Module 8: Chemical Bonding

Background

This module examines covalent bonding, which involves sharing electrons between non-metal elements. Organic chemistry and biochemistry study the molecules of life, most of which involve atoms covalently bonded to each other.

It has long been known that pure carbon occurs in different forms (allotropes) including graphite and diamonds. But it was not until 1985 that a new form of carbon was recognized: buckminsterfullerene, commonly known as a “buckyball.” This molecule was named after the architect and inventor R. Buckminster Fuller (1895–1983), whose signature architectural design was the geodesic dome, characterized by a lattice shell structure supporting a spherical surface. Experimental evidence revealed the formula, C60, and then scientists determined how 60 carbon atoms could form one symmetric, stable molecule. They were guided by bonding theory—the topic of this chapter—which explains how individual atoms connect to form more complex structures.

Three figures are shown. The left figure is a many-sides spherical ball composed of hexagonal rings which have carbon atoms at each corner. The center picture shows a soccer ball. The right picture shown as water tower with sides shaped like hexagonal rings.
Figure 1. Nicknamed “buckyballs,” buckminsterfullerene molecules (C60) contain only carbon atoms. Here they are shown in a ball-and-stick model (left). These molecules have single and double carbon-carbon bonds arranged to form a geometric framework of hexagons and pentagons, similar to the pattern on a soccer ball (center). This unconventional molecular structure is named after architect R. Buckminster Fuller, whose innovative designs combined simple geometric shapes to create large, strong structures such as this weather radar dome near Tucson, Arizona (right). (credit middle: modification of work by “Petey21”/Wikimedia Commons; credit right: modification of work by Bill Morrow)

Learning Objectives for Chemical Bonding

  1. Apply the concept of electronegativity to characterize a bond as covalent, ionic, or polar covalent.
  2. Employ the octet rule to draw Lewis structures.
  3. Recognize common exceptions to the octet rule (including radicals) and appraise their impacts on structure and reactivity.
  4. Draw resonance structures and evaluate their relative contributions to the resonance hybrid using formal charges and the octet rule.
  5. Relate trends in atom size and bond order to rationalize and support observed bond lengths and strengths.
  6. Use tabulated bond enthalpies to calculate approximate reaction enthalpies.

Why is this content important?

In the news:  Ozone (O3) occurs naturally in our atmosphere, and it has several important roles in our environment. Ozone levels can be affected by human activity. Check out this article from the Environmental Protection Agency about the health effects of ozone.

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Chem 103/104 Resource Book by crlandis and Chem 104 Textbook Team is licensed under a Creative Commons Attribution-NonCommercial-ShareAlike 4.0 International License, except where otherwise noted.