Properties of Atoms & Bonds

Introduction

This study guide covers foundational concepts in organic chemistry related to atomic structure, bonding, and molecular geometry. Understanding these principles is essential for predicting molecular behavior and reactivity in organic compounds.

Atomic Structure and Lewis Structures

Charges and Lone Pairs on Lewis Structures

Lewis structures are diagrams that represent the bonding between atoms and the distribution of electrons in a molecule. Correctly assigning charges and lone pairs is crucial for understanding molecular properties.

• Common Lewis Structures: Carbon, nitrogen, oxygen, and hydrogen have typical bonding patterns and common formal charges.

• Lone Pairs: Non-bonding pairs of electrons are shown as dots around atoms.

• Formal Charge: The charge assigned to an atom in a molecule, calculated by:

• Example: In CH3OH, oxygen has a formal charge of 0; in CH3NH3+, nitrogen has a formal charge of +1.

Bonding and Electron Distribution

Where Do the Electrons Reside?

Electrons in molecules are distributed in bonds and lone pairs. The arrangement of electrons determines molecular shape and reactivity.

• Covalent Bonds: Electrons are shared between atoms.

• Molecular Geometry: The 3D arrangement of atoms is influenced by electron pairs.

Electronegativity and Bond Polarity

Electronegativity

Electronegativity is a measure of an atom's ability to attract shared electrons in a chemical bond. Differences in electronegativity lead to bond polarity.

• Polar Bonds: More electronegative atoms attract electron density, creating partial charges (δ+ and δ-).

• Example: In H–C–F, fluorine is more electronegative, resulting in a partial negative charge on F and a partial positive charge on H.

Electronegativity Differences and Bond Types

The difference in electronegativity (ΔEN) between two atoms determines the bond type:

• : Non-polar covalent (electrons shared equally)

• : Polar covalent (electrons shared unequally)

• : Ionic (electrons transferred)

• Example: O–C () is polar covalent; F–Li () is ionic.

Polar and Ionic Mixtures

Some compounds contain both polar covalent and ionic bonds.

• Example: Sodium acetate (Na+ CH3COO-) has an ionic bond between Na+ and acetate, and polar covalent bonds within the acetate ion.

• Example: CH3NH3Cl contains both covalent and ionic interactions.

Bond Formation and Orbital Theory

How Do Bonds Form?

Bonds form when atomic orbitals overlap, allowing electrons to be shared or transferred between atoms.

• Orbital Overlap: Covalent bonds result from the overlap of atomic orbitals.

• Molecular Orbitals: When atomic orbitals combine, they form molecular orbitals that can be bonding or antibonding.

Phases and Nodes in Orbitals

Atomic orbitals have phases, similar to wave peaks and troughs. Nodes are regions where the probability of finding an electron is zero.

• s Orbitals: Spherical, with no nodes.

• p Orbitals: Dumbbell-shaped, with a node at the nucleus.

Bonding and Antibonding Orbitals

When orbitals overlap in the same phase, constructive interference occurs, forming a bonding orbital. Opposite phases result in destructive interference, forming an antibonding orbital.

• Bonding Orbital: Lower energy, increased electron density between nuclei.

• Antibonding Orbital: Higher energy, decreased electron density between nuclei.

Types of Covalent Bonds

Sigma (σ) and Pi (π) Bonds

Covalent bonds are classified based on the type of orbital overlap:

• Sigma (σ) Bond: Formed by head-on overlap of orbitals (s-s, s-p, or p-p). All single bonds are sigma bonds.

• Pi (π) Bond: Formed by side-to-side overlap of p orbitals. Present in double and triple bonds.

• Single Bond: One sigma bond.

• Double Bond: One sigma bond and one pi bond.

• Triple Bond: One sigma bond and two pi bonds.

Hybridization and Molecular Geometry

Hybridization Theory

Hybridization explains the shapes of molecules by combining atomic orbitals into new hybrid orbitals.

• sp Hybridization: Linear geometry, 180° bond angles (e.g., C2H2).

• sp2 Hybridization: Trigonal planar geometry, 120° bond angles (e.g., C2H4).

• sp3 Hybridization: Tetrahedral geometry, 109.5° bond angles (e.g., CH4).

• VSEPR Theory: Valence Shell Electron Pair Repulsion theory predicts molecular shapes based on electron pair repulsion.

Bond Rotation and Saturation

The type of bond affects molecular rotation and saturation:

• Free Rotation: Sigma bonds allow free rotation around the bond axis.

• Restricted Rotation: Pi bonds restrict rotation, leading to cis/trans isomerism.

• Saturated Fats: Contain only single (sigma) bonds.

• Unsaturated Fats: Contain double or triple bonds (pi bonds present).

Summary Table: Bond Types and Properties

Additional info: Some context and terminology have been expanded for clarity and completeness, including explicit definitions, formulas, and examples relevant to college-level organic chemistry.