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Chapter 16: Problem 28

A chemist prepares dilute solutions of equal molar concentrations of \(\mathrm{NH}_{3}, \mathrm{NH}_{4} \mathrm{Br}, \mathrm{NaF}\), and \(\mathrm{NaCl}\). Rank these solutions from highest \(\mathrm{pH}\) to lowest \(\mathrm{pH}\).

Short Answer

Expert verified
The ranking is: NH鈧� > NaF > NaCl > NH鈧凚r.

Step by step solution

01

Identify the Nature of Each Solution

Each compound in the problem forms a solution with specific properties: - **Ammonia (\( \mathrm{NH}_3 \), weak base)**: It slightly increases the \( \mathrm{pH} \) of the solution.- **Ammonium bromide (\( \mathrm{NH}_4\mathrm{Br} \), acidic salt)**: \(\mathrm{NH}_4^+\) is an acidic ion, slightly lowering \(\mathrm{pH}\).- **Sodium fluoride (\( \mathrm{NaF} \), salt of a weak acid and strong base)**: \(\mathrm{F}^-\) is a base ion, slightly increasing \(\mathrm{pH}\).- **Sodium chloride (\( \mathrm{NaCl} \), neutral salt)**: Has no significant effect on the \(\mathrm{pH}\).
02

Compare Basicity and Acidity

Evaluate the tendency of each compound to either increase or decrease the \(\mathrm{pH}\):- **\(\mathrm{NH}_3\):** Increases \(\mathrm{pH}\) due to its weak basic nature.- **\(\mathrm{NaF}\):** Increases \(\mathrm{pH}\) as \(\mathrm{F}^-\) acts as a weak base.- **\(\mathrm{NH}_4\mathrm{Br}\):** Decreases \(\mathrm{pH}\) because \(\mathrm{NH}_4^+\) acts as a weak acid.- **\(\mathrm{NaCl}\):** Does not change \(\mathrm{pH}\) much as it is neutral.
03

Rank the Solutions by pH

Based on whether each compound increases, decreases or does not change the \(\mathrm{pH}\), rank them from highest to lowest \(\mathrm{pH}\):1. \( \mathrm{NH}_3 \), the highest \(\mathrm{pH}\)2. \( \mathrm{NaF} \), slightly lower \(\mathrm{pH}\) than \( \mathrm{NH}_3 \)3. \( \mathrm{NaCl} \), neutral effect4. \( \mathrm{NH}_4\mathrm{Br} \), the lowest \(\mathrm{pH}\)

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Key Concepts

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

Weak Base
A weak base is a substance that partially dissociates in water to produce hydroxide ions, which slightly increases the pH of a solution. In our example, ammonia (\( \mathrm{NH}_3 \)) acts as a weak base. Unlike strong bases, which completely dissociate and greatly increase the pH, weak bases like \( \mathrm{NH}_3 \) only partially ionize.
The dissociation can be represented by the equation: \[\mathrm{NH}_3 + \mathrm{H}_2\mathrm{O} \rightleftharpoons \mathrm{NH}_4^+ + \mathrm{OH}^-\]
  • The production of \( \mathrm{OH}^- \) ions results in a basic solution.
  • Weak bases are common in household cleaning products due to their milder effects compared to strong bases.
Acidic Salt
An acidic salt is formed when a strong acid reacts with a weak base. In our example, ammonium bromide (\( \mathrm{NH}_4\mathrm{Br} \)) is an acidic salt. This is because it contains the ammonium ion (\( \mathrm{NH}_4^+ \)), which is the conjugate acid of the weak base \( \mathrm{NH}_3 \).
The ammonium ion reacts with water to form a slightly acidic solution, as shown in the reaction: \[\mathrm{NH}_4^+ + \mathrm{H}_2\mathrm{O} \rightleftharpoons \mathrm{NH}_3 + \mathrm{H}_3\mathrm{O}^+\]
  • The production of hydronium ions (\( \mathrm{H}_3\mathrm{O}^+ \)) lowers the pH, creating an acidic environment.
  • Acidic salts are often used in fertilizers and household products.
Neutral Salt
A neutral salt doesn't affect the pH of a solution significantly and results from the reaction of a strong acid and a strong base. \( \mathrm{NaCl} \), or sodium chloride, is a classic example of a neutral salt. In aqueous solutions, \( \mathrm{NaCl} \) dissociates completely into sodium ions (\( \mathrm{Na}^+ \)) and chloride ions (\( \mathrm{Cl}^- \)).

Since neither ion reacts with water to produce additional \( \mathrm{H}^+ \) or \( \mathrm{OH}^- \) ions, the pH of the solution tends to remain neutral, close to 7.
  • Neutral salts are commonly used in cooking and preserving foods.
  • They do not strongly alter the acid-base balance of a solution.
Basic Ion
Basic ions result from the dissociation of salts formed between a weak acid and a strong base. In the given context, fluoride ions (\( \mathrm{F}^- \)) in \( \mathrm{NaF} \) act as basic ions. They result from the ionization of sodium fluoride, a salt of strong base sodium hydroxide and weak acid hydrofluoric acid. \[\mathrm{NaF} \rightarrow \mathrm{Na}^+ + \mathrm{F}^-\]

The fluoride ion can react with water in the following way:\[\mathrm{F}^- + \mathrm{H}_2\mathrm{O} \rightleftharpoons \mathrm{HF} + \mathrm{OH}^-\]
  • The formation of \( \mathrm{OH}^- \) ions makes the solution slightly basic.
  • This increase in pH is lesser than that caused by a strong base but still noteworthy, indicating the weakly basic nature.

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Most popular questions from this chapter

A generic base, \(\mathrm{B}^{-}\), is added to \(2.25 \mathrm{~L}\) of water. The \(\mathrm{pH}\) of the solution is found to be \(10.10\). What is the concentration of the base \(\mathrm{B}^{-}\) in this solution? \(K_{a}\) for the acid \(\mathrm{HB}\) at \(25^{\circ} \mathrm{C}\) is \(1.99 \times 10^{-9}\).

A 0.288-g sample of an unknown monoprotic organic acid is dissolved in water and titrated with a \(0.115 M\) sodium hydroxide solution. After the addition of \(17.54 \mathrm{~mL}\) of base, a \(\mathrm{pH}\) of \(4.92\) is recorded. The equivalence point is reached when a total of \(33.83 \mathrm{~mL}\) of \(\mathrm{NaOH}\) is added. a. What is the molar mass of the organic acid? b. What is the \(K_{a}\) value for the acid? The \(K_{a}\) value could have been determined very easily if a pH measurement had been made after the addition of \(16.92 \mathrm{~mL}\) of \(\mathrm{NaOH}\). Why?

A chemist wanted to determine the concentration of a solution of lactic acid, \(\mathrm{HC}_{3} \mathrm{H}_{5} \mathrm{O}_{3} .\) She found that the \(\mathrm{pH}\) of the solution was \(2.51\). What was the concentration of the solution? The \(K_{a}\) of lactic acid is \(1.4 \times 10^{-4}\).

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?

Acrylic acid, whose formula is \(\mathrm{HC}_{3} \mathrm{H}_{3} \mathrm{O}_{2}\) or \(\mathrm{HO}_{2} \mathrm{CCH}=\mathrm{CH}_{2}\), is used in the manufacture of plastics. A \(0.10\) \(M\) aqueous solution of acrylic acid has a pH of \(2.63 .\) What is \(K_{a}\) for acrylic acid?
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