Chapter 16 - Key Terms

16.1 Classifications of Acids and Bases

Arrhenius Acids: Substances that increase the concentration of hydrogen ions (H⁺) or hydronium ions (H₃O⁺) in aqueous solutions.

Arrhenius Bases: Substances that increase the concentration of hydroxide ions (OH⁻) in aqueous solutions.

Hydronium Ion (H₃O⁺): A water molecule (H₂O) that has accepted a proton (H⁺), forming H₃O⁺.

Bronsted-Lowry Acid: A substance that donates a proton (H⁺) to another substance.

Bronsted-Lowry Base: A substance that accepts a proton (H⁺) from another substance.

Amphiprotic Substance: A substance that can act as both a Bronsted-Lowry acid (proton donor) and a Bronsted-Lowry base (proton acceptor).

Lewis Acid: A substance that accepts a pair of electrons during a chemical reaction.

Lewis Base: A substance that donates a pair of electrons during a chemical reaction.

16.2 Conjugate Acid-Base Pairs

Conjugate Base: The species that remains after an acid has donated a proton (H⁺).

Conjugate Acid: The species that forms when a base accepts a proton (H⁺).

Conjugate Acid-Base Pair: Two species that differ by only one proton (H⁺); one is an acid and the other is its conjugate base (or vice versa).

16.3 The Autoionization of Water

Autoionization: The process by which water molecules spontaneously ionize to form equal concentrations of hydrogen ions (H⁺) and hydroxide ions (OH⁻).

Ion-product Constant: The product of the concentrations of H₃O⁺ and OH⁻ ions in pure water, typically equal to 1 × 10⁻¹⁴ at 25°C.

16.4 The pH Scale

pH: A measure of the acidity or basicity of a solution, defined as the negative logarithm of the hydrogen ion concentration: pH = -log[H⁺].

16.5 Strong Acids and Bases

The Strong Acids: Acids that completely dissociate in water, releasing all of their protons (H⁺). Examples include HCl, H₂SO₄, and HNO₃.

The Strong Bases: Bases that completely dissociate in water, releasing all of their hydroxide ions (OH⁻). Examples include NaOH and KOH.

16.6 Weak Acids

Acid Dissociation Constant (Kₐ): A measure of the strength of an acid in solution, defined as the equilibrium constant for the dissociation of the acid into its conjugate base and a proton.

Percent Ionization: The percentage of an acid or base that dissociates into ions in solution.

Polyprotic Acids: Acids that can donate more than one proton (H⁺) per molecule, such as sulfuric acid (H₂SO₄) or phosphoric acid (H₃PO₄).

16.7 Weak Bases

Amines: Organic compounds containing a nitrogen atom bonded to one or more alkyl or aryl groups, often acting as weak bases by accepting a proton (H⁺).

Base-dissociation Constant (Kb): A measure of the strength of a base in solution, defined as the equilibrium constant for the dissociation of the base into its conjugate acid and hydroxide ion (OH⁻).

16.8 Relationship Between Ka and Kb

Equation: The relationship between the acid dissociation constant (Kₐ) and the base dissociation constant (Kb) for conjugate acid-base pairs is given by the equation:
Kₐ × Kb = Kw (Kw is the ion-product constant for water (1 × 10⁻¹⁴ at 25°C)).

16.9 Acid-Base Properties of Salt Solutions

Hydrolysis: The reaction of a salt with water to form either an acidic or basic solution, depending on the nature of the salt’s constituent ions.

16.10 Acid-Base Behavior and Chemical Structure

Oxyacids: Acids that contain oxygen, hydrogen, and another element (usually a non-metal). The acidic properties depend on the number of oxygen atoms attached to the central atom.

Carboxylic Acids: Organic acids that contain a carboxyl group (-COOH). They are weak acids that dissociate to release a proton (H⁺) from the carboxyl group.