An unknown weak acid, ha, it titrated with 1.2 m naoh. the ph at the halfway point of this titration was found to be 4.102. if the initial ph of the weak acid solution (before titration) has a ph of 2.308, what was the concentration of the weak acid solution?

0.311 M

Explanation:

HA, our weak acid, reacts with NaOH by the following equation given below:

This may also be represented by the net ionic equation for this acid-base reaction:

During the titration up to the equivalence point, we would have some HA remaining and some produced. That said, we have a weak acid and its conjugate base forming a buffer. We may apply the Henderson-Hasselbach equation for buffers here:

At midpoint, half of HA reacts to produce half of , meaning when half of equivalence volume of NaOH is added.

This implies that:

at midpoint. We may then find the acid ionization constant for this acid:

The initial molarity has to be found now. We know that the initial pH is:

Since , then

Using an ICE table, for some initial molarity of acid, , we may write the equilibrium constant expression as:

Solve for the initial molarity of HA:

in the given question, the partial pressure is calculated by using the ideal gas equation, that is, pv = nrt

here, p is pressure, v is volume = 25 l, n is number of moles = 0.0104 moles, r is gas constant = -0.0821 l.atm / mol.k

t is temperature = 273 k

p = nrt / v

p = (0.0104 mol × 0.0821 l.atm / mol.k × 273 k) / 25 l

= 0.0093 atm