Le Chatlier\'s principle is stated as follows: \" \"if a change is imposed on a
ID: 628373 • Letter: L
Question
Le Chatlier's principle is stated as follows: " "if a change is imposed on a system at equilibrium, the position of the equilibrium will shift in a direction that tends to reduce that change." The system N2 + 3H2 <-> 2NH3 is used as an example in which the addition of nitrogen gas at equilibrium results in a decrease in H2 concentration and an increase in NH3 concentration. In the experiment the volume is assumed to be constant. On the other hand, if N2 is added to the reaction system in a container with a piston so that the pressure can be held constant, the amount of NH3 actually could decrease and the concentration of H2 would increase as equilibrium is reestablished. Explain how this can happen. Also, if you consider this same system at equilibrium, the addition of an inert gas, holding the pressure constant does affect the equilibrium position. Explain why the addition of an inert gas to this system in a rigid container does not affect the equilibrium position. Thank you!Explanation / Answer
Le Chateliers Principle Le Chateliers principle states that any change in a system at equilibrium results in a shift of the equilibrium in the direction which minimises the change.
The law was originally applied only to pressure and was first published in a note as: Any system in stable chemical equilibrium, subjected to the influence of an external cause which tends to change either its temperature or its condensation (pressure, concentration, number of molecules in unit volume), either as a whole or in some of its parts, can only undergo such internal modifications as would, if produced alone, bring about a change of temperature or of condensation of opposite sign to that resulting from the external cause.
A good example of the law is the Haber Process Chemical Equation, a reaction that Le Chatelier worked on but subsequently abandoned. The project was continued by Haber and Claude who successfully produced Ammonia on a commercial scale using this process.
N2 (g) + 3H2 (g) <===> 2NH3 (g) + heat Increasing
H2 in system shifts equilibrium to right Decrease
NH3 to shifts equilibrium to the right
Increase Heat to shifts equilibrium to the left
1. I always got confused with temperature when we started this concept. I found it easier to add it to the equation. If it is exothermic you can just add the absolute value to the products of the forward reaction. Like so:
N2 (g) + 3H2 (g) --> 2NH3(g) + 92kJ
So if you increase the temperature, you increase the energy. The system will partially oppose the change, favouring the rate of the forward reaction. You would observe (smell is usually included), a system that becomes more colourless/faded, and a more pungent smell.
2. That would increase the amount of N2. The system would partially oppose the change, favouring the rate of the reverse reaction. The system becomes less pungent and more yellow.
3. Okay this is the same as increasing the pressure if you can picture it. In this case you have to look at all states of matter. All are gases which means this will make a difference (this has little to no difference on solids and liquids so they are usually ignored). On the left we have 4 moles of gases, and on the right we only have two. Increasing the pressure would therefore favour the left, and the system will partially oppose the change, shifting the equilibrium to the left. The observations would be the same as in number 2.
4. Catalysts have NO effect on the equilibrium
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