Chambers, Holliday. Modern inorganic chemistry

316 GROUPVIhTHE HALOGENS

which the central atom exhibits a high co-ordination number. The other larger halide ions show this tendency to a greatly diminished extent and the complexes formed are usually less stable, although certain metals (e.g. mercury) form iodo-complexes, for example [HgI4]2" which are more stable than fluoroor chloro-complexes. In certain cases there is insufficient space around the atom for as many iodine atoms as for other halogens, for example phosphorus forms pentahalides with fluorine, chlorine and bromine (and in the case of fluorine the ion [PF6] ~), but no pentaiodide. The large size of iodine also accounts for the fact that there are few complexes with more than four iodine ligands.

An important reason for low coordination of iodide ions is that high coordination Implies a high oxidation state of the central atom, which often (but not always) means high oxidising power—and this means oxidation of the easily oxidised iodide ligands. Thus the nonexistence of, for example, phosphorus(V) pentaiodide is to be explained by the oxidation of the iodide ligands and reduction of phosphorus to the +3 state, giving only PI3, not PI5.

OCCURRENCE AND EXTRACTION

FLUORINE

Fluorine occurs widely in nature as insoluble fluorides. Calcium

(PO4)2. Bones and teeth contain fluorides and some natural water contains traces.

Fluorine cannot be prepared directly by chemical methods. It is prepared in the laboratory and on an industrial scale byelectrolysis. Two methods are employed: (a) using fused potassium hydrogenfluoride, KHF2, in a cell heated electrically to 520-570 K or (b) using fused electrolyte, of composition KF :HF = 1:2, in a cell at 340-370 K which can be electrically or steam heated. Moissan, who first isolated fluorine in 1886, used a method very similar to (b) and it is this process which is commonly used in the laboratory and on an industrial scale today. There have been many cell designs but the cell is usually made from steel, or a copper-nickel alloy (4MoneF metal). Steel or copper cathodes and specially made amorphous carbon anodes (to minimise attack by fluorine) are used. Hydrogen is formed at the cathode and fluorine at the anode, and thehydrogen fluoride content of the fused electrolyte is maintained by passing in

GROUP VII: THE HALOGENS 317

hydrogen fluoride periodically. The fluorine obtained is almost pure, containing only a little hydrogen fluoride, which is removed by passage of the gas over sodium fluoride :

NaF -f HF -> NaHF2

Fluorine boils at 85 K to give a greenish-yellowdiatomic gas.

CHLORINE

The most common compound of chlorine is sodium chloride, NaCl, and this occurs widely in nature. Large deposits are found in Cheshire and these are extracted by the use of water although some is mined as rock salt. In many parts of the world sodium chloride is obtained from sea water. Other chlorides are found in small

oxidising agents are required for the oxidation and amongst those commonly used are manganese(IV) oxide, MnO2, potassium dichromate(VI), K2Cr2O7, both of which need to be heated with concentrated hydrochloric acid, and potassium manganate(VII), KMnO4, which evolves chlorine at room temperature when treated with concentrated hydrochloric acid :

MnO2 + 4HC1 -> MnCl2 -f C12 + 2H2O

14H+ + Cr2O?~ + 6C1" -> 2Cr3+ + 7H2O + 3C12

16H+ + 2MnO + 10C1" -> 2Mn2+ + 8HO + 5C1

Alternatively a mixture of almost any solid chloride and manganese- (IV) oxide will yield chlorine when warmed with concentrated sulphuric acid. These are the most common laboratory methods but there are many others.

On a large scale chlorine is obtained in several ways.

1.By the electrolysis of concentrated sodium chloride solution; this process was initially used primarily for the production of sodium hydroxide but the demand for chlorine is now so great that the chlorine is a primary and not a by-product.

2.By the electrolysis of fused magnesium chloride or fused sodium chloride.

