Equilibrium — Part 2: Ionic equilibrium, pH, buffers and solubility product: visual lesson, CBSE Class 11 Chemistry

Lesson

Ionic Equilibrium

Part 2 — equilibria of acids, bases, salts and sparingly soluble compounds in water.

In this lesson you will

Distinguish Arrhenius, Brønsted–Lowry and Lewis acids and bases, and identify conjugate pairs
Calculate pH from Ka, Kb, degree of ionisation and Kw
Explain the common ion effect, salt hydrolysis and buffers (Henderson equation)
Use Ksp to find solubility and to predict precipitation

Table of four questions of Part 2 and the tool used to answer each
Question	Tool
How acidic is this solution?	pH and Ka
Why does a salt solution turn litmus?	Hydrolysis
How can a solution resist a pH change?	Buffer
Will a precipitate form on mixing?	Qsp against Ksp

Concept

Acids and bases: three definitions

Each definition widens the last: Arrhenius (in water), Brønsted–Lowry (proton transfer), Lewis (electron-pair transfer).

Table comparing the Arrhenius, Brønsted–Lowry and Lewis definitions of acids and bases with examples
Theory	Acid	Base	Example
Arrhenius	Gives $\ce{H+}$ in water	Gives $\ce{OH-}$ in water	$\ce{HCl}$ , $\ce{NaOH}$
Brønsted–Lowry	Proton donor	Proton acceptor	$\ce{NH3 + H2O <=> NH4+ + OH-}$
Lewis	Electron-pair acceptor	Electron-pair donor	$\ce{BF3}$ accepts a pair from $\ce{NH3}$

Concept

Conjugate acid–base pairs and amphoteric species

When an acid loses $\ce{H+}$ it forms its conjugate base; when a base gains $\ce{H+}$ it forms its conjugate acid. A strong acid has a very weak conjugate base. Species that can do both (water, $\ce{HCO3-}$ ) are amphoteric.

Table of acids with conjugate bases and bases with conjugate acids
Acid	Conjugate base	Base	Conjugate acid
$\ce{HCl}$	$\ce{Cl-}$	$\ce{NH3}$	$\ce{NH4+}$
$\ce{CH3COOH}$	$\ce{CH3COO-}$	$\ce{H2O}$	$\ce{H3O+}$
$\ce{H2O}$	$\ce{OH-}$	$\ce{HCO3-}$	$\ce{H2CO3}$
$\ce{H2PO4-}$	$\ce{HPO4^2-}$	$\ce{CO3^2-}$	$\ce{HCO3-}$

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