First of all,
Le Chatelier's Principle states that
when a system at equilibrium is disturbed, the system will adjust itself to minimise the changes.
"Favours forward reaction" means that when equilibrium is reached, the concentration of products will be higher than the concentration of reagents.
"Favours reverse reaction" means that when equilibtium is reached, the concentration of reagents will be higher than the concentration of products.
Temperature
Say the reaction is A + B
C + D, and the forward reaction is exothermic.
This can be rewritten as A + B
C + D +
heat, i.e. heat is another product.
When the system is at equilibrium and the temperature is increased, heat is added. The equilibrium is hence disturbed and the only way to re-establish equilibrium conditions is to get rid of the heat.
You can see that C + D +
heat
A + B (same equation written backwards), which means to get rid of the excess heat, C will need to react with D and form A and B. This increases the concentration of A and B, and decreases the concentration of C and D - the equilibrium has shifted to the left.
Pressure
Say the reaction is 2E + F
4G, with all reagents and products being gaseous.
You should know that one mole of
any gas has a specific volume at standard conditions. This means regardless of which gases are present, the more moles of gas there are, the more volume they occupy.
On the left side of the equation there are 3 moles of gas (2E + F) and on the right side there are 4 (4G).
If we increase the pressure (i.e. decrease volume), the gases become a little more cramped. The equilibrium has been disturbed. To re-establish the equilibrium the gases will need to "uncramp" themselves.
Written backwards, 4G
2E + F. So 4 moles of gas can convert itself to 3 moles of gas if the reaction proceeds backwards. This reduces the effect the decrease in volume has upon the equilibrium. So the equilibrium shifts to the left.
I hope that helped a bit, it does take awhile to understand equilibrium but you'll definitely get it sooner or later.
