Brønsted–Lowry Theory Basics

Overview and Key Concepts

Brønsted–Lowry Theory: Broadens the traditional Arrhenius model by classifying:
- Acids as proton (H⁺) donors.
- Bases as proton (H⁺) acceptors.
- Applicable in aqueous and non-aqueous environments, enhancing its relevance.

Definitions and Chemical Equations

Acid: A species that donates a proton (H⁺).
Base: A species that accepts a proton (H⁺).

infoNote

Acid: A species that donates a proton (H⁺).
Base: A species that accepts a proton (H⁺).

Example Reactions:
- Proton release from HCl:
  $\mathrm{HCl} \rightarrow \mathrm{H}^+ + \mathrm{Cl}^-$
- Interaction between water and ammonia:
  $\mathrm{H}_2\mathrm{O} + \mathrm{NH}_3 \rightarrow \mathrm{NH}_4^+ + \mathrm{OH}^-$ $H_{2} O + NH_{3} \to NH_{4}^{+} + OH^{-}$
  - This illustrates ammonia acting as a base.

Conjugate Acid-Base Pairs

Concept and Importance:
- Formed when an acid donates a proton, and a base accepts it.
- Crucial for understanding acid-base reactions.
Examples:
- HCl/Cl⁻:
  - HCl loses a proton to form Cl⁻.
- NH₄⁺/NH₃:
  - NH₄⁺ donates a proton to become NH₃.

Illustration of proton transfer in HCl/Cl⁻ pair.

General proton exchange mechanism.

Dissociation Constants (Ka) and pKa

Overview

Dissociation Constant (Ka): Indicates the acid strength and degree of dissociation in water.
pKa: A logarithmic scale to express acid strength, where lower pKa values correspond to stronger acids.

Detailed Explanation

Ka: Definition and Explanation

Ka quantifies acid dissociation behaviour: $K_a = \frac{[\mathrm{H}^+][\mathrm{A}^-]}{[\mathrm{HA}]}$

infoNote

Strong Acids: High Ka, indicating complete dissociation (e.g., HCl).
Weak Acids: Low Ka, indicating partial dissociation (e.g., Acetic Acid).

pKa: Introduction and Calculation

pKa simplifies the expression of acid strength: $pK_a = -\log_{10}(K_a)$

Example Calculation:

For Ka = 1.0 × 10⁻⁵, calculate pKa:
- Solution:
  - pKa = -log₁₀(1.0 × 10⁻⁵)
  - pKa = 5.00

Flowchart explaining Ka and dissociation dynamics.

Calculate Ka and pKa:

Problem 1:
- Given [H⁺] = 0.01 M, [A⁻] = 0.01 M, [HA] = 0.98 M
- Calculate Ka:
$K_a = \frac{[0.01][0.01]}{0.98}$
- Solution:
  - Ka = (0.01 × 0.01) ÷ 0.98
  - Ka = 0.0001 ÷ 0.98
  - Ka = 1.02 × 10⁻⁴
  - pKa = -log₁₀(1.02 × 10⁻⁴) = 3.99

Steps from Ka to pKa calculations.

Visualising and Interpreting Titration Curves

Equivalence Points

Equivalence Point Identification:
- Characterised by a pronounced pH change, marking neutralisation.
- The nature of pH transitions, whether sharp or gradual, indicates acid/base strength.

Titration curves with equivalence points.

Neutralisation

Neutralisation: Occurs when an acid reacts with a base to produce water.
Equation: $\mathrm{H}^+ + \mathrm{OH}^- \rightarrow \mathrm{H}_2\mathrm{O}$

chatImportant

Key Reaction: Hydrogen ions combine with hydroxide ions to yield water.

Neutralisation reaction diagram.

Problem-Solving Strategies

Acid Strength vs. Concentration

Acid Strength: Determined by the degree of ionisation.
Concentration: Indicates the amount of solute present in a solution.

Strength vs. concentration diagram.

Common Pitfalls

chatImportant

Mistakes to Avoid:

Misinterpreting similar reactions.
Overlooking common ion effects that influence equilibria.

List of common exam mistakes in acid-base questions.

Addressing Common Misconceptions

Misconception: Equating stoichiometry with acid/base strength.
Utilise visual aids to demonstrate the true differences.

Conclusion

Grasping the Brønsted–Lowry Theory and dissociation constants like Ka and pKa is essential for:

Understanding chemical reactions thoroughly.
Designing effective buffer systems.
Executing precise titrations.

Brønsted–Lowry Theory Basics Simplified Revision Notes for SSCE HSC Chemistry