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

216 GROUPV

(since the pressure is high) and the unconverted hydrogen and nitrogen are returned to the inlet and passed again over the catalyst.

In the laboratory ammonia is obtained when any ammonium salt is heated with an alkali, either solid or in solution:

NH: + OH' -> NH3t + H2O

It is best prepared by heating an intimate mixtureof solidammonium chloride and quicklime:

2NH4C1 -h CaO -> CaCl2 + 2NH3 + H2O

After drying over quicklime, calcium oxide CaO, the ammonia is collected by upward delivery. (N.B. Both of the common drying agents, calcium chloride and concentrated sulphuric acid, combine with the gas.)

Ammonia is also produced when an ionic nitride is hydrolysed, for example magnesium nitride, produced when magnesium burns in nitrogen:

Mg3N2 4- 6H2O -> 3Mg(OH)2 + 2NH3T

PROPERTIES

Ammonia is a colourless gas at room temperature and atmospheric pressure with a characteristic pungent smell. It is easily liquefied either by cooling (b.p. 240 K) or under a pressure of 8-9 atmospheres at ordinary temperature. Some of its physical and many of its chemical properties are best understood in terms of its structure. Like the other group head elements, nitrogen has no d orbitals available for bond formation and it is limited to a maximum of four single bonds. Ammonia has a basic tetrahedral arrangement with a lone pair occupying one position:

/"N

/\

\1

X

Because of the lone pair of electrons, ammonia has a dipole moment (high electron density at the lone pair) and this concentration of negative charge can attract (positive) hydrogen atoms in adjacent molecules giving fairly strong intermolecular forces, i.e. hydrogen bonding. Consequently ammonia has a high latent heat of vaporisation and a relatively high boiling point (see Table 9.2 and p. 52),

G R O U P V 217

facts at one time made use of in refrigeration employing ammonia. The great solubility of ammonia in water (1 volume of water dissolves 1300 volumes of ammonia at 273K) can be attributed to hydrogen bonding between ammonia and water molecules. (N.B. Concentrated ammonia solution has a density of 0.880 gem"3 and contains 35 % of ammonia.) The reaction :

NH3 4- H 2 O^NH 3 . H 2 O

is exothermic and can easily be reversed by heat, all the ammonia being evolved on boiling.

A second competing reaction also occurs:

NH3 .H2 O -^NH^ + OH~

For this second reaction K298 = 1.81 x 10~5 and hence pK6 for ammonia solution is 4.75. The entity NH3 . H2O is often referred

to as ammonium hydroxide, NH4OH, a formula which would imply that either nitrogen has a covalency of five, an impossible arrangement, or that NH4OH existed as the ions NH^ and OH~. It is possible to crystallise two hydrates from concentrated ammonia solution but neither of these hydrates is ionic. Hence use of the term "ammonium hydroxide' is to be discouraged in favour of 'ammonia solution'.

CHEMICAL PROPERTIES OF AMMONIA

These may, for convenience, be divided into a number of topics but all are closely related depending very largely on the presence of the lone pair of electrons on the nitrogen atom.

Ammonia as a donor molecule. Because of the presence of the lone pair of electrons on the nitrogen atom, ammonia can behave as an electron pair donor. For example, ammonia abstracts a proton from a water molecule producing the tetrahedral ammonium, NH^, ion and forms the compounds H3N—>A1C13 and H3N—>BC13.

The commonly observed behaviour of ammonia as a ligand is due to the lone pair of electrons on the nitrogen atom, and ammonia forms numerous complex ammines with both transition elements and typical metals; the bonding varies from weak ion-dipole attraction to strong covalent bonding. (For examples of ammonia as a ligand, see pp. 46, 363.) The formation of the ammine CaCl2 . 8NH3 explains why calcium chloride cannot be used to dry ammonia gas.

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Ammonia as a base. The ammonia molecule has a powerful affinity for protons and hence.

