Acids, Bases, and Salt Solutions: Study Notes for General Chemistry I

Acids and Bases

Introduction to Acids and Bases

Acids and bases are fundamental chemical species that participate in a wide range of chemical reactions. Understanding their properties and behavior is essential for predicting the outcomes of reactions and the characteristics of solutions.

• Acids are substances that donate protons (H+) in aqueous solution.

• Bases are substances that accept protons or donate hydroxide ions (OH-).

• Strong acids and strong bases dissociate completely in water, while weak acids and weak bases only partially dissociate.

• Examples: HNO3 (strong acid), HF (weak acid), NaOH (strong base), NH3 (weak base)

Mixtures of Acids

When two acids are mixed in solution, their individual dissociation reactions contribute to the overall concentration of hydronium ions (H3O+), affecting the solution's pH.

• Key reactions: Each acid dissociates according to its own equilibrium constant (Ka).

• Example: Mixing 0.10 mol L-1 HNO3 (strong acid) with 0.10 mol L-1 HF (weak acid). HNO3 dissociates completely, while HF partially dissociates.

• Equilibrium constants: Ka for HF = $6.3 \times 10^{-4}$; Ka for HNO3 = $5.6 \times 10^{1}$

• Determining dominant contribution: The strong acid (HNO3) will contribute most to the H3O+ concentration.

Relationship Between Ka and Kb

The acid dissociation constant (Ka) and the base dissociation constant (Kb) are related for any acid-base conjugate pair.

• Acid dissociation: $HA(aq) + H_2O \rightarrow H_3O^+(aq) + A^-(aq)$

• Base dissociation: $A^-(aq) + H_2O \rightarrow HA(aq) + OH^-(aq)$

• Relationship: $K_a \times K_b = K_w$ where $K_w$ is the ion product of water ($1.0 \times 10^{-14}$ at 25°C).

• pK values: $pK_a + pK_b = 14$

pH of Salt Solutions

Classification of Salt Solutions

Aqueous salt solutions can be neutral, acidic, or basic depending on the nature of the ions produced upon dissolution.

• Neutral salts: Produced from strong acids and strong bases (e.g., NaCl).

• Acidic salts: Produced from strong acids and weak bases (e.g., NH4Cl).

• Basic salts: Produced from weak acids and strong bases (e.g., NaCH3COO).

Table: Classification of Cations and Anions in Salt Solutions

Cation - Non-reacting	Acidic	Basic
Li+, Na+, K+, Rb+, Cs+, Ba2+	NH4+, certain metal ions (Fe3+, Al3+, Cu2+)	NONE

Anion - Non-reacting	Acidic	Basic
Cl-, Br-, I-, NO3-, ClO4-	HSO4-, H2PO4-, anions of certain polyprotic organic acids	ALL THE REST!

Predicting pH of Salt Solutions

To determine whether a salt solution is acidic, basic, or neutral, analyze the acid-base properties of the ions formed.

• Example: NaCl forms Na+ and Cl-, both non-reacting, so the solution is neutral.

• Example: NH4Cl forms NH4+ (acidic) and Cl- (non-reacting), so the solution is acidic.

• Example: NaCH3COO forms Na+ (non-reacting) and CH3COO- (basic), so the solution is basic.

Calculating pH and Ionization

Steps to Calculate pH

Calculating the pH of a solution depends on the nature of the solute (strong acid/base, weak acid/base, or salt).

• Strong acid/base: Assume complete dissociation; use concentration directly.

• Weak acid/base: Use equilibrium expressions and Ka or Kb values.

• Salt solutions: Consider hydrolysis of ions and their effect on pH.

• General formula: $pH = -\log[H_3O^+]$

Percent Ionization

Percent ionization measures the fraction of acid molecules that ionize in solution.

• Formula: $\text{Percent Ionization} = \frac{[H_3O^+]_{eq}}{[HA]_{initial}} \times 100\%$

• Example: For a 2.5 mol L-1 HNO2 solution, use Ka and equilibrium calculations to find [H3O+].

Comparing Acid Strengths

The strength of an acid is determined by its Ka or pKa value. Lower pKa means a stronger acid.

• Example: Given three acids with different pKa values, the one with the lowest pKa will produce the most acidic solution at the same concentration.

Extra Practice Problems

Sample Calculations

• Calculate the pH of a solution of aspirin (weak acid) with Ka = $3.16 \times 10^{-4}$ and concentration 0.25 mol L-1.

• Calculate the pH of a solution of sodium benzoate (common food preservative) with pKa = 4.19 and concentration 0.25 mol L-1.

For each, set up the equilibrium expression, solve for [H3O+], and calculate pH.

Summary Table: Key Equations

Concept	Equation (LaTeX)
pH	$pH = -\log[H_3O^+]$
Relationship between Ka and Kb	$K_a \times K_b = K_w$
Percent Ionization	$\text{Percent Ionization} = \frac{[H_3O^+]_{eq}}{[HA]_{initial}} \times 100\%$
pKa and pKb	$pK_a + pK_b = 14$

Additional info: These notes expand on brief points and tables from the original file, providing definitions, examples, and formulas for clarity and completeness.