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IIT JEE Exam Ionic Equilibrium notes

Fundamentals of Acids, Bases & Ionic Equilibrium

Acids & Bases

When dissolved in water, acids release H+ ions, base release OH– ions.

Arrhenius Theory

When dissolved in water, the substances which release
(i) H+ ions are called acids (ii) OH− ions are called bases

Bronsted & Lowry Concept

Acids are proton donors, bases are proton acceptors
Note that as per this definition, water is not necessarily the solvent.
When a substance is dissolved in water, it is said to react with water e.g.
HCl + H2O ® H3O+ + Cl− ; HCl donates H+ to water, hence acid.
NH3 + H2O ® NH4 + + OH− ; NH3 takes H+ from water, hence base.
For the backward reaction, NH4
+ donates H+, hence it is an acid; OH− accepts H+, hence it is base.
NH3 (base) & NH4+ (acid) from conjugate acid base pair.

Conjugate acid and bases
To get conjugate acid of a given species add H+ to it. e.g. conjugate acid of N2H4 is N2H5
+.
To get conjugate base of any species subtract H+ from it. e.g. Conjugate base of NH3 is NH2
−.
Note: Although Cl− is conjugate base of HCl, it is not a base as an independent species. In fact,
anions of all strong acid like Cl−, NO3
−, ClO4
− etc. are neutral anions. Same is true for cations
of strong bases like K+, Na+, Ba++ etc. When they are dissolved in water, they do not react
with water (i.e. they do not undergo hydrolysis) and these ions do not cause any change in
pH of water (others like CN− do).

Some examples of :

Basic Anions : CH3COO−, OH−, CN− (Conjugate bases of weak acids)
Acid Anions: HSO3
−, HS− etc. Note that these ions are amphoteric, i.e. they can
behave both as an acid and as a base. e.g. for H2PO4
− :
HS− + H2O S2− + H3O+ (functioning as an acid)
HS− + H2O H2S + OH− (functioning as a base)
Acid Cations : NH4
+, H3O+ etc.(Conjugate acids of weak bases)
Note : Acid anions are rare.

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