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

336 GROUP VII: THE HALOGENS

exposed to ultraviolet light or by the action of ozone on chlorine dioxide:

6C1O2 + 2O3 -> 3C12O6

It is a liquid at room temperature, melting point 276.5 K. The molecular weight, determined in carbon tetrachloride, indicates the dimefic formula, but magnetic measurements show the presence of small quantities of the paramagnetic monomer C1O3 in the pure liquid. It is rather an unstable compound and decomposes slowly even at its melting point, and more rapidly on heating, forming finally oxygen and chlorine. It is a powerful oxidising agent and reacts violently even with water with which it forms a mixture of chloric(V) and chloric(VII) acids.

Dichlorine htptoxide, C12O7, is the most stable of the chlorine oxides. It is a yellow oil at room temperature, b.p. 353 K, which will explode on heating or when subjected to shock. It is the anhydride of chloric(VIl) acid (perchloric acid) from which it is prepared by dehydration using phosphorus(V) oxide, the acid being slowly reformed when water is added.

Bromine oxides

These are all unstable substances and little is known about them.

Dibromine monoxide, Br2O, is prepared, similar to the corresponding dichlorine compound, by the action of a solution of bromine in carbon tetrachloride on yellow mercury(II) oxide:

2HgO + 2Br2 -» Hg2OBr2 + Br2O

It is a dark brown liquid, m.p. 256 K, which decomposes rapidly at room temperature.

Tribromine octoxide, Br3O8, is a white solid obtained when ozone and bromine react together at 273 K at low pressure. It is unstable above 200 K in the absence of ozone. It is known to exist in two forms, both soluble in water.

Bromine dioxide^ BrO2, is prepared by passing an electric discharge through a mixture of oxygen and bromine at low temperature and pressure. It is a yellow solid, stable only below 230 K, decomposing above this temperature to give oxygen and bromine.

GROUP VII: THE HALOGENS 337

Iodine oxides

fhere appears to be only one true oxide of iodine, diiodine pentoxide, I2O5. It is a white solid prepared by heating iodic acid(V) to 450 K:

2HIO3 -* H2O + I2O5

As the equation indicates, it is the anhydride of iodic-acid(V), which is re-formed when water is added to the pentoxide. Mixed with concentrated sulphuric acid and silica, it is a quantitative oxidising agent for carbon monoxide at room temperature:

SCO + I2O5 -> 5CO2T + I2

OXO-ACIDS AND THEIR SALTS

For many years it was thought that fluorine did not form any oxoacids or oxo-acid anions. Recent work, however, indicates the existence of fluoric(I) acid (hypofluorous acid), HFO, formed by the reaction of fluorine with water at 273 K. The acid forms colourless crystals, m.p. 156K, is very unstable and has, as expected, very strong oxidising properties.

The acids of chlorine(I), bromme(T) and iodine(I) are weak acids, the pKa values being 7.4, 8.7 and 12.3 respectively.They are good oxidising agents, especially in acid solutions. The acids decrease in stability from chloric(I) to iodic(I).

Only chlorine forms a +3 acid, HC1O2. This is also a weak acid and is unstable. The +5 acids, HXO3, are formed by chlorine, bromine and iodine; they are strong acids. They are stable compounds and all are weaker oxidising agents than the corresponding +1 acids.

The existence of chloric(VII) (perchloric) and iodic(VII) (periodic) acids has long been known but bromic(VII) acid has only recently been prepared.

Halic(I) acids of chlorine, bromine and iodine

The amount of halic(I) acid formed when the halogen reacts reversibly with water decreases from chlorine to iodine and the concentration of iodic(I) acid in a saturated solution of iodine is negligible. However the equilibrium

2H2O + X2 ^ HXO + H3O+ 4- X~

GROUP VII: THE HALOGENS 339

Halic(IH) acids, HXO2

Only chloric(III) acid, HC1O2, is definitely known to exist. It is formed as one of the products of the reaction of water with chlorine dioxide (see above). Its salts, for example NaClO2, are formed together with chlorates(V) by the action of chlorine dioxide on alkalis. Sodium chlorate(III) alone may be obtained by mixing aqueous solutions of sodium peroxide and chlorine dioxide:

2C1O2 + Na2O2 -> 2NaClO2 + O2t

A solution of the free acid may be obtained by using hydrogen peroxide, instead of sodium peroxide.

