Formic acid: HCOOH

1. Formic acid: HCOOH Acetone

2. Kekulé structures Resonance structures; each point corresponds to a CH Each C is sp2 hybridized, one of the sp2 forming a s-bond with H 1s orbital and the other two forming s-bonds with adjacent C sp2 orbitals. Benzene C6H6

3. The un-hybridized p orbital on each C is available for p-bonding with p orbitals on either of the adjacent C atoms

4. Actual structure of benzene is a resonance hybrid of the two alternating bond patterns; the 6 C atoms are identical, and the electrons in the p-bonds spread around the entire ring This lowers the energy of the molecule - resonance adds stability to a molecule

5. Characteristics of p bonds bonds Energy of C=C is < 2 x energy of C-C bond Energy of CC is < 3 x energy of C-C bond C, N, O form double bonds with one another and with elements from later periods Double bonds are rarely found between elements in period 3 are below - atoms are too large for effective side-by-side overlap. Molecules with alternate double-single bonds - conjugated molecules

6. Isomers: Molecules with the same molecular formula but different structures cis-1,2-dichloroethylene trans-1,2-dichloroethylene Rotation can occur about a single sigma bond Rotation is restricted about a double bond; isomers are a consequence

7. Change of shape triggers a signal along the optic nerve

8. Molecular Orbital Theory VB theory: localized bond VB theory provides the basis of calculating electron distributions in molecules but cannot explain the properties of some molecules. O2: VB theory O: Is2 2s2 2p4 sp2 hybridized O, one sp2 from each forms s-bond and the other two are occupied with the lone pairs. The un-hybridized p on each forms the p-bond Indicates that in O2 molecule, all electrons are paired. However O2 was observed to be paramagnetic

9. VB theory assumes that the electrons are localized between the two bonding atoms Molecular orbital theory: electrons are spread throughout the entire molecule; electrons are delocalized over the whole molecule. Pure atomic orbitals combine to produce molecular orbitals that are spread out, delocalized, over an entire molecule Molecular orbitals are built by adding together -superimposing - atomic orbitals belonging to the valence shell of the atoms in the molecules.

10. H2: wavefunction representing the molecular orbitals (MOs) for H2 can be represented by combining the two atomic orbitals (AOs) for the separated H atoms. Wavefunction of the H2 MO y+ = yA1s + yB1s yA1s or yB1s 1s orbital centered on one of the H atom(A or B) The molecular orbital, y, is a linear combination of atomic orbitals Any molecular orbital formed from a superposition of atomic orbitals is called a LCAO-MO. y+ is a bonding orbital; energy of y+ is lower than that of either AO In H2, the contribution from each AO to the MO is equal

11. The two AOS are waves centered on different nucleii. Bonding orbital: AO wavefunctions interfere constructively - MO wavefunction in blue.

12. N AOs overlapping will form N MOs Two H AOs overlapping form two Mos; one of which is the bonding orbital, y+. The wavefunctions of the two H AOs can also interfere destructively - anti-bonding MO of higher energy than each of the AOs y- = yA1s - yB1s Node between two nuclei Probability of finding electrons between nuclei reduced; nuclei repel each other http://www.shef.ac.uk/chemistry/orbitron/index.html

14. Molecular Orbital Energy Level Diagram Energy of bonding MO < AO Energy of anti-bonding MO > AO

15. Diatomic Molecules Build all possible MOs from available valence AOs Then accommodate valence electrons in molecular orbitals using the aufbau principles 1) Electrons occupy the lowest energy MOs first, then orbitals of increasing energy 2) Pauli exclusion principle: each orbital can occupy up to two electrons; if two electrons in an orbital must be paired 3) Hund’s rule: if more than one orbital of the same energy is available electrons enter them singly with parallel spinds.

16. Lowest unoccupied MO (LUMO) H2 molecular orbital energy-level diagram or correlation diagram Bonding MO - s1s anti-bonding MO s *1s Highest occupied MO (HOMO) H2 Ground state electron configuration (s1s)2

17. Bond Order = 0.5(number of electrons in bonding MOs - number of electrons in anti-bonding MOs) H2+ bond order = 0.5 (s1s)1 H2 bond order = 1 (s1s)2 He2 bond order = 0 (s1s)2 (s*1s)2

18. Period 2 elements In period 2 elements each atom has one 2s and three 2p valence AOs; expect to form eight MOs The two 2s orbitals (one from each atom) overlap to form a s2s bonding MO and a s*2s antibonding MO The six 2p orbitals (three from each atom) overlap to form six MOs The two 2p-orbitals directed toward each other form a bonding s-orbital (s2p) and an anti-bonding s*-orbital (s*2p) Two 2p orbitals that are perpendicular to the internuclear axis overlap side by side to form two bonding p and two anti-bonding p* orbitals.

19. Anti-bonding s and s* orbitals formed from p AOs Bonding p and p* orbitals formed from p AOs

20. MO diagram for homonuclear diatomic molecules Li2 through N2 MO diagram for homonuclear diatomic molecules O2 and F2