Understanding the arrangement of valence electrons in atoms is the first step to predicting how they will bond.

Lewis Structures (Electron Dot Structures): Lewis structures are simple, yet powerful, diagrams that represent the valence electrons of atoms within a molecule or polyatomic ion. They provide a visual framework for understanding how electrons are shared in covalent bonds and where non-bonding (lone) pairs of electrons are located. The goal in drawing Lewis structures is generally to achieve stable electron configurations, typically an octet (eight valence electrons) for most atoms, and a duet (two valence electrons) for hydrogen.

Steps to Draw Lewis Structures:
1. Count the total number of valence electrons: Sum the valence electrons for all atoms in the molecule or ion. For polyatomic ions, add one electron for each negative charge and subtract one electron for each positive charge.
2. Determine the central atom: This is usually the least electronegative atom, as it is most likely to share electrons with multiple other atoms. Hydrogen can never be a central atom as it only forms one bond.
3. Draw single bonds: Connect the central atom to all the outer (terminal) atoms with a single covalent bond. Each single bond uses two valence electrons.
4. Complete octets on outer atoms: Distribute the remaining valence electrons as lone pairs on the outer atoms until each outer atom (except hydrogen, which needs only a duet) has an octet.
5. Place remaining electrons on the central atom: If any valence electrons are still unaccounted for, place them as lone pairs on the central atom.
6. Form multiple bonds if necessary: If, after placing all electrons, the central atom does not have an octet, convert one or more lone pairs from an adjacent outer atom into double or triple bonds between the central and outer atoms. This increases the number of shared electrons to satisfy the central atom's octet.

VSEPR Theory (Valence Shell Electron Pair Repulsion Theory): Once a Lewis structure is drawn, VSEPR theory provides a straightforward method for predicting the three-dimensional geometry (shape) of molecules and polyatomic ions. The fundamental principle of VSEPR theory is that electron domains around a central atom will arrange themselves as far apart as possible in three-dimensional space to minimize repulsion between them.

Common Electron Domain Geometries (and corresponding Molecular Geometries when no lone pairs are present):
- 2 Electron Domains: The electron domains arrange linearly, resulting in a linear molecular geometry (e.g., carbon dioxide, CO2; beryllium chloride, BeCl2). Bond angle is 180°.
- 3 Electron Domains: The electron domains arrange in a trigonal planar fashion, leading to a trigonal planar molecular geometry (e.g., boron trifluoride, BF3; sulfur trioxide, SO3). Bond angle is 120°.
- 4 Electron Domains: The electron domains arrange tetrahedrally, giving a tetrahedral molecular geometry (e.g., methane, CH4; silicon tetrachloride, SiCl4). Ideal bond angle is 109.5°.
- 5 Electron Domains (HL): The electron domains arrange in a trigonal bipyramidal pattern, resulting in a trigonal bipyramidal molecular geometry (e.g., phosphorus pentachloride, PCl5). This geometry has two distinct positions: axial and equatorial, with bond angles of 90° and 120°.
- 6 Electron Domains (HL): The electron domains arrange octahedrally, leading to an octahedral molecular geometry (e.g., sulfur hexafluoride, SF6). Bond angles are 90°.

Predicting Molecular Geometry with Lone Pairs: Lone pairs of electrons occupy more space around the central atom than bonding pairs because their electron density is concentrated closer to the nucleus of the central atom and is not shared between two nuclei. Consequently, lone pairs exert stronger repulsive forces on other electron domains. This increased repulsion distorts the ideal bond angles predicted by the electron domain geometry, leading to distinct molecular geometries.

Examples with 4 electron domains (illustrating lone pair effects):
- Methane (CH4): The central carbon atom has 4 bonding pairs and 0 lone pairs. Both electron domain and molecular geometry are tetrahedral with ideal bond angles of 109.5°.
- Ammonia (NH3): The central nitrogen atom has 3 bonding pairs and 1 lone pair. The electron domain geometry is tetrahedral, but the lone pair's greater repulsion pushes the three N-H bonding pairs closer together, resulting in a trigonal pyramidal molecular geometry with bond angles of approximately 107°.
- Water (H2O): The central oxygen atom has 2 bonding pairs and 2 lone pairs. The electron domain geometry is tetrahedral. The two lone pairs exert even stronger repulsive forces, pushing the two O-H bonding pairs even closer, leading to a bent or V-shaped molecular geometry with bond angles of approximately 104.5°.

Molecular Polarity: The overall polarity of a molecule is a critical property that influences its physical behavior, such as solubility and boiling point. Molecular polarity depends on two interconnected factors:
1. The polarity of individual bonds: Determined by the electronegativity difference between the bonded atoms. Polar bonds create individual bond dipoles.
2. The molecular geometry: The arrangement of these individual bond dipoles in three-dimensional space determines whether they cancel each other out or add up to create an overall molecular dipole moment. If a molecule is perfectly symmetrical and the bond dipoles cancel each other out, the molecule is non-polar (e.g., carbon tetrachloride (CCl4) and carbon dioxide (CO2)). If the bond dipoles do not cancel out due to an asymmetrical arrangement, the molecule is polar (e.g., water (H2O) and ammonia (NH3)).