Chambers, Holliday. Modern inorganic chemistry

326 GROUP VII: THE HALOGENS

In aqueous solution sulphur dioxide (sulphurous acid) is oxidised to sulphuric acid :

SO2 + C12 + 2H2O -* H2SO4 4- 2HC1

Chlorine reacts with some metallic oxides to yield chlorides, for example

2Fe2O3 + 6C12 -» 4FeCl3 + 3O2

Bromine has many oxidising reactions (E^ = + 1.07 V) and like chlorine it will oxidise sulphur dioxide in aqueous solution to sulphuric acid, and hydrogen sulphide to sulphur.

Iodine has the lowest standard electrode potential of any of the common halogens (E^ = +0.54 V) and is consequently the least powerful oxidising agent. Indeed, the iodide ion can be oxidised to iodine by many reagents including air which will oxidise an acidified solution of iodide ions. However, iodine will oxidise arsenate(III) to arsenate(V) in alkaline solution (the presence of sodium carbonate makes the solution sufficiently alkaline) but the reaction is reversible. for example by removal of iodine,

_ ~ + I2 + 2OH~ ^ AsOr + 2I~ 4- H2C arsenate(III) arsenate^V)

The oxidation of the thiosulphate ion S2Oj" to tetrathionate ion, S4Ol^, is used to estimate iodine:

The disappearance of iodine at the end point is detected by the addition of fresh starch solution which gives a blue complex as long as iodine is present.

HALOGEN COMPOUNDS

THE HYDRIDES (HYDROGEN HALIDES)

Physical properties

All the halogens form hydrides by direct combination of the elements. The hydrogen halides formed are covalently bonded, and when pure are colourless gases at room temperature. Some important physical properties of the hydrogen halides are given in Table 11.3 below. The data in Table 113 clearly reveal unexpected properties for hydrogen fluoride. A graph of atomic number of the halogen against b.p. for the hydrogen halides has been given on

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Table 11.3

PROPERTIES OF THE HYDROGEN MAUDES

In the solid state hydrogen fluoride exists as an infinite zig-zag chain of molecules. Association also occurs in the liquid and gaseous phases and in the latter phase, investigations indicate the presence of (HF)2 molecules and also more highly associated forms existing not only as chains but also as rings, for example (HF)6.

The ability to form hydrogen bonds explains the formation of complex ions such as HF^ and H2p3 when a fluoride salt, for example potassium fluoride, is dissolved in aqueous hydrofluoric acid:

KF + HF ^ KHF2

This reaction can be reversed by heating and is a convenient method of obtaining anhydrous hydrogen fluoride from an aqueous solution.

The dipole moments of the hydrogen halides decrease with increasing atomic number of the hydrogen, the largest difference occurring between HF and HC1, and association of molecules is not an important factor in the properties of HC1, HBr and HI. This change in dipole moment is reflected in the diminishingpermittivity (dielectric constant) values from HF to HI.

THERMAL STABILITY OF HYDROGEN HALIDES

The enthalpies of formation and hydrogen-halogen bond strengths are given in Table 113. The formation of hydrogen fluoride from its elements occurs with explosive violence; the hydrogen-fluorine bond produced is extremely strong (H—F = 566 kJ mol"1, cf. C—C in diamond 356 kJ mol~ *) and stable to heat up to very high temperatures. Both chlorine and bromine undergo a photochemical chain reaction with hydrogen. The hydrogen-halide bond strength

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The liquid, like water, has a high dielectric constant (permittivity) and is weakly conducting. It is a good solvent for many inorganic and organic substances, to give conducting solutions. Substances which move the equilibria to the right when dissolved in hydrogen fluoride, by taking up the fluoride ions, are 4acids'. For example, boron trifluoride forms the tetrafluoroborate anion in a solution of hydrogen fluoride:

2HF + BF3 ^H2 F+ + BFJ tetrafluoroborate ion

However, many substances, notably alcohols, have a greater proton affinity than the hydrogen fluoride molecule, and so behave as bases, for example ethanol:

C2H5OH + HF ^ C2l _

Even nitric acid will do this, i.e.:

HNO3 + HF ^ H2NOa 4- F"

Thus nitric acid behaves as a base in hydrogen fluoride. Hence increases of conductivity when substances dissolve in hydrogen fluoride may be due to "acidic' or 'basic' behaviour.

