Chemical equilibrium manik

Reversible Reactions
Not all chemical reactions proceed to completion. In most reactions two or more
substances react to form products which themselves react to give back the original
substances. Thus A and B may react to form C and D which react together to reform A and B.
A + B → C + D (Forward reaction)
A + B ← C + D (Reverse reaction)
A reaction which can go in the forward and backward direction simultaneously is
called a reversible reaction.
A + B ⇌ C + D
Such a reaction is represented by writing a pair of arrows between the reactants and
products. The arrow pointing right indicates the forward reaction, while that pointing left shows
the reverse reaction.
A few common examples of reversible reactions are listed below:
2NO2 (g) ⇌ N2O4 (g)
H2 (g) + I2 (g) ⇌ 2HI (g)
PCl5 (s) ⇌ PCl3 (s) + Cl2 (g)
CaCO3 (s) ⇌ CaO (s) + CO2 (g)

Chemical Equilibrium
Let us consider the reaction,
A + B ⇌ C + D
If we start with A and B in a closed vessel, the forward reaction
proceeds to form C and D. The concentrations of A and B decrease
and those of C and D increase continuously. As a result the rate of
forward reaction also decreases and the rate of the reverse reaction
increases. Eventually, the rate of the two opposing reactions equals
and the system attains a state of equilibrium.
Thus, Chemical equilibrium may be defined as: the state of a
reversible reaction when the two opposing reactions occur at the
same rate and the concentrations of reactants and products do
not change with time.
Furthermore, the true equilibrium of a reaction can be attained from
both sides. Thus, the equilibrium concentrations of the reactants and
products are the same whether we start with A and B, or C and D.

Depending on the phases of reaction species, chemical equilibrium can be
divided into two types. They are-
1) Homogenous Equilibrium:The equilibrium reaction in which all the
reactants and products remain in the same phase is called a
homogeneous equilibrium. For example
a) H2(g) + I2(g) ⇌ 2HI(g)
b) N2O4(g) ⇌ 2NO2(g)
c) CH3COOH(l) + C2H5OH(l) ⇌ CH3COOC2H5(l) + H2O(l)
2) Heterogeneous Equilibrium :The equilibrium reaction in which
all the reactants and products are not in the same phase is called a
heterogeneous equilibrium. For example,
CaCO3(s) ⇌ CaO(s) + CO2(g)
NH4Cl(s) ⇌ NH3(g) + HCl(g)
3Fe(s) + 4H2O(g) ⇌ Fe3O4(s) + 4H2(g)

We know that when a reaction say A + B ⇌ C + D attains equilibrium,
the concentrations of A and B as also of C and D remain constant with
time. Apparently it appears that the equilibrium is dead. But it is not so.
The equilibrium is dynamic. Actually, the forward and the reverse
reactions are taking place at equilibrium but the concentrations remain
unchanged.
The dynamic nature of chemical equilibrium can be easily understood
on the basis of the kinetic molecular model. The molecules of A and
B in the equilibrium mixture collide with each other to form C and D.
Likewise C and D collide to give back A and B. The collision of
molecules in a closed system is a ceaseless phenomenon. Therefore
collisions of A and B giving C and D (Forward reaction) and collisions
of C and D giving back A and B (reverse reaction) continue to occur
even at equilibrium, while concentrations remain unchanged.

(1) Constancy of concentrations
When a chemical equilibrium is established in a closed vessel at
constant temperature, concentrations of the various species in the
reaction mixture become constant.
The reaction mixture at equilibrium is called equilibrium mixture.
The concentrations at equilibrium are called equilibrium
concentrations.
The equilibrium concentrations are represented by square
brackets [ ].Thus, [A]denotes the equilibrium concentration of
substance A in moles per litre.
(2) Equilibrium can be initiated from either side
The state of equilibrium of a reversible reaction can be
approached whether we start with reactants or products.
For example, the equilibrium is established if we start the reaction with
H2 and I2 or 2HI.
H2(g) + I2(g) ⇌ 2HI(g)
2HI(g) ⇌ H2(g) + I2(g)
H2(g) + I2(g) ⇌ 2HI(g)

