn. (2) can explain neatly both the above facts. These very dilute solutions absorb only in the 7000 cm-1

region of the spectrum and it has been postulated therefore that the absorption band at 15,000 cm-1 shown by the concentrated


solutions is due to electron transitions associated with the dimer species. The amount of expanded metal monomer, M(NH3)6, present at any one concentration has been calculated to be very small but is still sufficient to account for the large shifts noted in the 23Na nuclear magnetic resonance studies. Calculations have shown that the volume increase noted in the formation of dilute metal solutions is well accounted for in terms of the electron-cavity model. However, a volume expansion is also found even in the concentrated solutions but this can be explained since an entity like M(NH3)6 occupies a greater volume than the corresponding ion M(NH3)6+

M -> M(NH3)6+ + e~olv dilute solution

ï ' M(NH3)6 concentrated solution

Having obtained fairly satisfactory explanations for the observed phenomena, one of our main concerns in the next chapter will be with the chemical properties of these alkali metal-ammonia solutions.

Further Reading HILDEBRAND, J. H., Ammonia as a solvent, / . Chem. Educ. 25, 74 (1948). FRANKLIN, E. C , The Nitrogen System of Compounds, Reinhold, New York,

1935. SYMONS, M. R. C , The nature of metal solutions, Quart. Rev. Chem. Soc. 13,

99 (1959). JOLLY, W. L., Metal-ammonia solutions, Prog, lnorg. Chem. 1, 235 (1959).

CHAPTER III

REACTIONS IN LIQUID AMMONIA

1. Reactions of Metal-Ammonia Solutions Since these solutions consist essentially of metal cations and

electrons which are to some extent "free", they all behave as highly versatile reducing agents in both inorganic and organic chemistry. The particular metal cation present does not usually affect the course of the reaction, except when it forms an insoluble salt with the reduction product.

(a) Reduction of Elements In general, elements which are reduced are those of Groups IV,

V and VI of the Periodic Table. Reduction of some transition metal salts to the metal (see Section 1 (b)) may be followed by further reduction of the finely divided metal, but the usual effect here is to catalyse the decomposition of the metal-ammonia solution to give the amide and hydrogen.

Some examples of reactions with the Group IV, V and VI elements are given in Table 2; reactions with the other elements are similar. The notable feature is the formation of polyanions, for example Pb9

4~ and S72_. Thus with sulphur, the normal

sulphide Na2S is first formed as a white precipitate, but slowly dissolves (if excess sulphur is present) to a red solution of the polysulphides N a ^ . Lead behaves similarly, but dissolves very slowly; the polyplumbides are more conveniently prepared by reaction of excess alkali metal with a solution of lead iodide in liquid ammonia; conductometric and potentiometric studies (p. 33) of these reactions indicated the various stages of polyanion formation.

21


The interesting point here is that some of these compounds have compositions similar to those of intermetallic compounds which separate as phases of definite composition from binary mixtures of fused metals. Compounds such as Na4Pb9 seem to be intermediate between true intermetallic compounds of sodium and lead (for example Na5Pb2) and the polyanionic salts such as polysulphides and the polyhalides of Group VII. The colours which compounds like Na4Pb9 give in liquid ammonia solution may be due to aggregates of the polyanions; the latter react with added cations, e.g.

Pb7*- + 2Pb2+ -> 9Pb|

and electrolysis of a polyplumbide deposits lead at the anode.

TABLE 2

Group

IV

V

VI

Element

Sn Pb

P As Bi

S Se

Metal

Na Na

K Na Na

Na Na Na

Products

Na4Sn9, red solution Na4Pb7, green solution Na4Pb9, green solution

KP5.3NH3 Na3Asx(A:= 1,3, 5,7) NaaBi* (JC = 1, 3, 5)

Na2Oa; Na02 Na2Sx (x = 1 to 7) NaaSe* (JC = 1 to 6)

Although alkali metals and ozone probably give ozonides in liquid ammonia, these compounds are too unstable for isolation and study; however ammonium ozonide has been prepared by the low temperature ozonization of ammonia, and the reaction of lithium hydroxide with ozonized oxygen in the presence of ammonia gives the compound Li(NH3)403, which is soluble in liquid ammonia without decomposition.

REACTIONS IN LIQUID AMMONIA 23

(b) Reduction of Salts Here, the most common course of the reaction is to produce the

metal, where the salt is a simple compound and not a complex; thus silver and nickel salts are reduced to the respective metals. However, the production of finely divided metal by the reduction catalyses the alkali metal-ammonia reaction to give an alkali metal amide, and hence with a nickel salt the reaction sequence is, using potassium metal

NiBr2 + 2K -> Ni + 2KBr

2K + 2NH3 — U 2KNH2 + H2

2KNH2 + NiBr2 + 2NH3 -> Ni(NH2)2.2NH3 + 2KBr

Examples of anion reduction are known. Iodates are reduced to iodides and permanganates to manganates, but here also side reactions may occur as, for example, in the permanganate reduction, where the formation of traces of manganese dioxide will catalyse the alkali metal amide production.

The reactions of complex metal salts with alkali metal-liquid ammonia solutions are important since an unusually low oxidation state of the metal may be produced in the reduction. Hence reduction of a complex nickel(II) cyanide with potassium gives, with excess of the cyanide, the bright red tetracyanonickelate(I)

Ni(CN)42- + er -> [Ni(CN)J3-

and with excess alkali metal, the yellow tetracyanonickelate(O) is precipitated

Ni(CN)42- + 2e- -> [NiCCISOJ4-

This anion is isoelectronic with nickel(O) tetracarbonyl, Ni(CO)4; in water, hydrogen is evolved and the tetracyanonickelate(I) ion obtained. Similar reactions to this yield the complex cyano-ions

[Co^CN^]4-, [Mn°(CN)6]6- and [Cr°(CN) J · -


In certain cases lower oxidation states than zero can be produced; for example the metal-metal bond in dimanganese(O) deca-carbonyl, Mn2(CO)10, is readily cleaved by an alkali metal to yield the corresponding pentacarbonylmanganate(—1):

(CO)5Mn—Mn(CO)5 + 2Na -> 2NaMn(CO)5

(c) Reduction of Covalent Compounds A wide range of reactions are possible here; they are considered

as follows :

(/) Charge transfer without bond fission. Both nitric oxide and carbon monoxide react with alkali metals in liquid ammonia, giving sparingly soluble or insoluble compounds of empirical formulae MJNO and ΜΧ(Χ), i.e. apparently containing the anions NO~ and CO~. Recent spectroscopic studies of these substances lead to the conclusion that MxNO is to be formulated as M^NgOa, containing the ion N202

2~, which is an isomer — the eis isomer — of the hyponitrite ion N202

2~, i.e.

Ö /

N=N /

O

eis hyponitrite trans

The position with regard to "MICO" is less clear, though it seems that ultimately this is converted to the salt of hexahydroxy-benzene, e.g. K6C606; there is spectroscopic and X-ray diffraction evidence for the intermediate formation of acetylene diolates MI

2C202 with the structure M2[0—C^C—O], and the caesium salt Cs2C202 yields glyoxal, (CHO)2, on acid hydrolysis, which supports this formulation. Other compounds intermediate between MI

aC2Oa and MI6C606 may also be formed.

N=N / \

Ό O"


(iï) Fission of hydrogen atoms. Covalent hydrides of both metals and non-metals react with metal-ammonia solutions by replacement of one or more hydrogen atoms :

PH3 + K -> KPH2 + iH2

SnH4 + Na -> NaSnH3 + iH2

SnH4 + 2Na -> Na2SnH2 + H2

With an organo-metal hydride, for example Me3GeH, the hydrogen atom is removed in preference to the organic group and MIGeMe3 is formed. The products of these hydrogen abstraction reactions are of importance synthetically (see p. 30).

Special considerations apply when the hydride has acceptor properties, since the reacting entity will not be the simple hydride alone. Diborane and ammonia, for example, form the solid compound "B2H6.2NH3"; the reaction of this with sodium in liquid ammonia assisted in establishing its structure as a boro-hydride, i.e. as

[H3N.BH2.NH3]BH4

which reacted thus: [H3N.BH2.NH3]BH4 + Na ->

NaBH4 + BH2NH2 + £H2 + NH3

Tetramethyldiborane, Me4B2H2, is split into ammonia-dimethyl-borane (H3NBMe2H) and the salt Na2HBMe2 when treated with sodium in liquid ammonia.

The simplest hydride-ammonia adduct is, of course, the ammonium ion and as expected ammonium salts (which behave as acids in liquid ammonia) react rapidly with metals to give hydrogen and the corresponding metal salt, for example

NH4C1 + Na -> NH3 + JH2 + NaCl

(Hi) Fission of alky I groups. The reaction of tetramethyltin or tetramethyllead with metal-ammonia solutions occurs thus: (M = Sn or Pb),

MMe4 + er -> MMe3~ + CH3


When M = Sn the reaction stops at this point, the methyl radicals appearing as ethane. When M = Pb and potassium is the reducing species, solvolysis of the PbMe3~ ion takes place giving first ions such as PbMeaNH2~, and finally lead imide, PbNH, and amide ions; with lithium or sodium, the insolubility of their amides promotes the reaction:

PbMe2NH2~ + Na+ -> PbMe2 + NaNH2

which leads to the production of dimethyllead. Some reduction and solvolysis of the methyl radicals formed in the initial stages of the reaction takes place to yield methane

NH3 + CH3 + er -> NH2~ + CH4

A similar type of reaction occurs with tetramethylammonium salts, but the ammonolysis is rather different.

NMe4+ + er -> NMe3 + CH3

CH3 + er -> CH3-CH3- + NH3 -> CH4 + NH2-

With a tetraethylammonium salt, ethylene is produced by the reaction

NEt4+ + NH2- -> NEt3 + C2H4 + NH3

As with hydrides, alkyl compounds which are acceptor molecules will form adducts with the solvent; hence alkyl groups are not removed from a trialkylboron by alkali metals in liquid ammonia, instead the reaction

Me3BNH3 + er -> Me3BNH2" + £H2

occurs to give stable solid salts, for example KtMegBNHJ.

(iv) Fission of either atoms or groups. Although a great variety of reactions between metal-ammonia solutions and organic compounds containing covalently attached atoms or groups have been reported, they are outside the scope of this book. Few reactions of purely inorganic co valent compounds in

REACTIONS IN LIQUID AMMONIA 27 this category are known, since most of these compounds undergo ammonolysis to some extent.

Boron trifluoride does not ammonolyse in pure liquid ammonia, but in the presence of potassium, fission of a fluoride ion results in ammonolysis

H3NBF3 + K -* H2NBF2 + KF + iH2

More complicated products are obtained if sodium or lithium are used; with lithium, for example, the product is

(NH2)2BNHB=NH Some organo-metallic halides undergo reaction with metal-

ammonia solutions without extensive ammonolysis. One example of a general type of reaction is that of trialkyltin(I V) halides

Me3SnBr + 2Na -> Me3SnNa + NaBr

The product can be used to prepare hexamethyldistannane, Me6Sn2

Me3SnNa + Me3SnBr -> Me3SnSnMe3 + NaBr or mixed alkyls of tin, Me3SnR

Me3SnNa + RX -> Me3SnR + NaX; X = halogen This kind of reaction may be used similarly, with some triaryl (or dialkyl or diaryl)-halogen compounds of tin and other Group IV metals such as germanium and lead, see page 30.

Dimethylbromostibine, Me2SbBr, with sodium in ammonia yields only tetramethyldistibine

2Me2SbBr + 2Na -> Me2SbSbMe2 + 2NaBr Unlike the antimony-antimony bond in the tetramethyldistibine

formed above, certain metal-metal bonds are cleaved by the alkali metals. One such cleavage which has been mentioned previously is the formation of sodium pentacarbonylmanganate(— 1) from dimanganese decacarbonyl and sodium. This type of behaviour is particularly common among the Group IV elements and under the correct conditions all the elements of the group will partake


in the cleavage, even carbon : C2R6 + 2M -> 2MCR3; R = C6H5

Probably the most studied is the reaction of the silicon-silicon bond with alkali metals. Hexa-aryldisilane, for example, will undergo a smooth reaction to give the metal triarylsilyl in high yield

Si2R6 + 2M -> 2MSiR3

Such products are very useful reagents in synthetic chemistry since other novel metal-metal bonds can be formed from them, e.g.

HgCl2 + 2MSiR3 -> Hg(SiR3)2 + 2MC1 Cleavage accounts for the fact that no distannanes (or di-

plumbanes) are formed in the reaction of excess alkali metal with the trialkyl(aryl)tin (or lead) halides

R3SnCl + 2Na -> R3Sn - SnR3 U 2R3SnNa

2. Oxidation Reactions These are rare in liquid ammonia. In water, the "electron

transfer" required in an oxidation-reduction reaction is often achieved by a mechanism of atom or group transfer, e.g. by transfer of a hydroxyl radical OH; the appearance of hydrogen peroxide as a by-product of oxidation reactions is an indication of this. The corresponding formation of NH2 radicals in liquid ammonia appears to be difficult, and only recently has hydrazine, H2NNH2, been prepared by electrolysis of liquid ammonia solutions (see below).

Potassium amide is slowly oxidized by air in liquid ammonia, and potassium nitrite and potassium hydroxide are formed. Oxidation may also be achieved with potassium nitrate

3NH2- + 3N03- -* 30H- + N2 + NH3 + 3N02~

In both reactions, the formation of hydroxyl ions is to be noted as requiring the removal of hydrogen from the amide and transfer

REACTIONS IN LIQUID AMMONIA 29 of it to the oxygen of the oxidant. One apparently genuine electron transfer oxidation is known, viz.

[Sn(NH2)6;r + I2 -> [Sn(NH2)6]2- + 2I~ Sn(II) Sn(IV)

A curious reaction is that between zirconium(III) bromide and potassium amide, in which the zirconium is oxidized to an average oxidation state of 3-5 (appearing as a polymeric imido-amido zirconium complex) and the potassium ions are reduced to give metallic potassium.

Ammonium nitrate in liquid ammonia is not an oxidizing agent like aqueous nitric acid, and oxidizing agents such as permanganates or chromâtes have very limited power in ammonia.

The production of hydrazine from liquid ammonia is not achieved by conventional electrolysis with metal-liquid ammonia electrodes, but by glow-discharge electrolysis where a layer of vapour separates anode and liquid; in this vapour it is considered that positive gaseous ions, for example NH3+ or NH2+, are produced and that these move into the liquid and there produce NH2 radicals, e.g. by the reaction

NH3+ + NH3 -> NH4+ + NH2

hydrazine being produced by combination of these radicals.

3. Metathetic Reactions These reactions are numerous and many are extremely useful in

synthesis. Broadly, they depend either (a) on the combination of solvated protons (ammonium ions) with amide ions to give solvent molecules, or (b) on the association of a cation and an anion which yield a salt insoluble in ammonia, i.e.

(a) NH4X + MNH2 -> MX + 2NH3

(b) AB + CD -> AC + BD| Reactions of type (a) are not of general use, since it is often

easier to prepare salts MX than ammonium salts NH4X; however, the reaction


NH4N3 + KNH2 -> KN3 + 2NH3

has been used to obtain a solution of potassium azide in ammonia. Many reactions of type (b) were investigated by Franklin; thus,

since silver chloride is soluble and alkali metal chlorides are insoluble in ammonia, reactions of the type

AgCl + KNH2 -> AgNH2 + KClj occur readily. More usefully, reactions with salts such as KPH2 (p. 25) can be used to obtain substituted hydrides, e.g.

KPH2 + EtCl -* EtPH2 + KClj A similar type of reaction yields dihydrogenphosphides of other metals [Co(NH3)6](N03)3 + 3KPH2 -> Co(PH2)3 + 3KN03 + 6NH3

Since some alkali metal borohydrides are soluble in ammonia, they can be used to prepare other borohydrides by metathesis

NH4C1 + KBH4 -> NH4BH4 + KClj The borohydrides [Cr(NH3)6] [BH4]3 and [Mg(NH3)6] [BH4]2 have been prepared similarly. When trimethyllead chloride and potassium borohydride are allowed to react at low temperature in ammonia the expected metathesis

Me3PbCl + KBH4 -> Me3PbBH4 + KC1 j occurs; but on warming, the ammonia solvent abstracts a BH3 group by electron-pair-donation, and the product is the hydride Me3PbH:

Me3PbBH4 + NH3 -> Me3PbH + H3NBH3

The hydride also unites reversibly with ammonia, thus Me3PbH + NH3 ^ NH4PbMe3

Metathetic reactions have proved useful in the synthesis of organo-metallic compounds containing two (or more) metals, e.g. the reaction,

(C6H5)3GeNa + BrSnMe3 -> (C6H5)3GeSnMe3 + NaBr\


4. Titrations in Liquid Ammonia From the postulated self-dissociation of liquid ammonia

2NH3 ^ NH4+ + NH2-Frankhn defined acids as substances which dissolved to give cations characteristic of the solvent self-ionization either by direct dissociation or by reaction with the solvent

NH4C1 -> NH4+ + CIUCI + NH3 -> NH4+ + Cl-

Bases, on the other hand, give anions characteristic of the solvent self-ionization

KNH2 - Ä * K+ + NH2-cf.

KOH —°-> K+ + OH~ Franklin was then able to demonstrate acid-base neutralization

reactions such as NH4C1 + KNH2 -> KC1 + 2NH3

Calorimetrie studies have since shown that the heat of neutralization for

NH4+ + NH2- -> 2NH3

is virtually constant at about 26 kcal in many ammonium salt-alkali metal amide reactions. These experiments suggest that acid-base titrations are also possible in liquid ammonia but no colour indicators are known which give satisfactory, reversible colour changes at the end point. Ammonia has a greater affinity for protons than water so that all the usual indicators which behave as weak acids in water are fully ionized in liquid ammonia due to the levelling action of the solvent (see p. 44); similarly, indicators which are weak bases will be completely undissociated. Consequently, indicators will exhibit the same colour in liquid ammonia as they do in basic aqueous solutions and will be unaffected by the addition of either ammonium salts or amides.


It is possible to detect the end point of an acid-base titration by carrying out the reaction in a conductivity cell and noting the change in conductivity of the solution as base (or acid) is added, the solution being stirred vigorously after each addition. The end point of the titration is shown up as a distinct minimum when the conductivity is plotted against the amount of added base (or acid). Normally conductometric titrations in liquid ammonia are carried out at about —34° to keep the pressure below atmospheric, and stringent precautions are taken to exclude the entry of water and carbon dioxide into the apparatus.

trt

o +-Ü to

c

% a> Si ~jö c V 4-o Q-

,

(

f (a)

f o - o ^ /

X

[(b) X

Ì X x /

-**x-^

1 KO 2-0 3Ό

Ratio moles of KNH2/acid

FIG. 3. Potentiometric titration of (a) an ammonium salt, and (b) a guanidinium salt with potassium amide in liquid ammonia.

The same titration can also be followed potentiometrically by placing a platinum reference electrode in the amide solution and noting the change in potential, as the ammonium salt is added, between this electrode and a reference electrode composed of a mercury pool in contact with a saturated ammonia solution of mercuric chloride.

Many compounds which behave as acids in ammonia can be titrated in the same way, for example the guanidinium ion,


H2NC(NH)NH3+, when titrated with potassium amide has three breaks in the potentiometric titration curve (Fig. 3) showing that the ion behaves as a tribasic acid in liquid ammonia.

H2NC(NH)NH3 NH2-H2NC(NH)NH2 + NH2-H2NC(NH)NH- + NH2-

H2NC(NH)NH2 + NH3

H2NC(NH)NH- + NH3

[HNCCNHiNH]2- + NH3

1-4 £ o >

. 1-3 <Λ 0>

O

+-O 2 1-2 ÜJ

c 0) Φ * ! l 4- l'I Φ

-O

"B '■£

Φ 1-0

a.

_ _ 1 — 1 I 1

1

— rW^"^ 1

- 1

1

1 Λ

(\ r\ £ n - Q ^^

1 I : 1 I 1 — OS ΙΌ 1-5 2Ό

Ratio potassium /Cd 2 + 2-5

FIG. 4. Potentiometric titration of a cadmium salt with potassium in liquid ammonia.

However, the volatility of the ammonia is a great drawback to the full exploitation of the technique and such titrations are normally carried out more easily at room temperature in amine solutions (see p. 36).

Potentiometric titration is also a useful method by which to study reduction reactions involving alkali metal-ammonia solutions and has been used to investigate the possibility of transitory lower oxidation states during the reduction of metallic salts. Zinc, cadmium and mercury(II) salts are reduced directly to the metal and no M(I) states can be detected. A typical titration


curve is shown for cadmium in Fig. 4. The first break in the curve corresponds to the end point of the reaction

Cd2+ + 2K -* Cdj + 2K+

whilst the second break signifies the formation of an intermetallic compound KCd3.

The formation of the intermetallic Na/Pb compounds described on page 22 was studied by the potentiometric titration of a

2-4

£ 2·2|

I 2-Of-

a> o 1-eU

• 4-<■> <D

LJ \-6\-c % ''4

4-a> * 1-2

£ \'0\-

K3Co(CN)6

_L· 05 1-0 1-5 2-0

Ratio potassium/cyanometallate

2-5

FIG. 5. Potentiometric titration curves obtained when potassium is added gradually to nickel and cobalt cyanometallates in liquid ammonia.

solution containing 15-5 mg of sodium in 75 ml of liquid ammonia with a solution of 170 mg lead iodide in 69-7 ml of liquid ammonia; at — 50°C the reaction follows the equation

(4 + 2x)Na + xPbI2 -► NaJPb* + 2xNaI

and inflexions in the titration curve at 49-9 and 52-3 ml of titrant correspond to formation of Na4Pb7 and Na4Pb9 respectively.


The reduction of the tetracyanonickelate(II) and hexacyano-cobaltate(III) ions with potassium-ammonia solutions at — 33°C results in the formation of Ni(I) and Co(I) cyanometallates as shown both by a break in the potentiometric titration curves (Fig. 5), and the fact that the two air-sensitive compounds can be isolated for analysis

K2Ni(CN)4 + K -> K3Ni(CN)4

K3Co(CN)6 + 2K -> K3Co(CN)4 + 2KCN

«*U-£ •4-

o > 0 r

<0 -S 02 o v 0-4 Ld

« 0-6 0) $ 2 0-8 "Ô t -1-0 jD "o a.

b - N > _ _

^ ^ ^ - < ^ ^ ^ χ ^ ^ Α y s

(

_

1 1

>

p

V

i I l 1 0-25 0-5 075

Ratio moles of K 2 N Ì ( C N ) 4 / K

1-0

FIG. 6. Potentiometric titration curve obtained when potassium tetracyanonickelate(II) is added gradually to a solution of potassium in

liquid ammonia.

