Acid-Base Equilibria: Study Notes for General Chemistry

Chapter 16: Acid-Base Equilibria

16.1 Acids and Bases – A Brief Review

This section introduces the fundamental definitions of acids and bases, highlighting the differences between the Arrhenius, Brønsted-Lowry, and Lewis concepts. Understanding these definitions is essential for classifying substances and predicting their behavior in chemical reactions.

• Arrhenius Definition: Acids produce H+ (hydronium ions) in aqueous solution; bases produce OH- ions.

• Brønsted-Lowry Definition: Acids are proton (H+) donors; bases are proton acceptors. This definition is broader than Arrhenius and does not require water as a solvent.

• Lewis Definition: Acids are electron pair acceptors; bases are electron pair donors. This is the most general definition.

• Monoprotic acids: Acids that donate one proton per molecule (e.g., HCl, HNO3).

• Polyprotic acids: Acids that can donate more than one proton per molecule (e.g., H2SO4).

Example: HCl is an Arrhenius acid because it produces H+ in water; NH3 is a Brønsted-Lowry base because it accepts a proton.

16.2 Brønsted-Lowry Acids and Bases

The Brønsted-Lowry theory expands the concept of acids and bases to include reactions outside aqueous solutions. It introduces the idea of conjugate acid-base pairs.

• Acid: Substance that donates a proton (H+).

• Base: Substance that accepts a proton (H+).

• Conjugate Acid-Base Pair: Two species that differ by a single proton. The acid becomes its conjugate base after donating a proton, and the base becomes its conjugate acid after accepting a proton.

Example: In the reaction NH3 + H2O ⇌ NH4+ + OH-, NH3 is a base, H2O is an acid, NH4+ is the conjugate acid, and OH- is the conjugate base.

• Hydrolysis of water: $\mathrm{H_2O + H_2O \rightleftharpoons H_3O^+ + OH^-}$

Key Concept: The strength of an acid is inversely proportional to the strength of its conjugate base.

16.3 The Autoionization of Water

Water can act as both an acid and a base, undergoing autoionization to produce hydronium and hydroxide ions. This process is fundamental to understanding pH and acid-base equilibria in aqueous solutions.

• Autoionization Reaction: $\mathrm{2H_2O \rightleftharpoons H_3O^+ + OH^-}$

• Ion Product Constant for Water (Kw): $K_w = [H_3O^+][OH^-] = 1.0 \times 10^{-14}$ at 25°C

• Neutral Solution: $[H_3O^+] = [OH^-] = 1.0 \times 10^{-7}$ M

• Acidic Solution: $[H_3O^+] > [OH^-]$

• Basic Solution: $[H_3O^+] < [OH^-]$

Example: If $[H_3O^+] = 1.0 \times 10^{-5}$ M, the solution is acidic.

16.4 The pH Scale

The pH scale is a logarithmic measure of the concentration of hydronium ions in solution. It is widely used to express the acidity or basicity of a solution.

• Definition: $\mathrm{pH = -\log[H_3O^+]}$

• pOH: $\mathrm{pOH = -\log[OH^-]}$

• Relationship: $\mathrm{pH + pOH = 14.00}$ at 25°C

• Neutral solution: pH = 7.00

• Acidic solution: pH < 7.00

• Basic solution: pH > 7.00

Example: If $[H_3O^+] = 1.0 \times 10^{-4}$ M, then pH = 4.00.

• Measuring pH: pH can be measured using indicators (color changes), pH paper, or electronic pH meters.

16.5 Strong Acids and Bases

Strong acids and bases dissociate completely in aqueous solution. This section lists common strong acids and bases and provides guidance for calculations involving their solutions.

• Strong Acids: HCl, HBr, HI, HNO3, HClO4, H2SO4 (first dissociation only)

• Strong Bases: Group 1A and 2A metal hydroxides (e.g., NaOH, KOH, Ca(OH)2)

• Calculation: For strong acids and bases, $[H_3O^+]$ or $[OH^-]$ equals the initial concentration of the acid or base.

Strong Acids	Strong Bases
HCl	NaOH
HBr	KOH
HI	LiOH
HNO3	Ca(OH)2
HClO4	Ba(OH)2
H2SO4 (1st step)	Sr(OH)2

Example: Calculate the pH of 0.01 M HCl: $\mathrm{pH = -\log(0.01) = 2.00}$

16.6 Weak Acids

Weak acids only partially dissociate in solution. Their strength is quantified by the acid dissociation constant, Ka.

• General Dissociation: $\mathrm{HA \rightleftharpoons H^+ + A^-}$

• Acid Dissociation Constant: $K_a = \frac{[H^+][A^-]}{[HA]}$

• Relative Strength: The larger the Ka, the stronger the acid.

Example: Acetic acid (CH3COOH) has $K_a = 1.8 \times 10^{-5}$, indicating it is a weak acid.

Summary Table: Acid and Base Strengths

Type	Examples	Dissociation
Strong Acid	HCl, HNO3	Complete
Weak Acid	CH3COOH	Partial
Strong Base	NaOH, KOH	Complete
Weak Base	NH3	Partial

Additional info: Some slides included practice problems and visual aids (e.g., pH indicators, pH meters) to reinforce concepts. For calculations, always use the appropriate formula and check units. The pH scale is logarithmic, so each unit change represents a tenfold change in [H+].