Chloric(III) acid is a fairly weak acid, and is an oxidising agent, for example it oxidises aqueous iodide ion to iodine. Sodium chlorate(III) (prepared as above) is used commercially as a mild bleaching agent; it bleaches many natural and synthetic fibres without degrading them, and will also bleach, for example, oils, varnishes and beeswax.

Chlorates(III) disproportionate on heating, or on boiling the aqueous solution, thus:

3C1O2 -> 2C1OJ + Cl~

chlorate(V) chloride

Halic(V) acids

Chlorine, bromine and iodine form halic(V) acids but only iodic(V) acid, HIO3, can be isolated. Solutions of the chloric(V) and bromic(V) acids can be prepared by the addition of dilute sulphuric acid to barium chlorate(V) and bromate(V) respectively, and then filtering (cf. the preparation of hydrogen peroxide). These two acids can also be prepared by decomposing the corresponding halic(I) acids, but in this case the halide ion is also present in the solution.

Attempts to concentrate chloric(V) and bromic(V) acids beyond certain limits lead to decomposition which may be violent.

lodic(V) acid is prepared by oxidising iodine with concentrated nitric acid:

312 4- 10HNO3 -» 6HIO3 + lONOt + 2H2O

The iodic acid(V) and some diiodine pentoxide separate out and the iodic(V) acid is purified by recrystallization from hot water.

All the halic(V) acids are strong acids and their salts are not appreciably hydrolysed in aqueous solution. They are also powerful oxidising agents (seebelow).

340 G R O U P V I I : THE HALOGENS

HALATE(V) SALTS

Generally the solubility of a given metal halate decreases from chlorate(V) to iodate(V) and many heavy metal iodates(V) are quantitatively insoluble. Like their parent acids, the halates(V) are strong oxidising agents, especially in acid solution: their standard electrode potentials are given below (in volts):

Unexpectedly we find that the bromate(V) ion in acid solution (i.e. effectively bromic(V) acid) is a more powerful oxidising agent than the chlorate(V) ion, C1OJ. The halates(V) are thermally unstable and can evolve oxygen as one of the decomposition products. Potassium chlorate(V), when heated, first melts, then resolidifies due to the formation of potassium chlorate(VII) (perchlorate):

4KC1O3 -* 3KC1O4 + KC1

but a further, stronger heating will make the chlorate(VII) decompose, evolving oxygen:

KC1O4 -» KC1 + 2O2

The decomposition of potassium chlorate(V) is catalysed by manganese(IV) oxide, MnO2, and oxygen is evolved on heating the mixture below the melting point of the chlorate(V),

The ability of the solid chlorates(V) to provide oxygen led to their use in matches and fireworks. Bromates(V) and iodates(V) are used in quantitative volumetric analysis. Potassium hydrogen diiodate(V), KH(IO3)2, is used to standardise solutions of sodium thiosulphate(VI) since in the presence of excess potassium iodide and acid, the reaction

IOJ + 51" + 6H+ -» 3I2 + 3H2O

occurs quantitatively. The liberated iodide is then titrated using the thiosulphate solution of which the concentration is required:

GROUP VII: THE HALOGENS 341

Haifc(VII) acids

The existence of chloric(VII), (perchloric), HC1O4, and several periodic(VII) acids has long been established. Bromie(VII) acid and the bromate(VII) ion have only recently been discovered.

These acids differ so greatly in their properties that they will be considered separately.

CHLORIC(VIl) ACID AND CHLORATES(VIl)

Chloric(VII) acid is prepared by carefully distilling potassium chlorate(VII) with concentrated sulphuric acid under reduced pressure:

KC1O4 + H2SO4 -> KHSO4 + HC1O4

It is a liquid, b.p.363 K, but if heated it decomposes and hence must be distilled under reduced pressure; decomposition may occur with explosive violence and this can occur even at room temperature if impurities are present. Combustible material, for example paper and wood, ignite spontaneously with explosive violence on contact with the acid, and it can produce painful blisters on the skin.

Chloric(VII) acid fumes in moist air and is very soluble in water, dissolving with the evolution of much heat. Several hydrates are known; the hydrate HC1O4. H2O is a solid at room temperature and has an ionic lattice [H3O+] [C1OJ].

