Chapter 4: Acids & Bases/Electron Flow – Organic Chemistry Study Notes

Chapter 4 – Acids & Bases/Electron Flow

Introduction

Acids and bases are foundational concepts in organic chemistry, influencing nearly every reaction mechanism. Understanding their definitions, strengths, and the structural factors that affect their behavior is essential for mastering organic reactions.

• Bronsted-Lowry and Lewis acid-base reactions are central to organic chemistry.

• Reviewing introductory chemistry concepts (Arrhenius, Bronsted-Lowry acids/bases) is recommended.

4.1 – Acids & Bases: Three Definitions

Definitions of Acids and Bases

There are three major definitions for acids and bases, each with unique criteria and applications in chemistry.

• Arrhenius Acid: Produces H+ ions in aqueous solution.

• Bronsted-Lowry Acid: Proton donor (gives up H+).

• Lewis Acid: Electron pair acceptor.

• Arrhenius Base: Produces OH- ions in aqueous solution.

• Bronsted-Lowry Base: Proton acceptor.

• Lewis Base: Electron pair donor.

Type	Definition	Example
Arrhenius Acid	Produces H+ in water	$\mathrm{HCl} \rightarrow \mathrm{H}^+ + \mathrm{Cl}^-$
Bronsted-Lowry Acid	Donates a proton	$\mathrm{HCl} + \mathrm{H}_2\mathrm{O} \rightarrow \mathrm{H}_3\mathrm{O}^+ + \mathrm{Cl}^-$
Lewis Acid	Accepts an electron pair	$\mathrm{H}^+ + :\mathrm{OH}^- \rightarrow \mathrm{H}_2\mathrm{O}$

Type	Definition	Example
Arrhenius Base	Produces OH- in water	$\mathrm{NaOH} \rightarrow \mathrm{Na}^+ + \mathrm{OH}^-$
Bronsted-Lowry Base	Accepts a proton	$\mathrm{OH}^- + \mathrm{H}^+ \rightarrow \mathrm{H}_2\mathrm{O}$
Lewis Base	Donates an electron pair	$:\mathrm{NH}_3 + \mathrm{H}^+ \rightarrow \mathrm{NH}_4^+$

4.2 – Brønsted-Lowry Acid-Base Reactions: A Detailed Analysis

Arrow-Pushing Formalism

Electron flow in acid-base reactions is depicted using curved arrows, which show the movement of electron pairs during bond formation and breaking.

• Arrows start at an electron pair (usually on the base) and point toward the acidic proton.

• Negatives (electron-rich) attack positives (electron-deficient).

• Mechanisms combine the half-reactions of acid and base.

Example:

• In the reaction $\mathrm{NH}_3 + \mathrm{HBr} \rightarrow \mathrm{NH}_4^+ + \mathrm{Br}^-$, the arrow starts at the lone pair on nitrogen and points to the hydrogen of HBr.

4.3 – Measuring Acid Strength

Acidity Constants and pKa

The strength of an acid is quantified by its acid dissociation constant ($K_a$) and its logarithmic counterpart, pKa.

• $K_a$ measures the tendency of an acid to donate a proton to water.

• Higher $K_a$ means a stronger acid.

• $\mathrm{p}K_a = -\log K_a$; lower pKa means a stronger acid.

• Weak acids have strong, unstable conjugate bases; strong acids have weak, stable conjugate bases.

Example: Hydrogen gas (H2) has a high pKa, indicating it is a weak acid and does not readily participate in acid-base reactions.

4.4 – What Makes an Acid Acidic or a Base Basic? Structural Effects on Acidity/Conjugate Base Stability

Structural Factors Affecting Acidity

Several molecular features influence the stability of conjugate bases and thus the acidity of their parent acids.

• Electronegativity: More electronegative atoms stabilize negative charge, increasing acidity.

• Atomic Size: Larger anions disperse negative charge over a greater volume, stabilizing the base and increasing acidity.

• Resonance: Delocalization of electrons stabilizes the conjugate base, increasing acidity.

• Inductive Effect: Electron-withdrawing groups (e.g., halogens) near the acidic hydrogen stabilize the conjugate base via induction, increasing acidity.

• Hybridization: Greater s-character (sp > sp2 > sp3) holds electrons closer to the nucleus, stabilizing the base and increasing acidity.

Order of Importance:

• Atomic size

• Resonance

• Electronegativity

• Induction

• Hybridization

Example: Trichloroacetic acid (with electron-withdrawing Cl groups) is much more acidic than acetic acid due to induction.

4.5 – Lewis Acid-Lewis Base Definition

Lewis Acids and Bases in Organic Chemistry

The Lewis definition focuses on electron pair transfer, explaining most organic reactions.

• Lewis base: Electron pair donor (nucleophile).

• Lewis acid: Electron pair acceptor (electrophile).

• Nucleophiles attack electrophiles, forming new bonds.

Example: Water ($:\mathrm{OH}_2$) acts as a Lewis base, donating an electron pair to a carbocation (Lewis acid).

When drawing mechanisms, always show all lone pairs and use curved arrows to indicate electron flow from base to acid.

Summary Table: Acid/Base Definitions

Definition	Acid	Base
Arrhenius	Produces H+ in water	Produces OH- in water
Bronsted-Lowry	Proton donor	Proton acceptor
Lewis	Electron pair acceptor	Electron pair donor

Key Equations

• $K_a = \frac{[\mathrm{A}^-][\mathrm{H}_3\mathrm{O}^+]}{[\mathrm{HA}]}$

• $\mathrm{p}K_a = -\log K_a$

Applications and Examples

• Predicting acid/base strength using pKa values.

• Using electron-pushing formalism to draw mechanisms.

• Analyzing molecular structure to determine acidity/basicity.

Additional info: These notes expand on the slide content by providing full definitions, equations, and context for each acid/base concept, as well as examples and tables for clarity.