The further reduction of these complexes to give Ni(0) and Co(0) tetracyanometallates is too slow to be followed potentio-metrically. However, when the titration of tetracyanonickelate(II) with potassium is carried out in the reverse order, i.e. by adding the nickel(II) complex in small amounts to the potassium solution the titration curve shows a break when only 0-5 mole of complex


has been added due to the immediate formation of tetracyano-nickelate(O) (see Fig. 6).

5. Anhydrous Amines as Solvents The extraordinary usefulness of liquid ammonia as an ionizing

solvent stimulated investigations into the possible use of aliphatic amines as solvents. However, it soon became apparent that the amines were greatly inferior to ammonia in solvent power and interest in them dwindled rapidly (see however Section 6).

The ability of ammonia to dissolve the alkali metals giving deep blue solutions is only feebly copied by methylamine and ethyl-amine which both dissolve lithium to a small extent; the secondary and tertiary amines do not dissolve any metals. Sodium and potassium have also been shown to dissolve in methylamine which contains a trace (ca. 0·1 per cent) of ammonia as impurity.

Many inorganic compounds which are soluble in aliphatic amines are covalent (e.g. S, Br2 and I2) or have large ions and therefore presumably low lattice energies (e.g. KSCN, Hg(CN)2 and Hgl2); lithium chloride is a notable exception to these two generalizations. The solvent power of an amine for a given salt often decreases markedly with increasing alkyl chain length as is demonstrated by potassium iodide which is very soluble in liquid ammonia, less soluble in methylamine and almost totally insoluble in ethylamine.

The conductivity of salts dissolved in primary amines is often less than for the same salts dissolved in ammonia but still indicates a slight ionization. Although the secondary amines are much poorer solvents, solutions of salts in both primary and secondary amines show roughly the same conductivity whereas the tertiary amines are greatly inferior in both solvent power and ionizing ability. No salt has yet been found which has any appreciable solubility and conductivity in, for example, trimethylamine. However, ethylenediamine has been studied quite extensively and will now be briefly discussed.


Ethylenediamine The commercial product, a 95 per cent mixture of ethylene

diamine and water, is best dehydrated by allowing it to stand over sodium hydroxide and barium oxide for several days and then treating it with metallic sodium for a further day; a fractional distillation from freshly activated alumina finally affords the pure solvent. Ethylenediamine is considerably more basic than ammonia and readily absorbs carbon dioxide and water from the atmosphere. All ethylenediamine solutions must therefore be handled under pure, dry nitrogen since even a brief exposure of a solution to the atmosphere is sufficient to cause a significant decrease in the freezing point and a corresponding increase in conductivity.

TABLE 3

Physical properties of anhydrous ethylenediamine

Melting point, °C Boiling point, °C Viscosity, cP (25°C) Conductivity, ohm"1 cm"1 (25°C) Dielectric constant (25°C)

11-3 117-3

1-7 9 x 10-8

12-9

Many inorganic salts are soluble in ethylenediamine. Usually chlorides are insoluble whilst thiocyanates, bromides, iodides, nitrates and perchlorates are among the most soluble salts. Ethylenediamine readily forms solvates which in several cases are made by simply dissolving a salt hydrate in pure ethylenediamine and recrystallizing the solution

LiI.3H20 + ien -> Lil3en + 3H20 en = H2NC2H4NH2

Other examples of solvates are Nal.3e«, BaCl2.4e«, CaCl2.6e«, and SrCl2.6e« of which the last three are virtually insoluble in ethylenediamine.


Potentiometric, conductometric and cryoscopic measurements have shown that all salts act as weak electrolytes in ethylene-diamine and are only slightly dissociated, a typical example being silver nitrate which has a dissociation constant of about 6 X 10-4. This behaviour is, of course, to be expected in a solvent of low dielectric constant.

Sodium, potassium and lithium are soluble to give rather unstable blue solutions. Lithium, besides being the most soluble, is the only metal which gives paramagnetic solutions at low concentrations owing to the presence of solvated electrons. So far these metal-ethylenediamine solutions have not been used at all extensively for preparative purposes even though the more useful liquid range of ethylenediamine makes experimentation somewhat easier than when liquid ammonia is used.

6. Titrations in Anhydrous Basic Solvents The anhydrous amines are finding increasing use as solvents for

acid-base titrations since being strongly basic they are able to "level" (p. 44) the acid strength of weak acids such as carboxylic acids and phenols. The standard bases used in these titrations are normally alkali metal alkoxides or quaternary ammonium hydroxides dissolved in benzene-alcohol mixtures or, less often, in the amine used for dissolution of the acid. Besides extending the list of titratable substances to those outside the range of the more normal aqueous volumetric analysis, the use of "organic" solvents such as the amines also adds to this list many organic acids which are insoluble in water.

Titrations in non-aqueous solvents are simple to carry out and differ only slightly from their counterpart in aqueous solution. The main difference is that water and carbon dioxide must be excluded from the titration vessel otherwise the end point of the titrations is not sharp, especially when very weak acids are being used.

In butylamine (chosen mainly for its useful liquid range, m.p. —50-5°C, b.p. 77-8°C), carboxylic acids and phenols can be

REACTIONS IN LIQUID AMMONIA 39 determined to the visual end point (yellow to blue) using thymol blue as the indicator and sodium methoxide in benzene-methanol as the standard base. The end point can also be detected potentio-metrically using a glass-antimony electrode system.

Antimony indicator electrode

ntimony reference electrode

Potentiometer —J

FIG. 7. Typical apparatus used for a potentiometric titration in a non-aqueous solvent; arrangements for obtaining the necessary dry and

inert atmosphere in the vessel are not shown.

Sodium aminoethoxide, NaOC2H4NH2, and sodium methoxide are both soluble in ethylenediamine so that the titration of acids can be carried out wholly in this solvent. However, the glass electrode does not function in ethylenediamine solutions containing sodium ions so that the titrations are usually followed with an antimony-antimony electrode system (see Fig. 7). Phenol, alkylated phenols, carboxylic acids, o- and m-hydroxybenzoic acids and amino-acids have all been titrated successfully in this solvent; the buiTering action of the amino group in amino-acids

4 0 NON-AQUEOUS SOLVENTS IN INORGANIC CHEMISTRY

is so effectively eliminated in a basic solvent like ethylenediamine that these acids can be titrated as readily as any other carboxylic acid.

Mixtures of a carboxylic acid and a phenol show two breaks in the potentiometric titration curve, the first corresponding to the neutralization of the carboxylic acid and the second, slightly less pronounced break to the neutralization of the phenol.

Volume of base Volume of base Volume of base

(a) (b) (c)

FIG. 8. Potentiometric titration of acids in ethylenediamine: (a) a carboxylic acid; (b) mixture of a carboxylic acid and a phenol, or a hydroxybenzoic acid such as salicylic acid; (c) a dihydroxyphenol.

A simple titration of phenol against either sodium or potassium methoxide using a colour indicator can be carried out in the following manner. Two drops of the indicator o-nitroaniline (0-15 g in 100 ml of dry benzene) are added to about 25 ml of anhydrous ethylenediamine and titrated to the red colour with sodium methoxide dissolved in ethylenediamine in order to neutralize any carbon dioxide present in the solvent. Phenol is then dissolved in the neutral solvent and titrated with methoxide until the indicator colour changes from clear yellow to orange-red. The methoxide solution can be standardized against pure benzoic acid in a similar manner.

s

\


Further Reading WATT, G. W., Reactions of inorganic substances with solutions of metals in

liquid ammonia, Chem. Rev. 46,289 (1950); Reactions in liquid ammonia, / . Chem. Educ. 34, 538 (1957).

FOWLES, G. W. A. and NICHOLLS, D., Inorganic reactions in liquid ammonia, Quart. Rev. Chem. Soc. 16, 19 (1962).

BECKETT, A. H. and TINLEY, E. H., Titrations in Non-aqueous Solvents (3rd edition), British Drug Houses Ltd., 1960.

CHAPTER IV

PROTONIC SOLVENTS

THIS chapter is devoted to a study of the three most common protonic solvents, pure sulphuric acid, anhydrous hydrofluoric acid and glacial acetic acid. The title "acid" is, of course, derived from their role in the aqueous solvent system and is strictly incorrect unless one states in which medium the compounds behave as acids; e.g. acetic acid behaves as a weak acid in water but in pure sulphuric acid it is a base. With this in mind we will use the more correct title of anhydrous hydrogen fluoride for the second solvent but unfortunately the other two have no more suitable names.

Among these three solvents there is an interesting gradation in dielectric constant, that of sulphuric acid being the highest (ca. 120). The force of interaction (F) between two charges +e± and — e2 (e.g. two ions) is given by

e ra

where c is the dielectric constant of the medium in which the ions are situated and r is the distance between them. In sulphuric acid, therefore, there is the least ion-ion interaction and all electrolytes are found to be fully ionized.

The formation of solvates by inorganic salts often gives an indication of solubility of the salt in the particular solvent, since in the solvate certain of the ions making up the crystal are surrounded by co-ordinated solvent which has the effect of reducing the lattice energy of the crystal. The greater the dielectric constant of the co-ordinated solvent the more the lattice energy

42

PROTONIC SOLVENTS 43

will be reduced making the possibility of solution more favourable when the solvate is placed in the solvent. Looking at solvation in another way, it can perhaps be regarded as the formation of a very concentrated (often solid) solution of the salt — when further solvent is added, the only effect is one of dilution as the ions are separated by more and more solventf.

As will be seen, sulphuric acid and liquid hydrogen fluoride, though they are good solvents for inorganic compounds, are very reactive and extensive solvolysis of dissolved salts often occurs. Acetic acid has the lowest dielectric constant of the three and therefore salts on the whole are relatively insoluble. On the other hand, since a low dielectric constant indicates that the acetic acid molecules are only slightly polar (to a first approximation the dielectric constant is some measure of the polarity of the solvent molecules), covalent, non-polar compounds show a high solubility in acetic acid following the familiar "law" of like dissolves like.

This "law" arises from the fact that solvent molecules which are highly polar tend to interact with each other (for example, hydrogen bonding is a direct outcome of this polarity). Therefore, a non-polar molecule cannot break up the molecular interaction of the solvent in order to enter into solution and one finds such molecules are normally insoluble. This is the case with sulphuric acid where only a very few covalent molecules are soluble. On the other hand, if the solvent-solvent interaction is weak, the non-polar solute is readily accepted into solution since its molecules are little different from those of the solvent.

Water is an impurity common to all acidic solvents; being comparatively basic it undergoes protonation quite readily

HaO + HX ^ H3O+ + X-and gives the solvents a spuriously high electrical conductivity.

t In references to "solvates" elsewhere in this book, the term will be used to describe any compound, formed by dissolution of a salt in a particular solvent, which contains "elements of solvent" held either by ion-dipole forces, covalent bonds, or as part of a complex ion, e.g. in KF.BrF8 (i.e. KBrF4, p. 88).


Obviously, for all accurate work, water must be rigorously excluded both from the solvent and from all the apparatus with which the solvent is likely to come in contact. The high affinity of protonic solvents for water is probably mainly due to the ease with which water, either as H20 or H30+, can enter into hydrogen bond formation with the solvent molecules.

Conductivity measurements on these protonic solvents indicate that self-ionization, involving proton transfer between solvent molecules, occurs according to the general equation

HA + HA ^ H2A+ + A-L _ H + _ i l_H+_t

acid 1 base 2 acid 2 base 1

where A~ = HS04-, F~ or CH3COO_, although in sulphuric acid the position is more complicated due to the simultaneous occurrence of other equilibria. Hence hydrogen sulphates, fluorides and acetates act as bases in the solvents in the same way that hydroxides are the bases of the aqueous system

H20 + H20 ^ H3O+ + OH-L H + _t L H U

acid base acid base

Levelling Action Acids in these solvents are proton donors; the stronger the acid

the more easily it donates a proton to the solvent. If the solvent (AH) has a high proton affinity (i.e. is basic) complete proton exchange will occur on the addition of a strong proton donor, HX

AH + HX -> AH2+ + X-base acid acid base

and all strong acids (proton donors) in this solvent will display the same acidity since they will all react completely with the solvent to produce the solvated proton, H2A+. In other words, the


strongest acid which can exist in a protonic solvent is that which is produced by the self-ionization of the solvent

2AH ^ H2A+ + A-acid base

cf. 2H20 ^ H3O+ + OH-

and all strong acids are "levelled" to the strength of this ion, H2A+. For example, the mineral acids hydrochloric, nitric and sulphuric, are all levelled to the acid strength of the oxonium ion, H3O4-, in aqueous solution. In liquid ammonia, these acids are again levelled but now to the acidity of the ammonium ion, NH4+.

When dissolved in a solvent of weaker proton affinity, a strong acid will have more difficulty in protonating the solvent molecules and the number of H2A+ ions produced will be proportional to the acid's strength; clearly in such a solvent, acids are "differentiated" according to their proton-donating ability. Sometimes the solvent may be so acidic (e.g. pure sulphuric acid) that only very few solutes are capable of protonating the solvent molecules.

A similar situation occurs with bases. Bases are completely protonated by an acidic solvent and are therefore levelled to the basicity of the solvent anion, a typical case being that of ammonia dissolved in glacial acetic acid

NH3 + CH3COOH -> NH4+ + CH3COO-

In a less acidic solvent like water, many bases are differentiated since their protonation by the solvent is not complete

NH3 + H 2 0 ^ NH4+ + OH-

However, very strong bases like the ethoxide ion, C2H50~, and the hydride ion, Hr, are completely protonated even by a weakly acid solvent like water in which they are levelled to the basicity of the hydroxyl ion

C2H50- + H 2 0 -^ C2H5OH + OH-


Such compounds are obviously best handled in a non-protonating (basic) solvent. Conversely, in a highly acidic medium like liquid hydrogen fluoride weak bases (e.g. Ch and Br~) are levelled to the common basicity of the fluoride ion by the protonation (solvolytic) reaction

Cl- + HF -> F- + HC1 The levelling action of solvents has, therefore, a profound

influence on the choice of a solvent for a particular reaction. We saw in the previous chapter that compounds which behave as weak acids in water can easily be titrated in amine solvents ; the amines, being more basic than water, are readily protonated by acids

RNH2 + HX -> RNH3+ + X-the strengths of which are therefore levelled to that of the RNH3

+

ion. On the other hand, since amines are poor proton donors, very strong bases such as the methoxide ion are quite stable in solution and can be used as standard bases in titrations

RNH3+ + [OR']" -> RNH2 + ROH

Similarly, in an acid solvent many bases are levelled by protonation to the basicity of the solvent anion, including those classed as "weak" in the aqueous solvent system. It is therefore possible to titrate these weak bases in an acid solvent, acetic acid being the one most frequently employed mainly because of its ready availability and ease of manipulation.

Sometimes it is desirable to study the relative acid strengths of a series of compounds, in which case a solvent is required which has only weak proton-acceptor properties and is able to differentiate between the various proton donors. Studies of this nature, carried out in glacial acetic acid, have shown the strengths of several acids to be in the order

HC104 > HBr > H2S04 > HCl > HN03

Similar measurements can, of course, be made on bases dissolved in basic solvents.


Sulphuric Acid Large quantities of pure sulphuric acid are manufactured

annually on an industrial scale. Sulphur dioxide is oxidized to sulphur trioxide, using oxygen and a catalyst, and the trioxide dissolved in either water to give sulphuric acid or in concentrated sulphuric acid to give oleum. Exactly 100 per cent sulphuric acid is obtained by diluting oleum with aqueous sulphuric acid until the melting point is 10-37°C.

TABLE 4

Physical properties of sulphuric acid

Melting point, °C Boiling point, °C Dielectric constant Specific conductivity, ohm-1 cm-1

Viscosity, cP

10-37 290-317

/100 (25°) \120(10°)

1-04 x 10~2(25°) 24-5 (25°)

As is apparent from Table 4, sulphuric acid has a useful liquid range. The high dielectric constant (cf. water, 80) leads one to expect many inorganic salts to be soluble and ionic interaction to be very low since most electrolytes will be fully ionized. In actual fact, most inorganic salts other than the hydrogen sulphates (containing the ion HS04~) suffer extensive solvolysis when dissolved in pure sulphuric acid.

The main drawback to the usefulness of sulphuric acid as a solvent is its high viscosity (about 25 times as high as water at the same temperature) which makes dissolution of solutes a slow process unless the temperature is raised when the probability of solvolysis, already high, increases still further. In such viscous solutions metathetic reactions may not always be instantaneous and, furthermore, any solvent adhering to the precipitated product is very difficult to remove during purification.


The molepular association in sulphuric acid which is indicated by the high viscosity is thought to be due to hydrogen bonding. Thus most covalent compounds are insoluble in sulphuric acid because they are unable to break up the intermolecular hydrogen bonds of the solvent in order to enter into solution; sometimes, however, solvation of a covalent solute can take place initially and then the solvate may enter into solution either ionically or by forming hydrogen bonds with the solvent.

The dielectric constant of sulphuric acid supports an extensive self-ionization which gives rise to the high value of the specific conductivity. The main self-ionization process in the pure liquid is thought to be

2H2S04 ^ H3S04+ + HS04-for which the ionic product [H3S04

+] [HS04~] is calculated as 1-7 X 10-4 at 10°C; the corresponding value [H30+] [OH"] for water is about 3 X 10~15 at the same temperature.

Like water and ammonia, sulphuric acid is an amphoteric solvent in that it may act either as an acid or as a base. The high value of the ionic product implies extensive proton transfer in the pure liquid

HaS04 + H2S04 ^ H3S04+ + HS04-t _ H+ — J L _ H+ — * base acid acid base

A sulphur trioxide transfer reaction also takes place in pure sulphuric acid to give disulphuric acid and water

2H2S04 ^ H2S207 + H20 both of which are virtually completely ionized at this concentration

H2S207 + H2S04-> H3S04+ + HS207-H20 + H2S04 -► H30+ + HS04-

The latter three equations reduce to 2H2S04 ^ H30+ + HS207-


which has been called the ionic self-dehydration reaction. The total concentration of ions produced in all the self-dissociation reactions is 0-43 molai which is a high enough concentration to depress the freezing point of pure sulphuric acid by 0-26°. Therefore, if no ionization of any kind took place, the freezing point of sulphuric acid would be 10-37 + 0-26 = 10-63°C.

Conductivity Measurements Whilst the ions HS04~ and H3S04

+ are found to conduct electricity, the transport numbers of other ions dissolved in sulphuric acid are lower by a factor of about 100. The reason for these low ionic mobilities is believed to be due to the high viscosity of the solvent which allows only very slow migration of ions in an electric field; a mechanism of "proton-jumping", made relatively easy by intermolecular hydrogen bonding, is postulated to account for the apparently high mobilities of the HS04~ and H3S04

+ ions: e.g. for the sulphuric acidium ion, H3S04

+,

HO \

S /

HO

O " ^

\ OH-.

+ °N

oy

OH * /

S ' \

OH

o OH

OH

HO

HO

\ O OH

S / \

L-HO OH-J- - -

O /

OH

cathode

OH-

\

HO O O OH \ • \ /

S S / \ / \

HO O HO 0 - -

O OH \ /

S / \

-HO OH-.

This phenomenon had been used to determine the number of H3S04

+ or, more usually, HS04~- ions in a given concentration D


of solute in sulphuric acid. If the solute is water, one mole of hydrogen sulphate ion is produced by each mole of water

H 2 0 + H2S04 -> H3O+ + HS0 4 -

Hence the conductivity of the solution above that of the pure solvent is due to a known concentration of HS04"~ ions. If another solute such as dinitrogen trioxide, N203, gives a conductivity of three times the value of water at the same molar concentration, then it may be assumed that 3 moles of hydrogen sulphate ion have been released into solution

3H2S04 + N 2 0 3 -> 2NO+ + 3HS04~ + H30+

Such conductivity measurements, coupled with cryoscopic and Raman spectral studies, have helped enormously in understanding solute behaviour in sulphuric acid solution.

Solvate Formation Freezing point-composition diagrams of the system M2S04-

H2S04 (where M is an alkali metal or NH4+) show the formation of solvates such as K2S04.H2S04, K2S04.3H2S04 and Li2S04. 9H2S04, some of which can be isolated as crystalline solids. They all contain the ion HS04~ so that they are better formulated as hydrogen sulphates (bisulphates), some of which are solvated: thus the examples quoted above become KHS04, KHS04.H2S04 and LiHS04.4H2S04. The alkaline earth metal sulphates also give similar, but less soluble, hydrogen sulphates when dissolved in sulphuric acid, e.g. Ba(HS04)2.2H2S04 and MgiHSO^. 2H2S04. Sulphates of Group III and IV metals like aluminium and tin are virtually insoluble in sulphuric acid and form no solvates.

Density and viscosity measurement of hydrogen sulphate solutions in sulphuric acid also indicate association of solvent with the cation although the solvation numbers are generally greater than in the isolable solvates: KHS04.2H2S04, Ca(HS04)2. 8H2S04 and NaHS04.3H2S04.


