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Whenever we hear the word Equilibrium immediately a
picture arises in our mind an object under the influence of two
opposing forces. For chemical reactions also this is true. A
reaction also can exist in a state of equilibrium balancing
forward and backward reactions.
Reversible and Irreversible reactions
A chemical reaction is said to have taken place when the
concentration of reactants decreases, and the concentration of
the products increases with time. The chemical reactions are
classified on the basis of the extent to which they proceed, into
the following two classes;
(1) Reversible reactions : Reaction in which entire
amount of the reactants is not converted into products is termed
as reversible reaction.
(i) Characteristics of reversible reactions
(a) These reactions can be started from either side ,
(b) These reactions are never complete ,
(c) These reactions have a tendency to attain a state of
equilibrium, in which Free energy change is zero (G = 0),
(d) This sign (⇌) represents the reversibility of the
reaction,
(ii) Examples of reversible reactions
(a) Neutralisation between an acid and a base either of
which or both are weak e.g.,
CH COOH NaOH 3
⇌ CH COONa HO 3 2
(b) Salt hydrolysis, e.g.,
Fe Cl 3 3 H 2 O ⇌ Fe OH 3 HCl 3
(c) Thermal decomposition, e.g.,
PCl 5 ( g ) ⇌ PCl (^) (^3) ( (^) g ) Cl (^2) ( g ) Q
(d) Esterification, e.g.,
CH 3 COOH C 2 H 5 OH^ ⇌ CH 3 COOC 2 H 5 H 2 O
(e) Evaporation of water in a closed vessel, e.g.,
H 2 O ( l ) ⇌ H 2 O ( g ) Q
(2) Irreversible reactions : Reaction in which entire
amount of the reactants is converted into products is termed as
irreversible reaction.
(i) Characteristics of irreversible reactions
(a) These reactions proceed only in one direction (forward
direction),
(b) These reactions can proceed to completion,
(c) In an irreversible reaction, G < 0,
(d) The arrow () is placed between reactants and
products,
(ii) Examples of irreversible reactions
(a) Neutralisation between strong acid and strong base
e.g.,
NaOH HCl NaCl H 2 O 13. 7 kcal
(b) Double decomposition reactions or precipitation
reactions e.g.,
BaCl (^) (^2) ( (^) aq ) H 2 SO (^4) ( aq ) BaSO (^4) ( s ) 2 HCl ( aq )
(c) Thermal decomposition, e.g.,
() 2
, (^2 3) () 2 3
2 KClO KCls O
MnO s
(d) Redox reactions, e.g.,
SnCl (^) 2 ( aq ) 2 FeCl 3 ( aq ) SnCl 4 ( aq ) 2 FeCl 2 ( aq )
Chemical Equilibrium
Chapter
Equilibrium and Its dynamic nature
“ Equilibrium is the state at which the concentration of
reactants and products do not change with time. i.e.
concentrations of reactants and products become constant.”
The important characteristics of equilibrium state are,
(1) Equilibrium state can be recognised by the constancy
of certain measurable properties such as pressure, density,
colour, concentration etc. by changing these conditions of the
system, we can control the extent to which a reaction proceeds.
(2) Equilibrium state can only be achieved in close vessel.
(3) Equilibrium state is reversible in nature.
(4) Equilibrium state is also dynamic in nature.
(5) At equilibrium state,
Rate of forward reaction = Rate of backward reaction
(6) At equilibrium state, G = 0, so that H = T S.
Law of mass action and Equilibrium constant
On the basis of observations of many equilibrium
reactions, two Norwegian chemists Goldberg and Waage
suggested (1864) a quantitative relationship between the rates
of reactions and the concentration of the reacting substances.
This relationship is known as law of mass action. It states that
“ The rate of a chemical reaction is directly proportional
to the product of the molar concentrations of the reactants at a
constant temperature at any given time. ”
The molar concentration i.e. number of moles per litre is also
called active mass. It is expressed by enclosing the symbols of
formulae of the substance in square brackets. For example, molar
concentration of A is expressed as [ A ].
Consider a simple reversible reaction
aA bB ⇌ cC dD (At a certain temperature)
According to law of mass action
Rate of forward reaction
a b f
a b [ A ] [ B ] k [ A ][ B ]
Rate of backward reaction
c d b
c d [ C ][ D ] k [ C ][ D ]
At equilibrium ,
Rate of forward reaction = Rate of backward reaction
c d b
a b k (^) f [ A ][ B ] k [ C ][ D ]
a b
c d
c b
f
A B
C D
K
k
k
[ ][ ]
[ ][ ]
Where, Kc is called equilibrium constant.
In terms of partial pressures, equilibrium constant is
denoted by p
K and
b B
a A
d D
c C p P P
P P
K
In terms of mole fraction, equilibrium constant is
denoted by x
K and
b B
a A
d D
c C x X X
X X
K
Relation between Kp , Kc and Kx
n Kp KcRT
( )
n Kp KxP
()
n = number of moles of gaseous products – number of
moles of gaseous reactants in chemical equation.
