This document covers various topics related to chemical equilibrium including: 1. Irreversible and reversible reactions, and examples of each. 2. Types of equilibrium including homogeneous, heterogeneous, physical, and chemical equilibrium. 3. The law of mass action and how equilibrium constants are calculated. 4. How changing conditions like temperature, pressure, and concentration affects chemical equilibria. 5. Additional topics like acid-base theories, buffer solutions, and solubility products are also briefly discussed.

1. Ranjan Kumar Mohanty

2. 1. Irreversible and reversible Reaction
2. Types of Equilibrium
3. Other classifications of Equilibrium
4. Law of Mass Action
5. Equilibrium Constant
6. Reaction Quotient or Mass Action ratio
7. Le-Chartlier Principle
8. Autoprotolysis of Water
9. Logarithm Identities
10. Logarithm Table
11. PH and POH
12. Arrhenius theory of electrolysis
13. Lewis Theory of acid and base
14. Bronsted & Lowery Theory
15. Conductor and Electrolyte
16. Solubility Product
17. Common ion effect
18. Buffer Solution

3. The reaction in which reactant only form the product or the reaction which takes
place in one direction is called irreversible reaction.
BaSO4 + 2NaCl
White ppt.
Example : BaCl2 + Na2 SO4
The reaction which takes place in both the direction is called reversible reaction.
In this type of reaction, reactants combine to form product & again product
give back reactant.
Example : N2 + 3H2
2NH3
( aq )
( g ) ( g )
( aq ) ( s ) ( aq )
( g )

4. Equilibrium is state of reversible reaction when rate of forward reaction is
equal to the rate of backward reaction.
2NH3
Example : N2 + 3H2
( g ) ( g ) ( g )
 Homogeneous equilibrium
 Heterogeneous equilibrium

5. The equilibrium between species having same physical state is called
homogeneous equilibrium.
Example : N2 + 3H2 2NH3
( g ) ( g )
The equilibrium between species having different physical state is called
homogeneous equilibrium.
Example : CaCO3 CaO +CO2
( g ) ( s )
( g )
( g )

6. Equilibrium between the same compound & different phases is called physical
equilibrium.
Example : HO
HO
22Equilibrium between reactant & product is known as chemical equilibrium.
Example : N2 + 3H2 2NH3
The equilibrium between ions & unionized molecule is known as ionic
equilibrium.
Ag+ + Cl --
( l ) ( g )
( g ) ( g ) ( g )
Example : AgCl
( s ) ( aq ) ( aq )

7. Rate of reaction is directly proportional to the molar concentration of reactants
having power equal to their coefficients in the reactions
xA + yB + zC Product
Rate [A]x [B]y [C]z
Rate = k[A]x [B]y [C]z
ȣ
[k = Rate constant]

8. It is the ratio of molar concentration of product to the molar concentration of
reactant having power equal to their coefficient in the reaction.
2NH3
Example : N2 + 3H2
( g ) ( g ) ( g )
(mol/lit)2 1
Unit of K = (mol/lit)(mol/lit)3 = (mol/lit)2 = mol
-2
lit
2
[NH3]2
K = [N2][H2]3

9. The value of K obtain by putting the concentration of reactants & products of
any stage of reaction (other than equilibrium) into the equilibrium constant
expression is called mass action or Reaction Quotient. It may be Qc or Qp .
 If Q < K, it means forward reaction becomes faster i.e. more of the product
are formed.
 If Q > K, then the backward reaction become faster it means more
reactants are formed.
 If Q = K, it means reaction is in equilibrium.

10. If external stress is applied over equilibrium then the equilibrium shift in such
a direction, so that external stress become minimize.
On increasing the temperature, endothermic reaction become fast and exothermic
reaction become slow.
2NH3
Example : N2 + 3H2
( g ) ( g ) ( g )
On increasing the temperature forward reaction become slow and backward
reaction becomes fast.
OR
On increasing the temperature rate formation of ammonia decreases and rate of
dissociation of ammonia increases.

