Acids, Bases, and Acid-Base Equilibria in Organic Chemistry (Chapter 4)

Acids and Bases in Organic Chemistry

Arrhenius Acids and Bases

The Arrhenius definition is one of the earliest ways to classify acids and bases, focusing on their behavior in water.

• Arrhenius Acid: A substance that ionizes in water to produce H+ ions.

• Arrhenius Base: A substance that ionizes in water to produce OH- ions.

• H+ and H3O+ are used synonymously (H+ in water forms the hydronium ion).

$\mathrm{H}^+ (aq) + \mathrm{H}_2\mathrm{O}(l) \rightarrow \mathrm{H}_3\mathrm{O}^+ (aq)$

Brønsted-Lowry Acids and Bases

The Brønsted-Lowry definition broadens the concept of acids and bases beyond aqueous solutions.

• Brønsted-Lowry Acid: A substance that donates a proton (H+).

• Brønsted-Lowry Base: A substance that accepts a proton.

Conjugate Acid-Base Pairs

Acid-base reactions involve the transfer of a proton from an acid to a base, resulting in conjugate acid-base pairs.

• When an acid transfers a proton to a base, it becomes a conjugate base.

• When a base accepts a proton, it becomes a conjugate acid.

Example:

Acid	Base	Conjugate Base	Conjugate Acid
HCl	H2O	Cl-	H3O+

Acid-Base Mechanisms and Curved Arrows

Curved Arrow Notation

Curved arrows are used to show the movement of electrons during acid-base reactions.

• Arrows start at the electron source (lone pair or bond) and point to the electron acceptor.

• Correct use of arrows is essential for depicting proton transfers and electron flow.

• Never use arrows to indicate atom movement directly; always show electron movement.

Example: The reaction between methanol and hydroxide ion:

$\mathrm{CH}_3\mathrm{OH} + \mathrm{OH}^- \rightarrow \mathrm{CH}_3\mathrm{O}^- + \mathrm{H}_2\mathrm{O}$

Water as a Reactant

• Water is not always needed as a reactant in acid-base reactions.

• Other acids and bases can participate, such as acetic acid and ammonia.

Example:

Acid	Base	Conjugate Base	Conjugate Acid
CH3COOH	NH3	CH3COO-	NH4+

Resonance and Acidity

Resonance Structures in Acid-Base Chemistry

Resonance delocalization stabilizes conjugate bases, increasing acidity.

• When multiple acceptor sites exist, the site with greater delocalization is favored.

• Draw all possible resonance structures to determine the most stable form.

Example: Protonation of acetic acid at different oxygen atoms:

• Protonation at the carbonyl oxygen is favored due to greater delocalization.

π Electrons as Brønsted-Lowry Bases

π electrons in double and triple bonds can act as bases and accept protons from strong acids.

• Proton transfer to π bonds can generate carbocations.

• The stability of the resulting carbocation depends on its structure (e.g., secondary vs. primary).

Example: 2-butene reacts with HBr to form a sec-butyl cation.

Acid Strength and Equilibrium

Acid Dissociation Constant (Ka)

The strength of an acid is quantified by its acid dissociation constant, Ka.

$\mathrm{HA} + \mathrm{H}_2\mathrm{O} \rightleftharpoons \mathrm{A}^- + \mathrm{H}_3\mathrm{O}^+$

$K_a = \frac{[\mathrm{A}^-][\mathrm{H}_3\mathrm{O}^+]}{[\mathrm{HA}][\mathrm{H}_2\mathrm{O}]}$

Because [H2O] is large and nearly constant, the expression simplifies to:

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

pKa and Acid Strength

• pKa is the negative logarithm of Ka: $pK_a = -\log K_a$

• Larger pKa = weaker acid; smaller pKa = stronger acid.

• The stronger the acid, the weaker its conjugate base, and vice versa.

Example: Acetic acid has a pKa of 4.76, indicating it is a weak acid.

Table: Common Acids and Their pKa Values

The following table summarizes the pKa values of various acids and their conjugate bases:

Acid	pKa	Conjugate Base
Hydrochloric acid (HCl)	-7	Cl-
Acetic acid (CH3COOH)	4.76	CH3COO-
Ammonium ion (NH4+)	9.2	NH3
Water (H2O)	15.7	OH-
Methane (CH4)	50	CH3-

Additional info: Table values are representative; refer to your textbook for a comprehensive list.

Acid-Base Equilibria and Calculations

• At equilibrium, the stronger acid and base react to form the weaker acid and base.

• The position of equilibrium can be predicted using pKa values:

$\Delta pK_a = pK_{a,\text{product acid}} - pK_{a,\text{reactant acid}}$

$K_{eq} = 10^{\Delta pK_a}$

Example Calculation

Given: Ethanol, pKa = 15.9; Carbonic acid, pKa = 6.36. Which is the stronger acid?

• Carbonic acid is the stronger acid (lower pKa).

Thermochemistry and Mechanisms of Acid-Base Reactions

Thermochemistry examines the energy changes during acid-base reactions.

• A reaction mechanism shows the stepwise process of bond breaking and forming.

• The transition state is the highest energy point along the reaction coordinate.

• Gibbs free energy ($\Delta G$) relates to the equilibrium constant:

$\Delta G^\circ = -RT \ln K_{eq}$

• R = 8.31 J/(mol·K)

• Standard temperature and pressure: 298 K and 1 atm

Molecular Structure and Acidity

The structure of a molecule affects its acidity:

• Greater electronegativity of the atom bearing the negative charge increases acidity.

• Stabilization of the conjugate base (e.g., by resonance or inductive effects) increases acidity.

• Methanol forms the most stable anion among the examples shown, making it the strongest acid.

Summary Table: Key Concepts

Concept	Description
Arrhenius Acid/Base	Produces H+/OH- in water
Brønsted-Lowry Acid/Base	Proton donor/acceptor
Conjugate Acid-Base Pair	Acid and base differing by one proton
pKa	Measure of acid strength; lower pKa = stronger acid
Resonance	Delocalization stabilizes conjugate base, increases acidity
Curved Arrows	Show electron movement in mechanisms

Additional info: For practice, draw Lewis structures and use curved arrows to show electron flow in acid-base reactions. Refer to your textbook for more detailed mechanisms and examples.