Answer to Question #92684 in Biochemistry for kayley

Question

You need to conduct an experiment in the laboratory. This requires that you prepare 500 ml of 0.5 M sodium acetate buffer with pH = 4.30. In laboratory you have 10 M solution of CH3COOH (pKa = 4.75), and a stock of CH3COONa.3H2O (MW=136.082 g/mol). Using the Henderson-Hasselbach equation and the mass balance and the ionization equation below, calculate how to prepare this buffer.
Assume that the preparation consists of weighing out the necessary amounts of the compounds, or pipetting the required volume of the stock solution available. These are subsequently dissolved in water, and the buffer is prepared in a volumetric flask. (8)
CH3COOH +H2O CH3COO- + H30
+

Answer 1

$pH=pK_a + log \frac {[Acetate^-]} {[HAcetate]}$

$4.30 = 4.75 + log \frac {[Acetate^-}{[HAcetate]}$

$\frac {[Acetate^-]} {HAcetate]} = 10^{-0.45} = 0.35$

$[Acetate^-] + [HAcetate] = 0.5 M$

So, we have an equation with two unknowns. By doing simple math we find concentrations as:

[Acetate^-] = 0.37M

[HAcetate] = 0.13 M

Now we can find moles of acetic acid as:

$0.13M \cdot 0.5L = 0.065 moles$

$\frac {0.065 moles} {0.5 M} = 0.13 L = 130 mL$

and sodium acetate:

$0.37 M \cdot 0.5L = 0.185 moles$

$0.185 moles \cdot 136.082 g/mol = 25.2 g$

So, to prepare buffer solution we need to take 130 mL of 10M acid and 25.2 g of sodium acetate in a volumetric flask and dilute with DI water up to 500 mL.