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Chapter 17: Problem 20

You have the following solutions, all of the same molar concentrations: \(\mathrm{KBr}, \mathrm{HBr}, \mathrm{CH}_{3} \mathrm{NH}_{2}\), and \(\mathrm{NH}_{4} \mathrm{Cl}\). Rank them from the lowest to the highest hydroxide-ion concentrations.

Short Answer

Expert verified
HBr < 狈贬鈧凛濒 < KBr < 颁贬鈧僋贬鈧� for hydroxide-ion concentration.

Step by step solution

01

Identify the Nature of Each Solution

First, identify the acidic or basic nature of each compound based on their chemical formulas: - **KBr** is a neutral salt, formed from a strong acid (HBr) and a strong base (KOH). - **HBr** is a strong acid. - **颁贬鈧僋贬鈧�** is a weak base (methylamine). - **狈贬鈧凛濒** is an acidic salt, formed from a weak base (NH鈧�) and a strong acid (HCl).
02

Understand Hydroxide-ion Concentrations

The hydroxide-ion concentration ( [ ext{OH}^- ] ) depends on whether the compound acts as an acid or base: - Strong acids like HBr do not contribute to hydroxide ions as they completely dissociate into hydrogen ions. - Strong electrolytes like KBr do not alter [ ext{OH}^- ] significantly after dissociation. - Weak bases like 颁贬鈧僋贬鈧� increase hydroxide ions upon dissociation. - Acidic salts like 狈贬鈧凛濒 tend to decrease [ ext{OH}^- ] since they provide hydrogen ions.
03

Rank Based on Hydroxide-ion Concentrations

Considering the ionization and resulting hydroxide-ion concentrations: - **HBr** will have the lowest [ ext{OH}^- ] because it is a strong acid. - **狈贬鈧凛濒** will have low [ ext{OH}^- ] due to the formation of NH鈧刕+ which releases H^+ . - **KBr** remains mainly neutral, resulting in a minor presence of [ ext{OH}^- ] . - **颁贬鈧僋贬鈧�** being a weak base, will maximize [ ext{OH}^- ] .
04

Write the Final Ranking

Based on the analyses above, the order from lowest to highest hydroxide-ion concentration is: 1. **HBr** 2. **狈贬鈧凛濒** 3. **KBr** 4. **颁贬鈧僋贬鈧�**

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Key Concepts

These are the key concepts you need to understand to accurately answer the question.

Acid-Base Chemistry
Acid-base chemistry is a fundamental part of understanding how solutions behave in water. When acids or bases dissolve in water, they may ionize, introducing hydrogen ions (H鈦�) or hydroxide ions (OH鈦�) into the solution. This process affects the solution's pH, an indication of its acidity or basicity.
To identify whether a compound will yield an acidic or basic solution, consider the strength of the acid or base. Strong acids or bases dissociate fully in water:
  • Strong acids like hydrochloric acid (HCl) will significantly affect the pH by introducing more hydrogen ions, lowering the pH and making the solution more acidic.
  • Weak bases, like ammonia (NH鈧�) or methylamine (颁贬鈧僋贬鈧�), do not fully dissociate and slightly raise the hydroxide ions, increasing the solution's basicity.
In the context of our exercise, understanding these behaviors is key to predicting the outcome in terms of hydroxide-ion concentration.
Solution Ranking
In solution ranking, we must understand how each compound influences hydroxide-ion concentration. Each solution behaves differently based on its acid-base nature, which influences the presence of OH鈦� in the solution.
Let's look at the given compounds:
  • KBr: It is a neutral salt formed from a strong acid and a strong base. It has little to no effect on hydroxide-ion concentration, remaining largely inert.
  • HBr: Being a strong acid, it completely dissociates, increasing H鈦� ions and thus has the minimal effect on hydroxide ions.
  • 颁贬鈧僋贬鈧�: As a weak base, it partially dissociates in water to release some OH鈦�, making it the solution with the highest hydroxide-ion concentration among the choices.
  • 狈贬鈧凛濒: This is an acidic salt due to NH鈧勨伜 which releases H鈦� upon dissociation, thereby further reducing OH鈦� concentration.
This detailed analysis helps in ranking these solutions based on their hydroxide-ion presence: HBr < 狈贬鈧凛濒 < KBr < 颁贬鈧僋贬鈧�.
Weak Bases
Weak bases, unlike their strong counterparts, do not fully dissociate in water. As a result, they only partially ionize, which limits their ability to increase hydroxide ions significantly.
Such bases establish an equilibrium between the undissociated compounds and the ions in the solution. Consider methylamine (颁贬鈧僋贬鈧�):
  • 颁贬鈧僋贬鈧� in water partially dissociates to form CH鈧僋H鈧冣伜 and OH鈦�.
  • This reaction increases the number of hydroxide ions compared to neutral solutions.
Understanding weak bases is crucial for predicting their impact on the solution's alkalinity. Their partial dissociation contrasts starkly with strong acids, which fully ionize, leading to notably different hydroxide-ion concentrations.
Strong Acids
Strong acids like hydrobromic acid (HBr) are unique because they are fully dissociative in aqueous solutions. This complete dissociation results in increased hydrogen ions in the solution, without contributing any OH鈦� ions.
This is what happens with strong acids:
  • They disassociate completely in water, releasing a high concentration of H鈦� ions.
  • The presence of more H鈦� ions shifts the equilibrium of water鈥檚 auto-ionization toward more formation of H鈧侽 and less OH鈦�.
Due to this, strong acids are ranked lowest when considering hydroxide-ion concentration. They do not add OH鈦� to the solution but rather drive the reaction toward a decrease in its concentration, enhancing the solution's acidity.

