# Chemical Equilibria

Chemical equilibria refers to dynamic equilibria. This means that both the forwards and reverse reactions
are still occurring, but they are occurring at the same rate, so there is no net change in the concentrations
of reactants or products. For a reaction to reach dynamic equilibrium, it must be in a closed system.
Le Chateliers Principle;
If you change the conditions of a reactions equilibrium, the equilibrium will shift to oppose the change
 Temperature -> Increase favours endothermic, Decrease favours exothermic
 Pressure -> Increase favours side with fewest moles, Decrease favours side with most moles
 Removing a chemical = equilibrium shifts to produce more
 Adding a chemical = equilibrium shifts to remove the excess
State that the shift in equilibrium is to oppose the change, and then the change that is being opposed
(i.e. an increase in temperature, decrease in pressure etc….)
Equilibrium Constant Kc;
Kc is the products, raised to their coefficients, multiplied
together divided by the reactants, raised to their coefficients,
multiplied together.
Kc only works with homogenous reactions, where all the
reactants and products are in the same state.
[ ] = concentration in mol/dm3
. You must calculate this before
working out Kc.

If the volume has not been given to you, just use x as the volume to find concentration. The x will cancel
out in the equation so you do not require the volume.
You may have to use an ICE table to calculate the equilibrium moles of reactants/products
Units of Kc:
To calculate the units of Kc, you write out the units of all terms in the Kc expression, then cancel out as
many of these units as possible and simplify to find the units of Kc
The only factors that affects Kc is temperature. Pressure will temporarily alter Kc, but eventually the
proportions of reactants : products will return to normal, so pressure doesn’t affect Kc. Temperature
however will permanently affect the value of Kc
Equilibrium Constant Kp:
This is similar to Kc, but is calculated with partial pressures rather than concentrations;
 Find the equilibrium moles from the ICE table
 Find mole fractions (moles of reactant / total # moles)
 Find the partial pressures (mole fraction * total pressure)
 Use partial pressures in Kp expression, raising each Partial Pressure to the power of its coefficient in
the given equation
 Find the units of Kp in the same way as in Kc (kPa/Pa/atm rather than mol/dm3
)
As we are not using concentrations, you should never use [ ] in the expression for Kp. Again, Kp is only
affected by temperature, exactly like Kc