IONIC EQUILIBRIA - Notes : Class 12th Chemistry Chapter 3 : Maharashtra State Board New Syllabus : Free Net Notes

Maharashtra State Board New Syllabus - Chemistry Chap 3 - Ionic Equilibria

Notes, Important Points, Definitions, Formulae..

Introduction

What is Ionic Equilibrium ? 

Answer - 

Definition : - The equilibrium between ions and unionized molecules in solution is called ionic equilibrium.

Example : - 

CH3COOH (aq)   ⇌  CH3COO- (aq)+ H⊕ (aq)

GENERAL FORM - AB ⇌ A+ + B– 

What is an Electrolyte & a Non-Electrolyte ? 

Answer -  

Definition Electrolyte : - The substances which give rise to ions when dissolved in water are electrolytes.

Definition Non-Electrolyte : - The non electrolytes are those which do not ionize and exist as molecules in aqueous solutions.

CLASSIFICATION OF ELECROLYTE BASED ON - Extent of Ionisation in Dilute Aqueous Solutions

Strong Electrolyte - 

Definition : - The electrolytes ionizing completely or almost completely are strong electrolytes. 

Example : - strong acids, strong bases and salts.

Weak Electrolyte - 

Definition : - The electrolytes which dissociate to a smaller extent in aqueous solution are weak electrolytes.

Note : An equilibrium is established between ions and unionized molecules (ionic equilibrium) when weak electrolyte is added in aqueous solutions.

Example : - Weak acids and weak bases.

Degree of Dissociation 

Definition : The degree of dissociation of an electrolyte is defined as a fraction of total number of moles of the electrolyte that dissociates into its ions when the equilibrium is attained.

Denoted by - ∝ (alpha)

ACIDS AND BASES

Arrhenius Theory of Acids and Bases : - 

Definitions According to Arrhenius : -

Acid : Acid is a substance which contains hydrogen and gives rise to H⊕

ions in aqueous

solution. 

 

Base : Base is a substance that contains OH group and produces hydroxide ions (OH ) ions in aqueous solution. 

Limitations of Arrhenius Theory :-

• It is applicable only to aqueous solutions.
• It does not account for basicity of NH3 and Na2CO3 which do not have OH group.

Bronsted-Lowry Theory ( Bronsted-Lowry Proton Transfer Theory )

Acid : Acid is a substance that donates a proton (H+) to another substance .

Base : Base is a substance that accepts a proton (H+) from another substance .

Conjugate Base : The base produced by accepting the proton from an acid is the conjugate base of that acid.

Conjugate Acid : The acid produced when a base accepts a proton is called the conjugate acid of that base.

Conjugate Acid-Base Pair : A pair of an acid and a base differing by a proton is said to be a conjugate acid-base pair.

Lewis Theory 

Acid :  Any species that accepts a share in an electron pair is called Lewis Acid.

Base : Any species that donates a share in an electron pair is called Lewis Base.

Amphoteric Nature of Water : Water has the ability to act as an acid as well as a base. Such behaviour is known as amphoteric nature of water.

Ionisation of Acids and Bases.

CLASSIFICATION OF ACIDS AND BASES ARE BASED ON THEIR EXTENT OF DISSOCIATION 

: -

Here are some examples of strong acids, strong bases , weak acids and weak bases.

1) Strong Acids :- EXAMPLE - HCl, HNO3, H2SO4, HBr AND HI

2) Strong Bases :- EXAMPLE - NaOH AND KOH.

3) Weak Acids : - EXAMPLE - HCOOH, HF, H2S

4) Weak Bases :- EXAMPLE - Fe(OH)3, Cu(OH) 

Note : 

1) Stong acids and Strong bases are almost completely dissociated in water. 

2) Weak Acids and weak bases are partially dissociated in water.

Dissociation costant for Weak Acids and Weak Bases

Defination : The dissociation constant of a weak acid or weak base is defined as the equilibrium constant for dissociation equilibrium of weak acid or weak base , respectively.

Ostwald's Dilution Law :-