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Chapter 15: Problem 1

What are the major species in solution after \(\mathrm{NaHSO}_{4}\) is dissolved in water? What happens to the \(\mathrm{pH}\) of the solution as more \(\mathrm{NaHSO}_{4}\) is added? Why? Would the results vary if baking soda \(\left(\mathrm{NaHCO}_{3}\right)\) were used instead?

Short Answer

Expert verified
When NaHSO4 is dissolved in water, the major species present are Na+, HSO4-, H+, and OH-. The solution is initially acidic due to HSO4- dissociation. As more NaHSO4 is added, the solution becomes more acidic due to an increase in H+ ions. On the other hand, if NaHCO3 (baking soda) were used instead, the resulting solution would be less acidic because HCO3- ions can act as both weak acid and weak base, neutralizing some H+ ions.

Step by step solution

01

When sodium hydrogen sulfate (NaHSO4) is dissolved in water, it dissociates into its constituent ions: sodium (Na+) and hydrogen sulfate (HSO4-). Water itself can also dissociate to a small extent, forming hydrogen (H+) and hydroxide (OH-) ions. Therefore, the major species in the solution would be Na+, HSO4-, H+, and OH-. #Step 2: Determine initial acidity or basicity of the NaHSO4 solution#

HSO4- is a weak acid because it can donate a proton (H+) to the solution. As HSO4- dissociates, it forms H+ and SO4^2- ions. The initial NaHSO4 solution would therefore have an acidic pH because of the acid dissociation. #Step 3: Analyze the effect of adding more NaHSO4 on the pH#
02

When more NaHSO4 is added to the solution, it will dissociate and increase the concentration of HSO4- ions. As HSO4- ion concentration increases, so does its dissociation into H+ and SO4^2-, causing the concentration of H+ ions to rise. The pH of the solution, which is defined as \(-\log [\mathrm{H}^+]\) in water, will decrease (become more acidic) due to an increase in H+ ions. #Step 4: Compare the results with sodium hydrogen carbonate (NaHCO3)#

Now, let's analyze what would happen if sodium hydrogen carbonate (baking soda) or NaHCO3 is used instead. When NaHCO3 dissolves in water, it dissociates into Na+ and HCO3- ions. The HCO3- ion can both donate and accept protons; it can act as both a weak acid and a weak base. Since it can act as a weak base, adding NaHCO3 to the solution can neutralize some of the H+ ions present due to the dissociation of HSO4- ions. Consequently, the solution's acidity will decrease, and its pH will increase, meaning it would be less acidic in comparison to the NaHSO4 solution. In summary: 1. The major species in the NaHSO4 solution are Na+, HSO4-, H+, and OH-. 2. The initial NaHSO4 solution is acidic. 3. Adding more NaHSO4 makes the solution more acidic due to an increase in H+ ions. 4. If baking soda (NaHCO3) were used instead, the resulting solution would be less acidic than the original NaHSO4 solution due to its weak acid/base properties.

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Most popular questions from this chapter

Consider a solution formed by mixing \(50.0 \mathrm{~mL}\) of \(0.100 \mathrm{M}\) \(\mathrm{H}_{2} \mathrm{SO}_{4}, 30.0 \mathrm{~mL}\) of \(0.100 \mathrm{M} \mathrm{HOCl}, 25.0 \mathrm{~mL}\) of \(0.200 \mathrm{M} \mathrm{NaOH}\). \(25.0 \mathrm{~mL}\) of \(0.100 \mathrm{M} \mathrm{Ba}(\mathrm{OH})_{2}\), and \(10.0 \mathrm{~mL}\) of \(0.150 \mathrm{M} \mathrm{KOH}\). Calculate the \(\mathrm{pH}\) of this solution.

Two drops of indicator \(\operatorname{HIn}\left(K_{\mathrm{a}}=1.0 \times 10^{-9}\right)\), where HIn is yellow and \(\mathrm{In}^{-}\) is blue, are placed in \(100.0 \mathrm{~mL}\) of \(0.10 \mathrm{M} \mathrm{HCl}\). a. What color is the solution initially? b. The solution is titrated with \(0.10 M \mathrm{NaOH}\). At what \(\mathrm{pH}\) will the color change (yellow to greenish yellow) occur? c. What color will the solution be after \(200.0 \mathrm{~mL} \mathrm{NaOH}\) has been added?

The active ingredient in aspirin is acetylsalicylic acid. A \(2.51-g\) sample of acetylsalicylic acid required \(27.36 \mathrm{~mL}\) of \(0.5106 M\) \(\mathrm{NaOH}\) for complete reaction. Addition of \(13.68 \mathrm{~mL}\) of \(0.5106 \mathrm{M}\) \(\mathrm{HCl}\) 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.

Calculate the \(\mathrm{pH}\) at the halfway point and at the equivalence point for each of the following titrations. a. \(100.0 \mathrm{~mL}\) of \(0.10 \mathrm{M} \mathrm{HC}_{7} \mathrm{H}_{5} \mathrm{O}_{2}\left(K_{\mathrm{a}}=6.4 \times 10^{-5}\right)\) titrated by \(0.10 \mathrm{M} \mathrm{NaOH}\) b. \(100.0 \mathrm{~mL}\) of \(0.10 \mathrm{M} \mathrm{C}_{2} \mathrm{H}_{3} \mathrm{NH}_{2}\left(K_{\mathrm{b}}=5.6 \times 10^{-4}\right)\) titrated by \(0.20 \mathrm{M} \mathrm{HNO}_{3}\) c. \(100.0 \mathrm{~mL}\) of \(0.50 \mathrm{M} \mathrm{HCl}\) titrated by \(0.25 \mathrm{M} \mathrm{NaOH}\)

A friend asks the following: "Consider a buffered solution made up of the weak acid HA and its salt NaA. If a strong base like \(\mathrm{NaOH}\) is added, the HA reacts with the \(\mathrm{OH}^{-}\) to form \(\mathrm{A}^{-}\). Thus the amount of acid (HA) is decreased, and the amount of base \(\left(\mathrm{A}^{-}\right)\) is increased. Analogously, adding \(\mathrm{HCl}\) to the buffered solution forms more of the acid (HA) by reacting with the base \(\left(\mathrm{A}^{-}\right) .\) Thus how can we claim that a buffered solution resists changes in the \(\mathrm{pH}\) of the solution?' How would you explain buffering to this friend?
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