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

306 GROUP VI

Sulphur hexafluoride, SF6, is chemically unreactive, resembling nitrogen, and is unaffected by heat, water, fused alkalis, and many heated metals. This stability is attributed to the high S—F bond strength and to the inability of attacking reagents, such as water, to coordinate to the covalently saturated sulphur (see SF4 below). It finds a use as a high-voltage gaseous insulator.

Both selenium hexafluoride and tellurium hexafluoride are more reactive than sulphur hexafluoride. Tellurium hexafluoride is slowly hydrolysed bywater to telluric(VI) acid and on heating it decomposes to fluorine and the tetrafluoride.

The tetrafluorides of the elements can be prepared. They are all less stable than the corresponding hexafluorides and are hydrolysed readily by water. They can all be used as fluorinating agents and sulphur tetrafluoride is extensively used for this purpose, for example the fluorination of organic carbonyl groups:

c=o

The structure of sulphur tetrafluoride, and probably also SeF4 and TeF4, is trigonal bipyramidal with one position occupied by a lone pair of electrons :

F

Chlorides

Sulphur and selenium form the chlorides disulphur dichloride S2C12 and diselenium dichloride Se2Cl2. They are made by the direct combination of the elements. Both are covalent, yellow liquids which are readily hydrolysed by water:

S2C12 + 3H2O -> 2HC1 + H2S + SO^ + 2H+

(Further reaction between hydrogen sulphide and the sulphite ion yields sulphur together with thionic acids):

2Se2Cl2 + 3H2O -> H2SeO3 + 3Se 4- 4HC1

GROUP VI 307

Diselenium dichloride acts as a solvent for selenium. Similarly disulphur dichloride is a solvent for sulphur and also many other covalent compounds, such as iodine. S2C12 attacks rubber in such a way that sulphur atoms are introduced into the polymer chains of the rubber, so hardening it. This product is known as vulcanised rubber. The structure of these dichlorides is given below:

X= S or Se (cf. H2O2, p. 279).

Sulphur and tellurium form a chloride of formula XC12. Sulphur

dichloride SC12 is a red liquid at room temperature whilst the corresponding tellurium compound is a black solid.

A number of bromides and iodides are known but there are no sulphur iodides.

Halide oxides

A number of halide oxides are formed by sulphur and selenium but only one is considered here.

SULPHUR DICHLORIDE OXIDE, THIONYL CHLORIDE, SOC12

This is an important laboratory reagent and has the structure shown below :

a a

It is prepared by heating together phosphorus pentachloride and a sulphite, for example calcium sulphite :

2PC15 + CaSO3 -» 2POC13 + CaCl2 4- SOC12

The oxide dichloride, b.p. 351 K, is separated from the less volatile phosphorus oxychloride by a fractional distillation.

Sulphur oxide dichloride is a colourless liquid which fumes in moist air. It is hydrolysed by water to give a mixture of sulphurous

hydrochloric acids :

308 GROUP VI

SOC12 + 2H2O ^ 4H+ + SOr +2C1~

Hence on warming, sulphur dioxide is evolved.

Sulphur oxide dichloride is used as a chlorinating agent inorganic chemistry, for example in the preparation of acid chlorides:

CH3COOH + SOC12 -> CH3COC1 + SO2f + HClt

The advantage of the method, readily seen from the equation, is that the other products of the reaction are gaseous and escape. Hence equimolar quantities of reactants are used.

A somewhat similar reaction is the power of sulphur oxide dichloride to remove water ofcrystallisation from hydrated chlorides, the hydroxyl groups of the water molecule reacting as do those in the acid molecules in the above reaction.

The action is a general one and may be written thus:

MCln.xE2O + xSOC!2 -> MCI, + xSO2t 4- 2xHClT

The reaction provides a valuable method of preparing anhydrous chlorides of metals. It has been used to prepare the anhydrous chlorides of copper(II), zinc, cadmium, chromium(III), iron(III), cobalt(II) and nickel.

In both reactions above, the oxide dichloride is refluxed with the acid or the hydrated chloride; the sulphur dioxide and hydrogen chloride pass off and any unused sulphur oxide dichloride is distilled off in vacua.

TESTS FOR SULPHUR

Oxidation of a sulphur compound with concentrated nitric acid yields sulphuric acid or a sulphate, which can be tested for with barium chloride. This can be used to estimate the sulphur.

QUESTIONS

1. How would you obtain a sample of pure ozone? Account for the conditions used in your method of preparation. What is the arrangement of oxygen atoms in an ozonide and what evidence would you cite in support of the structure you suggest?

(L,A)

2. Comment on and, where you are able, suggest reasons for the following observations:

GROUP VI 309

(a)Na2O dissolves in water to give an alkaline solution: C12O dissolves in water to give an acidic solution.

(b)C12O is a gaseous oxide, its molecule being V-shaped: Na2O is an ionic compound which has an infinite 3-dimensional lattice structure.

(c)A12O3 forms a hydrated oxide which is basic, but the addition of alkali produces a solution containing the aluminate anion, A1O2-

(d)SiO2 and CO2 are both acidic oxides. SiO2 is a solid of high melting-point, whereas CO2 is a gas.

(e)N2O is a gaseous, neutral oxide, its molecule being linear.

(CA)

3. Give an explanation of the following observations:

(a)An aqueous solution of sodium sulphide smells of hydrogen sulphide.

