Students learn to interpret the equilibrium
constant expression (no units required) from the chemical equations
of equilibrium reactions.
Part 2: Just to further illustrate these points and introduce the equilibrium
constant, here are some more examples.
There exists an equilibrium reaction for the production of ammonia
from nitrogen and hydrogen gas:
Le Châtelier's Principle holds here as well.
If we were to add ammonia to this system once it was already
at equilibrium, the system would move to produce more of the
reactants.
Equilibrium Constant:
Just to revise where the expression for the equilibrium constant,
K, comes from, let's consider this general equation for an
equilibrium reaction of a moles of compound A with b moles of
compound B to yield c moles of compound C and d moles of
compound D:
The equilibrium constant is equal to the ratio of concentrations, raised
to the power of their stiochiometric coefficient, between products
and reactants. This value will be constant (for a given fixed temperatue)
and can be represented by the following equation:
So for the previous example of ammonia formation the expression is
given as:
So looking at the equation you can see if more ammonia were added to the system
the value of the ratio would increase. So to return the value to the original
value of K, the system must convert ammonia into nitrogen and hydrogen.
