Revision for the IB SL exam

1. Equilibrium
By: Ali Wagih

2. Dynamic Equilibrium
Definition:
• Equilibrium is dynamic (reactions doesn’t stop)
• Equilibrium is achieved in a closed system
• At equilibrium there's no change in macroscopic properties (no
change in color and density)
• Equilibrium can be reached from either direction (doesn’t matter if
started with all reactants or all products)
Physical Systems:
• There will come a time when the rate of the forward reaction is
equal to the rate of the reverse reaction. At this point, the system
has reached equilibrium.
Chemical Systems:
• The concentrations of both reactants and products remain constant
over time

3. Equilibrium Constant
• For the hypothetical homogeneous chemical reaction:
aA+bB cC+dD
The equilibrium constant is defined as
KC = [C]^c [D]^d , where [A] signifies the molar
concentration of species
[A]^a [B]^b
• Note that the expression for the equilibrium constant
includes only solutes and gases; pure solids and liquids
do not appear in the expression.

4. Magnitude of the Equilibrium Constant
• When Kc >> 1, the reaction goes almost to completion
• When Kc << 1, the reaction hardly proceeds
If the value of Kc is large (Kc >> 1 ):
• the equilibrium lies to the right hand side - i.e. there is
a much greater concentration of the products than the
reactants
If the value of Kc is small (Kc << 1):
• the equilibrium lies to the left hand side - i.e. there is a
much greater concentration of reactants than
products.

5. Le Chatelier’s Principle
• States: when a stress is brought to bear on a
system at equilibrium, the system will react in
the direction that serves to relieve the stress.

6. Temperature
• The only thing which can change the value of Kc for a given
reaction is a change in temperature.
• The effect of a change of temperature on a reaction will
depend on whether the reaction is exothermic or
endothermic
• When the temperature increases, Le Chatelier's principle
says the reaction will proceed in such a way as to
counteract this change, i.e. lower the temperature.
• Endothermic reactions (Hrxn > 0 )will move forward, and
exothermic reactions (Hrxn < 0 )will move backwards
• The reverse is true for a lowering of temperature.

7. Pressure
• According to Le Chatelier, the position of equilibrium
will move in such a way as to counteract the change.
That means that the position of equilibrium will move
so that the pressure is reduced again.
• Increasing the pressure on a gas reaction shifts the
position of equilibrium towards the side with fewer
molecules.
• In the case of the same number of molecules on both
sides of the equilibrium reaction, increasing the
pressure has no effect whatsoever on the position of
the equilibrium.

8. Concentration
• When the concentration of a product is
increased, the reaction proceeds in reverse to
decrease the concentration of the products.
a+b c+d , reaction moves backwards
• When the concentration of a reactant is
increased, the reaction proceeds forward to
decrease the concentration of reactants.
a+b c+d , reaction moves forward

9. Value of the Equilibrium Constant
Equilibrium constants aren't changed if you
change the pressure of the system.
Equilibrium constants aren't changed if you add
a catalyst.
Temperature:
• Equilibrium constants are changed if you
change the temperature of the system
• As the temperature increases, the value of Kc
falls.

10. Catalyst
• Adding a catalyst makes absolutely no
difference to the position of equilibrium
• Le Chatelier's Principle doesn't apply to them.
• A catalyst speeds up the rate at which a
reaction reaches dynamic equilibrium.

11. The Haber Process
• Ammonia: (Haber Process)
• N2 + 3H2 2NH3 -92kJ/mol
• 400-450 degrees Celsius is not a low
temperature, which defies the Le Chatelier’s
principle, as it would be expected to have a lower
temperature to produce more ammonia.
However, this was done to speed up the rate of
the process
• 200 atm is not a very high pressure, but it was
used because very high pressures are expensive
to achieve

12. Sulphur Trioxide Formation
• 2SO2(g) + O2(g) <=> 2SO3(g) ∆H = -196 kJ/mol
Conditions required for the reaction:
• 1. Temperature: 450 oC
• 2. Pressure : 1 atm
• 3. Catalyst: Vanadium (V) oxide