This document discusses the key concepts of chemical equilibrium. It defines reversible reactions as those that can proceed in both the forward and backward directions simultaneously. At equilibrium, the rates of the forward and reverse reactions are equal and the concentrations of reactants and products remain constant. Several examples of reversible reactions are provided. Characteristics of chemical equilibrium include the constancy of concentrations at equilibrium and the independence of the equilibrium constant from the initial concentrations. Le Chatelier's principle is introduced, which states that if a system at equilibrium experiences a change, it will shift its position to counteract that change. The effects of changing concentration, pressure, temperature, and adding a catalyst are described based on this principle. Industrial processes for maximizing yields of important chemicals

1. CHEMICAL EQUILIBRIUM

2. REVERSIBLE REACTIONS
• A reaction which can go in forward &
backward direction simultaneously is called
reversible reaction.
Depicting Equilibrium
In a system at equilibrium, both the forward and
reverse reactions are running simultaneously.
We write the chemical equation with a double
arrow:

3. SOME EXAMPLES OF REVERSIBLE
REACTIONS
CaCO3(s) CO2 (g) + CaO(s)
H2 (g) + I2 (g) 2 HI(g)
N2(g) + 3H2(g) 2NH3(g)
CH3COOC2H5(aq) + H2O(aq) CH3COOH(aq) + C2H5OH(aq)

4. Chemical Equilibrium
(Definitions)
A chemical system where the concentrations of
reactants and products remain constant over time.
On the molecular level, the system is dynamic: The
rate of change is the same in either the forward or
reverse directions.
As a system approaches equilibrium, both the
forward and reverse reactions are occurring.
At equilibrium, the forward and reverse reactions
are proceeding at the same rate.

6. A System at Equilibrium
• Once equilibrium is achieved, the amount of
each reactant and product remains constant.
Concentrations become constant Rates become equal

7. CHARACTERISTICS OF CHEMICAL
EQUILIBRIUM
1- Constancy of concentrations
When a chemical reaction is established in a
closed vessel at constant temperature conc. of
various species in reaction mixture become
constant & the reaction mixture at equilibrium
is called Equilibrium mixture & the conc. At
equilibrium is called equilibrium conc.
UNITS: moles/litre

8. 2-Equilibrium can be initiated from
either side
N2(g) + 3H2(g) 2NH3(g)
Add nitrogen and hydrogen gases together in anyAdd nitrogen and hydrogen gases together in any
proportions. Nothing noticeable occurs.proportions. Nothing noticeable occurs.
Add heat, pressure and a catalyst, you smellAdd heat, pressure and a catalyst, you smell
ammonia => a mixture with constant concentrationsammonia => a mixture with constant concentrations
of Nof N22 , H, H22 and NHand NH33 is produced.is produced.
Start with just ammonia and catalyst. NStart with just ammonia and catalyst. N22 and Hand H22 willwill
be produced until a state of equilibrium is reached.be produced until a state of equilibrium is reached.
As before, a mixture with constant concentrations ofAs before, a mixture with constant concentrations of
nitrogen, hydrogen and ammonia is produced.nitrogen, hydrogen and ammonia is produced.

9. • No matter what the starting composition of
reactants and products, the same ratio of
concentrations is realized when equilibrium is
reached at a certain temperature and pressure.

10. 3-Equilibrium cannot be attained in
open vessel
• Equilibrium can be established only if the
reaction vessel is closed & no part of the
reactants or products is allowed to escape out.
• In an open vessel, the gaseous reactants or
products may escape out into the atmosphere
leaving behind no possibility of attaining
equilibrium.
• However equilibrium can be attained when all
the reactants & products are in the same phase
e.g,ethanol & ethanoic acid.

11. 4-A catalyst cannot change the
equilibrium point
• Catalysts increase the rate of both the forward
and reverse reactions. Equilibrium is achieved
faster, but the equilibrium composition remains
unaltered. A catalyst lowers the activation
energy barrier of a reaction. The diagram
shows that the catalyst lowers the activation
energy for the forward and reverse reactions by
the same amount. Consequently a catalyst
does not affect the position of the equilibrium
but it does speed up the rate at which
equilibrium is reached.

13. 5-Equilibrium constant is
independent of the initial conc. of
reactants
6- At equilibrium the Gibbs free energy (G) is
the minimum & any change taking place at
equilibrium proceeds without change in free
energy i.e; ∆G=0

14. Law of Mass actionLaw of Mass action
(EquiLibriuM ExprEssion)(EquiLibriuM ExprEssion)
• Two Norwegian chemists, Guldberg & Waage
studied experimentally a large no of equilibrium
reactions. In 1864, they postulated a
generalization called LAW OF MASS ACTION
which states that:
• “The rate of chemical reaction is proportional to
the active masses of reactants”
• ACTIVE MASS: means the molar conc. i.e. no. of
moles per liter expressed by enclosing formula of
substance in square brackets.

15. EquiLibriuM
constant:EquiLibriuM Law
• For a general equilibrium
the equilibrium expression is:
where K is the Equilibrium Constant. (Units for K will vary.)

16. • EQUILIBRIUM CONSTANT can be defined as:
• The product of equilibrium conc. Of products
divided by the product of equilibrium conc. Of
reactants, with each conc. Term raised to a
power equal to the coefficient of substance in
the balanced equation.

