Chemical Principles

  • Peter R. Bergethon

Abstract

We will now explore the principal methods for describing and calculating the electronic distribution in molecules. The importance of a knowledge of a biological molecule’s electron distribution cannot be overstated, for the electronic distribution in a molecule is responsible for its chemical properties. As we will see, the formation of chemical bonds with the consequent bond strengths, lengths, force constants, and angles defines the shape and reactivity of molecules. The polarity of bonds determines charge distribution and gives rise to the dipole moment of a molecule. The interaction of the electronic distribution with photons determines the optical character and spectral properties of molecules and the source of color and photochemistry. Many of the interactional forces on a large scale, including dipole and quadruple interactions, dispersion forces, and hydrogen bonding, all depend on the molecules’ electronic structure. Our appreciation of the chemistry and chemical biology of living organisms is built on an understanding of the electronic structure.

Keywords

Atomic Orbital Internuclear Distance Antibonding Orbital Hybrid Orbital Molecular Orbital Theory 
These keywords were added by machine and not by the authors. This process is experimental and the keywords may be updated as the learning algorithm improves.

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Further Reading

History of Chemistry

  1. Partington J. R. (1989) A Short History of Chemistry. Dover Publications, New York.Google Scholar

General

  1. Gillespie R. J., Spencer J. N. and Moog R. S. (1996a) An Approach to Reaction Thermodynamics through Enthalpies, Entropies and Free Energies of Atomization. J. Chem. Ed., 73: 631–37.CrossRefGoogle Scholar
  2. Gillespie R. J., Spencer J. N. and Moog R. S. (1996b) Bonding and Molecular Geometry Without Orbitals: The Electron Domain Model. J. Chem. Ed., 73: 622–27.CrossRefGoogle Scholar
  3. Gillespie R. J., Spencer J. N. and Moog R. S. (1996c) Electron Configurations from Experiment. J. Chem. Ed., 73: 617–22.CrossRefGoogle Scholar
  4. Spencer J. N., Moog R. S. and Gillespie R. J. (1996) Ionization Energies, Electronegativity, Polar Bonds and Partial Charges. J. Chem. Ed., 73: 627–31.CrossRefGoogle Scholar

Intermolecular Forces

  1. Campanario J. M., Bronchalo E., and Hidalgo M. A. (1994) An Effective Approach for Teaching Intermolecular Interactions. J. Chem. Ed., 71:761–66. Discusses the use of potential energy maps in understanding chemical properties and interactions. Examples are predominantly biochemical applications.Google Scholar
  2. Israelachvili J. (1992) Intermolecular and Surface Forces, 2d ed. Academic Press, London. The first section of this book provides a good exploration of the intermolecular forces in chemical systems.Google Scholar

Bonding

  1. Atkins P. W. and Friedman R. S. (1996) Molecular Quantum Mechanics, 3rd ed. Oxford University Press, New York.Google Scholar
  2. Gillespie R. J. (1992) Multiple Bonds and the VSEPR Model. J. Chem. Ed., 69: 116–20.CrossRefGoogle Scholar
  3. Pauling L. (1960) The Nature of the Chemical Bond and the Structure of Molecules and Crystals: An Introduction to Modern Structural Chemistry, 3rd ed. Cornell University Press, Ithaca, New York.Google Scholar
  4. Scheiner S. (1994) Ab Initio Studies of Hydrogen Bonds: The Water Dimer Paradigm. Annu. Rev. Phys. Chem., 45: 23–56.PubMedCrossRefGoogle Scholar

Biochemistry

  1. Stryer L. (1995) Biochemistry, 4th ed. W. H. Freeman, New York.Google Scholar

Copyright information

© Springer Science+Business Media New York 1998

Authors and Affiliations

  • Peter R. Bergethon
    • 1
  1. 1.Department of BiochemistryBoston University School of MedicineBostonUSA

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