The protonic definition of acids and bases given by Bronsted can be extended to the reactions occurring in non-aqueous solvents containing hydrogen such as NH3, N2H4,HF, H2SO4, CH3COOH, HCN and alcohols.
In an attempt to have a more general definition of acids and bases applicable to protonic and non-protonic solvents, several definitions have been proposed. One of these is due to Cady and Elsey (1928) according to whom an acid is solute that, either by direct dissociation or by reaction with the solvent gives the anion characteristic of the solvent and a base is a solute that, either by direct association or by reaction with the solvent, gives the cation characteristic of the solvent. If for example, we consider the solvent H2O, its characteristic cation and anion are H3O+ and OH– respectively as shown below:
Thus all those compounds which can give H3O+ ions in H2O will act as acids and all the compounds which can give OH– ions in H2O will behave as bases.
Similarly in N2O4 as solvent substance such as NOCl which yield NO+ ions are acids and the substances such as NaNO3 which yield NO3– ions are bases.
Evidently this definition of acids applies equally well to protonic and non-protonic solvents.
The auto-ionisation of some protonic and non-protonic solvents is shown below.
Just as with the Arrhenius definition, neutralisation is a reaction between an acid and a base to produce a salt and the solvent. Neutralisation reaction in some non-aqueous solvents are given below
|In liq. NH3||:||NH4Cl||+||NaNH2||⇔||NaCl||+||2NH3|
|In liq. SO2||:||SOCl2||+||[N(CH3)4]2 SO3||⇔||2[N(CH3)4]Cl||+||2SO2|
It may be seen from the following reactions that there is a complete analogy between the solvolytic acid and amphoteric behaviours in aqueous solvents.
|Solvolytic behaviour||In liq. NH3:||AlCl3 + NH3 ¾¾ → [Al(NH2)]2+ + H+ + 3Cl–|
|In H2O:||: AlCl3 + H2O ¾¾ → [Al(OH)]2+ + H+ + 3Cl–|
|Amphoteric behaviour||In liq. NH3:
|Zn(NH2)2 + 2NH2– ¾¾ → [Zn(NH2)4]2–
Zn(OH2) + 2OH– ¾¾ → [Zn(OH)4]2–
occurring in aqueous solvents (protonic and non-protonic both). It also includes many non-aqueous acid-base systems.
i) This theory does not consider a number of acid base reactions included in the protonic definition.
ii) It limits acid base phenomena to solvent systems only. Thus it does not explain the acid – base reactions which may occur in the absence of solvent.
iii) It cannot explain the neutralisation reactions occurring without the presence of ions. Thus this theory can simply be said to be an extension of the Arrhenius water – ion system.
Lewis Concept : The Electron – Donor – Acceptor System
This theory explains the acid-base phenomena not in terms of ionic reactions but in terms of electronic structure of the acid and base along with the formation of a coordinate covalent bond.
According to Lewis (1923), an acid is any species (molecule, radical or ion) that can accept an electron-pair to form a coordinate covalent bond and a base is any species that can donate an electron-pair to the formation of a coordinate covalent bond. Thus, in the Lewis system, an acid is an electron pair-acceptor and a base is an electron pair-donor.
Thus according to Lewis theory, the process of neutralisation is simply the formation of a coordinate bond between an acid and a base. The neutralisation product, termed as coordinate complex or adduct, may be either non-ionisable or may undergo dissociation or condensation reaction depending on its stability.
Now consider the reaction between a proton (H+) and : NH3 molecules as shown below.
Evidently in the above reaction proton (H+) accepts one electron pair from :NH3 molecule and is, therefore, acid, whereas :NH3 molecule which donates an electron pair, is a base. The adduct is NH4+ ion.
Lewis bases and Bronsted-Lowry bases are the same substances, since any molecule or ion which accepts protons does so because of the presence of an unshared pair of electrons. In the above example NH3 molecules is a proton acceptor (i.e., Bronsted-base) and an electron pair donor (i.e., Lewis – base).
Bronsted and Lewis theories are thus identical as far as bases are concerned except that the wording used for definition of the bases is different in both the theories. Thus NH3, H2O, OH–, Cl–, CN– etc. are the bases on the Bronsted as well as on the Lewis systems. There are however, few compounds such as amides, ethers, nitriles, C2H4,C2H2, C6H6 etc. which have little or no tendency to accept protons but react readily with Lewis – acids.
