i don't understand buffers (1 Viewer)

[.ben]

Here's teh example

HA + H₂O <--> H₃O⁺ + A^-

-- If an acid is added, the hydronium ions react with A^- therefore equilibrium shifts to the right to replace A^-

--If a base is added, the hydroxide ions react with H₃O⁺ and thus equilibrium shifts to right to replace the H₃O⁺

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What's wrong with these explanations of buffers?

 

[Riviet]

In adding an acid, you are increasing the [H₃0⁺]. By Le Chatelier's principle, the base A^- will react with the extra H₃0 to form HA so that the change in [H₃0⁺] is minimised. Therefore the equilibrium shifts to the left.

Your explanation for adding a base is correct. Hope that helps.

 

[.ben]

Why is adding acid increasing acid while addin base is decreasing acid?

 

[Riviet]

That is something that I'm unsure about... just remember it as the concentration of the hydronium changing when adding an acid/base.

 

[.ben]

ok, well thanks anyway

 

[tristambrown]

it';s because the buffer reacts with whatever you are adding to the solution

eg bicarboante can function as an weak acid (see it donating a proton to the water as it moves to the right hand side [note it can also accept protons and act as a base, but we will ignore it for this example]). Carbonate functions as a base

Bicarbonate <---> Carbonate
HCO3^- <----> H+ CO3^2-

these two establish an equilibruim

if you add a strong acid like HCL it will react with carbonate and disturb the equilibrium.

CO3^2- + 2HCL --> H2CO3 + Cl^- ions (this is carbonic acid, a MUCH weaker acid than HCL as it remains mostly unionised in it's aqueous state, hence the ph does not fall very much from the addition of small amts of HCL)

as a result of the carbonate being used up neutralising the HCL more of the bicarbonate will transform into carbonate (al la le chatelliers [spelling?] principle) which in turn will continue to react with the remaining HCL, neutralising the strong acid hence buffering the solution- that is - keeping the ph relatively constant.

If too much HCL is added however the carbonate, being constantly removed by neutralisation of the acid, will be constantly formed until all the bicarbonate it is being formed from is gone. At this point the PH will spike and with no more buffer to hold the ph stable the HCL will do it's job as it normally would

this system works in reverse for bases, neutralising the bases with the weak acid bicarbonate. (note this definition as "acid" is only applicable when using a B& L definition)