Solvolysis We saw in the last section that alkali and alkaline earth metal

sulphates give hydrogen sulphate ions on dissolution in sulphuric acid. This solvolysis is analogous to that observed when their oxides dissolve in water to form hydroxides

K2S04 + H2S04 -> 2KHS04

K20 + H20 -> 2KOH

The widespread occurrence of solvolysis among salts dissolved in sulphuric acid is in part an outcome of the extensive self-ionization of the solvent (even in water, where the self-ionization is extremely minute, hydrolysis is quite a usual phenomenon). In several cases, solvolysis may be regarded as a metathetic reaction in which one source of reacting ions is the solvent and one of the products is either unionized or insoluble in sulphuric acid:

(a) Sodium chloride is readily soluble in sulphuric acid but even below room temperature insoluble hydrogen chloride is evolved from the solution

NaCl + H2S04 -> HClf + NaHS04

Cl- + H3S04+ -> HClf + H2S04

(b) In the case of ammonium perchlorate the solvolysis product, perchloric acid, is a weak acid in sulphuric acid solution and so solvolysis approaches completion by the removal of perchlorate ions from the solution

NH4C104 + H2S04 -> NH4+ + HC104 + HS04" C104- + H3S04+ -> HC104 + H2S04

The last reaction illustrates the use of cryoscopy (a study of the freezing point depression of a solvent caused by added solute) in the study of reactions in pure sulphuric acid. As in water, the cryoscopic constant of sulphuric acid has first to be found by using an unionized solute; as mentioned above, these compounds


are rare but sulphuryl chloride, S02C12, has been used since it dissolves in pure sulphuric acid without increasing the conductivity of the latter. The freezing point depression produced by a known concentration of another solute can then be used to calculate the number of separate particles produced on dissolution of the solute after various corrections have been applied, such as allowance for the amount of solvent used up by self-ionization. When applied to ammonium perchlorate, cryoscopy showed that (in accordance with the equation given above) for every molecule of ammonium perchlorate added to the solvent, three particles were released into solution of which conductivity showed one was HSO4-; these facts do not support a simple ionic dissociation

NH4C104 -> NH4+ + C104-Raman spectra have also proved useful in the study of solvolysis

products. Though undissociated sulphuric acid has a strong and characteristic Raman spectrum, it is still possible to observe part of the spectrum of the hydrogen sulphate ion, so that its build-up can easily be followed by observing the Raman line at about 1050 cm-1. For example, nitric acid and metallic nitrates dissolve in sulphuric acid to give a conducting solution the Raman spectrum of which has lines characteristic of HS04~ ions (though hydrogen sulphate ions are produced by the self-ionization of sulphuric acid, their concentration is too low to be detected by Raman spectroscopy). The spectrum of the nitrate solution also has a further single line at 1400 cm-1 which was shown to belong to the nitronium ion N02

+

HNO3 + 2H2S04 -> N02+ + 2HS04- + H30+ NO3- + 3H2S04 -> N02+ + H3O+ + 3HS04-

Since the nitronium ion exhibits only one Raman line when in solution, Group Theory shows that it must be a centrosymmetric (triatomic) molecule, i.e. it must have a linear structure O—N—0+. Thus Raman spectra, besides acting as "fingerprints" for the particles present in solution, can also give their structure in favourable cases.


Nitronium hydrogen sulphate is also produced when dinitrogen pentoxide dissolves in sulphuric acid

N205 + 3H2S04 -> 2N02+ + 3HS04~ + H30+ which is in keeping with the formulation of N205 as nitronium nitrate, N02

+N03_ in the solid state.

Metallic nitrites react with pure sulphuric acid to give the nitrosyl ion NO+ which was identified by Raman spectroscopy, whilst the stoichiometry of the reaction

N02- + 3HaS04 -> NO+ + 3HS04- + H30+ was proved by cryoscopy. The same ion is obtained on dissolving dinitrogen trioxide in sulphuric acid

N203 + 3H2S04 -► 2NO+ + 3HS04" + H30+ whereas dinitrogen tetroxide gives a mixture of both nitrosyl and nitronium ions (see page 106)

N204 + 3H2S04 -> NO+ + N02+ + 3HS04~ + H30+

Protonation Reactions It is of some interest to describe separately the reactions in

which solvolysis takes place by the addition of a proton to the solute molecule. Such solutes behave as bases in sulphuric acid owing to the simultaneous production of hydrogen sulphate ions.

When water is dissolved in pure sulphuric acid two particles are released into solution, one of which was shown to be the hydrogen sulphate ion by the Raman spectral studies. This is in accordance with the equation

H20 + H2S04 -> H30+ + HS04-Several compounds which behave as acids in water are protonated in sulphuric acid solution instead of dissociating into ions and thereby act as bases in this solvent

HA + H2S04 -> H2A+ + HS04-e.g. H3P04 + H2S04 -> H4P04+ + HS04~

phosphoric acid


Sulphuric acid has also been used to produce transition metal-hydrogen bonds by similar protonation reactions. The hydrides are, in the main, only stable in solution and their existence has been proved by proton nuclear magnetic resonance. Typical hydrides which have been made this way are

[(C6H5)3PFe(CO)4H]+, [^-C5H5W(CO)3]2H+, and

[7T-C5H5Fe(CO)2Mn(CO)5]H+

In the latter case, where the hydrogen is thought to be bound to the iron atom, a stable hexafluorophosphate of the cation can be made by pouring the sulphuric acid solution into ice-cold water and adding ammonium hexafluorophosphate

[7T-C5H5Fe(CO)2Mn(CO)5]H+ + NH4PF6

-> [7T-C5H5Fe(CO)2Mn(CO)5]HPF6 + NH4+

Acids and Bases in Sulphuric Acid We have seen that by virtue of solvolysis and protonation many

compounds behave as bases in sulphuric acid by increasing the concentration of hydrogen sulphate ions. On the other hand, compounds which behave as acids in sulphuric acid are rare. One of the first to be recognized was disulphuric acid, H2S207, which ionizes to give the sulphuric acidium ion, H3S04

+ by the reaction

H2S207 + H2S04 ^ H3S04+ + HS207-

Cryoscopic and conductometric measurements showed, however, that disulphuric acid was only partially ionized in solution and was therefore a weak acid.

Other substances which act as strong acids in water, e.g. HC1, HC104 and HN03, are either unstable or undissociated in sulphuric acid. Hydrogen fluoride is similar in that it reacts with the solvent to form virtually undissociated fluorosulphonic acid

HF + 2H2S04 -> FS03H + HS04" + H30+


Boric acid and boric oxide are solvolysed by sulphuric acid according to the equations

H3BO3 + 6H2S04 -> B(HS04)4- + 3H30+ + 2HS04~ B203 + 9H2S04 -> 2B(HS04)4- + 3H30+ + HS04"

Both the resulting solutions are basic. However, if boric acid (or oxide) is dissolved in oleum a solution is formed which contains the sulphuric acidium ion

H3BO3 + 3H2S207 -> H3S04+ + B(HS04)4- + H2S04

since the tetrasulphatoboric acid HB(HS04)4 is a strong enough acid to protonate the H2S04 molecule. This reaction can be used to make acid solutions of known strength by adding boric acid to oleum until the disulphuric acid is just used up. A conductometric titration of these solutions against potassium hydrogen sulphate shows a pronounced minimum when 1 mole of base has been added:

B(HS04)4- + H3S04+ + K+ + HS04--> 2H2S04 + KB(HS04)4

Similar titrations can be performed using ammonium or alkaline earth hydrogen sulphates but it has proved impossible to isolate solid salts of the composition MB(HS04)4 or M'[B(HS04)4]2; apparently sulphuric acid is readily eliminated to form poly-hydrogensulphatoborates which probably contain B—O—B linkages, similar to the polyborates (e.g. borax, Na2B4O7.10H2O) known in the aqueous system.

Weak hydrogensulphato-acids are formed by tin and lead when their tetraaryls or tetraacetates are dissolved in sulphuric acid :

(C6H5)4Sn + 14H2S04 -+ H2Sn(HS04)6 + + 4C6H5S03H + 4H30+ + 4HS04~

Pb(OOCCH3)4 + 10H2SO4 -» HgPKHSOJe + + 4CH3COOH2+ + 4HS04-


However, tetramethyltin loses only one methyl group (as methane) on reaction with sulphuric acid and behaves as a base:

(CH3)4Sn + H2S04 -> (CH3)3Sn+ + HS04" + CH4

Radical Formation in Sulphuric Acid Many aromatic hydrocarbons dissolve in sulphuric acid to give

paramagnetic radical cations in which the unpaired electron is delocalized over the whole molecule:

CioH8 + H2S04 -> C10H8+

naphthalene "naphthalene positive ion"

The pseudo-aromatic molecule tetrasulphur tetranitride, S4N4, when dissolved in sulphuric acid gives rise to a radical thought to be SN2

+ formed by rupture of the tetrasulphur tetranitride molecule

S N-7^S /

I / λ N / / N +H2S04 ►S+SNa y/ i

/ sy— N s /

Both sulphur and iodine react with oleum (but not pure sulphuric acid) to give paramagnetic solutions, which in the case of sulphur are green in colour and contain two unidentified radicals. Iodine dissolves only slightly in pure sulphuric acid to give a pink solution, but if the sulphuric acid contains about 65 per cent of sulphur trioxide the solubility of iodine rises to over 10 molar, the solutions being blue and paramagnetic due to the presence of the iodine cation, I+, which is probably stabilized by solvation.

Other, diamagnetic, iodine complexes can be made in sulphuric acid. For example silver sulphate and iodine form a 1 : 1 complex


which has the stoichiometry Agl2+. The brown solutions of iodine in 30 per cent oleum contain the triiodine cation I3

+ which potassium iodate is able to oxidize to the 1+ cation

HIO3 + 2I3+ + 8H2S04 -> 71+ + 3H3O+ + 8HS04-

The other halogens do not form similar X+ cations either in concentrated oleum or in the presence of potassium iodate.

Hydrogen Fluoride Commercially, hydrogen fluoride is prepared by heating

calcium fluoride (in the form of good quality fluospar) with concentrated sulphuric acid. After condensation the hydrogen fluoride is stored, and transported, in steel tanks. A preliminary

TABLE 5

Physical properties of hydrogen fluoride

Melting point, °C Boiling point, °C Specific conductivity, ohm-1 cm-1

Viscosity, cP

Dielectric constant

Surface tension, dyne/cm

- 8 3 19-8

2-6 - 5-7 x 10-e (0°) / 0-26 (0°) I 0-57 (-50°) f 84 (0°)

4 134 (-42°) L 175 (-73°)

8-9 (19°)

purification is often carried out by making the solvate NaF.HF which is vacuum dried at 200°C for several hours before being heated to 450°C when the combined hydrogen fluoride is released and may be condensed out. Immediately prior to use this partially purified hydrogen fluoride is fractionated in an efficient still and collected directly in a conductivity cell or a reaction vessel. The pure liquid is so reactive that a single transfer in a sealed, metal or plastic apparatus is sufficient to cause a significant rise in the conductivity to 10~5 or 1(H ohm-1 cm-1 at 0°C.


It should be stressed from the start that liquid hydrogen fluoride is a dangerous reagent and on contact with the skin it causes very painful wounds which are slow to heal. Hydrogen fluoride also attacks any material containing silicon, so that glass has to be excluded from the list of possible handling materials. Of the metals, copper, iron and nickel may be used for reaction vessels but brass is readily attacked and the zinc removed. In recent years fluorinated plastics, such as Teflon (C2F4)n, and Kel-F (C2F3C1)W, have been used to construct vacuum lines, conductivity cells and reaction vessels capable of handling the pure solvent.

The low viscosity and surface tension of liquid hydrogen fluoride greatly assist in the mixing of heterogeneous reactants and also in the settling of precipitates from the metathetic reactions. The high dielectric constant would imply high solubility of many inorganic salts but unfortunately solvolysis often occurs. Liquid hydrogen fluoride has a high affinity for water and this, coupled with the widespread occurrence of solvolysis, makes it very difficult to remove water from the solvent. To date, there is no known chemical which will dehydrate it; the two most commonly used drying reagents, phosphorus(V) oxide and phosphorus pentachloride, are solvolysed:

P205 + 3HF -> HP02F2 + H2P03F PC15 + 5HF -> PF5 + 5HC1

It is possible to dehydrate hydrogen fluoride electrolytically, the dissolved water being preferentially electrolysed to give oxygen, oxygen difluoride, OF2, and hydrogen. Unfortunately, during the electrolysis severe anodic polarization occurs causing a fall in current and a corresponding increase in drying time, which may be as long as three days. Recently it has been found that if the current is reversed about every 30 sec extensive depolarization takes place which allows a higher average current to flow through the solvent and so greatly reduces the drying time.

The conductivity of liquid hydrogen fluoride has been explained by assuming the self-ionization reaction

2HF ^ H2F+ + F-


so that fluorides act as bases; no compound has yet been discovered which ionizes directly to the fluoronium ion, H2F+, although several covalent fluorides interact with the solvent to give this ion, e.g.

SbF5 + 2HF ^ H2F+ + SbF6-and thereby act as acids in liquid hydrogen fluoride.

Solvate Formation The melting point-composition diagram of the system H20-

HF indicates the formation of three solvates H2O.HF, H20.2HF and H20.4HF, of which the former has been shown to exist as oxonium fluoride, H30+F-, in the solid state. Water dissolved in hydrogen fluoride conducts electricity showing that oxonium ions are probably present in the solution also

H20 + HF -> H3O+ + F-Solvate formation has been demonstrated for the alkali and

alkaline earth metal fluorides, typical examples being KF.HF, 2KF.5HF, KF.3HF, KF.4HF and BaF2.HF. These solvates are somewhat unusual in that they contain solvated anions, F(HF)W~. Normally, as in the water and ammonia solvent systems, it is the cation which is solvated. The structures of two of these solvated anions, determined by X-ray diffraction studies on single crystals, show that the solvent is bound to the fluoride ion by hydrogen bonds:

F—H—F

Linear ion with equal H—F bond Two forms of H2F3_ in the crystal

lengths The H—F distances are unknown

The structure of the H2F3~ ion is very similar to that of solid hydrogen fluoride:


It is worthy of note that in solid hydrogen fluoride the hydrogen bonding gives rise to a two-dimensional polymer whereas in water, sulphuric acid and ammonia the resulting structure is three-dimensional. Like these three solvents, hydrogen fluoride is thought to be still extensively hydrogen-bonded in the liquid state, the two-dimensional structure explaining, in part, the comparatively low viscosity.

It follows from the above discussion that the self-ionization of liquid hydrogen fluoride ought strictly to be written

(2 + AZ)HF ^ H2F+ + F(HF)n-

but for clarity, the solvation of the fluoride ion will not be stressed in future pages.

Transport measurements made on solutions of acids and bases in liquid hydrogen fluoride show that the fluoride ion has a mobility greater than any other ion in this solvent. This is thought to indicate that a special process obtains for the transference of fluoride ions through the solution. The fluoride ion is strongly solvated in solution and hydrogen fluoride molecules are no doubt continually attaching themselves to and breaking away from the fluoride ion group. During electrolysis it is possible for a hydrogen fluoride molecule to be taken into the chain at one end and another molecule to be split away from the other:

F F / \ /

H H H / \ , / ? F

-* anode

This would result in a relatively rapid transference of the fluoride ion through the solution, especially as the high degree of solvation would make a purely ionic transfer a slow process. Since the solvated proton has only a low mobility in hydrogen fluoride, a similar mechanism for the transference of H2F+ ions cannot occur.

? F H HI \ / \

+ \ -> / H H ■ H F F


Solvolysis As stated previously, solvolysis occurs extensively in liquid

hydrogen fluoride. The acidic nature of the solvent causes the alkali and alkaline earth metal oxides and hydroxides to dissolve readily with the production of the metal fluoride, e.g.

M20 + 2HF -> 2MF + H20

MOH + HF -> MF + H20

Similarly, many metal carbonates effervesce in liquid hydrogen fluoride due to the evolution of carbon dioxide

M2C03 + 2HF -> 2MF + H20 + C02 M = alkali metal

M'C03 + 2HF -> M'F2 + H20 + C02

M' = Mg, Zn, Cd, Pb, Co, Ni, Cu and alkaline earth metals

The more electropositive metals such as lithium, calcium and even thallium dissolve in hydrogen fluoride with the evolution of hydrogen:

2T1 + 2HF -> 2T1F + H2

thallium(I) fluoride

The affinity of hydrogen fluoride for water often induces attack of hydroxyl groups covalently bonded to many elements to produce water (which enters solution as H30+) and an M—F bond. A typical example of this behaviour is the formation of fluorosulphonic acid from sulphuric acid; only one hydroxyl group is attacked

F \

(HO)2S02 + HF -> H20 + SO, HO


Many metal halides are rapidly solvolysed to the insoluble hydrogen halide which escapes from solution, the reaction being a valuable method for producing pure metal fluorides in solution. Unlike the solvolyses in bromine trifluoride (see Chapter V) no oxidation occurs during the fluorination and the element retains its original oxidation state:

VC14 + 4HF -> VF4 + 4HC1

MXn + «HF -> MF„ + nUX M = alkali metals, Ag, Zn, Hg, Fe X = Cl, Br, I

TaCl5 + 5HF -> TaF5 + 5HC1

The solvolysis of antimony pentachloride has been studied in detail. At — 75°C the chloride dissolves with the formation of H2FSbCl5F:

SbCl5 + 2HF -> H2F+ + SbCl5F-

Above — 40°C, chlorine-fluorine exchange occurs relatively quickly for three of the chlorine atoms whilst a fourth is substituted above 0°C :

SbCl5F- + 4HF -> SbF5Cl- + 4HC1

The final chlorine atom is exchanged only very slowly at room temperature:

SbF6Cl- + HF = 9 Ä . SbFe- + HC1

Acids and Bases in Hydrogen Fluoride All compounds giving rise to fluoride ions (solvated fluoride

ions) behave as bases in liquid hydrogen fluoride. The fluoride ions may be produced by direct ionization of the solute

KF + «HF -> K+ + F(HF)„-

PROTONIC SOLVENTS 63 or may arise from protonation and solvolytic reactions

NH3 + HF -> NH4+ + F-

BaCl2 + 2HF -> Ba2+ + 2F~ + 2HC1

No known compounds ionize directly to give fluoronium ions, therefore the only compounds which behave as acids in liquid hydrogen fluoride are those which interact with the solvent to give solvated protons:

H2F+A- -> H2F+ + A- \ ) not known

HA + HF -> H2F+ + A" J

SbF5 + 2HF -> H2F+ + SbFe~ known

Tracer experiments using radioactive 18F show a rapid exchange of fluorine between antimony pentafluoride and the solvent, which is consistent with the above equation. Furthermore, the dilute solutions of antimony pentafluoride conduct electricity, verifying the presence of ions, whilst infrared and Raman spectra prove the identity of SbF6~. In concentrated solutions ion pairs, thought to be H2F+SbF6~, are indicated by both spectral and conductometric studies :

2HF + SbF5 ^ HaF+SbF6- ^ H2F+ + SbF6~ Ionization drops rapidly as con- Complete ionization when concen-centration of SbF5 is increased tration of HF is above 80 per cent

above 20 per cent

Acidic solutions are also formed when phosphorus and arsenic pentafluorides are dissolved in liquid hydrogen fluoride

PF5 + 2HF -> H2F+ + PF6-(As) (As)

These solutions react as typical acids dissolving many metals with the evolution of hydrogen. With solutions or suspensions of


metal fluorides acid-base neutralization reactions occur: KF + H2FPFe -> KPF6 + 2HF

NdF3 + 3H2FSbF6 -> Nd(SbF6)3 + 6HF

In the latter case the addition of a strong base reprecipitates the neodymium trifluoride from solution :

3NaF + Nd(SbF6)3 -> NdF3 j + 3NaSbF6

Amphoteric behaviour is demonstrated by several elements of which aluminium is typical:

A1F3 + 3NaF —-> Na3AlF6

Na3AlF6 + 3H2FSbF6 -* A1F3 + 3NaSbF6 + 6HF

It was thought originally that boron trifluoride dissolved in liquid hydrogen fluoride to form tetrafluoroboric acid:

BF3 + 2HF -> H2F+ + BF4~

but more recent work has shown that boron trifluoride obeys Henry's law when dissolved in hydrogen fluoride indicating that little or no reaction with the solution occurs to give tetrafluoro-borate ions. However, the boron trifluoride in solution acts as a strong Lewis acid on the addition of fluoride ion so that some of its reactions superficially resemble Br0nsted-Lowry acid-base neutralizations involving a hypothetical acid H2F+BF4-:

AgF + BF3 -> AgBF4

cf. AgF + H2F+BF4 -> AgBF4 + 2HF

Basic solutions are formed when nitrosyl fluoride or dinitrogen trioxide dissolve in hydrogen fluoride:

NOF -> N O + F-N203 + 2HF -> 2NO+ + H20 + 2F~

The presence of the nitrosyl ion is demonstrated by the ready precipitation of nitrosyl perchlorate on the addition of perchloric

PROTONIC SOLVENTS 65 acid to the solution. An interesting reaction of the nitrosyl ion is its formation of a violet solution (called violet hydrogen fluoride) when treated with nitric oxide under pressure. The species responsible for the colour is thought to be the radical ion N202

+. The same ion is formed by reacting nitric oxide with the nitrosyl ion dissolved in sulphuric acid.

A solution of nitrosyl fluoride in liquid hydrogen fluoride is a very reactive fluorinating agent as shown by the following typical reactions:

NOCI — - * NOF — - NO+ + F(HF)n- (n = 3, 6) solvolysis

NOF + Be -> (NOF)2BeF2; perhaps (NO)2BeF4

NOF + B -> NOBF4

NOF + Al -> A1F3

NOF + SOCl2 -> SOF2

Oxidation Reactions Hydrogen fluoride is a useful solvent in which to prepare

metallic fluorides. We have seen that solvolysis by liquid hydrogen fluoride produces many fluorides in which the metals are in their original oxidation state. It is possible, by the use of appropriate oxidizing agents, to produce elements in a higher oxidation state. Elemental fluorine diluted with gaseous nitrogen will oxidize thallium(I) to thallium(III) :

T1F + F2 — -> TIF3 and silver(I) to silver(II) :

AgF + £F2 - ^ - * AgF2

Chlorine trifluoride dissolved in liquid hydrogen fluoride is a more convenient reagent and with it metallic tellurium can be oxidized to the hexafluoride which crystallizes out of solution at — 78°C as large clear crystals. Red phosphorus and metallic


arsenic give hexafluorophosphate and hexafluoroarsenate ions respectively, e.g.

P + CIF3/HF -> H2F+ + PF6-

As + CIF3/HF -> H2F+ + AsFe-

Electrochemical Fluorination Electrolysis of liquid hydrogen fluoride produces fluorine

at the anode and hydrogen at the cathode. If an organic compound is present in the electrolytic cell and the hydrogen fluoride electrolysed using a potential insufficient to discharge fluorine, complete fluorination of the organic compound occurs at the nickel anode as though "nascent" fluorine is formed during the electrolysis, e.g. R20 -> RF20 R = saturated or unsaturated alkyl

or aryl R3N -> RF3N RF = saturated perfluoroalkyl (linear

or cyclic) RH -> RFF

Similarly, the electrolysis of ammonium fluoride as a 5-10 per cent solution in hydrogen fluoride gives nitrogen trifluoride, NF3:

NH4+ + (F) -> NF3

This is somewhat similar to the preparation of nitrogen trichloride, where chlorine is bubbled through an aqueous solution of ammonium salt :

NH4+ + Cl2 -> NC13

During the electrolysis of ammonium fluoride two by-products, eis and trans dinitrogen difluoride, are formed:

F /

N=N N = N / \ /

F F F eis trans

PROTONIC SOLVENTS 67 The solvolysis of sulphuric acid by hydrogen fluoride stops

when one hydroxyl group has been replaced, the product being fluorosulphonic acid, HS03F. It is possible to replace the remaining hydroxyl group by electrolysing a solution of fluorosulphonic acid in liquid hydrogen fluoride:

HOS02F -> F2S02 sulphuryl fluoride

When sodium perchlorate is electrolysed as a 10 per cent solution in hydrogen fluoride, perchloryl fluoride, C103F, is formed at the nickel anode together with oxygen difluoride and traces of oxygen and chlorine whilst hydrogen is evolved at the iron cathode:

C104- + HF + H+ — d 4 CIO3F + H2 + i0 2 + OF2

As nitrous oxide obeys Henry's law when it is dissolved in liquid hydrogen fluoride no interaction with the solvent must occur and indeed the solvent shows no increase in conductivity as the gas enters solution. Because of this, a conductivity additive, either sodium or potassium fluoride, is added to the solution prior to the electrolysis which gives low yields (based on current) of nitrogen trifluoride, oxygen difluoride and pernitrylfluoride:

HF electrolysis N20 z^^:^ NF3 + OF2 + N03F

2% 2% 1 %

Extensive attack of the nickel anode accounted for 50-80 per cent of the current used in the experiment.