3.By the oxidation of hydrogen chloride. A mixture of hydrogen chloride with air or oxygen is passed over a catalyst of copper(II)

318 GROUP VII: THE HALOGENS

chloride containing one or more chlorides of rare-earth metals on a silica support at a temperature of 600-670 K; the reaction is exothermic:

4HC1 + O2 ==* 2H2O 4- C12

The equilibrium constant for this reaction decreases with increase in temperature but the higher temperature is required to achieve a reasonable rate of conversion. Hydrogen chloride is now being produced in increasing quantities as a by-product in organic chlorination reactions and it is economic to re-convert this to chlorine.

Chlorine has a boiling point of 238 K and is a greenish-yellow diatomic gas at room temperature. It can be liquefied by cooling or by a pressure of a few atmospheres at room temperature.

BROMINE

Bromides of sodium, potassium, magnesium and calcium occur in sea water (about 0.07 % bromine) but the Dead Sea contains much more (5% bromine). Salt deposits (e.g. at Stassfurt) also contain these bromides. Silver bromide, AgBr, is found in South America.

In the laboratory, bromine is prepared by oxidation of bromide ion; the oxidation is carried out by mixing solid potassium bromide with manganese(IV) oxide and distilling with concentrated sulphuric acid:

2KBr + MnO2 + 3H2SO4 -> Br2 + 2KHSO4 + MnSO4 + 2H2O

The bromine is condensed and collected in a water-cooled receiver as a dark-red liquid.

On the industrial scale, bromine is obtained from sea water by using the displacement reaction with chlorine (the reaction by which bromine was discovered):

2Br" + C12 -> 2CP + Br2

The sea water is first treated with chlorine in acid solution (sulphuric acid is added) and very dilute bromine is obtained by blowing air through the solution. This is mixed with sulphur dioxide and the gases passed up a tower down which water trickles:

SO2 + Br2 + 2H2O -> 2HBr + H2SO4

The mixture of the two acids (now much richer in bromine than the sea water) is then treated with chlorine again, and bromine

GROUP VII: THE HALOGENS 319

obtained. The bromine may be freed from chlorine by bubbling it through iron(III) bromide solution, which retains the chlorine. Last traces of bromine from the process can be removed by passing over moist iron filings. Bromine is a dark-red heavy liquid, boiling point 332K, appreciably volatile at ordinary temperatures. It is soluble in organic solvents, for example chloroform, and they can be used to extract bromine from aqueous solutions (see Tests, p. 349).

IODINE

Iodine occurs to a minute extent (less than 0.001 %) in sea water, but is found in greater concentration, combined in organic form, in certain seaweeds, in oysters and in cod livers. Crude Chile saltpetre, or caliche contains small amounts of sodium iodate, NaIO3, from which iodine can be obtained (see below). Some insoluble iodides. for example tiiose of silver and mercury(II), occur in Mexico. Iodine is found in the human body in the compound thyroxin in the thyroid gland; deficiency of iodine in diet causes enlargement of this gland (goitre).

Iodine is rarely prepared in the laboratory; the method used is the oxidation of an iodide by manganese(IV) oxide and sulphuric acid, for example with sodium iodide:

2NaI + MnO2 + 3H2SO4 -> MnSO4 + 2NaHSO4 + I2 + 2H2O

The iodine distils off and can be collected on a cooled surface. It may be purified by sublimation in vacuo.

This reaction is also used on a large scale, to obtain iodine from seaweed. The ash from burnt seaweed ("kelp1) is extracted with water, concentrated, and the salts other than iodides (sulphates and chlorides) crystallise out. The more soluble iodides remain and the liquor is mixed with sulphuric acid and manganese dioxide added; the evolved iodine distils off and is condensed.

