Federal Board of Intermediate and Secondary Education

1. Chapter 7:
Chemical Equilibrium
Naveed Mallana
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2. Reduced Syllabus For 2021
 7. CHEMICAL EQUILIBRIUM
 Introduction
 7.1.1 Law of Mass Action
 7.1.2 Examples of Equilibrium Constant expression
 7.1.3 Units of equilibrium constant
 7.1.4 Equilibrium expression including partial pressure, no of moles and mole fraction.
 7.1.5 Types of equilibrium
 7.1.7 Applications of equilibrium constant
 7.2 Factors Affecting Equilibrium ( Le-Chatellier Principle)
 7.2.1 Effect of Change in Concentration
 7.2.2 Effect of Change in Pressure or Volume
 7.2.3 Effect of Change in Temperature
 7.3 Industrial Application of Le-Chatellier Principle (Haber’sProcess)
 7.5 Common Ion Effect
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3. Reversible Reactions and
Dynamic Equilibrium
When a reaction starts, the
reactants are consumed and
products are made.
• forward reaction = reactants 
products
• Therefore, the reactant
concentrations decrease and the
product concentrations increase.
• As reactant concentration
decreases, the forward reaction
rate decreases.
Eventually, the products can react
to re-form some of the reactants.
• reverse reaction = products 
reactants
• assuming the products are not
allowed to escape
• As product concentration
increases, the reverse reaction
rate increases.
Processes that proceed in both the
forward and reverse direction are
said to be reversible.
• reactants  products
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4. Law of Mass Action
 Two chemists C.M Guldberg and P Wage in 1864 proposed the law of
mass action as a
 It states that "the rate at which a substance reacts is proportional to Its
active mass and the rate of a chemical reaction is proportional to the
product of the active masses of the reacting substances".
 It can also be defined as
 The rate of chemical reaction is proportional to the product of molar
concentrations of each reacting substance raised to a power
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6. Assignment: Find the value of Kc 6

7. Conditions for Equilibrium
• Kc applies only at equilibrium.
• Kc is independent of initial concentration of reactants and products but depends upon temperature. At a given temperature,
it has only one value. Whether we start reaction with pure reactants or pure products or any composition in between, the
value of Kc remains unchanged.
• Kc is related to the coefficients of the balance chemical equation. The concentration of the products are placed in the
numerator and those of reactants in the denominator. Each concentration is raised to a power equal to its coefficient in the
balance chemical equation.
• The magnitude of Kc indicates the position of equilibrium.
It is important to recognize the difference between the equilibrium constant expression and rate law of a reaction. You will
learn about rate law expression in chapter on chemical kinetics. In both these expressions, concentration terms are raised to
powers. The rate law describes how the rate of a reaction changes with concentration. It cannot be written from the balanced
chemical equation. Whereas the equilibrium expression describes the concentration of reactants and products when the net
rate of reaction is zero. It can be written from a balance chemical equation.
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8. Units of Equilibrium Constant:
 1. Write equilibrium constant expression.
 2. Write mol.dm^-3 as units of concentration of each species within
square brackets.
 3. Simplify the expression.
 Equilibrium constant may or may not have units. Equilibrium constant has
no units if the
 number of moles of the reactants are equal to the number of moles of
the products.
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9. Equilibrium expression including partial
pressure, no of moles and mole fraction
 Consider the general gaseous equation
 Equilibrium constant 𝐾 𝑝 in terms of partial pressures is given by:
 𝐾 𝑝 is related with 𝐾𝑐 by the following equation.
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10.  When equilibrium concentrations of reactant and products are expressed
in terms of moles, the equilibrium 𝐾 𝑛 can be written as.
 𝐾 𝑝 is related with 𝐾 𝑛
 When equilibrium concentrations of reactant and products are expressed
in terms of mole fractions, the equilibrium 𝐾𝑥 can be written as.
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11. Types of Equilibrium
Homogeneous Equilibrium:
An equilibrium system in which all of the reactants and products are in the same phase.
Heterogeneous Equilibrium:
Equilibria which involve more than one phases are called heterogeneous equilibria.
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12. Applications of Equilibrium Constant:
• Direction of the chemical
reaction
• Extent of the chemical reactions
• Effect of changes in condition of
the chemical reaction on
equilibrium.
Equilibrium
constant for a
reaction can
be used to
predict many
important
features of the
reaction.
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13. 13
Predicting the Direction of Reaction:
• The [products] will decrease and [reactants]
will increase.
If Q > K, the reaction will
proceed fastest in the reverse
direction.
• The [products] will increase and [reactants] will
decrease.
If Q < K, the reaction will
proceed fastest in the forward
direction.
• The [products] and [reactants] will not change.
If Q = K, the reaction is at
equilibrium.
If a reaction mixture contains just reactants, Q = 0, and the reaction will proceed in the forward
direction.
If a reaction mixture contains just products, Q = ∞, and the reaction will proceed in the reverse
direction.

14. The Extent of Chemical Reaction
The extent of a chemical reaction can be
predicted by considering the magnitude of
equilibrium constant. There are three
possibilities.
• 𝑲 𝒄 is very large
• 𝑲 𝒄 is very small
• 𝑲 𝒄 is neither very small nor very large
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15. 𝑲 𝒄 is very large:
 Very large value of 𝐾𝑐 indicates that the reaction goes virtually to
completion.
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16. 𝑲 𝒄 is very Small:
 The small value of Kc indicates that the reaction has very little tendency to
move in forward direction.
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17. Kc is neither very small nor very large 17

18. Factors Affecting Equilibrium:
Factors that control the position of chemical
equilibrium are called factors affecting equilibrium.
These factors are
• The Effect of change in concentration
• The effect of Pressure change
• The effect of Change in temperature
• The effect of Addition of Catalyst
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Le Chatelier’s Principle: It states that if a change is imposed on a system at
equilibrium, the position of the equilibrium will shift in a direction which tend
to reduce that change.

19. The effect of Change in Concentration
When the concentration of
one of the reactants or
products in equilibrium
mixture is disturbed then the
system will shift to
accommodate the substance
added or removed and
restore equilibrium again.
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20. The effect of Pressure change
When pressure on a gaseous
at equilibrium is increased, the
system tends to reduce the volume
to undo or minimize the effect of
increased pressure. This is done by
decreasing the total number of
gaseous molecules in the system.
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21. The Effect of change in Temperature
 Exothermic reactions favor reverse reaction with the increase of
temperature and vice versa.
 Endothermic reaction favor forward direction with the increase of
temperature and vice versa.
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22. The effect of Addition of Catalyst
 Catalyst added to a reaction mixture speeds up both the forward and
reverse reaction to the same degree and has no effect on the equilibrium.
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23. Industrial Application of Le-Chatelier’s Principle
(Synthesis of Ammonia by Haber’s Process)
 The manufacturing of Ammonia by Haber’s process is presented y the
following equation.
 This equation provide the following information.
 The reaction is exothermic.
 The reaction proceeds with a decrease in number of molecules or moles.
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24.  Le-Chatelier principle suggests three ways to get maximum yield of
ammonia.
 Low temperature
 High pressure
 Continual Removal of Ammonia
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25. Common Ion Effect
The phenomenon in which
the degree of ionization or
solubility of an electrolyte
is suppressed the addition
of highly soluble electrolyte
containing a common ion
is called common ion
effect.
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