Calculate the concentrations of all species present and the pH in 0.10 M solutions of the following substances. See Appendix C for values of equilibrium constants. (b) Sodium acetate, Na1CH3CO22

Verified step by step guidance

Identify the species present in the solution: Sodium acetate (NaCH3COO) dissociates in water to form Na^+ and CH3COO^- ions.

Recognize that CH3COO^- is the conjugate base of acetic acid (CH3COOH) and can undergo hydrolysis in water to form CH3COOH and OH^-.

Write the hydrolysis equilibrium expression for CH3COO^-: CH3COO^- + H2O \rightleftharpoons CH3COOH + OH^-.

Use the equilibrium constant for the hydrolysis reaction, which is related to the K_a of acetic acid: K_b = \frac{K_w}{K_a}, where K_w is the ion-product constant of water.

Set up an ICE (Initial, Change, Equilibrium) table to determine the concentrations of CH3COOH, OH^-, and CH3COO^- at equilibrium, and use the K_b expression to solve for the concentration of OH^- to find the pH.

Verified video answer for a similar problem:

This video solution was recommended by our tutors as helpful for the problem above.

Video duration:

Was this helpful?

Key Concepts

Here are the essential concepts you must grasp in order to answer the question correctly.

Acid-Base Equilibria

Acid-base equilibria involve the transfer of protons (H+) between species in solution. In the case of sodium acetate, it acts as a weak base in water, where it can hydrolyze to produce acetate ions (CH3COO-) and hydroxide ions (OH-). Understanding the equilibrium between these species is crucial for calculating pH and concentrations.

Hydrolysis of Acetate Ion

The acetate ion (CH3COO-) can undergo hydrolysis in water, reacting with water to form acetic acid (CH3COOH) and hydroxide ions (OH-). This reaction is essential for determining the pH of the solution, as it establishes the equilibrium that dictates the concentration of hydroxide ions, which in turn affects the pH.

Equilibrium Constants

Equilibrium constants (K) quantify the ratio of concentrations of products to reactants at equilibrium for a given reaction. For the hydrolysis of acetate, the equilibrium constant (K_b) can be used to calculate the concentrations of all species in solution. This is vital for determining the pH, as it allows for the calculation of the extent of the hydrolysis reaction.

Recommended video:

Guided course

01:14

Equilibrium Constant K

Related Practice

Textbook Question

Write a balanced net ionic equation for the principal reaction in solutions of each of the following salts. In each case, identify the Brønsted–Lowry acids and bases and the conjugate acid–base pairs. (a) Na2CO3

Textbook Question

Classify each of the following ions according to whether they react with water to give a neutral, acidic, or basic solution. (a) F-

Textbook Question

Classify each of the following salt solutions as neutral, acidic, or basic. See Appendix C for values of equilibrium constants. (a) Fe1NO323

Textbook Question

Calculate the pH and the percent dissociation of the hydrated cation in 0.020 M solutions of the following substances. See Appendix C for values of equilibrium constants. (a) Fe1NO322

Textbook Question

Calculate Ka for the cation and Kb for the anion in an aqueous NH4CN solution. Is the solution acidic, basic, or neutral?