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Dot point notes on Le Chatelier’s principle:
−↽− −⇀− represents equilibrium sign (which looks like this: )
2.2.3 Define Le Chatelier’s principle
The symbol for equilibrium is −↽−⇀− and simply means that, in a closed system, the rate of the forwards reaction is equal to the backwards reaction. This simple means that the reactants are converting to products at the same rate that the products are converting back into the reactants. Whilst there appears to be no change on a macroscopic level, the system is continually changing on a microscopic level. This process, known as dynamic equilibrium, results in the concentration of the substances in the system remaining constant.
According to Le Chatelier’s principle, if a system at equilibrium is disturbed, then the system will adjust itself in order to minimise the disturbance. However, note that the effects of the disturbance are never fully removed. They are only minimised, or lessened to a degree.
2.2.4 Identify factors which can affect the equilibrium in a reversible reaction
Le Chatelier’s principle is one which plays a crucial role in the HSC Chemistry course. Thus, a sound understanding of it is important, and it may appear again in this subject depending upon what Option you do. For this reason, a treatment sounder than required for this dotpoint will be provided.
Several factors can affect the equilibrium in a reversible reaction. These disturbances to the system can be in the form of changes in concentration, pressure, volume, or temperature.
Imagine a system in equilibrium of four compounds, A, B, C, and D. A + B −↽− −⇀− C + D
The simplest way of visualising changes in concentration is simply seeing Le Chatelier’s principle as working to minimise any changes made to the equilibrium. As more of A or B is added, then the system will try to minimise the change by converting more A and B into C and D. As such, the equilibrium shifts to the right.
Conversely, if more of C or D is added, increasing the concentration of the products, then the system will convert more C and D into A and B, shifting the equilibrium to the left.
Note that a system can only minimise a disturbance. It cannot completely undo it.
Imagine a system in equilibrium of four compounds, A, B, C, and D. Unlike the example used to illustrate changes in concentration, the four compounds in this example are gases, and the number of moles of A is two rather than one.
2 A(g) + B(g) −↽−⇀− C(g) + D(g)
Determining the affect of changes in the pressure of a system is simply an exercise in counting moles of gases. In the equilibrium above, there are three moles of gas on the left side, and 2 moles of gas on the right. Any increase in pressure will result in the system trying to relieve the pressure by ‘leveling’ the moles of gas within the system. As such, in the above system, an increase in pressure will lead to a shift in the equilibrium to the right. This occurs simply because the system is essentially counteracting the fact that three moles of gas are becoming two moles of gas.
Conversely, a decrease in pressure will shift the above equilibrium to the left in an attempt to increase pressure once again.
Changes in pressure affect only gases. Increasing the pressure in the following system will lead to equilibrium shifting to the right, as there are two moles of gas on the left side and only one on the right.
A(g) + B(g) −↽−⇀− C(g) + D(s)
Any change in volume in a gaseous equilibrium is simply a change in pressure. As such, treat increases in volume as decreases in pressure, as there are more moles of gas in the fixed space, and treat decreases in volume as increases in pressure.
The effect of Le Chatelier’s principle with changes in temperature can often be confusing. However, simply thinking of heat as either a product or reactant greatly simplifies any problems, as shown in the equilibrium below, where the reaction is endothermic (Absorbs heat in order for the reaction to occur) rather than exothermic (Releases heat).
A + B + H e a t −↽− −⇀− C + D
In the above endothermic equilibrium, an increase in temperature will result in the system working to reduce the temperature by shifting the equilibrium to the right, converting A and B into C and D in order to reduce temperature.
Conversely, a decrease in temperature will shift the equilibrium to the left, converting C and D into A and B in order to produce more heat.
In the case of an exothermic reaction, the equation will be of the form A + B −↽− −⇀− C + D + H e a t
As shown above, treating heat energy as an actual item in the equilibrium is a much simpler method of thinking of a problem. Simply determine whether a reaction is exothermic forwards, i.e. the heat is placed on the right, or endothermic forwards, i.e. the heat is placed on the left.
Remember- Changes in concentration, pressure, volume and temperature will all disturb a system in equilibrium.
From the Student’s Guide to HSC Chemistry. Licensed for free distribution under the GFDL.
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