Chemical Bonding: Making and Breaking Bonds

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Chemical Bonding

Introduction

Chemical bonding is a foundational concept in organic chemistry, describing how atoms combine to form molecules. Understanding the nature of chemical bonds, their formation, and their breaking is essential for predicting molecular structure, reactivity, and properties.

Atomic Structure

Subatomic Particles and Electron Shells

Nucleus: Contains protons and neutrons; most of the atom's mass is concentrated here.
Electrons: Much smaller than protons/neutrons, occupy regions of space (orbitals) around the nucleus.
Quantum Mechanics: Electrons exist in quantized energy levels (shells), defined by principal quantum numbers (n = 1, 2, 3, ...).
Valence Shell: The outermost electron shell, crucial for chemical bonding.

Bonding & Non-Bonding Electrons

Valence Electrons and the Octet Rule

Valence Electrons: Electrons in the outermost shell, involved in bonding.
Lone Pairs: Pairs of electrons occupying the same orbital, not involved in bonding.
Octet Rule: Main group elements tend to gain, lose, or share electrons to achieve eight electrons in their valence shell, attaining noble gas configuration.
Example: Oxygen (O) has 6 valence electrons and tends to form two bonds to complete its octet.

Hypervalency

Expanded Octet and Main Group Elements

Hypervalency: The ability of some elements (typically in the 3rd period and beyond) to accommodate more than 8 electrons in their valence shell (expanded octet).
Example: Sulfur in SF6 can have 12 valence electrons.
Debate: The role of d-orbitals in hypervalency is debated; alternative models include ionic and 3-center-4-electron bonding.
Possible Electron Counts: 10, 12, or even 16 electrons in the valence shell for hypervalent compounds.

Hypervalency & Biological Systems

Biological Examples of Hypervalency

Phosphorus: Found in phosphate anions, ATP, and nucleic acids, often exhibits hypervalency.
Phosphate Groups: Key components in DNA, RNA, and energy transfer molecules like ATP.
Post-Translational Modifications (PTMs): Phosphorylation of proteins involves hypervalent phosphorus species.

Covalent Bonding

Formation and Theoretical Models

Covalent Bond: Formed when two atoms share a pair of electrons.
Born-Oppenheimer Approximation: Assumes nuclei are stationary while electrons move, simplifying calculations of molecular energy.
Molecular Potential Energy Curve: Shows how the energy of a system varies with bond length; minimum energy corresponds to bond formation.

Bond Formation Process

Unpaired valence electrons on two atoms approach each other.
Electron clouds overlap, attracted by nuclei, lowering potential energy.
At optimal distance, a stable molecule forms (e.g., H2).
If atoms get too close, repulsion increases energy, destabilizing the system.

Bonding Theories

Valence Bond Theory vs. Molecular Orbital Theory

Valence Bond Theory	Molecular Orbital Theory
Electrons localized between two atoms; bonds form via orbital overlap.	Atomic orbitals combine to form molecular orbitals (MOs) delocalized over the molecule.
Explains shape and reactivity well; cannot explain all phenomena (e.g., photochemical reactions).	Conceptually challenging; explains delocalization and reactivity; widely used in theoretical chemistry.

From Atomic Orbitals to Molecular Orbitals

Types of Molecular Orbitals

Sigma (σ) Orbitals: Formed by head-on overlap of atomic orbitals; strong, single bonds.
Pi (π) Orbitals: Formed by side-on overlap; present in double and triple bonds.
Bonding vs. Antibonding: Constructive interference forms bonding MOs; destructive interference forms antibonding MOs.

Polarity of Covalent Bonds

Electronegativity and Bond Classification

Electronegativity: Tendency of an atom to attract electrons in a bond.
Non-Polar (Homopolar) Bonds: Electronegativity difference < 0.5 (e.g., H-H).
Polar Bonds: Electronegativity difference between 0.5 and 2.0 (e.g., H-O).
Permanent Dipoles: Polar bonds can result in molecules with permanent dipole moments.

Ionic Bonds

Formation and Properties

Ionic Bond: Formed when electrons are transferred from a metal to a non-metal, creating cations and anions.
Electronegativity Difference: > 2.0 (e.g., NaCl).
Electrostatic Attraction: Ionic bonds are due to attraction between oppositely charged ions.

Dative Bonds

Coordinate Covalent Bonds

Dative (Coordinate) Bond: Both electrons in the shared pair come from the same atom.
Chemically Identical: Dative bonds are indistinguishable from regular covalent bonds once formed.
Biological Example: Haem group in hemoglobin binds oxygen via dative bonds.

Dative Bonds & Coordination Complexes

Coordination Number and Geometry

Coordination Complex: Central metal ion surrounded by ligands (molecules/ions donating electron pairs).
Coordination Number: Number of ligands attached to the metal; determines geometry (e.g., 2 = linear, 4 = tetrahedral, 6 = octahedral).
Kepert Model: Predicts shape of complexes, ignores non-bonding electrons.
Chelation: Ligands that form more than one bond to the metal are called chelators (e.g., EDTA, MDSA).

Double Bonds & Conjugation

Conjugation and Stability

Conjugation: Delocalization of electrons across adjacent double bonds, increasing stability.
Example: Benzene, where π electrons are delocalized over six carbon atoms.
Conjugated Systems: Double bonds separated by a single bond; electrons can move freely across the system.

Conjugation & Resonance

Resonance Structures

Resonance: Delocalization of electrons in molecules where bonding cannot be described by a single Lewis structure.
Resonance Hybrid: Actual structure is a hybrid of all possible resonance forms.
Example: Carbonate ion (CO32−) and benzene.

Resonance Forms in Nucleic Acids & Peptides

Biological Relevance of Resonance

DNA/RNA Bases: Guanine and cytosine exhibit resonance, contributing to their flat, aromatic structures.
Peptide Bond: Resonance between C=O and C-N bonds restricts rotation, giving proteins their structural rigidity.

Breaking Bonds

Homolytic and Heterolytic Fission

Homolytic Fission (Homolysis): Each atom retains one electron from the bond, generating radicals.
Bond-Dissociation Energy (BDE): Energy required to break a bond homolytically; measured in kJ/mol.
Heterolytic Fission (Heterolysis): Both electrons go to one atom, generating ions.
Electronegativity: Determines which atom receives electrons in heterolysis.

Summary Table: Types of Chemical Bonds

Bond Type	Electron Sharing	Example	Key Features
Covalent	Shared equally or unequally	H2, H2O	Strong, directional, forms molecules
Ionic	Transferred	NaCl	Electrostatic attraction, forms crystals
Dative (Coordinate)	Both electrons from one atom	NH4+	Common in metal complexes
Resonance/Conjugation	Delocalized	Benzene, peptide bond	Stabilizes structure, restricts rotation

Key Equations

Bond Dissociation Energy:

Electronegativity Difference: