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A Review of Bonding Theory

A Review of Bonding Theory. Chapter 3. Lewis Electron-Dot. Valence electrons and the octet rule. Resonance Multiple resonance structures usually indicates a lower overall electronic energy. Expanded shells d -orbitals are made available to ‘expand’ the central atom to accept more electrons

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A Review of Bonding Theory

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  1. A Review of Bonding Theory Chapter 3

  2. Lewis Electron-Dot • Valence electrons and the octet rule. • Resonance • Multiple resonance structures usually indicates a lower overall electronic energy. • Expanded shells • d-orbitals are made available to ‘expand’ the central atom to accept more electrons • Formal charge • Formal charge assignment may be helpful is assigning a bonding arrangement of various possibilities.

  3. VSEPR • Molecules adopt geometries in which their valence electron pairs position themselves as far from each other as possible (Figure 3-8). • Provides shapes not a picture of bonding • Lone pair influence on geometry • CH4, NH3, and H2O • lp/lp>bp/lp>bp/bp • SF4 and ClF3 Angles involving lone pairs cannot be determined experimentally. • Multiple bonds • Double and triple bonds have slightly greater repulsive effects than single bonds (C2H4 and OIF4-1)

  4. Electronegative Effects on Geometry • Electronegativity is the measure of an atom’s ability to attract electrons from a neighboring atom to which it is bonded. • Fluorine is the highest (except for a couple of noble gases). Why? • Atoms with high electronegativities tend to draw electron density away from the center atom. This will allow a lone pair to reduce the bond angle. • PI3, PBr3, and PF3 The situation can be reversed if the center atom has a higher electronegativity. • H2O and H2S

  5. Electronegative Effects on Geometry • As the size of central atom increases, he other atoms can be at greater distances. This allows the lone pair to have a larger repulsive effect and decrease the bond angle. • PI3, AsI3, and SbI3

  6. Polar Molecules • Bonds resulting from two different atoms are termed as polar. • Partial negative and positive charges • =Qr • For a molecule polar bonds, vector addition is commonly used to determine the net dipole. • BCl3 and BCl2H. • Calculating the net dipole accurately, however, is more complex than simply adding vectors. • Spartan calculations (CH3Cl, CH2Cl, CHCl3, and CH4) • Effect of lone pairs (NH3 versus NF3, H2O (Spartan)

  7. Polar Molecules • Molecules with permanent dipoles are attracted to each other. Large attractive forces produce high melting and boiling points. • Nonpolar molecules also have intermolecular attractive forces due to momentary fluctuations in the electron density producing small temporary dipoles. The dipoles induce temporary dipoles in adjacent molecules. • Termed as London or dispersion forces.

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