The preparation and reactions of hydrogen halides

HYDROGEN FLUORIDE, ANHYDROUS HYDROFLUORIC ACID, HF

Hydrogen fluoride is the most important compound of fluorine. It is prepared in the laboratory, and on the large scale, by the reaction of calcium fluoride with concentrated sulphuric acid:

CaF2 + H2SO4 -» CaSO4 + 2HF?

The reaction is carried out in a lead retort; one suitable for the laboratory can be made from a piece of lead piping, bent like a retort and closed at the shorter end. This is charged with fluorspar and the acid and heated, and the hydrogen fluoride is distilled into a polythene vessel.

Anhydrous hydrogen fluoride (as distinct from an aqueous solution of hydrofluoric acid) does not attack silica or glass. It reacts with metals to give fluorides, for example with heated iron the anhydrous iron(II) fluoride is formed; the same product is obtained by displacement of chlorine from iron(II) chloride:

Fed, + 2HF -» FeF? + 2HC1T

330 GROUP VII: THE HALOGENS

Hydrogen fluoride also effects replacement reactions in organic compounds. For example, carbon tetrachloride yields a mixture of chlorofluoromethanes CC13F, CC12F2 and so on. Like all the other hydrogen halides, hydrogen fluoride adds on to olefms, for example:

CH2=CH2 + HF -> CH3CH2F

Aqueous hydrogen fluoride is a weak acid (see above) and dissolves silica and silicates to form hexafluorosilicic acid; hence glass is etched by the acid, which must be kept in polythene bottles.

In addition to the abnormal properties already discussed, aqueous hydrofluoric acid has the properties of a typical acid, attacking metals with the evolution of hydrogen and dissolving most metallic hydroxides and carbonates.

Uses of hydrogen fluoride

By far the largest use of hydrogen fluoride is in the manufactureof fluorocarbons which find a wide variety of uses including refrigerants, aerosol propellants and anaesthetics. Hydrogen fluoride is also used in the manufacture of synthetic cryolite, Na3AIF6, and the production of enriched uranium.

HYDROGEN CHLORIDE

Hydrogen chloride is formed:

1.By the direct union of hydrogen and chlorine. Very pure hydrogen chloride is made by direct union of pure hydrogen and chlorine in a quartz vessel.

2.As the product of the hydrolysis of many substances in which chlorine is covalently bound, for example:

SOC12 + 2H2O -> H2SO3 + 2HC1

PC13 4- 3H2O -> H3PO3 + 3HC1

It is prepared in the laboratory by warming sodium chloride with concentrated sulphuric acid:

NaCl + H2SO4 -> NaHSO4 + HClt

The gas is dried by passage through concentrated sulphuric acid and collected over mercury.

On the large scale, hydrogen chloride can be produced by the

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same reaction, which is usually carried a stage further by stronger heating, i.e.

NaCl + NaHSO4 -> Na2SO4 + HClt

Anhydrous hydrogen chloride is not particularly reactive, either as a gas at ordinary temperatures, or a liquid (b.p. 188 K) and does not react with metals such as iron or zinc, nor with dry oxides. A few reactive metals such as sodium, will burn in the gas to give the chloride and hydrogen :

2Na + 2HC1 -> 2NaCl + H2

However, if heated hydrogen chloride is passed over heated metals, the chloride is formed ; in the case of a metal exhibiting variable oxidation state, the lower chloride is obtained :

Aqueous hydrochloric acid

In aqueous solution, hydrogen chloride forms hydrochloric acid. The concentrated acid contains about 40% hydrogen chloride (about 12 M). A graph of the boiling point of hydrogen chloridewater mixtures against composition shows a maximum at about 20 % HC1; hence if either the concentrated or dilute acids be distilled, then either hydrogen chloride or water respectively distil over, leaving behind "constant boiling-point' acid.

