Introduction

class="introduction"

Valence Bond Theory
Hybrid Atomic Orbitals
Multiple Bonds
Molecular Orbital Theory

class="key-equations" title="Key-Equations"

class="summary" title="Chapter Summary"

class="exercises" title="Exercises"

class="references" title="References"

A pitcher is shown pouring liquid oxygen through the gap between two magnets, where it has formed a solid disk. A call out box near the stream of liquid oxygen shows an image of six pairs of spheres, spread apart from one another. Another call out box near the solid disk shows ten pairs of spheres much closer together.

We have examined the basic ideas of bonding, showing that atoms share electrons to form molecules with stable Lewis structures and that we can predict the shapes of those molecules by valence shell electron pair repulsion (VSEPR) theory. These ideas provide an important starting point for understanding chemical bonding. But these models sometimes fall short in their abilities to predict the behavior of real substances. How can we reconcile the geometries of s, p, and d atomic orbitals with molecular shapes that show angles like 120° and 109.5°? Furthermore, we know that electrons and magnetic behavior are related through electromagnetic fields. Both N₂ and O₂ have fairly similar Lewis structures that contain lone pairs of electrons.

Two Lewis diagrams are shown. The diagram on the left shows two nitrogen atoms, represented by the letter N connected by three lines and with a lone pair of electrons on each end of the structure. The diagram on the right shows two oxygen atoms, depicted by the letter O, connected by two lines. Two pairs of electrons surround each oxygen to the top and ends of the structure. Yet oxygen demonstrates very different magnetic behavior than nitrogen. We can pour liquid nitrogen through a magnetic field with no visible interactions, while liquid oxygen (shown in [link]) is attracted to the magnet and floats in the magnetic field. We need to understand the additional concepts of valence bond theory, orbital hybridization, and molecular orbital theory to understand these observations.

This work is licensed under a Creative Commons Attribution 4.0 International License.

You can also download for free at http://cnx.org/contents/85abf193-2bd2-4908-8563-90b8a7ac8df6@12.1

Attribution:

For questions regarding this license, please contact partners@openstaxcollege.org.
If you use this textbook as a bibliographic reference, then you should cite it as follows: OpenStax College, Chemistry. OpenStax CNX. http://cnx.org/contents/85abf193-2bd2-4908-8563-90b8a7ac8df6@12.1.
If you redistribute this textbook in a print format, then you must include on every physical page the following attribution: "Download for free at http://cnx.org/contents/85abf193-2bd2-4908-8563-90b8a7ac8df6@12.1."
If you redistribute part of this textbook, then you must retain in every digital format page view (including but not limited to EPUB, PDF, and HTML) and on every physical printed page the following attribution: "Download for free at http://cnx.org/contents/85abf193-2bd2-4908-8563-90b8a7ac8df6@12.1."