Chemical Properties of Vanadium

Pure vanadium is stable and retains its lustre in damp air. On being rapidly heated in a stream of oxygen the powdered metal burns, forming vanadium pentoxide, V₂O₅; the reaction is, however, incomplete. At a bright red heat the metal combines with nitrogen to form a nitride. In excess of chlorine, vanadium burns to form a tetrachloride, VCl₄, which is also produced by the action of carbonyl, sulphuryl, thionyl, and sulphur chlorides at 600° C. When heated in hydrogen the gas is absorbed. Vanadium is not attacked by solutions of alkali chlorides, bromine water, or cold hydrochloric acid, whether dilute or concentrated. Hydrochloric acid gas, however, at 300° to 400° C. gives rise to the trichloride, VCl₃. Vanadium is slowly attacked by hydrofluoric acid and by hot, concentrated sulphuric acid. A specimen of vanadium which contained 8.66 per cent, of carbon and 1.6 per cent, of other impurities, when treated with concentrated sulphuric acid at 330° C., gave vanadium pentoxide, V₂O₅, with evolution of sulphur dioxide. At lower temperatures the dioxide, VO₂, was formed, but this was converted into the pentoxide when the temperature was raised, thus:

(i) 2V + 4H₂SO₄ = 2VO₂ + 4SO₂ + 4H₂O.
(ii) 2VO₂ + H₂SO₄ ⇔ V₂O₅ + SO₂ + H₂O.

Vanadium is readily attacked in the cold by dilute nitric acid and by concentrated nitric acid or aqua-regia, giving vanadic acid. This latter acid is also formed by the action of other oxidising agents on vanadium, as, for example, chloric acid, perchloric acid, bromic acid, potassium iodate. On being fused with sodium carbonate, caustic potash, or potassium nitrate, vanadates of sodium or potassium are produced. Vanadium reduces solutions of mercuric chloride, cupric chloride, and ferric chloride to mercurous chloride, cuprous chloride, and ferrous chloride respectively, and precipitates the metal from solutions of gold chloride, silver nitrate, platinum chloride, iridium chloride. Carbon monoxide attacks vanadium between 500° and 800° C. with the formation of a carbide. Glass and porcelain vessels absorb vanadium at high temperatures. The metal can be rendered " passive," as in the case of iron, by immersion in oxidising agents, e.g. chromic acid, nitric acid, or by making the metal the anode in an electrolytic bath of various salts. Cathodic treatment reconverts the vanadium to the " active " state. Marino, however, could not effect this change, but this may have been due to differences in the purity of the vanadium specimens used.

The compounds of vanadium are numerous and to some extent complicated; this is due to the variable valency of the element. It forms a series of oxides - VO, V₂O₃, VO₂, V₂O₅, and although the oxide V₂O₇ has not hitherto been isolated, pervanadates derived from it are well defined. A monoxide having the formula V₂O has also been reported, but its existence is doubtful; compounds containing monovalent vanadium are unknown. As is usual in the case of any one element, the acidity of the oxide increases with increasing oxygen content, and basic properties gradually become less marked.

The difficulty that is experienced in reducing any of the oxides of vanadium to the metal is attributed partially to their very strongly exothermic nature. Ruff and Friedrich obtained the following figures from combustions carried out in a bomb calorimeter -

1. 2V + 5(½O₂) = V₂O₅ + 437,000±7,000 calories.
2. 2V + 3(½O₂) = V₂O₃ + 302,000±10,000 calories

Previous determinations of the heat of formation of the pentoxide gave 313,030 calories and 250,315 calories.

Mixter investigated the thermal changes that ensue when the lower oxides are converted into the pentoxide, and from these data calculated the following -

1. 2V + 5(½O₂) = V₂O₅ + 441,000 calories.
2. 2V + 2O₂ = 2VO₂ + 412,800 calories
3. 2V + 3(½O₂) = V₂O₃ + 353,200 calories
4. 2V + O₂ = 2VO + 208,600 calories

The last figure is in fair agreement with an indirect determination effected by Slade and Higson, who obtained

2V + O₂ = 2VO + 222,000 calories.