Hydrochloric acid is a strong monobasic acid, dissolving metals to form salt and evolving hydrogen. The reaction may be slow if the chloride formed is insoluble (for example lead and silver are attacked very slowly). The rate of attack on a metal also depends on concentration ; thus aluminium is attacked most rapidly by 9 M hydrochloric acid, while with other metals such as zinc or iron, more dilute acid is best.

Electrolysis of hydrochloric acid yields hydrogen at the cathode and oxygen at the anode from the dilute acid, but chlorine at the anode (of carbon) from the concentrated acid. Electrolysis of the concentrated acid is used on the large scale to recover chlorine.

If tetramethylammonium chloride is dissolved in hydrochloric acid, the unstable salt [(CH3)4N] [HC12], can be crystallised out; here chlorine is showing weak hydrogen bonding (cf. F----H—F~ which is stable, and C1--H—Cl~ which is unstable).

332 GROUP VII: THE HALOGENS

Uses of hydrogen chloride—Hydrogen chloride is sometimes used in the preparation of an ester, for example ethyl benzoate, where it acts as both an acid catalyst and a dehydrating agent. Hydrochloric acid is used primarily to produce chlorides, for exampleammonium chloride. It is extensively used in the manufactureoi aniline dyes, and for cleaning iron before galvanising and tin-plating.

HYDROGEN BROMIDE, HBr

Hydrogen bromide cannot be prepared readily by the action of sulphuric acid on a bromide, because the latter is too easily oxidised by the sulphuric acid to form bromine. It is therefore obtained by

Moist violet phosphorus

on glass beads

Bromine

. = >. —3

Hydrogen bromide

Moist violet phosphorus

Figure 113, Preparation of hydrogen bromide

the hydrolysis of a covalent bromide; a convenient one is phosphorus tribromide. By dropping bromine on to a paste of violet phosphorus and water, phosphorus tribromide is formed and immediately hydrolysed thus:

PBr3 + 3H2O -» H3PO3 + BHBrt

Any free bromine can be removed by passing the evolved gas through a U tube packed with glass beads covered with moist violet phosphorus (Figure 113).

Hydrogen bromide may also be prepared by dropping bromine into benzene containing aluminium powder, which acts as a catalyst to the reaction:

C6H6 + Br2-^-C6H5Br + HBrt

Hydrogen bromide is a colourless gas similar in properties to hydrogen chloride. It is very soluble in water, giving hydrobromic

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acid. The latter may be prepared directly by slow hydrolysis of a covalent bromide: a convenient one is disulphur dibromide, S2Br2, made by dissolving sulphur in excess liquid bromine. The mixture is then hydrolysed, and hydrobromic acid distilled off:

S2Br2 + 2Br2 + 4H2O -> 6HBr -f H2SO4 + Si

The acid which conies over is a constant boiling mixture containing about 47% hydrogen bromide (density = 1.46gem"3).

Hydrobromic acid is rather easily oxidised when exposed to light and becomes brown due to the bromine liberated. Otherwise, its properties are those of a strong acid, similar to hydrochloric acid.

HYDROGEN IODIDE, HI

Hydrogen iodide is prepared in a similar way to hydrogen bromide, by the action of water on a mixture of iodine and violet phosphorus. The hydrogen iodide evolved may be collected by downward delivery or may be condensed (b.p. 238 K); it reacts with mercury and so cannot be collected over the latter.

An aqueous solution of hydrogen iodide, up to 50% strength, may be prepared by passing hydrogen sulphide (or sulphur dioxide) into a suspension of iodine in water:

H2S + I2 -*2H+ + 21- + Si SO|" + I2 + H2O -» 2H+ + 2I~ + SO|~

These reactions illustrate the oxidising action of iodine.

In the first reaction, sulphur may be filtered off, leaving only hydriodic acid.

Properties—Hydrogen iodide is a colourless gas. It is very soluble in water and fumes in moist air (cf. hydrogen chloride), to give hydriodic acid. Its solution forms a constant boiling mixture (cf. hydrochloric and hydrobromic acids). Because it attacks mercury so readily, hydrogen iodide is difficult to study as a gas, but the dissociation equilibrium has been investigated.