Most iodine produced commercially comes from the sodium iodate(V) remaining after sodium nitrate has been crystallised from Chile saltpetre. The iodate(V) is first reduced to iodide by blowing sulphur dioxide into the solution (or by addition ofsodium sulphite):

ioj + ssor ->r + ssoj-

More iodate is then added, and with the sulphuric acid formed (or added if sodium sulphite is used), iodine is liberated :

IO + 51" + 6H+ -* 3I2 4- 3H2O

320 GROUP VII: THE HALOGENS

Alternatively, the iodide is precipitated as copper(I) iodide by addition of copper(II) sulphate, in presence of sulphite, thus:

21" + 2Cu2+ + SOi- -h H2O -> 2CuI 4- SOj" + 2H+

The iodine is then liberated by heating the copper(I) iodide with sulphuric acid and iron(III) oxide:

2CuI + 6H2SO4 -f 2Fe2O3 -> 2CuSO4 + 4FeSO4 4- 6H2O 4- I2

The copper(II) sulphate is recovered and used to precipitate more copper(I) iodide.

Iodine and its compounds are. relative to the other halogens, costly substances.

Iodine is a dark-coloured solid which has a glittering crystalline appearance. It is easily sublimed to form a bluish vapour in vacno. but in air, the vapour is brownish-violet. Since it has a small vapour pressure at ordinary temperatures, iodine slowly sublimes if left in an open vessel; for the same reason, iodine is best weighed in a stoppered bottle containing some potassium iodide solution, in which the iodine dissolves to form potassium tri-iodide. The vapour of iodine is composed of I2 molecules up to about 1000 K; above this temperature, dissociation into iodine atoms becomes appreciable.

Like bromine, iodine is soluble in organic solvents, for example chloroform, which can be used to extract it from an aqueous solution. The iodine imparts a characteristic purple colour to the organic layer; this is used as a test for iodine (p. 349). NB Brown solutions are formed when iodine dissolves in ether, alcohol, and acetone. In chloroform and benzene a purple solution is formed, whilst a violet solution is produced in carbon disulphide and some hydrocarbons. These colours arise due to charge transfer (p. 60) to and from the iodine and the solvent organic molecules.

CHARACTERISTIC REACTIONS OF THE HALOGENS WITH HYDROGEN

All the halogens combine directly with hydrogen, the reaction generally occurring with less vigour in the series F2, C12, Br2,12-

The rate of reaction between fluorine and hydrogen varies a great deal with conditions. Solid fluorine and liquid hydrogen explode even at 21 K but mixing of the gases at room temperature in the dark may preclude any reaction; however a reaction can

GROUP VII: THE HALOGENS 321

occur with explosive violence. A chain mechanism is likely for the reaction.

Mixtures of chlorine and hydrogen react only slowly in the dark but the reaction proceeds with explosive violence in light. A suggested mechanism for the photochemical chain reaction is:

C12 + hv -» 2Cr

Cl* + H2 ->HC1 + H-

H' + C12 -> HC1 4- Cl* and so on.

In the presence of charcoal, chlorine and hydrogen combine rapidly, but without explosion, in the dark. A jet of hydrogen will burn in chlorine with a silvery flame and vice versa.

The affinity of chlorine for hydrogen is so great that chlorine will react with many compounds containing this element, for example hydrocarbons (a wax taper burns in chlorine).

Chlorine substitutes the hydrogen of methane giving successively the chlorides CH3C1, CH2C12, CHC13 and CC14. It is to be noted that if a hydrocarbon is unsaturated, chlorine atoms will first add

Bromine, like chlorine, also undergoes a photochemical chain reaction with hydrogen. The reaction with bromine, however, evolves less energy and is not explosive.

Like chlorine, bromine can displace hydrogen from saturated hydrocarbons, though not as readily, and adds on to unsaturated ones.

Iodine and hydrogen react reversibly to give hydrogen iodide:

H2 + I2 ^ 2HI

This equilibrium has been extensively studied by Bodenstein. Unlike the other halogen-hydrogen reactions, it is not a chain reaction but a second order, bimolecular, combination.

Iodine does not replace hydrogen from saturated hydrocarbons directly, as do both chlorine and iodine.

