HBr and C5H5N react in a 1:1 mole ratio:
C5H5N + HBr
C5H5NH+Br-
Since the molarities of both reactants is 0.200 M, then it will
take 25.0 mL of HBr to titrate the 25.0 mL of
C5H5N.
The Kb for C5H5N is 1.7 x
10-9.
Ka C5H5NH+ = (1 x
10-14) / Kb = (1 x 10-14) / (1.7 x
10-9) = 5.9 x 10-6
pKa = -log Ka = -log (5.9 x 10-6) = 5.23
a)
[OH-] = sqrt (Kb x [B]0) = sqrt ((1.7 x
10-9)(0.200)) = 2.0 x 10-5
pOH = -log [OH-] = -log (2.0 x 10-5) = 4.70
pH = 14.00 - pOH = 14.00 - 4.70 = 9.30
b)
(b) 12.5 mL represents the halfway point of the titration (i.e. the
equivalence point is at 25.0 mL HBr) where
[C5H5N] =
[C5H5NH+]. Using the
Henderson-Hasselbalch equation for buffers,
pH = pKa + log ([buffer base] / [buffer acid]) = 5.23 + log
([C5H5N] /
[C5H5NH+])
pH = 5.23 + log 1 = 5.23 + 0 = 5.23
(c) Since the equivalence point is at 25.0 mL HBr, at 22.0 mL HBr
you have converted 22.0 / 25.0 of the C5H5N
to C5H5NH+. So there are 22.0
parts C5H5NH+ and 3.0 parts
C5H5N. So pH can be calculated by using
Henderson-Hasselbalch equation .
pH = 5.23 + log (3.0 / 22.0) = 5.23 - 0.865 = 4.365
(d) We are now at the equivalence point. The only thing present
is the salt C5H5NH+Br-.
It acts as a weak acid like so:
C5H5NH+ + H2O
H3O+ + C5H5N
The pH of a weak acid solution may be calculated using
[H3O+] = sqrt ([HA] x Ka)
initial mmoles C5H5N = M
C5H5N x mL C5H5N =
(0.200)(25.0) = 5.00 mmoles C5H5N. Since all
of it was converted to C5H5NH+ in
the titration, then at the equivalence point there are 5.00 mmoles
of C5H5NH+.
[C5H5NH+] = 5.00 mmoles / 50.0 mL
= 0.1 M
[H3O+] = sqrt ((0.10)(5.9 x 10-6))
= 7.6 x 10-4
pH = -log [H3O+] = -log (7.6 x
10-4) = 1.88
lculate the pH for each of the following cases in the titration of 25.0 mL of...
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