3: Simple Bonding Theory
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- 3.1: Lewis Electron-Dot Diagrams
- This page provides a detailed explanation of Lewis electron dot diagrams, introduced by Gilbert Lewis in 1916, which illustrate the bonding between atoms in a molecule. The text describes how valence electrons are represented as dots and outlines the octet rule where atoms form bonds to achieve eight electrons in their valence shell, except for some like H and He.
- 3.2: Valence Shell Electron-Pair Repulsion
- The Valence Shell Electron Repulsion (VSEPR) model can predict the structure of most molecules and polyatomic ions in which the central atom is a nonmetal; it also works for some structures in which the central atom is a metal. VSEPR builds on Lewis electron dot structures and together can predict the geometry of each atom in a molecule. The main idea of VSEPR theory is that pairs of electrons (in bonds and in lone pairs) repel each other.
- 3.3: Molecular Polarity
- Dipole moments occur when there is a separation of charge. They can occur between two ions in an ionic bond or between atoms in a covalent bond; dipole moments arise from differences in electronegativity. The larger the difference in electronegativity, the larger the dipole moment. The distance between the charge separation is also a deciding factor into the size of the dipole moment. The dipole moment is a measure of the polarity of the molecule.
- 3.4: Hydrogen Bonding
- A hydrogen bond is an intermolecular force (IMF) that forms a special type of dipole-dipole attraction when a hydrogen atom bonded to a strongly electronegative atom exists in the vicinity of another electronegative atom with a lone pair of electrons. Hydrogen bonds are are generally stronger than ordinary dipole-dipole and dispersion forces, but weaker than true covalent and ionic bonds.
