If a chemical system at equilibrium experiences a change in concentration, temperature, pressure or volume, the equilibrium will be shifted to counteract this change, and a new equilibrium will be established.Now, let's break this statement down:
Change in Concentration
A + B ⇌ C + D
In the above reaction, if the concentration of C were increased, then the equilibrium must shift to counteract this change. In order to do so, the equilibrium will have to shift toward the reverse direction so that the excess C can be consumed.
If, instead, the concentration of C were decreased, then the opposite will happen. The reaction will have to shift toward the forward direction so that there can be more C.
Change in Temperature
A + B ⇌ C + D ΔH=-46KJ
In the above reaction, you can see that the reaction is exothermic because the ΔH is negative. In other words, the forward reaction releases heat.So you can sort of treat heat (energy) as a product, and rewrite the equation like this:
A + B ⇌ C + D + energy
Then, everything is simple. If you were to increase the temperature, you'd be adding more of the product into the equilibrium, and therefore the equilibrium will have to shift toward the left to consume this product. If you were to decrease the temperature, you'd be removing some of the product, and the equilibrium will have to shift toward the right to make up for this removal of product.
Change in Partial Pressure
Note: A change in pressure will only have an effect on reactions that involve gas molecules
2A(g) + B(g) ⇌ C(g)
In the above reaction, you can see that there are three moles of gas on the left side of the equation, and one on the right side.
If the pressure is increased, then the volume will decrease. As a result, the equilibrium will have to shift toward the right because there are fewer moles of gas on the right side (and thus it saves space).
Similarly, if the pressure is decreased, then the volume will increase. The equilibrium will have to shift to the left in order for all the "extra space" to be occupied.