Chapter 1: Introduction

Structure and Bonding in Organic Chemistry (A Brief Review)

Atomic orbitals can combine to generate new orbitals:

Hybridization: combination of atomic orbitals of the same atom producing new orbitals of lower energy

sp, sp² and sp³ These orbitals explain the geometry of molecules

Molecular Orbitals:

Sigma bonding (σ): Hydrogen molecule

Sigma bond: is covalent (most common bond inorganic molecules)

2 other types of bonds in organic molecules:

Pi (π) bond (DOUBLE BOND)

Hydrogen-bond (H-bond) (will discuss more later)

All single bonds are sigma bonds and all double or triple bonds contain only one sigma bond

Sigma bonds can be formed from atomic orbitals, hybridized or native, or a combination

π-bond: cannot exists if a σ-bond is not already present

Result from the overlap of p orbitals of two atoms. The π-bond is always perpendicular to the sigma bond connecting the nuclei

Occur in sp² and sp hybridized atoms (double/triple bonds)

Double bond:

4 electrons in the bonding region between the nuclei

first pair forms the sigma bond

Second pair forms the pi bond

Triple Bond:

6 electrons in the bonding region between the nuclei

first pair forms the sigma bond

Second/third pairs form the pi bonds

Molecular shapes and 3-D drawings

Molecular shapes are related to hybridization of central atom sp³: only single bonds; 109.5^o angle (methane) sp²: double bond; 120^o angle (ethylene) sp: triple bond; 180^o angle (acetylene)

Only 4 general shapes are normally found in organic molecules:

Linear (sp hybridized atoms)

Trigonal planar (sp² hybridized)

Trigonal pyramid or tetrahedral: (sp³ hybridized)

Names of Organic Compounds

Names of Organic molecules are based on the names of Hydrocarbons.

The simplest hydrocarbons are the alkanes, all their names end by the suffix “ane” and is preceded by a prefix indicating the number of carbon atoms in the molecules.

Example:

Properties of Organic Molecules

Polarity of molecules Bond polarity can range from:

Non-polar covalent (C-C)

Ionic (NH₄⁺^-Cl)

Polar covalent (C-O)

Example:

the dipole moment of a bond is the measure of its polarity (in Debye: D)

The molecular dipole moment is the dipole moment of the molecule taken as a whole. It is a good indicator of the overall polarity of molecules.

Its value is the vector sum of individual bond dipole moments

Examples:

Attractions and Repulsions: 3 major kinds of attractive forces cause molecules to associate into solids or liquids:

dipole-dipole forces

polar molecules

London forces (van der Waals)

affect all molecules

Hydrogen bonds

for molecules with NH or OH groups

Dipole-dipole forces:

attractive intermolecular forces resulting from the attraction of the positive and negative ends of polar molecules

result: polar molecules have higher boiling/melting point

London dispersion forces:

In non-polar molecules (alkanes), this is the principal attractive force

They arise from temporary dipole moments created by the proximity of other molecules

Result: molecules with larger surface area will have a higher boiling/melting point

Hydrogen bonding:

Not a true bond, but a very strong form of dipole-dipole attraction. An H atom can participate in H-bonding if it is bonded to O or N (-OH, -NH)

The O-H and N-H bonds are strongly polarized and the positive charge is located on H

Hydrogen bonds form between the hydrogen and the lone pair of the heteroatom from another molecule:

Example: Methanol

result: the stronger the H-bond, the higher the boiling/melting point

alcohols (OH) have higher boiling points than amines (NH) of similar molecular weights, because the OH bond is more strongly polarized than NH

Medicinal Chemistry

Table of Contents for Medicinal Chemistry