2. Introduction
Based on the direction of the reaction. The chemical reactions classified into two types.
1. Irreversible reactions. 2. Reversible reactions.
2. Irreversible reactions: If the reaction takes place only in one direction then the reaction is irreversible.
These reaction undergo 100% completion
Examples:
1.Neutralization reaction : HCl(aq) + NaOH(aq) → NaCl(aq) + H2O(l)
2. Precipitation reactions: AgNO3 (aq) + NaCl(aq) → AgCl(s) + NaNO3
3.Decomposition reaction: CaCO3(s) → CaO(s) + CO2(g)
(In open container only) (It is reversible in closed container)
4. Combustion reactions: C2H5OH(l) + 3 O2 (g)→ 2 CO2 (g)+ 3 H2O(g)
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3. 2. Reversible reactions: The chemical reaction which takes place in two opposite directions simultaneously under identical
conditions are said to be reversible reactions.
The reaction takes place from left to right is called forward reaction
The reaction takes place from right to left is called backward reaction
In reversible reactions the concentration of none of the reactants will become zero even at infinite time.
In reversible reaction a stage will be reached at which no further change in concentration of either reactants or products will be observed.
Examples:
N2 (g)+ 3H2 (g) 2 NH3(g)
H2 (g)+ I2 (g) 2HI(g)
2SO2 (g)+ O2 (g) 2SO3(g)
PCl5 (g) PCl3(g) + Cl2(g)
N2 (g)+ O2(g) 2NO(g)
CaCO3(s) CaO(s) + CO2(g) (It is reversible in closed container)
NH4HS(s) NH3 (g) + H2S(g) (It is reversible in closed container)
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4. Characteristics of Chemical Equilibrium
1.The equilibrium can be attained only if the system is closed.
2. The measurable properties of the system become constant at equilibrium and remain
unchanged with time so long as the external factors remain the same.
3. The equilibrium can be approached from either direction.
4. Catalyst has no net effect on equilibrium
5. At equilibrium DG = 0 .
6.At equilibrium , rate of forward reaction(rf)
rate of backward reaction(rb ) is equal.
rf = rb
5. Types of Equilibrium
1.Homogeneous equilibrium:
The reactions, in which all the reactants and the products are in the same phase, are called
homogeneous equilibrium reactions.
Examples:
N2 (g)+ 3H2 (g) 2 NH3(g)
H2 (g)+ I2 (g) 2HI(g)
2SO2 (g)+ O2 (g) 2SO3(g)
PCl5 (g) PCl3(g) + Cl2(g)
N2 (g)+ O2(g) 2NO(g)
2. Heterogeneous equilibrium:
The reactions, in which the reactants and the products are present in different phases, are called heterogeneous
equilibrium reactions.
Examples:
CaCO3(s) CaO(s) + CO2(g)
NH4HS(s) NH3 (g) + H2S(g)
H2O(l) H2O(g)
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6. Law of mass action
It is proposed by guldberg and waage
According to this rate of chemical reaction is directly proportional to product of
active masses of all reactants which are raised to the power equals to its
stoichiometric coefficients in the balanced equation.
Rate of reaction(r) = change in concentration of R(or)P / change in time
Rate of reaction(r) = DC/Dt
Ex: R P , r = - D[R]/Dt = D[P]/Dt
Active mass(a):
For pure solids, pure liquids a = 1(unity)
(because the active mass of pure solids and liquids depends on the
density and molecular mass. The density and molecular of a mass of
pure liquids and solids are constant)
In case of solutions , a = molar concentration (Molarity)
In case of gases , a = Partial pressure (or) molar concentration
(Molarity)
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7. Equilibrium Constant
Let us consider a hypothetical reaction
aA + bB cC + dD
according to the law of mass action,
Rate of the forward reaction, rf α [A]a [B]b
or Rate of the forward reaction, rf = kf [A]a [B]b
kf = rate constant of forward reaction
Similarly,
Rate of the backward reaction, rb α [C]c [D]d
or Rate of the backward reaction, rb = kb[C]c [D]d
kb = rate constant of backward reaction
At equilibrium, the two rates become equal, i.e.,
Rate of the forward reaction(rf) = Rate of the backward reaction(rb)
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8. Equilibrium Constant
So that, at equilibrium
kf [A]a [B]b = kb [C]c [D]d
But, kf / kb = Equilibrium constant, Kc
So,
The above equation is the law of chemical equilibrium.
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9. Relationshipbetween Kc and Kp:
aA(g) + bB(g) cC(g) + dD(g)
Let us assume A,B,C,D are ideal gases.
From ideal gas equation,
PV = nRT , P = (n/V)RT molar concentration = n/V.
PA = [A]RT , PB = [B]RT ,PC = [C]RT, PD = [D]RT
So that,
∆n = [(Number of moles of gaseous products)-(Number of moles of
gaseous reactants)]
Kp = Kc (RT)∆n
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10. Units of Equilibrium Constant
H2 (g) + I2 (g) 2HI(g)
Equilibrium constant is given by
N2 (g)+ 3H2 (g) 2 NH3(g)
Equilibrium constant is given by,
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15. Factors affecting equilibrium
Lechateliers principle: When a system at equilibrium undergoes stress the equilibrium shifts in such a direction in
order to undo the stress
There are 3 type of stress:
1) Changing concentration:
R P
At equilibrium,
If concentration of reactants increases()
equilibrium shifts towards products side(Right)
So that, [Reactants] decreases , [Products] increases
Similarly ,
If concentration of products increases ()
equilibrium shifts towards reactants side(left)
So that, [Reactants] increases , [Products] decreases
2) Changing pressure.
3) Changing Temperature.
