1. Molecular Geometries
1.1. A solid wedge stands for a bond that projects toward you a “hashed” wedge for one that points away from you, and a simple line for a bond that lies in the plane of the paper.
1.2. Valence shell electron-pair repulsion (VSEPR) model.
1.2.1. The VSEPR model rests on the idea that an electron pair, either a bonded pair or an unshared pair, associated with a particular atom will be as far away from the atom’s other electron pairs as possible.
1.2.1.1. angles of 109.5°, a value referred to as the tetrahedral angle.
1.2.1.2. The H—O—H angle in water (105°) and the H—N—H angles in ammonia (107°)
1.2.1.3. Trigonal planar molecule. These three bonded pairs are farthest apart when they are coplanar, giving bond angles of 120°.
1.2.1.4. A multiple bond (double or triple) is treated as a unit in the VSEPR model
1.3. Bonded pairs take up less space than unshared pairs. A bonded pair feels the attractive force of two nuclei and is held more tightly than an unshared pair localized on a single atom. Thus, repulsive forces increase in the order.
1.3.1. A linear arrangement of atoms in carbon dioxide allows the electrons in one double bond to be as far away as possible from the electrons in the other double bond.
2. Molecular Dipole Moments
2.1. The molecular dipole moment is the resultant of all of the individual bond dipole moments of a substance.
2.1.1. Some molecules, such as carbon dioxide, have polar bonds, but lack a dipole moment because their geometry causes the individual CO bond dipoles to cancel.
2.2. From the Lewis formula of a molecule, we can use electronegativity to tell us about the polarity of bonds and combine that with VSEPR to predict whether the molecule has a dipole moment.
3. Resonance and Curved Arrows
3.1. Sometimes more than one Lewis formula can be written for a molecule, especially if the molecule contains a double or triple bond.
3.1.1. Resonance concept: when two or more Lewis formulas that differ only in the distribution of electrons, no single Lewis formula is sufficient. The true structure is said to be a resonance hybrid of the various Lewis formulas.
3.2. Lewis formulas show electrons as being localized (shared in a bond or unshared belonging to a single atom).
3.2.1. In reality, electrons distribute themselves leading to their most stable arrangement. This means that a pair of electrons can be delocalized, or shared by several nuclei.
3.3. We use a double-headed arrow to signify resonance and read it to mean that the Lewis formulas shown contribute to, but do not separately describe, the electron distribution in the molecule.
3.3.1. Curved arrows to keep track of delocalized electrons they show the origin and destination of a pair of electrons.
3.3.2. Ozone, for example, has a single structure; it does not oscillate back and forth between two contributors