introduction-to-inorganic-chemistry

16.5 Further aspects

Aqueous ammonia as a base

Aqueous ammonia is a base in the classical sense, but poses problems on the stoicheiometric scheme. On this scheme, it is necessary to postulate the formation of ammonium hydroxide by the reaction

NH3 + H2O → NH4OH

the NH4OH being the base and NH3 the base-former. Unfortunately, no evidence for the formation of a species of formula NH4OH has been obtained.

On the electrolytic scheme, however, it is not necessary to postulate the intermediate formation of NH4OH. The properties of aqueous ammonia can be well accounted for in terms of the equilibrium

NH3 + H2ONH4+ + OH–

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On this scheme, NH3 is again a base-former - it is the solution that is basic, as if there were a base of formula NH4+OH–.

Note, however, that tertiary alkyl amines, NR3, on treatment with an alkyl halide, RX, give quaternary ammonium salts, [NR4]X, analogous to ammonium salts, [NH4]X, and these, on treatment with moist silver oxide, give the corresponding hydroxides

2[NR4]X + Ag2O + H2O → 2[NR4]OH + 2AgX

These are very strong alkalis.

On the nature of the hydrogen ion in solution

While the hydration of all ions is acknowledged, special attention has been paid to the hydration of the hydrogen ion, for the following reason:

Anhydrous acids, unlike anhydrous bases and salts, are not ionic. For example, hydrogen chloride is a gas, and its molecules have a low polarity (H0.2+Cl0.2−, Chap. 5). While molten sodium hydroxide and sodium chloride are good electrolytic conductors, liquid hydrogen chloride scarcely conducts electricity at all:

A few anhydrous acids are good conductors (e.g. 100% sulfuric acid), but this is unusual.

Thus, in the case of the hydrogen ion, it looks as though hydration is essential for its formation. Put another way, it looks as if the H+ ion cannot exist on its own, but only in association with one or more water molecules.

The number of water molecules that it is associated with in a particular solution is difficult to determine. That one molecule is enough for its formation, however, can be demonstrated by examining the hydrates of strong acids like hydrochloric, nitric, and sulfuric. The lowest hydrates of these are as follows:

In the liquid state, these of course conduct electricity, and their structures in the solid state have been shown to be H3O+Cl–, H3O+NO3–, and H3O+HSO4– respectively.

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In view of the above, some chemists insist that the hydrogen ion in water should not be called the hydrogen ion, but the hydroxonium ion, and should not be written as H+ but as H3O+. In doing this they recognize that the H3O+ ion will be hydrated further, but believe that no formulation less than this will do justice to the difference in properties between anhydrous and aqueous acids. In practice, however, most chemists continue to write H+, and to call it the hydrogen ion. Whether the latter approach is lazy or the former pedantic is a matter for discussion.

Some chemists further insist that compounds like HCl, HNO3, and H2SO4 should not strictly be called acids. The true acids are (H3O)Cl, (H3O)NO3, and (H3O)HSO4, and HCl, HNO3, and H2SO4 are really acid-formers. On this basis, the latter should not be called hydrochloric acid, nitric acid, and sulfuric acid, but hydrogen chloride, hydrogen nitrate, and hydrogen sulfate. Few chemists would go as far as this, though “hydrogen chloride” has become the generally accepted name for anhydrous HCl.

IUPAC nomenclature

For other than very common oxoacids, IUPAC prefer to name acids as compounds of H+, e.g. hydrogen chloride. For common oxoacids, however, they recommend the traditional name, or a more systematic alternative, e.g.

Most users stick to the traditional names.

16.6 Lowry-Brønsted conception of acids and bases

This is a broader conception than those we have discussed so far. It is one that gets over the distinction between acids and acid-formers or bases and base-formers, though it does so at the expense of other difficulties.

On this scheme,

an acid is a molecule or ion that supplies protons (i.e. a proton donor);

a base is a molecule or ion that accepts protons (i.e. a proton acceptor).

This means that the molecules HCl, HNO3, and H2SO4 are now acids, not acid-formers:

HCl + H2O → H3O+ + Cl– etc.

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Similarly, the species CO32– and NH3 are now bases and not base-formers:

CO32– +H2O HCO3– + OH– etc.