Related electrolyses of nitrosyl fluoride dissolved in hydrogen fluoride give only nitric oxide and nitrosyl fluoride as the gaseous products:

N O + _cathode_ N O l -> NOF

2 F - - ^ U F 2 J


A considerable fraction of the liberated fluorine attacked the nickel anode which explains the presence of free nitric oxide in the effluent gases.

The Other Hydrogen Halides as Ionizing Solvents The low melting and boiling points of hydrogen chloride,

bromide and iodide suggest that little or no association occurs in the liquid phase, quite unlike hydrogen fluoride which is extensively hydrogen bonded. This is further reflected in their low dielectric constants which are, to a first approximation, a good measure of the polarity of the solvent molecules. As might be expected salts have only a limited solubility in these solvents whilst certain covalent compounds, such as the boron, tin(IV) and phosphoryl halides and many organic compounds, are quite soluble. Hydrogen chloride has been studied the most extensively and will be discussed briefly; hydrogen bromide and hydrogen iodide are very similar in solvent properties to hydrogen chloride but contamination of solutes with iodine formed by disproportion-ation of hydrogen iodide is often noticed.

TABLE 6

Physical properties of the hydrogen halides

Melting point, °C Boiling point, °C Dielectric constant Electrical conductivity,

ohm-1 cm-1

HC1

-112 -83-7 10-6

3-5xl0-9(-85°)

HBr

-88-5 - 6 7 0

70 lxl0- 1 0(-84°)

HI

-50-8 -35-4 (4 atm)

3-4 very low

As can be seen from Table 6, hydrogen chloride has a narrow liquid range which makes experimentation difficult; if vacuum manipulation of the solvent is used the temperature must be closely controlled at about — 90°C to prevent an excessive rise in


pressure. The conductivity may be explained in terms of the self-ionization

3HC1 ^ H2C1+ + HC12-The hydrogen dichloride ion can be isolated in solvates formed by recrystallizing quaternary ammonium chlorides from their hydrogen chloride solution

NR4C1 + HC1 -> NR4HC12

No acids have been found which ionize directly to the H2C1+ ion; however, certain co valent halides apparently dissolve to give solvated protons, e.g.

BClg + 2HC1 ^ H2C1+ + BC14-

whilst protonation or chloride ion exchange reactions involving the solvent produce basic solutions:

C5H5N + 2HC1 -> C5H5NH+ + HC12-pyridine

PC15 + HC1 -> PC14+ + HC12-Conductometric titration techniques can be used to follow neutralization reactions between these acids and bases, a few examples being

C5H5N + BCI3 + HC1 -> C5N5NHBC14

pyridinium tetrachloroborate

Et4NCl + BF3 ► Et4NBClF3

2Me4NCl + SnCl4 ► (Me4N)2SnCl6

When a strong electron pair donor (i.e. a Lewis base) dissolves without ionization in hydrogen chloride the addition of an electron acceptor (i.e. a Lewis acid) produces the expected donor-acceptor adduci, as for example in the reaction

CI3PO + BCI3 -> CI3PO—BCI3 Lewis base Lewis acid adduct


Although these simple reactions can be extended to other analogous systems, the chemistry of the liquid hydrogen halide solvents will apparently never be very extensive.

Acetic Acid Glacial acetic acid is prepared industrially on a large scale by

several methods of which the following are typical:

(i) 2CH3CHO + 0 2 j ^ + 2CH3COOH acetaldehyde air

/■■\ /".u /~>tr , /~./~w activated charcoal „ „ „ „ „ „ (11) CH3OH + CO — ^ . - ^ — ^ CH3COOH

(iii) Q H 1 0 + 0 2 ^ ^ ^ CH3COOH Butane air

TABLE 7

Physical properties of glacial acetic acid

Melting point, °C Boiling point, °C Dielectric constant Specific conductivity, ohm-1 cm-Viscosity, cP

16-62 1181

6-2 (25°) 2-4 - 3-0 x 10-9 (25°)

1-15 (25°)

The 99-7 per cent commercial acid is first boiled with chromium-(VI) oxide, Cr03, to remove oxidizable impurities, and then dehydrated by refluxing with a small amount of pure acetic anhydride for 2 or 3 hours:

(CH3CO)20 + H20 -> 2CH3COOH

a fractional distillation at atmospheric pressure in a well-dried apparatus affords the anhydrous acid, the purity of which may be checked by the melting point.

PROTONIC SOLVENTS 71 The physical constants of acetic acid permit its handling in open

vessels at room temperature although precautions should be taken to stop the entry of moisture, for example by working in a dry-box. Certain experiments, in which traces of moisture are of no consequence, may even be carried out in the open laboratory provided the vessels are normally closed with corks, rubber bungs or well-fitting ground glass stoppers, and are only opened to the atmosphere for brief periods during manipulations. When corks or bungs are used they must be covered by thin polythene sheet otherwise the water absorbed in them will be extracted into the solvent.

Although the low dielectric constant cannot be expected to support a very extensive self-ionization, the specific conductivity of the pure liquid suggests that slight ionization does occur:

2CH3COOH ^ CH3COOH2+ + CH3COO-

for which the ionic product [CH3COOH2+] [CH3COOi is ca. IO-14 at room temperature.

The many salts which are soluble in glacial acetic acid include nitrates, cyanides, thiocyanates and halides, particularly iodides. Acetates, which act as bases in acetic acid, are also often soluble. A useful rule for solubility is that compounds insoluble in water are also insoluble in acetic acid, although the reverse of the rule does not necessarily apply. It is therefore not surprising that metathetic reactions such as

AgN03 + KI -* Agl| + KNO3

AgN03 + KSCN -> AgSCNj + KN03

ZnCl2 + Na2C204 -> ZnC204| + 2NaCl sodium oxalate

proceed smoothly in acetic acid. However, though metathetic reactions do take place in glacial acetic acid, extensive ion-pair


formation always occurs even for the strongest electrolytes; for example, the "ion-pair dissociation constants" for potassium bromide and potassium acetate are 1 X 10~7 and 3 X 10~"7

respectively. Such low values are, of course, a natural outcome of the low dielectric constant of the solvent. Moreover, spectro-photometric studies have shown that in certain cases ion-triplets and ion-quadruplets can occur at concentrations above 0-2 molar.

X-Y+X- Y+X-Y+ Y+X-Y+X- X-Y+ Y+X-

ion-triplets ion-quadruplets

A further useful observation about solubilities is that all metallic sulphates, unlike pure sulphuric acid, are insoluble in glacial acetic acid, thus allowing the solvent to be used for the preparation of anhydrous metal sulphates which are precipitated when solutions of sulphuric acid and a hydrated metal salt are mixed :

Cu(N03)2«H20 + H2S04 -► CuS04j Ni(N03)2wH20 + H2S04 -> NiS04 j Co(N03)2«H20 + H2S04 -» CoS04|

Water, ammonia and hydrogen sulphide and several inorganic covalent halides, such as stannic chloride, germanium tetra-chloride and arsenic trichloride, are quite soluble in acetic acid as are the common acids, perchloric, sulphuric, orthophosphoric and hydrochloric. No silver chloride is precipitated on mixing solutions of arsenic trichloride or stannic chloride with silver nitrate, showing that no ionization (or solvolysis to hydrogen chloride) occurs on dissolution. As in aqueous solution, hydrogen sulphide precipitates heavy metal sulphides when passed through a solution of their salts.

MtOOCCHa)* + |H2S -> MS* j + xCH3COOH 2

The addition of a strong acid, such as hydrogen chloride in acetic acid, causes the precipitated sulphide to redissolve.

PROTONIC SOLVENTS 73 Comparatively few salts form stable acetic acid solvates. The

metal acetates give the most examples, e.g. KOOCCH3.HOOCCH3, KOOCCH3.2HOOCCH3, Ba(OOCCH3)2.2HOOCCH3, Ba(OOCCH3)2.3HOOCCH3, and Cu(OOCCH3)2.HOOCCH3 ; HgCl2.HOOCCH3 is one of the few metal halide solvates. Apparently all these compounds are simple solvates and contain no complex ions.

Solvolysis Glacial acetic acid is an acidic solvent so that solvolysis may be

expected to occur with those solutes which are capable of reacting with the solvent protons. For example, sodium carbonate effervesces fairly vigorously evolving carbon dioxide:

Na2C03 + 2CH3COOH -> 2NaOOCCH3 + C02 + H20

and sodium hydroxide dissolves to give sodium acetate: NaOH + CH3COOH -> NaOOCCH3 + H20

Electropositive metals, including thallium, evolve hydrogen when placed in glacial acetic acid:

2T1 + 2CH3COOH -> 2T100CCH3 + H2

In the case of zinc metal, only slight reaction occurs to give hydrogen and a precipitate of zinc acetate; however, if base (sodium acetate) is added to the solution the rate of dissolution of the zinc increases, no doubt due to the formation of a soluble acetatozincate complex :

Zn(OOCCH3)2 + 2NaOOCCH3 -> Na2Zn(OOCCH3)4

It might have been predicted that, when dealing with a protonic solvent like glacial acetic acid, covalent halides such as stannic chloride and arsenic trichloride would be extensively solvolysed to give hydrogen chloride (compare their reaction with water, pure sulphuric acid and liquid ammonia). However, as stated above these two chlorides do not solvolyse although stannic


chloride gives an octahedral solvate SnCl4.2HOOCCH3 which behaves as an acid in acetic acid (p. 76).

Boron trichloride, on the other hand, is rapidly solvolysed to give mainly hydrogen chloride and tetraacetyl diborate:

2BC13 + 6HOOCCH3

-> 6HC1 + [(CH3COO)2B]20 + (CH3CO)20

The infrared spectrum of the diborate shows that two types of C = 0 bond are present in the molecule, one C = 0 group being co-ordinated to a boron atom, the other being "free". The spectrum does not distinguish, however, between the two most probable alternatives :

CH3 I

H3CCOO OOCCH3 O — C = 0

I I / \ O — B — O—B—O CHsCOOB O BOOCCH3 I t t i \ /

H3CC = 0 0 = C C H 8 0 = C — O I

CH3 An interesting solvolytic reaction is that shown by potassium

hydrogen phthalate, KOOCC6H4COOH, which gives potassium acetate and undissociated phthalic acid.

KOOCC6H4COOH + CH3COOH -* KOOCCH3 + C6H4(COOH)2

Potassium hydrogen phthalate is therefore useful as a standard for preparing basic solutions of accurately known strength since the alkali metal acetates are deliquescent, making accurate weighing difficult to carry out.

Complex Formation Many of the complexes studied in acetic acid have been of the

acetatometallate type and often demonstrate the amphoteric nature of the metal.


Sodium acetate first precipitates zinc acetate from a solution of a zinc salt in acetic acid

ZnCl2 + 2NaOOCCH3 -> 2NaCl + Zn(OOCCH3)2 j but on the addition of an excess of acetate ion the precipitate redissolves due to the formation of sodium tetracetatozincate.

Zn(OOCCH3)2 + 2NaOOCCH3 -> Na2Zn(OOCCH3)4

Similarly, cupric acetate is soluble in potassium acetate solutions. An interesting variation is found when excess of ammonium acetate is added to a cupric salt. The sparingly soluble green cupric acetate first formed redissolves:

Cu(N03)2 + 2NH4OOCCH3 -> 2NH4N03 + Cu(OOCCH3)2| Cu(OOCCH3)2 + 4NH4OOCCH3 -> (NH4)4Cu(OOCCH3)6

and from the solution, the ammonium acetatocuprate(II) can be isolated as a solid solvate (NH4)4Cu(OOCCH3)6.4HOOCCH3. If the acetatocuprate(II) solution is heated to 100°C the colour changes from blue to the intense violet-blue of the ion Cu(NH3)4

2+. The ammonia, which presumably forms a stronger complex than the acetate group, arises from the dissociation of ammonium acetate at about 100°C:

CH3COOH NH4OOCCH3 N

N NH3 + HOOCCH3 100°

When the anhydrous rare earth acetates dissolve in glacial acetic acid, conductance measurements indicate that uncharged species, which contain chelating acetate groups, occur in the solutions; e.g. neodymium acetate:

CH3 / C \ o o

* * N d ' \ HsCC i N Q CCH3 \ 0 / \ 0 '


The solubility of neodymium acetate is increased by the addition of potassium acetate and a bluish coloured complex is formed. Under the influence of an electric field the coloured species migrates towards the anode as expected for an acetato-neodimiate(III) ion:

Nd(OOCCH3)3 + «KOOCCHg -> KnNd(OOCCH3)3+n

When stannic chloride dissolves in acetic acid heat is evolved and a solvate, SnCl4.2HOOCCH3, is formed. The solution conducts electricity and exhibits weakly acidic properties, which have been explained by assuming the further solvation reaction:

SnCl4.2HOOCCH3 + HOOCCH3 -> [H2OOCCH3]+ [HSnCl4(OOCCH3)2]-

% [H2OOCCH3]+ + [HSnCl4(OOCCH3)2]-

Zinc chloride forms a similar solvate on dissolution in acetic acid:

ZnCl2 + 2HOOCCH3 -> ZnCl2.2HOOCCH3

When the solution is treated with an excess of hydrogen chloride, the solvate decomposes and is replaced by the tetrachloro-zincate(II) ion:

ZnCl2.2HOOCCH3 + 2HC1 -> H2ZnCl4 + 2HOOCCH3

\ HOOCCHs [H2OOCCH3]2+ZnCl4

2-

Cobalt(II) halides dissolve in acetic acid to give blue-coloured solutions which are thought to contain a tetrahedral, neutral cobalt solvate, CoX2.2CH3COOH. The addition of lithium halide to these solutions results in the formation of two anionic cobalt complexes:

CoX2.2CH3COOH + X- -> CoX3.CH3COOH- + CH3COOH CoX3.CH3COOH- + X- -> CoX4

2~ + CH3COOH X = Cl,BrorSCN


Cobalt(II) acetate is undissociated in glacial acetic acid solution although the addition of small amounts of water to the solution increases the dielectric constant sufficiently to allow a partial ionization:

Co(OOCCH3)2.2CH3COOH

^ CoOOCCH3.2CH3COOH+ + CH3COO"

When lithium acetate is added to the solution of cobalt(II) acetate in glacial acetic acid, anionic acetatocobaltate(II) complexes are formed due to the amphoteric nature of cobalt:

Co(OOCCH3)2.2CH3COOH + CH3COO~

Cti COO -

-> [Co(OOCCH3)3CH3COOH]- — > Co(OOCCH3)42"

Acids and Bases in Glacial Acetic Acid From the dissociation reaction of acetic acid

2HOOCCH3 ^ H2OOCCH3+ + OOCCH3"

acetates can be recognized as bases while compounds producing solvated protons will act as acids. The strength of an acid, as measured by the concentration of solvated protons in the solution, will depend on two major factors. In the first case, the dissociation of an acid HX will be determined both by the tendency of HX to split off protons and by the affinity of acetic acid for protons :

HX + HOOCCH3 -> H2OOCCH3+ + X-

Secondly, if the acid is strong, this reaction will go to completion; but it should be remembered that, due to the low dielectric constant of acetic acid, ion-pair formation will be extensive causing a substantial lowering of the concentration of "free" solvated protons:

A- + H2OOCCH3+ -» [A]- [H2OOCCH3]+


The dissociation of inorganic acids in acetic acid, as determined by conductivity studies, decreases in the order

HC104 > HBr > H2S04 > HCl > HN03 ratio 400 160 30 9 1

The strongest acid, perchloric, is present as ions and ion-pairs even in very dilute solutions whereas the weak acids are considered to be present as three entities: ions; ion-pairs; and undissociated molecules. Bases such as the alkali metal acetates or certain amines also form ion-pairs even at high dilution.

It is possible, even though the ionization of both acid and base is not extensive, to carry out neutralization reactions in acetic acid which may be followed conductometrically. For example, the reaction

HC104 + NaOOCCHg -> NaC104 + HOOCCH3

may be studied by placing a weighed amount of potassium hydrogen phthalate into a known weight (or volume) of glacial acetic acid contained in a well-stoppered conductivity cell fitted with platinum electrodes and a magnetic stirrer:

KH(COO)2C6H4 + HOOCCH3-^ KOOCCH3 + (HOOC)2C6H4

The conductivity of the solution is measured after each addition of a standard solution of perchloric acid in acetic acid from a weight burette. During the neutralization reaction, the conductivity of the solution will fall virtually linearly as acetate ions are consumed; when an equivalent quantity of perchloric acid has been added the conductivity will rise with each further addition due to the increasing concentration of acetic acidium ions. The intercept of the two branches of the conductivity curve gives the neutralization point.

When a similar titration is carried out by adding a standard solution of sodium acetate to sulphuric acid, the conductivity curve shows two breaks corresponding to the neutralization of each of the two acidic protons. If the titration is followed potentiometrically, however, the only end point which can be

PROTONIC SOLVENTS 79 detected corresponds to the neutralization of one proton, and sodium hydrogen sulphate is formed:

NaOOCCH3 + H2S04 -> NaHS04 + HOOCCH3

The conductivity curve of the titration of perchloric acid against the lanthanide (rare earth) triacetates shows a break when three equivalents of acid have been added :

Ln(OOCCH3)3 + 3HC104 -> Ln(C104)3 + 3HOOCCH3

0-5 1-0 Equivalents of perchloric acid

1-5

FIG. 9. Conductometric titration of potassium acetate with perchloric acid in glacial acetic acid.

If a weaker acid such as sulphuric acid is used for the titration, the break in the conductivity curve occurs when only two acetate groups have been neutralized, the product precipitating out of solution:

Ln(OOCCH3)3 + H2S04 -> Ln(OOCCH3)S04| + 2HOOCCH3

Many metal di- and tri-acetates give very poor end points when titrated with perchloric acid because of their high degree of association in glacial acetic acid. However, if acetonitrile, CH3CN,


is added to the solution prior to the titration, sharp, well-defined end points are obtained for the acetates of those metals which are known to form strong cyanide complexes: Cu, Co, Ni, Zn and Hg. It is thought that acetonitrile complexes with the metal ions, so freeing the acetate ions for reaction with the perchloric acid:

M(OOCCH3)2 + *CH3CN -> MCCHaCN)/* + 2CH3COO-CH3COO- + HC104 -> CH3COOH + C104-

I I I I L_ 0 K) 2-0 3-0 40

Equivalents of acid used

FIG. 10. Conductometric titration of a rare earth metal triacetate with either perchloric or sulphuric acid in glacial acetic acid.

Amines behave as bases in acetic acid by virtue of their reaction with the solvent which produces the acetate ion, i.e. they are levelled to the basicity of the acetate ion:

NH3 + HOOCCH3 -> NH4+ + OOCCH3-(CH3)2NH + HOOCCH3 -> (CH3)2NH2+ + OOCCFLf

C5H5N + HOOCCH3 -> C5H5NH+ + OOCCH3" pyridine

Since acetic acid is a more acidic solvent than water the above protonation reactions will occur to a much greater extent than the corresponding formation of hydroxyl ions in water:

NH3 + H20 ^ NH4+ + OH-

PROTONIC SOLVENTS 81 It follows that amines behave as stronger bases in acetic acid than in water and for this reason glacial acetic acid has received much attention from analytical chemists as a medium for the quantitative determination of many organic amines. For example, it is possible, by a suitable procedure, to determine the constituents of an aliphatic amine mixture such as trimethylamine, dimethylamine and monomethylamine. A potentiometric titra-tion of an aliquot of the mixture against perchloric acid in glacial acetic acid gives the total amine content. To a second aliquot, acetic anhydride is added and the mixture allowed to stand at room temperature for 3 hr when the dimethylamine and monomethylamine are acylated to non-basic products; a titration against perchloric acid gives the trimethylamine content. Finally, salicylaldehyde is added to a third aliquot when a Schiff's base is formed with the primary amine, monomethylamine:

+ CH3NH2-

C = NCH,

Not titratable

which allows the determination of the combined dimethylamine and trimethylamine content. From the three titrations the concentration of each amine in the original mixture can be calculated.

As an interesting variation on the above theme cupric 8-quinolinolate may be analysed by the following procedure. The complex is first broken down by treatment with a glacial acetic acid-acetic anhydride mixture:

/"Λ

N\ // N

- * - 2 Cu

+ Cu(00CCH,)2

OOCCH,

E


and the basic nitrogen group of the acylated quinoline titrated potentiometrically with perchloric acid :

f W / = ■(

\\ t~ acetic

+ HC104 ———► H acid —OOCCH3

ft vNH+

\u \ //

+ C104

OOCCH,

Finally, the addition of acetonitrile to the solution releases the acetate groups bound to the copper and enables them to be titrated with more of the perchloric acid. A typical titration curve for this analysis is shown in Fig. 11.

1-0 2-0 3-0 4-0 Ratio perchloric acid/copper complex

5-0

FIG. 11. Potentiometric titration of copper(II) 8-quinolinolate with perchloric acid in glacial acetic acid.

