Printer Friendly

Stretching conceptions of chemical bonds.

Stretching conceptions of chemical bonds

No one has directly observed a chemical bond, so scientists who try to envision such bonds must rely on experimental clues and their own imaginations.

The models resulting from these mental exercises provide windows onto chemical phenomena that might otherwise go unnoticed, and they help scientists predict how molecules might behave without actually making them. But the prevailing models can also hinder scientists from recognizing concepts or phenomena that don't fit into them, contends Richard P. Messmer, a physicist at General Electric's Research and Development Center in Schenectady, N.Y.

In the Jan. 16 JOURNAL OF THE AMERICAN CHEMICAL SOCIETY, Messmer argues that an unconventional bonding scheme, known as the generalized valence bond (GVB) theory, gracefully accounts for a class of molecules that conventional theories can't portray without slipping in counterintuitive concepts and adjustments.

Sulfur trioxide ([SO.sub.3]) exemplifies these "hypervalent" molecules. The traditional picture of its bonds relies on the pre-1920 Lewis-Langmuir octet rule, which requires that eight electrons surround each of [SO.sub.3]'s four atoms. This scheme portrays the molecule as a set of three equivalent "resonant structures," each using a different oxygen atom to form a double bond with the sulfur atom while the other oxygen atoms form single bonds, for a total of four bonds.

In the 1930s, Linus Pauling developed the valence bond (VB) theory, reformulating the octet rule according to quantum mechanical principles. In this model, electrons occupy specific regions around atomic nuclei. When electrons on adjacent atoms pair up, they form a bond localized between the atoms. For hypervalent molecules, however, the VB portrait can become complicated and computationally cumbersome. Another widely applied model, called the molecular orbital (MO) theory, eases those computations, but only by including hard-to-visualize concepts such as bonding electronic orbitals that aren't confined between bonded atoms.

The GVB theory, developed more than 20 years ago by William A. Goddard III of Caltech in Pasadena, combines the computational ease of the MO theory with the conceptual clarity of the VB theory, Messmer says. In this scheme, [SO.sub.3]'s sulfur atom forms six equivalent bonds, two with each oxygen atom, yielding a unique structure in which each bond comprises a pair of localized orbitals. There's no need to invoke resonant structures or "delocalized" orbitals, Messmer says.

The theory has gained few followers so far, but Goddard says he suspects that will change as chemists show that it can resolve seeming anomalies such as hypervalent molecules and can yield useful predictions about other types of chemical behavior. Goddard, Messmer and others are now working to demonstrate just that.

"These [GVB concepts] really are extensions of old ideas," Messmer says. "By adding flexibility, we can describe things more simply. And that should allow people to think about molecules in a new way."
COPYRIGHT 1991 Science Service, Inc.
No portion of this article can be reproduced without the express written permission from the copyright holder.
Copyright 1991, Gale Group. All rights reserved. Gale Group is a Thomson Corporation Company.

Article Details
Printer friendly Cite/link Email Feedback
Author:Amato, Ivan
Publication:Science News
Date:Feb 2, 1991
Previous Article:Computer elevates Venus to new heights.
Next Article:Cancer war escalates to genetic weapons.

Related Articles
Coming 'round to old views of benzene.
Partners for a noble element.
Sculpting light to maneuver molecules.
Lasers offer surgical control over reactions.
Dynamic properties of rubber.
Breaking bonds reveals their strength.
Chemistry Nobel spotlights fast reactions.
Israel Bonds testimonial luncheon to recognize engineer Irwin Cantor.
Longest carbon-carbon bonds discovered.

Terms of use | Copyright © 2017 Farlex, Inc. | Feedback | For webmasters