Can someone answer part 2 This pre-lab assignment is mandatory. You will receive

Question

Can someone answer part 2

This pre-lab assignment is mandatory. You will receive a zero for the experiment if the pre-lab assignment is not completed before the beginning of the lab session. Calculate the masses of the conjugate acid and base needed to prepare each of the six buffer solutions listed in the table on the left, (pK_3 of H_2PO^-_4 = 6.70) Assume the availability of the phosphate salts listed in the table on the right. Fill out the worksheet below. Show all work on separate pages for full credit. Suppose you added 2.00 mL of 0.250 M HCl to the pH 6.70 hydrogen phosphate buffer that was made in Question 1. What is the pH of the resulting buffer solution? Show all work.

Explanation / Answer

According to Henderson's Equation,

pH = pKa + log [salt]/[acid]

salt = HPO42-

acid = H2PO4-

H2PO4-   ---> HPO42- + H+

When pH = pKa = 6.7 , [salt]=[acid]

[HPO42-]=[ H2PO4-]

The buffer solution contains total 0.25 M acid and salt. salt + acid= 0.25 M = 0.25*100/1000 = 0.025 moles

So, if salt present at pH = 6.7 is x moles, then acid present = 0.025-x moles

x = 0.025 - x

x = 0.0125 moles

So, [HPO42-]=[ H2PO4-] = 0.0125 moles

2 mL of 0.25 M HCl contains = 0.25*2/1000 = 0.0005 moles

Added HCl reacts with HPO42- to produce H2PO4-.

As a result, concentration of H2PO4- increases and concentration of decreases HPO42-.

Amount of HPO42- after reaction = 0.0125-0.0005 moles = 0.012 moles

[HPO42-] = 0.012*1000/total volume of solution = (0.012*1000/102) M = 0.117 M

Amount of H2PO4- after reaction = 0.0125+0.0005 moles = 0.013 moles

[H2PO4-] = 0.013*1000/total volume of solution = (0.013*1000/102) M = 0.127 M

pH = pKa + log [salt]/[acid]

= 6.7 + log 0.117/0.127

= 6.66

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