1.3.1: Lewis Electron-Dot Diagrams The bonding between atoms in a molecule can be topically modeled though Lewis electron dot diagrams. Creating Lewis diagrams is rather simple and requires only a few steps and some accounting of the valence electrons on each atom. Valence electrons are represented as dots. When two electrons are paired (lone pairs), they are represented by two adjacent dots located on an atom, and when two paired electrons are shared between atoms (bonds), they are shown as lines. 1.3.1.1: Resonance 1.3.1.2: Breaking the octet rule with higher electron counts (hypervalent atoms) 1.3.1.3: Formal Charge 1.3.1.4: Lewis Fails to Predict Unusual Cases - Boron and Beryllium 1.3.2: 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. 1.3.2.1: Lone Pair Repulsion 1.3.2.2: Multiple Bonds 1.3.2.3: Electronegativity and Atomic Size Effects 1.3.2.4: Ligand Close Packing 1.3.3: 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. 1.3.4: 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. 1.3.5: Problems