Molecular Geometry and Bonding Theories

1. Molecular Geometry and Bonding Theories

2. Molecular Shapes • The overall shape of a molecule is determined by its bond angles. The bond angles and bond distances accurately define molecular geometry. • In carbon tetrachloride the molecule forms a pyramidal tetrahedron. With a carbon-chlorine bond distance of 1.78 A and a chlorine-carbon-chlorine bond angle of 109.5 degrees. • There are 5 basic shapes of molecules with the basic formula of Abn Linear, Trigonal Planar, Tetrahedral, Trigonal bipyramidal and Octahedral. (figure 9.3 page 343)

3. Why do we care?

4. VSEPR • The valence-shell eletron-pair repulsion (VSPER) model is used to explain the shapes of these basic molecules. • Remember that a single covalent bond is formed between two atoms when a pair of electrons occupies the space between the atoms. • A bonding pair of electrons thus defines a region in space where the electrons are most likely to be found. This region is called the electron domain. • A nonbonding pair of electrons is found most likely on one atom.

5. VSEPR (cont) • The VSEPR is based on the consequence of electron-electron repulsion. • The best arrangement of a given number of electron domains is the one that minimizes the repulsion between them. (Table 9.1 p 345) • The shapes of different ABn molecules depends upon the number of electron domains and the necessity to minimize the interaction between the domains. • The arrangement of electron domains about the central atom is called it electron domain geometry. • The molecular geometry is the arrangement of only the atoms, thus non-bonding pairs of electrons are not part of the description.

6. The effect of Non-bonding electrons and Multiple Bonds on Bond Angles • Consider the bond angles in the molecules CH4, NH3, H2O. • Based on what you know now what would the molecular geometry of each of these molecules look like. • The H-C-H angle in methane is 109.5, the H-N-H angle in ammonia is 107, the H-O-H angle in water is 104.5, why? • The non-bonding pair of electrons takes up more room than the bonding pair. Pushing the molecules closer together. (p 348)

7. Covalent Bonding and Orbital Overlap • The VSEPR model explains geometries of molecules but it fails to explain why bonds exist in the first place. • By combining Lewis notion of electron pair bonds and atomic orbitals we can derive the valence-bond model of molecular bonding. • This model develops the idea that valence orbitals from one atom overlap with the valence orbitals of another atom. The overlap of oribitals allows two electrons of opposite spin to share the common space between nuclei, forming a covalent bond.

8. Hybrid Orbitals • To produce certain geometries the central atom of a molecule will combine orbitals to create hybrid orbitals. • The common hybrids are sp, sp2, sp3,sp3d and sp3d2 • Example, BeF2 F is 1s22s22p5 indicating an unpaired electron in the 2p orbital of F • However Be is 1s22s2 in which all electrons are paired. • If Be had an electron promoted into the p-orbital it would then have two unpaired electrons to bond. (p 357-358)

9. Carbon • Carbon hybridizes its orbitals easier than any other atom which explains its vast array of molecular compounds. • Carbon very easily forms all three s and p orbital possibilities sp, sp2, sp3 which is why carbon (organic) chemistry is richer than the chemistry of all other elements.

10. Multiple Bonds • Sigma, σ, bonds are covalent bonds in which the electron density is concentrated along the internuclear axis. • It is formed by the overlap of two s orbitals as in H2, an s and p orbital as in HCl or two p orbitals as in Cl2. • Multiple bonds result from the overlap of two p orbitals oriented perpendicularly to the internuclear axis. • This results in a pi (π) bond which is a covalent bond in which the orbital overlaps lie above and below the internuclear axis.

11. Multiple bonds (cont) • Single bonds are always sigma bonds. • Double bonds contain one sigma bond and one pi bond. • Triple bonds contain one sigma and two pi bonds. • Figure 9.24 page 363

12. Resonance Structures, Delocalization and π Bonding • Benzene (C6H6) is a molecule in which the pi bonds are not localized. That is the pi bonds are not associated with a particular set of carbon atoms. • The carbon atoms in benzene have an sp2 hybridization, so that the H-C-C bond angle is 120 degrees. • Since the π bonds in benzene are localized between C-C atoms the π bonds are said to be delocalized.

13. Molecular Orbital Theory • Whenever two atomic orbitals overlap, two molecular orbitals form. • The overlap of the 1s orbitals of two hydrogen atoms to form H2 produces two Molecular Orbitals (p 368). • One of the molecular orbitals lies lower in energy. • The lower energy molecular orbital is called the bonding molecular orbital, the other the antibonding molecular orbital. • The interaction between two 1s atomic orbitals is represented in and energy level diagram.

14. Bond Order • The bond order represents the stability of a covalent bond and is defined as half the difference between the number of bonding electrons and the number of antibonding electrons. • Bond Order = ½(# bonding electrons - # antibonding electrons) • A bond order of 1 represents a single bond, a bond order of 2 represents a double bond, etc. • Because Molecular Orbital theory also treats molecules containing an odd number of electrons, bond orders of ½, 3/2 or 5/2 are possible. • Example 9.8 (p 370)