Dissociation

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If pH is > pKa, then acid is deprotonated.

1. What about pKb?

2. What is the proof for that statement? Meaning why is acid deprotonated when pH > pKa?

I assume pH > pKa means the solution is more basic than the solute so, in an attempt to equilibriate, the solute will protonate all the base around it, thus becoming deprotonated.

If pH > pKb? I think you would have to calculate the pKa and see what the relationship is between pH and pKa.
 
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The reaction is:

HA <-> H+ + A-

pKa is -log[H+] for the concentration of the acid when the acid/base equilibrium is established. if pH > pKa, the concentration of [H+] is lower than the equilibrium concentration and the reaction will move to the right.

pKb applies to the conjugate base - if pOH > pKb the base will be protonating since [OH] is too low.

Added: If you want a more precies proof, you should consider that Ka=[H][A]/[HA] at equilibrium. But for [HA]>>[H] we have that [A]/[HA] = 1, and Ka = [H] at equilibrium or pKa = pH at equilibrium.

The interesting assumption here is that [HA] >> [H], which is correct most of the time. It is still means that at the extreme cases the above will be incorrect.
 
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Cool thanks for the helpful responses. Well understood now!

This chapter is taking way too long...
 
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