1. ammonia gas will react with gaseous hydrogen containing compounds which are acidic, for example hydrogen chloride:

NH3 + HC1 c°o! NH4Cl(i.e. NH + CP)

heai

(N.B. A trace of water is required to make the forward reaction proceed at a realistic rate.)

2. ammonia will react with aqueous acids, for example

2NH3 + H2SO4(aq) -> (NH4)2SO4

which is more correctly written

2NH3 + 2H3O+ + SOJ- -> 2NH^ + 2H2O + SOJ-

Aqueous ammonia can also behave as a weak base givinghydroxide ions in solution. However, addition of aqueous ammonia to a solution of a cation which normally forms an insoluble hydroxide may not always precipitate the latter, because (a) the ammonia may form a complex ammine with the cation and (b) because the concentration of hydroxide ions available in aqueous ammonia may be insufficient to exceed the solubility product of the cation hydroxide. Effects (a) and (b) may operate simultaneously. The hydroxyl ion concentration of aqueous ammonia can be further reduced by the addition of ammonium chloride; hence this mixture can be used to precipitate the hydroxides of, for example, aluminium and chromium(III) but not nickel(II) or cobalt(II).

Because of ammine formation, when ammonia solution is added slowly to a metal ion in solution, the hydroxide may first be precipitated and then redissolve when excess ammonia solution is added; this is due to the formation of a complex ammine ion, for example with copper(II) and nickel(II) salts in aqueous solution.

Ammonia as a reducing agent. Ammonia gas will not burn in air but it does burn in oxygen with a yellowish flame after ignition. A convenient apparatus is shown in Figure 9.3, By reversing the gas supplies it can easily be shown that oxygen will also burn in ammonia.

In the presence of catalyst, usually platinum, ammonia is oxidised by oxygen (and air) to nitrogen oxide. NO. This reaction, used to obtain nitric acid from ammonia (p. 238), can be demonstrated in the laboratory using the apparatus shown in Figure 9.4; the oxygen rate should be slow.

220 G R O U P V

Using the apparatus shown in Figure 9.3 it can be shown that ammonia gas will burn in chlorine gas, the ignition being spontaneous in this case:

2NH3 + 3C12 -> N2 + 6HC1

6HC1 + 6NH3 -» 6NH4C1

If ammonia is used in large excess and the chlorine diluted with nitrogen, chloramine, NH2C1, is formed:

NH3 + C12 -> NH2C1 4- HC1

When chlorinegas is in excess a highlyexplosive substance,nitrogen trichloride, NC13, is formed:

2NH3 + 6C12 -» 2NC13 4- 6HC1

When chlorine is passed into aqueous ammonia, ammonium chloride and nitrogen are formed. If, however, sodium chlorate(I) (hypochlorite) is used instead of chlorine, chloramine is first formed:

NH3 + OC1" -> NH2C1 + OH"

Normally the chloramine immediately undergoes further reaction, giving off nitrogen:

2NH2C1 4- OCr + 2OH~ -» N2T + 3C1" + 3H2O

but in the presence of glue or gelatin the chloramine reacts with more ammonia to give hydrazine:

NH2C1 + NH3 4- OH" -> N2H4 4- Cl~ + H2O

It is thought that the function of the glue or gelatin is to combine with very slight traces of heavy metal cations, for example Cu2 +, which are known to catalyse the nitrogen-forming reaction.