It is of interest to note that, commencing with vanadium pentoxide, the amounts of heat absorbed in the successive formation of the next lower oxide rapidly increase as the metal is reached. The heat changes in the following equations have been calculated from Mixter's figures, given above:

1. V₂O₅ - ½O₂ = V₂O₄ - 28,200 calories.
2. V₂O₄ - ½O₂ = V₂O₃ - 59,600 calories
3. V₂O₃ - ½O₂ = 2VO - 144,600 calories
4. 2VO - O₂ = 2V - 208,600 calories

The heats of formation of a few other oxides are here inserted for the sake of comparison:

1. 2P + 5(½O₂) = P₂O₅ + 365,300 calories.
2. 2Fe + 3(½O₂) = Fe₂O₃ + 195,600 calories
3. 2Cr + 3(½O₂) = Cr₂O₃ + 267,000 calories
4. 2Al + 3(½O₂) = Al₂O₃ + 392,600 calories

The comparatively high heat of formation of vanadium pentoxide and the tendency of aluminium to alloy with metallic vanadium explain the non-success of the application of the thermite process for the production of pure vanadium from the pentoxide and from vanadates.

Vanadyl salts are salts of tetravalent vanadium, and contain the divalent [VO]^•• radical. Many vanadium compounds are known which appear to contain a [VO] group, but the vanadium is either trivalent or pentavalent. Throughout this book the term vanadyl is restricted to compounds of tetravalent vanadium, that is, to salts of the oxide VO₂. Hence, for example, the compound VOCl₃, which contains pentavalent vanadium, is called vanadium oxy trichloride, and not by the more usual but less logical name " vanadyl chloride."

As is usual with salts of metals which exhibit variable valency, those of vanadium are coloured in solution. The colour varies with the valency; salts of V₂O₅ are yellow, those of VO₂ are blue, those of V₂O₃ are green, and those of VO are lavender (see table). Remarkable colour changes can be observed by diluting considerably the reddish solution obtained by dissolving the pentoxide, V₂O₅, in hydrochloric or sulphuric acid and then adding metallic zinc. Under the influence of the nascent hydrogen produced the solution passes through all shades of blue and green, and finally assumes a lavender tint. The same effect can be produced by electrolytic reduction of the hydrochloric acid solution. These characteristic colours cannot definitely be attributed to the existence of penta-, tetra-, tri-, and di-valent vanadium cations, since, as has been indicated, the various vanadium salts readily undergo hydrolysis in contact with water to give rise most probably to the following cations - [VO]^••• or [VO₂]^• from V^V, [VO]^•• from V^IV, and [VO]^• from V^III. These oxygenated radicals cannot be without influence on the different colours observed.

Several reactions, the velocities of which are affected by the presence of vanadium salts, have been quantitatively investigated. It appears to be established that the compounds employed usually function as oxygen carriers, and that their effect depends, therefore, on the ease with which they undergo oxidation and reduction. To give two instances - (a) The reduction of chloric acid, HClO₃, by hydriodic acid, HI, is accelerated by the addition of a vanadous salt, because chloric acid is much more rapidly reduced by a vanadous salt than by hydriodic acid; on the other hand, the reduction of persulphuric acid, H₂S₂O₈, with HI is not appreciably affected by addition of a vanadous salt, because the last named reduces persulphuric acid but slowly. (b) Vanadium pentoxide accelerates the oxidation of sucrose to oxalic acid by nitric acid, that of ethyl alcohol to acetaldehyde and acetic acid by air, that of potassium iodide to iodine by hydrogen peroxide, and that of stannous salts by nitric acid. In these reactions the vanadium pentoxide gives up its oxygen to the oxidisable substance, being itself reformed at the expense of the oxidising agent.

In some cases, however, the modus operandi is modified. In the oxidation of hydriodic acid with chromic acid, the data indicate that while liberation of iodine takes place, the vanadous or hypovanadic salt employed as the catalyst also undergoes oxidation to vanadate. The vanadium compound here belongs to the class of catalysts known as inductors, and the reaction is comparable to the oxidation in aqueous solution of sodium sulphite with sodium arsenite, whereby both sodium sulphate and sodium arsenate are produced.

