Acids, Bases And Salts - Theories

Limitations of Arrhenius theory:

A)The definition of acids and bases are limited to aqueous solution only.

B)Acidic nature of substances which do not contain H⁺ and basic nature of substances which do not contain OH^- were not explained. 
 Example: CO₂, SO₂(acidic oxides), NH₃ ,Na₂O(Bases)

C)There are some substance which do not contain H⁺ and OH^- but can neutralise acids and a base.
            \( CO_2 + 2NaOH\xrightarrow{{}}Na_2 CO_3 + H_2 O \)    
            \( NH_3 + HCl\xrightarrow{{}}NH_4 Cl \) 
          \( CaO + CO_2 \xrightarrow{{}}CaCO_3 \)     
 
D)Studies on solutions showed that H⁺ ions do not have independent existence and exist as H₃O⁺. This is in contradiction to Arrhenius theory.
    \( (H^ + + H_2 O\xrightarrow{{}}H_3 O^ + ) \)

Modified version of Arrhenius theory:
It rectifies most of the above limitations
(i)  Water is weak electrolyte and ionises to a very weak extent.
        H₂O      \( \rightleftharpoons \)          H⁺ + OH^– 
        H⁺  +  H₂O   \( \rightleftharpoons \)   H₃ O+ 
       
       H₂O  +  H₂O  \( \rightleftharpoons \)  H₃ O⁺ + OH^– 

Above reaction is called Autoionisation or selfionisation of water.
(ii)  Water is neutral in nature i.e.[H₃O⁺] = [OH^–]
(iii) The substances which increase the H₃O⁺ ion concentration act as acids and while those which increase OH^– ion concentration act as bases.
Ex.A)    SO₂  + H₂O \( \to \) H₂SO₃  \( \rightleftharpoons \) H₃O⁺ + HSO₃^–
                Acid 
 B)    NH₃  + H₂O \( \to \) NH₄OH  \( \rightleftharpoons \) NH₄⁺ + OH^–
            Base
In order to account for these drawbacks, an advanced theory was introduced the protonic concept of acids and bases or Lowry - Bronsted theory of acids and bases.