(3) Equilibrium cannot be attained in an open vessel.
The equilibrium can be established only if the reaction vessel is closed and no
part of the reactants or products is allowed to escape out. In an open vessel, the
gaseous reactants and/or products may escape into the atmosphere leaving
behind no possibility of attaining equilibrium.
However, the equilibrium can be attained when all the reactants and products
are in the same phase. e.g. ethanol and ethanoic acid.
(4) A catalyst cannot change the equilibrium point.
When a catalyst is added to a system in equilibrium, it speeds up the rate of both
the forward and the reverse reaction to an equal extent. Therefore, a catalyst
cannot change the equilibrium point except that it is achieved earlier.
This enhances the rate of the reaction.
(5) Value of equilibrium constant does not depend upon
the initial concentration of reactants.
It has been found that equilibrium constant must be the same when the
concentrations of reacting species are varied over a wide range.

Two Norwegian chemists, Guldberg and Waage, studied experimentally
a large number of equilibrium reactions. In 1864, they postulated a
generalisation called the law of mass action.
It states that: the rate of a chemical reaction is proportional to
the active masses of the reactants.
By the term ‘active mass’ is meant the molar concentration i.e.,
number of moles per litre. It is expressed by enclosing the formula of the
substance in square brackets.
Mathematical Expression of Law of Mass Action
Let us consider a general reaction
A + B ⇌ C + D
And let [A], [B], [C] and [D] represent the molar concentrations of A, B, C
and D at the equilibrium point.

According to the Law of Mass action,
Rate of forward reaction α [A] [B] = k1 [A] [B]
Rate of reverse reaction α [C] [D] = k2 [C] [D]
Where, k1 and k2 are rate constants for the forward and reverse reactions.
At equilibrium, rate of forward reaction = rate of reverse reaction.
Therefore, k1 [A] [B] = k2 [C] [D]
At any specific temperature k1/k2 is constant since both k1 and k2 are constants.
The ratio k1/k2 is called Equilibrium constant and is represented by the symbol
Kc, or simply k. The subscript ‘c’ indicates that the value is in terms of
concentrations of reactants and products. The equation (1) may be written as
)1...(....................
[A][B]
[C][D]
,
2
1

k
k
or
ionsConcentratReactant-[A][B]
ionsConcentratProduct-[C][D]
,constantEuilibrium ck

Consider the reaction,
Coefficient of A 2A ⇌C+D
Here, the forward reaction is dependent on the collisions of each of two
A molecules. Therefore, for writing the equilibrium expression, each
molecule is regarded as a separate entity i.e.,
A + A ⇌ C + D
Then the equilibrium constant expression is-
As a general rule, if there are two or more molecules of the same
substance in the chemical equation, its concentration is raised to
the power equal to the numerical coefficient of the substance in
the equation.
AofntCoeffficietoEqualPower-[A]
[C][D]
[A][A]
[C][D]
2
ck

The general reaction may be written as-
aA + bB ⇌ cC + dD
where a, b, c and d are numerical coefficients of the
substances A, B, C and D respectively. The equilibrium
constant expression is-
Where, Kc is the equilibrium constant. The general definition of
the equilibrium constant may thus be stated as: the product of
the equilibrium concentrations of the products divided by the
product of the equilibrium concentrations of the reactants,
with each concentration term raised to a power equal to the
numerical coefficient of the substance in the balanced
equation.
ba
dc
[B][A]
[D][C]
ck

PROBLEMS
PROBLEM-1: Give the equilibrium
constant expression for the
reaction, N2(g) + H2(g) ⇌NH3(g)
PROBLEM-2: Write the equilibrium
constant expression for the
reaction, N2O5(g) ⇌ NO2(g) + O2 (g)