Further, the ion H3O+ is itself an acid, the ion OH– is a base, and the H2O molecule is both an acid and a base. These functions can all be seen in the self-ionization

2H2OH3O+ + OH–

On the Lowry-Brønsted scheme, an acid-base reaction effectively involves a competition for protons representable by an equilibrium of the type

HXx + Yy Xx–1 + HYy+1

where x and y represent the charges on HX and Y respectively. Here Xx–1 is called the “conjugate base” of the acid HXx, and HXx the “conjugate acid” of the base Xx–1. For example, NH4+ is the conjugate acid of NH3, and CH3COO– is the conjugate base of CH3COOH, and the reaction between them involves the competition between the two bases (NH3, CH3COO–) for the proton:

NH4+ + CH3COO–NH3 + CH3COOH

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Advantages of the Lowry-Brønsted scheme are as follows:

(i)It takes in acid-formers and base-formers.

(ii)It is applicable to other solvents besides water. For example, in liquid ammonia, NH4+ acts as a proton donor, NH2– acts as a proton acceptor, and NH3 functions in both ways. The LowryBrønsted definitions thus apply unchanged.

Disadvantages of the scheme are:

(i)It loses the category of salt. The scheme is essentially dualistic in character, and has no place for a third category. Sodium chloride is regarded merely as an aggregate of ions, one of which (Na+) is neutral, and the other (Cl–) a base - the conjugate base of HCl.

(ii)The scheme also loses the idea of neutralization. Though the latter could be defined as the complete transfer of a proton from an acid to a base (cf. the stoicheiometric conception), it would include some processes for which the name neutralization hardly seems appropriate, e.g.

HCl + H2O → H3O+ + Cl–

In view of the disadvantages, it can be questioned whether it is worth trying to extend the ideas of acid and base in the way that Lowry and Brønsted do. Chemistry might be better served if the classification of molecules and ions according to their ability to accept or donate protons – which is undoubtedly a very useful classification – were to be divorced from the classification of substances into acids and bases, and be given a terminology of its own. In fact, such terminology lies readily to hand, the terms “proton donor” and “proton acceptor” being entirely suitable for this purpose.

Note

IUPAC are trying to dissuade chemists from calling H+ a proton. A proton is strictly 1H+. Ordinary samples of acid contain 99.985% of 1H+ and 0.015% of 2H+. IUPAC recommend the following names:

Exercise

Classify each of the following compounds on each of the four schemes given above (classical, stoicheiometric, electrolytic, Lowry-Brønsted):

(i) NaCN; (ii) NH4Cl; (iii) NaHCO3.

(Click here for answers http://bit.ly/hStxuL or see Appendix 1)

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16.7 Solvent-system conception of acids and bases

The electrolytic conception of acids and bases can be extended to other ionizing solvents by means of the following definitions:

An acid is a compound that ionizes to give the cation natural to the solvent.

A base is a compound that ionizes to give the anion natural to the solvent.

The ions “natural” to a solvent are those produced in its self-ionization. For some of the more important ionizing solvents, these are as follows.

Thus, NH4Cl is an acid in liquid ammonia, KNH2 is a base, KHSO4 is a base in 100% sulfuric acid, etc.

This extension of the ideas of acids and bases is a very useful one, and does not stretch the original conception too much. Note, however, that substances can change their character in going from one solvent to another. Thus, NH4Cl is a salt in water, but an acid in liquid ammonia. Similarly, KHSO4 is an acid salt in water, and a base in 100% H2SO4. A third example is CH3COOH, which is an acid-former in water

CH3COOH + H2OH3O+ + CH3COO–

and a base-former in sulfuric acid

CH3COOH + H2SO4H3CO2H2+ + HSO4–

This being so, it is best to specify the solvent as well as the class of substance, at least when the solvent is not water, e.g.

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16.8 Other conceptions of acids and bases

The remarks made at the end of the section on the Lowry-Brønsted definition apply with even greater force to two other definitions that have been proposed for acids and bases, namely:

These take us a long way from the original conceptions of acid and base, and might be better divorced from these ideas completely. At the very least, they should only be used in a qualified way, as in the expressions “Lewis acid” and “Lewis base”.

For further details about these two definitions, see textbooks.

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17 Oxidation and reduction

17.1 Basic theory

The word “oxidation” originally meant the gain of oxygen by an element or compound, and “reduction” the loss of oxygen, as in the processes:

There are, however, other processes that are similar to the above, for example:

There is a close parallel between reactions (1a) - (1d) and (2a) - (2d), not only in terms of the equations for the reactions, but also in terms of the nature of the reactants and products. Thus in (1a) and (2a) magnesium metal combines with a nonmetal to form a product that is electrically salt-like.

