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Chapter 17: Problem 23

A buffer is prepared by adding \(20.0 \mathrm{~g}\) of sodium acetate \(\left(\mathrm{CH}_{3} \mathrm{COONa}\right)\) to \(500 \mathrm{~mL}\) of a \(0.150 \mathrm{M}\) acetic acid \(\left(\mathrm{CH}_{3} \mathrm{COOH}\right)\) solution. (a) Determine the \(\mathrm{pH}\) of the buffer. (b) Write the complete ionic equation for the reaction that occurs when a few drops of hydrochloric acid are added to the buffer. (c) Write the complete ionic equation for the reaction that occurs when a few drops of sodium hydroxide solution are added to the buffer.

Short Answer

Expert verified
The pH of the buffer solution is approximately 5.76. The complete ionic equation for the reaction between hydrochloric acid and buffer components is \(H^{+}(aq) + CH_3COO^{-}(aq) \rightarrow CH_3COOH(aq)\), and for the reaction between sodium hydroxide and buffer components is \(OH^{-}(aq) + CH_3COOH(aq) \rightarrow CH_3COO^{-}(aq) + H_2O(l)\).

Step by step solution

01

1. Calculate the moles of sodium acetate and acetic acid in the buffer solution

To calculate the moles of sodium acetate, we can use the given mass and molecular weight of the substance: Moles of sodium acetate = (20.0 g) / (molecular weight of CH₃COONa) Molecular weight of CH₃COONa = (12.01 * 2) + (1.01 * 3) + (16.00 * 2) + (22.99 * 1) = 82.04 g/mol Moles of sodium acetate = 20.0 g / 82.04 g/mol ≈ 0.244 mol Now, let's calculate the moles of acetic acid: Moles of acetic acid = (0.150 mol/L) x (0.500 L) = 0.075 mol
02

2. Use the Henderson-Hasselbalch equation to determine the pH of the buffer

The Henderson-Hasselbalch equation is: pH = pKa + log ([A⁻] / [HA]) We're given that the acetic acid (CH₃COOH) has a pKa = 4.75. We have moles of sodium acetate (A⁻) and moles of acetic acid (HA), so we can calculate the pH: pH = 4.75 + log (0.244 mol / 0.075 mol) ≈ 4.75 + 1.01 ≈ 5.76 Thus, the pH of the buffer solution is approximately 5.76.
03

3. Write the complete ionic equation for the reaction between hydrochloric acid and the buffer constituents

When hydrochloric acid (HCl) is added to the buffer, it reacts with the acetate ion (CH₃COO⁻) from sodium acetate to form acetic acid: H⁺(aq) + Cl⁻(aq) + CH₃COO⁻(aq) -> CH₃COOH(aq) + Cl⁻(aq) The complete ionic equation is: H⁺(aq) + CH₃COO⁻(aq) -> CH₃COOH(aq)
04

4. Write the complete ionic equation for the effect that occurs when sodium hydroxide is added to the buffer

When sodium hydroxide (NaOH) is added to the buffer, it reacts with the undissociated acetic acid (CH₃COOH) to form sodium acetate: Na⁺(aq) + OH⁻(aq) + CH₃COOH(aq) -> CH₃COO⁻(aq) + H₂O(l) + Na⁺(aq) The complete ionic equation is: OH⁻(aq) + CH₃COOH(aq) -> CH₃COO⁻(aq) + H₂O(l) To summarize, the pH of the buffer solution is approximately 5.76. The complete ionic equation for the reaction between hydrochloric acid and buffer components is H⁺(aq) + CH₃COO⁻(aq) -> CH₃COOH(aq), and for the reaction between sodium hydroxide and buffer components is OH⁻(aq) + CH₃COOH(aq) -> CH₃COO⁻(aq) + H₂O(l).

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Most popular questions from this chapter

The value of \(K_{s p}\) for \(\mathrm{Cd}(\mathrm{OH})_{2}\) is \(2.5 \times 10^{-14}\) (a) What is the molar solubility of \(\mathrm{Cd}(\mathrm{OH})_{2} ?(\mathbf{b})\) The solubility of \(\mathrm{Cd}(\mathrm{OH})_{2}\) can be increased through formation of the complex ion \(\mathrm{CdBr}_{4}^{2-}\left(K_{f}=5 \times 10^{3}\right) .\) If solid \(\mathrm{Cd}(\mathrm{OH})_{2}\) is added to a NaBr solution, what would the initial concentration of \(\mathrm{NaBr}\) need to be in order to increase the molar solubility of \(\mathrm{Cd}(\mathrm{OH})_{2}\) to \(1.0 \times 10^{-3}\) moles per liter?

\(\begin{array}{llll}& \text { (a) If the molar solubility of } & \mathrm{CaF}_{2} & \text { at } & 35^{\circ} \mathrm{C} & \text { is }\end{array}\) \(1.24 \times 10^{-3} \mathrm{~mol} / \mathrm{L},\) what is \(K_{s p}\) at this temperature? (b) It is found that \(1.1 \times 10^{-2}\) of \(\mathrm{SrF}_{2}\) dissolves per \(100 \mathrm{~mL}\) of aqueous solution at \(25^{\circ} \mathrm{C} .\) Calculate the solubility product for \(\mathrm{SrF}_{2}\). (c) The \(K_{s p}\) of \(\mathrm{Ba}\left(\mathrm{IO}_{3}\right)_{2}\) at \(25^{\circ} \mathrm{C}\) is \(6.0 \times 10^{-10} .\) What is the molar solubility of \(\mathrm{Ba}\left(\mathrm{IO}_{3}\right)_{2}\) ?

A sample of \(0.2140 \mathrm{~g}\) of an unknown monoprotic acid was dissolved in \(25.0 \mathrm{~mL}\) of water and titrated with \(0.0950 \mathrm{M} \mathrm{NaOH}\). The acid required \(27.4 \mathrm{~mL}\) of base to reach the equivalence point. (a) What is the molar mass of the acid? (b) After \(15.0 \mathrm{~mL}\) of base had been added in the titration, the \(\mathrm{pH}\) was found to be 6.50 . What is the \(K_{a}\) for the unknown acid?

A \(35.0-\mathrm{mL}\) sample of \(0.150 \mathrm{M}\) acetic acid \(\left(\mathrm{CH}_{3} \mathrm{COOH}\right)\) is titrated with \(0.150 \mathrm{M} \mathrm{NaOH}\) solution. Calculate the \(\mathrm{pH}\) after the following volumes of base have been added: (a) \(0 \mathrm{~mL},(\mathbf{b})\) \(17.5 \mathrm{~mL},(\mathrm{c}) 34.5 \mathrm{~mL},(\) d) \(35.0 \mathrm{~mL},\) (e) \(35.5 \mathrm{~mL},\) (f) \(50.0 \mathrm{~mL} .\)

Consider the titration of \(30.0 \mathrm{~mL}\) of \(0.050 \mathrm{M} \mathrm{NH}_{3}\) with \(0.025 \mathrm{M}\) HCl. Calculate the pH after the following volumes of titrant have been added: (a) \(0 \mathrm{~mL},(\mathbf{b}) 20.0 \mathrm{~mL},(\mathbf{c}) 59.0 \mathrm{~mL},(\mathrm{~d}) 60.0 \mathrm{~mL},\) (e) \(61.0 \mathrm{~mL}\) (f) \(65.0 \mathrm{~mL}\).
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