11. On increasing the pressure, such reaction becomes fast in which number of moles
decreases, while those reaction becomes slow in which number of moles increases.
2NH3
On increasing the pressure forward reaction become fast and backward reaction
becomes slow.
OR
Example : N2 + 3H2
1mol 3mol
4mol
2mol
On increasing the pressure rate formation of ammonia increases and rate of
dissociation of ammonia decreases.

12. On increasing the concentration of reactants, forward reaction becomes fast and
on increasing the concentration of product, backward reaction become fast.
Example : N2 + 3H2 2NH3
( g ) ( g ) ( g )
On increasing the concentration of N2 and H2 , forward reaction becomes fast
and on increasing the concentration of NH3 ,backward reaction becomes fast.

13. Catalyst is a substance which affect the rate of reaction but it will not effect the
equilibrium state.
Ques : State the conditions which help in increasing the formation of Sulpur -
trioxide gas in the following reaction?
2SO2 + O2
2SO3 + Heat
Ans : Formation of SO3 can be increased by :
1.Decreasing the temperature
2.Increasing the pressure
3.Increasing the concentration of SO2 and O2.

14. Autoprotolysis means self ionization of water. Application of autoprotolysis is
that water can react with acid as well as base.
H2O + HOH H3O+ + OH-
( l ) ( l ) ( aq ) ( aq )
Kw = [H3O+][OH-]
[H3O+][OH-] = 1 x 10-14
[H+ ][OH-] = 10-14 [H3O+ = H+ ]

15. 1. log A x B = log A+ log B
2. log A / B = log A - log B
3. log An = n log A
Example : log 28 x 10-4
= log 28 + log 10-4
= log 28 – 4log 10
= 1.4472 – (4 x 1)
= – 2.5548

17. It is reciprocal of logarithm of hydronium ion concentration.
1
PH = log[H3O+]3
PH= -log[H3O+]3
PH = 7 Neutral
PH < 7 Acidic
PH > 7 Alkaline
It is reciprocal of logarithm of hydroxide ion concentration.
POH = - log[OH- ]3
For water :
PH + POH = 14

18. Ques : Calculate H3O+, OH-, PH & POH of 0.0001 M HCl solution?
( aq ) ( aq )
0.0001M
Ans : HCl + H2
0.0001M
( g ) ( l )
0.0001M
[H3O+] =10– 4 M ( [H3O+]=[H+] )
PH + POH = 14
4 + POH = 14
POH = 14 – 4
POH = 10
[H+] [OH–] =10–14
[10- 4] [OH–] =10– 4
[OH–] =10–10 M
PH = - log [H+]
PH = - log 10– 4
PH = + 4 log 10
PH = + 4(log 10 = 1)
ACIDIC MEDIUM
H+ + Cl –

19. 1. When an electrolyte dissolve in water then it conduct electricity and split up
into its ions. This process is known as ionization.
2.Ions have charge over them. Positively charged ions is known as cations and
negatively charged ions is known as anions.
3.Total positive charge is equal to the negative charge carried by the ions.
4.The extinct of dissociation of an electrolyte is called as degree of dissociation
or degree of ionization.
Degree of dissociation = number of molecular ionise
total no. of molecule taken
5.Degree of ionization for strong electrolyte is 1 because strong electrolyte are
completely ionized but the degree of dissociation of weak electrolyte are
partially ionized.