Hydriodic acid is a strong acid, reacting with bases to give iodides, containing the ion I~. It is also a strong reducing agent (so also is hydrogen iodide, particularly at high temperatures, when dissociation into hydrogen and iodine is considerable). Thus, it reduces sulphuric acid to a mixture of sulphite, sulphur and hydrogen sulphide, the last reaction predominating:

H,SO4 -f 8HI -> H2S -f 41, + 4H2O

334 GROUP VII:THE HALOGENS

Hence hydrogen iodide cannot be produced by the reaction of sulphuric acid with an iodide. Hydriodic acid is slowly oxidised by air (more rapidly in light) liberating iodine:

4HI + O2 -» 2H2O + 2I2

Other examples of its reducing action are:

1. Reduction of dinitrogen oxide to ammonia (which gives the ammonium ion with the acid):

N2O 4- 10HI -> 2NH^ + 21" + H2O + 412

2. Reduction of nitric to nitrous acid:

HNO3 + 2HI -> HNO2 -f I2 -f H2O

OXIDES

None of the halogens reacts directly with oxygen but all form oxides by indirect methods.

Fluorine oxides

The oxides of fluorine are more correctly called oxygen fluorides because of the greater electronegativity of fluorine.

Oxygen difluoride OF2 is obtained when a rapid stream of gaseous fluorine is passed through 2 % caustic soda solution:

2F2 -f 2NaOH -» 2NaF + OF2 + H2O

It is a gas at room temperature with a boiling point of 128 K. It is a strong oxidising agent, some reactions occurring with explosive violence. Water hydrolyses it slowly at room temperature, but the reaction evolving oxygen is rapid in the presence of a base, and explosive with steam:

OF2 -h H2O -> O2 + 2HF

Fluorine is known to form three other oxides, O2F2, O3F2 and O4F2 but all these decompose below 200 K.

Chlorine oxides

Chlorine forms several very reactive, unstable oxides. Dichlorine monoxide C12O is a yellowish gas at room temperature, the liquid

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boiling at 275 K. It is prepared by treating freshly prepared yellow mercury(II) oxide with either chlorine gas, or with a solution of chlorine in tetrachloromethane (carbon tetrachloride):

2HgO + 2C12 -> HgO .HgCU + C12O

On heating (and sometimes at ordinary temperatures) it explodes, yielding chlorine and oxygen—this decomposition also being catalysed by light. It dissolves in water to give an orange-yellow liquid containing some chloric(I) acid of which dichlorine monoxide is the formal anhydride. It is a strong oxidising agent converting many metals to a mixture of their oxides and chlorides.

Liquid chlorine dioxide, C1O2, boils at 284 K to give an orangeyellow gas. A very reactive compound, it decomposes readily and violently into its constituents. It is a powerful oxidising agent which has recently found favour as a commercial oxidising agent and as a bleach for wood pulp and flour. In addition, it is used in water sterilisation where, unlike chlorine, it does not produce an unpleasant taste. It is produced when potassium chlorate(V) is treated with concentrated sulphuric acid, the reaction being essentially a disproportionation of chloric(V) acid :

3KC1O3 + 3H2SO4 -» 3KHSO4 + 3HC1O3

3HC1O3 -» 2C1O2 -h HC1O4 + H2O

The reaction usually proceeds with explosive violence and a better method of preparation is to heat, gently, moist crystals of ethanedioic acid (oxalic acid) and potassium chlorate(V) :

2KC1O3 4- 2H2C2O4 -» K2C2O4 + 2H2O + 2CO2 + 2C1O2

Industrially an aqueous solution of chlorine dioxide can be prepared by passing nitrogen dioxide up a packed tower down which sodium chlorate(V) flows :

C1OJ + NO2 -> NOa + C1O2

The aqueous solution is safe to handle, the dissolution being essentially physical. On standing in sunlight the solution slowly decomposes to a mixture of acids. In alkaline solution a mixture of chlorate(III), C1OJ, and chlorate(V), CIOJ, ions is rapidly produced. Chlorine dioxide is paramagnetic, the molecule containing an odd electron and having a structure very like that of NO2 (p. 231).

Dichlorine hexoxide. C17O*. is formed when chlorine dioxide is