More recently the conversion of benzene into maleic acid in the presence of vanadium oxides as catalysts has been studied with a view to throwing light on the mechanism of such oxidations. The data obtained seem to show that the action depends on an " oscillation " between the two oxides V₂O₅ and V₂O₄, the dissociation of the former supplying activated oxygen for the reaction; but it is also shown that the nature of the products of the oxidation is a function of some other property of the catalyst not yet clearly understood.

The presence of a vanadium salt in dilute sulphuric acid solution has also been found to improve the catalytic action of platinum in the combination of hydrogen and oxygen.

Vanadium precipitates the metal from solutions of salts of gold, silver, platinum, and iridium, and reduces solutions of mercuric chloride, cupric chloride and ferric chloride to mercurous chloride, cuprous chloride, and ferrous chloride, respectively. In these reactions the vanadium passes into solution as the tetravalent ion. No precipitation or reduction ensues, however, when vanadium is added to solutions of divalent salts of zinc, cadmium, nickel, and lead. From these reactions it has been estimated that the electrolytic potential of the change, vanadium (metal) → tetravalent ions, is about –0.3 to -0.4 volt, which is approximately equal to the electrolytic solution pressure of copper. This figure is a little uncertain through the difficulty of securing pure vanadium.

When an electrolyte which is without action on vanadium at ordinary temperatures (for example, dilute solutions of mineral acids, of oxalic acid, or of potassium halides) is electrolysed with a vanadium anode, a complex tetravalent vanadium ion is produced. Similarly, electrolysis at 100° C. and in molten chlorides of sodium or zinc gives rise to complex tetravalent vanadium ions. The E.M.F. in each case is found to be independent of the nature of the electrolyte. When, however, solutions of caustic soda or of caustic potash are employed, the vanadium dissolves as a pentavalent ion, irrespective of variations in concentration, temperature, or current density. The pentavalent ion is electro-negative; the tetravalent ion is strongly hydrolytic, and readily gives rise to the vanadyl ion [VO]^••, which is weakly electropositive. The trivalent vanadium ion displays more marked electropositive properties, which again increase with the formation of divalent ions.

In 1898, Cowper-Coles claimed to have successfully effected the electrolytic reduction of an acid solution of vanadium pentoxide to metallic vanadium, but the product was subsequently shown by Fischer to have been a deposit of platinum hydride. Fischer, in a series of over three hundred experiments, varied the temperature, current density, cathode material, concentration, electrolyte, addition agent, and construction of cell, but in not one instance was the formation of any metallic vanadium observed. In most cases reduction ceased at the tetravalent state (blue). At temperatures above 90° C. reduction appeared to proceed to the divalent state (lavender). The use of carbon electrodes led to the trivalent state (green), but only lead electrodes produced the trivalent state at temperatures below 90° C. Platinum electrodes reduced the electrolyte to the blue vanadyl salt below 90° C.; using a divided cell and temperatures above 90° C. the lavender salt was obtained.

Electrolytic reduction of pentavalent and tetravalent vanadium salts has frequently been employed for the preparation of vanadium compounds of lower valency. Bleecker has also prepared vanadium pentoxide and several vanadates electrolytically.

Roscoe, in 1870, found that vanadium " absorbs or combines with " up to 1.3 per cent, of its weight of hydrogen, according to the state of division of the metal, when heated in a current of the gas, and confirmed the observation in the following year. Muthmann, Weiss, and Riedelbauch subsequently reported that the amount of hydrogen taken up by the vanadium increases with increase of temperature and duration of contact; they stated that at 1300° C. a stable hydride of vanadium was produced containing 16.1 per cent, of hydrogen. This is described as a black powder which is unaffected by air, hot water, or boiling hydrochloric acid. Prandtl and Manz, however, were unable to obtain any compound of vanadium and hydrogen, and state that the previously observed increases in weight were due either to (a) absorption of oxygen and nitrogen, or (b) absorption of hydrogen by impurities in the vanadium. More recent investigations with vanadium of 90 per cent, purity which had been preheated for one hour in a vacuum at 1100° C. have shown that the quantity of hydrogen absorbed varies with the temperature and the pressure. One gram of the metal absorbs 122.6 c.c. of hydrogen at ordinary temperatures and pressures and 2.01 c.c. at 1100° C. It is, therefore, improbable that a definite compound is produced.

