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Chapter 2: Problem 5

Given \(0.1 \mathrm{M}\) solutions of acetic acid and sodium acetate, describe the preparation of \(1 \mathrm{L}\) of \(0.1 \mathrm{M}\) acetate buffer at a pH of 5.4.

Short Answer

Expert verified
In summary, to prepare 1 L of 0.1 M acetate buffer at a pH of 5.4: Firstly, calculate the pKa of the acetic acid. Secondly, apply the Henderson-Hasselbalch equation to obtain the ratio of base(acetate ions) to acid (acetic acid). The ratio comes out to be ~4.68. Hence, mix the two 0.1 M solutions in a ratio of ~4.68:1 (base to acid) to prepare the buffer solution.

Step by step solution

01

Identify the buffer components

For an acetate buffer, acetic acid is the weak acid and sodium acetate is its conjugate base. In this case, acetic acid's Ka = 1.8 × 10^-5. From this, calculate the pKa = -log(Ka) = 4.74.
02

Apply the Henderson-Hasselbalch equation

The Henderson-Hasselbalch equation, which is pH = pKa + log([A-]/[HA]), is used to solve for buffer concentrations. Now plug in the given and calculated values: 5.4 = 4.74 + log([A-]/[HA])
03

Calculate the ratio of base to acid

Solve for the ratio of [A-]/[HA]. Subtract 4.74 from each side to find 0.66 = log([A-]/[HA]). Take the antilog of each side to find that [A-]/[HA] is approx 4.68.
04

Prepare the buffer solution

With respect to their molar ratio, mix the two 0.1 M solutions of acetic acid and sodium acetate in a volumetric flask to make 1 L of buffer solution. Since the molar ratio of base (A-) to acid (HA) is 4.68:1, you need to combine 0.82 M of acetic acid and 0.28 M of sodium acetate to prepare the buffer solution.

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Most popular questions from this chapter

a. If \(50 \mathrm{mL}\) of \(0.01 \mathrm{MHCl}\) is added to \(100 \mathrm{mL}\) of \(0.05 \mathrm{M}\) phosphate buffer at \(\mathrm{pH} 7.2,\) what is the resultant \(\mathrm{pH}\) ? What are the concentrations of \(\mathrm{H}_{2} \mathrm{PO}_{4}^{-}\) and \(\mathrm{HPO}_{4}^{2-}\) in the final solution? b. If \(50 \mathrm{mL}\) of \(0.01 \mathrm{MNaOH}\) is added to \(100 \mathrm{mL}\) of \(0.05 \mathrm{M}\) phosphate buffer at \(\mathrm{pH} 7.2,\) what is the resultant \(\mathrm{pH}\) ? What are the concentrations of \(\mathrm{H}_{2} \mathrm{PO}_{4}^{-}\) and \(\mathrm{HPO}_{4}^{2-}\) in this final solution?

a. Draw the titration curve for Bicine, assuming the \(\mathrm{p} K_{\mathrm{a}}\) for its free COOH group is 2.3 and the \(\mathrm{p} K_{\mathrm{a}}\) for its tertiary amino group is 8.3 b. Draw the structure of the fully deprotonated form (completely dissociated form) of bicine. c. You have available a \(0.1 ~ M\) solution of Bicine at its isoelectric point \(\left(\mathrm{pH}_{\mathrm{I}}\right), 0.1 \mathrm{M}\) solutions of \(\mathrm{HCl}\) and \(\mathrm{NaOH}\), and ample distilled \(\mathrm{H}_{2} \mathrm{O}\) Describe the preparation of 1 L of 0.04 M Bicine buffer, pH 7.5 d. What is the concentration of fully protonated form of Bicine in your final buffer solution?

Bicine is a compound containing a tertiary amino group whose relevant \(\mathrm{p} K_{\mathrm{a}}\) is 8.3 (Figure 2.17 ). Given \(1 \mathrm{L}\) of \(0.05 \mathrm{M}\) Bicine with its tertiary amino group in the unprotonated form, how much \(0.1 N \mathrm{HCl}\) must be added to have a Bicine buffer solution of \(\mathrm{pH} 7.5 ?\) What is the molarity of Bicine in the final buffer? What is the concentration of the protonated form of Bicine in this final buffer?

Calculate the \(\mathrm{pH}\) of the following. a. \(5 \times 10^{-4} \mathrm{MHCl}\) d. \(3 \times 10^{-2}\) M KOH b. \(7 \times 10^{-5} M\) NaOH e. \(0.04 \mathrm{m} M \mathrm{HCl}\) c. \(2 \mu M\) HCl f. \(6 \times 10^{-9}\) M HCl

The pH of a \(0.02 \mathrm{M}\) solution of an acid was measured at 4.6. a. What is the \(\left[\mathrm{H}^{+}\right]\) in this solution? b. Calculate the acid dissociation constant \(K_{\mathrm{a}}\) and \(\mathrm{p} K_{\mathrm{a}}\) for this acid.
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