Чамберс К., Холлидей А.К. Современная неорганическая химия, 1975

236 GROUP V

value. However, in the presence of water vapour the line powder soon becomes covered with a layer of glassy trioxophosphoric acid, and this reduces the rate at which drying can occur. For this reason, gases are better dried by passing them through loosely-packed kpentoxide\ rather than merely over the surface.

Oxides of arsenic

Arsenic forms two important oxides, As4O6 and As4O10.

ARSENIC(lIl) OXIDE, As4O6

This is formed when arsenic burns in air (cf. phosphorus which gives P4O10). It can exist in two crystalline modifications ; the stable one at room temperature, which also occurs naturally as arsenolite, has an octahedral form. Solid arsenic(III) oxide is easily reduced, for example by heating with charcoal, when arsenic deposits as a black shiny solid on the cooler parts of the tube.

Arsenic(III) oxide is slightly soluble in water, giving a solution with a sweetish taste—but as little as 0.1 g can be a fatal dose! (The antidote is freshly-precipitated iron(III) hydroxide.) The solution has an acid reaction to litmus, due to the formation of arsenic(III) acid:

As4O6 + 6H2O ^ 4H3AsO3

Arsenic(III) acid is an extremely weak acid; in fact, the oxide is amphoteric, since the following equilibria occur :

Hence arsenic(III) oxide dissolves readily in alkalis to give arsenates(III), for example

As4O6 + 6CO^~ -> 4AsO^~ + 6CO2T

but in strong acid solution tripositive arsenic ions may be formed. This reaction indicates very clearly the increased electropositive character of arsenic.

In aqueous solution arsenic(III) oxide is a reducing agent being oxidised to arsenate(V) by halogens, chlorate(I), nitric acid and even iron(III) chloride.

G R O U P V 237

ARSENIC(V) OXIDE, As4O10

Unlike phosphorus pentoxide, this oxide cannot be made directly. Arsenic(V) acid, H3AsO4 (strictly, tetraoxoarsenic acid), is first prepared by oxidising arsenic(III) oxide with concentrated nitric acid or some other strong oxidising agent:

2H3AsO3 + 2HNO3 -> 2H3AsO4 + NOT + NO2T + H2O

On concentrating the solution, a solid of formula As4O10 .8H2O (which may be composed by hydrated arsenic(V) acid) is obtained, and this, on fairly prolonged heating to 800 K, loses water and leaves arsenic(V) oxide. No compounds corresponding to the other acids of phosphorus are formed, but salts are known.

Arsenic(V) oxide is a white deliquescent solid, which liberates oxygen only on very strong heating, leaving the (III) oxide:

As4O10 -> As4O6 + 2O2

It dissolves in water to give arsenic(V) acid, and in alkalis to form arsenates( V}.

Oxides of antimony

Antimony forms both a + 3 and a + 5 oxide. The -f 3 oxide can be prepared by the direct combination of the elements or by the action of moderately concentrated nitric acid on antimony. It is an amphoteric oxide dissolving in alkalis to give antimonates(III) (for example sodium 'antimonite', NaSbO2), and in some acids to form salts, for example with concentrated hydrochloric acid the trichloride, SbCl3, is formed.

Antimony(V) oxide can be prepared by treating antimony with concentrated nitric acid. It is an oxidising agent and when gently heated loses oxygen to form the trioxide. (The change in oxidation state stability shown by antimony should be noted since it corresponds to increasing metallic character.)

Unlike the amphoteric +3 oxide, the +5 oxide is acidic and dissolves only in alkalis to give hydroxoantimonates which contain the ion [Sb(OH)6J~. A third oxide, Sb2O4, is known but contains both antimony(Ili) and antimony(V), Sbm(SbvO4), cf. Pb3O4.

Oxides of bismuth

Bismuth forms both + 3 and + 5 oxides. The + 3 oxide, unlike the corresponding oxides of the other Group V elements, is insoluble

238 GROUPV

in alkalis, and dissolves only in acids (when bismuth salts are formed), a clear indication of the more metallic nature of bismuth.

