Acid Base Equilibrium Full Topic Video

Acid-Base Equilibrium

• Definitions of Acids and Bases:

  • Arrhenius Definition:

    • Acid: A substance that increases the concentration of hydrogen ions (H⁺) when dissolved in water.

    • Base: A substance that increases the concentration of hydroxide ions (OH⁻) when dissolved in water.

  • Brønsted-Lowry Definition:

    • Acid: A proton donor (donates H⁺).

    • Base: A proton acceptor (accepts H⁺).

Examples of Acids and Bases

• Common Acids:

  • Hydrochloric Acid (HCl)

  • Nitric Acid (HNO₃)

  • Sulfuric Acid (H₂SO₄)

• Common Bases:

  • Sodium Hydroxide (NaOH)

  • Potassium Hydroxide (KOH)

  • Calcium Hydroxide (Ca(OH)₂)

Dissociation of Acids and Bases

• Hydrochloric Acid in Water:

  • HCl → H⁺ + Cl⁻

  • Produces hydronium ions (H₃O⁺): H⁺ + H₂O → H₃O⁺

• Proton Acceptance by Bases:

  • Example with Ammonia (NH₃):

    • NH₃ + H₂O ⇌ NH₄⁺ + OH⁻

    • NH₃ gains a proton, water loses a proton.

Conjugate Acid-Base Pairs

• Conjugate Acid:

  • Formed from a base after it accepts a proton.

• Conjugate Base:

  • Formed from an acid after it donates a proton.

• Example:

  • From NH₃, the conjugate acid is NH₄⁺; from H₂O (acting as an acid), the conjugate base is OH⁻.

Amphiprotic and Amphoteric Substances

• Amphiprotic:

  • A substance that can act as both an acid and a base (e.g., water).

• Amphoteric:

  • A substance that can exhibit both acidic and basic properties (also e.g., water).

Strong vs. Weak Acids and Bases

• Strong Acids:

  • Completely dissociate in water (e.g., HCl).

• Weak Acids:

  • Partially dissociate (e.g., acetic acid).

• Strong Bases:

  • Completely dissociate in water (e.g., NaOH).

• Weak Bases:

  • Partially dissociate (e.g., NH₃).

Autoionization of Water

• Chemical Reaction:

  • 2 H₂O ⇌ H₃O⁺ + OH⁻

• Both hydronium and hydroxide ions are present in pure water at equal concentrations (1.0 x 10^-7 M at 25°C).

• Ion Product Constant of Water (K_w):

  • K_w = [H₃O⁺][OH⁻] = 1.0 x 10^-14 at 25°C.

pH and pOH Calculations

• pH Definition:

  • pH = -log[H₃O⁺]

• pOH Definition:

  • pOH = -log[OH⁻]

• Relationship:

  • pH + pOH = 14 at 25°C.

Calculating pH from Ion Concentrations

• Given [H₃O⁺], calculate pH using the formula pH = -log[H₃O⁺].

• Calculate [OH⁻] using [OH⁻] = K_w / [H₃O⁺].

Practice Problems and Applications

• Practice identifying conjugate acid-base pairs and calculating pH or K_a/K_b using given concentrations and conditions.

• Focus on weak acids and bases including percent ionization for deeper understanding.

Equilibrium Constant Expressions

• For Weak Acid (HA):

  • HA ⇌ H⁺ + A⁻

  • K_a = [H⁺][A⁻] / [HA]

• For Weak Base (B):

  • B + H₂O ⇌ BH⁺ + OH⁻

  • K_b = [BH⁺][OH⁻] / [B]

Relationships between Acid and Base Constants

• K_a x K_b = K_w:

• Useful in determining the dissociation characteristics of weak acids and bases.

Neutralization Reactions

• Understand stoichiometry involved in neutralization reactions and the effect on pH calculations depending on the concentration of reactants.