Module 4-2

Module Information

  • Prepared by: Analytical Chemistry Academic Group, Institute of Chemistry, College of Science, UP Diliman

  • Module Leads:

    • C Grace Gregorio, PhD

    • Lilibeth Coo, PhD

    • Rheo Lamorena-Lim, PhD

    • Florenda Valera, PhD

  • Topic: Acid-Base Equilibria

Learning Outcomes

  • Recall the concept of acids and bases

  • Understand Bronsted-Lowry theory (conjugate acid/base pairs)

  • Distinguish between strong and weak acids and bases

  • Recall equilibrium expressions for acid/base dissociation reactions

  • Derive pH and pOH from the autoprotolysis of water, Kw

  • Calculate the pH and pOH of strong and weak acids/bases

  • Determine α values for polyprotic acids

  • Explain the preparation of buffer solutions

  • Apply the Henderson-Hasselbalch equation for buffer pH calculations

  • Explain buffer action and buffer capacity

The Nature of Acids and Bases

Definitions

  • Acid: Contains hydrogen and dissociates in water to yield hydronium ions (H3O+)

  • Base: Contains hydroxyl group and dissociates in water to yield hydroxide ions (OH–)

  • Neutralization Reaction: Reaction between H+ (H3O+) from acid and OH– from base, releasing energy (exothermic, ~56 kJ per mole)

Acid-Base Strength

  • Proton Acceptor/Donor:

    • Acids: Proton donors; e.g., HNO3, H2PO4. All Arrhenius acids are Bronsted-Lowry acids.

    • Bases: Proton acceptors; e.g., NH3, CO3 2-, F–, OH–. Not all Bronsted-Lowry bases are Arrhenius bases, but all Arrhenius bases contain OH–.

Bronsted-Lowry Theory

  • Highlights solvent role in acid-base behavior

  • Amphiprotic solvent (water) can act as both acid and base

  • Conjugate acid-base pairs (e.g., acid1-base1 and acid2-base2):

    • stronger acids have weaker conjugate bases and vice versa

Acid Strength Comparisons

  • Strong Acids: Completely dissociate in water

    • Equation Example: HA + H2O ↔ H3O+ + A-

  • Weak Acids: Slightly dissociate in water

    • Equation Example: HA + H2O ⇌ H3O+ + A-

Equilibrium Constant Expressions

Strong vs. Weak Acids

  • Strong Acids:

    • Qc = [H3O+][A-]/[HA] at equilibrium, Ka >> 1

  • Weak Acids:

    • Qc = [H3O+][A-]/[HA] at equilibrium, Ka << 1

Autoprotolysis of Water

  • Reaction: H2O + H2O ⇌ H3O+ + OH-

  • Ion-product constant (Kw = [H3O+][OH-] = 1.0 x 10^-14 at 25°C)

Calculation of pH

Strong Acids and Bases

  1. When strong acid/base is the major supplier of H3O+:

    • pH = -log[H3O+]

  2. When both acid/base and water contribute to H3O+:

    • Solve using Kw

  3. For weak acids/bases, equilibrium expressions are necessary to calculate [H3O+].

Weak Monoprotic Acids

  1. Strong concentrations can be derived if Ka × [HA] is significant compared to Kw.

  2. Quadratic formula might be necessary for calculating low concentrations.

Buffers and Henderson-Hasselbalch Equation

  • Back-ground: Weak acids and conjugate bases resist drastic pH changes.

  • Henderson-Hasselbalch Equation:

    • pH = pKa + log [A-]/[HA]

    • Rearrangement of the Ka equilibrium expression.

Buffer Solutions

  • Resulting solutions from weak acid and conjugate base combinations resist pH changes.

  • Limited strong acid/base additions are selectively neutralized.

Summary

  • Understanding acid-base equilibria is based on quantitative aspects influencing H3O+ concentration.

  • Ka and Kb values: Offer insight into acid/base strength.

  • Auto-protolysis of water is crucial in equilibria contexts.

  • The strength of conjugate bases is inversely related to their weak acid counterparts.

  • Buffers mixture resist abrupt pH changes, highlighting the importance of equilibrium principles.