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Chapter 14: Problem 7

You have a solution of the weak acid HA and add some HCl to it. What are the major species in the solution? What do you need to know to calculate the \(\mathrm{pH}\) of the solution, and how would you use this information? How does the \(\mathrm{pH}\) of the solution of just the HA compare with that of the final mixture? Explain.

Short Answer

Expert verified
The major species present in the solution of weak acid HA and strong acid HCl are HA, H+ (from both HA and HCl), A- (conjugate base of HA), and Cl- (from HCl). To calculate the pH of the solution, the initial concentrations of HA and HCl, as well as the dissociation constant Ka, are needed. After calculating the total initial concentration of H+ ions and using charge balance equation, we can find [H3O+] from the equilibrium equation and calculate the pH by applying the formula: \[\mathrm{pH} = -\log[H_3O^+]\]. The pH of the final mixture will be lower (more acidic) than the pH of the initial HA solution due to the increased concentration of H+ ions from HCl.

Step by step solution

01

Identify the major species in the solution

Initially, the weak acid HA will dissociate into its ions H+ and A-, and HCl, being a strong acid, will be fully dissociated into H+ and Cl-. Since HCl is a strong acid, it will donate H+ ions to the solution, which will further ionize HA molecules. Therefore, the major species present in the solution are: 1. HA (weak acid) 2. H+ (from HA and HCl) 3. A- (conjugate base of HA after dissociation) 4. Cl- (from the dissociation of HCl)
02

Determine the information needed to calculate the pH of the solution

In order to calculate the pH of the solution, we need to know the following information: 1. The initial concentration of the weak acid HA ([HA]₀) 2. The initial concentration of the strong acid HCl ([HCl]₀) 3. The dissociation constant of the weak acid, Ka (which can be obtained from a table or using the pKa value)
03

Using the information to calculate the pH

To calculate the pH of the solution, we can follow these steps: 1. Write the equilibrium expression for HA as it donates a proton to water: \[HA + H_2O \rightleftharpoons H_3O^+ + A^-\] 2. Write the Ka expression for HA: \[Ka = \frac{[H_3O^+][A^-]}{[HA]}\] 3. Calculate the total initial concentration of H+ ions due to both HA and HCl: \[[H_3O^+]_0 = [H^+]_{HA} + [H^+]_{HCl}\] 4. Use the charge balance equation to relate the concentrations of A- and H3O+: \[[A^-] = [H_3O^+] - [Cl^-]\] 5. Substitute the concentrations (in terms of initial concentrations) and Ka value into the Ka expression. 6. Calculate the [H3O+] from the equilibrium equation and find the pH using the formula: \[\mathrm{pH} = -\log[H_3O^+]\]
04

Compare the pH of the HA solution with the final solution

The pH of the pure HA solution would be higher (meaning it is less acidic) than the final solution since the addition of HCl contributes to the increase of H+ ion concentration, making the solution more acidic. By finding the pH of the HA solution before the addition of HCl using the Ka expression and comparing it with the pH of the final mixture after adding HCl, one can observe the difference between the two pH values. The final mixture will have a lower pH than the initial HA solution.

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Most popular questions from this chapter

Consider the titration of \(100.0 \mathrm{mL}\) of \(0.100 \mathrm{M}\) HCN by \(0.100 M \mathrm{KOH}\) at \(25^{\circ} \mathrm{C} .\left(K_{\mathrm{a}} \text { for } \mathrm{HCN}=6.2 \times 10^{-10} .\right)\) a. Calculate the \(\mathrm{pH}\) after \(0.0 \mathrm{mL}\) of \(\mathrm{KOH}\) has been added. b. Calculate the \(\mathrm{pH}\) after \(50.0 \mathrm{mL}\) of \(\mathrm{KOH}\) has been added. c. Calculate the \(\mathrm{pH}\) after \(75.0 \mathrm{mL}\) of \(\mathrm{KOH}\) has been added. d. Calculate the \(\mathrm{pH}\) at the equivalence point. e. Calculate the pH after 125 mL of KOH has been added.

The active ingredient in aspirin is acetylsalicylic acid. A \(2.51-\mathrm{g}\) sample of acetylsalicylic acid required \(27.36 \mathrm{mL}\) of \(0.5106 M\) NaOH for complete reaction. Addition of \(13.68 \mathrm{mL}\) of \(0.5106 \mathrm{M}\) HCl to the flask containing the aspirin and the sodium hydroxide produced a mixture with \(\mathrm{pH}=3.48 .\) Determine the molar mass of acetylsalicylic acid and its \(K_{\mathrm{a}}\) value. State any assumptions you must make to reach your answer.

Mixing together solutions of acetic acid and sodium hydroxide can make a buffered solution. Explain. How does the amount of each solution added change the effectiveness of the buffer?

A buffered solution is made by adding \(50.0 \mathrm{g} \mathrm{NH}_{4} \mathrm{Cl}\) to 1.00 \(\mathrm{L}\) of a 0.75-M solution of \(\mathrm{NH}_{3}\). Calculate the pH of the final solution. (Assume no volume change.)

Which of the following can be classified as buffer solutions? a. \(0.25 M\) HBr \(+0.25 M\) HOBr b. \(0.15 M \mathrm{HClO}_{4}+0.20 \mathrm{M} \mathrm{RbOH}\) c. \(0.50 M\) HOCl \(+0.35 M\) KOCl d. \(0.70 M \mathrm{KOH}+0.70 \mathrm{M} \mathrm{HONH}_{2}\) e. \(0.85 M \mathrm{H}_{2} \mathrm{NNH}_{2}+0.60 \mathrm{M} \mathrm{H}_{2} \mathrm{NNH}_{3} \mathrm{NO}_{3}\)
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