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How Does Pka Relate To Ka

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Why does the Ka value change for polyprotic acids?
I was wondering how and why the value for Ka of the first hydronium ion in a polyprotic acid differs from the Ka value of the second, and so on. Wikipedia is erroneous in this respect! Please only give an answer if you are sure :)

to understand this you have to relate ka or pka to stability. the stronger acid has a more stable conjugate base. as the acid gives up more H+'s, it becomes more -. putting all that negative charge means it becomes less stable.

How do you find the half equivalence point on a titration curve?
For lab, I have already found the equivalence point on the titration curve I've constructed from my data, but for my report, I have to show a half equivalence point. Can anyone tell me how to find this? Can someone also explain how the half equivalence point relates to anything for finding Ka or Kb? Thanks so much in advance!

The half equivalence point occurs when [HA]=[A-] during the buffer region of your titration curve. At the 1/2 eq. pt the Pka=pH of the solution, and using pKa= -log [ka], using antilog can give you the ka via ka=10^-pKa.

What is the relationship between pKa and pH?
Please no one tell me what they both are, since I know very well. Instead I'd like to know (in simple English terms) how the terms are related. What are both of them used for in relationship?

Suppose you have a weak acid, HA. In solution, it will ionize slightly by the reaction: HA + H2O <---> H3O+ + A- The equilibrium constant for this reaction, Ka, is equal to: Ka = [H3O+][A-]/[HA] This constant is unique and different for every weak acid. Now, pKa is the -log of Ka. pH is the - log [H3O+] in a solution. By doing a little algebra with the equation for Ka, you can show that: pH = pKa + log [A-]/[HA] This is the Henderson-Hasselbalch equation, and is used to relate the pH of a buffer solution to the pKa and the concentrations of the weak acid and its conjugate base.

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