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APPLICATIONS OF CHEMICAL EQUILIBRIUM (LIVE) 19 MAY 2015Section A: Summary Notes
What is equilibrium?
Reactions that take place in both the forward and reverse directions simultaneously are
calledreversible reactions.
Observable macroscopic changes stop, while microscopic changes continue as reactants
change to products, and products change back into reactants.
When the rate of the forward reaction equals the rate of the reverse reaction, we say a state
of dynamic equilibrium has been reached.
Le Chatelier’s Principle
If the conditions of an equilibrium system are changed by changing temperature, pressure or
concentration, a process takes place which tends to oppose the effect of the change.
An equilibrium may be disturbed by changing any one (or more) of the factors for the
equilibrium:
o Temperature
o Concentration (gases and solutions)
o Pressure (gases only)
Changing Equilibrium Conditions
N2 (g) + 3H2 (g) ⇌ 2NH3 (g) (∆H < 0)
Concentration
An increase in concentration of any reactant will cause an increase in the reaction rate of the
forward reaction.
Increasing the concentration of the N2 or H2 would, therefore, increase the rate of the forward
reaction, hence favouring the shift of the equilibrium towards the forward reaction.
Temperature
An increase in temperature causes an increase in the rate of both reactions.
∆H refers to the forward reaction. If it is negative, the reaction is exothermic (energy is
liberated).
Increasing the temperature will favour the endothermic (reverse) reaction.
Pressure
Pressure can be increased by decreasing the volume of the container.
When the volume of the container decreases, the total concentration of all gases increases.
According to Le Chatelier’s, the reaction that will decrease the total number of gas moles in
the space will be favoured. (favours side with lowest number of gas moles)
So to ensure maximum yield of ammonia:
Use catalyst to reach equilibrium quickly.
Once equilibrium is reached:
o Drop temperature to 450 °C
o Increase conc. of N2 & H2& decrease conc. of NH3
o Increase pressure in container by reducing volume.
, Graph examples
Comparing reaction rates of the forward and reverse reactions
[A] (square brackets mean concentration of A) and [B] initially decrease.
[C] and [D] initially increase.
The rate of the forward reaction becomes constant and becomes equal to the rate of the
reverse reaction.
At this point (equilibrium) the concentrations of reactants and products remain constant.
Comparing concentrations of products and reactants graphically.
the [products] increases until equilibrium is
N2O4 (g)⇌ 2NO2(g) established and [products] remains constant
light brown dark brown
An increase in temperature
causes more NO2 to form.
This suggests that the forward
reaction must be endothermic.
(ΔH > 0)
Also increasing pressure
favours side with least gas
moles (reactants) therefore
reverse reaction (1 mole as
opposed to 2 moles of gas) is the [reactants] decreases until equilibrium is
favoured. established and [reactants] remains constant
28
Section B: Practice Questions
Question 1
(Taken from Preparatory Exam 2008)
William wants to determine the equilibrium constant for the decomposition of calcium carbonate
3
(CaCO3). He seals 2,0 g of CaCO3 in an evacuated 1,0 dm metal flask and connects a pressure
o
gauge to the flask. The flask is placed in an oven and heated to a temperature of 800 C at which
equilibrium was reached according to the following equation:
𝐶𝑎𝐶𝑂3 𝑠 ⇌ 𝐶𝑎𝑂 𝑠 + 𝐶𝑂2 𝑔 ∆𝐻 > 0
The graph obtained for pressure versus time for the decomposition of calcium carbonate is shown
below:
1.1. How does the rate of the reverser reaction change from to to t1? (2)
1.2. What is the reason for the horizontal line between t2 and t3? (1)
1.3. Draw a sketch graph to show how the mass of CaCO3 changes for the period to to t3. (4)
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