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

ACIDS AND BASES: OXIDATION AND REDUCTION 107

150

100

070

Figure 4.7. Potentiometric titration of Fe(II) with Ce(lV)

number of redox reactions; quantitatively they have several advantages over ordinary indicator methods.

Thus, for example, an analysis using coloured solutions can be carried out, where an indicator cannot be used. Moreover, it is not easy to find a redox indicator which will change colour at the right point. Potentiometric methods can fairly readily be made automatic.

TESTS FOR REDUCING AND OXIDISING AGENTS

The redox properties of all reagents are relative and a given reagent may be both a reducing and an oxidising agent depending upon the reaction in which it is involved. Thus, for example, sulphur dioxide in aqueous solution is an oxidising agent with respect to hydrogen sulphide, but a reducing agent with respect to acidified potassium dichromate(VI) solution. Similarly hydrogen peroxide in acidic solution is an oxidising agent relative to iron(II) ions but a reducing agent relative to manganate(VII) ions in aqueous solution. However, it is convenient to establish approximate 'reference points' for laboratory reagents, which can then be loosely classified as follows:

Reagents are reducing if they:

1.Decolorise a solution of potassium manganate(VII) acidified with dilute sulphuric acid.

2.Turn a solution of potassium dichromate(VI) acidified with dilute sulphuric acid from orange to green.

108 ACIDS AND BASES: OXIDATION AND REDUCTION

3. Change a solution of iron(III) in aqueous solution to iron(II).

Reagents are oxidising if they:

1.Liberate iodine from a potassium iodide solution acidified with dilute sulphuric acid.

2.Convert iron(II) to iron(III) in aqueous acid solution.

QUESTIONS

1.(a) The following are standard redox potentials in volts in 1 N acid solution for the reactions

Mn+ + xe~ -> M(n"x)+ (symbolised as Mfl+/M("~-x)+), where, for example, the process

Na+ + e~ -> Na (symbolised as Na+/Na) is defined as having a large negative potential:

Cr2+/Cr - 0.9 V, Mn2+/Mn - 1.2 V, Cr3+/Cr2+ -0.4V, Mn3+/Mn2+ + 1.5V,

Fe2 + /Fe -0.4V,

Fe3+/Fe2+ + 0.8V. Use these data to comment upon:

(i)the stability in acid solution of Fe3 + towards reducingagents as compared to that of either Cr3+ or Mn3+ ;

(ii)the ease with which metallic iron can be oxidised to iron(II) (ferrous) ions compared to the similar process for either metallic chromium or metallic manganese;

(iii)the result of treating a solution containing either chromium(II) (chromous) or manganese(II) (rnanganous) ions with a solution containing iron(III) (ferric) ions.

(b) The following equations represent four chemical reactions involving redox processes:

(i)3N2H4 + 2BrO3^ -> 3N2 + 2Br" + 6H2O,

(ii)5As2O3 + 4MnO4 + 12H+ -> 5As2O5 + 4Mn2+ -f 6H2O,

(iii)SO2 + I2 + 2H2O -^ H2SO4 + 2HI,

(iv)VOJ" + Fe2+ + 6H+ -> VO2+ -f Fe3+ + 3H2O

112 HYDROGEN

molecule and allows reaction to proceed. The reactions of hydrogen will now be examined in more detail.

REACTIONS WITH ELECTROPOSITIVE METALS

These give ionic or salt-like hydrides, for example

2Na + H2 -> 2NaH

These solid ionic hydrides (having an ionic lattice and containing the hydride ion H ~) react with water, for example

CaH2 + 2H2O -» Ca(OH)2 + 2H2

i.e.

+ H

We can see that the hydride ion H ~ functions as a very strong base (p, 89) withdrawing a proton from the water molecule and uniting with it to give H2, i.e. H~ + H^ -> H2, a highly exothermic process. It follows that we cannot use these ionic hydrides in aqueous solutions; however, some of them (notably lithium hydride. LiH) can be used in suspension in organic solvents as reducing agents, and others can be converted to complex hydrides which can be used in solution (see below),

The existence of the hydride ion is shown by electrolysis of the fused salt when hydrogen is evolved at the anode. If calcium hydride is dissolved in another fused salt as solvent, the amount of hydrogen

REACTIONS WITH TRANSITION METALS

Most of these metals only react with hydrogen on heating; the first stage of reaction is the taking of hydrogen on to the metal surface, whereby the hydrogen molecules become attached as hydrogen atoms—a process known as chemisorption, With some metals reaction can proceed further, and hydrogen atoms penetrate into the metal lattice and occupy positions between the metal atoms— interstitial positions, as shown in Figure 5.1.