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16. Factors affecting equilibrium
The state of equilibrium of a reversible chemical reaction is
mainly influenced by the following factors:
1. Concentration
2. Pressure
3. Temperature
4. Catalyst
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17. Effect of Concentration
A + B C + D
Forward reaction: A + B C + D ( Left to right)
Backward reaction: C + D A + B (Right to left)
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Component Concentration Direction in which equilibrium
shifts
Effect
Reactants
Increases Equilibrium shifts towards
right(products side)
Forward reaction
favorable
Decreases Equilibrium shifts towards left
(reactants side)
Backward reaction
favorable
Products
Increases Equilibrium shifts towards left
(reactants side)
Backward reaction
favorable
Decreases Equilibrium shifts towards
right(products side)
Forward reaction
favorable
18. Effect of Pressure
It is applicable for reversible reactions having at least one gaseous component
Dn = (No. of mole of gaseous products - No. of moles of gaseous reactant)
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Example Pressure Direction in which
equilibrium shifts
Effect
If Dn = 0 H2 (g)+ I2 (g) 2HI(g)
Dn = 2 – (1+1) = 0
Increase/
Decrease
No effect No effect
If
Dn = +Ve PCl5 (g) PCl3(g) + Cl2(g)
Dn = (1+1) - 1 = 1
Increases Equilibrium shifts
towards left
(reactants side)
Backward
reaction
favorable
Decreases Equilibrium shifts
towards right
(products side)
Forward reaction
favorable
19. Effect of Pressure
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Example Pressure Direction in which
equilibrium shifts
Effect
If
Dn = -Ve N2 (g) + 3H2 (g) 2 NH3(g)
Dn = 2 – (1+3) = -2
Increases Equilibrium shifts
towards right
(products side)
forward
reaction
favorable
Decreases Equilibrium shifts
towards left
(reactants side)
Backward
reaction
favorable
Note: In case of Dn (+Ve/-Ve) , If pressure increases equilibrium shifts
towards less No .of moles side.
If pressure decreases equilibrium shifts towards more No . of moles side.
20. Effect of temperature
Enthalpy
Change (DH)
Example Temperatur
e
Direction in which
equilibrium shifts
Effect
DH = - Ve
forward
reaction is
exothermic
N2 (g) + 3H2 (g)
2 NH3(g)
DH = - 93. kj
Increases Equilibrium shifts towards
left (reactants side)
Backward reaction favorable
Decreases Equilibrium shifts towards
right (products side)
forward reaction favorable
DH = + Ve
forward
reaction is
endothermic
N2 (g)+ O2(g)
2NO(g)
DH = +45 k.cal
Increases Equilibrium shifts towards
right (products side)
forward reaction favorable
Decreases Equilibrium shifts towards
left (reactants side)
Backward reaction favorable
DH = 0 No effect Increase/
Decreases
No effect No effect
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21. Effect of Catalyst
• The effect of catalyst is same for both forward
and backward reaction. Hence, catalyst has no
net effect on equilibrium state.
• No effect on equilibrium
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22. Application of Le Chatelier’s principle
Physical equilibrium: Liquid - vapor system
Ex: water water vapor
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Factor Change Direction of equilibrium Effect
Pressure Increases Equilibrium shifts towards left
(water)
Condensation of
vapors into water
Decreases Equilibrium shifts towards
Right (water vapor)
Vaporisation of
water into vapor
Temperature Increases Equilibrium shifts towards
Right (water vapor)
Vaporisation of
water into vapor
Decreases Equilibrium shifts towards left
(water)
Condensation of
vapors into water
23. Application of Le Chatelier’s principle
Ex: Ice water
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Factor Change Direction of equilibrium Effect
Pressure Increases Towards decrease of volume,
Equilibrium shifts towards Right (water)
More ice melts
Decreases Towards increase of volume,
Equilibrium shifts towards left (ice )
Water converts
into ice
Tempera
ture
Increases Equilibrium shifts towards Right (water) More ice melts
Decreases Equilibrium shifts towards left (ice) Water converts
into ice
24. Manufacture of ammonia (Haber's Process):
Ammonia can be synthesized from nitrogen and hydrogen in accordance
with the reaction.
N2 (g) + 3H2 (g) 2 NH3(g) ; DH = - 93. kj
The characteristics of the reaction are:-
1. The most favorable pressure range for the production of ammonia is found to be 200 -
900 atm.
2. Since at low temperature, the reactions tend to be slow due to kinetic effects, an
optimum temperature (450 0C) should give the most favorable results.
3.Iron used as Catalyst , molybdenum(Mo) (or) (K2O + Al2O3) used as promoters
4.Increase in the concentration of N2,H2 greater would be the formation of NH3. moreover
NH3 formed should be removed continuously by liquefaction so that, equilibrium
shifts in forward direction.
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Pressure(
atm)
1 100 300 1000
NH3(%) 15.3 81.5 90.0 99.3
Temperat
ure(C)
200 400 600 700
NH3(%) 90.0 47.7 13.8 7.3
25. Manufacture of sulphur trioxide (Contact Process):
During the manufacture of sulphuric acid (Contact process), sulphur trioxide is prepared by the
oxidation of sulphur dioxide(SO3)
SO2(g) + O2(g) SO3(g) ; DH = - 190 kj
The characteristics of the reaction are:-
1. This reaction proceeds with a decrease in volume. Therefore, high pressure will favor the forward
reaction. Usually a pressure of 1.5-1.7 atm serves the purpose.
2. The formation of SO3 being exothermic is favored by low temperatures accordance with
Lechateliers principle . Usually an optimum temperature of 673-723 K is used.
3. finely divided platinum(or)vanadium pent oxide(V2O5) is used as catalyst to attain the equilibrium
state rapidly.
4. Higher the concentration of SO2,O2 greater would be the yield of SO3.
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