Structure of the Atom

The Nucleus

The atom consists of a dense, positively charged nucleus containing protons and neutrons, which accounts for most of the atom's mass. Surrounding the nucleus is a region of space occupied by electrons, which are much lighter and move rapidly in defined regions called orbitals.

• Protons: Positively charged particles in the nucleus.

• Neutrons: Neutral particles in the nucleus.

• Electrons: Negatively charged particles occupying the space around the nucleus.

The electron density is highest near the nucleus and decreases with distance.

Atomic Orbitals

Types and Shapes of Orbitals

Electrons occupy regions of space called atomic orbitals, each with a characteristic shape and energy.

• s orbital: Spherical in shape; can hold up to 2 electrons.

• p orbital: Dumbbell-shaped; three mutually perpendicular p orbitals (px, py, pz) per energy level, each holding up to 2 electrons.

• d orbital: More complex, cloverleaf shapes; five d orbitals per energy level starting from the third shell.

Each orbital can hold a maximum of two electrons with opposite spins.

Electron Configuration and the Aufbau Principle

Electrons fill atomic orbitals in order of increasing energy, following the Aufbau principle. The order is determined by the relative energies of the orbitals:

• 1s → 2s → 2p → 3s → 3p → 4s → 3d → 4p → 5s → 4d → 5p → 6s → 4f → 5d → 6p → 7s → 5f → 6d → 7p

The maximum number of electrons per shell is:

• 1st shell: 2 electrons (1s)

• 2nd shell: 8 electrons (2s, 2p)

• 3rd shell: 18 electrons (3s, 3p, 3d)

Hund's Rule and the Pauli Exclusion Principle also govern electron arrangement:

• Hund's Rule: Electrons occupy degenerate orbitals singly before pairing.

• Pauli Exclusion Principle: No two electrons in an atom can have the same set of four quantum numbers.

Elements in Organic Chemistry

Common Elements

Organic compounds primarily contain the following elements:

• Carbon (C)

• Hydrogen (H)

• Nitrogen (N)

• Oxygen (O)

• Phosphorus (P)

• Sulfur (S)

• Halogens: Fluorine (F), Chlorine (Cl), Bromine (Br), Iodine (I)

These elements are highlighted in the periodic table for their prevalence in organic molecules.

Chemical Bonding Theory

Valence Bond Theory

Valence Bond (VB) Theory explains covalent bond formation as the overlap of atomic orbitals from two atoms, resulting in a shared pair of electrons.

• Bond strength is maximized when orbitals overlap directly (head-on), forming a sigma (σ) bond.

• Sideways overlap of p orbitals forms a pi (π) bond, which is weaker than a sigma bond.

Example: In the H2 molecule, two 1s orbitals overlap to form a σ bond, releasing energy:

• Bond dissociation energy for H–H: 436 kJ/mol

• Bond length: 0.74 Å

Molecular Orbital Theory

Molecular Orbital (MO) Theory describes bonds as the combination of atomic orbitals to form molecular orbitals that are delocalized over the entire molecule.

• Bonding MO: Lower energy, electrons stabilize the molecule.

• Antibonding MO: Higher energy, electrons destabilize the molecule.

For H2:

• Bonding MO: Constructive overlap of 1s orbitals.

• Antibonding MO: Destructive overlap, higher energy, usually unfilled.

Hybridization and Molecular Structure

sp3 Hybridization: Methane (CH4)

In methane, carbon forms four equivalent bonds by hybridizing its 2s and three 2p orbitals to create four sp3 hybrid orbitals.

• Tetrahedral geometry

• Bond angles: 109.5°

sp3 Hybridization: Ethane (C2H6)

Each carbon in ethane is sp3 hybridized, forming a sigma bond between the two carbons and three sigma bonds to hydrogens.

• All bond angles approximately 109.5°

sp2 Hybridization: Ethylene (C2H4)

In ethylene, each carbon atom is sp2 hybridized, using three hybrid orbitals for sigma bonds and one unhybridized p orbital for the pi bond.

• Planar structure

• Bond angles: 120°

• Double bond consists of one sigma and one pi bond

sp Hybridization: Acetylene (C2H2)

Each carbon in acetylene is sp hybridized, forming a linear molecule with a triple bond (one sigma and two pi bonds).

• Bond angle: 180°

Hybridization of Nitrogen, Oxygen, Phosphorus, and Sulfur

These elements can also hybridize their orbitals to form bonds in organic molecules:

• Nitrogen: Typically sp3 (ammonia), sp2 (imines), or sp (nitriles)

• Oxygen: Typically sp3 (water, alcohols), sp2 (carbonyls)

• Phosphorus: Often sp3 (phosphines, phosphates)

• Sulfur: Can be sp3 (thiols, sulfides), sometimes expanded octet

Representing Chemical Structures

Lewis (Electron-Dot) and Line-Bond Structures

Chemical structures can be represented in several ways:

• Lewis structures: Show all valence electrons as dots.

• Line-bond (Kekulé) structures: Bonds are shown as lines; lone pairs may be omitted for clarity.

• Condensed structures: Atoms and bonds are written in a linear fashion, omitting some or all bonds.

• Skeletal structures: Only carbon skeleton and heteroatoms are shown; hydrogens on carbons are implied.

Example: Methane (CH4) can be represented as:

• Lewis: H:C:H (with dots for electrons)

• Line-bond: H–C–H

• Condensed: CH4

Bonding Patterns and Octet Rule

Atoms tend to form bonds to achieve a stable octet (eight valence electrons):

• Hydrogen: 1 bond

• Carbon: 4 bonds

• Nitrogen: 3 bonds, 1 lone pair

• Oxygen: 2 bonds, 2 lone pairs

• Halogens: 1 bond, 3 lone pairs

Examples of Organic Molecules

Representative Structures

Organic molecules can be simple or complex. Examples include:

• Cholesterol: A complex sterol with multiple rings and functional groups.

• Benzylpenicillin: An antibiotic with a β-lactam ring and aromatic group.

• Carvone: A terpene responsible for the odor of spearmint.

Table: Common Elements in Organic Compounds

Key Equations

• Electron configuration order (Aufbau principle):

• Bond energy (example for H2):

Additional info: Some context and explanations have been expanded for clarity and completeness, including the typical bonding patterns of elements and the nature of hybridization in common organic molecules.