Ionic Equilibria: Acids and Bases (Chapter 18 Study Notes)

Ionic Equilibria: Acids and Bases

Introduction

This chapter explores the equilibria of acids and bases in aqueous solutions, focusing on the distinction between strong and weak electrolytes. Understanding these concepts is essential for predicting the behavior of acids, bases, and salts in water.

Review of Strong Electrolytes

Definition and Classification

• Strong electrolytes are substances that ionize or dissociate completely in water, producing ions that conduct electricity efficiently.

• Weak acids and bases ionize or dissociate only partially, much less than 100%.

• There are three main classes of strong electrolytes:

  1. Strong Water Soluble Acids

  2. Strong Water Soluble Bases

  3. Water Soluble Salts

Strong Water Soluble Acids

• Examples include nitric acid ($\mathrm{HNO_3}$), which ionizes completely in water:

$\mathrm{HNO_3_{(aq)} + H_2O_{(l)} \xrightarrow{\approx 100\%} H_3O^+_{(aq)} + NO_3^-_{(aq)}}$

Or equivalently: $\mathrm{HNO_3_{(aq)} \xrightarrow{\approx 100\%} H^+_{(aq)} + NO_3^-_{(aq)}}$

Strong Water Soluble Bases

• Examples include strontium hydroxide ($\mathrm{Sr(OH)_2}$), which dissociates completely:

$\mathrm{Sr(OH)_2_{(s)} \xrightarrow{H_2O} Sr^{2+}_{(aq)} + 2OH^-_{(aq)}}$

Water Soluble Salts

• Examples include sodium chloride and calcium nitrate:

$\mathrm{NaCl_{(s)} \xrightarrow{H_2O} Na^+_{(aq)} + Cl^-_{(aq)}}$

$\mathrm{Ca(NO_3)_2_{(s)} \xrightarrow{H_2O} Ca^{2+}_{(aq)} + 2NO_3^-_{(aq)}}$

Calculating Ion Concentrations in Strong Electrolyte Solutions

• Since strong electrolytes dissociate completely, the concentration of each ion equals the stoichiometric amount from the formula.

• Example 1: In 0.050 M $\mathrm{HNO_3}$, $[H_3O^+] = 0.050$ M and $[NO_3^-] = 0.050$ M.

• Example 2: In 0.020 M $\mathrm{Sr(OH)_2}$, $[Sr^{2+}] = 0.020$ M and $[OH^-] = 2 \times 0.020 = 0.040$ M.

The Autoionization of Water

Equilibrium in Pure Water

• Water undergoes a very slight autoionization:

$\mathrm{2H_2O_{(l)} \rightleftharpoons H_3O^+_{(aq)} + OH^-_{(aq)}}$

• The equilibrium constant for this reaction is:

$K_c = [H_3O^+][OH^-]$

Ion-Product Constant for Water ($K_w$)

• At 25°C, experimental measurements show $[H_3O^+] = [OH^-] = 1.0 \times 10^{-7}$ M.

• Thus, the equilibrium constant is:

$K_c = [H_3O^+][OH^-] = (1.0 \times 10^{-7})(1.0 \times 10^{-7}) = 1.0 \times 10^{-14}$

• This constant is called the ion-product constant for water and is given the symbol $K_w$:

$K_w = [H_3O^+][OH^-] = 1.0 \times 10^{-14}$

Significance

• The value of $K_w$ is fundamental for understanding the pH and pOH scales, as well as acid-base equilibria in aqueous solutions.

• In pure water at 25°C, $[H_3O^+] = [OH^-] = 1.0 \times 10^{-7}$ M, which corresponds to a neutral pH.

Summary Table: Strong Electrolyte Dissociation

Type	Example	Dissociation Equation	Ion Concentrations (for 0.10 M solution)
Strong Acid	$\mathrm{HNO_3}$	$\mathrm{HNO_3 + H_2O \rightarrow H_3O^+ + NO_3^-}$	$[H_3O^+] = 0.10$ M, $[NO_3^-] = 0.10$ M
Strong Base	$\mathrm{Sr(OH)_2}$	$\mathrm{Sr(OH)_2 \rightarrow Sr^{2+} + 2OH^-}$	$[Sr^{2+}] = 0.10$ M, $[OH^-] = 0.20$ M
Strong Salt	$\mathrm{NaCl}$	$\mathrm{NaCl \rightarrow Na^+ + Cl^-}$	$[Na^+] = 0.10$ M, $[Cl^-] = 0.10$ M

Key Concepts

• Strong electrolytes dissociate completely, making calculation of ion concentrations straightforward.

• Autoionization of water is a fundamental equilibrium, with $K_w$ as its equilibrium constant.

• Understanding these principles is essential for further study of acid-base equilibria, pH calculations, and buffer systems.