You find the Ka of your unknown acid is 6.3x10
-5
. If you are being asked to make a buffer at pH 4.00, what
is the appropriate ratio of A-
to HA to be combined in your flask?
To create a buffer at pH 4.00 using the given Ka value, we need to use the Henderson-Hasselbalch equation:
pH = pKa + log([A-]/[HA])
Given that the pH is 4.00, we can rearrange the equation to solve for [A-]/[HA]:
[A-]/[HA] = 10^(pH - pKa)
Using the given Ka, which is 6.3x10^(-5), and plugging in the values, we have:
[A-]/[HA] = 10^(4.00 - (-log10(6.3x10^(-5))))
[A-]/[HA] = 10^(4.00 + 4.2)
[A-]/[HA] = 10^8.2
[A-]/[HA] = 158,489.32
Therefore, the appropriate ratio of A- to HA to be combined in your flask is approximately 1:158,489 (or rounded to the nearest whole number).
To determine the appropriate ratio of A- to HA for the buffer solution, you need to use the Henderson-Hasselbalch equation, which relates the pH of the buffer solution to the concentrations of the conjugate acid (HA) and conjugate base (A-).
The Henderson-Hasselbalch equation is given as:
pH = pKa + log([A-]/[HA])
In this equation, pKa is the negative logarithm of the acid dissociation constant (Ka) of the weak acid. In this case, pKa = -log(Ka) = -log(6.3x10^-5).
To make a buffer at pH 4.00, you need to rearrange the Henderson-Hasselbalch equation to solve for the ratio [A-]/[HA]:
pH - pKa = log([A-]/[HA])
4.00 - (-log(6.3x10^-5)) = log([A-]/[HA])
4.00 + 4.20 = log([A-]/[HA])
8.20 = log([A-]/[HA])
To find the ratio [A-]/[HA], we need to take the antilog (inverse logarithm) of both sides of the equation:
[A-]/[HA] = 10^8.20
[A-]/[HA] = 1.58x10^8
Therefore, the appropriate ratio of A- to HA to be combined in your flask is approximately 1.58x10^8 : 1. This means that for every one molecule of HA, you should combine 1.58x10^8 molecules of A- to create the buffer solution at pH 4.00.
Use the HH equation.
pH = pKa + log (base)/(acid)