Formic acid HCOOH Acetone Benzene C 6 H

Formic acid: HCOOH Acetone

Benzene C 6 H 6 Kekulé structures Resonance structures; each point corresponds to a CH Each C is sp 2 hybridized, one of the sp 2 forming a s-bond with H 1 s orbital and the other two forming s-bonds with adjacent C sp 2 orbitals.

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

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

Characteristics of p 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

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

Change of shape triggers a signal along the optic nerve

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. O 2: VB theory O: Is 2 2 p 4 sp 2 hybridized O, one sp 2 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 O 2 molecule, all electrons are paired. However O 2 was observed to be paramagnetic

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.

H 2: wavefunction representing the molecular orbitals (MOs) for H 2 can be represented by combining the two atomic orbitals (AOs) for the separated H atoms. Wavefunction of the H 2 MO y+ = y. A 1 s + y. B 1 s y. A 1 s or y. B 1 s 1 s 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 H 2, the contribution from each AO to the MO is equal

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

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- = y. A 1 s - y. B 1 s Node between two nuclei Probability of finding electrons between nuclei reduced; nuclei repel each other http: //www. shef. ac. uk/chemistry/orbitron/index. html

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

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.

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

Bond Order = 0. 5(number of electrons in bonding MOs - number of electrons in anti-bonding MOs) H 2+ bond order = 0. 5 (s 1 s)1 H 2 bond order = 1 (s 1 s)2 He 2 bond order = 0 (s 1 s)2 (s*1 s)2

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

Antibonding Bonding s and s* orbitals formed from p AOs p and p* orbitals formed from p AOs

MO diagram for homonuclear diatomic molecules Li 2 through N 2 MO diagram for homonuclear diatomic molecules O 2 and F 2