# Chemistry

Given a diprotic acid, H2A, with two ionization constants of Ka1=3.9E-2 and Ka2=5.9E-8.

Calculate the pH and molar concentration of each protonated form for a 0.135 M solution of:
(a) H2A
(b) NaHA
(c) Na2A

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1. For H2A it follows the usual form of ionizsation as if you had a monoprotic acid.
..........H2A ==> H^+ + HA^-
I.......0.135M....0......0
C.........-x......x......x
E........0.135-x..x.......x

Ka1 = (H^+)(HA^-)/(H2A)
Substitute and solve for x = (H^+) = (HA^-) and convert to pH. (H2A) = 0.135-x. I strongly suspect you will need t use the quadratic equation to solve this because ka1 is relatively large.

For NaHA.
(H^+) = sqrt [(k2(HA^-)+Kw)/(1+(HA^-)/k1)] and convert to pH.

For Na2A, set up an ICE chart for the hydrolysis of A^2-.

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2. I'm not getting the right answers for the pH, [H2A], and [HA^-] in part a (0.135 M solution for H2A)

For H2A, I did
3.9E-2=x^2/(.135-x)

x^2+.039x-.005265=0
x=.06725
[H^+]=[HA^-]=.06725 M

pH=-log(.06725)=1.17

Finding [H2A]=.135 M-.06725 M=.06775 M

Finding [A^2-]=Ka2*[HA^-]/[H^]=5.9E-8 M

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