1. HF H: 1s 1 F: 1s 2 2s 2 2p 5 Overlap between the valence orbital of H (1s) and valence orbital of F (2p) to form a  bonds Note: electron spin is paired in the  orbital By definition: z is the direction along the internuclear axis

2. N 2 N: 1s 2 2s 2 2p 3 The two p z orbitals from each N can overlap to form a  orbital. The p x and p y orbitals are perpendicular to the internuclear axis

3.  bond - overlap of two p z orbitals  bond - overlap of two p x orbitals and/or two p y orbitals

4. In a  bond, electron density has a nodal plane that contains the bond axis

5. According to the VB theory A single bond is a  -bond A double bond is  -bond plus a  -bond A triple bond is a  -bond and two  -bonds. VB theory: assumes bonds form when unpaired electrons in valence shell atomic orbitals pair the atomic orbitals overlap end to end to form  -bonds or side by side to form  -bonds.

6. Hybridization of Orbitals C: Is 2 2s 2 2p 2 VB theory, as described so far, would predict that C can form just two bonds

7. In CH 4, C forms four bonds. C needs four unpaired electrons so that each can pair with a H atom - need to revise valence-bond theory “Promote” a 2s electron to a 2p orbital - this requires energy. But now C has four unpaired electron and since bonding releases energy the cost of promoting is overcome by the lowering of energy on bond formation

8. CH 4 Promoting a 2s electron to 2p allows C to have four unpaired electrons. All bonds on CH 4 are equivalent “Mix” the 2s and the three 2p orbitals to form four hybrid orbitals all of the same energy and spatial distribution - hybridization. One s + three p = four sp 3 orbitals

10. Hybrid orbitals are constructed on an atom to reproduce the electron arrangement of the experimentally determined shape of the molecule. In CH 4 : each sp 3 orbital has one unpaired electron Each overlaps with a 1s orbital of H to form  -bond The four resulting  -bonds point towards the corners of a tetrahedron. All four  -bonds are identical http://www.whfreeman.com/chemicalprinciples/

11. Ethane: C 2 H 6 Each C has four sp 3 hybrid orbitals, pointing towards the corner of a tetrahedron, each with one electron Three of these four overlap with three H atoms forming  - bonds (C sp 3, H 1s). The C-C bond is formed by an overlap of the remaining sp 3 orbital on each C forming a  -bond (C sp 3, C sp 3 ).

12. NH 3 H: Is 1 N: Is 2 2s 2 2p 3 Hybridize the 2s and 2p orbitals in N to form four sp 3 hybrid orbitals. One of the sp 3 has two paired electrons - the lone pair on N The three other sp 3 orbitals form s-orbitals with each of the three H 1s orbitals

13. H 2 O H: Is 1 O: Is 2 2s 2 2p 4 Hybridize the 2s and 2p orbitals in O to form four sp 3 hybrid orbitals. O   2p 2s O  sp 3 Two of the sp 3 have two paired electrons - the two lone pairs The two other sp 3 orbitals overlap with H 1s orbitals

14. An s orbital and two p orbitals can hybridize to form three sp 2 hybrid orbitals which point to the corners of an equilateral triangle - trigonal planar geometry Example: BF 3

15. An s and a p orbital can hybridize into two sp orbitals that point in opposite directions - linear geometry

16. PCl 5 P: [Ne] 3s 2 3p 3 Cl: [Ne] 3s 2 3p 5 P  Cl  3p 3p 3s 3s Promote a 3s electron to the 3d orbital P  ___ _ sp 3 d empty 3d Valence shell expansion - expansion to include d orbitals along with s and p orbitals

17. One 2, three p, and one d orbital form five sp 3 d hybrid orbitals, each pointing towards a corner of a trigonal bipyramid

18. One 2, three p, and two d orbital form six sp 3 d 2 hybrid orbitals, each pointing towards a corner of a octahedron

19. SF 6 S: [Ne] 3s 2 3p 4 F: [He] 2s 2 2p 5 S  3p  3s Include two 3d orbital and hybridize one s, three p and two d S   __ _ sp 3 d 2 empty 3d

21. Multiple Bonds Ethylene: CH 2 CH 2 Experimental data: all six atoms lie in the same plane and the H-C-H and C-C-H bond angles are 120 o. Trigonal planar geometry indicates that each C is sp 2 hybridized For each C: two of the sp 2 orbitals bond with two H 1s orbitals to form  -bonds, The third Csp 2 bond on each bond with each other to form a C-C  -bond

22. The “pure” 2p orbitals on each C overlap to form a  -bond between the two C atoms The electron density of this  -orbital lies above and below the axis of the C-C  -bond http://www.whfreeman.com/chemicalprinciples/

23. Acetylene: C 2 H 2 Linear molecule; each C is sp hybridized, leaving two pure p orbitals on each C http://www.whfreeman.com/chemicalprinciples/

24. Multiple bonds are formed when an atom forms a  -bond by using an sp or sp2 hybrid orbital and one or more  -bonds by using un-hybridized p orbitals

25. Formic acid: HCOOH