The variable valency of vanadium is well displayed in its halides and oxyhalides. These are set out in the following table -

Halides and oxyhalides of Vanadium

Valency	Fluorine	Chlorine	Bromine	Iodine
Divalent	VF2	VCl2
Trivalent	VF3 VF3.3H2O	VCl3 VCl3.3H2O VOCl	VBr3 VBr3.6H2O VOBr	VI3.6H2O
Tetravalent	VF4 VOF2	VCl4 VOCl2	VOBr2
Pentavalent	VF5 VOF3 VO2F	VOCl3	VOBr3
Doubtful valency.		VO2Cl2.8H2O V2O2CL4.5H2O V2O3CL2.4H2O	VO2Br3.5H2O

The table shows that the stability of the halides decreases with increase in the atomic weight of the halogen. All the halides are hygroscopic and show a very strong tendency to undergo hydrolysis, a tendency which increases with the valency. The tetrabromide and tetriodide have not been isolated; VF₄ and VCl₄ can perhaps be regarded as salts of the very weak base VO₂. They are easily fusible compounds, and undergo hydrolysis so readily that they evolve the gaseous halogen acid and "fume " in moist air; they are therefore comparable with the tetrahalides of titanium, germanium, tin and lead.

The preparation of the anhydrous halides of vanadium is possible only in the dry way, since attempts to remove the water from the hydrated halide result in the formation of a basic salt, as in the cases of the halides of iron, chromium, aluminium, bismuth, etc. The general methods of preparation employed are:

1. Halogenation of the oxides or sulphides.
2. Reduction of a higher halide either directly or by means of hydrogen.
3. Direct synthesis.

Since fluorine is difficult to prepare and manipulate, the anhydrous fluorides and oxyfluorides are prepared by the action of anhydrous hydrogen fluoride on other halides or oxyhalides of vanadium.

The double fluorides of vanadium and other double halogen salts in most cases can be regarded as in accordance with Werner's theory of co-ordinated compounds. The vanadium has a co-ordination number six.

The hydrated fluoride, VF₃.3H₂O, may be regarded as an aquo-salt, . The gradual replacement of aquo-water by fluorine gives the following -

; ; [VF₆](NH₄)₃;
i.e.
VF₃.NH₄F.2H₂O; VF₃.2NH₄F.H₂O; VF₃.3NH₄F. Similarly, the double potassium salt, VF₃.2KF.H₂O, is . The alkali metal can be replaced by aniline to produce salts of the same type -

and [VF₆](NH₃C₆H₅)₃.

The constitution of the acid radical of the potassium salt containing one molecule of water corresponds to that of those double salts which contain seven molecules of water. It therefore follows that the acid radical in the latter is also co-ordinatively saturated. The salt VF₃.ZnF₂.7H₂O thus becomes . Replacement of the zinc atom by atoms of other metals gives the corresponding cobalt, nickel, and cadmium compounds, which have similar constitution.

Double salts of the other vanadium halides are much less stable and therefore less numerous than those containing fluorine. They are represented on the co-ordination theory by formulae which are analogous to those set out above. Unlike vanadium trifluoride, which forms VF₃.3H₂O, the molecules of the corresponding hydrates of the other halogen compounds are found to possess 6H₂O, and are written -

[V(H₂O)₆]Cl₃; [V(H₂O)₆]Br₃; [V(H₂O)₆]I₃.

It is of interest to note that Meyer and Backa, by treating vanadium trichloride and tribromide with liquid ammonia, have recently obtained hexammine derivatives -

[V(NH₃)₆]Cl₃; [V(NH₃)₆]Br₃.

These are comparable in their reactions to the hexammine of ferric chloride, [Fe(NH₃)₆]Cl₃. Hexammino-vanadium trichloride loses its chlorine on being treated with nitric acid and forms the corresponding nitrate, [V(NH₃)₆](NO₃)₃.

The double oxy-salts may be regarded as belonging to the following general series -

; ; .

Thus, VOF₂.2NH₄F.H₂O can be written ; VOF₂.ZnF₂.7H₂O as ; VO₂F.3NH₄F as ; VO₂F.ZnF₂.7H₂O as . It will be observed that the coordination number six is maintained by bringing a molecule of water into the co-ordinated complex.