A somewhat similar reaction scheme has been used for the determination of nickel which is first precipitated and isolated as nickel dimethylglyoxime, Ni(dmg)2. The precipitate, after being digested in a glacial acetic acid-acetic anhydride mixture, is titrated with perchloric acid in the presence of acetonitrile:


Ni(dmg)2 + 2HOOCCH3

CH3CN Y

Ni(CH8CN)s*- + 2CH3COO-

Oxidation Reactions Glacial acetic acid is a useful medium for studying compounds

in solution which are readily solvolysed by other solvents such as water. An example of this is the preparation of stannic iodide by the oxidation of tin dichloride using iodine dissolved in a mixture of acetic acid and acetic anhydride. The anhydride is added to the solvent to remove water of crystallization from the tin dichloride crystals. The oxidation of stannous tin to stannic takes place readily at the reflux temperature and on cooling, crystals of the orange stannic iodide separate out leaving stannic chloride in the solution

SnCl2 + I2 -> SnCl2I2

2SnCl2I2 -> SnCl4 + Snl4

As in aqueous solution, the oxidation of the oxalate ion by a eerie salt can be studied quantitatively by electrometric titration. The usual method used to determine the equivalence point of the titration is to place two electrodes in the solution, apply a potential difference between them and note changes in the current passing during the addition of oxalate to the eerie salt, ammonium hexanitratocerate. As the end point of the redox reaction is approached the current passing falls rapidly with each small addition of oxalate, till at the end point it becomes zero. Such titrations show that 2 moles of eerie ion react with a mole of oxalate to produce 2 moles of carbon dioxide:

2Ceiv + caCV- " " ^ 2Cem + 2CO.


Interestingly, tracer experiments using radioactive 14C showed that both moles of carbon dioxide originated from the solvent and not from the oxalate ion as might have been expected. The cerate solution in acetic acid is light sensitive and the volumetric work is therefore carried out with amber-coloured burettes and the solutions stored in amber bottles.

Using a similar experimental technique it is possible to study the oxidation of ferrous iron by sodium permanganate in the presence of perchloric acid. The permanganate solution is very unstable in acid solution and so the titration is normally carried out by adding the permanganate in small amounts to a stirred solution of the ferrous salt containing perchloric acid. The same reaction occurs in both acetic acid and water:

2NaMn04 + 16H+ + 10Fe2+ -> 10Fe3+ + 2Na+ + 2Mn*+ + 8H20

The normality of the permanganate, as determined by this reaction, may be checked by titration against sodium oxalate in acetic acid, the end point being the first permanent tinge of colour caused by excess of the permanganate:

2NaMn04 + 16H+ + 5C2042~

perchloric acid

-> 10CO2 + 2Na+ + 2Mn2+ + 8H20 The ferrous solution used in these experiments was prepared by dissolving the very soluble ferrous perchlorate hexahydrate in glacial acetic acid containing enough acetic anhydride to destroy the water of crystallization. The solutions must be stored in an atmosphere of nitrogen since oxidation of the ferrous iron occurs readily in the presence of air. This oxidation occurs more quickly in acetic acid than in water, probably owing to the increased solubility of oxygen in the former solvent.

The oxidation of cerium(III) acetate by lead(IV) acetate has been studied kinetically and shown to be first order dependent on the concentrations of both Ce(III) and Pb(IV). The postulated mechanism for the reaction is


Ce(III) + Pb(IV) -> Pb(III) + Ce(IV) rate determining

Pb(III) + Ce(III) -> Pb(II) + Ce(IV) rapid

There was no direct evidence for a Pb(III) species but the proposed scheme gave the best fit to the kinetic data.

Further Reading AUDRIETH, L. F. and KLEINBERG, J., Non-aqueous Solvents, Wiley, New York,

1953, Chapters 8, 9 and 10. BECKETT, A. H. and TTNLEY, E. H., Titrations in Non-aqueous Solvents (3rd

edition), British Drug Houses Ltd., 1960. GILLESPIE, R. J. and ROBINSON, E. A., The sulphuric acid solvent system,

Advances in Inorganic Chemistry and Radiochemistry, 1, 385 (1959). SIMONS, J. H., Fluorine Chemistry, 1 (Academic Press Inc., 1950), p. 225. A

brief survey of the history, physical and chemical properties of hydrogen fluoride.

CHAPTER V

NON-PROTONIC SOLVENTS

IN ALL the solvents so far discussed, the base analogues have varied with each solvent whereas the solvated proton was, in every case, responsible for acidity, the solvents being referred to as protonic solvents. There are, however, many solvents which contain no ionizable hydrogen. The examples of typical non-protonic solvents which will be briefly reviewed in this chapter are bromine trifluoride, sulphur dioxide, dinitrogen tetroxide and phosphorus oxychloride.

The chemistry carried out in these solvents can sometimes be systematized by using Franklin's definition of "acids" and "bases" (see Chapter I) as substances which give rise, respectively, to the cation and anion characteristic of the solvent under consideration. This may be exemplified by considering the case of bromine trifluoride which is thought to ionize in the following way:

2BrF3 ^ BrF2+ + BrF4~

In such a system, salts like potassium or silver tetrafluoro-bromates(III), KBrF4 and AgBrF4, will behave as bases whereas a compound BrF3.SbF5, formed when antimony pentafluoride dissolves in bromine trifluoride, can behave as the acid [BrFJ+fSbFß]". By a suitable experimental procedure, the neutralization reaction

AgBrF4 + BrF2SbF6 -> AgSbF6 + 2BrF3

can be followed conductometrically in much the same manner as its counterpart in water.

86

NON-PROTONIC SOLVENTS 87

However, recent work on isotope exchange reactions between liquid sulphur dioxide and either thionyl chloride or metallic sulphites, has shown that the self-ionization of sulphur dioxide, if it occurs at all, is so slight as to be of no significance. Obviously in such a case it is impossible to recognize either acids or bases using the definitions given above and the solvent can be assumed to take no part in ionic reactions other than to act as the reaction medium.

Liquid dinitrogen tetroxide, like pure sulphuric acid, has several possible dissociation mechanisms, one of which involves the neutral radical nitrogen dioxide, N0 2 . Furthermore, because of its very low dielectric constant, liquid dinitrogen tetroxide is a poor ionizing solvent but, as will be seen, its chemistry can be greatly extended by the addition of a suitable co-solvent.

Phosphorus oxychloride is a little-studied solvent and consequently its chemistry is not yet very extensive. As with certain other chlorides (for example, sulphur monochloride, arsenic trichloride and carbonyl chloride) there is still considerable doubt as to whether the pure liquid under oes any self-ionization since direct evidence of such a process is always extremely hard to find. However, it is instructive to see how a new solvent is investigated and the final outcome serves as a warning: premature guesses should not be made as to the mode of self-ionization of a pure non-protonic solvent just to explain or systematize the experimental results.

Bromine Trifluoride This highly reactive compound is prepared directly from the

elements and impurities such as bromine, hydrogen fluoride, bromine fluoride and bromine pentafluoride, removed by a fractional distillation in a steel apparatus at atmospheric pressure. The bromine trifluoride distilling over between 126-128°C (760 mm) is then pure enough to be handled in a quartz apparatus. A final vacuum distillation into the reaction vessel is usually carried out before each experiment.


It is well to bear in mind when discussing the solvent properties of bromine trifluoride that even asbestos reacts with incandescence and typical organic solvents like benzene and carbon tetrachloride explode when exposed to the reagent.

TABLE 8

Physical properties of bromine trifluoride

Melting point, °C Boiling point, °C Specific conductivity, ohm-1 cm-1

Trouton constant Dielectric constant Viscosity, cP

- 9 127

8 x 10-8(15°) 25-3

7-9 (25°) 2-2 (25°)

The specific conductivity of the pure liquid has been explained by assuming a self-ionization:

2BrF3 ^ BrF2+ + BrF4~

while the fact that the conductivity decreases with increasing temperature over the range 15-60°C is thought to be due to one or both of the ionic species being thermally unstable.

The alkali and alkaline earth metal fluorides dissolve readily in liquid bromine trifluoride and from the solutions it is possible to isolate compounds such as KF.BrF3 and BaF2.2BrF3 which, at first glance, might be thought to be simple solvates. However, physical and chemical evidence has shown that they contain the planar tetrafluorobromate(III) ion, BrF4~, and are the base analogues of the bromine trifluoride solvent system. Other fluorides, such as antimony pentafluoride, dissolve in bromine trifluoride, may form the bromofluoronium ion, BrF2

+, and act as the acid analogues of this solvent.

The chemistry in liquid bromine trifluoride, though quite extensive, can be summarized under two main headings. First, the solvent may be used purely as a vigorous fluorinating agent to make a fluoride of the element under investigation; when this


happens the element is invariably oxidized to its highest valency state. The excess of bromine trifluoride is driven off by heating in a vacuum to leave the desired fluoride in a pure state.

Second, the fluoride may be left in solution after its formation and then reacted with another fluoride which is capable of forming a complex with it, often by means of an acid-base neutralization.

Acid-Base Reactions The formation of metal tetrafluorobromates(III) is not confined

to the dissolution of the appropriate metal fluoride in liquid bromine trifluoride. Any alkali or alkaline earth halide or carbonate (or in the case of silver, the pure metal) will dissolve in bromine trifluoride to give tetrafluorobromates(III), an intermediate in the reaction presumably being the metal fluoride formed by solvolysis:

BrF, BrF8 Ag ► AgF ► AgBrF4

BrF BrF Li2COa ^ C02 + £02 + LiF i LiBrF4

Similarly an acid analogue such as BrF2SbF6 may be prepared from almost any other antimony compound besides the penta-fluoride :

Sb203 i 3/202 + SbF5 — i- BrF2SbFe

SbCl5 i FC1 + SbF5 X BrF2SbF6

All the undesired products formed in these solvolytic reactions are volatile and may be removed from the solution by distillation.

As was mentioned at the beginning of this chapter it is possible to carry out a conductometric titration between the base analogue AgBrF4 and the acid analogue BrF2SbF6. This may be accomplished by starting the titration with a known amount of silver metal dissolved in bromine trifluoride. The titration is carried out directly inside a conductivity cell, the conductivity at the


start of the experiment being due to ions derived from silver tetrafluorobromate(III) and from the solvent. As known amounts of antimony trioxide are added to the solution the conductivity drops as the large, less mobile SbF6~ ions replace the smaller BrF4~ ions. After an equivalent amount of antimony trioxide has been added, the conductivity rises with each further addition due to the presence of excess BrF2

+ and SbF6~ ions. When the titration is carried out by adding tin dichloride to a

known amount of silver tetrafluorobromate(III), the minimum in the conductivity curve occurs at a ratio of 2Ag : ISn, since the acid derived from tin is dibasic:

SnCl2 -> SnF4 -> (BrF2)2SnF6

(BrF2)2SnF6 + 2AgBrF4 -> Ag2SnF6 + 4BrF3

On careful analysis of several salts formed by titrations similar to the two described above, it was often found that "bromine" was present as impurity. This was shown to be due, in the main, to solvolysis (which will cause an incomplete reaction between acid and base) by use of the following reaction sequences:

Titanium dioxide dissolves in bromine trifluoride to give (BrF2)2TiF6; this was shown to be the case by treating the solution with nitrosyl fluoride, NOF, and analysing the product, nitrosyl hexafluorotitanate, which was free from "bromine":

(BrF2)2TiF6 + 2NOF -> (NO)2TiF6 + 2BrF3

However, when equivalent quantities of potassium bromide and titanium dioxide were mixed in liquid bromine trifluoride and excess of the solvent removed under vacuum the product analysed as K2TiF6.0-95BrF3. X-Ray powder photographs showed the bromine tiifluoride to be present as potassium tetra-fluorobromate(III) indicating the reversibility of the reaction :

2KBrF4 + (BrF2)2TiF6 -> K2TiF6 + 4BrF,


Moreover, when a pure sample of potassium hexafluorotitanate (made by treating potassium fluoride and titanium dioxide with liquid hydrogen fluoride) was dissolved in bromine trifluoride a product, K2TiF6.llBrF3, was obtained on evaporation of the solvent. X-Ray studies again showed large quantities of the solvolysis product, potassium tetrafluorobromate(III) to be present.

Dissolution of Salts in Bromine Trifluoride We have seen that alkali metal halides and carbonates are

readily converted to the tetrafluorobromates(III). It is interesting to study the products which have been obtained from the solvolysis of their other salts.

When an oxyacid salt is used which contains an element capable of forming a complex fluoro-anion, the product o solvolysis is the alkali metal derivative of that fluoro-anion. For example, on dissolution of borax, a sodium borate of formula Na2B4O7.10H2O, in bromine trifluoride the following reaction sequences occur:

BrF Na2B4O7.10H2O ^ 2NaF + 4BF3 + 17/2 0 2 + 20HF

NaBrF4 + BrF2BF4-> NaBF4 + 2BrF3

On evaporation of the solution under a vacuum all of the products are volatile except sodium tetrafluoroborate which remains behind in a state of high purity. If, as in the case of sodium ortho-phosphate, Na3P04, all the solvolysis products are not volatile, then the complex alkali metal fluoride left on evaporation of the bromine trifluoride is impure :

Na3P04 + BrF3 -> 2NaBrF4 + NaPF6 + 2 0 2

In this case the contaminant of the sodium hexafluorophosphate is the involatile sodium tetrafluorobromate(III).

Solvolysis of alkali metal sulphates in bromine trifluoride stops when only one sulphur-oxygen bond has been attacked and the corresponding fluorosulphonates are produced:

K2S04 + BrF3 -> | 0 2 + KS03F + KBrF4


The alkali metal nitrates act as sources of both metal and nitronium (N02

+) tetrafluorobromates(III) : KNO3 + BrF3 -> KBrF4 + N02BrF4 + i 0 2

and on addition of an acid (for example BrF2BF4 in the form of boric acid, B(OH)3) both potassium and nitronium tetrafluoro-borates are produced: KBrF4 + N02BrF4 + BrF2BF4 -> KBF4 + N02BF4 + 4BrF3

It is possible to prepare pure nitronium derivatives by mixing dinitrogen tetroxide and an acid in bromine trifluoride:

N204 + BrF3 -> 2N02+ + 2BrF4~ N204 + Ρ2Οδ + BrF3 -> 2N02PF6

N204 + 2Au + BrF3 -> 2N02AuF4 nitronium tetrafluoroaurate

while the closely related nitrosyl (NO+) complexes are formed by reacting nitrosyl chloride, NOCI, with the required acid in bromine trifluoride :

NOCI + Au + BrF3 -> NOAuF4

2NOC1 + Ge02 + BrF3 -> (NO)2GeFe nitrosyl hexafluorogermanate

Formation of Transition Metal Fluorides Liquid bromine trifluoride has found considerable use in the

preparation of transition metal fluorides, many of which are difficult to obtain in a pure state by other methods. Gold, for example, readily dissolves in bromine trifluoride with the evolution of bromine; evaporation of excess solvent at 50°C under a vacuum leaves a brown-coloured powder of composition AuF6Br, which on heating to 120°C loses bromine trifluoride and leaves behind pure auric fluoride:

Au + 2BrF3 -^ BrF2AuF4 + Br2

120° BrF2AuF4 ► AuF3 + BrF3

NON-PROTONIC SOLVENTS 93 The acid fluorobromonium tetrafluoroaurate, BrF2AuF4, can

be neutralized by a base such as silver tetrafluorobromate(III) to give a precipitate of silver tetrafluoroaurate :

AgBrF4 + BrF2AuF4 -> AgAuF4 + 2BrF3

which is done most simply by adding a 1 : 1 mixture of the two metals to the solvent

Ag + Au + BrF3 -> AgAuF4

Similar reactions can be carried out using tantalum, niobium, platinum, palladium and ruthenium derivatives, e.g.

PtCl4 -► PtF4(BrF2)2 —δ-> Ag2PtF6 | heat

PtF4

Nb -* NbFeBrF2 4* KNbFe | heat

NbF5

An interesting effect noted on the dissolution of many oxides in bromine trifluoride is that the combined oxygen is often released quantitatively as gaseous oxygen which can easily be collected and measured in a conventional vacuum system. This reaction might well lend itself to the direct determination of oxygen in oxides, a procedure which is very difficult to accomplish by normal analytical methods. Typical examples of this behaviour are

Se02 + BrF3 -* SeF6 + Oaf U03 + BrF3-^ UF6 + 3/2 02f I205 + BrF3 -> 2IF5 + 5/2 02f

However, in certain cases not all the oxygen is released and an oxide-fluoride of the element results

V A + BrF3 -> 2VOF3 + 3/2 02f


The Use of Other Halogen Compounds as Solvents Other liquid interhalogen compounds (and also iodine) have

been studied as ionizing solvents, for example, iodine penta-fluoride, iodine monochloride and iodine monobromide. They all exhibit an electrical conductivity in the range 10~4 to 10~5 ohm-1

cm-1 in the liquid phase which has prompted theories of self-ionizations such as

2IF5 ^ IF4+ + IF6-

2IC1 ^ 1+ + IC12-

212 ^ 1+ + I3-

from which it is possible to recognize acid and base analogues and also typical neutralization reactions

KC1 + ICI -> K+IC12- (base)

SbCl5 + ICI -> I+SbCl6- (acid)

KIC12 + ISbCl6 -> KSbCl6 + 2IC1 base acid salt solvent

Sulphur Dioxide Liquid sulphur dioxide is important historically since it was one

of the first non-aqueous solvents to be studied in detail. The commercially available material, manufactured by roasting sulphur or metallic sulphides in air, contains only small quantities of sulphur trioxide and water as impurities. The sulphur trioxide and most of the water can be removed by passing the sulphur dioxide gas through concentrated sulphuric acid, whilst the rest of the water is removed by passing the gas over phosphorus(V) oxide. The melting point of the solvent may then be taken as a criterion of purity.

The low boiling point of sulphur dioxide demands either the use of pressure vessels, when work is carried out in closed systems at room temperature, or that reactions must be studied at low


TABLE 9

Physical properties of sulphur dioxide

Melting point, °C Boiling point, °C Dielectric constant

Specific conductivity, ohm-1 cm-1

Viscosity, cP

Vapour pressure, atm

1 -

{ {

-75-5 -10-2

17-3 (-16-5°) 24-6 (-68-8°)

- 5 x 10-7(-65°to0°) 0-39 (0°) 0-43 (-10°) 1-53(0°) 3-23 (20°)

temperatures in, for example, a vacuum system similar to the type used for liquid ammonia research. The low dielectric constant will lead to reduced solubility of ionic compounds when compared to the water system and when solubility does occur the extensive formation of ion-pairs in the solution will be expected, as evidenced by low equivalent conductance of salts in liquid sulphur dioxide.

Those ionic compounds which are soluble have, in the main, large ions such as iodide or thiocyanate, which imply low lattice energies. An exception to this rule is lithium fluoride and it is possible in this case that the small, highly polarizing lithium ion has a high solvation energy which can offset to a certain extent the large lattice energy of the salt.

Whilst correlating with the relative insolubility of ionic salts, the low dielectric constant enhances the solubility of many covalent compounds, making liquid sulphur dioxide a valuable solvent for boron trichloride, thionyl derivatives, iodine mono-chloride and bromine, all of which are miscible with the solvent.

For many years the specific conductivity of liquid sulphur dioxide was thought to indicate the self-ionization :

2S02 ^ S02+ + S0 32 -

However, this "sulphito-theory" has been challenged recently mainly on the evidence of tracer studies using 1 80 and radioactive

96 NON-AQUEOUS SOLVENTS IN INORGANIC CHEMISTRY 35S. For example, there is no exchange between thionyl chloride and liquid sulphur dioxide, i.e. thionyl chloride does not behave as an acid in sulphur dioxide since if it gave cations characteristic of the solvent, rapid exchange of both sulphur and oxygen would occur :

2S02 ^ S02+ + S032-

% SOCl2 ^ S02+ + 2C1-

Sulphur trioxide does not exchange sulphur with liquid sulphur dioxide but on the addition of a sulphite to the solution exchange of sulphur between trioxide and dioxide occurs readily. This experiment demonstrates very effectively that the concentration of sulphite ions arising from the ionization of pure sulphur dioxide must be exceedingly small. In fact, it is now thought that liquid sulphur dioxide exhibits no self-ionization at all. It is of interest to point out also that sulphites cannot be recovered unchanged from liquid sulphur dioxide because they react with the solvent to form pyrosulphites which contain the assymetric ion [03S—S02]2-.

Early workers assumed that reactions such as

PC15 + S02 — ^ POCI3 + SOCl2

60° wci6 + so2 —► woci4 + soci2

were solvolytic in nature. This cannot be the case if solvofysis is assumed to involve the participation of ions derived from the self-dissociation of the solvent.

Solvate Formation Many compounds, including salts, form coloured solvates with

sulphur dioxide, the amount of co-ordinated solvent varying between 0-5 and 14 moles per mole of solute, e.g. NaI.4S02, Li.I.2S02 and CsSCN.0-5SO2. In general, the sulphur dioxide solvates are thermally unstable and have dissociation pressures of


several centimetres of mercury at room temperature. For example, potassium iodide crystallizes from sulphur dioxide solution below — 10°C as the pale red KI.4S02 but when the crystals are held at 0°C they lose sulphur dioxide continuously until pure potassium iodide remains; in other cases the coordinated sulphur dioxide is lost in a step-wise manner as lower, more stable solvates are formed. Potassium iodide slowly reacts with sulphur dioxide at room temperature:

KI + S02 -> K2S04 + S + I2

so that the solvate KI.4S02 is unstable when kept at room temperature even if placed in a sealed tube to stop the loss of sulphur dioxide.

Complex Formation Closely allied to the solvates discussed above are the compounds

MF.S02 formed when alkali metal or quaternary ammonium fluorides dissolve in liquid sulphur dioxide. X-Ray powder photographs show they are not simple solvates but contain the ion S02F~, the salts being isomorphous with the corresponding metal chlorates, MC103. The acid from which these fluorosulphinates are derived, HS02F, is thought to be the unstable compound HF.S02 indicated in the freezing point-composition diagrams of HF/S02 mixtures.

The alkali fluorosulphinates are useful fluorinating agents and in many cases the reactions can be carried out in liquid sulphur dioxide:

AsCl3 + 3MS02F -> AsF3 + 3MC1 + 3S02

CeH5COCl + MS02F -> C6H5COF + MCI + S02

With nitrosyl chloride in liquid sulphur dioxide the fluorosulphinates form nitrosyl sulphinate:

NOCI + MS02F -+ NOS02F + MCI


which reversibly decomposes into nitrosyl fluoride and sulphur dioxide at room temperature.