The oxidising properties of the aqueous solutions of chloric(VII) acid change dramatically with temperature and the concentration of the acid. Cold dilute solutions have very weak oxidising properties and these solutions will react, for example, with metals, producing hydrogen without reduction of the chlorate(VII) ion occurring:

Zn + 2HC1O4 -» Zn(ClO4)2 + H2T

Hot concentrated solutions of chloric(VII) acid and chlorates(VII), however, react vigorously and occasionally violently with reducing agents.

Chloric(VII) acid is one of the strongest acids known, and it behaves as such even when dissolved in solvents with poor proton affinity; thus it can be used as an acid in pure ethanoic acid as a solvent:

CH3COOH + HC1O4 ^ CH3COOH^ + CIO4

342 G R O U P V I I ; THE HALOGENS

CHLORATES(VII)

These can be prepared by electrolytic oxidation of chlorates(V) or by neutralisation of the acid with metals. Many chlorates(VII) are very soluble in water and indeed barium and magnesium chlorates- (VII) form hydrates of such low vapour pressure that they can be used as desiccants. The chlorate(VII) ion shows the least tendency of any negative ion to behave as a ligand, i.e. to form complexes with cations, and hence solutions of chlorates (VII) are used when it is desired to avoid complex formation in solution.

The chlorate(VII)ion, C1O~, is isoelectronic with the sulphate(VI) ion, SO^", and has a similar tetrahedral symmetry.

lodic(VII) acids

These are acids which can be regarded, in respect of their formulae (but not their properties) as hydrates of the hypothetical diiodine heptoxide, I2O7. The acid commonly called 'periodic acid; I2O7 . 5H2O, is written H5IO6 (since the acid is pentabasic) and should strictly be called hexaoxoiodic(VII) acid. It is a weak acid and its salts are hydrolysed in solution. It can be prepared by electrolytic oxidation of iodic(V) acid at low temperatures :

IO3" + 2H2O + OH" -» H5IO 6

The Aperiodic acids' and "periodates' are powerful oxidising agents and they will oxidise manganese to manganate(VII), a reaction used to determine small quantities of manganese in steel.

HALIDES

The rigid classification of halides into covalent and ionic can only be an oversimplification, and the properties of the halides of a given element can very greatly depend upon the halogen. Thus the classification is only one of convenience.

General methods of preparation

Many salt-like halides can be prepared by the action of the hydrohalic acid, HX, on the metal or its oxide, hydroxide or carbonate. The halides prepared by this method are often hydrated, particularly when a less electropositive metal is involved, for example zinc, iron.

GROUP VII: THE HALOGENS 343

Anhydrous halides, however, are obtained when the metal is heated with the dry hydrogen halide or the halogen. In the case of elements with more than one oxidation state, the hydrogen halide produces a lower halide and the halogen a higher halide, for example

Sn + 2HC1 -» SnCl2 + H2T

Sn + 2C12 -» SnCl4

The higher iodides, however, tend to be unstable and decomposition occurs to the lower iodide (PI5 -» PI3). Anhydrous chlorides and bromides of some metals may also be prepared by the action of acetyl (ethanoyl) halide on the hydrated ethanoate (acetate) in benzene, for example cobalt(II) and nickel(II) chlorides:

Co(CH3COO)2 + 2CH3COC1 + 2H2O -> CoCl2i + 4CH3COOH

Sulphur dichloride oxide (thionyl chloride) on the hydrated chloride can also be used to produce the anhydrous chloride in certain cases, for example copper(II) chloride and chromium(III) chloride:

CrCl3. 6H2O + 6SOC12 -> 6SO2t + 12HC1T + CrCl3

Halides of non-metals are usually prepared by the direct combination of the elements. If the element exhibits more than one oxidation state, excess of the halogen favours the formation of the higher halide whilst excess of the element favours the formation of the lower halide (e.g. PC15 and PC13).

Ionic (salt-like) halides

These are halides formed by highly electropositive elements (for example those of Groups I and II, except for beryllium and lithium). They have ionic lattices, are non-volatile solids, and conduct when molten; they are usually soluble in polar solvents in which they produce conducting solutions, indicating the presence of ions.