Equilibrium Constant Expression in Terms of Partial Pressures
When all the reactants and products are gases, we can also formulate
the equilibrium constant expression in terms of partial pressure. The
relationship between the partial pressure (p) of any one gas in the
equilibrium mixture and the molar concentration follows from the
general ideal gas equation
The quantity n/V is the number of moles of the gas per unit volume and is simply
the molar concentration.
Thus, P = (Molar concentration)RT. i.e., the partial pressure of a gas in the
equilibrium mixture is directly proportional to its molar concentration at a
given temperature. Therefore, we can write the equilibrium constant expression
in terms of partial pressure instead of molar concentrations.
For a general reaction,
lL(g) + mM(g) ⇌ yY(g) + zZ(g)
Here, Kp is the equilibrium constant, the subscript p referring to partial pressure.
Partial pressures are expressed in atmospheres.
RT
V
n
Por,nRTPV 






m
M
l
L
z
Z
y
Y
PP
PP
][][
][][
KP 

Let us consider a general reaction,
jA(g) + kB(g) ⇌ l C(g) + mD(g)
Where, j, k, l and m are numerical coefficients of the substance, A, B, C and D
respectively. And let [A], [B], [C] and [D] represent the molar concentrations of A, B,
C and D at the equilibrium point.
The equilibrium constant expression is,
Where, all reactants and products are gases. We can write the equilibrium
constant expression in terms of partial pressures as-
Assuming that all these gases constituting the equilibrium mixture obey the ideal
gas equation, the partial pressure (p) of a gas is-
Where, n/V is the molar concentration. Thus, the partial pressures of individual
gases A, B, C and D are: PA = [A] RT; PB = [B] RT; PC = [C] RT; PD = [D] RT
)1.......(....................
[B][A]
[D[C]
K kj
ml
C 
)2.......(..........k
B
j
A
m
D
l
C
P
][P][P
][P][P
K 
RT
V
n
P 







Substituting these values in equation (2), we have-
Where, ∆n = (l + m) – (j + k), is the difference in the sums of
the coefficients for the gaseous products and reactants.
From the expression (3) it is clear that when ∆n = 0, Kp = Kc.
kj
ml
kj
ml
P
kkjj
mmll
P
[RT][RT]
[RT][RT]
[B][A]
[D][C]
K,
[RT][B][RT][A]
[RT][D][RT][C]
K


or
)3......([RT]KK
[RT]KK,
[RT]KK,
[RT]
[RT]
[B][A]
[D][C]
K,
P
P
k)(j-m)(l
P
kj
ml
kj
ml
P
n
n
c
cor
cor
or










In the equilibrium expression for a particular reaction, the concentrations
are given in units of moles/litre or mol/L, and the partial pressure are
given in atmospheres (atm). The units of Kc and Kp, depend on the
specific reaction.
(1) When the total number of moles of reactants and products are
equal.
In the equilibrium expression of these reactions, the concentration or
pressure terms in the numerator and denominator exactly cancel out.
Thus, Kc or Kp for such a reaction is without units.
Taking example of the reaction,
H2(g) + I2(g) ⇌ 2HI(g)
)(
l/L)(mol/L)(mo
(mol/L)
]][I[H
[HI]
K
2
22
2
C unitsNo
)(
(atm)(atm)
(atm)
]][P[P
][P
K
2
I2H2
2
HI
P unitsNo

(2) When the total number of moles of the reactants and products
are unequal.
In such reactions Kc will have units (mol/litre)n and Kp will have units
(atm)n, where n is equal to the total number of moles of products
minus the total number of moles of reactants.
Thus, for the reaction: N2O4(g) ⇌ 2NO2(g)
So, the units for KC are mol/L and
KP units are atm.
For the reaction : N2(g) + 3H2(g) ⇌2NH3(g)
Units for KC are and KP may be found as follows
Thus,
units of KC are mol–2 L2 and
units of Kp are atm–2.
(mol/L)
(mol/L)
(mol/L)
]O[N
][NO
K
2
42
2
2
C 
(atm)
(atm)
(atm)
][P
][P
K
2
ON
2
NO
P
42
2

