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.

Significance of ΔG & ΔG0
= 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

Cα2is constant.

Ka & Kbare 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.