As a general rule, the concentration of pure solids and
pure liquids are not included when writing an equilibrium
equation.
Value
of n
Relation
between Kp and
Kc
Units of Kp Units of Kc
0 Kp = Kc No unit No unit
0 Kp > Kc ( atm )n^ ( mole l –^1 )n
<0 Kp < Kc ( atm )n^ ( mole l –^1 )n
Characteristics of equilibrium constant
(1) The value of equilibrium constant is independent of
the original concentration of reactants.
(2) The equilibrium constant has a definite value for
every reaction at a particular temperature. However, it varies
with change in temperature.
(3) For a reversible reaction, the equilibrium constant for
the forward reaction is inverse of the equilibrium constant for
the backward reaction.
In general,
backwardreaction
forwardreaction
K
K
(4) The value of an equilibrium constant tells the extent
to which a reaction proceeds in the forward or reverse
direction.
(5) The equilibrium constant is independent of the
presence of catalyst.
(6) The value of equilibrium constant changes with the
change of temperature. Thermodynamically, it can be shown
that if K 1 and K 2 be the equilibrium constants of a reaction at
absolute temperatures T 1 and T 2
. If H is the heat of reaction
at constant volume, then
2 1
2 1
1 1
- 303
log log R T T
H K K (Van’t Hoff equation)
Concentration
Reactants
Time
Products
Equilibrium state
Rate of reaction
Equilibrium state
Time
G= 0
Table : 8.1 Homogeneous equilibria and equations for equilibrium constant ( Equilibrium pressure is P atm in a V
L flask )
n 0 ; Kp Kc n 0 ; Kp Kc p c n 0 ; K K
()
2 g
H +
()
2 g
I ⇌
()
2 g
HI
( )
2 ( )
g g
N H ⇌
( )
g
NH
( )
2 ( )
g g
SO O ⇌
( )
g
SO
()
5 g
PCl ⇌
( )
2 ( )
3 g g
PCl Cl
Initial mole 1 1 0 1 3 0 2 1 0 1 0 0
Mole at
Equilibrium
(1– x ) (1– x ) 2 x (1– x ) (3– 3 x ) 2 x (2– 2 x ) (1– x ) 2 x (1– x ) x x
Total mole at
equilibrium
2 (4 – 2 x ) (3 – x ) (1 + x )
Active
masses
V
1 x
V
1 x
V
2 x
V
1 x
V
1 x 3
V
2 x
V
2 2 x
V
1 x
V
2 x
V
1 x
V
x
V
x
Mole fraction
2
1 x
2
1 x
2
2 x
x
x
22
1
x
x
2
1
2
3
( 2 x )
x
x
x
3
2 2
x
x
x
x
3
2
3
1
x
x
1
1
x
x
1
x
x
1
Partial
pressure
2
1 x p
2
1 x p
2
2 x p 2 ( 2 )( 2 )
3 ( 1 )
2 ( 2 )_
1
x
Px
x
x P x
x P
x
x P 3
2 2
x
x P 3
1
x
x P 3
2
x
x P 1
1
x
x P 1
x
x P 1
Kc
2
2
x
x
4
2 2
x
xV
^
3
2
1 x
xV
x V
x
2
Kp
2
2
x
x
4 2
2 2
x P
x x
3
2
P x
x x
2
2
1 x
Px
Table : 8.2 Heterogeneous equilibria and equation for equilibrium constant (Equilibrium pressure is P atm)
NH 4 HS ( s ) ⇌ NH 3 ( g )+ H 2 S ( g ) C ( s ) CO 2 ( g )⇌ 2 CO ( g ) NH 2 CO 2 NH 4 ( s )⇌ 2 NH (^) 3 ( g ) CO 2 ( g )
Initial mole 1 0 0 1 1 0 1 0 0
Mole at equilibrium (1– x ) x x (1– x ) (1– x ) 2 x (1– x ) 2 x x
Total moles at equilibrium
(solid not included)
2 x (1+ x ) 3 x
Mole fraction
2
1
2
x
x
2
1
x
x
1
1
x
x
1
2
3
2
3
1
Partial pressure
2
P
2
P
x
x P 1
1
x
x P 1
2
3
2 P
3
P
p
K
2 P
2
2
x
Px
3 P
Relationship between equilibrium constant and
G °
G for a reaction under any condition is related with G °
by the relation, G G 2. 303 RT log Q
Standard free energy change of a reaction and its
equilibrium constant are related to each other at temperature T
by the relation, G RT K
o 2. 303 log
For a general reaction aA bB ⇌ cC dD
b B
a A
d D
c C
a a
a a K ( )( )
Where a represent the activity of the reactants and
products. It is unit less.