22-
3
2
3
22
2
3
C /Lmol
l/L)(mol/L)(mo
(mol/L)
]][H[N
][NH
K 
2-
3
2
3
NH
2
NH
P (atm)
(atm)(atm)
(atm)
]][P[P
][P
K
22
3


Liquid Systems (Liquid equilibria)
The chemical equilibrium in which all the reactants and products
are in the liquid phase, are referred to as the liquid equilibria.
Like the gas-phase equilibria, the liquid equilibria are also called
homogeneous equilibria.
For example, alcohols and acids react to form esters and water.
CH3COOH(l) + C2H5OH(l) ⇌ CH3COOC2H5(l) + H2O(l)
Let us start with a mole of acetic acid and b moles of alcohol.
If x moles of acetic acid react with x moles of ethyl alcohol, x
moles of ester and x moles of water are produced when the
equilibrium is established.

Now the moles present at equilibrium are:
CH3COOH = (a – x) moles
C2H5OH = (b – x) moles
CH3COOC2H5 = x moles
H2O = x moles
CH3COOH(l) + C2H5OH(l) ⇌ CH3COOC2H5(l) + H2O(l)
(a – x) (b – x) x x
If V litter be the total volume of the equilibrium mixture, the concentrations of the various species
are:
The equilibrium constant expression may be written as
It may be noted that the volume terms V in the numerator and denominator cancel out. In liquid
systems when there is a change in the number of moles as a result of the reaction, it is
necessary to consider the volume V while calculating the equilibrium constant K.
V
x
O][H
V
x
]HCOOC[CH
V
x-b
OH]H[CH
V
x-a
COOH][CH
2
523
52
3




x)x)(b(a
x
or,K
V
b-x
V
a-x
V
x
V
x
K





2

Chemical equilibria in which the
reactants and products are not all in the
same phase is called heterogeneous
equilibria.
Example: The decomposition of calcium
carbonate upon heating to form calcium
oxide and carbon dioxide.
If the reaction is carried in a closed vessel,
the following equilibrium is established.
CaCO3(s) ⇌ CaO(s) + CO2(g)
Other examples:
NH4Cl(s) ⇌ NH3(g) + HCl(g)
3Fe(s) + 4H2O(g) ⇌ Fe3O4(s) + 4H2(g)

In 1884, the French Chemist Henry Le Chatelier proposed a general principle which
applies to all systems in equilibrium. The principle states that-
“When a system is at equilibrium a change in any one of the factors upon
which the equilibrium depends will cause the equilibrium to shift in a direction
such that the effect of the change is diminished.”
There are three ways in which the change can be caused on a chemical equilibrium:
(1) Changing the concentration of a reactant or product.
(2) Changing the pressure (or volume) of the system.
(3) Changing the temperature.
Thus, when applied to a chemical reaction in equilibrium, Le Chatelier’s principle can
be stated as: if a change in concentration, pressure or temperature is caused to
a chemical reaction in equilibrium, the equilibrium will shift to the right or the
left so as to minimize the change.

Effect of a Change
in Concentration
We can restate Le Chatelier’s principle for the
special case of concentration changes: when
concentration of any of the reactants or
products is changed, the equilibrium shifts in a
direction so as to reduce the change in
concentration that was made.
A change in the concentration of a reactant or
product can be effected by the addition or removal
of that species.
Let us consider a general reaction,
A + B ⇌ C
When a reactant, say, A is added at equilibrium, its
concentration is increased.
The forward reaction alone occurs momentarily. According to Le Chatelier’s
principle, a new equilibrium will be established so as to reduce the
concentration of A. Thus, the addition of A causes the equilibrium to shift to
right. This increases the concentration (yield) of the product C.

