Chemical Bonding

Scientists are constantly discovering new compounds and modifying their theories to explain the behaviors of these compounds.
Upon completion of this unit, you will:
- Understand the Kössel-Lewis approach to chemical bonding, which emphasizes that matter is composed of different types of elements.
- Explain the octet rule and its limitations.
- Draw Lewis structures of simple molecules.
- Describe various bonding types and theories:
  - Formation of different types of bonds.
  - Valence Shell Electron Pair Repulsion (VSEPR) theory.
  - Valence Bond (VB) theory and Molecular Orbital (MO) theory.
  - Concept of hydrogen bonding.
The attractive force between atoms that allows them to form compounds is called a chemical bond, and its exploration raises many questions:
- Why do atoms combine?
- Why are certain combinations possible?
- Why do some combinations result in definite shapes?

In 1916, Kössel and Lewis provided explanations based on the inertness of noble gases and valence electrons.
Each element has a specific number of valence electrons, which determines its bonding behavior.
Formation of ions:
- Negative ions form by gaining electrons.
- Positive ions form by losing electrons, leading to stable noble gas electronic configurations.
Octet Rule: Atoms tend to achieve a stable configuration by having eight electrons in their outer shell.
- This observation is particularly applicable to noble gases.
During chemical reactions:
- Atoms may transfer or share electrons to achieve stable configurations as seen in NaCl (sodium chloride) formation.

Lewis Symbols: Visual representation of valence electrons in an atom, where each dot represents a valence electron.
Lewis symbols help illustrate how atoms bond to form molecules.
For example:
- Sodium (Na) transfers one electron to Chlorine (Cl) resulting in NaCl:
  - Lewis representation:
    - Na → Na⁺ + e⁻
    - Cl + e⁻ → Cl⁻
The electrostatic attraction between ions forms an ionic bond, termed an electrovalent bond.

The octet rule is helpful but not universally applicable:
- Some atoms can have an incomplete octet.
- Examples include atoms like BOr and elements that can expand their octet by utilizing d-orbitals.
Some elements demonstrate odd-electron configurations, leading to unpaired electrons in cases like nitrogen oxides.

The Valence Shell Electron Pair Repulsion (VSEPR) theory predicts the geometry of molecules based on repulsion between electron pairs surrounding a central atom.
Electron pairs can be bonded or non-bonded (lone pairs), influencing the molecular shape.

Covalent Bonds: Formed by sharing electrons between atoms, often represented using Lewis structures.
Hybridization: The mixing of atomic orbitals to form new hybrid orbitals that better describe bonding:
- Types of Hybridization:
  - sp hybridization: Linear molecules
  - sp2 hybridization: Trigonal planar molecules
  - sp3 hybridization: Tetrahedral molecules
Each hybrid orbital holds either a lone pair of electrons or shared pairs with another atom.

Provides a more sophisticated view by combining atomic orbitals from different atoms to form molecular orbitals:
- Sigma (σ) Bonds: Formed by head-on overlap of orbitals.
- Pi (π) Bonds: Formed by side-to-side overlap of p orbitals.
Bond Order: Determines bond strength; it is defined as the difference between the number of bonding and antibonding electrons divided by two.

A special type of dipole-dipole attraction occurring when a hydrogen atom covalently bonded to an electronegative atom (e.g. N, O, F) exhibits an attraction with another electronegative atom.
Hydrogen bonds are crucial for determining the properties of compounds like water, influencing boiling points, solubility, and molecular structure.

An understanding of chemical bonding theories enables the interpretation of molecular structures and interactions between atoms. Through Lewis structures, VSEPR theory, and molecular orbital theory, students can comprehend how atoms combine to form stable compounds.