Hybridization of Orbitals Creating Bonds that are neither pure s nor p nor d!

Valence Bond Model Linus Pauling developed valence bond theory/model in the 1930’s and won the Nobel Prize for his work in 1954 A covalent bond is a pair of electrons with opposite spins 1s 2s 2px 2py 2pz F (in HF)

So to form a covalent bond, an atom needs an unpaired electron However, most atoms don’t have unpaired electrons. . . . 1s 2s 2px 2py 2pz C We assume that as the bonding atoms approach one another, the orbitals undergo a dramatic change: they hybridize

How do I predict the hybridization of the central atom? Count the number of lone pairs AND the number of atoms bonded to the central atom # of Lone Pairs + # of Bonded Atoms Hybridization Examples 2 sp BeCl2 3 sp2 BF3 4 sp3 CH4, NH3, H2O 5 sp3d PCl5 6 sp3d2 SF6

Practice Problems PCl5 SO3 Protocol: Lewis dot structure, count pairs around central atom Atoms bond hybrid orbitals repulsions Take a picture shape bond angles polarity

Multiple Bonds Geometry is not effected by multiple bonds (double, triple) This is because the extra pair of electrons to make a multiple bond are not hybridized! Ethylene acetylene H H H—C C—H C C H H

Sigma and Pi bonds Single bonds In the valence bond theory, single bonds are made up of atomic orbitals that are cylindrically symmetric about the line joining the bonded atoms (the inter-nuclear axis). Such bonds are called sigma (s) bonds. Overlap of 2 s orbitals in H2 Overlap of an s and a p orbital in HF Overlap of 2 p orbitals in F2 Overlap of an s or p orbital with an spy hybrid orbital - BeCl2, CH4, PCl5, etc. Multiple bonds The second (or third) bond in a multiple bond involves overlap of two p orbitals that are perpendicular to the internuclear axis. The inter-nuclear axis contains a nodal plane. Electron density is above and below the nodal plane. These bonds are called pi (p) bonds.

Bonding in Ethylene Sigma bonds on carbon result from sp2 hybrid orbitals. The resulting geometry is trigonal planar for each carbon. The unhybridized p orbital on each carbon form the second C–C bond, which is a pi bond.

Free Rotation Single Bonds (sigma bonds) allow free rotation about the bond. Multiple Bonds do not allow the molecule to rotate about the bond, since the degree of overlap of the pi-bonded orbitals would be changed. http://www.enter.net.au/~fairsci/chemistry/yr12/alkane.htm

Sigma (s) and Pi (p) Bonds 1 sigma bond Single bond Double bond 1 sigma bond and 1 pi bond Triple bond 1 sigma bond and 2 pi bonds How many s and p bonds are in the acetic acid (vinegar) molecule CH3COOH? C H O s bonds = 6 + 1 = 7 p bonds = 1

cis and trans Isomerism Rotation about the internuclear axis of C-C single bonded atoms is relatively easy at room temperature. Rotation about the internuclear axis of C=C double bonded atoms is not very easy at room temperature: p-bond energy ~ 260 kJ/mol The result can be that two structural forms are “frozen out” Each form is actually a different compound because the atoms have different arrangements in 3-space Each can have different physical properties such as boiling temperature, molecular polarity or color Isomers are compounds that have the same molecular formula but different arrangements of the atoms in 3-space Example: cis- and trans- 1,2-dichloroethene

Valence Bond Theory Delocalized p bonds: When two or more resonance structures can be written involving p bonds, the p bonds can be represented as being spread out over the molecule where the p bonds occur. The electrons and the bonds containing them are delocalized. Example: Benzene The s framework of C-C and C-H bonds is based on sp2 hybridized C atoms. a. After accounting for s bonding, an unhybridized p orbital remains on each C atom, with one electron per p orbital b. The p orbitals overlap to form p bonds that form a continuous p electron cloud above and below the plane of the ring Delocalized p bonds

Summary of Valence Bond Results Bonded atoms share one or more pairs of electrons. At least one sigma bond exists between each pair of bonded electrons. Sigma bonds are cylindrically symmetric along the inter-nuclear axis and electrons are concentrated - localized - between the bonded atoms. An appropriate set of hybrid atomic orbitals is formed to form s bonds. The set of hybrid orbitals depends on the number of s bonds to be formed, the number of non-bonded electron pairs and the geometry of the overall molecule. Atoms sharing more than one pair of electrons form pi bonds by sideways overlap of p atomic orbitals. pi bonds have a nodal plane containing the inter-nuclear axis In pi bonds, electron density is concentrated above and below the nodal plane. Molecules with two or more resonance structures can have pi bonds delocalized over more than two atoms.

Practice Problem What is the hybridization of each indicated atom in the following molecule? How many sigma and pi bonds are in the molecule?