# Buffers Equilibrium

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#1
I think my concept of how equilibrium shifts is wrong when acid/base is added in small amounts to buffer solution.

Heres one of the questions I keep getting the equilibrium shifts wrong for:
Explain the buffering action of the system H2CO3 + H2O <==> HCO3- + H3O+ when H+ ions are added to the system.

I get that the H+ ions react with the HCO3- ions to form H2CO3 and that there are large amounts of acid and conjugate base to provide HCO3- ions to react with the H+ ions but what I don't understand is why equilibrium shifts left. And by reacting the H+ concentration does not change hence minimises the change in pH.

In my mind when equilibrium shifts left it is increasing the concentration of H2CO3 (is that wrong does it decrease the H2CO3 concentration) as otherwise there would be even more H2CO3 at equilibrium as well as the products from the H+ reacting with the HCO3- ions.

I am thinking that to compensate for the decrease in HCO3- ions equilibrium shifts right to increase the HCO3- concentration. (However this doesn't even make complete sense to me as H+ ions concentration would increase altering the pH and equilibrium will shift to restore Ka).

Why does equilibrium shift left and not right?
0
4 years ago
#2
You can also think of it as the H+ ions reacting with the water, H2O, in turn forming hydroxonium ions (H3O+). Consequently, the concentration of the products increase and thus the equilibrium shifts to the left to reduce the concentration of the hydroxonium ions. In turn, balancing out the changes in proton concentration and minimising changes in pH. -> this is what happens in reality i.e in our blood (we use a bicarbonate buffer).

This can be more easily deduced from the below dissociation of the weak acid H2CO3:

0
4 years ago
#3
(Original post by Liv_123)
I think my concept of how equilibrium shifts is wrong when acid/base is added in small amounts to buffer solution.

Heres one of the questions I keep getting the equilibrium shifts wrong for:
Explain the buffering action of the system H2CO3 + H2O <==> HCO3- + H3O+ when H+ ions are added to the system.

I get that the H+ ions react with the HCO3- ions to form H2CO3 and that there are large amounts of acid and conjugate base to provide HCO3- ions to react with the H+ ions but what I don't understand is why equilibrium shifts left. And by reacting the H+ concentration does not change hence minimises the change in pH.

In my mind when equilibrium shifts left it is increasing the concentration of H2CO3 (is that wrong does it decrease the H2CO3 concentration) as otherwise there would be even more H2CO3 at equilibrium as well as the products from the H+ reacting with the HCO3- ions.

I am thinking that to compensate for the decrease in HCO3- ions equilibrium shifts right to increase the HCO3- concentration. (However this doesn't even make complete sense to me as H+ ions concentration would increase altering the pH and equilibrium will shift to restore Ka).

Why does equilibrium shift left and not right?
Kozmo's answer explains the concept well from a different perspective.

In my mind, you are missing a vital piece of information:

"there are large amounts of acid and conjugate base to provide HCO3- ions to react with the H+ ions"

"I am thinking that to compensate for the decrease in HCO3- ions equilibrium shifts right to increase the HCO3- concentration"

. The Acid and CB in the equilibrium do provide a source of H+ and A-, correct. However these sources are very small as it's a weak acid.

In order for a buffer to properly 'mop-up' the H+, you need to add a salt of the conjugate base so there's an excess of the A- or lots available. In this case, you would add a slightly higher concentration of NaHCO3 (than H2CO3) which immediately dissociates to give you lots of HCO3- anions and some Na+ spectator counterions.

If you didn't have the Sodium bicarbonate 'stockpile', the equilibrium would go very acidic indeed because it's a weak acid that's present. So there's little dissociation of HCO3- ions originating from the H2CO3.

After you add the Sodium Bicarbonate, the equilibrium will immediately move to the left slightly but not very significantly (due to some HCO3- anions reacting with the present H+).

Now that you have a stockpile of HCO3- anions. The buffer equilibrium is ready to do battle!

The question is adding some H+, you now have plenty of HCO3- ions to shoot them down from your Sodium Bicarbonate!

Now you are reacting HCO3- anions so more H2CO3 will be produced, (concentration increases) and the equilibrium will move left. (Think of it like a see saw).

Now the goal is to minimise the changes in pH so if you dump loads of H+ in then obviously the pH will increase no matter how much the equilibrium attempts to react off the H+ ions!

The equilibrium constant (if you are asked about it), wouldn't change as it's a ratio I.E

4/8 = 2/4 = 1/2. You can change the numbers but they are the same ratio 0.5!

The only thing that changes Kc is temperature.

Revise Le Chatelier's Principle and the theory behind buffers.

This may help: http://www.chemguide.co.uk/physical/...elier.html#top

Also: http://www.chemguide.co.uk/physical/...a/buffers.html
0
#4
(Original post by _NMcC_)
Kozmo's answer explains the concept well from a different perspective.

In my mind, you are missing a vital piece of information:

"there are large amounts of acid and conjugate base to provide HCO3- ions to react with the H+ ions"

"I am thinking that to compensate for the decrease in HCO3- ions equilibrium shifts right to increase the HCO3- concentration"

. The Acid and CB in the equilibrium do provide a source of H+ and A-, correct. However these sources are very small as it's a weak acid.

In order for a buffer to properly 'mop-up' the H+, you need to add a salt of the conjugate base so there's an excess of the A- or lots available. In this case, you would add a slightly higher concentration of NaHCO3 (than H2CO3) which immediately dissociates to give you lots of HCO3- anions and some Na+ spectator counterions.

If you didn't have the Sodium bicarbonate 'stockpile', the equilibrium would go very acidic indeed because it's a weak acid that's present. So there's little dissociation of HCO3- ions originating from the H2CO3.

After you add the Sodium Bicarbonate, the equilibrium will immediately move to the left slightly but not very significantly (due to some HCO3- anions reacting with the present H+).

Now that you have a stockpile of HCO3- anions. The buffer equilibrium is ready to do battle!

The question is adding some H+, you now have plenty of HCO3- ions to shoot them down from your Sodium Bicarbonate!

Now you are reacting HCO3- anions so more H2CO3 will be produced, (concentration increases) and the equilibrium will move left. (Think of it like a see saw).

Now the goal is to minimise the changes in pH so if you dump loads of H+ in then obviously the pH will increase no matter how much the equilibrium attempts to react off the H+ ions!

The equilibrium constant (if you are asked about it), wouldn't change as it's a ratio I.E

4/8 = 2/4 = 1/2. You can change the numbers but they are the same ratio 0.5!

The only thing that changes Kc is temperature.

Revise Le Chatelier's Principle and the theory behind buffers.

This may help: http://www.chemguide.co.uk/physical/...elier.html#top

Also: http://www.chemguide.co.uk/physical/...a/buffers.html
Thank you so much makes complete sense now! Thank you again!
0
4 years ago
#5
(Original post by Liv_123)
Thank you so much makes complete sense now! Thank you again!
No problem, practice practice practice
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