Bonding Theories in Gen. Chem.

1. Bonding Theories in Gen. Chem. • Assumes a basic (in-class) background of… • Lewis Structures • VSEPR • Valence Bond Theory • Will briefly touch on these to lay groundwork

2. Bonding Theories in General • Chemists have advanced theories on bonding to explain chemical phenomena • Some work well…some discarded entirely. • Plusses/minuses for many GC theories • Simple but incomplete (Lewis, VSEPR) • More complete, but high-level math and abstract thinking required

3. Enter…Valence Bond Theory • Accounts for orbitals in bonding • Basic tenets • Bonds form when orbitals overlap • Better overlap leads to a stronger bond • Works well for H2 and some other simple molecules

4. VB Approach to p-orbitals Works for simple HF molecule. Overlap of the s and p orbitals However…approach breaks down for CH4

5. 109.5 v 90—a clash of angles The three p-orbitals (below) orthogonal The bond angles in methane aren’t 90 These atomic orbitals must change shape

6. VB “Hybridization” of AO’s For methane, the s and p orbitals ‘mix’ and form orbitals with different shapes (naturally

7. VB hybrid orbitals, overlap and CH4 VSEPR, Lewis structures and VB Theory provide a firm foundation for viewing chemical bonding. Can’t expain ‘everything’ though…O2?

8. Moving into MO Theory • VB Theory fails to predict the attraction of liquid oxygen to a magnet. • Both theories rely on ‘mixing’ atomic orbitals • VB forms ‘hybrid’ orbitals • MO theory forms ‘molecular orbitals’ • Distinction comes from orbital location

9. Mixing of AO’s to form MO’s • When two atomic orbitals mix (from different atoms), two molecular orbitals form • Bonding orbital (lower in energy than either orbital) • Anti-bonding orbital (higher in energy). • Like VB Theory, the shapes of orbitals change

10. Starting Simple…just some H’s Electrons occupy lower energy orbital (so H2 forms) Considering the simplest case…two 1s orbitals from hydrogens.

11. Doubling the ‘complexity’? Nah Electrons occupy lower AND higher energy orbitals—so there’s no benefit to forming He2. Time to introduce “Bond Order” So…we established H2 forms…can MO Theory explain why He2 does not? Sure…

12. Complicating things just a smidge • Beyond H/He, the orbitals become a bit more complex—due to the # and types of • Some guidelines—to help us out • AO’s in = MO’s out • AO’s with similar energies/shapes mix better • MO’s accommodate just 2 e’s • Electrons fill lowest energy orbitals first

13. Mixing two p-orbitals ‘side-on’ • Two of the p-orbitals overlap ‘side-on’ (px and py) and the resulting orbitals look like • Referred to as “pi” orbitals

14. Mixing two p-orbitals ‘head-on’ This ‘head-on mixing forms a sigma bond, and the lower energy orbital is along the bond axis.

15. Energy levels of the mixed orbitals The figure on the right shows how these orbitals align electronically Note (only second row—1s omitted)

16. Filling the orbitals for each element Diagrams for Li, Be, B, and C Predictions? Be2 unstable, B2 paramagnetic!

17. Moving along to N2 and O2 Nitrogen…bond order of three (as expected) And oxygen is paramagnetic

18. MO Successes…not limited to O2 Emission 650-680 nm N2 absorbs cosmic energy, becomes ‘excited’, then emits this energy…back to ‘regular’ N2 Check out the ‘transition’ below.