20. If external stress is applied over equilibrium then the equilibrium
shift in such a direction, so that external stress become minimize.
Electron deficient species or the species which accept the electron pair
is known as Lewis acid.
Example : BF3 , Al2Cl3 , H+, H3O+
The species which can donate electron pair or electron rich species is
known as Lewis Base.
Example : NH3 , HOH , Cl –

21. The substance which donate proton (H+) in water is known as acid.
Example : H3O+, HCl
The substance which gain proton (H+) is known as base.
Example : NH3, Cl –
HCl + HOH HO + + Cl –
Acid 1 Base 2 Acid 32
Base 1
NH+ HO NH3 24
Base 1 Acid 2
+ + OH
Acid 1 Base 2

22. 2- and NH3
Ques : Write the conjugate acid of the CO3
2- HCO3
Ans : 1. CO3
-
+
2. NH3 NH4
- and NH4
Ques : Write the conjugate base of the HCO3
- CO3
Ans : 1. HCO3
2-
+ NH3
2. NH4

23. It is a substance which allows the electric current to pass through it
but itself can’t decompose into ions.
Example : Cu, Al, etc..
Electrolyte is a substance which allows the electric current to pass
through it in molten and aqueous state & itself decompose into ions.
Example : NaCl, dil. HCl
Types of Electrolyte :
1. Strong electrolyte
2.Weak electrolyte

24. An electrolyte which completely dissociate into its ions.
Examples :
Strong Acid
HCl H+ + Cl –
HSO2H + + SO2-
24 4
HNOStrong Base
H+ + NO- 3 3
NaOH Na+ + OH-KOH
K+ + OH-Strong
Salt
NaCl Na+ + Cl -
KNO3 K+ + NO3
-
( aq )
( aq )
( aq )
( aq )
( aq )
( aq )
( aq )

25. An electrolyte which completely dissociate into its ions.
EWxeaakmpAlecsid:
CHOOH CHCOO- + H+
33HCO2H+ + CO2-
23 3
HCN H+ + CN- Weak Base
NH4OH NH4
+ + OH-Mg
+ +2OH-Weak
(OH)2 Mg2
Salt
+ + CO3
(NH4) CO3 2NH4
2-
( aq )
( aq )
( aq )
( aq )
( aq )
( aq )

26. When a partially soluble salt is dissolved in water to make its
saturated solution then an equilibrium exist between undissolved
solid and its ion in the solution. In this case equilibrium constant is
known as solubility product (Ksp).
Condition for precipitation : Ionic product > Solubility product
Ist case : AB
( s )
( aq ) ( aq ) A+ + B –
s s
Ksp = [A+] [B-]
Ksp = s x s (s = solubility)
Ksp = s2 Example : AgCl

27. ( aq ) ( aq ) A2+ + B –
Ksp = [A2+] [B-]2
Ksp = s x (2s)2
Ksp = 4s3
Example : PbCl2
( aq ) ( aq ) A3+ + B –
Ksp = [A2+] [B-]2
Ksp = s x (3s)3
Ksp = 27s4
Example :Fe(OH)3
IInd case : A( s )B2
s 2s
Ques : Find the solubility product of AB3?
Ans : AB( s ) 3
s 3s

28. Dissociation of weak electrolyte minimize by adding strong
electrolyte which gives one of the common ion in the solution.
Example :
1.Weak Electrolyte : HS
2H+ + S2-
2Strong Electrolyte : HCl H+ + Cl -
Common
2.Weak Electrolyte : NHOH NH44
+ + OH-Strong
+ + Cl -
Electrolyte : NH4Cl NH4
Common

29. The solution whose PH remains unchanged on adding little amount
of the acid and base is known as buffer solution.
Example : Blood is also a buffer solution having PH 7.4
Acidic buffer solution
Basic buffer solution

30. PKa = - log Ka
(equilibrium constant of acid)
[salt]
PH= PKa +
[acid]
PKb = - log Kb
(equilibrium constant of base)
[salt]
POH= PKb +
[base]

31. It is prepared by mixing equimolar amount of weak acid and its
salt of strong base.
Example : CH3COOH and CH3COONa
It is prepared by mixing equimolar amount of weak base and its
salt of strong acid.
Example : NH4OH and NH4Cl

32. Ranjan Kumar Mohanty