Lecture 20 Week 8

Lecture 20 Topics: Acid and Base

Properties and Definitions

  • Acids and Bases are fundamental concepts in chemistry, characterized by distinct physical and chemical properties.

Arrhenius Theory

  • The Arrhenius theory defines acids and bases in terms of their behavior in water.

    • Acids produce H+ ions (protons) in aqueous solutions.

    • Bases produce OH- ions (hydroxide ions) in aqueous solutions.

Bronsted-Lowry Theory

  • The Bronsted-Lowry definition expands upon the Arrhenius theory:

    • Acids are defined as proton (H+) donors.

    • Bases are defined as proton (H+) acceptors.

Strength and Dissociation

  • The strength of acids and bases is determined by their degree of dissociation in solution:

    • Strong acids and bases completely dissociate into their ions in solution (indicated by a single arrow in chemical equations).

    • Weak acids and bases only partially dissociate (indicated by an equilibrium arrow) and are characterized by their ionization constants (Ka).

Properties of Acids and Bases

  • Properties of Acids:

    • Sour Taste: Common characteristic; citric acid in lemons exemplifies this.

    • Ability to Dissolve Metals: Many acids can react with metals, producing hydrogen gas.

    • Neutralization of Bases: Acids can neutralize bases to form salt and water.

    • Litmus Test: Acidic solutions turn blue litmus paper red.

  • Properties of Bases:

    • Bitter Taste: Often associated with many household bases.

    • Slippery to the Touch: A property of many strong bases.

    • Neutralization of Acids: Bases can neutralize acids, also leading to the formation of salts and water.

    • Litmus Test: Basic solutions turn red litmus paper blue.

Amphoteric Substances

  • Amphoteric substances are capable of acting as both acids and bases, depending on the environment.

    • Example: Water (H2O) can either donate a proton or accept a proton, making it amphoteric.

Examples of Acids and Bases

  • Classification:

    • HCl: (Arrhenius acid, Bronsted-Lowry acid)

    • NaOH: (Arrhenius base, Bronsted-Lowry base)

    • Br-: (Neither)

    • H2O: (Both - amphoteric)

    • HSO4-: (Bronsted-Lowry acid)

    • NH3: (Bronsted-Lowry base)

Conjugate Acid-Base Pairs

  • In acid-base reactions, acids donate protons to bases, forming conjugate acid-base pairs:

    • Reactions:

      • NH3(aq) + H2O(l) → NH4+(aq) + OH−(aq)

        • Here, NH3 acts as a base and H2O as an acid.

      • H2SO4(aq) + H2O(l) → HSO4−(aq) + H3O+(aq)

        • H2SO4 is the acid and H2O is the base.

      • HCO3−(aq) + H2O(l) → H2CO3(aq) + OH−(aq)

        • HCO3− acts as a base.

Acid Dissociation and Ka

  • Dissociation in Aqueous Solutions:

    • Acids dissociate into H3O+ ions and their corresponding anions.

    • Strong Acids: Completely dissociate in solution (single arrow); examples include HCl, HBr, HI, HNO3, HClO4, and H2SO4.

    • Weak Acids: Only partially dissociate (equilibrium arrow) and are characterized by their Acid Ionization Constant (Ka), which reflects the strength of the acid.

Strong Acids and Bases

  • Strong Acids:

    • Hydrochloric Acid (HCl)

    • Hydrobromic Acid (HBr)

    • Hydroiodic Acid (HI)

    • Nitric Acid (HNO3)

    • Chloric Acid (HClO)

    • Perchloric Acid (HClO4)

    • Sulfuric Acid (H2SO4)

  • Strong Bases:

    • Lithium Hydroxide (LiOH)

    • Sodium Hydroxide (NaOH)

    • Potassium Hydroxide (KOH)

    • Rubidium Hydroxide (RbOH)

    • Cesium Hydroxide (CsOH)

    • Calcium Hydroxide (Ca(OH)2)

    • Strontium Hydroxide (Sr(OH)2)

    • Barium Hydroxide (Ba(OH)2)

Proton Count and Dissociation

  • Monoprotic, Diprotic, Triprotic Acids:

    • Monoprotic: Have one acidic proton (HA).

    • Diprotic: Have two acidic protons (H2A).

    • Triprotic: Can release three protons (H3A).

  • Dissociation Reactions:

    • Monoprotic: HA (aq) + H2O (l) → H3O+ (aq) + A− (aq)

    • Diprotic:

      • 1st: H2A (aq) + H2O (l) → H3O+ (aq) + HA− (aq)

      • 2nd: HA− (aq) + H2O (l) → H3O+ (aq) + A2− (aq)

    • Triprotic: General form with three successive dissociations.

Protonation Examples

  • Determine Protonation Type:

    • H3PO4: Triprotic

    • H2CO3: Diprotic

    • HF: Monoprotic

    • HC2H3O2: Monoprotic

    • H2SO4: Diprotic

    • HCN: Monoprotic

Acid Ionization Constant (Ka)

  • Weak Acid Ka Values and Expressions:

    • General form: HA (aq) + H2O (l) ⇌ H3O+ (aq) + A− (aq)

    • For phosphoric acid: Include reaction and the corresponding Ka expression related to its dissociation steps.