a. A 0.037 M solution of a weak acid is 4.67 percent ionized in
solution. What is the Ka for this acid?
b. The pOH of a 0.29 M weak acid solution is 9.7. What is the Ka
for this weak acid?
consider a 1.00 M solution of a weak acid, HA. The pH of the
solution is found to be 3.85. A) calculate the [H3O+] in the
solution. This would be the equilibrium concentration of H3O+ in
the solution. B) write out an ICE table as before. Here, we don’t
know the numerical value of Ka but we know the [H3O+] at
equilibrium which you should see from your ICE table easily relates
to the value of “x” in your table...
A weak acid HA (pKa = 5.00) was titrated with 1.10 M KOH. The
acid solution had a volume of 100.0 mL and a molarity of 0.107 M.
Find the pH at the following volumes of base added: Vb = 0.00,
1.00, 5.00, 9.00, 9.90, 10.00, 10.10, and 12.00 mL. (Assume Kw =
1.01 ✕ 10−14.)
A weak acid (HA) has a pKa of 4.009. If a solution of this acid
has a pH of 4.024, what percentage of the acid is not ionized?
(Assume all H in solution came from the ionization of HA.)
A weak acid (HA) has a pKa of 4.666. If a solution of this acid
has a pH of 4.294, what percentage of the acid is not ionized?
(Assume all H in solution came from the ionization of HA.)
Please help me out...I am seriously not sure where to begin. I
converted pka to Ka but then I'm not sure where to go from there or
if I am on the right track. Any detailed help of how to solve...
What is the pH of a buffer solution that is composed
of a weak acid, HA (Ka=8.02×10–9), and the
conjugate base, A–, after 3.24 mL of 0.089 M HCl
solution is added. The initial concentrations of the 142 mL buffer
solution are [HA]=0.58 M and [A–]=0.61 M. Enter your
value to two (2) decimal places.
A 37.00 mL solution of 0.390 M weak acid Ha(aq) (Ka=2.25x10^-5)
is titrate with 0.390 M NaOH(aq). Calculate the pH of the solution
14.00 mL past the end point.