NOS02F ^ NOF + S02

Like the alkali fluorosulphinates, NOS02F can act as a fluori-nating agent in liquid sulphur dioxide but in certain cases the reaction goes several stages further to give nitrosyl complexes : 6NOS02F + Assida -> NOAsvF6 + 2NO + 3NOC1 + 6S02

BC13 + 3NOS02F -> BF3 + 3NOC1 + 3S02

BF3 + NOS02F -> NOBF4 + S02

Liquid sulphur dioxide is an excellent medium for the production of halogeno-complexes such as the hexachloroantimonates. Antimony pentachloride, unlike phosphorus pentachloride, dissolves without reaction in sulphur dioxide. On the addition of tetramethylammonium chloride to the solution there is a ready formation of tetramethylammonium hexachloroantimonate(V) :

Me4NCl + SbCl5 -> Me4NSbCl6

Similar reactions occur with acetyl chloride: CH3COCI + SbCl5 -> CH3COSbCl6

and with nitrosyl chloride: NOCI + SbCl5 -> NOSbCl6

Metathetic reactions between the various halogeno complexes also take place readily in liquid sulphur dioxide :

NOSbCl6 + Me4NPF6 -> NOPF6| + Me4NSbCl6

Nitronium hexachloroantimonate, N02SbCl6, formed by reacting nitryl chloride and antimony pentachloride together in liquid chlorine, gives an electrically conducting solution in liquid sulphur dioxide which can be used to prepare other nitronium compounds by metathetic reactions such as:


N02SbCl6 + Me4NC104 -* N02C104| nitroneum perchlorate

N0 2 SbCl 6 + Me4NBF4 -> N 0 2 B F 4 | nitronium tetrafluoroborate

An interesting series of reactions, which on the sulphito-theory was regarded as demonstrating amphoteric behaviour, is that between aluminium chloride and tetramethylammonium sulphite:

2A1C13 + 3(Me4N)2S03 -> A12(S03)3J + 6Me4NCl the precipitate of aluminium sulphite slowly dissolves in the presence of excess quaternary sulphite to give tetramethylammonium trisulphitoaluminate:

A12(S03)3 + 3(Me4N)2S03 -> 2(Me4N)3Al(S03)3

The addition of thionyl chloride results in the reprecipitation of aluminium sulphite : 2(Me4N)3Al(S03)3 + 3SOCl2-> 6Me4NCl + A12(S03)3| + 6S02

Metathetic Reactions As metal chlorides are insoluble in liquid sulphur dioxide it is

generally arranged in metatheses that one of the products is a metal chloride which then precipitates out during the reaction. A simple example might be the addition of thionyl chloride to a solution of caesium sulphite :

Cs2S03 + SOCl2 -> 2CsCl j + 2S02

Similar metathetic reactions have been used to attempt the synthesis of thionyl derivatives, for example

2AgOOCCH3 + SOCl2 -> 2AgCl + SO(OOCCH3)2 silver acetate thionyl acetate

Thionyl acetate is stable in solution but if the solvent is evaporated off, decomposition sets in to give acetic anhydride and sulphur


dioxide. Thionyl thiocyanate, when formed by the addition of thionyl chloride to potassium thiocyanate solution

SOCl2 + 2KSCN -> SO(SCN)2 + 2KC1 j

can be crystallized out of solution at —80° but begins to decompose rapidly at temperatures above —10° to give polymeric thiocyanogen

2xSO(SCN)2 ~> ^SCN)* + *S + *S02

Thionyl iodide is too unstable to exist even in solution at low temperature:

SOCl2 + KI -> KC1 + [SOI2] -> I2 + S + S02

Many covalent chlorides, which are solvolysed by solvents containing reactive hydrogen atoms, readily dissolve unchanged in liquid sulphur dioxide, making this medium very useful for reactions of the type

SiCl4 + 4KSCN -> 4KC1 j + Si(NCS)4

POCl3 + 3KSCN -> 3KC1J + PO(NCS)3

BCI3 + 3KSCN -> 3KC1J + B(NCS)3

The infrared spectra of all these products are consistent with an isothiocyanato-formulation. Addition of donor molecules, such as pyridine and trimethylamine, to the solution of tris(isothio-cyanato)boron gives typical donor-acceptor addition compounds:

B(NCS)3 + C5H5N -> C5H5N-B(NCS)3

B(NCS)3 + (CH3)3N -> (CH3)3N-B(NCS)3

Dinitrogen Tetroxide Dinitrogen tetroxide is often prepared in the laboratory by

reacting metals, such as copper, with concentrated nitric acid; the manufacture on an industrial scale involves the aerial oxidation of nitric oxide, NO. After distillation to remove the other

NON-PROTONIC SOLVENTS 101 oxides of nitrogen and drying over phosphorus(V) oxide, the solvent is ready for use.

Dinitrogen tetroxide has the lowest dielectric constant of all the solvents studied so far in this book (cf. the dielectric constant of benzene, 2-6). From this we might expect low solubility of inorganic salts, high solubility of co valent compounds and a very low degree of self-ionization. However, much of the chemistry carried out in liquid dinitrogen tetroxide can be understood in terms of slight self-ionization:

N204 ^ N O + NO3-

TABLE 10

Physical properties of dinitrogen tetroxide

Melting point, °C Boiling point, °C Dielectric constant Conductivity, ohm-1 cm- {

-11-3 21-1

2-4 (18°) 1 x IO"12 (17°)

2-5 x 10-18 (21°)

for which the ionic product is about 10~24 g equiv2/l2. The rapid exchange of 15N between the solvent and labelled ammonium nitrate, N(CH3)4

15N03, is probable confirmation of this ionization.

In both the solid and liquid states dinitrogen tetroxide has a structure with all N—O bonds of the same length; resonance forms such as

O O-N N

O- O

contribute to the structure, and it is clear that the dissociation into nitrosyl and nitrate ions cannot take place by any single-step


mechanism. It has been postulated, therefore, that the dissociation takes place in two stages, the first of which is the heterolytic cleavage of the nitrogen-nitrogen bond

N204 ^ N02+ + N02-

followed by the transfer of an oxygen atom from the positive to the negative ion

N02+ + N02- ^ NO+ + NO3-

The position is complicated even further due to the homolytic dissociation which occurs at the same time:

N 2 0 4 ^ 2 N 0 2 Dissociation = 0-13 per cent at 21°C

The paramagnetic nitrogen dioxide molecules impart a deep brown colour to the normally pale yellow solvent.

Very few simple inorganic salts are in fact found to be soluble in dinitrogen tetroxide so that acid-base relationships are relatively unimportant; of the co valent compounds studied, bromine is miscible whilst iodine and sulphur are readily soluble.

Complex Formation When inorganic salts do dissolve in liquid dinitrogen tetroxide

it is usually found that complex formation has occurred. A typical example of this behaviour occurs on the dissolution of zinc nitrate in dinitrogen tetroxide. On evaporation of the solvent, a solid of composition Zn(N03)2.2N204 remains from which the adhering dinitrogen tetroxide can be removed by heating. However, the compound is not a simple solvate but is a nitrosyl salt, (NO)2Zn(N03)4. The addition of ethylammonium nitrate to a solution of this nitrosyl salt results in the formation of ethyl-ammonium tetranitratozincate which can be regarded as a rare example of an acid-base neutralization in liquid dinitrogen tetroxide:

(NO)2Zn(N03)4 + EtNH3N03 -> (EtNH3)2Zn(N03)4 + 2N204 acid base salt solvent

NON-PROTONIC SOLVENTS 103 A similar example of complex formation is found when uranyl nitrate, U02(N03)2, dissolves in dinitrogen tetroxide to give the nitrosyl complex, NO[U02(N03)3].

The formation of these nitrosyl compounds represents cases in which the self-ionization of dinitrogen tetroxide, N204 ^ NO+ + NO3-, is enhanced by the continued removal of nitrate ions by the formation of complex anions.

In other cases, the presence of a strong electron acceptor such as boron trifluoride was thought to arrest the self-ionization of dinitrogen tetroxide at the stage N204 ^ N02+ + N02" by the production of nitronium complexes

3BF3 + 2N204 -> N02[ONOBF3] + N02[N(OBF3)2]

A recent careful study of the Raman and infrared spectra of the products shows, however, that these are principally nitrosonium and nitronium tetrafluoroborates formed, possibly, by the reaction

3N204 + 8BF3 -> 3NOBF4 + 3N02BF4 + B203

Iron pentacarbonyl, Fe(CO)5, loses all 5 moles of carbon monoxide on solvolysis by liquid dinitrogen tetroxide and on evaporation of the excess solvent a complex of ferric nitrate remains :

Fe(CO)5 + N204 -> 5CO + Fe(N03)3N204

The infrared spectrum of this yellow solid indicates it is a nitrosyl derivative NO[Fe(N03)4] and this has been corroborated by the precipitation of other salts such as (CH3)4NFe(N03)4, e.g. by the reaction

(CH3)4NC1 + 4N204 + FeCl3 -> (CH3)4NFe(N03)4 + 4NOC1

When heated in a vacuum the NOFe(N03)4 loses dinitrogen tetroxide and a sublimate of unknown structure, Fe(N03)3. 0·9Ν2Ο4, is formed.


Dinitrogen Tetroxide-Solvent Mixtures Though liquid dinitrogen tetroxide is a very poor solvent when

used alone, its chemical usefulness can be greatly increased by adding other suitable co-solvents to it. The technique of mixing liquids to obtain a solvent whose properties have been "tailored" to suit a particular problem is not often encountered in inorganic chemistry. Liquid dinitrogen tetroxide forms a useful model with which to study this effect because the dinitrogen tetroxide molecule can dissociate in several different ways dependent on the properties of the added co-solvent.

The addition of nitromethane, CH3N02 (dielectric constant, e = 37) to liquid dinitrogen tetroxide produces a large increase in the electrical conductivity due to the enhanced self-ionization

N204 ^ NO+ + NO3-

caused by the higher dielectric constant of the solvent mixture compared to pure dinitrogen tetroxide. Copper metal, though not reacting with liquid dinitrogen tetroxide, readily evolves nitric oxide when nitromethane is present in the solvent

Cu + 2N204 -> Cu(N03)2 + 2NOf

the reaction proceeding most rapidly in those solvent-mixtures which exhibit maximum electrical conductivity (90 mole per cent nitromethane), i.e. when the concentration of nitrosyl ions is at a maximum.

A similar increase in the rate of dissolution of copper metal is also noticed on the addition of certain organic solvents, such as ether or ethyl acetate, to the dinitrogen tetroxide. These liquids do not have a high dielectric constant so that the increased rate of reaction must be due to a different mechanism than that described above for nitromethane. It has been found that the organic compounds complex with nitrosyl ions (and also in several cases with the undissociated dinitrogen tetroxide molecules) to produce a significant increase in the ionization of dinitrogen tetroxide

NON-PROTONIC SOLVENTS 105 «EtOOCCH3 + N204 ^ (EtOOCCH3)„N204

ethyl acetate ^ (EtOOCCH3)nNO+ + NO3-

When dinitrogen tetroxide is diluted with glacial acetic acid, the rate of dissolution of metallic copper is again increased but the reaction product obtained depends on the composition of the mixture. Thus solvents rich in dinitrogen tetroxide give the deep blue solvate Cu(N03)2N204 which loses the adhering solvent on heating:

85° Cu(N03)2N204 > Cu(N03)2 + N204f (or 2N02)

If the solvent mixture contains about 40 mole per cent of acetic acid, however, the product is Cu(OOCCH3)2.HOOCCH3, from which anhydrous cupric acetate can be obtained by heating. As neither glacial acetic acid nor pure dinitrogen tetroxide attack copper, it appears that each liquid must increase the reactivity of the other, possibly by complex formation.

In liquid bromine trifluoride, dinitrogen tetroxide ionizes in a completely different manner to give only nitronium (N02

+) ions and the solutions can be used to prepare nitronium derivatives by reaction with suitable fluoroacids :

Au + N204 + BrF3 -> N02AuF4

B203 + N204 + BrF3 -> N02BF4

These reactions possibly proceed via the formation of an unstable base, N02BrF4, by the nitronium ion

N204 + BrF3 -> N02F -> N02BrF4

N02BrF4 + BrF2BF4 -> N02BF4 + 2BrF3

Raman spectra of dinitrogen tetroxide-concentrated nitric acid mixtures indicate extensive ionization to the nitrosyl cation

N204 ^ NO+ + N03-

therefore when boron trifluoride is added to the mixture, nitrosyl


tetrafluoroborate is one of the products

BF3 + HNO3 - J Î A BF4-

BF4- + NO+ -> NOBF4j

However, when boron trifluoride is added to a mixture of concentrated sulphuric acid and dinitrogen tetroxide, the reaction product contains both nitrosyl and nitronium tetrafluoroborates, the latter compound appearing due to the instability of nitrate ions in concentrated sulphuric acid:

N204 ^ NO+ + NO3-

NO3- + H2S04 -> N02+ (see page 52)

NO+ + N02+ + 2BF4- -* NOBF4 + N02BF4

Phosphorus Oxychloride Phosphorus oxychloride, made by reacting together phosphorus

pentachloride and phosphorus pentoxide, contains phosphoric acid and hydrogen chloride as the main impurities, both of which are removed by distillation. Before accurate physical measurements are taken in phosphorus oxychloride, the solvent is often distilled from either potassium metal or activated silica gel directly into the apparatus, care being taken to exclude moisture since the oxychloride is readily hydrolysed to hydrogen chloride.

The liquid range of phosphorus oxychloride besides being usefully wide makes investigation possible at normal temperature and pressure. As might be expected from the low value of the dielectric constant, few inorganic salts are soluble although many covalent compounds which hydrolyse rapidly in water can be studied in this medium.

The low specific conductivity of phosphorus oxychloride suggests a small, but significant, self-ionization in the pure liquid:

POCI3 ^ POCl2+ + Cl-

NON-PROTONIC SOLVENTS

TABLE 11

Physical properties of phosphorus oxy chloride

107

Melting point, °C Boiling point, °C Dielectric constant Specific conductivity, ohm-1 cm-1

Viscosity, cP

1-25 105-8 (753 mm)

13-9(22°) 2 x IO"8 (20°)

111 (22°)

with the added possibility of either one or both ions being solvated, POCl3.POCl2+ or POCl4-. Though there appears to be no direct evidence in support of the above self-ionization, mobility of chloride ions in liquid phosphorus oxychloride has been demonstrated by exchange reactions using a soluble ionic chloride, like tetramethylammonium chloride labelled with radioactive chlorine (36C1). The rapid chlorine exchange between solute and solvent, though not proving the above self-ionization occurs in phosphorus oxychloride, is consistent with such a process.

In other solvent systems it has been the practice to study solvate-formation, since in many cases solvates contain ions characteristic of the parent solvent, thus lending strong support to the postulated self-ionization mechanisms. Many examples of phosphorus oxychloride solvates are known, e.g. BCl3POCl3, AsCl3POCl3, FeCl3POCl3, (TiCl4POCl3)2, SbCl3POCl3, SbCl5POCl3 and GaCl3POCl3. Unfortunately, all the solvates which have been subjected to physical measurements, such as Raman and infrared spectroscopy and X-ray diffraction studies, are not ionic. They are normally donor-acceptor complexes with oxygen acting as the donor atom:

ci CL CL

cr ^ C L

opcL3

a /

CL

ΛΚ

CL

CL

^CLv

"CL

OPCL3

CL

ÔL


However, in a few cases (e.g. AsCl3POCl3 and SbCl3POCl3) infrared spectra indicate that the solvates are merely held together by dipole-dipole interactions even when the molecules undergoing solvation can function as electron acceptors.

Conductivity measurements on solutions of these solvates in phosphorus oxychloride confirm that they are largely undisso-ciated. However, the conductivity, though being of a low value, is still much higher than that of the pure solvent and has been accepted as proof of a slight ionization: e.g.

BCl3POCl3 ^ POCl2+ + BC14 . > dissociation constant = 10-· vl· SbCl5POCl3 ^ POCV + SbCl,

Since the ionization processes produce cations characteristic of the solvent, the solutions will be weakly acid and capable of titration against a soluble chloride; for example the neutralization reactions

Et4NCl + POCl2BCl4 -> Et4NBCl4 + POCl3

Et4NCl + POCl2SbCle -> Et4NSbCle + POCl3

can be followed either potentiometrically or conductometrically and the products isolated for analysis.

The neutralization of POCl3FeCl3 with tetraethylammonium chloride has been studied spectrophotometrically. When an equivalent amount of chloride has been added to the ferric complex the resulting yellow solution exhibited the spectrum characteristic of the tetrachloroferrate(III) ion, FeCl4~. The same spectrum is exhibited by very dilute solutions ( < 0-01 molar) of ferric chloride in phosphorus oxychloride and it might be thought that essentially complete ionization of the FeCl3POCl3 complex has occurred:

POCl3FeCl3 -> POCl2+ + FeCl4~

In more concentrated solutions, the solvent molecules are shown to be covalently co-ordinated to the ferric chloride as in the solid solvate.

NON-PROTONIC SOLVENTS 109 Unfortunately later workers proved, again spectrophoto-

metrically, the presence of tetrachloroferrate(III) ions in solutions of anhydrous ferric chloride in triethyl phosphate, OP(OC2H5)3. This evidence shows that, as it is obviously not necessary to obtain a chloride ion from the solvent to form FeCl4~ ions, the self-ionization

POCI3 ^ POCl2+ + Cl-

need not be invoked to explain the reactions of ferric chloride (and possibly many other compounds) carried out in phosphorus oxychloride. For the specific case of ferric chloride, the equilibria present in the solution are now considered to be

FeCl3 + CI3PO -* Cl3FeOPCl3

^ [FeCl3-3(OPCl3)irfl]+ + *FeCl4-%

Fe(OPCl3)w3+ + 3FeCl4-

Furthermore, the reaction of tetramethylammonium chloride with either ferric chloride or antimony trichloride, to give the ions FeCl4~ or SbCl4~ can be followed conductometrically in liquid triethyl phosphate thus proving that the chloride ion transfer does not require the participation of solvent ions. This was interpreted as strong support for using Lewis acid-base mechanisms in both solvents.

Liquid phosphorus oxychloride has proved to be a useful medium in which to make complexes of the reactive Group III and Group V halides. The dropwise addition of selenium oxychloride to a boiling solution of phosphorus pentachloride in phosphorus oxychloride results in the immediate formation of selenium tetra-chloride, which reacts further with excess of the pentachloride to produce a white solid complex

SeOCl2 + PC15 — ^ SeCl4 + POCl3

SeCl4 + PC15 — - * SeCl4.PCl5


Other chlorides, such as boron trichloride, tin tetrachloride and auric chloride, form similar complexes all of which give electrically conducting solutions and are thought to be ionic, containing the tetrahedral tetrachlorophosphonium ion, PC14+, e.g. PC14+BC14-and (PCl4

+)2SnCl62-.

Noble Gas Compounds as Ionizing Solvents Recently xenon hexafluoride (m.p. 46°C; b.p. 100°C) has been

studied as a solvent, although work with it is both dangerous (due to the possible formation of the explosive xenon trioxide by accidental hydrolysis) and difficult (e.g. all the metal electrodes yet tried are readily attacked, making conductance measurements impossible). However, caesium fluoride is somewhat soluble at 50-55°C giving caesium heptafluoroxenate(VI):

Cs + XeF6 -> CsXeF7 (yellow)

CsXeF7 _*™*^ ** 5 0 ° c K Cs2XeF8 + XeF6f ' absence of solvent ft ° ft^ stable to 400°C

Rubidium fluoride behaves similarly. Antimony pentafluoride on the other hand gives XeSbFu (m.p. 200°C) which might be formulated as XeF5

+SbFe~. If the self-dissociation of the solvent is assumed to be 2XeF6 ^ XeF5

+ + XeF7~, then XeSbFu and CsXeF7 can be recognized as acids and bases respectively.

Further Reading AUDRBETH, L. F. and KLEINBERG, J., Non-aqueous Solvents, Wiley, New

York, 1953, Chapters 11, 12 and 13. ADDISON, C. C , The Use of Non-aqueous Solvents in Inorganic Chemistry.

Royal Institute of Chemistry Monograph No. 2, 1960. ELVING, P. J. and MARKOWITZ, J. M., Chemistry of solutions in liquid sulphur

dioxide, / . Chem. Educ. 37, 75 (I960). GRAY, P., The Chemistry ofDinitrogen Tetroxide, Royal Institute of Chemistry

Monograph No. 4, 1958. GUTMANN, V., Reactions in some non-aqueous solvents, Quart. Rev. Chem.

Soc. 10, 451 (1956). GUTMANN, V., Ionic reactions in solutions of certain chloride and oxy-

chlorides, / . Phys. Chem. 63, 378 (1959). NORRIS, T. H., Isotopie exchange reactions in liquid sulphur dioxide and

related non-aqueous systems, / . Phys. Chem. 63, 383 (1959).

CHAPTER VI

HIGH TEMPERATURE SOLVENTS

IN THIS chapter we shall consider solvent systems where the liquid range is from about 100 to 1500°C. Inorganic solvents in this range have been used for a long time, but more commonly on an industrial rather than a laboratory scale. Thus, the large-scale extraction of either very electropositive elements (e.g. the alkali metals) or very electronegative elements (e.g. fluorine or chlorine) is carried out by electrolysis of a fused salt, because these elements would reduce or oxidize any low-temperature solvent which might otherwise be suitable. The fused salt may be a simple binary electrolyte such as sodium chloride, and this then represents the simplest type of fused salt solvent; or it may contain other salts (i.e. solutes) added to lower the melting point or to produce more favourable electrochemical conditions. These changed conditions or properties may arise due to solvent-solute interactions, e.g. by formation of complex ions, and clearly such interactions deserve fundamental study.

Two other large-scale processes which may involve reactions in melts are worth mentioning, viz. element displacement reactions and oxide displacement reactions. In the former, one element (usually a metal) is displaced by another, i.e.

MX + M' ^ M + M'X where MX is usually an oxide (or it may be a sulphide or halide) and M' may be carbon or another metal. EUingham and others have calculated the free energy changes, AG9 as a function of temperature for the processes

M X ^ M + X M'X ^ M' + X

111


(where X = oxide or sulphide) for a large number of elements ; these values enable the direction of the displacement reaction to be predicted at any given temperature. Changes of state of reactants or products will produce discontinuities in the AG/T plots; so far as simple melting of the oxides or metals is concerned, these do not usually effect significant changes in the direction of the reaction, but addition of other substances to the melts may be very important. Thus, in theory, the reduction of magnesium oxide by silicon is not predicted — the free energy calculations predict reduction of silica by magnesium — but in fact magnesium oxide is readily reduced by ferrosilicon at about 1500°C, for here silica is removed as molten iron silicate, and the dissolving of the silica reduces its activity sufficiently to reverse the expected reaction. More generally, the removal of traces of oxides and other impurities in the melting or refining of metals, by adding fluxes (silica to form silicates with basic impurities, calcium oxide for acidic oxides such as those of phosphorus) is a process which produces a molten slag; this can be separated from the molten metal if the two are mutually insoluble, have different densities, and if the slag is not too viscous. The successful operation of such processes clearly demands a physico-chemical study of the metal-slag system, e.g. composition and specific heat measurements.