WITH ELEMENTS OTHER THAN HYDROGEN

Fluorine is exceedingly reactive and combines vigorously with most elements. Some ignite spontaneously in gaseous fluorine at room

322 GROUP VII: THE HALOGENS

temperature, for example K, B, Si P. S, I. Other elements ignite when gently warmed in the gas. for example Ag and Zn, and even gold, platinum and xenon are attacked if heated strongly. Graphite is attacked slowly—hence the use of special electrodes in the extraction of fluorine—and diamond only above 950 K. Some metals, for example copper and nickel alloys, become coated with a superficial layer of fluoride. This prevents further reaction and hence vessels of these materials are used for the preparation and storage of fluorine. Oxygen and nitrogen do not combine directly with fluorine.

Chlorine reacts with most elements, both metals and non-metals except carbon, oxygen and nitrogen, forming chlorides. Sometimes the reaction is catalysed by a trace of water (such as in the case of copper and zinc). If the element attacked exhibits several oxidation states, chlorine, like fluorine, forms compounds of high oxidation state, for example iron forms iron(III) chloride and tin forms tin(IV) chloride. Phosphorus, however, forms first the trichloride, PC13, and (if excess chlorine is present) the pentachloride PC15.

Bromine has a lower electron affinity and electrode potential than chlorine but is still a very reactive element. It combines violently with alkali metals and reacts spontaneously with phosphorus, arsenic and antimony. When heated it reacts with many other elements, including gold, but it does not attack platinum, and silver forms a protective film of silver bromide. Because of the strong oxidising properties, bromine, like fluorine and chlorine, tends to form compounds with the electropositive element in a high oxidation state.

Iodine, though generally less reactive than bromine, combines directly with many elements, for example silver, gold and aluminium, forming iodides. Mercury is also attacked and mercury(I) iodide. Hg2I2, is first formed but in the presence of excess iodine this is oxidised to mercury(II) iodide, HgI2. Iodine and phosphorus (red and white)react in the presence of water to form first phosphorus(III) iodide, PI3, which is then hydrolysed to yield hydrogen iodide (p. 333). Iodine reacts with the other halogens to form interhalogen compounds (p. 345).

WITH COMPOUNDS

The reactions with water

The oxidising power of fluorine is seen in its reaction with water: in the liquid phase, water reacts to give hydrogen peroxide and some

GROUP VII: THE HALOGENS 323

fluorine monoxide (see below); in the gas phase ozone and oxygen are produced

3H2O + 3F2 ^6HF + O3

Recent work indicates the existence offluoric(l) acid. HFO, formed by the reaction of fluorine and water at 273 K. The acid forms colourless crystals, m.p. 156K, is very unstable, and has, as expected, very strong oxidising properties.

Chlorine and bromine are both moderately soluble in water, and on crystallisation these solutions give solid hydrates with the halogen molecules occupying cavities within a modified ice lattice. Iodine is only slightly soluble in water in which it forms a brown solution (brown solutions are also formed in ether, alcohol and acetone). The aqueous solutions of chlorine and bromine are good oxidising agents. Chlorine, and to a lesser extent bromine, reacts reversibly with water to give a mixture of acids, for example :

The presence of chloric(I) acid makes the properties of "chlorine water' different from those of gaseous chlorine, just as aqueous sulphur dioxide is very different from the gas. Chloric(I) acid is a strong oxidising agent, and in acid solution will even oxidise sulphur to sulphuric acid; however, the concentration of free chloric(I) acid in 'chlorine water' is often low and oxidation reactions are not always complete. Nevertheless when "chlorine' bleaches moist litmus, it is the chloric(I) acid which is formed that produces the bleaching. The reaction of chlorine gas with aqueous bromide or iodide ions which causes displacement of bromine or iodine (see below) may also involve the reaction

2r + HC1O + HC1 -> 2CP + I2 + H2O

since water is present to produce the two acids. Chlorine water loses its efficiency as an oxidising agent on standing because the chloric(I) acid decomposes. There are two possible ways of decomposition :

324 GROUP VII: THEHALOGENS

aqueous hydrogen peroxide). Bromine water undergoes a similar decomposition in sunlight and oxygen is evolved but in general it is more stable than chlorine water and the equilibrium

Br2 + H2O ^ HBr + HBrO

lies further to the left.