Bismuth(V) oxide is not easy to prepare; the (III) oxide (or better a suspension of the hydroxide) must be oxidised with a strong oxidising agent such as the peroxodisulphate ion. When this is carried out, the bismuthate ion, [Biv(OH)6]", is formed. On evaporation, the sodium salt, for example, has the formula NaBiO3. Addition of acid to a solution of a bismuthate precipitates the (V) oxide, Bi2O5, but this loses oxygen rapidly and forms the trioxide. The bismuthate ion is an extremely strong oxidising agent, for example the manganese(II) ion Mn~^ is oxidised to manganate(VII)

OXOACIDS AND THEIR SALTS

Nitrogen

NITRIC(V') ACID, HNO3

Nitric acid is prepared in the laboratory by distilling equal weights of potassium nitrate and concentrated sulphuric acid using an air condenser, the stem of which dips into a flask cooled by tap water. The reaction is:

H2SO4 + KNO3 -» KHSO4 + HNO3

The temperature is kept as low as possible to avoid decomposition of the nitric acid to (brown) nitrogen dioxide. The nitric acid condenses out as a fuming liquid; it may be purified by redistillation with concentrated sulphuric acid. If the nitric acid is condensed at room temperature, it gives off dinitrogen pentoxide, N2O5 (which fumes with the atmospheric moisture), and so becomes diluted somewhat. Only if it is frozen out at 231 K (the melting point) does it form pure nitric acid, HNO3. "Concentrated9 nitric acid contains about 67 % of the pure acid—this is the constant boiling mixture formed by distilling a solution of any concentration. Hence concentrated nitric acid is not pure nitric acid.

On the large scale, nitric acid is now made in large quantities by the catalytic oxidation of ammonia, employing the reaction:

4NH3 + 5O2 -> 4NO46H2O: AH - - 120 kJmol"l

The process is as follows: ammonia gas (made by the Haber process) is liquefied under pressure, to freeze out any water, and the anhydrous gas is then passed together with dust-free air through a

GROUP V 239

converter (Figure 9.6). This contains a gauze of platinum, or platinum-rhodium, heated at first electrically, then maintained at red heat by the exothermic reaction which takes place on it. The air-ammonia mixture must only remain in contact with the catalyst for a fraction of a second, otherwise the nitrogen oxide decomposes to give nitrogen and oxygen. From this converter, the nitrogen oxide is mixed with more air, to convert it to nitrogen dioxide.This reaction is also exothermic and the heat from it may be used to pre-heat the air stream entering the converter.

Pt gauze

Air

Figure 9.6. Manufacture oj nitric acid

The nitrogen dioxide is then passed up a water-cooled steel tower, fitted with baffles down which water flows. Here the nitrogen dioxide dissolves to give nitric acid and nitrogen oxide ; air is also passed up the tower to oxidise the latter to give more nitrogen dioxide, which is absorbed in turn, so that ultimately almost complete conversion of the nitrogen oxides to nitric acid is complete; the acid is collected, at a strength of 50-65%, at the base of the tower.

Properties. Pure nitric acid isa colourless liquid,density 1 .52 g cm~3, dissociating slightly above its melting point into dinitrogen pentoxide and water, as already mentioned : on boiling, more oxides of nitrogen are formed and the liquid obtained is then the constant boiling-point acid, density 1.41gcm~3; hence this latter acid ('concentrated nitric acid*) is usually yellow in colour due to dissolved oxides formed during distillation. The colour deepens on exposure to daylight because nitrogen dioxide is formed in solution by the photochemical reaction :

-g-^ 4NO + 2HO -h O

240 GROUP V

A similar decomposition occurs if nitric acid is subjected to a temperature above its boiling point.

The chemical properties of nitric acid require us to consider the structure first. The vapour of pure nitric acid (i.e. anhydrous) is probably composed of molecules of 'hydrogen nitrate', which structurally is a resonance hybrid of such forms as :

>

In liquid nitric acid, hydrogen bonding gives a loose structure similar to that of hydrogencarbonateions. However, although pure nitric acid does not attack metals readily and does not evolve carbon dioxide from a carbonate, it is a conducting liquid, and undergoes auto-ionisation thus :

The second equilibrium is the more important, giving rise to the nitronium ion, NOJ, already mentioned as a product of the dis sociation of dinitrogen tetroxide. Several nitronium salts have been identified, for example nitronium chlorate(VII), (NO2)+(C1O4)~. If pure nitric acid is dissolved in concentrated sulphuric acid, the freezing point of the latter is depressed to an extent suggesting the formation of four ions, thus:

HNO3 + 2H2SO4 ^ NO2+ + H3O+ + 2HSO;

It is the nitronium ion which is responsible for nitrating actions in organic chemistry which are carried out in a mixture of nitric and sulphuric acids. When nitric acid is dissolved in water, its behaviour is that of a strong acid, i.e.:

