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

Chapter 4 – Acids & Bases/Electron Flow

Introduction

The chemistry of acids and bases forms the foundation of organic and general chemistry. Understanding acid-base reactions is essential for interpreting many chemical processes, especially in organic chemistry. This chapter reviews key definitions and concepts, including Arrhenius, Brønsted-Lowry, and Lewis acids and bases, and explores how electron flow underpins acid-base mechanisms.

4.1 – Acids & Bases: Three Definitions

Definitions of Acids and Bases

There are three major definitions of acids and bases, each with its own criteria and applications:

• Arrhenius Acids and Bases: An Arrhenius acid increases the concentration of H+ ions in aqueous solution. An Arrhenius base increases the concentration of OH- ions in aqueous solution.

• Brønsted-Lowry Acids and Bases: A Brønsted-Lowry acid is a proton (H+) donor, while a Brønsted-Lowry base is a proton acceptor.

• Lewis Acids and Bases: A Lewis acid is an electron pair acceptor, and a Lewis base is an electron pair donor.

These definitions are summarized in the following tables:

Acid Type	Definition	Example
Arrhenius acid	Produces H+ in water	$\mathrm{HCl} \rightarrow \mathrm{H}^+ + \mathrm{Cl}^-$
Brønsted-Lowry acid	Proton donor	$\mathrm{HCl} + \mathrm{H}_2\mathrm{O} \rightarrow \mathrm{H}_3\mathrm{O}^+ + \mathrm{Cl}^-$
Lewis acid	Electron pair acceptor	$\mathrm{H}^+ + :\mathrm{OH}^- \rightarrow \mathrm{H}_2\mathrm{O}$

Base Type	Definition	Example
Arrhenius base	Produces OH- in water	$\mathrm{NaOH} \rightarrow \mathrm{Na}^+ + \mathrm{OH}^-$
Brønsted-Lowry base	Proton acceptor	$\mathrm{OH}^- + \mathrm{H}^+ \rightarrow \mathrm{H}_2\mathrm{O}$
Lewis base	Electron pair donor	$:\mathrm{NH}_3 + \mathrm{H}^+ \rightarrow \mathrm{NH}_4^+$

Example: Hydrochloric acid (HCl) can be classified as an Arrhenius, Brønsted-Lowry, or Lewis acid depending on the context of the reaction.

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

Arrow Pushing Formalism

Arrow pushing formalism (also called electron pushing or curved arrow formalism) is used to illustrate the movement of electrons during chemical reactions. In acid-base reactions, arrows always start at an electron pair (usually on the base) and point toward the acidic proton or electron-deficient atom.

• Arrows start at the electron pair of the base and point toward the acidic proton.

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

Example: In the reaction between ammonia and hydrogen bromide:

$\mathrm{NH}_3 + \mathrm{HBr} \rightarrow \mathrm{NH}_4^+ + \mathrm{Br}^-$

The arrow starts at the lone pair on nitrogen and points toward the hydrogen atom of HBr.

4.3 – Measuring Acid Strength

Acidity Scale: $K_a$ and $pK_a$

The strength of an acid is measured by its acid dissociation constant ($K_a$) or its logarithmic counterpart, $pK_a$:

• $K_a$ is a measure of an acid's willingness to donate a proton to water.

• The higher the $K_a$, the stronger the acid.

• $pK_a = -\log K_a$; the lower the $pK_a$, the stronger the acid.

Relationship between acid and base strength:

• Weak conjugate acids are paired with strong Brønsted-Lowry bases.

• Strong conjugate acids are paired with weak Brønsted-Lowry bases.

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

Base Strength and $pK_a$

• $pK_a$ values indirectly measure base strength.

• A weak acid has a strong, unstable conjugate base.

• A strong acid has a weak, stable conjugate base.

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

Structural Effects on Acidity

Several structural factors influence the acidity of a molecule by affecting the stability of its conjugate base:

• Electronegativity: Within a row of the periodic table, higher electronegativity stabilizes the conjugate base, increasing acidity.

• Atomic Size: Within a group, larger anions spread negative charge over a greater volume, stabilizing the conjugate base and increasing acidity.

• Resonance: Electron delocalization by resonance stabilizes the conjugate base, making the acid more acidic.

• Inductive Effect: Electron-withdrawing groups (such as halogens) near the acidic hydrogen stabilize the conjugate base via induction, increasing acidity. The closer the group, the greater the effect.

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

Order of importance: Atomic size > Resonance > Electronegativity > Induction > Hybridization

Example: Trichloroacetic acid is more acidic than acetic acid due to the inductive effect of three chlorine atoms stabilizing the conjugate base.

4.5 – Lewis Acid-Lewis Base Definition

Lewis Acids and Bases

The Lewis definition focuses on electron flow rather than proton transfer:

• Lewis base: Electron pair donor

• Lewis acid: Electron pair acceptor

Lewis acids are also called electrophiles (electron-loving), and Lewis bases are called nucleophiles (nucleus-loving).

Example: In the reaction $:NH_3 + BF_3 \rightarrow NH_3BF_3$, ammonia acts as the Lewis base (donor), and boron trifluoride acts as the Lewis acid (acceptor).

Drawing Lewis Acid-Base Mechanisms

• Identify bonds formed or broken.

• Determine electron redistribution.

• Always show lone pairs and use curved arrows to indicate electron movement.

Summary Table: Comparison of Acid/Base Definitions

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

Additional info: These concepts are foundational for understanding organic reaction mechanisms and predicting the behavior of acids and bases in chemical systems.