Oxide displacement reactions are exemplified by the displacement of phosphorus(V) oxide by silica from calcium phosphate

Ca3(P04)2 + 3Si02 -> 3CaSi03 + P205

The calcium silicate forms a molten slag, and similar problems of separation occur in such cases as those for metal-slag mixtures. The phosphorus(V) oxide is subsequently reduced to phosphorus by heating with carbon. A later reference is made to this reaction.

The above discussion applies to processes which have been in use for a long time. More recent uses of molten salts, e.g. in nuclear fuel processing, as reactor moderators, or in fuel cells, have stimulated studies of their properties. Advances in spectro-scopic methods have aided the older methods (e.g. cryoscopy, conductivity) of identifying solute species present in molten salt

HIGH TEMPERATURE SOLVENTS 113

solvents, and often much physico-chemical information is available. As yet, however, actual reactions in molten salts have not received extensive study; one reason for this is certainly the considerable experimental difficulties encountered in high temperature work, and we must now consider the experimental methods available.

Experimental Methods

Over the wide temperature range to be considered, wide variations of method exist; rather elaborate methods may have to be used at the higher temperatures. For heating, resistance- or induction-heated furnaces are usually employed, often in the form of a tube of refractory material (horizontally or vertically mounted), in the centre of which the temperature can be controlled with reasonable precision with respect to both uniformity and time. Resistance heating gives better control but induction heating is often simpler and the reaction vessel (e.g. a crucible) can itself be the heating element. Good thermal insulation of the furnace is of course essential in effectively controlling temperature. The latter may be measured by a thermocouple (these can now be used up to 2400°C) or by a radiation pyrometer, or by thermal expansion of a rod of refractory material. The most important choice to be made is that of the refractory material for the reaction vessel. This must obviously be as stable as possible under the conditions of experiment, but at higher temperatures some attack on the containing vessel is almost inevitable; the rate of attack depends upon diffusion of the melt into the refractory and therefore increases with decreasing melt viscosity. Oxides are very often used, e.g. silica (quartz) or alumina, but borides, carbides and sulphides may also be used; graphite and some metals (e.g. platinum) are also useful, especially at temperatures which are not too high. With oxides, reduction by the hydrogen formed from decomposition of water vapour at high temperatures is always a possibility, and rigorous drying of salts to be melted must be undertaken. Oxidation (to a higher oxide) is a less common

F


difficulty. Where it is necessary to preserve a clear surface in the containing vessel, e.g. as an optical window for absorption spectroscopy, the melt may be supported vertically by inert gas pressure above the lower window of the container, so avoiding contact with it; the top surface of the melt itself can then act as the parallel upper surface. In Raman spectroscopy, the melt surface acts as the only "window" necessary, but rigorous drying is still essential to prevent turbidity in the melt.

Aspirator bulb

Refractory brick

A K

-Steel tube

-Copper sampling tube

FIG. 12. Metal sampling pipette.

For density measurements, the change in weight of a float immersed in the melt may be measured, or a pyknometric method used. Viscosity can be determined by a rotating cylinder method, and vapour pressures by the Knudsen effusion technique. Cryo-scopic measurements are made by using the same general type of freezing-point apparatus as is used at ordinary temperatures, with a thermocouple to record the temperature changes. Careful

HIGH TEMPERATURE SOLVENTS 115 temperature control of the cooling jacket is essential to avoid excessive supercooling.

The measurement of quantities relating to chemical change is of interest, e.g. the concentration of a constituent at a given time. A common method here is to use quenching; the reaction vessel, or some part of it, is cooled rapidly by either dropping into a suitable liquid (e.g. mercury) or by a blast of cold air, and the solid dissolved in a conventional solvent and analysed by ordinary methods. In the case of a molten metal covered by (e.g.) a molten layer of silicate slag, a metal pipette may be used (Fig. 12). This consists of a steel tube with an aspirator bulb at the upper end; a copper sampling tube at the lower end is greased before insertion into the melt, and vaporization of the grease keeps the end clear for entry of the metal as the pipette is lowered into the reaction vessel. The metal sample is sucked into the copper tube by using the aspirator and the pipette is withdrawn and quenched ; the metal sample can then be taken out of the copper sampling tube without difficulty.

Methods used to study metal solubilities in fused salts are of interest. If the metal concerned has a cation which is that of the fused salt, electrolysis of the latter is often the best method of obtaining the metal-salt system, since production of the metal in situ avoids possible oxidation or contamination before addition to the melt. After being brought to equilibrium at a given temperature, the metal-fused-salt mixture can be quenched as a whole and analysed; but some separation of metal from salt invariably occurs even with drastic quenching. To overcome this, devices such as those shown in Fig. 13 may be used. The container here is a capsule of stainless steel, sealed after filling (by welding), and with thermocouples spot-welded along its length. It is heated horizontally in a furnace which can be rocked to ensure salt-metal equilibrium; the tube is then tilted to a vertical position to separate the phases, the steel ball being in the side-arm. Further tilting allows the ball to fall from the side-arm and trap a sample of the salt in the bottom of the container (b) ; where it remains until quenched and analysed. A tube of slightly different shape


may be used to sample the metal which is decanted into the side-arm by tilting (c). These containers have been used to determine the solubilities of alkali metals in their salts. In another method, use is made of the fact that the potential of a carbon electrode immersed in a melt depends upon the concentration of dissolved metal. A refractory metal electrode immersed in the liquid metal serves as reference electrode. This method has been used to determine the solubilities of cadmium and lead in their respective chlorides.

r^ r^\ r*~\

Metal

Salt

(a) (b) (c)

FIG. 13. Stainless steel capsules for studies of alkali metals and their salts at temperatures up to 1050°C.

Electrochemical measurements in molten salts may be made in borosilicate glass, silica or mullite vessels up to 1000°C; above this temperature few materials are sufficiently inert, and above 1600°C only conductance and transport have been studied. The necessity of using a closed system (to prevent access of carbon dioxide or water vapour to the melt) presents problems in the

HIGH TEMPERATURE SOLVENTS 117 mounting of electrodes, and special seals must be used. Conductance measurements at lower temperatures are made straightforwardly, often using platinum electrodes; for higher temperatures (e.g. in molten silicates and borates) measurements using containers and electrodes of molybdenum have been used up to 1800°C, and electrodes of molten platinum have been employed for molten oxides. Platinum dissolves anodically in many melts, and here graphite may have to be used, though graphite can be attacked by electrolytic products (e.g. fluorine attacks it to give (CF)n). A reference electrode, required for electrode potential measurements, is not easy to obtain; those of the type M-MCL,. (molten), joined to the melt by a liquid junction, have two disadvantages, viz. dilution of the MCl by the melt (and hence contamination of the latter), and a variable junction potential. Thus, Ag-AgCl used with an asbestos plug will slowly introduce silver ions into a melt. By using a silver chloride eutectic mixture, better results can be achieved for a silver electrode with a glass diaphragm to prevent diffusion of silver ions, and here no significant junction potential is introduced.

For polarography, the dropping mercury electrode can be used until the vapour pressure of mercury becomes too great — at about 150°C.

A successful electrode for voltammetry in melts up to 600°C is shown in Fig. 14; a tungsten microelectrode is sealed into matching borosilicate glass and is provided with a gas-flushing device which alternately makes and breaks contact of the electrode with the melt; at bubble rates of > 30 per minute reproducible "polarograms" are obtained.

Chromatography in molten salts has been carried out on an alumina column; solutions of (e.g.) anhydrous transition metal chlorides in an alkali metal nitrate eutectic melt form well-defined bands on such a column and can be eluted as chloro-complexes by using the same melt, with an added alkali metal chloride. Ion-exchange using a zeolite and molten nitrates has been carried out; the sodium ions from sodium nitrate on the zeolite will undergo partial exchange with the T1(I) ions of thallous nitrate.


In the laboratory, there are clearly very great advantages in keeping the temperature as low as possible when fused salt systems are studied. Some pure salts have low melting points, e.g. ammonium nitrate (170°C) and potassium thiocyanate (173°C). Other inorganic metal compounds, with a character intermediate between true salts and covalent compounds, notably the halides of mercury(II), are also of low melting point, but are often distinctive in their fused state from the true salt halides.

-Parallel ground glass joints

Inert gas

-Tungsten wire

Ground flat and flush

FIG. 14. Electrode for voltammetry in melts up to 600°C.

Some of the alkali or alkaline earth metal nitrates or chlorides have many advantages for physico-chemical studies; they do not readily decompose, or attack the container, they are non-hygroscopic and they are easily obtained in a pure and anhydrous state. The difficulties arising from their relatively high melting points can often be avoided by using eutectic mixtures, a few of which are shown in Table 12.


TABLE 12

Eutectic mixtures

Salts

L1NO3-KNO3 NaN03-KN03 LiCl-KCl LiN08-NaN08-KN08

Melting point,

125 218 350 120

The Structure of Fused Salts and Oxides Simple Salts and Oxides

The fused alkali metal halides have the simplest structure and will be discussed first. They obey Faraday's laws of electrolysis and exhibit high electrical conductivity (for sodium chloride, the equivalent conductance at 908°C is 152-5 ohm^cm^quiv 1

compared with 12-65 ohm-1cm2equiv-1 for a solution at infinite dilution). These facts clearly indicate that the melt is composed of ions, and this is confirmed by X-ray and neutron diifraction studies which also show that only small structural changes in the immediate neighbourhood of a given ion occur on melting. Thus, the co-ordination number for nearest neighbours (cation-anion) falls somewhat, from 6 to about 4, but the cation-anion distance actually decreases slightly; cation-cation and anion-anion distances may increase or decrease (Table 13).

TABLE 13

Interionic distances and co-ordination numbers (C.N.)

Salt (AB)

CsCl

Nal

State

Solid Liquid Solid Liquid

A-B

(A) 3-57 3-53 3-35 3-15

C.N.

( A - B ) 6 4-6 6 40

A-A or B-B

(A) 505 4-87 4-74 4-80


Long-range changes are more significant; the molar volume of most alkali halides increases by 20-25 per cent on melting, and hence it must be assumed that long-range order is lost and that extra free volume appears, as holes or vacant lattice sites, in the melt. A small number of ion-pairs may exist but in general the alkali halides are properly described as entirely ionic, and there are no covalent or molecular species present to make the ionic charges more diffuse.

FIG. 15. Cation sites about the ions NOs~ or COs2-.

When the other common alkali metal salt melts, e.g. the nitrates, are examined, the situation is slightly different because the anion is no longer spherical. Studies of the vibrational spectra of molten nitrates suggest that the nitrate ion is a planar, disk-like entity with a volume about equal to that of the chloride ion. In some nitrates, loss of degeneracy of certain vibrational modes in the spectra may be attributed to contact between anion and cation at certain preferred positions. X-ray and neutron diffraction studies support this idea, and indicate that positions X and Y (Fig. 15) are most probable for the sodium ions around the nitrate ion in molten sodium nitrate. Studies of the ultraviolet spectra of


several alkali nitrates again indicate shortening of the anion-cation (nearest neighbour) distances on melting. Similar considerations apply to molten carbonates and sulphates; here the cations (except Li+) tend to occupy corner sites (e.g. analogous to Z in Fig. 15).

Alkaline earth metal salts are generally similar in behaviour to the alkali metal compounds, showing good conductance when fused; but some of the halides of Group lib (Zn, Cd, Hg) show only low conductivity in the molten state. Thus, zinc chloride is almost a non-conductor near its melting point (262°C) but shows appreciable conductance on heating to ^600°C; cadmium iodide, on the other hand, has a reasonably high conductance at its m.p. (388°C) and this increases sharply with rise of temperature. Mercuric bromide is almost non-conducting over the range from its m.p. (238°) to 320°C. Many of these Group lib halides possess layer-type structures in the solid state (e.g. cadmium chloride contains layers of linked CdCl6 groups), and on heating breakdown of the lattices may occur to give either simple ions, complex ions, or discrete co valent molecules. The vibrational spectra of molten mercuric chloride and bromide suggests only the presence of linear X—Hg—X species (X = Cl, Br); but the small electrical conductance, and evidence from the heat and entropy of fusion, suggest presence of ionic species HgX+, HgX3~ and possibly HgX4

2~, and these may be attributed to a small degree of self-ionization, e.g.

2HgX2 ^ HgX+ + HgX3-

Chemical evidence, to be discussed later in this chapter, supports this view, and it must be concluded that the ionic species are present in concentrations too small to be detectable spectro-scopically. By contrast, Raman spectral studies of molten gallium dihalides GaX2 showed these to be composed of Ga+ and GaX4~ ions before the existence of these ions in the solid GaX2 compounds was verified by X-ray diffraction methods; in consequence, the molten compounds possess high conductivity,


Most other fused metal halides which have been investigated either have behaviour intermediate between that of an alkali halide and a solid covalent halide (e.g. lead(II) chloride) or are volatile, fully covalent halides (e.g. tin(IV) chloride). Knowledge of other molten salts of metals outside the alkali and alkaline earth groups is sparse. Oxide melts may be "network" liquids, e.g. silica, phosphorus(V) oxide, or molecular liquids, e.g. osmium tetroxide (both these classes have low conductivity); or they may be ionically conducting metal oxides M2

I0, or electronically conducting, e.g. FeO; the predominant conduction mechanism is not always easy to determine.

Mixtures of Salts Phase rule studies (melting point-composition curves) for two-

component systems are familiar to most chemists, the presence of maxima or minima indicating compound and eutectic formation respectively. Consider first a mixture of two ionic compounds, e.g. lithium and potassium chlorides. Here, the melting point curve indicates simple eutectic formation and no solid solution, and the equivalent conductance-composition curve also shows a minimum. These observations are explained as follows: if the conductance is assumed to be essentially cationic then, as lithium chloride is added to potassium chloride, the lithium ions, having a higher charge density than those of potassium, are a more effective "glue" for the chloride ions (which form the lattice) and the molar volume decreases. At low lithium chloride concentrations, potassium ions must still be responsible for most of the conductance, and as these find it more difficult to move through the contracting lattice, the conductance decreases, until, at a sufficiently high lithium chloride concentration, lithium ions take over the conductance and this rises finally to the value for pure lithium chloride.

Spectroscopic studies of an equimolar fused mixture of silver nitrate and potassium nitrate give vibrational frequencies intermediate between those of the pure nitrates, and the spectrum retains the same general features as for pure silver nitrate,


Mixtures containing various sodium nitrate-lithium nitrate ratios show a linear variation in vl9 the symmetrical (nitrate) stretching frequency, as the composition is changed. These data suggest no changes other than one of lattice dimensions on mixing.

In mixtures of ionic and low-conductance ("covalent") halides, many experimental observations confirm complex formation. Thus molten mercuric chloride-alkali chloride systems show vibrational lines characteristic of the species HgCl3~ and HgCl42~, these ions being the same as those present in aqueous solutions of the two chlorides. Addition of potassium chloride to pure fused mercuric chloride therefore results in a reduction of the intensity of the stretching frequency vl9 due to HgCl2 molecules, and increases in the frequencies due to the species HgCl3~ and HgCl42_\ In a solution of mercuric chloride in fused thallous nitrate, however, there is little evidence for ions such as HgCl2N03~ and the mercuric chloride probably remains essentially molecular here. Another mixed chloride system, KC1-CdCl2, has been studied by a variety of methods — spectroscopic, heat capacity, conductance, molar volume, e.m.f.— and all these show the formation of complex ions CdCl3~, CdCl42~" and CdCle4-; the salts KCdCl3 and K4CdCl6 can be crystallized from the melt. Another method, also used with this system, is that of cryoscopic analysis; here, quantities of the order of only 0-1 per cent mole fraction of the added constituent are used, and the effect on the freezing point 7} is determined. We then have

Tf = nKfm where Kf is the freezing-point constant, m the mole fraction of the added constituent and n is the number of "foreign particles", i.e. the total number of particles per added molecule minus the number of common particles produced per molecule of solute. In practice, Tf is plotted against concentration of solute and the value ofn which fits the plot most closely is evaluated. Hence in sodium chloride as solvent, n = 1 for most sodium salts, n = 2 for potassium fluoride, n = 3 for barium fluoride, and so on. A solution of sodium hexafluoroaluminate Na3AlF6 in fused sodium


chloride gives n = 7 when very dilute, but n decreases with increasing solute concentration because of the equilibria.

Na3AlF6 ^ 3Na+ + Al3+ + 6F" (n = 7)

3Na+ + AlFß3- (n = 1)

In fused potassium thiocyanate as solvent, complex formation would be expected; thus nickel bromide as solute gives values of« corresponding to formation of the complex anions [Ni(CNS)4]2~ (green) and [Ni(CNS)6]4_ (blue). The cryoscopic method is valid only in dilute solution, and species which appear only at greater solute concentrations will not be detected.

The investigation of transition metal salts in fused salt solvents has been of much interest in recent years. If such a metal halide MXX is dissolved in an alkali halide MJX as solvent, then complex formation due to the common anion, giving complexes ( M X ^ ^ -

can be expected. The only other species present will be the M1

cations, and these are unlikely to interfere. The situation is therefore more simple than in an aqueous solution, where water molecules may also behave as ligands, replacing X~ ligands; such water ligands may also be acidic and complexes containing OH~ ions may result. The absence of solvent interference enables studies of halo-complexes to be made in fused halide solvents, and their structures and stereochemistry can be interpreted in terms of ligand field theory. Gruen and McBeth, using absorption spectra in the visible and ultraviolet ranges, have made many of these studies. Thus, the spectrum of Co2+ salts in fused LiCl-KCl eutectic indicates that the cobalt is present as the tetrahedral CoCl42~ ion over the whole temperature range accessible to measurement. Evidence for the tetrahedral configuration of the NiCl42~ ion, present in solutions of nickel(II) chloride in LiCl-KC1 eutectic, was obtained by comparing the spectrum of the melt with that of solid C s ^ n ^ N i ^ C ^ . In the salt Cs2ZnCl4, the zinc is known to be tetrahedrally co-ordinated to the four chloride ions, and some of the zinc ions can be replaced by nickel

HIGH TEMPERATURE SOLVENTS 125 without change of structure, i.e. isomorphously. The similarity of melt and solid spectra provided the evidence for tetrahedral co-ordination around the nickel. Examination of the spectrum of Co(II) in fused LiN03-KN03 as solvent showed octahedral (CoiNOg)^4- ions to be present; addition of halide ions X-

produced tetrahedral species, e.g. (CoX2(N03)2)2-, and finally CoX4

2~. In chloride melts, Gruen and McBeth found that tripositive ions (e.g. Ti34", Cr34") form 6-co-ordinate complexes more readily than dipositive ions, suggesting that an electrostatic factor is important in determining the co-ordination number ; however, they concluded that crystal field stabilization energies also play an important part in determining both configuration and co-ordination number.

Recently, several studies of formation constants of complexes or ion-pairs have been made in molten salt solvents. Thus, the degree of association of silver and chloride ions to form ion-pairs Ag+Cl" or complexes AgCl2~ has been determined in fused nitrate mixtures by potentiometric methods. The association constant for Ag+Cl- is found to depend upon the nature of the salt nitrate used as solvent: thus, the stability of Ag+Ch is greater in fused caesium nitrate than in fused potassium nitrate. This observation is explained by the "reciprocal coulomb" effect which predicts coulombic stabilization of the association of the smaller cation with the smaller anion and the larger cation with the larger anion. Similar measurements have been made for lead(II) and cadmium cations with halide ions in nitrate melts, and again the expected size effects operate. An interesting experimental approach here is the use of distribution constant measurements, e.g. for lead(II) chloride distributed between silver chloride and potassium nitrate,

_ (PbCl2)KNO, _ n 7 9

* " (PbCl2)Agci -

K depending on the distribution constant of molecules of PbCl2 between the two phases and on the association constants of Pb2+

and Cl~ ions in the potassium nitrate,


Solutions of Elements in Fused Salts Metals

Observations on metals in their molten salts were made by Davy more than 150 years ago; he noted the production of colour in the melts and the difficulty of recovering metals liberated by electrolysis of alkali hydroxides. This problem of recovery is of major industrial importance, for, in general, only metals dissolving in their own halides show no decomposition of the solvent, and most molten oxy-salts are attacked. Solutions of one metal in the salt of another are rarely sufficiently stable for precise physico-chemical study.

For metals in solution in their own fused halides, there are two extreme possibilities. The first, broadly represented by the alkali and alkaline earth metals, is solution with little solvent interaction. The dissolved metal behaves as a (cation + electron(s) ), the electrons being accommodated in ways rather similar to those in alkali metal-liquid ammonia solutions, i.e. as solvated metallic electrons or substituted for anions in cavities. At the other extreme, broadly represented by the post-transition metals, there is dissolution with definite chemical reaction, leading to the formation of "subhalides", i.e. where the metal exhibits low oxidation states. In some cases, crystallization of a definite sub-halide is possible, in others the low oxidation state is stable only in the melt. Between the two extremes, there are many ill-defined systems and the state of the electrons (solvated or attached to the metal cation) may vary with temperature and composition.

The alkali and alkaline earth metals show varying solubilities in their halides (e.g. K-KC1 at 800°C, 1 mole per cent; Ba-BaCl2 at 900°C, 30 mole per cent). The solutions are usually highly coloured and the absorption spectrum is attributed to an electron with a wave function distributed over several metal ions (i.e. e~(M+)x), similar to an F-centre. Electrical (specific) conductivities increase markedly with increasing metal concentration, but the rate of increase varies widely for different metal-halide systems.

When one alkali metal is dissolved in the halide of another, an

HIGH TEMPERATURE SOLVENTS 127 equilibrium is set up, as shown in Table 14. (A later reference is made to these values.)

Little is known about the nature of solutions of transition metals in their molten halides ; solubilities are generally very low, although nickel dissolves in fused nickel(II) chloride to the extent of 9 mole per cent, but no subchloride is formed. Solubility increases at the ends of the transition series, zinc, cadmium and mercury exhibiting appreciable solubility in their respective halides, and the "subhalides" of mercury, Hg2Cl2 and Hg2I2, are well known. Zinc in zinc chloride gives species (ZnCl)2. In Group V, bismuth dissolves in its chloride to form "bismuth

TABLE 14

Na + KX ^ K + NaX [Na] [KX] [K] [NaX]

X F Cl Br I

t°C

1000 800 800 800

C

0-29 12-2 29 56

monochloride" of approximate composition BiCl. It has now been shown that this compound is made up of Bi95+, BiCl52_ and Bi2Cl82~ polyions. The polycation Bi95+ is isoelectronic with Pb9

5~, found in solutions of lead + sodium in liquid ammonia (p. 22), and the two may also be isostructural.