HNO3 + H2O ^ H3O+ + NOs

because of the proton affinity of water. The majority of the reactions of nitric acid are oxidations due to the nitrate ion in the presence of hydrogen ions—and the corresponding reduction product (from the nitrate ion) depends upon the hydrogen ion concentration and upon the nature of the substance oxidised ; it may be nitrogen dioxide, nitrogen oxide, dinitrogen oxide, nitrogen,hydroxylamine (NH2OH) or ammonia (as ammonium ion in acid solution). The following are some typical examples :

GROUPV 241

(1) Non-metals:—These are often oxidised to the corresponding oxoacid, and nitrogen oxide is formed. For example, sulphur gives sulphuric acid with cold concentrated nitric acid:

S + 2HNO3 -> H2SO4 + 2NO

Iodine gives iodic(V) acid with hot concentrated acid:

3I2 + 10HNO3 -+ 6HIO3 + 10NO + 2H2O

Fluorine, however, gives the substance 'fluorine nitrate', NO3F:

HNO3 + F2 -> NO3F + HF

Violet phosphorus is oxidised to phosphoric(V) acid.

(2) Metals:—Nitric acid reacts with all common metals except gold and platinum, but some are rendered passive by the concentrated acid, for example aluminium, iron, cobalt, nickel and chromium. With the very weakly electropositive metals such as arsenic, antimony and tin, the oxide of the metal in its higher oxidation state is obtained, for example antimony yields the oxide antimony(V) oxide, Sb2O5 (in hydrated form). With more electropositive metals the nitrate of the metal is always formed, and the other products vary greatly. Metals which do not liberate hydrogen from dilute acids form nitrogen oxide or nitrogen dioxide, according to conditions. For example, copper in cold nitric acid (1:1) reacts thus:

3Cu + 8HNO3 -> 3Cu(NO3)2 + 2NO? + 4H2O

In concentrated nitric acid (when warmed) the reaction is:

Cu + 4HNO3 -> Cu(NO3)2 + 2NO2 + 2H2O

Metals which do liberate hydrogen from dilute acids, for example zinc, magnesium, can react with nitric acid to give dinitrogenoxide, for example:

4Zn 4- 10HNO3 -* 4Zn(NO3)2 + N2O + 5H2O

and if the hydrogen ion content of the nitric acid is further increased, by adding dilute sulphuric acid, hydroxylamine or ammonia is formed.

With very dilute nitric acid and magnesium, some hydrogen is evolved.

With a nitrate in alkaline solution, ammonia is evolved quantitatively by Devarda's alloy (Al, 45%; Cu, 50%; and Zn, 5%). This reaction can be used to estimate nitrate in absence of ammonium ions (see below):

242 GROUP V

NO" + 4Zn + 15OH- + 6H2O -» NHJ + 4Zn(OH)^

(3) Cations:—Some of these are oxidised to a higher state by nitric acid. For example, iron(II) (in presence of sulphuric acid) is quantitatively oxidised to iron(III):

3Fe2+ -h NOJ + 4H+ -> 3Fe3+ + NOT + 2H2O

Tin(II) chloride, in presence of hydrochloric acid, is oxidised to tin(IV) chloride, the nitrate ion in this case being reduced to hydroxylamine and ammonia.

The noble metals such as gold and platinum, although almost insoluble in nitric acid, are very ready to form chloro-complexes, for example gold gives the [AuCl4]~ ion very readily. Hence they can be dissolved by aqua regia, a mixture of 3 volumes of concentrated hydrochloric acid and 1 volume of concentrated nitric acid. The latter oxidises the gold to the auric fgold(III)) state (Au3+), which then appears as the ion (AuQ4)~ (p. 431).

NITRATES

Hydrated nitrates, and anhydrous nitrates of very electropositive metals (for example Na, K), contain the ion NO^ which has the structure:

o

resonance hybrids

with the three N—O distances identical. In other anhydrous metal nitrates, prepared as on p. 233, the nitrate groups may be bonded covalently to the metal, thus: M—ONO2 (for example Cu(NO3)2,

p.413).

Nitrates are prepared by the action of nitric acid on a metal or its

oxide, hydroxide or carbonate. All nitrates are soluble in water. On heating, the nitrates of the alkali metals yield only oxygen and the nitrite:

2KNO3 -> 2KNO2 + O2T

Ammonium nitrate gives dinitrogen oxide and steam:

NH4NO3 -> 2H2O + N2Ot

GROUP V 243

The nitrates of other metals give nitrogen dioxide, oxygen and the metal oxide, unless the latter is unstable to heat, in which case the metal and oxygen are formed (for example from nitrates of silver and mercury):

2Cu(NO3)2 -> 2CuO + 4NO2 + O2

2AgNO3 -> 2Ag -h 2NO2 + O2

Nitrates are detected by:

1. The action of heat on the solid (above).

2.By the brown ring test with iron(II) sulphate and cold con centrated sulphuric acid.