Following the same line of argument, a decrease in the concentration of A by its
removal from the equilibrium mixture, will be undone by shift to the equilibrium
position to the left. This reduces the concentration (yield) of the product C.
Let us illustrate the effect of change of concentration on a system at equilibrium by
taking example of the ammonia synthesis reaction:
N2(g) + 3H2(g) ⇌ 2NH3(g)
When N2 (or H2) is added to the equilibrium already in existence (equilibrium I), the
equilibrium will shift to the right so as to reduce the concentration of N2 (Le Chatelier’s
principle). The concentration of NH3 at the equilibrium II is more than at equilibrium I.
The results in a particular case after the addition of one mole/litre are given below.
Obviously, the addition of N2 (a reactant) increases the concentration of NH3, while
the concentration of H2 decreases. Thus, to have a better yield of NH3, one of the
reactants should be added in excess.

To predict the effect of a change of pressure, Le Chatelier’s principle may be stated
as: when pressure is increased on a gaseous equilibrium reaction, the
equilibrium will shift in a direction which tends to decrease the pressure.
The pressure of a gaseous reaction at equilibrium is determined by the total
number of molecules it contains. If the forward reaction proceeds by the reduction of
molecules, it will be accompanied by a decrease of pressure of the system and vice
versa.
Let us consider a reaction,
A + B ⇌ C
The combination of A and B produces a decrease of number of molecules while the
decomposition of C into A and B results in the increase of molecules. Therefore, by the
increase of pressure on the equilibrium it will shift to right and give more C. A decrease
in pressure will cause the opposite effect. The equilibrium will shift to the left when C
will decompose to form more of A and B.

The reactions in which the number of product molecules is
equal to the number of reactant molecules, are unaffected
by pressure changes.
For example, H2 (g) + I2 (g)⇌ 2HI (g)
In such a case, the system is unable to undo the increase
or decrease of pressure.
In light of the above discussion, we can state a general
rule to predict the effect of pressure changes on chemical
equilibria.
The increase of pressure on a chemical equilibrium shifts
it in that direction in which the number of molecules
decreases and vice-versa.

Chemical reactions consist of two opposing reactions.
If the forward reaction proceeds by the evolution of heat
(exothermic), the reverse reaction occurs by the absorption of
heat (endothermic).
Both these reactions take place at the same time and equilibrium exists
between the two. If temperature of a reaction is raised, heat is added to
the system. The equilibrium shifts in a direction in which heat is
absorbed in an attempt to lower the temperature. Thus, the effect of
temperature on an equilibrium reaction can easily be predicted by the
following version of the Le Chatelier’s principle.
When temperature of a reaction is increased, the equilibrium
shifts in a direction in which heat is absorbed.

Let us consider an exothermic reaction,
A + B ⇌ C + heat
When the temperature of the system is increased, heat is supplied to it from
outside. According to Le Chatelier’s principle, the equilibrium will shift to the
left which involves the absorption of heat. This would result in the increase
of the concentration of the reactants A and B.
In an endothermic reaction the increase of temperature will shift the
equilibrium to the right as it involves the absorption of heat.
X + Y + heat ⇌ Z
This increases the concentration of the product Z.
In general, we can say that the increase of temperature favors the reverse
change in an exothermic reaction and the forward change in an endothermic
reaction.

Formation of Ammonia from N2 and H2
The synthesis of ammonia from nitrogen and hydrogen is an
exothermic reaction.
N2 + 3H2 ⇌ 2NH3 + 22.2 kcal
When the temperature of the system is raised, the equilibrium
will shift from right-to-left which absorbs heat (Le Chatelier’s
principle). This results in the lower yield of ammonia.
On the other hand, by lowering the temperature of the system, the
equilibrium will shift to the right which evolves heat in an attempt to
raise the temperature. This would increase the yield of
ammonia.
But with decreasing temperature, the rate of reaction
is slowed down considerably and the equilibrium is reached slowly.
Thus, in the commercial production of ammonia, it is not feasible to use tempe
rature much lower than 500°C. At lower temperature, even in the presence of
a catalyst, the reaction proceeds too slowly to be practical.