The formation of lower oxidation states by metals in melts may be assisted by the addition of other substances. Thus, there is some evidence for the formation of the species Cd2

2+ in solutions of cadmium in cadmium chloride, but on addition of aluminium chloride to the melt, A1C14~ ions are formed and the diamagnetic compound Cd2(AlCl4)2 can be obtained. The similar compound Bi3AlCl4 has also been isolated.


The first evidence of the ability of an intermetallic compound to exist in a molten salt was obtained recently; the compound Li3Bi dissolves in molten LiCl or LiCl-LiF mixtures to give a red solution over the temperature range 500-1000°C.

Non-metals Solutions of sulphur in molten salts are often blue due to the

formation of S4 molecules, and spectroscopic studies of the blue fused thiocyanates NaSCN and KSCN have shown that the colour here is due to S4 as impurity. Solutions of iodine in melts are known, and reference to reactions of iodine in melts is made later in this chapter.

Reactions in Fused Salts "Acid-Base" Reactions

It is appropriate to see how far the various acid-base concepts, discussed in the preceding chapters, can be extended to cover reactions in fused salts.

(a) Protonic acids. Proton abstraction by the solvent is usually a necessary prerequisite for acidic behaviour in hydrogen-containing substances. In a melt, the proton affinity of the anion is often low, so that proton-acidic behaviour is restricted to substances where the proton is already associated with some donor molecule to form a cation, e.g. as NH4

+, pyH+ (py = pyridine). Fused "onium" salts do in fact show acidic behaviour, e.g. by dissolving oxides; fused ammonium nitrate dissolves oxides such as CuO, MgO, ZnO, to give the corresponding metal nitrates, ammonia and water. Fused pyridinium chloride also dissolves metallic oxides, and many metals (e.g. Al, Cd, Mg) will dissolve in it with evolution of hydrogen and formation of a metal chloride and pyridine. Ammonium nitrate will dissolve metals but, like nitric acid, may also behave as an oxidizing agent, e.g. with copper:

Cu + 3NH4N03 -> Cu(N03)2 + N2 + 2NH3 + 3H20

HIGH TEMPERATURE SOLVENTS 129 Fused ammonium nitrate will also dissolve some metal carbonates. The old-established reaction in which a hydrated metal chloride is heated with ammonium chloride to give the anhydrous chloride is well known (e.g. the preparation of anhydrous magnesium chloride by fusion of the double salt NH4Cl.MgCl2.6H20). This reaction can be interpreted as the acid ammonium chloride preventing hydrolysis (with consequent basic salt formation) of the magnesium chloride. Halides of the lan-thanides have been prepared by reactions of their oxides with fused ammonium halides, excess of the latter being removed by subsequent heating under reduced pressure.

(b) Non-protonic acids. Reference has already been made (p. 121) to the self-dissociation of fused mercuric halides, e.g.

2HgBr2 ^ HgBr+ + HgBr3-

Mercuric bromide, m.p. 238°C and b.p. 320°C, has a convenient liquid temperature range and with a dielectric constant of 10 and a viscosity about three times that of water, some solvent power can be expected for polar substances. In it mercuric salts behave as acid anhydride analogues, e.g.

Hg(C104)2 + HgBr2 ^ 2(HgBr)C104

and (HgBr)NH2 and (HgBr)SCN behave as acid analogues. Addition of an alkali bromide gives basic behaviour since the equilibrium (e.g.)

KBr + HgBr2 ^ KHgBr3

is well to the right. "Acid-base" reactions are therefore represented by equations such as

(HgBr)C104 + KHgBr3 -> KC104 + 2HgBr2

Many mercurous salts are soluble in molten mercuric bromide as are also silver nitrate and tripotassium orthophosphate; mercurous sulphate, lead(II) perchlorate and the alkaline earth bromides are slightly soluble and most nitrates (other than silver) and sulphates are insoluble.


More recently, "acid-base" reactions have been studied in fused mercuric chloride, e.g.

HgtClO^ + 2NaCl -> 2NaC104 j + HgCl2

and complex-forming reactions have also been investigated, e.g. with lead(II) chloride the reactions

PbCl2 + NaCl -> NaPbCl3

PbCl2 + 2NaCl -> Na2PbCl4

occur, according to the ratio of the reactants. Here, therefore, chloro-solute complexes appear to be formed in preference to chloro-solvent complexes such as HgCl3~ and HgCl4

2-. These reactions do not as yet appear to have important synthetic

applications, perhaps because mercury complexes are so readily obtainable in conventional solvents. Between these mercuric salts and the low-temperature, non-protonic acid-base systems considered in Chapter V are the fused "semi-salts", e.g. arsenic(III) tribromide and antimony(III) trichloride. In the latter solvent, for example, a dissociation

2SbCl3 ^ SbCl2+ + SbCl4-is assumed and "acid-base" behaviour is then illustrated by reactions of the type

MiCl + (SbCl2)(MniCl4) -> Mi(MmCl4) + SbCl3

(Ml = alkali metal, ΜΠΙ = (e.g.) Al) Antimony(III) chloride has also been used as a solvent for the complex-forming reaction

2MIC1 + TiCl4 ^ i M2iTiCl6

(c) "Amortie" acids and bases. In 1939 Lux suggested that acid-base systems could be set up with reference to oxide ions instead of protons, viz.

base ^ acid + x02~~ e.g. CaO ^ Ca2+ + O2-

so42- ^ so3 + o2-


These equilibria were treated quantitatively, in terms of "acid strengths", by Flood, using the constant (a = activity)

#base

A recent determination for the equilibrium C032~ ^ C02 + O2""

in fused KN03-NaN03 gives K= 4 X 10-6, and other determinations give increasing "acid strength" in the series Si03

2~ < B02~ < PO8-

The Lux-Flood classification may be contrasted with the wider Lewis theory of acids and bases, in which an acid is an acceptor and a base the donor of an electron pair. On the latter theory, the oxide ion in the equilibria just discussed would be a base (Table 15).

TABLE 15

Acid

SiOa B203 or B02-s2o7

2-

Base

o2-o2-0 2 -

Product

Si082" or Si04

4~ B08

8-so4

2-

The Lux-Flood and Lewis acid-base approaches must be carefully distinguished; but both (so far as fused salt solvents are concerned) deal with equilibria involving oxide ions, and these equilibria are important in a variety of reactions on the industrial and laboratory scale. The reaction of silica with metallic oxides is responsible for slag formation, and the extent of this reaction governs the ease or otherwise of reduction to the metal, since the activity of the silica is altered (p. 112). In the laboratory, an analogous reaction is that between boric oxide (from borax) and metal salts to give, on fusion, coloured borates (borax bead test). These reactions are rather complicated, but one essential step appears to be the displacement of an oxide (e.g. sulphur trioxide, from a sulphate) by boric oxide to give the borate


anions and this can be explained in terms of the greater (Lux-Flood) "acid strength" of the boric oxide. Again on the large scale, the displacement reaction already mentioned (p. 112)

Ca3(P04)2 + 3Si02 -* 3CaSi03 + P205

can be interpreted as displacement of the weaker Lux-Flood acid "P205" by the stronger acid Si02.

Another group of reactions in melts which involve oxides are depolymerization reactions, e.g. the conversion of the S207

2~ to the S04

2~ ion, the conversion of insoluble polysilicates to simple silicates and the conversion of dichromates to chromâtes. The equilibrium constant for the latter reaction,

Cr2072~ + O2- ^ 2Cr04

2", K = 1-8 X 1012

has been found in molten potassium nitrate as solvent, by poten-tiometric titration of potassium dichromate with sodium monoxide using a copper-copper(II) oxide electrode to determine the oxide ion concentration. A similar method has been used for the equilibrium P207

4_ + O2- ^ 2P043". In another series of experi

ments, in fused NaN03-KN03 mixtures, the depolymerization reaction

S2072" + N03- ^ N02+ + 2S04

2-

was used (by adding a pyrosulphate to the melt) to furnish a known concentration of nitronium (N02

+) ions and so determine the extent of the dissociation

N03- ^ N02+ + O2-

the concentration of O2- being determined coulometrically. The depolymerization of alkali metal polyphosphates by effectively raising the M20/P205 ratio (by fusion with the oxide M20) is well known; but depolymerization can also be effected by nitrate, sulphate, or carbonate anions. The general reaction can be written


0 o 1 I

-P—O—P-I I o o

+ 02-

o 4-0

I o +

o 4-e

o where y + z = x + 2.

At 700°C, alkali metal perchlorates when fused with poly-metaphosphates yield oxygen and chlorine in amounts indicating a decomposition

MiC104 -* |M20 + £C12 + {θ%

and the M20 then effects depolymerization as indicated, the order of depolymerization being Li > Na > K. This decomposition of a perchlorate to (effectively) M20 rather than to MCI implies that the perchlorates are Lux-Flood "bases" with respect to the polymetaphosphate anions as "acids". It is to be noted that there is no oxidation or reduction in this case; by contrast, chromium(III) oxide is oxidized by fused alkali metal nitrates to chromium(VI) (as chromate). Here, therefore, the nitrate ion is both a Lux-Flood "base" (i.e. an oxide donor) and an oxidizing agent (electron acceptor) while the chromium(III) is an "acid" (oxide acceptor) and reducing agent (electron donor).

A solution of a dichromate in fused KN03/NaN03 eutectic is converted to a chromate by the reaction

C r ^ 2 " + 2N03- -* N205 + 2Cr042"

If the nitrate ion is assumed to dissociate to give N02+ and O2"

ions, the reaction can be written

2N03- ^ 2N02+ + 202-

CraOj2- + 202- -> 2Cr042-


Support for this mechanism is obtained by adding a bromide to the fused nitrate-dichromate mixture; the products are then nitrogen dioxide and bromine rather than dinitrogen pentoxide because of the reaction

N02+ + Br- -> N02Br

I N02 + ±Br2

On the Lewis acid-base theory, donor anions other than oxide should function as "bases". Fluorides have been used industrially to decrease the viscosities of silicate melts, presumably by a depolymerization process; the lesser tendency of fluorine to act as a bridging atom (compared with oxygen) implies that fluorides might be more effective depolymerizers than oxides. Alkali metal monofluoro-phosphates have been prepared by fusion of the appropriate fluoride and polymetaphosphate; otherwise little appears to have been done in reactions of fluorides with oxyanions.

Lewis acid-base equilibria of the type A1F3 + 3F- ^ A1F6

3-BeCl2 + 2C1- ^ BeCl42"

TiF4 + 2F- ^ TiF62-

are familiar in ordinary chemistry, but are also of importance in fused systems; the first is used to provide the conducting medium in the electrolysis of bauxite, the second is used similarly to enhance the conductivity of fused beryllium chloride for electrolytic extraction of beryllium, and the third permits the reduction of the TiFe

2~ ion by sodium to the metal, when the TiF4 itself is not reduced. The equilibrium

AICI3 + CI- ^ Alexis more readily studied in a fused chloride melt than in aqueous solution, for in the latter aquo-cations are formed too easily;

HIGH TEMPERATURE SOLVENTS 135 salts of the type MPAICI with tetrahedral A1C14~ anions, have long been known in chloride melts containing aluminium chloride. A melt of cadmium and aluminium chlorides contains not only Cd(AlCl4)2 but the cadmium(I) salt Cd2(AlCl4)2 already referred to; and melts of cobalt(II) and aluminium chlorides produce the bright blue Co(AlCl4)2 in the solid form of which — rather unexpectedly — the cobalt is octahedrally co-ordinated.

Oxidation and Reduction Reactions (a) By metals. The free-energy-temperature diagrams for the

reduction of oxides and sulphides have been referred to at the beginning of this chapter. They show, broadly, that the effectiveness of a metal in reducing an oxide or sulphide is directly related to its electrode potential in aqueous solution. Clearly, no standard electrode potentials E0 for metals in fused salts generally can be determined, because there is no common standard solvent — for example, one can, for metal-ion couples in (say) chloride melts, draw up a table of E0 values and this will not be the same as in (say) nitrate melts. However, tabulated E0 values in chloride melts at 1000°C show a very similar order to those in water and hence the reducing power is much as would be expected from the E0 values in aqueous solution. The alkali metals are the most effective reducing agents, the noble metals the least.

Another way of approaching simple reduction reactions of the type

MX + M' ^ M + M'X is suggested by the data of Table 14 on p. 127, for the reactions of alkali metals with other alkali metal halides. These equilibria seem to be governed by the reciprocal coulomb effect; with fluoride, the position of equilibrium is such that sodium fluoride predominates over potassium fluoride in the melt, whereas with iodide, the potassium iodide predominates.

Although reactions involving the reduction of substances containing simple monatomic anions such as 02"~, S2", or halides, can be discussed in the above terms, reduction of polyatomic


anions is more complicated. The reaction of molten sodium with fused sodium hydroxide is one of some importance; the main reaction is

NaOH + 2Na -> Na20 + NaH

but at the temperature range used, some dissociation of the sodium hydride occurs to give the metal and hydrogen, and in a closed system the latter will react thus:

Na20 + H2 -* NaOH + NaH

The final stoichiometry may therefore be considerably more complicated than that of the main reaction. The reactions of alkali metals with other oxyanions may be violent in fused systems and this is particularly true with nitrates, which oxidize a wide range of metals to their oxides.

The formation of complex ions is as important in fused systems as it is in aqueous systems, as a factor controlling the nature of a reduction. Aluminium has been used to reduce uranium(IV) to uranium(III) in alkali metal chloride melts; with magnesium, reduction proceeds further to metallic uranium. In the U(III) — Al(III) — Q - system produced by reduction of uranium(IV), the uranium exists mainly as (UC16)3~ and the aluminium(III) as AICI4-, A12C17- and other chloro-species.

A solution of iodine in molten LiCl-KCl eutectic is reduced by most metals to iodide, e.g.

I2 + Mg^Mg2+ + 2I-

(b) Other reductions. Some anions can act as reducing agents in fused salts. The solution of iodine in LiCl-KCl just mentioned is reduced also by sulphide ions to give iodide and sulphur. The cyanide ion in fused potassium cyanide has reducing properties which are enhanced by addition of mercuric cyanide; the following changes have been observed.


K4Moiv(CN)8 — ^ (CN)2 + K2Mo"i(CN)5

KCN K2Nin(CN)4 ·. K4Ni°(CN)4

K3Mnni(CN)e 1 ^ ) - K4Mnn(CN)e

The potassium hexacyanomanganate(II) on heating in KCN changes to the tricyano-compound KMnII(CN)3; the latter can also be prepared by the oxidation reaction

Mn + Hg(CN)2 + KCN -> KMnii(CN)3 + Hg

(c) Oxidations. These are commonly effected by either oxidizing oxy-acid anions, e.g. nitrate, perchlorate, or by oxides of metals in higher oxidation states, or by the halogens, especially chlorine. Many metals are, as already noted, oxidized readily (and sometimes violently) to their oxides by fused nitrates. In LiN03-NaN03 at 175°C, uranium(IV) is oxidized to uranium(VI), e.g.

U*+ + 2N03- -> U022+ + 2N02

but if salts, containing the U022+ cation are dissolved in LiCl-KCl

melts, the UC^2* is reduced by the chloride to give U02+ (i.e. U(V)) and chlorine; and U02C1 then, further, yields U02 and more chlorine when heated in vacuo.

Lithium perchlorate (m.p. 236°C) will, when fused and in presence of an ammonium salt (as "acid"), dissolve many metal salts, some of which are then oxidized; thus all uranium salts give UVI03, and when rhenium is added, chlorine dioxide is evolved and a perrhenate is possibly formed. In alkali metal chloride melts, chromium(VI) oxide oxidizes iodine initially to the +1 state and then to iodate; with chlorine, oxidation stops at the +1 state, with formation of iodine monochloride.


Metathetic Reactions Sundermeyer and his co-workers have used metathetic reactions

in fused halide solvents to prepare organo-silicon derivatives, e.g.

R3SiCl + NaN3 Z n ( ^ 0

K C ^ R3S1N3 + NaCl

(R = CH3, CeH5)

LiCl-KCl 2Me3SiCl + K2S ^Γ^ (Me3Si)2S + 2KC1

SiCl4 + 4KSCN a l k*h"e e a r t^ Si(SCN)4 + 4KC1 hahdes

«Me2SiCl2 + 2«MeOH ZnCla KCj» (Me2SiO)n + + 2«MeCl + «H20

The dehydration to give a polysiloxane in the last reaction is a characteristic of fused zinc chloride, and the ZnCl2-KCl solvent also dehydrates (e.g.) alcohols to olefins. In the same solvent, hydrochlorination can be effected, e.g. ethylene to ethyl chloride. The metathetic reactions with thiocyanate in alkaline earth halides can also be used for carbon compounds, e.g.

EtCl + KSCN -> EtSCN + KC1

In the metatheses listed, the tendency appears to be to pair large cations-large anions and small cations-small anions, i.e. the reciprocal coulomb effect operates.

Very few metathetic reactions with simple salts have been reported. In a fused mixture of calcium, potassium and magnesium chlorides as solvent, there is X-ray evidence for the reaction

CaO + MgCl2 -> MgO + CaCl2


Decomposition Reactions Pure crystalline metallic oxides have been prepared by heating

thermally unstable salts of the metals in solution in fused alkali metal salts. Temperatures up to 1000°C have been used; as examples a solution of cupric sulphate in fused potassium fluoride gave cupric oxide, and a solution of nickel(II) sulphate in fused sodium sulphate gave nickel(II) oxide ; for the preparation of zirconium(IV) oxide, a solution of the salt K2ZrF6 in sodium sulphate was used.

Further Reading BOCKRIS, J. O'M., WHITE, J. L. and MACKENZIE, J. D., Physico-chemical

Measurements at High Temperature, Butterworths, London, 1959. DELIMARSKII, IU. K. and MARKOV, B. F., Electrochemistry of Fused Salts

(English translation), Sigma Press, Washington, 1961. JANZ, G. J., SOLOMONS, C. and GARDNER, H. J., Physical properties and con

stitution of molten salts, Chem. Rev. 58, 461 (1958). BLOOM, H., Recent progress in the high temperature chemistry of inorganic

salt systems, Pure Appi. Chem. 7, 389 (1963). BLANDER, M., (ed.) Molten Salt Chemistry, Interscience, New York, 1964. WAIT, S. C. and JANZ, G. J., Vibrational spectra of ionic melts, Quart. Rev.

Chem. Soc. 17, 225 (1963). ADDISON, C. C, The Use of Non-aqueous Solvents in Inorganic Chemistry,

Royal Institute of Chemistry, Monograph No. 2, 1960.

INDEX

Acetic acid acids and bases in 77 complex formation by 74-7 conductometric titrations in

78-81 ionization of 71 oxidation reactions in 83-5 physical properties 70 potentiometric titrations in 81-3 preparation 70 solubilities in 71-2 solvates 73 solvolysis by 73

Alkali metal-ammonia solutions absorption spectra 17 conductivity 16, 17 nature of 18-20 nuclear magnetic resonance

spectra 17 paramagnetism 16 properties 16

Alkaline earth metal-ammonia solutions 16

Amines as solvents 36 titrations in 38

Ammonia, liquid manipulation of 8-10 metathetic reactions in 29-30 oxidation reactions in 28 physical properties 8 solubilities in 10-12 titrations in 31-6

Ammonium ozonide 22 Ammonolysis 12-15

Bromine trifluoride acid-base reactions in 89

as fluorinating agent 88-9 conductometric titrations in 89 physical properties 88 preparation 87 self-ionization 88 solubilities in 88 solvolysis by 91-2 transition metal fluorides in 92

Carbon tetraiodide, reaction with ammonia 14

Dinitrogen tetroxide complex formation in 102-103 dissociation of 101-102 in bromine trifluoride 105 physical properties 101 preparation 100 -solvent mixtures 104 structure 101

Electrochemical fluorination 66 Ethylenediamine, as solvent 37 Eutectic salt mixtures 119

Franklin, acid-base definitions 3, 86

Fused oxides, structure of 122 Fused salts

acid-base reactions in 129-35 anionic acids and bases in 130 complex ions in 121, 123, 125 decomposition reactions in 139 depolymerization reactions in

132-3 141

142 INDEX

Fused salts — continued effect of added substances 112 element displacement reactions

111-112 metathetic reactions in 138 mixtures of 122-5 non-protonic acids in 129-35 oxidation reactions in 137 protonic acids in 128-9 reduction

by metals in 111,135 reactions in 136

solutions of metals in 15,126-7 of non-metals in 128

structure of 119-121

Halides, ammonolysis of 14-15 High temperature solvents

chromatography in 117 containers for 113 electrochemical measurements in

116-117 ion-exchange in 117 metal

sampling 114-115 solubilities 115-116

spectroscopy in 114 temperature control of 113

Hydrogen chloride, liquid acids and bases in 69 conductivity of 69

Hydrogen fluoride acids and bases in 62 conductivity of 59 dehydration of 58 electrolyses in 66-8 manipulation of 58 oxidation reactions in 65 physical properties 57-8 preparation 57 solvates 59-60 solvolysis by 61-2 structure 60

Hydrogen halides, properties of liquid 68

Hydrolysis 2

Interhalogen compounds as solvents 94

Levelling action 44-6 Lewis, acid-base definitions 5, 131,

134 Lowry-Br0nsted, acid-base theory

4 Lux-Flood, acid-base definitions 5,

130-1

Mercuric halides reactions in 129-30 self-ionization of 121, 129

Metal-ammonia solutions fission of

alkyl groups by 25-6 atoms or groups by 26-8 hydrogen atoms by 25

reduction of covalent compounds 24-8 elements 21-2, 33-4 salts 23-4, 35-6

Noble gas compounds as solvents 110

Phosphorus oxychloride complexes in 109 physical properties 107 preparation 106 self-ionization 87, 106-107 solvates 107

Proton affinity 4 Protonic solvents, self-ionization 44

Self-ionization 3 Solubility, factors affecting 1-2 Solvates, definition of 43 Sulphur dioxide, liquid

complex formation in 97 conductivity of 95 exchange of solute with 96

INDEX 143

Sulphur dioxide, liquid — continued metathetic reactions in 99-100 physical properties 95 preparation 94 self-ionization 87, 95-6 solubilities in 95 solvates 96-7

Sulphuric acid acids and bases in 54-6 conductivity of 49-50 effect of high viscosity 47-8

physical properties 47 preparation 47 protonation by 53-4 radical formation in 56-7 self-ionization 48 solvates 50 solvolysis by 51-3

Xenon hexafluoride as solvent 110

non-aqueous solvents in inorganic chemistry

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