Question

1. Show the calculations required for preparing the buffers in this exercise. The pKa values for H3PO4 are 2.15, 6.82, 12.32. Note that you’ll be using the pKa value of 6.82. Why?

Method 1 a. Calculate the grams of NaH2PO4*H2O needed to prepare 100.0 mL of a 100mM solution. b. Calculate the amount of 1.00M NaOH you expect to add to adjust the pH of your NaH2PO4 solution to 7.0.

Method 2 c. Using the Henderson-Hasselbalch equation, calculate the g of NaH2PO4*H2O and g of Na2HPO4 needed to prepare 100.0 mL of 100.0mM sodium phosphate buffer, pH 7.0.

2. Describe how you would prepare the following solutions: a. 100.0 mL of 1-butanol:formic acid:water (100:35:25) b. 100.0 mL of 0.025M potassium hydrogen phosphate, 0.025 M disodium hydrogen phosphate, 0.100M NaCl

Help Please!

Answer 1

1) a. 100 mM = 100 mmol NaH2PO4.H2O/1 L solution

i.e. 10 mmol NaH2PO4.H2O is present in 100 mL solution

The molar mass of NaH2PO4.H2O = 138 g mol^-1

i.e. The mass of NaH2PO4.H2O = 10*10^-3 mol * 138 g mol^-1 = 1.38 g

b and c. According to Henderson-Hasselbulch equation:

pH = pKa2 + (n_NaOH/n_NaH2PO4.H2O)

Here, you need to adjust pH to 7.

i.e. 7 = 6.82 + Log(10 mmol/x mL*1 mmol/mL)

i.e. Log(10/x) = 7 - 6.82

i.e. 10/x = 10^0.18

i.e. x = 10/1.5

i.e. The volume of 1 M NaOH needed to adjust pH of the solution to 7 is 6.6 mL

Note: Here, you need to use pKa = 6.82 because the equilibrium exists between NaH₂PO4 and Na2HPO4

i.e. H₂PO₄^- HPO₄^2-