With the help of Le Chatelier’s principle we can work out the optimum conditions for securing the
maximum yield of products in industrial processes.
Synthesis of Ammonia (Haber Process)
The manufacture of ammonia by Haber process is represented by the equation-
N2(g) + 3H2(g) ⇌ 2NH3(g) + 22.0 kcal
A look at the equation provides the following information:
(a) The reaction is exothermic
(b) The reaction proceeds with a decrease in the number of moles.
(1) Low temperature. By applying Le Chatelier’s principle, low temperature will shift the equilibrium to
the right. This gives greater yield of ammonia. In actual practice a temperature of about 450°C is used
when the percentage of ammonia in the equilibrium mixture is 15.
(2) High pressure. High pressure on the reaction at equilibrium favors the shift of the equilibrium to the
right. This is so because the forward reaction proceeds with a decrease in the number of moles. A
pressure of about 200 atmospheres is applied in practice.
(3) Catalyst. As already stated, low temperature is necessary for higher yield of ammonia. But at
relatively low temperatures, the rate of reaction is slow and the equilibrium is attained in a long time.
To increase the rate of reaction and thus quicken the attainment of equilibrium, a catalyst is used. Finely
divided iron containing molybdenum is employed in actual practice.
Molybdenum acts as a promoter that increases the life and efficiency of the catalyst.
Pure N2 and H2 gases are used in the process. Any impurities in the gases would poison the catalyst
and decrease its efficiency.

Manufacture of Sulphuric Acid (Contact Process)
The chief reaction used in the process is-
2SO2(g) + O2(g) ⇌ 2SO3(g) + 42 kcal
Following information is revealed by the above equation:
(a) The reaction is exothermic.
(b) The reaction proceeds with a decrease in number of moles.
On the basis of Le Chatelier’s principle, the conditions for the maximum yield can be worked out
as below:
(1) Low temperature. Since the forward reaction is exothermic, the equilibrium will shift on the
right at low temperature. An optimum temperature between 400-450°C is required for the
maximum yield of sulphur trioxide.
(2) High pressure. Since the number of moles is decreased in the forward reaction, increase of
pressure will shift the equilibrium to the right. Thus, for maximum yield of SO3, 2 to 3 atmosphere
pressure is used.
(3) Catalyst. At the low temperature used in the reaction, the rate of reaction is slow and the
equilibrium is attained slowly. A catalyst is, therefore, used to speed up the establishment of the
equilibrium. Vanadium pentoxide, V2O5, is commonly used and it has replaced the earlier catalyst
platinum asbestos which was easily poisoned by the impurities present in the reacting gases. All
samples, SO2 and O2 used for the manufacture of sulphuric acid must be pure and dry.

Manufacture of Nitric Acid (Birkeland-Eyde Process)
Nitric acid is prepared on a large scale by making use of the reaction
N2 (g) + O2 (g) ⇌ 2NO (g) – 43.2 kcal
The equation tells us that:
(a) The reaction proceeds with no change in the number of moles.
(b) The reaction is endothermic and proceeds by absorption of heat.
The favorable conditions for the maximum yield of NO are:
(1) High temperature. Since the forward reaction is endothermic, increase of
temperature will favor it (Le Chatelier’s principle). Thus, a high temperature of
the order of 3000°C is employed to get high yield of nitric acid.
(2) No effect of pressure. Since the forward reaction involves no change in
the number of moles, a change in pressure has no effect on the equilibrium.
(3) High concentration. The formation of nitric oxide is favored by using high
concentrations of the reactants i.e. N2 and O2.

Chemical equilibrium manik

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