Organic Chemistry Chapter 3: Acids and Bases

Organic Chemistry Fifth Edition by David Klein - Chapter 3: Acids and Bases

3.1 Bronsted-Lowry Acids and Bases

  • Brønsted-Lowry Definition:
    • Acids donate a proton (H⁺).
    • Bases accept a proton (H⁺).
    • Conjugates:
    • A conjugate acid forms when a base accepts a proton.
    • A conjugate base forms when an acid donates a proton.
    • Visual Representation:
    • Conjugate Acid of Acid ↔ Conjugate Base of Base.

3.1 Conjugate Acids and Bases

  • Practice Exercise: Identify the acid, base, and conjugates in the reaction:
    • HCl + NaOH → NaCl + H₂O
    • Acid: HCl (donates proton)
    • Base: NaOH (accepts proton)
    • Conjugate Acid: H₂O (from NaOH)
    • Conjugate Base: Cl⁻ (from HCl)

3.2 Curved Arrows in Reactions

  • Introduction:

    • The making and breaking of bonds involves electron movement.
    • Curved Arrows:
    • Curved arrows show the flow of electron density.
    • Used similarly to drawing resonance structures, but indicate real electron movement.
    • Learning to draw mechanisms is a crucial skill in this course.

  • Specific Example:

    • Mechanism of an acid/base reaction:
    • The base attacks the acid using a pair of electrons.
    • Both the bond breaking and bond forming occur simultaneously in a single step.
    • IMPORTANT: Proton transfer reactions illustrate these concepts with multiple electron movements.

3.3 Quantifying Acidity

  • Introduction:

    • Understanding the difference between “strong” and “weak” acids and bases:
    • Strong acids/bases ionize completely in solution.
    • Weak acids/bases only partially ionize.
    • Using pKa values to compare strengths:
    • Quantitative Analysis: Strength of acids assessed by numerical values.
    • Qualitative Analysis: Stability of structures compared qualitatively.

  • Ka (Acid Dissociation Constant):

    • Measures an acid’s strength in water as the solvent.
    • Strong acids have a value of K_a > 1.

  • pKa:

    • Values range from 10−5010^{-50} to 101010^{10}.
    • The relationship is pK<em>a=−extlog(K</em>a)pK<em>a = - ext{log}(K</em>a).
    • This transforms large or small Ka values into manageable numbers (range from -10 to 50).
    • Lower pKa indicates a stronger acid.

  • Table 3.1 - pKa Values of Common Compounds:

    • Example: H₂SO₄ (pKa = -9) is 100 times stronger than HCl (pKa = -7).

3.4 Qualifying Acidity

  • Introduction to ARIO:

    • Stability Factors: Four main factors determine the stability of a conjugate base which affects acidity:
    • A: Type of atom - Larger atoms better stabilize the negative charge.
    • R: Resonance - Delocalization stabilizes lone pairs and increases acidity.
    • I: Induction - Electron-withdrawing groups stabilize negative charges.
    • O: Orbital hybridization - More s-character correlates with increased stability.

  • Atom:

    • Compare electronegativity and size when assessing stability.
    • Larger atoms like Br or I can stabilize negative charges better than smaller atoms like C or N.

  • Examples:

    • Assessing $ ext{C}$ versus $ ext{O}$ for stability based on electronegativity in the same period.

  • Resonance:

    • Discuss how delocalization spreads out charge and affects stability.

  • Induction:

    • Electron-withdrawing groups enhance stability of negative charge.

  • Orbitals:

    • Cationic versus anionic forms: sp² hybridized are more stable than sp³.

3.6 Predicting Equilibrium Position

  • Understanding equilibrium in acid/base reactions:

    • Predictions can be made based on:
    1. The pKa values of the acids involved (higher pKa = weaker acid favored).
    2. The relative stability of the conjugate bases (more stable base favored).
    • Cationic Acids: Stability assessed similarly; stable cationic acids will be less acidic.

  • Levels of Induction and Resonance:

    • The presence of nearby electronegative atoms affects the acidity of cationic acids as compared with anionic ones.

3.7 Leveling Effect

  • Water demonstrates a leveling effect on strong acids and bases:
    • Only acids weaker than hydronium (H₃O⁺) and bases weaker than hydroxide (OH⁻) can exist in aqueous solution without undergoing reactions with water.
  • Examples:
    • Sulfuric acid in water results in significant protonation of water, limiting its availability as a reagent.

3.9 Counterions

  • Counterions (Spectator Ions): Essential for charge balance in solutions but usually not shown in reactions.
    • Example Reaction: NaNH₂ + H₂O → NH₃ + NaOH

3.10 Lewis Acids and Bases

  • Definitions:
    • Lewis Acid: Accepts a pair of electrons.
    • Lewis Base: Donates a pair of electrons.
    • Brønsted-Lowry and Lewis definitions provide overlapping insights on acid/base behavior but focus on different aspects.

Practice Questions

  • Several practice exercises emphasize identification of acid/base strength, drawing mechanisms, and understanding equilibrium through pKa values.
    These exercises enhance comprehension and application of Chapter 3 concepts.