17. The Equilibrium Constant
Forward reaction: Reverse reaction:
Rate LawRate law

18. At equilibrium
Rearranging gives:

19. Equilibrium Expression
• 4 NH4 NH33(g) + 7 O(g) + 7 O22(g)(g) ↔↔ 4 NO4 NO22(g) + 6H(g) + 6H22O(g)O(g)
• Write the Equilibrium Expression for theWrite the Equilibrium Expression for the
reaction. The expression will have eitherreaction. The expression will have either
concentration units of mol/L (M), or units ofconcentration units of mol/L (M), or units of
pressure (atm) for the reactants and products.pressure (atm) for the reactants and products.
What would be the overall unit for K usingWhat would be the overall unit for K using
Molarity and atm units respectively.Molarity and atm units respectively.
K=
NO HO
NH O
2
2
2
4 6
3
4 7

20. SAMPLE EXERCISE: Write the equilibrium expression
for Kc for the following reactions:

21. What Does the Value of K Mean?
• If K >> 1, the reaction is
product-favored; product
predominates at
equilibrium.
• If K << 1, the reaction is
reactant-favored;
reactant predominates at
equilibrium.

22. The Equilibrium Constant in terms
of Partial Pressures
Because pressure is proportional to concentration
for gases, the equilibrium expression can also be
written in terms of partial pressures (instead of
concentration):

23. Relationship between Kc and Kp
From the ideal gas law we know that:
Substituting P=[A]RT into the expression for Kp for
each substance, the relationship between Kc and Kp
becomes: Kp = Kc (RT)∆n

24. Where:
• Partial pressures are proportional to
concentration via PV = nRT.
• Thus, for gas reactions, partial pressures can
be used in place of concentrations.
∆n = (moles of gaseous product) – (moles of gaseous
reactant)
• To be able to convert between Kc and Kp,
we need a relationship between
concentration and pressure.

25. b
B
a
A
d
D
c
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P
RT
P
RT
P
RT
P
K
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26. Where,∆n=(c+d)-(a+b),the difference in
sums of coefficients of gaseous products &
reactants.
• The Δn is the change in the number of moles
of gas when going from reactants to products.
• When does Kp equal Kc?
When ∆n=0, Kp=Kc
• Equilibrium constants are really defined in
terms of activity, not concentration.

27. PROBLEM
• For the reaction 2SO3(g)  2SO2(g) + O2(g), we
can write two equilibrium expressions
• [SO2]= 0.27 mol/L [O2] = 0.40mol/L [SO3] =
0.33mol/L, calculate the value of Kc?
2
SO
O
2
SO
2
3
2
2
2
3
22
or
]SO[
]O[]SO[
P
PP
KK pc ==

28. LIQUID SYSTEMS
• The chemical equilibrium in which all the
reactants & products are in liquid phase.
• Also called Homogeneous equilibrium.
• For-example alcohols & acids react to form
esters & water.
CH3COOH(aq) + C2H5OH(aq) CH3COOC2H5(aq) + H2O(aq)

29. HETEROGENEOUS EQUILIBRIUM
•If one or more reactants or products are in aIf one or more reactants or products are in a
different phase, the equilibrium isdifferent phase, the equilibrium is
heterogeneous.heterogeneous.
The conc. Of pure solids & liquids are notThe conc. Of pure solids & liquids are not
included in equilibrium constant expression asincluded in equilibrium constant expression as
the conc of pure solid/liquid is fixed & cannotthe conc of pure solid/liquid is fixed & cannot
vary. Kvary. Kcc= [CO= [CO22]]
CaCO3 (s) CO2 (g) + CaO(s)

30. Le ChateLier’s
PrinCiPLe
• “If a system at equilibrium is disturbed by a
change in temperature, pressure, or the
concentration of one of the components, the
system will shift its equilibrium position to the
right or to the left so as to counteract the
effect of the disturbance.”
• “When a stress is applied on a system in
equilibrium the system tends to adjust itself
so as to reduce the stress.”

31. EFFECT OF CHANGE IN
CONCENTRATION
• When the conc. Of any of the reactants or
products is changed the equilibrium shifts in a
direction so as to reduce the change in conc.
That was made.
• Addition of inert gas does not affect the
equilibrium position.

32. EFFECT OF CHANGE IN
PRESSURE
• When pressure is increased on a gaseous
equilibrium reaction, the equilibrium will shift
in direction which tends to decrease the
pressure.

33. EFFECT OF CHANGE IN
TEMPERATURE
• When temperature of a reaction is increased
the equilibrium shifts in a direction in which
heat is absorbed.
• The increase of temperature favors the
reverse change in an exothermic reaction &
the forward change in an endothermic
reaction.

34. EFFECT OF CATALYST
• A catalyst lowers the activation energy barrier
for any reaction….in both forward and reverse
directions!
• A catalyst will decrease the time it takes to
reach equilibrium.
• A catalyst does not effect the composition of
the equilibrium mixture

35. CONDITIONS FOR MAXIMUM
YIELD IN INDUSTRIAL PROCESSES
• Synthesis of ammonia (HABER PROCESS)
• Exothermic reaction
• Low Temp (4500
C)
• High Pressure (200atm)
• Catalyst (finely divided iron containing
molybdenum)
N2(g) + 3H2(g) 2NH3(g)

36. This apparatus helps
push the equilibrium
to the right by
removing the
ammonia (NH3) from
the system as a liquid.

37. Manufacture of Sulphuric acid
(CONTACT PROCESS)
• 2SO2(g) + O2(g)  2SO3(g)
• Exothermic reaction
• Low temperature(400-5000
C)
• High pressure (2-3atm)
• Catalyst (Pt & V2O5)

38. Manufacture of Nitric acid
(Birkland eyde process)
• Endothermic reaction
• High temperature (30000
C)
• N2+O2  NO