Ionic equilibrium acids and bases theory

Ionic equilibrium acids and bases  7.12 Acids and bases 7.12.1 Acids and bases according to arrhenius theory Definition of a...

Ionic equilibrium acids and bases 

7.12 Acids and bases

7.12.1 Acids and bases according to arrhenius theory

Definition of arrhenius acid: Hydrogen containing compounds which ionise in aqua solution to produce H+ ions as the only cation are called acids.

Example The hydrogen containing compounds such as HCl, HNO₃ , H₂SO₄, CH₃COOH etc belong to the class of acids because the compound get organised in aqua solution to form H+ ions.

In fact, bare H⁺ ion (or proton) has no separate entity or existence in aqueous solution. Due to very small size of H⁺ ion, its charge density is very high. Hence in aqueous solution, due to strong attraction between H⁺ ion and negative end (i.e., oxygen atom) of H₂0 molecule, H+ ion gets stabilised by hydration and exists as hydronlum or hydroxonlum ion (H₃O⁺ ).

H⁺ + H₂O= H₃O⁺  

Acids get ionised in aqueous solution to form H₃O⁺  (aq) ion.

Definition of arrhenious base : A base is a substance which on ionisation in aqueous solution produces hydroxyl ions (OH-) as the only anion.

Limitations of Arrhenius theory

1. According to Arrhenius theory, presence of water is compulsory for any compound to exhibit acidic or basic properties. It cannot be explained with the help of this theory that these are inherent characteristics of a compound.
2. Acidic or basic property of a substance insoluble in water cannot be explained by Arrhenius theory.
3. Acidity or basicity of any compound in non-aqueous solvents cannot be explained by Arrhenius theory. For example: acidity of NH4CI or basicity of NaNH2 in liquid ammonia cannot be explained by this theory.
4. According to Arrhenius theory, only the compounds containing hydroxyl lons(OH-) are considered as bases. Consequently, basicity of ammonia (NH3), methylamine (CH3NH2), aniline (C6H5NH2) ete, cannot be explained by this theory.

7.12.1 Acids and bases according to Bronsted Lowry Concept (Protonic Theory

)

Definition of acid : An acid is a substance which has a tendency to donate proton (or H+ ion).
Definition of base : A base is a substance which has a tendency to accept proton (or H+ ion).
According to this theory, it can be stated that an acid is a proton donor and a base is a proton acceptor.
Examples

H2O and [H2PO4]- acts as acid as well as base

Concept of conjugate acid-base pair

Definition:  A pair of species having a difference of one protons between themselves is called a conjugate acid-base pair.
Acids donates proton to produced conjugate base and base accept proton to produced conjugate acid.

Strength of conjugate acid and conjugate base in aqueous solution

Higher the tendency of an acid to donate a proton, greater will be its strength. Similarly, the basicity of a base increases with its proton accepting capacity.

[1] The acids HCI, HNO3, H2SO4 etc. undergo complete ionisation in aqueous solution to form H3O+ ions and the corresponding conjugate bases. Hence, these considered as strong acids in aqueous solution. In aqueous solution, the conjugate base produced from a strong acid, has less tendency than H2O to accept a proton. Therefore in aqueous solution, the conjugate base of a strong acid is found to be very weak in nature.

[2] The acids HF, HCN, CH3COOH, HCOOH undergo slight ionisation in aqueous solution to produce HO- ions and the corresponding conjugate bases. Their ability to donate protons is very low due to partial ionisation in aqueous solution and hence are called weak acids. In aqueous solution, the conjugate base produced from a weak acid has more tendency than H20 to accept a proton. Therefore in aqueous solution, the conjugate base of a weak acid is found to be very strong in nature.

[3] In aqueous solution, strong bases like [NH2]- , 02-, H¯ etc. react completely with water to form their corresponding conjugate acids and OH ions. The strength of these conjugate acids is much less than that of H20. Hence, in aqueous solution, the conjugate acid of a strong base is very weak in nature. On the other hand, in aqueous solution, weak bases like NH3, CH3NH2 etc. react partially with water to produce corresponding conjugate acids and OH- ions. The strength of these conjugate acids is greater than H20. Therefore, in aqueous solution, a weak base always produces a strong conjugate acid.

With increasing basicity of a base, the strength L.e, the proton donating capacity of its conjugate acid decreases. On the contrary, with decreasing basicity, the strength of the corresponding conjugate acid of the base increases.

Limitations of Bronsted-Lowry concept

1. With the help of this theory, the reaction between the acid and the base is explained in terms of gain or loss of proton(s). But, there are many acid-base reactions where exchange of proton(s) does not take place. Such types of acid-base reactions cannot be explained by this theory.
2. Acidic properties of various non-metallic oxides (e.g.,C2O, SO2 etc.) as well as basic properties of various metallic oxides (e.g., CaO, Ba0 etc.) cannot be explained with the help of Bronsted-Lowry concept. Apart from this, the acidic properties of BF,, AlCl,, SnCl, etc. cannot be explained with the help of this theory.

7.12.3 Lewis concept of acids and bases

Definition of acid: An acid is a substance which can accept a pair of electrons.
Definition of base: An acid is a substance which can donate a pair of electrons.

• Lewis acid is an acceptor of lone pair of electrons and form a co-ordinate bond with Lewis base.
• Lewis base ls a donor of lone pair of electrons and form a co-ordinate bond with Lewis acid.