If 'chlorine water' is boiled the chloric(I) acid decomposes as above, but a little may break down into steam and the acid anhydride, dichlorine monoxide:

2HC1O ^ C12O + H2O

The smell of chlorine water, somewhat different from that of gaseous chlorine, may be due to minute amounts of evolved dichlorine monoxide:

The reactions with alkalis

Oxygen difluoride, OF2, is obtained when gaseous fluorine is passed through very dilute (2%) caustic soda solution:

2F2 + 2NaOH -» 2NaF + F2O + H2O

but with more concentrated alkali, oxygen is formed:

2F2 + 4NaOH -> 4NaF + 2H2O + O2

The reactions of the other halogens can be summarised in the two equations:

(Reaction (11.4) is really a disproportionation reaction of the halate(I) anion: 3XO~ -> 2X~ + XO~.) Reaction (11.3) is favoured by the use of dilute alkali and low temperature, since the halate(I) anions, XO~ are thermally unstable and readily disproportionate (i.e. reaction (11.4)). The stability of the halate(I) anion, XO~, decreases from chlorine to iodine and the iodate(I) ion disproportionates very rapidly even at room temperature.

The formation of halate(V) and halide ions by reaction (11.4) is favoured by the use of hot concentrated solutions of alkali and an excess of the halogen.

When chlorine is passed over molten sodium or potassium hydroxide, oxygen is evolved, the high temperature causing the chlorate(V) ion to decompose:

2CKK -+2CP +302

G R O U P V I I : THE HALOGENS 325 OTHER DISPLACEMENT AND OXIDATION REACTIONS

Many of the reactions of halogens can be considered as either oxidation or displacement reactions; the redox potentials (Table 11.2) give a clear indication of their relative oxidising power in aqueous solution. Fluorine, chlorine and bromine have the ability to displace hydrogen from hydrocarbons, but in addition each halogen is able to displace other elements which are less electronegative than itself. Thus fluorine can displace all the other halogens from both ionic and covalent compounds, for example

2NaCl 4- F2 ~» 2NaF + C12

2 - - 7 —Cl + F -> 2 ~C—F + C1

2

and oxygen from water and silica :

SiO2 4- 2F2 -» SiF4 + O2

The reaction with silica explains why fluorine reacts with glass and quartz, but if these are rigorously freed from adsorbed water, the reaction is very slow ; hence dry fluorine can be manipulated in dry glass apparatus but all glass taps must be lubricated with fluoroearbon grease since hydrocarbon greases would be attacked. The very strong oxidising properties of fluorine in aqueous systems are seen in reactions such as the conversion of chlorate(V) to chlorate- (VII), chromium(III) to dichromate(VI) and the oxidation of the hydrogensulphate ion, HSO^, to peroxodisulphate :

2HSO4 + F2 -> S2Oi~ + 2HF

Also, in anhydrous conditions, silver reacts with fluorine and forms silver difluoride AgF2 and cobalt gives cobalt(III) fluoride, CoF3, these metals showing higher oxidation states than is usual in their simple salts.

Chlorine has a lower electrode potential and electronegativity than fluorine but will displace bromine and iodine from aqueous solutions of bromide and iodide ions respectively :

C12 + 2Br~ -> 2Cr -f Br2

Chlorine reacts directly with carbon monoxide to give carbonyl chloride (phosgene) :

CO + C12 -* COC12

and sulphur dioxide to give sulphur dichloride dioxide:

SO2 + Cl'2 -* SO2C12