Ammonia will reduce metallic oxides which are reduced by hydrogen (for example copper(II) oxide, CuO, lead(II) oxide, PbO), being itself oxidised to nitrogen:

2NH3 + 3PbO -> 3Pb + N2T 4- 3H2O

Reactions with electropositive metals. Ammonia gas reacts with strongly electropositive metals to form the amide, for example

2Na + 2NH3 -> 2NaNH2 + H2

This reaction also occurs slowly when sodium is dissolved in liquid ammonia; initially a deep blue solution is formed which then decomposes giving hydrogen and sodium amide.

many inorganic and organic substances as well as for the alkali metals. Liquid ammonia is slightly ionised:

2NH3 ^-NH + + NH2- (cf.2H2O ^H3O+ + OH")

Liquid ammonia, like water, is only a poor conductor of electricity. Ammonium salts dissolved in water behave as acids giving the ion NH^, whilst amides which give the ion NH^ behave as bases. Thus the reaction:

NH4C1 4- KNH2 -> KC1| + 2NH3 acid base salt solvent

is a neutralisation in liquid ammonia (p. 90).

Solutions of alkali metals in liquid ammonia are used in organic chemistry as reducing agents. The deep blue solutions effectively contain solvated electrons (p. 126), for example

Na -» Na+ + e~

e~ + xNH3-»e-(NH3)x

Ammonium salts. Ammonium salts can be prepared by the direct neutralisation of acid by ammonia. The salts are similar to alkali metal salts and are composed of discrete ions. Most ammonium salts are soluble in water. Since ammonia is volatile and readily oxidisable the behaviour of ammonium salts to heat is particularly interesting.

If the acid of the salt is also volatile, as in the chloride and the carbonate, dissociation occurs causing the salt to sublime:

NH4C1 ^ NH3 + HC1

The extent of dissociation at a given temperature can be determined by measuring the density of the vapour. Since anhydroussulphuric acid is less volatile than hydrogen chloride, ammonium sulphate does not readily sublime on heating; some ammonia is evolved to leave the hydrogensulphate:

(NH4)2SO4 -> NH4HSO4 4- NH3T

If the acid of the ammonium salt is an oxidising agent, then on heating the salt, mutual oxidation and reduction occurs. The oxidation products can be nitrogen or one of its oxides and the reactions can be explosive, for example:

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(NH4)2Cr207 -* N2 + 4H2O + Cr2O3

NH4NO3 -» N2O + 2H2O

The mixture of ammonium nitrate and powdered aluminium is an explosive known as ammonal.

Uses of ammonia and ammonium compounds. Most of the ammonia produced is used in the manufacture of nitrogenous fertilisers such as ammonium sulphate. Other uses include nitric acid and synthetic fibre and plastic manufacture.

DETECTION OF AMMONIA AND AMMONIUM SALTS

All ammonium salts evolve ammonia on heating with alkali. Ammonia may be detected by (a) its smell, (b) its action in turning red litmus blue and (c) the orange-brown colour produced with Nessler's reagent. This is a very sensitive test.

Ammonia may be estimated by dissolving the gas in a known volume of standard acid and then back-titrating the excess acid. In a method widely used for the determination of basic nitrogen in organic substances (the Kjeldahl method), the nitrogenous material is converted into ammonium sulphate by heating with concentrated sulphuric acid. The ammonia is then driven off by the action of alkali and absorbed in standard acid.

Ammonia present in very small quantities in solution may be estimated by comparing the intensity of colour produced with Nessler's reagent (p. 439) with standard colours, using a simple form of colorimeter called a 'Nessleriser'.

Hydroxylamine, NH2OH

Hydroxylamine is derived from ammonia by replacing one hydrogen atom by a hydroxylgroup. It is prepared by the electrolytic reduction of nitric acid, using a lead cathode :

HNO3 4- 6H+ + 6e~ -» NH2OH 4- 2H2O

Sulphuric acid is added to the electrolyte and the hydroxylamine is formed as hydroxylammonium sulphate, (NH3OH)2SO4 [cf. (NH4)2SO4]. Addition of barium chloride then precipitates barium sulphate and hydroxylammonium chloride, (NH3OH)C1, isobtained.

Pure hydroxylamine is a crystalline solid of low melting point (306 K) but is rarely prepared because it decomposes above 288 K

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above the surface (or it can be immersed if a high current density is used). The Raschig process—the reaction of ammonia with chloramine (p.220)—gives lower yields and the hydrazine isnot anhydrous.

