Unit 3 Chem 2

Arrhenius Theory

Defines acids as substances that produce H^+ ions in water, and bases as substances that produce OH^- ions in water.

Brønsted-Lowry Theory

Defines acid as a proton (H+) donor and a base as a proton (H+) acceptor.

Lewis Theory

Defines acids as electron pair acceptors and bases as electron pair donors.

Amphoteric

A substance that can act as either an acid or a base depending on what it reacts with, e.g., water.

K_w

The fundamental equilibrium constant for water, defined by the autoionization process.

Strong Acids

Acids that ionize 100% in water, e.g., HCl, HNO3, H2SO4.

Weak Acids

Acids that ionize only partially (<5%) in water, often existing mostly as molecules.

Monoprotic Acids

Acids that yield exactly one H^+ ion per molecule, e.g., HCl.

Diprotic Acids

Acids that yield up to two H^+ ions per molecule in a stepwise manner.

pH Calculation from [H^+]

pH = -log[H^+], used to find pH from hydrogen ion concentration.

Weak Acid Approximation

Shortcuts used for weak acids where [H^+] is assumed to be equal to [A^-] and [HA] remaining is approximately equal to the initial concentration.

Buffer Solution

A solution that resists significant changes in pH when small amounts of acid or alkali are added.

Neutralization Reaction

A reaction between an acid and a base that results in the formation of water and a salt.

Titration Curve

A graph showing how the pH of a solution changes as a titrant is added during a titration.

K_{sp}

The solubility product constant that describes the equilibrium between a solid and its ions in a saturated solution.

Common Ion Effect

The decrease in solubility of an ionic compound due to the presence of a common ion.

Strong Base

A base that dissociates completely in water to yield OH^- ions.

Autoionization of Water

The process by which water dissociates into H^+ and OH^- ions, establishing K_w.