# Le Chatlier's principle with a gas

• nobahar
In summary: Pi?In summary, if more PCl5 is added to a sealed container, the equilibrium shifts to the right. However, this would increase pressure.
nobahar
For:
PCl5 <-----> PCl3 + Cl2
in a sealed container, if I added more PCl5 the equilibrium should shift to the right, but this would increase the pressure. Does this mean that the quilibrium will also have a factor attempting to shift it to the left (because there's less moles on the left)? Or do I simply disregard this factor?
I'm not sure if I've understtod LCP correctly...
Any help appreciated, thanks in advance.

nobahar said:
For:
PCl5 <-----> PCl3 + Cl2
in a sealed container, if I added more PCl5 the equilibrium should shift to the right
Are you sure? The concentration of PCl5 remains unchanged, as it is a solid.

Oh...
Sorry, in the question booklet its quoted as a gas. If it was a gas, what would happen then?

All of the concentrations would change because the volume would change (assuming constant pressure). If pressure were allowed to change and the volume remained constant, you would only change the concentration of the PCl5. How would changing just the concentration of PCl5 affect the equilibrium? What do you think would happen if this were a constant pressure experiment?

According to Le Chatlier's principle, adding more PCl5 would shift the equilibrium to the right, this produces more moles and, in a closed (or isolated?) system allowing pressure to change, this would increase the pressure. Initially I thought that this is pushing the pressure value away from its equilibrium value, but it would seem that a different equilibrium has been established with a different pressure associated with it. Is this the case?

Concentration is the thing. Pressure doesn't figure into it unless the moles of something increases or decreases. What would happen to the equilibrium if you changed the pressure by injecting an inert gas? Would that alone change the concentration of the other components (assuming constant volume)?

If I injected an inert gas, it would push the equilibrium to the side with less moles in order to reduce the pressure, as the pressure of the system would increase. The concentrations would indeed alter.
If I introduced more of a reactant in a fixed volume, it's concentration would - even if temporarily - increase. Again, this would increase the overall pressure of the system. Given that pressure = nRT/V. All moles, whether partcipating in the reaction or not, contribute to the pressure. Since adding more moles of a recatant increases the pressure, I wondered if there would be an 'attempt' to counter the change in pressure.

nobahar said:
If I injected an inert gas, it would push the equilibrium to the side with less moles in order to reduce the pressure, as the pressure of the system would increase. The concentrations would indeed alter.
Not quite. Remember that concentration is moles/volume. Notice that pressure doesn't figure into that. If you add or subtract moles (and pressure changes as a result) and the volume remains constant, you will affect the equilibrium. Not so if you inject an inert gas.

If I introduced more of a reactant in a fixed volume, it's concentration would - even if temporarily - increase. Again, this would increase the overall pressure of the system. Given that pressure = nRT/V.
It is better to write this as Pi = niRT where "Pi" is a the partial pressure of species "i" and ni is the moles of "i".

nobahar said:
According to Le Chatlier's principle, adding more PCl5 would shift the equilibrium to the right, this produces more moles and, in a closed (or isolated?) system allowing pressure to change, this would increase the pressure. Initially I thought that this is pushing the pressure value away from its equilibrium value, but it would seem that a different equilibrium has been established with a different pressure associated with it. Is this the case?

BTW, all of this is correct except you should think of the equilibrium in terms of concentration rather than pressure since your understanding of what is happening will be more general then.

The overall pressure is the sum of the individual pressures. An inert gas under these circumstances would increase the pressure of the system, this was the reasoning I was working on; but since this does not appear to effect the equilibirum, then it is not the overall pressure which is important.
For a fixed volume, the individual concentrations will be constant. If I alter the volume I alter both the concentration and the pressure (overall pressure and individual pressures). If altering the overall pressure by injecting an inert gas does not effect the equilibrium, is it then that I should only concern myself with concentration in all circumstances? That in a reaction which involves gas the alteration of concentration just so happens to be coupled with an alteration in pressure (due to the volume change), and that this, as a consequence of alteration in concentration, can be used to work out what will happen to the equilibirum?
I think this is what you meant by the last post, I hope my attempts to understand are clearly put above. Thanks for all the help so far.

nobahar said:
The overall pressure is the sum of the individual pressures. An inert gas under these circumstances would increase the pressure of the system, this was the reasoning I was working on; but since this does not appear to effect the equilibirum, then it is not the overall pressure which is important.
Absolutely correct.

For a fixed volume, the individual concentrations will be constant. If I alter the volume I alter both the concentration and the pressure (overall pressure and individual pressures). If altering the overall pressure by injecting an inert gas does not effect the equilibrium, is it then that I should only concern myself with concentration in all circumstances?
Yes.

That in a reaction which involves gas the alteration of concentration just so happens to be coupled with an alteration in pressure (due to the volume change), and that this, as a consequence of alteration in concentration, can be used to work out what will happen to the equilibirum?
Yes, but you will need to do a complete analysis for each component in the mixture that shows up in the equilibrium expression ignoring the effect of any inert gases.

Many many thanks for the help chemistree. Your patience and time is much appreciated, I now have a far better understanding of what is happening. Thankyou

## 1. What is Le Chatelier's principle with a gas?

Le Chatelier's principle with a gas is a scientific principle that describes how a system at equilibrium will respond when subjected to a change in temperature, pressure, or concentration of a gas. It states that when a change is imposed on a system, the system will adjust in a way that minimizes the effect of the change and restores equilibrium.

## 2. How does Le Chatelier's principle apply to gases?

Le Chatelier's principle applies to gases by predicting how they will respond to changes in their environment. For example, if pressure is increased, the system will shift to decrease the number of gas molecules, and if pressure is decreased, the system will shift to increase the number of gas molecules.

## 3. What factors can affect the equilibrium of a gas according to Le Chatelier's principle?

Temperature, pressure, and concentration of a gas can all affect the equilibrium of a gas according to Le Chatelier's principle. Changes in these factors can cause the system to shift to restore equilibrium.

## 4. How can Le Chatelier's principle be used to predict the direction of a gas reaction?

Le Chatelier's principle can be used to predict the direction of a gas reaction by considering how changes in temperature, pressure, or concentration of a gas will shift the equilibrium. For example, if the concentration of a gas is increased, the reaction will shift to decrease the concentration of that gas.

## 5. Does Le Chatelier's principle only apply to gases?

No, Le Chatelier's principle can apply to any system at equilibrium, including solutions, solids, and gases. However, it is most commonly applied to gases due to their ability to easily expand and contract in response to changes in temperature and pressure.

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