The basis of the similarity of (1a) - (1d) and (2a) - (2d) is that chlorine is an electronegative element like oxygen. Thus (1a) - (1d) and (2a) - (2d) may be classed together if we extend the idea of oxidation and reduction as follows:

Oxidation: net gain of an electronegative group or net loss of an electropositive one.

Reduction: net loss of an electronegative group or net gain of an electropositive one.

Here “net” means “after allowance for difference in valency”, so that in

MgO + 2HCl → MgCl2 + H2O

there are no net gains or net losses (e.g., each magnesium atom gains two electronegative atoms, but loses one with twice the valency). On these definitions, assignments corresponding to those given for (1a) - (1d) may be made for (2a) - (2d). In addition, complementary assignments can be made for (1a) - (1c) and (2a) - (2c) corresponding to those for (1d) and (2d), e.g. for (1a)

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In (1d) and (2d), the oxygen and chlorine are unchanged: they gain an equivalent amount of magnesium to the hydrogen they lose.

In the case of (1a) and (2a), the similarity between the two reactions has an even deeper basis. Both MgO and MgCl2 are predominantly ionic compounds, so that (1a) and (2a) may be written:

2Mg + O2 → 2Mg2+O2– (1a')

Mg + Cl2 → Mg2+(Cl–)2 (2a')

In both processes, each magnesium atom loses two electrons, while each oxygen or chlorine atom gains electrons. This suggests the definitions:

Oxidation: loss of one or more electrons by an atom or radical.

Reduction: gain of one or more electrons by an atom or radical.

These definitions are of limited applicability, however, as can be seen from reactions (1b) and (2b). Neither H2O nor HCl is ionic, so that this new definition does not apply. Both molecules are, however, polarized, and if this polarization is imagined to be complete, the new definitions can then be used:

The dashed arrows here refer to the mental process of imagining the polarization being taken to the limit. The resulting charges are what we have previously called “polar numbers” or “oxidation numbers”. In terms of these, oxidation and reduction may be defined as follows:

Oxidation: an increase in the oxidation number of an atom or radical.

Reduction: a decrease in the oxidation number of an atom or radical.

These definitions are equivalent to those given earlier in terms of net gains or net losses of electronegative or electropositive species. They have the advantage of being easier to use, but suffer from being very artificial when applied to other than ionic compounds.

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Notes

(i)As now defined, oxidation and reduction always go together. Oxidation-reduction reactions are accordingly called “redox” reactions. The species that effects the oxidation is called the oxidizing agent, and is itself reduced in the reaction; the other species is called the reducing agent, and is itself oxidized.

(ii)On the new definitions, gain or loss of oxygen continues to be oxidation or reduction, except

(a)when the oxygen displaces, or is displaced by, another electronegative species, as in

MgO + 2HCl → MgCl2 + H2O

(b)when the oxygen is in combination with fluorine, where gain of oxygen corresponds to reduction of fluorine and loss to oxidation.

(iii)Organic chemists sometimes speak of loss of hydrogen as oxidation and gain of hydrogen as reduction. This is acceptable, as long as the hydrogen is electropositive relative to the species that is gaining or losing it, as is the case with organic compounds. In relation to metals like sodium or potassium, gain of hydrogen corresponds to oxidation.

(iv)The retention of the word “oxidation” in the broader definitions is unfortunate, especially when we have to speak of loss of oxygen as an oxidation, as we do in

2OF2 —heat→ O2 + 2F2

It might have been better to have used the words “elevation” and “reduction” to describe the broader categories, and “oxidation” and “deoxidation” to describe the original ones. As it is, whenever chemists wish to differentiate the latter from the former, they have to use the terms “oxygenation” and “deoxygenation”.

17.2 Half equations

When the compounds involved are ionic, the equation for a redox reaction can be divided into two half equations, one representing the reduction of the oxidizing agent and the other the oxidation of the reducing agent. Thus equation (2a) or (2a') can be divided into

Mg → Mg2+ (in MgCl2) + 2e–

Cl2 + 2e– → 2Cl– (in MgCl2)

Note that this can only be done when the compounds are ionic: it cannot be done in the case of an equation like (1b) or (2b).

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