Pure hydrazine is a colourless liquid, melting point 275 K, and boiling point 387 K. It is surprisingly stable for an endothermic compound (A/ff = + 50.6 kJ moP1). Each nitrogen atom has a lone pair of electrons and either one or both nitrogen atoms are able to accept protons to give N2H5h and the less stable N2H^. The base strength of hydrazine is, however, lower than that of ammonia. As might be expected, hydrazine is readily soluble in water from which the hydrate N2H4.H2O can be crystallised.

Hydrazine, unlike ammonia, will burn in air with evolution of much heat:

N2H4 + O2 -> N2 -f 2H2O

This reaction has been carefully studied with the aim of obtaining the enthalpy of combustion as electrical energy, and successful hydrazine-air fuel cells have been developed using potassium hydroxide as the electrolyte. The hydrazine fuel, however, has the disadvantage that it is expensive and poisonous.

In aqueous solution hydrazine can behave either as an oxidising or reducing agent. Powerful reducing agents such as zinc reduce hydrazine to ammonia, while chlorine oxidises it to give nitrogen:

N2H5+ 4- C12 -> N2T + 5H+ 4- 4CT

Hydrazine and its alkylated derivatives are used as rocket fuels; in organic chemistry, substituted phenylhydrazines are important in the characterisation of sugars and other compounds, for example aldehydes and ketones containing the carbonyl group ^c=O.

Hydrogen azide (hydrazoic acid), HN3

Hydrazoic acid has no resemblance to either ammonia orhydrazine. -i- _

It has a structure involving resonance between H—N=N=N and _ +

H—N—N=N. It is prepared by the oxidation of hydrazine in strongly acid solution; the oxidising agent used is usually nitrous acid (i.e. sodium nitrite is added to the acid solution of hydrazine):

N2HJ + HNO2 -> HN3 -f H+ + 2H2O

Pure hydrazoic acid is a colourless liquid, b.p. 310 K. It is very ready to detonate violently when subjected to even slight shock, and so is used in aqueous solution. It is a weak acid, reacting with alkali to give azides, which contain the ion NJ.

G R O UP V 225

Hydrazoic acid behaves as both an oxidising and reducing agent in solution. Thus it will oxidise hydrochloric acid to chlorine, the main products being nitrogen and ammonium ions:

HN3 + 3H+ + 2C1" -» C12T 4- NH^ + N2T

On the other hand, chloric(I) acid, for example, oxidises hydrazoic acid to nitrogen:

2HN3 + ocr -» 3N2T + cr -h H2o

The azides are salts which resemble the chlorides in solubility behaviour, for example silver azide, AgN3, is insoluble and sodium azide, NaN3, soluble in water. Sodium azide is prepared by passing dinitrogen oxide over molten sodamide:

2NaNH2 + N2O -> NaN3 + NaOH + NH3

All the azides are potentially dangerous, and liable to detonate on heating, but those of the alkali and alkaline earth metals can be heated with caution if pure; they then evolve pure nitrogen.

Hydrides of phosphorus

PHOSPHINE

Phosphine can be prepared by the reaction of a strong alkali with white phosphorus; potassium, sodium and barium hydroxidesmay be used:

P4 4- 3KOH + 3H2O -> 3KH2PO2 4- PH3T

potassium phosphinate (hypophosphite)

This reaction gives an impure product containing hydrogen and another hydride, diphosphane, P2H4.

Pure phosphine can be prepared by the reduction of a solution of phosphorus trichloride in dry ether with lithium aluminium hydride:

4PC13 + 3LiAlH4 -> 4PH3T + 3LiCl 4- 3A1C13

The reaction of potassium hydroxide solution with phosphonium iodide also gives pure phosphine:

PH4I + KOH -> KI 4- H2O + PH3T