3.By their oxidising action; heating with copper and concen trated sulphuric acid yields brown fumes of nitrogen dioxide.

4.By the evolution of ammonia with Devarda's alloy in alkaline solution in absence of ammonium ions; this is used quantitatively, the ammonia being absorbed in excess standard acid and the excess acid back-titrated.

NITROUS ACID

Nitrous acid, HNO2, is known as a gas, but otherwise exists only in solution, in which it is a weak acid. Hence addition of a strong acid to a solution of a nitrite produces the free nitrous acid in solution.

Nitrous acid is unstable, decomposing to give nitric acid and evolving nitrogen oxide :

3HNO -* NO + H O + 4- 2NO

It is an effective oxidising agent and can oxidise iodide to iodine, and the ammonium ion to nitrogen. The reduction products of nitrous acid vary greatly with conditions. For example, nitrogen oxide or ammonia may be formed when hydrogen sulphide is oxidised to sulphur, according to the acidity of the solution. Hydrazine is oxidised by nitrous acid to hydrogen azide. Nitrous acid can itself be oxidised to nitric acid, but only by strong oxidising agents such as manganate(VII). Nitrous acid is important in organic chemistry for its ability to diazotise primary aromatic amines—an important step in the manufacture of dyestuffs.

NITRITES

These all contain the ion NO^. They are much more stable than nitrous acid, and those of the alkali metals can be fused without

244 GROUP V

decomposition. They are usually prepared by heating the alkali metal nitrate, alone or with lead as a reducing agent—the latter method being the one used in the manufactureof sodium nitrite for use in the dye industry. Lead will also reduce nitrate to nitrite if present as lead sponge':

The addition of even a weak acid (such as ethanoic acid) to a nitrite produces nitrous acid which readily decomposes as already indicated. Hence a nitrite is distinguished from a nitrate by the evolution of nitrous fumes when ethanoic acid is added.

Phosphorus

Phosphorus forms a large number of oxoacids, many of which cannot be isolated but do form stable salts. In general, ionisable hydrogen is bonded to the phosphorus through an oxygen atom; hydrogen atoms attached directly to phosphorus are not ionisable.

THE + 3 ACIDS

Two of these are important:

X-Ray diffraction studies of the oxoacid anions indicate the following probable arrangements for the acids:

O O

In each case the P—O bonds have some multiple character. Phosphinic acid is a moderately strong monobasic acid. On heating the acid and its salts they disproportionate evolving phosphine:

4H2PO2~ -> 2PH3 -f 2

GROUP V 245

Both the acid and its salts are powerful reducing agents. They reduce, for example, halogens to halides, and heavy metal cations to the metal. Copper(II) ion is reduced further to give copper(I) hydride, a red-brown precipitate:

3H3PO2 + 3H2O + 2Cu2+ -> 2CuHi + 3H3PO3 + 7H+

Phosphonic acid, H3PO3, often called just 'phosphorous acid', is prepared by the hydrolysis of phosphorus trichloride; a stream of air containing phosphorus trichloride vapour is passed into ice-cold water, and crystals of the solid acid separate:

PC13 + 3H2O -> H3PO3 + 3HC1

The acid is dibasic (seestructure p. 244). Like phosphinic acid it disproportionates when heated :

4H3PO3 -» PH3 + 3H3PO4

and is a strong reducing agent. Also like phosphinic acid it reduces heavy metal ions to the metal, but copper(II) ions are not reduced to CuH.

THE + 5 ACIDS

The important phosphoric acids and their relation to the anhydride P4O10are:

(The formulae P4O10,xH2O are merely to illustrate the interrelationship and have no structural meaning.)

Tetraoxophosphoric acid, H3PO4:—This is prepared in the laboratory either by dissolving phosphorus(V) oxide in water (giving trioxophosphoric acid) and then heating to give the tetraoxo-acid; or by heating violet phosphorus with 33% nitric acid, which oxidises it thus:

4P 4- 10HNO3 + H2O -> 4H3PO4 + 5NO? + 5NO2T

Caution is required in both methods. In the second case, in particular, gentle heating only is essential once the reaction starts.