The change from ionic to covalent bonding is gradual in a given group or period; for a given halogen, as the size of the metal ion decreases and more especially as its charge increases, the degree of covalency increases. Thus, for example, in the chlorides of the four elements, potassium, calcium, scandium and titanium, i.e. KC1, CaCl2, ScCl3 and TiCl4, KC1 is essentially ionic, TiCl4 is essentially covalent.

When the several halides of a given element are considered, changes in bond character are also found. The fluoride is generally the most ionic with ionic character decreasing from fluoride to

344 GROUP VII: THE HALOGENS

iodide, for example aluminium trifluoride, A1F3, is ionic but the remaining aluminium halides are all essentially covalent.

When an element has more than one oxidation state the lower halides tend to be ionic whilst the higher ones are covalent—the anhydrous chlorides of lead are a good example, for whilst lead(II) chloride, PbCl2, is a white non-volatile solid, soluble in water without hydrolysis, lead(IV) chloride, PbCl4, is a liquid at room temperature (p. 200) and is immediately hydrolysed. This change of bonding with oxidation state follows from the rules given on p.49.

The solid anhydrous halides of some of the transition metals are often intermediate in character between ionic and covalent; their structures are complicated by (a) the tendency of the central metal ion to coordinate the halide ions around it, to form an essentially covalent complex, (b) the tendency of halide ions to bridge, or link, two metal ions, again tending to covalency (cf. aluminium chloride, p. 153 and iron(III) chloride, p. 394).

SOLUBILITY

Many ionic halides dissolve in water to give hydrated ions. The solubility of a given halide depends on several factors, and generalisations are difficult. Ionic fluorides, however, often differ from other halides in solubility. For example, calcium fluoride is insoluble but the other halides of calcium are highly soluble; silver fluoride, AgF, is very soluble but the other silver halides are insoluble.

Covalent halides

These are formed by less electropositive elements. They are characterised by the existence of discrete molecules which exist even in the solid state. They have generally lower melting and boiling points than the ionic halides, are more volatile and dissolve in non-polar solvents.

The melting and boiling points of a series of similar covalent halides of a given element are found to increase from the fluoride to the iodide, i.e. as the molecular weight of the halide increases. Thus, the trihalides of phosphorus have melting points PF3 = 121.5 K. PC13 = 161.2 K, PBra = 233 K, PI3 = 334K.

Most covalent halides are hydrolysed by water (carbon tetrachloride being a notable exception, p. 195) to give acidic solutions, by either method (a) (example FeCl3) or method (b) (example BC13)'

GROUP VII: THE HALOGENS 345

(a)FeCl3 4- 6H2O -> [Fe(H2O)6]3 + + 3C1"

[Fe(H2O)6]3+ + H2O ^ [Fe(H20)5(OH)]2+ + H3O+ etc.

(b)BC13 + 3H2O -> H3BO3 + 3HC1

The hydrolysis of phosphorus tribromide or triiodide is used in the preparation of hydrogen bromide and hydrogen iodide respectively:

PBr3 + 3H2O -» H3PO3 + 3HBrT

PI3 + 3H2O -> H3PO3 + 3HIT

Complex halides

Halogens can act as ligands and are commonly found in complex ions; the ability of fluorine to form stable complex ions with elements in high oxidation states has already been discussed (p. 316). However, the chlorides of silver, lead(II) and mercury(I)are worthy of note. These chlorides are insoluble in water and used as a test for the metal, but all dissolve in concentrated hydrochloric acid when the complex chlorides are produced, i.e. [AgCl2]~, [PbCl4]2~ and [HguCl3]~, in the latter case the mercury(I) chloride having also disproportionated.

INTER HALOGEN COMPOUNDS AND POLYHALIDES

There are four types of interhalogen compound:

Type XX^: IF7 (the only example), iodine heptafluoride

Iodine monochloride, IC1, monobromide, IBr, and trichloride, IC13, are solids at room temperature, the remainder being volatile liquids or gases. They are made by the direct combination of the halogens concerned. All are covalent with the larger halogen occupying a central position. With the exception of iodine pentafluoride, IF5, they are extremely reactive, behaving (like halogens) as oxidising agents and reacting with water. The two most important mterhalogen compounds are the trifluorides of chlorine, C1F3 (the