(b)When hydrogen sulphide is bubbled through an acidified solution of a cobalt(II) salt, no precipitate isformed,but a black precipitate is produced when the solution is made slightly alkaline.

(c)When hydrogen sulphide is bubbled through an aqueous solution of an aluminium(III) salt, a white precipitate of aluminium(III) hydroxide is obtained.

(d)Hydrogen sulphide (formula weight 34)is a gas, water (formula weight 18) is a liquid.

4.Describe one laboratory method for the preparation of a dilute solution of hydrogen peroxide.

In what way does a solution of hydrogen peroxide react with

(a)chlorine water, (b) potassium permanganate solution, (c) potassium dichromate solution, (d) hydrogen sulphide? 50 cm3 of an aqueous solution of hydrogen peroxide were treated with an excess

of potassium iodide and dilute sulphuric acid; the liberated iodine was titrated with 0.1 M sodium thiosulphate solution and 20.0 cm3

were required. Calculate the concentration of the hydrogen peroxide solution in g I"1.

(1MB, A)

Table 11.1

SELECTED PROPERTIES OF THE ELEMENTS

312 GROUP VI!: THE HALOGENS

1 Atomic radius ionic radius

0

g

.

o

400

*:

CL

cL 200

E

Ist ionisation energy

enthalpy of otornisation

o

E _o

Numerous ionic compounds with halogens are known but a noble gas configuration can also be achieved by the formation of a covalent bond, for example in halogen molecules, X2, and hydrogen halides, HX. When the fluorine atom acquires one additional electron the second quantum level is completed, and further gain of electrons is not energetically possible under normal circumstances, i.e.

GROUP VII: THE HALOGENS 313

promotion to 3s requires too much energy.Thus fluorine is normally confined to a valency of 1 although in some solid fluorides bridge structures M—F—M are known in which fluorine acquires a covalency of 2.

All the remaining halogens have unfilled d orbitals available and the covalency of the element can be expanded. Compounds and complex ions are formed both with other halogens and with oxygen in which the halogen can achieve a formal oxidation state as high as + 7, for example chlorine has formal oxidation states of +1 in the chlorate(I) anion CIO" ; -f 5 in the chlorate(V) anion CIO3, and + 7 in the chlorate(VII) anion C1OJ.

ELECTRODE POTENTIALS AND REACTIVITY OF THE HALOGENS

One surprising physical property of fluorine is its electron affinity which, at —333kJmol~l, is lower than that of chlorine, —364 kJmol"1, indicating that the reaction X(g) + e" -» X~(g) is more exothermic for chlorine atoms. In view of the greater reactivity of fluorine a much higher electron affinity might reasonably have been expected. The explanation of this anomaly is found when the steps involved in a complete reaction are considered. For example, with a Group I metal ion M+(g) the steps to form a crystalline solid are,

(3) X~(g) 4- M + (g) -> M+ X~(s) Lattice enthalpy

the overall reaction being

e~ +±X2(g) + M + (g)-M+X-(s)

The enthalpies for the reactions of chlorine and fluorine are shown graphically in Figure 112 as the relevant parts of a Born-Haber cycle. Also included on the graph are the hydration energies of the two halogen ions and hence the enthalpy changes involved in the reactions

iX2(g) + <?-^-

The very low bond dissociation enthalpy of fluorine is an important factor contributing to the greater reactivity of fluorine. (This low energy may be due to repulsion between non-bonding electrons on the two adjacent fluorine atoms.) The higher hydration and lattice enthalpies of the fluoride ion are due to the smaller size of this ion.

314

200 r

-200

-400

Enthalpy -600 (kJmol

-800

-IOOO -

-1200L

Figure 11.2. Formation of fluoride and chloride iom from the elements

Kev:

+ aq + e~ -> X'(aq)

Table 11.2

ENTHALPY DATA FOR HALiDE ION FORMATION IN AQUEOUS SOLUTION

GROUP VII: THE HALOGENS 315

Electron affinity and hydration energy decrease with increasing atomic number of the halogen and in spite of the slight fall in bond dissociation enthalpy from chlorine to iodine the enthalpy changes in the reactions

both decrease and the reaction becomes less exothermic. Hence the reactivity and the electrode potential (which is closely related to reaction (11.2) and indeed defined by it under standard conditions) decrease from fluorine to iodine. Table 11.2 gives the enthalpy change (kJ mol~*) for each halogen in reaction (1 1.2).

ELECTR ONEGATIYITY

The large value for fluorine, and the marked decrease from fluorine to iodine, are points to be noted. The high value for fluorine means that the bond between an element M and fluorine is likely to be more ionic (more polar) than a bond formed by M with any other elements. The low value for iodine indicates the possibility that iodine may be electropositive in some of its compounds.

OXIDISING PROPERTIES

For fluorine, the reaction

is energetically highly favourable for the formation both of X~ and of X~(aq). Hence gaseous fluorine is highly reactive towards metals, giving essentially ionic fluorides; and in solution (as its high electrode potential indicates) it is one of the most powerful oxidising agents, oxidising water very readily (p. 100).Hence the fluoride ion cannot be converted into fluorine in aqueous solution ; electrolysis of a found fluoride must be used. In contrast, iodide ions in solution are readily oxidised even by air (Table 4.3).

HALOGENS AS LIGANDS

The small fluoride ion shows a great tendency to act as a ligand and form complex ions, for example [A1F6]3~, [PF6]~, [FeF6]3~ in