If all these 'holes' were filled, the hydrogen-metal ratio would be a definite and fixed number; in practice, this rarely happens, and

HYDROGEN 113

these metal hydrides or interstitial hydrides may have variable composition (for example TiHx 7), depending on the uptake of hydrogen, i.e. they are non-stoichiometric. One further property in particular distinguishes these metal hydrides from the ionic hydrides; in the latter, uptake of hydrogen is not only quantitative but causes a contraction, i.e. the centres of the metal atoms (which become

Figure 5.1. Interstitial positions between layers of metal atoms

cations) move closer together—the metal lattice is, as it were,drawn together. In the metal hydrides, there is no such contraction, and, indeed, the metal atoms may move apart slightly. Hence formation of an ionic hydride leads to an increase in density, but formation of a metal hydride causes a decrease in density.

REACTIONS WITH NON-METALS AND WEAKLY ELECTROPOSITIVE METALS

Most of the elements of Groups III to VII form hydrides which are essentially covalent. Some examples are Group IV, methane CH4; Group V, phosphine PH3; Group VI, hydrogen sulphide H2S; Group VII, hydrogen chloride, HC1. There are several points to notice about these covalent hydrides. First, they are nearly all volatile liquidsor gases; but the simple hydridesNH3, H2O and HF, formed from the head elements of Groups V-VI1, show hydrogen bonding characteristics which make them less volatile than we should expect from the small size of their molecules (p. 52).

Secondly, the ability to form more than one hydride falls off as we go across a period. Thus, in Period 1. boron and carbon both form whole families of hydrides, nitrogen forms three (ammonia. NH3; hydrazine. N2H4; hydrazoic acid. N3H). oxygen two (H2O. H2O2) and fluorine one (HF). Again, as we descend a group, the energetic stability of the hydrides decreases—indeed, many hydrides are endothermic. and need indirect methods to supply the necessary energy for their preparation. In Group IV, methane is exothermic,

114 HYDROGEN

the others are endothermic and plumbane PbH4. the last hydridein the group, is almost too unstable to exist at all. (We shall note some of the methods needed to prepare these less stable hydrides in later chapters.) Since the stability of the typical hydride (i.e. that in which the element shows its group valency) falls off.it is hardly surprising to find that the lower elements in a group do not form families of hydrides (for example, in Group IV carbon and silicon form numerous hydrides, germanium forms a few.tin forms one (stannane. SnHJ and lead just manages to form PbH4).

The most important trend to be noted in the covalent hydrides is

This change in properties cannot be simply accounted for in terms of bond energies; the mean X—H bond energy increases from nitrogen to fluorine, and hydrogen fluoride has a large bonddissociation energy (566kJmol~1). But we note that in the CH4 molecule there are no lone pairs of electrons—all four valency electrons are involved in bonding. In ammonia, there is one lone pair, which as we have seen can be donated either to a proton (making ammonia a Lowry-Br0nsted base, NH3 + H+ ^NH^) or to another acceptor molecule (making ammonia a Lewis base, p. 91). The molecules H2O and HF have two and three lone pairs respectively; falling-off of base strength implies that the presence of more than one lone pair reduces the donor power of the molecule. But, obviously, the appearance of acidic behaviour implies that the bond X—H is more readily broken heterolytically i.e. to give X~ + H+. We may ascribe this to polarity of the bond, i.e. by saying that the pair of electrons in the covalent H—F bond is closer to the fluorine than to the hydrogen. Unfortunately, there is no very sure

whole is polar in, presumably, the sense H—F, but this does not necessarily tell us about the bond polarity). Another way of describing this trend towards acidity is to say that the electronegativity of the element increases from carbon to fluorine. We may simply note that this trend to acidity is also apparent in other periods, for example, in Period 3, silane SiH4 is non-acidic and non-basic.