Study Notes on Acids and Bases Definitions, Theories, and Applications

Definitions of Acids and Bases

• Acids: Substances that increase the concentration of H⁺ ions in water.
• Bases: Substances that increase the concentration of OH^- ions in water.

Conjugate Acids and Bases

• Conjugate Acid: A particle formed when a base gains a hydrogen ion (proton).
• Conjugate Base: A particle that remains after an acid donates a hydrogen ion.

Theories of Acids and Bases

There are three primary theories or definitions concerning acids and bases:

1. Arrhenius Theory
2. Brønsted-Lowry Theory
3. Lewis Theory

Definition #1: Arrhenius Theory (1903)

• Arrhenius Acid: A substance that, when added to water, increases the concentration of H⁺ ions.
• Arrhenius Base: A substance that, when added to water, increases the concentration of OH^- ions.

Key Points about Arrhenius Definition

• Arrhenius Acid Characteristics: Produces H⁺ in water.
• Arrhenius Base Characteristics: Produces OH^- in water.
• Limitations:
    - Restricted to aqueous solutions only.
    - Defines only one type of base (hydroxides).

Definition #2: Brønsted-Lowry Theory (1920’s)

• Brønsted-Lowry Acid: A proton donor; a substance that can give up a proton.
• Brønsted-Lowry Base: A proton acceptor; a substance that can accept a proton.

Important Concept

• A proton is essentially a hydrogen atom that has lost its electron.

Conjugate Acid-Base Pairs

• Conjugate Acid-Base Pair: Consists of two substances related by the loss or gain of a single hydrogen ion.
    - Example: For every acid that donates a proton, there corresponds a conjugate base that is formed.
    - Conjugate Acid: The particle formed when a base gains a hydrogen ion.
    - Conjugate Base: The particle that remains after an acid donates a hydrogen ion.
    - Example: In the ionization of water, OH^- is a conjugate base.

Examples of Reactions Involving Conjugate Acid-Base Pairs

1. NH₃ + HC₂H₃O₂ → NH₄⁺ + C₂H₃O₂^-
      - Here, NH₃ acts as a base because it accepts a proton from acetic acid, which acts as an acid.

2. H₂O + HC₂H₃O₂ → H₃O⁺ + C₂H₃O₂^-
      - Water acts as a base accepting a proton and forming H₃O⁺.

3. NH₃ + H₂O → NH₄⁺ + OH^-
      - NH₃ acts as a base and H₂O acts as an acid.

List of Common Acids

1. H₃PO₄ (Phosphoric acid)
2. HNO₃ (Nitric acid)
3. H₂CO₃ (Carbonic acid)
4. H₃PO₃ (Phosphorous acid)
5. HBr (Hydrobromic acid)
6. CH₃COOH (Acetic acid)

Strength of Acids and Bases

Identification of Weak or Strong Acids and Bases

• For each acid or base, determine the following:
    1. Hydrochloric acid (HCl):      - Weak/ Strong: [To be determined]
       - Ionization reaction: [To be determined]
       - 100% ionization or partial: [To be determined]
    2. Sulfuric acid (H₂SO₄):
       - Weak/ Strong: [To be determined]
       - Ionization reaction: [To be determined]      3. Ammonia (NH₃):
       - Weak/ Strong: [To be determined]
       - Ionization reaction: [To be determined]
    4. Acetic Acid (CH₃COOH):
       - Weak/ Strong: [To be determined]
       - Ionization reaction: [To be determined]

Conceptual Application

• Visualizing Acid and Base Dissolution:
    - Draw a particle diagram illustrating the dissolution of a strong acid (HA) and a weak acid (HB) in water.
    - Emphasize the difference in the degree of ionization of strong versus weak acids within the solution.