Acids and Bases Lecture Review

Vocabulary flashcards covering key definitions and concepts related to acids, bases, equilibrium, the pH scale, and buffers based on the lecture notes.

Arrhenius Acid

A substance that produces hydrogen ions (H+) when dissolved in water.

H+

Equivalent to a proton (hydrogen atom with no electron).

Hydrogen Chloride (HCl)

A common Arrhenius acid that ionizes in water to give H+ and Cl-.

Hydrochloric Acid

The chemical name for HCl.

Nitric Acid

The chemical name for HNO3.

Sulfuric Acid

The chemical name for H2SO4.

Acetic Acid

The chemical name for CH3COOH.

Arrhenius Base

A substance that produces hydroxide ions (OH-) when dissolved in water.

Sodium Hydroxide (NaOH)

A common Arrhenius base that dissolves in water to give Na+ and OH-.

Potassium Hydroxide

The chemical name for KOH.

Ammonia

NH3, a common Arrhenius base.

Bronsted-Lowry Acid

A substance that donates a proton (hydrogen ion, H+) in a chemical reaction.

Bronsted-Lowry Base

A substance that accepts a proton (hydrogen ion, H+) in a chemical reaction.

Proton Transfer

The process where a hydrogen ion (H+) moves between molecules or ions in acid-base reactions according to the Bronsted-Lowry definition.

Conjugate Acid-Base Pair

Molecules or ions that are related by the loss or gain of one H+.

Conjugate Acid

Formed when a Bronsted-Lowry base accepts one H+.

Conjugate Base

Formed when a Bronsted-Lowry acid loses one H+.

Hydronium Ion (H3O+)

Formed when a proton (H+) bonds to a water molecule.

Strong Acids

Acids that dissociate completely (100%) when dissolved in water, producing large amounts of H3O+ ions and anions; excellent proton donors.

Dissociate

To separate into ions when dissolved in a solvent.

Weak Acids

Acids that dissociate only slightly when dissolved in water, producing small amounts of H3O+ ions; poor proton donors.

Strong Bases

Bases that dissociate completely in water, producing large amounts of OH- ions; excellent proton acceptors.

Weak Bases

Bases that are poor acceptors of protons from water, producing small amounts of OH-.

Acid-Base Equilibrium

A state in acid-base reactions where the forward and reverse reactions proceed at equal rates, resulting in no net change in reactant and product concentrations.

Equilibrium

A dynamic state where the rates of the forward and reverse reactions are equal, leading to constant concentrations of reactants and products.

Le Chatelier's Principle

States that when a system at equilibrium is disturbed by a stress (e.g., change in concentration), the rates of the forward and reverse reactions will adjust to relieve the stress and re-establish equilibrium.

Amphoteric

A substance, like water, that can behave as both an acid and a base.

[H3O+]

Represents the concentration of hydronium ions, indicating the acidity of a solution.

[OH-]

Represents the concentration of hydroxide ions, indicating the basicity of a solution.

Ion-product constant for water (Kw)

The product of the concentrations of H3O+ and OH- ions in water, Kw = [H3O+][OH-], which is 1.0 x 10-14 at 25 °C.

Neutral Solution

An aqueous solution where the concentration of H3O+ is equal to the concentration of OH-.

Acidic Solution

An aqueous solution where the concentration of H3O+ is greater than the concentration of OH-.

Basic Solution

An aqueous solution where the concentration of OH- is greater than the concentration of H3O+.

pH Scale

A logarithmic scale from 0 to 14 that represents the H3O+ concentration of aqueous solutions.

pH

A measure of the acidity or basicity of a solution, calculated as -log[H3O+]; a pH < 7 is acidic, pH > 7 is basic, and pH = 7 is neutral.

Buffered Solution

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

Buffer Composition

Consists of an acid-base conjugate pair, typically a weak acid and a salt of its conjugate base, or a weak base and a salt of its conjugate acid.