CH110

Chapter 10 Notes

In Chapter 2, we learned how to write chemical formulas and line structures...In Chapter 9, we learned how to show the locations of electrons in Lewis structures.....how can we determine the actual shape of a molecule (the molecular geometry) from these things?

Valence-Shell Electron-Pair Repulsion Theory (VSEPR Theory)

Electron pairs (bonded and lone pairs) will orient themselves so that they are as far apart from one another as possible (the four electrons in a double bond and the six in a triple bond are each considered one "group")

A lone pair takes up more space around an atom than a bonded pair

Electron group geometry (see Fig. 10.2)

2 electron pairs- linear

3 electron pairs- trigonal planar

4 electron pairs- tetrahedral

5 electron pairs- trigonal bipyramidal

6 electron pairs- octahedral

VSEPR notation

central atoms are denoted "A"

terminal atoms are denoted "X"

lone pairs are denoted "E"

Molecular geometries (see Table 10.1)

AX₂

AX₃

AX₄

AX₅

AX₆

The presence of lone pairs affects ONLY the SHAPE of the molecule (which is described in terms of locations of atoms), NOT the electron group geometry (sometimes referred to as the molecular geometry) (which takes into account lone pairs)

AX₂E vs AX₃

AX₃E vs AX₄

AX₂E₂ vs AX₄

AX₄E vs AX₅

AX₃E₂ vs AX₅

AX₂E₃ vs AX₅

AX₅E vs AX₆

AX₄E₂ vs AX₆

For molecules with more than one central atom, the geometry and resulting shape around each atom must be evaluated

Electronegativity revisited- polarity and dipole moment

Molecular dipoles

Molecular shape and dipoles

Bonds are formed by the overlap of orbitals

Hybridization of atomic orbitals

Often, the number of equivalent bonds around a central atom cannot be adequately explained using "conventional orbitals"

sp³ orbitals

sp² orbitals

sp orbitals

d hybrids

Hybrid orbitals and multiple bonds

Geometric isomers

Molecular Orbital Theory (MO Theory)

Bonding and antibonding orbitals

Aromatic compounds

Band Theory (section 24.5-24.7)