Given a PH of buffer solution is 4.23 made from 0.35 formic acid. Calculate the concentration of formate.The constant dissociation for formic acid is 3.75
Use the Henderson-Hasselbalch equation.
4.23 = 3.75 + log(base)/(0.35) and solve for (base) = (formate) but before you do any of that you need to determine if that 0.235 is M, m, g/mL or just what. You should be M. Also, I assume that 3.75 is the pKa but you say in the post it is Ka.
To calculate the concentration of formate in a buffer solution with a known pH and concentration of formic acid, you can use the Henderson-Hasselbalch equation.
The Henderson-Hasselbalch equation is given as:
pH = pKa + log([A-] / [HA])
Where:
pH = the pH of the buffer solution
pKa = the acid dissociation constant of the acid component of the buffer
[A-] = concentration of the conjugate base
[HA] = concentration of the acid
In this case, formic acid (HCOOH) is a weak acid that partially dissociates into formate (HCOO-) and hydrogen ions (H+). The acid dissociation constant (Ka) for formic acid is 3.75. The pKa, which is the negative logarithm of Ka, is equal to -log(3.75).
Given that the pH of the buffer solution is 4.23, we can use the Henderson-Hasselbalch equation to calculate the concentration of formate.
First, calculate the concentration of formate ([A-]):
pH = pKa + log([A-] / [HA])
Rearrange the equation:
[A-] / [HA] = 10^(pH - pKa)
Substitute the values into the equation:
[A-] / 0.35 = 10^(4.23 - (-log(3.75)))
Simplify:
[A-] = 0.35 * 10^(4.23 - (-log(3.75)))
Using a scientific calculator, calculate the value within parentheses:
[A-] = 0.35 * 10^(4.23 + 0.575)
[A-] = 0.35 * 10^(4.805)
[A-] ≈ 0.35 * 6217.74
[A-] ≈ 2176.21
Therefore, the concentration of formate ([A-]) is approximately 2176.21 units.