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Most popular questions from this chapter

Tartaric acid is a weak diprotic fruit acid with \(K_{a 1}=\) \(1.0 \times 10^{-3}\) and \(K_{a 2}=4.6 \times 10^{-5}\) a. Letting the symbol \(\mathrm{H}_{2} \mathrm{~A}\) represent tartaric acid, write the chemical equations that represent \(K_{a 1}\) and \(K_{a 2} .\) Write the chemical equation that represents \(K_{a 1} \times K_{a 2}\) b. Qualitatively describe the relative concentrations of \(\mathrm{H}_{2} \mathrm{~A}\), \(\mathrm{HA}^{-}, \mathrm{A}^{2-}\), and \(\mathrm{H}_{3} \mathrm{O}^{+}\) in a solution that is about \(0.5 \mathrm{M}\) in tartaric acid. c. Calculate the \(\mathrm{pH}\) of a \(0.0250 \mathrm{M}\) tartaric acid solution and the equilibrium concentration of \(\left[\mathrm{H}_{2} \mathrm{~A}\right]\). d. What is the \(A^{2-}\) concentration?

What is the pH of the solution obtained by titrating \(1.24 \mathrm{~g}\) of sodium hydrogen sulfate, \(\mathrm{NaHSO}_{4}\), dissolved in \(50.0 \mathrm{~mL}\) of water with \(0.180 M\) sodium hydroxide until the equivalence point is reached? Assume that any volume change due to adding the sodium hydrogen sulfate or to mixing the solutions is negligible.

A quantity of \(0.25 M\) sodium hydroxide is added to a solution containing \(0.15\) mole of acetic acid. The final volume of the solution is \(375 \mathrm{~mL}\) and the \(\mathrm{pH}\) of this solution is \(4.45\). a. What is the molar concentration of the sodium acetate? b. How many milliliters of sodium hydroxide were added to the original solution? c. What was the original concentration of the acetic acid?

Each of the following statements concerns a \(0.10 M\) solution of a weak organic base, B. Briefly describe why each statement is either true or false. a. \([\mathrm{B}]\) is approximately equal to \(0.10 \mathrm{M}\). b. [B] is much greater than \(\left[\mathrm{HB}^{+}\right]\). c. \(\left[\mathrm{H}_{3} \mathrm{O}^{+}\right]\) is greater than \(\left[\mathrm{HB}^{+}\right]\). d. The \(\mathrm{pH}\) is 13 . e. \(\left[\mathrm{HB}^{+}\right]\) is approximately equal to \(\left[\mathrm{OH}^{-}\right]\). f. \(\left[\mathrm{OH}^{-}\right]\) equals \(0.10 \mathrm{M}\).

What is meant by the capacity of a buffer? Describe a buffer with low capacity and the same buffer with greater capacity.
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