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 Formal charge Formal charge assignment may be helpful is assigning a bonding arrangement of various possibilities.

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)

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

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

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)

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.