Ephraim has observed that the composition of the halogen double salts of vanadium and other metals appears to be dependent on the atomic weight of the second metal. This is shown by rewriting the formulae for some of the double salts - 2VF₃.6NH₄F, 2VF₃.5NaF, and 2VF₃.4KF. It has not been found possible to prepare double salts in which the number of molecules of alkali fluoride combined with two molecules of VF₃ is greater than 6, 5, and 4 respectively. Efforts to prepare a double salt having the composition 2VF₃.5KF, for instance, were unsuccessful. Similarly, rearranging the formulae for some double salts of vanadyl fluoride gives - 3VOF₂.9NH₄F, 3VOF₂.8NaF, and 3VOF₂.7KF.

The heats of formation of the three anhydrous chlorides of vanadium have been determined by combustion in a bomb calorimeter and are found to be as follows -

(i) V (solid) + Cl₂ (gas) = VCl₂ (solid) + 147,000 (±4000) calories.
(ii) V (solid) + 3/2Cl₂ (gas) = VCl₃ (solid) + 187,000 (±8000) calories
(iii) V (solid) + 2Cl₂ (gas) = VCl₄ (liquid) + 165,000 (±4000) calories

It is obvious that the combustion of a molecule of vanadium with an increasing number of molecules of chlorine is not. accompanied by a gradually increasing evolution of heat. The figures show that the formation of vanadium tetrachloride (liquid) from vanadium trichloride (solid) and chlorine (gas) proceeds endothermically -

(iv) VCl₃ (solid) + ½Cl₂ (gas) = VCl₄ (liquid) - 22,000 calories.

Vanadium tetrachloride is, in fact, stable only at high temperatures. The last figure is, however, unreliable, since it is considerably affected by (a) the experimental errors involved in the reactions (ii) and (iii) above, and (b) the heat of liquefaction of vanadium tetrachloride, which is at present unknown.

The heat of formation of vanadium oxytrichloride is given by the equation -

(v) V (solid ) + ½O₂ (gas) + 3/2Cl₂ (gas) = VOCl₃ (liquid) + 200,000 (±4000) calories.

Iodine does not react with finely divided vanadium, nor with the nitride, VN, or vanadous oxide, V₂O₃. Anhydrous vanadium triiodide has not as yet been prepared.

Normal chromates of vanadium have not been prepared. An ammonium vanadoclirornate, 2(NH₄)₂O.V₂O₅.2CrO₃.7H₂O, has been prepared in red crystals by dissolving vanadium pentoxide in ammonium chromate solution and evaporating at ordinary temperatures in vacuo.

Vanadium combines slowly with nitrogen when heated in the gas to high temperatures. A substance having the empirical formula V₂N was obtained at a red heat, but the existence of this compound must be regarded as doubtful, because the mononitride, VN, is formed at a white heat. The temperature at which combination takes place is also somewhat uncertain. Pure vanadium does not absorb nitrogen below 1250° C., but with a ferrovanadium alloy it was found that absorption of nitrogen by the vanadium takes place at an increasing rate from 800° to 1200° C., when a maximum absorption of 8 per cent, of nitrogen is reached; at higher temperatures the nitrogen content decreases rapidly.

When either vanadium trioxide, V₂O₃, or vanadium pentoxide, V₂O₅, is reduced with carbon in the electric furnace, the product contains carbon in proportions varying from 4 per cent, to 25 per cent., depending on the temperature attained and other conditions. If a mixture of the pentoxide and sugar charcoal is heated in the carbon tube of the electric furnace for ten minutes, a definite carbide having the formula VC is obtained. More recently this compound has been prepared by heating a mixture of vanadium trioxide and carbon in an atmosphere of hydrogen at 1100° C.

Normal carbonates of vanadium are unknown. An unstable ammonium vanadyl carbonate, of composition 3(NH₄)₂CO₃.7VO₂.5CO₂.16H₂O, has been obtained in small, violet crystals by the addition of ammonium carbonate to a neutral solution of vanadyl sulphate.

Vanadium and silicon are miscible in all proportions in the liquid state up to 60 per cent, vanadium. The eutectic point is at 1411° C. near to the silicon end of the series, while the composition at the maximum approximates to the composition of the disilicide, VSi₂. Two compounds of vanadium and silicon, V₂Si and VSi₂, have been prepared in the electric furnace.