Acids, Bases, Functional Groups, and Intermolecular Forces in Organic Chemistry

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Acids and Bases; Functional Groups

Introduction

This section introduces foundational concepts in organic chemistry, focusing on acids, bases, and functional groups. These topics are essential for understanding molecular structure, reactivity, and the behavior of organic compounds in various environments.

Acids are substances that can donate a proton (H+).
Bases are substances that can accept a proton.
Functional groups are specific groups of atoms within molecules that are responsible for characteristic chemical reactions.
Examples: Citric acid and ascorbic acid (vitamin C) are shown as organic acids with carboxyl and hydroxyl functional groups.

Bond Dipole Moments

Definition and Measurement

Bond dipole moments arise from differences in electronegativity between atoms in a covalent bond, resulting in unequal sharing of electrons and partial charges.

Bond dipole moment (μ) quantifies the separation of charge in a bond.
Measured in debyes (D).
Depends on both the amount of charge separation and the bond length.
Equation:

Where q is the magnitude of charge separation and d is the bond length.

Examples of Bond Dipole Moments

CH3—NH2 (methylamine): Moderate dipole due to N's electronegativity.
CH3—OH (methanol): Stronger dipole due to O's higher electronegativity.
CH3—Cl (chloromethane): Even stronger dipole due to Cl.
CH3—CN (methyl cyanide): Highest dipole among examples due to triple bond and N's electronegativity.

Bond Dipole Moments Table

The following table summarizes dipole moments for common covalent bonds:

Bond	Dipole Moment (μ, D)
H—C	0.3
H—N	1.31
H—O	1.53
H—F	1.56
C—N	1.48
C—O	2.4
C—F	1.29
C≡N	3.6

Additional info: The table also references electronegativity values for common elements, which influence dipole moments.

Molecular Dipole Moment

Definition and Calculation

The molecular dipole moment is the vector sum of all individual bond dipole moments in a molecule. It determines the overall polarity of the molecule.

Depends on both bond polarity and bond angles (molecular geometry).
Lone pairs of electrons also contribute to the dipole moment.

Examples

CH2Cl2 (dichloromethane): μ = 1.9 D, polar due to asymmetric arrangement of Cl atoms.
CHCl3 (chloroform): μ = 1.0 D, less polar than dichloromethane.
CCl4 (carbon tetrachloride): μ = 0 D, nonpolar due to symmetric arrangement.
cis-1,2-dibromethene: μ = 1.9 D, polar due to geometry.

Additional info: Symmetry in molecular structure can lead to cancellation of bond dipoles, resulting in nonpolar molecules even if individual bonds are polar.

Intermolecular Forces

Types and Effects

Intermolecular forces are attractions between molecules that influence physical properties such as melting point, boiling point, and solubility.

London dispersion forces: Weak, temporary dipole-induced attractions present in all molecules, especially nonpolar ones.
Dipole–dipole forces: Attractions between polar molecules due to permanent dipoles.
Hydrogen bonding: Strong dipole–dipole attraction occurring in molecules with N—H or O—H groups.

Dipole–Dipole Forces

Result from the approach of polar molecules with positive and negative ends.
Attractive when opposite charges are near; repulsive when like charges are near.
In liquids and solids, molecules orient to maximize attractive interactions.

London Dispersion Forces

Arise from temporary dipoles due to electron movement.
Strength increases with molecular size and polarizability.
Main force in nonpolar molecules.

Effect of Branching on Boiling Point

Long-chain isomers (e.g., n-pentane) have greater surface area and higher boiling points.
Increased branching makes molecules more spherical, reducing surface area and lowering boiling point.
Neopentane (most branched) has the lowest boiling point among pentane isomers.

Hydrogen Bonding

Requires N—H or O—H groups.
Hydrogen from one molecule is attracted to a lone pair on O or N of another molecule.
O—H bonds are more polar than N—H, so alcohols have stronger hydrogen bonding than amines.

Boiling Points and Intermolecular Forces

Hydrogen bonding significantly increases boiling points.
Example: Methanol (CH3OH) b.p. = 78°C vs. dimethyl ether (CH3OCH3) b.p. = –25°C.
Alcohols generally have higher boiling points than amines due to stronger hydrogen bonding.

Polarity Effects on Solubility

General Principles

Solubility is governed by the principle "like dissolves like": polar solutes dissolve in polar solvents, and nonpolar solutes dissolve in nonpolar solvents.

Molecules with similar intermolecular forces mix well.
Polar solute in polar solvent: dissolves, energy released, entropy increases.
Polar solute in nonpolar solvent: does not dissolve; solvent cannot break solute's intermolecular forces.
Nonpolar solute in nonpolar solvent: dissolves due to weak attractions being overcome.
Nonpolar solute in polar solvent (e.g., water): does not dissolve; would disrupt hydrogen bonding network.