Chapter 17: Acids & Bases – General Chemistry Study Notes

Acids & Bases

Introduction

Acids and bases are fundamental chemical species that play crucial roles in chemical reactions, biological systems, and industrial processes. Understanding their definitions, properties, and behaviors is essential for mastering general chemistry concepts.

Definitions of Acids & Bases

Arrhenius Definition

The Arrhenius definition is one of the earliest and simplest ways to classify acids and bases:

• Acid: Increases the concentration of hydrogen ions (H+) in aqueous solution.

• Base: Increases the concentration of hydroxide ions (OH-) in aqueous solution.

Example: HCl (hydrochloric acid) dissociates in water to produce H+ and Cl-; NaOH (sodium hydroxide) dissociates to produce Na+ and OH-.

Brønsted-Lowry Definition

The Brønsted-Lowry definition expands the concept of acids and bases beyond aqueous solutions:

• Acid: Proton (H+) donor.

• Base: Proton (H+) acceptor.

Example: In the reaction $\mathrm{HA} + \mathrm{H_2O} \rightleftharpoons \mathrm{A^-} + \mathrm{H_3O^+}$, HA donates a proton to water, forming A- and H3O+.

Lewis Definition

The Lewis definition is the most general and applies to a wide range of chemical reactions:

• Acid: Electron pair acceptor.

• Base: Electron pair donor.

Example: BF3 (boron trifluoride) is a Lewis acid because it can accept an electron pair; NH3 (ammonia) is a Lewis base because it can donate an electron pair.

Brønsted-Lowry Acid Reactions

General Acid Formula and Conjugate Pairs

Acids and bases interact in equilibrium reactions, forming conjugate acid-base pairs:

• General acid formula: HA

• Reaction: $\mathrm{HA} + \mathrm{H_2O} \rightleftharpoons \mathrm{A^-} + \mathrm{H_3O^+}$

• HA (acid) donates H+ to water, forming H3O+

• H2O (base) accepts H+ from HA

• A- is the conjugate base of acid HA

• H3O+ is the conjugate acid of base H2O

Acid Dissociation Constant ($K_a$)

Equilibrium Expression

The acid dissociation constant quantifies the extent to which an acid dissociates in water:

• General reaction: $\mathrm{HA} + \mathrm{H_2O} \rightleftharpoons \mathrm{A^-} + \mathrm{H_3O^+}$

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

• Equilibrium constant expression:

$K_a = \frac{[\mathrm{H^+}][\mathrm{A^-}]}{[\mathrm{HA}]}$

$K_a$ is called the acid dissociation constant.

Acid Strength

Relationship Between $K_a$ and Acid Strength

The strength of an acid is determined by its $K_a$ value:

• $K_a \gg 1$: Large amount of dissociation; strong acid.

• $K_a \ll 1$: Small amount of dissociation; weak acid.

Strong acids dissociate almost completely in water, while weak acids only partially dissociate.

Acid/Base Strength Table

Comparison of Common Acids and Their $K_a$ Values

The following table summarizes the strength of several acids based on their $K_a$ values:

Acid	HA	A-	Ka
Hydrochloric	HCl	Cl-	Very large
Sulfuric	H2SO4	HSO4-	Very large
Hydronium Ion	H3O+	H2O	1.0
Hydrofluoric	HF	F-	7.2 × 10-4
Acetic	CH3COOH	CH3COO-	1.8 × 10-5
Carbonic (1)	H2CO3	HCO3-	4.3 × 10-7
Ammonium Ion	NH4+	NH3	5.6 × 10-10
Carbonic (2)	HCO3-	CO32-	4.8 × 10-11
Water	H2O	OH-	1.0 × 10-14

Higher $K_a$ = Stronger Acid

Value of $K_a$

Interpreting $K_a$ in Acid-Base Equilibria

When comparing acids, the value of $K_a$ indicates relative strength:

• If HA is a stronger acid than H3O+, then $K_a > 1$ for HA.

• If HA is weaker, then $K_a < 1$.

This helps predict the direction of acid-base reactions and the position of equilibrium.

Auto-Ionization of Water

Water as Both Acid and Base

Water can act as both an acid and a base, undergoing auto-ionization:

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

• Alternatively: $\mathrm{H_2O} \rightleftharpoons \mathrm{H^+} + \mathrm{OH^-}$

• Equilibrium constant: $K_w = [\mathrm{H^+}][\mathrm{OH^-}]$

• At 25°C, $K_w = 1 \times 10^{-14}$

For any aqueous solution at 25°C, the product $[\mathrm{H^+}][\mathrm{OH^-}]$ is always $1 \times 10^{-14}$.

The pH Scale

Definition and Calculation

The pH scale is a logarithmic measure of hydrogen ion concentration, used to express acidity or basicity of a solution:

• pH = $-\log[\mathrm{H^+}]$

• pOH = $-\log[\mathrm{OH^-}]$

• pKa = $-\log K_a$

Example: For pure water at 25°C, $[\mathrm{H^+}] = 1.0 \times 10^{-7}$, so $\mathrm{pH} = 7$ (neutral solution).

Solutions with pH < 7 are acidic; pH > 7 are basic.

Summary Table: Acid/Base Strength

Acid	HA	A-	Ka
Hydrochloric	HCl	Cl-	Very large
Sulfuric	H2SO4	HSO4-	Very large
Hydronium Ion	H3O+	H2O	1.0
Hydrofluoric	HF	F-	7.2 × 10-4
Acetic	CH3COOH	CH3COO-	1.8 × 10-5
Carbonic (1)	H2CO3	HCO3-	4.3 × 10-7
Ammonium Ion	NH4+	NH3	5.6 × 10-10
Carbonic (2)	HCO3-	CO32-	4.8 × 10-11
Water	H2O	OH-	1.0 × 10-14

Higher $K_a$ = Stronger Acid

Additional info: The images provided show common household acids (e.g., vinegar, lemon juice) and bases (e.g., baking soda, ammonia), illustrating the practical relevance of acid-base chemistry in everyday life.