For pure solids and liquids: a 1.
For gases: a pressure of gas in atm.
For components in solution: a molar concentration.
Le-Chatelier's principle
Le-Chatelier and Braun (1884), French chemists,
made certain generalizations to explain the effect of changes in
concentration, temperature or pressure on the state of system in
equilibrium. When a system is subjected to a change in one of
these factors, the equilibrium gets disturbed and the system
readjusts itself until it returns to equilibrium. The generalization
is known as Le-Chatelier's principle****. It may be stated as :
“Change in any of the factors that determine the
equilibrium conditions of a system will shift the equilibrium in
such a manner to reduce or to counteract the effect of the
change.”
The principle is very helpful in predicting qualitatively
the effect of change in concentration, pressure or temperature
on a system in equilibrium.
Table : 8.3 The effect of varying conditions on the equilibrium a A + b B ⇌ c C + d D, n = ( c + d ) – ( a + b )
Change imposed on the
system in equilibrium
Equilibrium position moves Equilibrium constant Any other points
Conc. of A and/or B increased To right No change No change
Conc. of C and /or D
increased
To left No change No change
Pressure increased (^) To right if ( c d )( a b ), i.e. n ve
To left if ( c d )( a b ), i.e. n ve
No change if ( c d )( a b ), i.e. n 0
No change
No change
No change
Very little effect, if any, on
reactions in liquid solution.
Temperature increased (^) To left if (^) H ve (exothermic)
To right if H ve (endothermic)
Value decreased
Value increased
Equilibrium achieved faster
Addition of catalyst No change No change Equilibrium achieved faster
Application of Le-Chatelier's principle
The Le-Chateliers principle has a great significance for the
chemical, physical systems and in every day life in a state of
equilibrium.
(1) Applications to the chemical equilibrium
(i) Synthesis of ammonia (Haber’s process)
vol vol
N H
3
2 1
2 ^3 ⇌ NH kcal vol
2
3 (exothermic)
(a) High pressure( n 0 )
(b) Low temperature
(c) Excess of N 2 and H 2
(d) Removal of NH 3 favours forward reaction.
(ii) Formation of sulphur trioxide
vol vol
SO O
1
2 2
2 2 ⇌ SO kcal vol
2
3 (exothermic)
(a) High pressure( n 0 )
(b) Low temperature
(c) Excess of SO 2 and O 2 , favours the reaction in
forward direction.
(iii) Synthesis of nitric oxide
vol vol
N O
1
2 1
2 ^ ⇌ NO kcal vol
2 43. 2
2
(endothermic )
(a) High temperature
(b) Excess of N 2 and O 2
(c) Since reaction takes place without change in volume
i.e. , n 0 , pressure has no effect on equilibrium.
(iv) Formation of nitrogen dioxide
vol vol
NO O
1
2 2
2 ⇌ NO Kcal
vol
2
(a) High pressure
(b) Low temperature
(c) Excess of NO and O 2 favours the reaction in forward
direction.
(v) Dissociation of phosphours pentachloride
vol
PCl
1
5 ⇌ PCl^ Cl kcal vol vol
1
2 1
(a) Low pressure or high volume of the container, n 0
(b) High temperature (c) Excess of 5
PCl.
(2) Applications to the physical equilibrium
(i) Melting of ice (Ice – water system)
(GreaterVolume)
Ice ⇌ xkcal (Lesser Volume)
Water
(In this reaction volume is decreased from 1.09 c.c. to
1.01 c.c. per gm .)
(a) At high temperature more water is formed as it
absorbs heat.
(b) At high pressure more water is formed as it is
accompanied by decrease in volume.
(c) At higher pressure, melting point of ice is lowered,
while boiling point of water is increased.
(ii) Melting of sulphur : S ( s )⇌ S (^) ( l ) xkcal
(This reaction accompanies increase in volume.)
(a) At high temperature, more liquid sulphur is formed.
(b) At higher pressure, less sulphur will melt as melting
increases volume.
(c) At higher pressure, melting point of sulphur is
increased.
(iii) Boiling of water (water- water vapour system)
(Lowvolume)
Water ⇌ xkcal
(Higher volume)
Water Vapours
(It is accompanied by absorption of heat and increase in
volume.)
(a) At high temperature more vapours are formed.
(b) At higher pressure, vapours will be converted to liquid
as it decreases volume.
(c) At higher pressure, boiling point of water is increased
(principle of pressure cooker).
(iv) Solubility of salts : If solubility of a salt is
accompanied by absorption of heat, its solubility increases with
rise in temperature; e.g., NH 4 Cl , K 2 SO 4 , KNO 3 etc.
(a) H (^) 2 I 2
⇌ HI
(b) AgNO NaCl 3
AgCl NaNO 3
(c) CaCO 3 ⇌ CaO CO 2
(d) O 2 (^) 2 SO 2 ⇌ 2 SO 3