Classification of Lewis Acids
Any Lewis acid must contain at least one empty orbital in the valence shell of one of its atoms to accept an electron pair form a Lewis-base. Lewis – acids may be classified as:
(i) Molecules containing a central atom with an incomplete octet. Typical examples of this class of acids are electron deficient molecules such as alkyls and halides of Be, B and Al. Some reactions of this type of Lewis acid with Lewis bases are shown below:
Lewis Acid + Lewis base —> Adducts
(ii) Molecules containing a central atom with vacant d-orbitals. The central atom of the halides such as SiX4 ,GeX4 ,TiCl4 , SnX4 , PX3 ,PF3 ,SF4 ,SeF4 ,TeCl4 etc. have vacant d-orbitals. These substances can, therefore, accept an electron pair from the Lewis base to accommodate in their vacant d-orbital and can thus form adducts with a number of halide ions and organic bases.
These substances are, therefore, Lewis acids. These halides are vigorously hydrolysed by H2O to form an oxy acid or oxide of the central atom and the appropriate HX . The hydrolytic reactions take place presumably through the intermediate formation of unstable adducts with H2O. For example
(iii) Simple cations. Theoretically all simple cations are potential Lewis acids. Reactions of some cations as Lewis acids with Lewis Bases are shown below. It will be seen that these reactions are identical with those which produce Werner complexes.
Lewis acid + Lewis base —> Adduct or Addition compounds
Ammonation: Ag+ + 2(:NH3) —>[NH3 —> Ag —> NH3]+
The Lewis acid strength or coordinating ability of the simple cations which, according to Lewis, are Lewis acids, increases with
(a) an increase in the positive charge carried by the cation
(b) an increase in the nuclear charge for atoms in any period of the periodic table.
(c) a decrease in ionic radius.
(d) a decrease in the number of shielding electron shells.
Evidently the acid strength of simple cations increases for the element on moving from left to right in a period and from bottom to top in a group of periodic table. Thus:
Fe2+ < Fe3+ (positive charge increases, +2 —> +3)
K+ < Na+ (on moving from bottom to top in a group)
Li+ < Be2+ (on moving from left to right in a period)
——- strength of Lewis acids increasing ——–>
(iv) Molecules having multiple bond between atoms of dissimilar electro-negativity. Typical examples of molecules falling in this class of Lewis acids are CO2,SO2 and SO3.
In these compounds the oxygen atoms are more electronegative than S- or C- atom. As a result, the electron density of π-electrons is displaced away from carbon or sulphur atoms which are less electronegative than oxygen, towards the O-atom.
The C- or S- atom thus becomes electron deficient and is, therefore, able to accept an electron pair from a Lewis base such as OH– ions to from dative bond.
SO2 also reacts in the same manner with OH– ion
(v) Elements with an electron sextent. Oxygen and sulphur atoms contain six electrons in their valence shell and can, therefore, be regarded as Lewis acids. The oxidation of to ion by oxygen and to ion by sulphur are the acid-base reactions.
Utility of Lewis concept
(i) This concept also includes those reactions in which no protons are involved.
(ii) Lewis concept is more general than the Bronsted – Lowry concept (i.e. protonic concept) in that acid-base behaviour is not dependent on the presence of one particular element or on the presence or absence of a solvent.
(iii) It explains the long accepted basic properties of metallic oxides and acidic properties of non-metallic oxides
(iv) This theory also includes many reactions such as gas phase, high temperature and non-solvent reaction as neutralisation process.
(v) The Lewis approach is, however, of great value in case where the protonic concept is inapplicable, for example, in reaction between acidic and basic oxides in the fused state.
(i) Since the strength of Lewis acids and bases is found to depend on the type of reaction, it is not possible to arrange them in any order of their relative strength.
Thus, for example, experiments show that fluoride complex of Be2+ ions is more stable than that of Cu2+ ion, indicating that Be2+ ion is more acidic than Cu2+ ion. On the other hand amine complex of Cu2+ is more stable than that of Be2+ ion indicating that Cu2+ is more acidic than Be2+ ion.
(ii) According to the phenomenological criteria, an acid-base reaction should be a rapid reaction. There are, however, many Lewis acid-base reactions which are slow.