Equilibrium

## Physical Equilibrium

An example of physical Equilibrium is when the number of water particles evaporating and the number of condensing particles become equal . Physical Equilibrium is the state in which two reverse process occur at the same rate .

## Chemical Equilibrium

Chemical Equilibrium is a process in which two reverse reactions occur at the same rate . Thus , chemical equilibrium is possible only in reversible reactions .

The reaction taking place from Reactants to Products is called Forward Reaction . The reaction taking place from Products to Reactants is called Reverse reaction .

At equilibrium , the rate of Forward Reaction and Reverse reaction becomes same and the concentrations attain constant values . Thus , equilibrium is dynamic in nature , although the concentrations become static .

## Law of Mass Action

The Law of Mass Action states that the rate of a Reaction is directly proportional to the products of their concentrations raised to the power of their respective stoichiometric coefficients .

## Equilibrium Constant

Equilibrium Constant = Rate Constant of Forward Reaction / Rate Constant of Reverse Reaction

Kp = KcdnRT

## Le Chatelier's Principle

The Le Chatelier's Principle states that when a stress (additional change) is applied to a reacting system in equilibrium , the system shifts itself in such a way so as to reduce the strain caused by the stress .

# Ionic Equilibrium

Ionization is a reversible process .Thus Ionic equilibrium is the equilibrium associated with ionization .

## Electrolytes

Electrolytes are those solutions through which electricity can be passed . Strong Electrolytes dissociate completely while weak electrolytes dissociate incompletely . Thus , the dissociation of weak electrolytes id=s a reversible process .

## Dissociation

Degree of Dissociation

α = No. of moles dissociated / Total No. of moles

For strong electrolytes , the α = 1 .

## Acids and Bases

**Arrenhius Theory**

Acids are H+ donating species . Bases are OH- donating species

**Lowry - Brownsted Theory**

Acids are proton donating species . Bases are proton accepting species .

**Lewis Theory**

Acids are electron accepting species . Bases are electron accepting species .

**Strength of Acid and Base :-**

Ka = [H+][A-] / [HA] = Cα 2 / 1 - α (Ostwald's Dilution Law)

Kb = [OH-][B+] / [BOH] = Cα2 / 1 - α

Cα2 is constant .

Ka & Kb are just equilibrium constants and hence depend only on temperature . The strengths of acids and bases also depends upon the solvent used . If the solvent can accept protons easily , the acid will donate more protons .

## Concept of pH

pH stands for Power of Hydrogen . It is the negative Logarithm of Concentration of H+ ions to the base 10 .

pH = - log10[H+] OR pKa = - log10Ka

Since it is negative , lower the pH , greater is the power of hydrogen ; stronger is the acid

pOH is the negative Logarithm of Concentration of OH- ions to the base 10 .

pOH = -log10[OH-] OR pKb = - log10Kb

pH + pOH = 14

**Relative Acid Strength**

1) Polarity of bond to which Hydrogen is attached to **:** The more polar this bond is , the more easily protons are given away .

2) Bond Strength **:** The less the bond strength(bond with hydrogen) , the stronger the acid .

i) HF < HCl < HBr < HI

ii) H2O < HF

iii) HIO < HBrO < HClO

iv) HClO < HClO2 < HClO3 < HClO4

v) HSO4- < H2SO4

## Ionic Product of Water

For dissociation of Water ,

K = [H+][OH-] / [H2O]

K [H2O]= [H+][OH-]

Kw = [H+][OH-]

## Buffer Solutions

Buffer solutions are the solutions whose pH does not change considerably with addition of small amounts of acid or base . Buffer solutions are generally prepared by mixing a weak acid(e.g. CH3COOH and CH3COONa) and it's salt or weak base and it's salt (e.g. NH4OH and NH4Cl). i.e. weak acid (CH3COOH) and it's conjugate base(CH3COO-) or weak base (NH4OH) and its conjugate acid(NH4+) .

**Henderson - Hasselbalch Equation**

**Acidic Buffer :**

Ka = [H+][A-] / [HA]

[H+] = Ka  [HA] / [A-] = Ka [acid] / [salt]

**pH = pKa - log10{[acid] / [salt]}**

 A- is the concentration of salt , since it is completely dissociated .

**Basic Buffer :**

Kb = [OH-][B+] / [BOH]

[OH-] = Kb [BOH] / [B+] = Kb [base] / [salt]

**pOH = pKb - log10{[base] / [salt]}**

## Conjugate Acids and Base

A strong acid gives a weak conjugate base and a strong base gives a weak conjugate acid .

## Solubility Product

For dissociation of AgCl ,

K = [Ag+][Cl-] / [AgCl]

[AgCl] = K' = constant

KK' = [Ag+][Cl-]

Ksp = [Ag+][Cl-]

1. If **I.P. > S.P**.  ; Precipitate is formed

2. If **I.P. = S.P.**  ; Equilibrium

3. If **I.P. < S.P.**  ; No Precipitate formed

Order of Solubility depends upon concentration of one of the Ions .

## Common Ion Effect

Common Ion Effect states that when a weak electrolyte is dissolved in strong electrolyte having an ion common to each other ; the dissociation of the weaker electrolyte is suppressed by the strong electrolyte .

## Tips and Tricks

1. While dealing with Ksp , Ionic Product , ka and kb  , Rely on the equations and it's stoichiometry .