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Chapter 16: Problem 65

Define salt hydrolysis. Categorize salts according to how they affect the \(\mathrm{pH}\) of a solution.

Short Answer

Expert verified
Salt hydrolysis is the process where a salt reacts with water to form either an acidic or basic solution. Salts can be categorized based on their impact on pH into: salts from a strong acid and a strong base that form neutral solutions, salts from a strong acid and a weak base that form acidic solutions, and salts from a weak acid and a strong base that form basic solutions.

Step by step solution

01

Define Salt Hydrolysis

Salt Hydrolysis is the reaction of a salt with water to form either an acidic or basic solution. When salts are dissolved in water, they may undergo a reaction with water ions, referred to as hydrolysis, which can result in the formation of acidic or basic solutions.
02

Understanding Impact of Salt on pH

Salts can affect the \(pH\) of a solution because they can form acidic, basic, or neutral solutions when dissolved in water. This is dependent on the properties of the parent acid and base.
03

Categorizing Salts

Salts can be categorized into three types based on their effect on the pH of a solution:\n 1. Salts derived from a strong acid and a strong base (e.g., NaCl, KNO3): When such salts dissolve in water, the solution is neutral (\(pH = 7\)) because the cation and the anion do not react with water.\n 2. Salts derived from a strong acid and a weak base (e.g., NH4Cl): These salts yield an acidic solution (\(pH < 7\)) because the anion from the strong acid does not react with the water, but the cation from the weak base hydrolyzes to form H3O+.\n 3. Salts derived from a weak acid and a strong base (e.g., CH3COONa): These salts yield a basic solution (\(pH > 7\)) because the cation from the strong base does not react with water, but the anion from the weak acid harbors to form OH- .

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Most popular questions from this chapter

Identify the acid-base conjugate pairs in each of these reactions: (a) \(\mathrm{CH}_{3} \mathrm{COO}^{-}+\mathrm{HCN} \rightleftharpoons \mathrm{CH}_{3} \mathrm{COOH}+\mathrm{CN}^{-}\) (b) \(\mathrm{HCO}_{3}^{-}+\mathrm{HCO}_{3}^{-} \rightleftharpoons \mathrm{H}_{2} \mathrm{CO}_{3}+\mathrm{CO}_{3}^{2-}\) (c) \(\mathrm{H}_{2} \mathrm{PO}_{4}^{-}+\mathrm{NH}_{3} \rightleftharpoons \mathrm{HPO}_{4}^{2-}+\mathrm{NH}_{4}^{+}\) (d) \(\mathrm{HClO}+\mathrm{CH}_{3} \mathrm{NH}_{2} \rightleftharpoons \mathrm{CH}_{3} \mathrm{NH}_{3}^{+}+\mathrm{ClO}^{-}\) (e) \(\mathrm{CO}_{3}^{2-}+\mathrm{H}_{2} \mathrm{O} \rightleftharpoons \mathrm{HCO}_{3}^{-}+\mathrm{OH}^{-}\) (f) \(\mathrm{CH}_{3} \mathrm{COO}^{-}+\mathrm{H}_{2} \mathrm{O} \rightleftharpoons \mathrm{CH}_{3} \mathrm{COOH}+\mathrm{OH}^{-}\)

\(\mathrm{Al}^{3+}\) is not a Bronsted acid, but \(\mathrm{Al}\left(\mathrm{H}_{2} \mathrm{O}\right)_{6}^{3+}\) is Explain.

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The \(\mathrm{pH}\) of a solution is 6.7 . From this statement alone, can you conclude that the solution is acidic? If not, what additional information would you need? Can the \(\mathrm{pH}\) of a solution be zero or negative? If so, give examples to illustrate these values.

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