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Lecture

Inorganic Chemistry, Gary L. Miessler, Donald A. Tarr Textbook Chapter 3

26 Pages
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Department
Chemistry
Course Code
CHEM 212
Professor
Richard Oakley

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Description
We now turn from the use of quantum mechanics and its description of the atom to an elementary description of molecules Although most of the discussion of bonding in this book uses the molecular orbital approach to chemical bonding simpler methods that provide approximate pictures of the overall shapes and polarities of molecules are also very useful This chapter provides an overview of Lewis dot structures valence shell electron pair repulsion VSEPR and related topics The molecular orbital descriptions 5 and later chapters but the of some of the same molecules are presented in Chapter ideas of this chapter provide a starting point for that more modern treatment General chemistry texts include discussions of most of these topics this chapter provides a re them recently view for those who have not used Ultimately any description of bonding must be consistent with experimental data on bond lengths bond angles and bond strengths Angles and distances are most fre quently determined by diffraction Xray crystallography electron diffraction neutron diffraction or spectroscopic microwave infrared methods For many molecules there is general agreement on the bonding although there are alternative ways to describe it For some others there is considerable difference of opinion on the best way to describe the bonding In this chapter and Chapter 5 we describe some useful qualitative ap proaches including some of the opposing views I E f 31 1 Lewis electrondot diagrams although very much oversimplified provide a goodstarting point for analyzing the bonding in molecules Credit for their initial use goes LEWIS ELECTRON to G N ewis an American chemist who contributed much to thermodynamics and DOT DIAGRAMS chemical bonding in the early years of the 20th century In Lewis diagrams bonds between two atoms exist when they share one or more pairs of electrons In addition some molecules have nonbonding pairs also called lone pairs of electrons on atoms G N Lewis J Am Chem Soc 191638762 Valence and the Structure of Atonzs and Molecules Chemical Catalogue Co New York 1923 5 1 Chapter 3 Simple Bonding Theory 52 These electrons contribute to the shape and reactivity of the molecule but do not directly bond the atoms together Most Lewis structures are based on the concept that eight valence electrons corresponding to s and p electrons outside the noble gas core form a particularly stable arrangement as in the noble gases with sZp6 configurations An exception is hydrogen which is stable with two valence electrons Also some molecules require more than eight electrons around a given central atom A more detailed approach to electrondot diagrams is presented in Appendix D Simple molecules such as water follow the octet rule in which eight electrons surround the oxygen atom The hydrogen atoms share two electrons each with the oxy gen forming the familiar picture with two bonds and two lone pairs Shared electrons are considered to contribute to the electron requirements of both atoms involved thus the electron pairs shared by H and 0 in the water molecule are counted toward both the 8electron requirement of oxygen and the 2electron require ment of hydrogen Some bonds are double bonds containing four electrons or triple bonds contain ing six electrons 31 1 RESONANCE In many molecules the choice of which atoms are connected by multiple bonds is arbi trary When several choices exist all of them should be drawn For example as shown in Figure 31 three drawings resonance structures of 03 are needed to show the double bond in each of the three possible C0 positions In fact experimental evi dence shows that all the C0 bonds are identical with bond lengths 129 pm be tween doublebond and singlebond distances 1 16 pm and 143 pm respectively none of the drawings alone is adequate to describe the molecular structure which is a combi nation of all three not an equilibriun between them This is called resonance to signi fy that there is more than one possible way in which the valence electrons can be placed in a Lewis structure Note that in resonance structures such as those shown for 03 in Figure 31 the electrons are drawn in different places but the atomic nuclei remain in fixed positions The species CO NOs and SO3 are isoelectronic have the same electronic structure Their Lewis diagrams are identical except for the identity of the central atom When a molecule has several resonance structures its overall electronic energy is lowered making it more stable Just as the energy levels of a particle in a box are low ered by making the box larger the electronic energy levels of the bonding electrons are lowered when the electrons can occupy a larger space The molecular orbital descrip tion of this effect is presented in Chapter 5 G 7 2 G 72 d d 7 7 FIGURE 31 Lewis Diagrams for c C c c01 90 o 00
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