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Equilibrium.notebook
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May 31, 2019
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Le Chatelier's Principle at system at equilibrium when subjected to a change will respond in such a way as to minimize the effect of the change.
More simply put whatever we do to a system at equilibrium, the system will respond in the opposite way. The balance between the forward and reverse reaction will shift to offset any change and return the system to equilibrium.
The most common changes we can make are to change temperature, pressure, concentration or to add a catalyst.
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Effect of changing concentrationN2(g) + 3 H2(g) 2 NH3(g)
If we were to add more N2(g) to this reaction after it reached equilibrium, the reaction must USE UP the additional N2(g).
To do this, the excess N2(g) reacts with any remaining H2(g) to form more product! The reaction SHIFTS TO THE RIGHT, products are favoured.
WHAT DO YOU THINK HAPPENS IF WE REMOVE NH3(g) AS IT FORMS?
May 281:40 PM
Effect of changing temperature
This depends on whether a reaction is exothermic or endothermic. Generally speaking, we are either trying to use up extra heat added, or replace heat we have removed.
N2(g) + 3H2(g) 2 NH3(g) H = 92 kJ
What happens if we heat the reaction mixture? Cool it?
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Effect of changing pressure (applies to equilibrium systems included gases only)
4 HCl(g) + O2(g) 2 H2O(g) + 2 Cl2(g)We compare TOTAL MOLES of gaseous substances on each side. Increasing pressure will cause the equilibrium to shift to the side with FEWER moles of gas (as this will decrease total pressure, balancing our change). Decreasing pressure will cause the opposite!
In this case which way does equilibrium shift if we INCREASE pressure?
May 281:51 PM
Effect of introducing a catalyst
Catalysts give an alternate pathway for a reversible reaction to happen by LOWERING activation energy. However, the activation energy for the forward AND reverse reaction is changed by the same amount, meaning the rate of both forward and reverse reaction is changed by the same factor. THIS MEANS ADDING A CATALYST DOES NOT CAUSE A SHIFT IN EQUILIBRIUM OR CHANGE KC.
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Applications of Equilibrium in Industry HABER PROCESS
As chemists, we can use what we know about equilibrium and chemical kinetics to both increase the YIELD of our desired products and increase the RATE at which they are formed to maximize our economic gains. If rate and equilibrium work against each other, sometimes a compromise must be made.
Fritz Haber was a German chemist who figured out how to synthesize ammonia from its elements on a large scale. His work was initially related to a shortage of fertilizers, but eventually had a significant impact on the first world war as his process was used to manufacture ammonia for explosives.
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Haber Process Continued
N2(g) + 3 H2(g) 2 NH3(g) H = 92 kJ
How did Haber maximize production of ammonia? Taking advantage of Le Chatelier's principle.
What happens if we remove NH3 as it is produced?
What happens if we increase the pressure during the reaction?
What happens if we use a lower temperature? (Not too low or reaction gets too slow, since temperature also impacts rate?)
What happens if we add a catalyst?
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