- Created by: Sophie
- Created on: 02-06-10 17:59
1. Dynamic equilibria
Consider a reversible reaction A + B == C + D
As the reaction proceeds, the rate of the forward reaction decreases and the rate of the reverse reaction increases. Eventually, the reaction will reach a stage where both forward and backward reactions are proceeding at the same rate:
At this stage, a dynamic equilibrium has been reached. “Dynamic” means that the reaction has not stopped; it is simply moving in both directions at the same rate. “Equilibrium” means that the amount of reactants and products in the system is staying the same.This is in contrast to a static equilibrium, in which there is no movement in either direction. Chemical equilibria are in general dynamic rather than static.
All reactions are reversible in theory; although in practice many are not considered to be so. In some cases the reverse reaction is insignificant; in others, it is not allowed to take place.
2. Open and closed systems
A closed system is one from which reactants and products cannot escape. In closed systems the forward and reverse reactions continue until dynamic equilibrium is reached. All reactions in a closed system are thus reversible in theory, although they are only considered as such if both forward and reverse reactions occur to a significant extent.
Eg H+(aq) + OH-(aq) à H2O(l)
In this case the reverse reaction is not significant so the reaction is represented by single arrow.
Eg H2(g) + I2(g) 2HI(g)
In this case the reverse reaction is significant, so the reaction is represented by an equilibrium sign:
An open system is one from which reactants and products can escape. The open air or a fume cupboard is an example of an open system. In an open system, the products are removed as soon as they are formed, so the reverse reaction is not allowed to take place. Such reactions clearly never reach equilibrium, but proceed until all the reactions have been converted into products. Reactions which proceed under these conditions are clearly irreversible.
Eg H2O(l) à H2O(g)
This reaction would not be expected to proceed significantly under normal conditions, since water is more stable than steam at normal temperatures. However puddles will disappear completely if left for long enough. This is because the water vapour is removed by wind currents as soon as it is produced, and so the reverse reaction is not allowed to take place. Thus the system never reaches equilibrium and the reaction is irreversible.
Thus if a reaction is represented by an equilibrium sign, it is often assumed that:
- the system is closed
- the reverse reaction is significant
LE CHATELIER'S PRINCIPLE
If the conditions are changed after equilibrium has been established, the rates of the forward and reverse reactions may be affected by different amounts and the system may no longer be at equilibrium. If this is the case, it will move in one direction or…