Ch 1

Review of Basic Bonding Concepts

Atomic Structure and Terminology

Understanding atomic structure is fundamental to organic chemistry, as it underpins bonding and molecular behavior. Key terms and concepts are defined below.

• Atom: The smallest unit of an element, defined by its number of protons.

• Atomic Number (Z): Number of protons in an atom.

• Mass Number (A): Number of protons plus number of neutrons.

• Isotopes: Atoms of the same element (same number of protons) but different numbers of neutrons.

Atomic and Subatomic Dimensions

Atoms are extremely small, with dimensions typically measured in femtometers (fm) and angstroms (Å).

• 1 Å = 100,000 fm

• Helium atom nucleus and electron cloud illustrated to show scale.

Electron Shells and Orbitals

Electrons occupy shells and subshells (orbitals) around the nucleus. The arrangement determines chemical properties.

	1st Shell	2nd Shell	3rd Shell
Orbitals	s	s, p	s, p, d
Number of Atomic Orbitals	1	1, 3	1, 3, 5
Maximum Number of Electrons	2	8	18

Electron Configurations

Electron configurations describe the arrangement of electrons in atomic orbitals. The following table summarizes configurations for selected elements:

Atomic Number	1s	2s	2px	2py	2pz	3s	3px	3py	3pz
1	1
5	2	2	1
6	2	2	1	1
8	2	2	2	1	1
9	2	2	2	2	1
12	2	2	2	2	2	2	2	2

Principles Governing Electron Configuration

• Aufbau Principle: Electrons fill the lowest energy orbitals first.

• Pauli Exclusion Principle: Each orbital can hold a maximum of two electrons with opposite spins.

• Hund's Rule: Electrons occupy empty orbitals of the same energy before pairing up.

Examples: Electron Configurations

• Boron (5 electrons): $1s^2 2s^2 2p^1$

• Carbon (6 electrons): $1s^2 2s^2 2p^2$

• Magnesium (12 electrons): $1s^2 2s^2 2p^6 3s^2$

Lewis Structures and Formal Charge

Lewis Structures

Lewis structures represent valence electrons as dots or lines, showing bonding and lone pairs.

• NH3 (Ammonia): Nitrogen has one lone pair and forms three single bonds with hydrogen.

• CH3OH (Methanol): Oxygen has two lone pairs and forms bonds with carbon and hydrogen.

Formal Charge Calculation

Formal charge helps identify the distribution of electrons in molecules and ions.

• Formula: $\text{Formal Charge} = \text{Number of valence electrons} - (\text{number of lone pair electrons} + \text{number of bonds})$

• Example (Hydronium ion, H3O+): $6 - (2 + 3) = +1$

• Example (Carbocation, CH3+): $4 - (0 + 3) = +1$ Carbocations are important intermediates in organic reactions.

Assigning Formal Charges

Formal charges are assigned to each atom in a molecule to determine electron distribution and reactivity.

• Positive formal charge (+1) indicates electron deficiency.

• Negative formal charge (-1) indicates electron excess.

Covalent Bonds and Electronegativity

Electronegativity Values

Electronegativity is the tendency of an atom to attract electrons in a bond. The following table summarizes values for common elements:

Element	Electronegativity
H	2.1
B	2.0
C	2.5
N	3.0
O	3.5
F	4.0
Al	1.5
Si	1.8
P	2.1
S	2.5
Cl	3.0
Br	2.8

Bond Polarity and Electronegativity Difference

• Nonpolar Covalent Bonds: Electronegativity difference ($\Delta E$) < 0.5. Electrons are shared equally. Example: H-H ($\Delta E = 0$)

• Polar Covalent Bonds: $0.5 \leq \Delta E \leq 1.9$. Electrons are shared unequally. Example: C-OH ($\Delta E = 1$)

• Ionic Bonds: $\Delta E > 1.9$. Electrons are transferred. Example: Na-Cl ($\Delta E = 2.1$)

Bonding in Organic Compounds

Bond Types and Hybridization

Organic molecules form covalent bonds through orbital overlap. Hybridization explains the geometry and bonding properties.

• sp3 Hybridization: Four equivalent orbitals, tetrahedral geometry, bond angle 109.5° (e.g., methane).

• sp2 Hybridization: Three equivalent orbitals, trigonal planar geometry, bond angle 120° (e.g., ethene).

• sp Hybridization: Two equivalent orbitals, linear geometry, bond angle 180° (e.g., acetylene).

Formation of Sigma and Pi Bonds

• Sigma (σ) Bond: Formed by head-on overlap of orbitals (e.g., sp3-s in methane).

• Pi (π) Bond: Formed by side-on overlap of unhybridized p orbitals (e.g., in double and triple bonds).

Bond Length and Strength

• Bond length decreases as s-character increases in hybrid orbitals.

• Greater s-character brings bonding electrons closer to the nucleus, increasing bond strength.

Structural Representations in Organic Chemistry

Types of Structural Formulas

• Condensed Structure: Shows atoms and their connectivity in a compact form (e.g., CH3CH2OH).

• Kekulé Structure: Shows all bonds explicitly, including lone pairs.

• Skeletal (Bond-Line) Structure: Simplified representation; vertices represent carbon atoms, hydrogens are implied, heteroatoms are shown explicitly.

Examples of Structural Representations

• Condensed: CH3CH2OH

• Kekulé: Explicit bonds and lone pairs drawn

• Skeletal: Lines and vertices, hydrogens omitted, heteroatoms shown

Summary Table: Key Concepts

Concept	Definition/Example
Atom	Smallest unit of an element; defined by number of protons
Isotope	Same number of protons, different number of neutrons
Electron Configuration	Arrangement of electrons in orbitals (e.g., $1s^2 2s^2 2p^2$ for carbon)
Formal Charge	$\text{FC} = \text{valence electrons} - (\text{lone pair electrons} + \text{bonds})$
Electronegativity	Tendency to attract electrons; F is most electronegative
Bond Polarity	Nonpolar ($\Delta E < 0.5$), Polar ($0.5 \leq \Delta E \leq 1.9$), Ionic ($\Delta E > 1.9$)
Hybridization	sp3 (tetrahedral), sp2 (trigonal planar), sp (linear)
Sigma Bond	Head-on orbital overlap
Pi Bond	Side-on p orbital overlap
Structural Formula	Condensed, Kekulé, Skeletal

Additional info: Some context and examples have been expanded for clarity and completeness, including explicit definitions and formulae for formal charge and hybridization.