2.4 (a) Dynamic chemical equilibrium
Chemical reactions which take place in both directions
are called reversible reactions. Reactions like these however reach equilibrium.
This is when chemical reactions are still occurring but the rates of the
forward and reverse reactions are equal and so the concentration of each
species remains constant. This type of system is said to be in dynamic
equilibrium.
Example H2(g) + I2(g) <=>2HI(g)
(A dynamic equilibrium).
The proportion of products to reactants in the equilibrium
mixture is described as the equilibrium composition or as the position
of equilibrium:
aP + bQ <=> cR + dS
If the conversion of P and Q into R and S is small the
position of the equilibrium lies to the left. If the equilibrium
mixture is largely composed of R and S, the position of the equilibrium
lies to the right.
The position of equilibrium is dependant on the equilibrium
constant Kc and can also be altered by a change in external
conditions such as concentration or pressure. The equilibrium constant
is only affected by a change in temperature.
Task 2.4a With reference to the reaction between
hydrogen and iodine explain: (a) reversible reaction (b) equilibrium reaction
(c) dynamic equilibrium (d) position of equilibrium (e) shifting the position of
equilibrium to the right or left.
2.4 (b) Temperature, pressure and concentration and the
position of equilibrium
Le Chatelier's principle states that when a reaction at equilibrium is subjected
to a change in condition (temperature, pressure or concentration), the
equilibrium composition/position alters to reduce the effect of the
change. Homogeneous equilibria describe those in which all of the
substances are in the same phase e. g. all gases, all, liquids or all in aqueous
solution.
When there is a change in concentration
the
position of an equilibrium changes.
If a substance becomes more concentrated the position of the equilibrium
shifts to reduce the concentration of that substance. So for the reaction
below:
N2(g) + 3H2(g) <=> 2NH3(g) /\ H = -92kJmol-1
When the pressure is increased the equilibrium shifts
in the direction that tends to decrease the pressure. This is done by decreasing
the number of molecules present, by moving the position of equilibrium
from left to right. A higher yield of the product ammonia results
from the use of high pressure so typical plant operates at about 200 atmospheres
pressure.
For ammonia manufacture decreasing the temperature leads to a higher yield
of ammonia because the reaction from left to right is exothermic and causes
the temperature to rise again if ammonia is formed.
watch
animations of these effects
Task 2.4b State
and explain the effect of (a) increasing the temperature and (b) decreasing the
pressure (c) decreasing the temperature (d) increasing the pressure on the
following reactions:
2NO(g) + O2(g) = 2NO2(g) ; DH
= +57kJmol-1
2H2(g) + O2(g) = 2H2O(g) ; DH
= - 280kJmol-1
CH4(g) + H2O(g) = CO(g) + 3H2(g) ; DH
= + 40kJmol-1
2.4 (c) Catalyst and the position of equilibrium
Catalysts do not alter the equilibrium constant or the
position of equilibrium. They do affect the time needed for the system
to reach equilibrium.
2.4 (d) Temperature and optimum reaction conditions
In the Haber process a moderately high temperature
of around 500oC is used to speed the rate at which equilibrium
is reached. This temperature is chosen in spite of the fact that a lower
temperature gives a higher yield. The manufacture of
nitric acid involves 3 steps the first of which is the catalytic oxidation
of ammonia to nitrogen monoxide.