EquilibriumUsing and Controlling Reactions

1

Chemical Equilibrium

Most chemical reactions don’t go to completion.

Instead with the right conditions they will reach a balance between reactants and products.

This is called Equilibrium

2

Chemical Equilibrium

Forward reaction: Reactants Products

Back reaction: Products Reactants

Overall: Reactants Products Not a static state, but a dynamic

(moving) state (forward and back reactions are occurring)

Rate of forward reaction = Rate of back reaction 3

Chemical Equilibrium

Concentration of reactants and products is constant

No changes in macroscopic properties (colour intensity, pressure, pH)

CONDITIONS FOR EQUILIBRIUM Closed system. No gain or loss of

reactants or products to or from the surroundings eg. Solutions, Sealed containers

Constant Temperature

4

Chemical Equilibrium

5

Chemical Equilibrium

Initially the rate of forward reaction is high because the concentration of reactants is high.

Rate decreases as [reactants] decreases

Rate of back reaction is zero until some products form.

Rate increases as [products] increases Until at equilibrium rate of forward

reaction = rate of back reaction

6

Equilibrium constant

If reactants A and B reacted to form products C and D this could be represented by the following reaction. ( a, b, c and d refer to mole ratio in which they react) aA + bB cC + dD

At equilibrium the concentrations of these species are constant and can be used to calculate the equilibrium constant for the reaction.

7

Equilibrium constant

8

ba

dcc

BA

DCK

Kc has no units

Constant value at constant temperature

[A] [B] etc represent concentrations in molL-1

Equilibrium constant

The size of Kc indicates the yield of the products relative to the amount of reactants at equilibrium.

High Kc values > 10 indicate a high yield of products while Kc values < 0.1 indicate a low yield of products.

Kc is only affected by temperature changes not changes in pressure or concentrations.

9

Le Châtelier's Principle

Henry Le Châtelier was a French chemist.

Observed in 1888 how concentration changes, pressure changes and heat changes altered the position of the equilibrium.

10

Le Châtelier's Principle

If the conditions of equilibrium are changed then the system will respond in such a way as to counteract the introduced changes, if that is possible.

“Reaction proceeds to partially oppose the change”

11

Concentration Changes

Conditions: Constant volume, Constant temperature

If a reactant or product is added or removed, then the reaction will proceed in the direction to decrease/ increase the species that was altered.

12

Concentration Changes

Reactants Products

[Reactant] equilibrium moves to the right

[Product] equilibrium moves to the left

[Reactant] equilibrium moves to left

[Product] equilibrium moves to right

13

Concentration Changes

14

N2O4 2NO2 ΔH= +58kJmol-1

Pressure Changes

Pressure is to the total number of molecules in the gas phase.

Reduction in the total number of molecules lowers the internal pressure.

Increase in the total number of molecules raises the internal pressure.

The pressure of an equilibrium system can be changed by changing the volume of the system.

15

Pressure Changes

INCREASE PRESSURE Decrease volume. Concentration

of all species increases. Reaction proceeds to decrease pressure. Moves to reduce moles.

DECREASE PRESSURE Increase volume. Concentration of

all species decreases. Reaction proceeds to increase pressure. Moves to increase moles.

16

Pressure Changes

17

N2O4 2NO2 ΔH= +58kJmol-1

Pressure Changes

If there are equal number of molecules on both sides then an equilibrium can not adjust to pressure changes.

18

Temperature Changes

Quoted H value always refers to the forward reaction.

If the forward reaction of an equilibrium system is exothermic, then the back reaction is endothermic (same value).

19

Temperature Changes

Temperature increase causes equilibrium to shift in endothermic direction to absorb heat.

Temperature decrease causes equilibrium to shift in exothermic direction to release heat.

20

Temperature Changes

21

N2O4 2NO2 ΔH= +58kJmol-1

Temperature Changes

Temperature changes do affect Kc

EXOTHERMIC (Forward reaction) If Temp increases Kc decreases If Temp decreases Kc increasesENDOTHERMIC (Forward

reaction) If Temp increases Kc increases If Temp decreases Kc decreases

This means that calculated Kc values are specific for a given temperature.

Catalysts

Increase the rate of both forward and back reactions.

Equilibrium is reached more quickly.

Do not affect the position of the equilibrium.

23

Calculations: Example

For the reaction at 273oC;CO2(g) + H2(g) H2O(g) +

CO(g) 0.750 moles of H2 and 1.20 moles of

CO2 were allowed to come to equilibrium in a 2.00L flask. At equilibrium 0.600 moles of CO was present.

Calculate Kc24

Calculations: One for you to try

A 2.50L reaction flask containing 2.95 moles of nitrogen and 4.16 moles of hydrogen was allowed to come to equilibrium according to:

N2(g) + 3H2(g) 2NH3(g) At equilibrium 0.240 moles of

ammonia was present. Verify that the equilibrium constant is 0.00232 (3 sf) under these conditions.

25