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When a gaseous system in equilibrium is compressed or expanded, this may have
an effect on the position of the equilibrium. Consider the simple gas phase
reaction
- N2O4(g) < = > 2NO2(g)
If
this system is in equilibrium and we compress it, we will move the equilibrium
point. According to LeChatlier's
Principle, the system will move so as to partially compensate for the
change. Since the pressure has increase, the reaction will move in the direction
which will lower the pressure: since there are fewer molecules on the left than
on the right, the reaction will move to the left.
In general, the following rules apply for gaseous reactions
- When a equilibrium system is compressed, a reaction will occur in the
direction that decreases the total number of moles of gas
- When a equilibrium system is expanded, a reaction will occur in the
direction that increases` the total number of moles of gas
- If both sides of the reaction have the same number of moles of gas,
expansion or compression will not affect the position of the equilibrium.
Example: If you compress the following systems when they are at
equilibrium, what will happen?
- 2CO2(g) < = > 2CO(g) + O2(g)
- NO(g) + CO2(g) < = > NO2(g) +
CO(g)
Solution:Since we are compressing the systems, they will move in the
direction with fewer moles of gas.
- There are 2 moles on the left of reaction 1, 3 on the right. The
reaction will move to the left.
- There are 2 moles on the left of reaction 2 and 2 on the right. There
wil be no effect when the pressure changes
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