Acidity Constant + Equilibrium Shift Question

[DylanBurrowes]

Hey there,

I'm a student very much driven as to WHY certain things are - it makes things have a lot more practical sense and I can remember it 5000x easier.

So, two questions:

1) In the acidity and ionisation constants, how does the water just disappear from the equation (i.e. WHY does it count as "1")?
2) Why do equilibrium shifts actually occur due to Le Chatellier's principle? e.g. What makes a gaseous reaction favour a pressure decreasing reaction when the pressure is increased? Usually my teacher uses examples that give the chemicals 'emotions', such as "The gaseous solution doesn't like the added pressure so it wants to decrease it", but it is simply a devise used for explaining, not a scientifically-valid explanation.

If anyone would be able to help it would be so much appreciated! These two questions have been hanging over my head for the past three weeks!
Thanks!

[PB]

2) it is simply the system's most stable form. For example, if I squeezed a tennis ball i exert pressure on it, but once i release it, the ball will pop back into its normal spherical shape. The actual concept is a lot harder than this - which is the reason why Le Chatelier's principle was devised to simplify things. So yeh, just remember that when pressure is increase, the position of equilibirum would shift to the side with less molecules in order to reduce the pressure in the system.

[PB]

1) Firstly, [H2O] is NOT 1. it is a constant -to be precise (56M because in 1L of water there is 1000g of water, which is 56mols)   Therefore, if [H2O] is just going to be constant every single time, you can simplify the acidity constan by just removing it as it should have no effect on the equation.

[zvezda]

Hey there,

I'm a student very much driven as to WHY certain things are - it makes things have a lot more practical sense and I can remember it 5000x easier.

So, two questions:

1) In the acidity and ionisation constants, how does the water just disappear from the equation (i.e. WHY does it count as "1")?
2) Why do equilibrium shifts actually occur due to Le Chatellier's principle? e.g. What makes a gaseous reaction favour a pressure decreasing reaction when the pressure is increased? Usually my teacher uses examples that give the chemicals 'emotions', such as "The gaseous solution doesn't like the added pressure so it wants to decrease it", but it is simply a devise used for explaining, not a scientifically-valid explanation.

If anyone would be able to help it would be so much appreciated! These two questions have been hanging over my head for the past three weeks!
Thanks!

For your 2nd q:
Think about the mol ratio for a moment. The side of the equation with the most molecules per mol of equation will have a higher concentration as such (because your volume is constant). Therefore, the right side (for arguments sake) has a higher concentration at equilibrium. Now think about what happens when decreasing volume (increase in pressure and hence, concentration for both sides of the equation). But, this change will increase the concentration of the right side more than the left, meaning that its rate of reaction is increased more significantly. Given this, the reaction that begins from the right side will be favoured momentarily. Thus, a net back reaction will occur until equilibrium is once again established

[lzxnl]

Hey there,

I'm a student very much driven as to WHY certain things are - it makes things have a lot more practical sense and I can remember it 5000x easier.

So, two questions:

1) In the acidity and ionisation constants, how does the water just disappear from the equation (i.e. WHY does it count as "1")?
2) Why do equilibrium shifts actually occur due to Le Chatellier's principle? e.g. What makes a gaseous reaction favour a pressure decreasing reaction when the pressure is increased? Usually my teacher uses examples that give the chemicals 'emotions', such as "The gaseous solution doesn't like the added pressure so it wants to decrease it", but it is simply a devise used for explaining, not a scientifically-valid explanation.

If anyone would be able to help it would be so much appreciated! These two questions have been hanging over my head for the past three weeks!
Thanks!

For the first point, equilibrium constants are defined in terms of chemical activities, the log of which give the difference in chemical potential between a substance at a given time in existence and its potential in a standard state, usually at 298 K, pressures of 100 kPa and concentrations of 1 M. Now for pure water, the standard state is pure water. Its chemical potential is, therefore, generally pretty close to the potential to its standard state, so the difference is pretty much 0 and the activity is 1. Now, instead of having concentrations in equilibrium constants, we actually have chemical activities (read up on equilibrium constants on wiki if you want). If the activity of water is 1 and doesn't change much from it due to how much water there is in comparison to everything else, there's no need to include in in the equilibrium expression.

As for the second point, Le Chatelier's principle is really just a guide. Don't take it too literally. Let's just say that we have N2O4 and NO2 in equilibrium. If we increase the pressure, then the reaction quotient is affected. Q = p(NO2)^2/p(N2O4) if we take the reaction N2O4 <=> 2NO2 to be the forward reaction, and increasing the pressure means increasing the top more than the bottom; i.e. the reaction quotient is no longer equal to the equilibrium constant. When the reaction quotient equals the equilibrium constant, only then is the system in equilibrium again. This is literally where the system has its lowest potential to react; when the equilibrium constant and reaction quotients are equal.