Lecture 10. Chemical bonding: Introduction
Tuesday 20 February 2024
The importance of molecular shape and polarity. Electronegativity and bond polarity. Periodic trends in electronegativity. Representing polar covalent bonds. Dipole moment, partial charges, and percent ionic character. Introduction to Lewis structures: First principles and simple examples. Lewis structures of diatomic and triatomic molecules. A look ahead to how Lewis structures will inform us about molecular properties of shape and polarity.
Reading: Tro NJ. Chemistry: Structure and Properties (3rd ed.) - Ch.5, §5.1-5.3 (pp.215-222)
Summary
Bond polarity can be thought of as arising from a difference in electronegativity between the bonded atoms (this difference in electronegativety is denoted as ΔEN). A continuum of bonding types exists between 100% ionic bonding and perfectly covalent bonding.
It is important to practice drawing Lewis structures for molecules or polyatomic ions, since these symbolic representations are the basis for prediction of the properties of substances made up of these molecules or polyatomic ions. Initially, we'll build Lewis structures based on a careful (and therefore correct) count of total number of valence electrons shared by the atoms comprising the molecule or ion, and the so-called "octet rule", which states that atoms tend to lose, gain, or share electrons in order to reach a noble gas electron configuration (i.e. a completely filled valence shell), which for n = 2 or n = 3 constitutes a filled valence shell of eight electrons (electron configuration of ns2np6).
As we refine our practice of drawing correct Lewis structures, we'll encounter cases in which a given formula can give rise to Lewis structures for two different skeletal structures (or ways to connect the atoms with bonds), or in which a Lewis structure for a given skeletal structure. is just one of several possible valid forms. We'll next develop the tool of formal charge and the idea of resonance forms to address and interpret such cases.