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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

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 s bonds. Note: electron spin is paired in the s orbital. By definition: z is the direction along the internuclear axis . N 2 N: 1s 2 2s 2 2p 3.

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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

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  1. HF H: 1s1 F: 1s22s22p5 Overlap between the valence orbital of H (1s) and valence orbital of F (2p) to form a s bonds Note: electron spin is paired in the s orbital By definition: z is the direction along the internuclear axis

  2. N2 N: 1s22s22p3 The two pz orbitals from each N can overlap to form a s orbital. The px and py orbitals are perpendicular to the internuclear axis

  3. s bond - overlap of two pz orbitals p bond - overlap of two px orbitals and/or two py orbitals

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

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

  6. Hybridization of Orbitals C: Is2 2s2 2p2 VB theory, as described so far, would predict that C can form just two bonds

  7. In CH4, 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. CH4 Promoting a 2s electron to 2p allows C to have four unpaired electrons. All bonds on CH4 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 sp3 orbitals

  9. Hybrid orbitals are constructed on an atom to reproduce the electron arrangement of the experimentally determined shape of the molecule. In CH4: each sp3 orbital has one unpaired electron Each overlaps with a 1s orbital of H to form s-bond The four resulting s-bonds point towards the corners of a tetrahedron. All four s-bonds are identical http://www.whfreeman.com/chemicalprinciples/

  10. Ethane: C2H6 Each C has four sp3 hybrid orbitals, pointing towards the corner of a tetrahedron, each with one electron Three of these four overlap with three H atoms forming s-bonds (C sp3, H 1s). The C-C bond is formed by an overlap of the remaining sp3 orbital on each C forming a s-bond (C sp3, C sp3).

  11. NH3 H: Is1 N: Is2 2s2 2p3 Hybridize the 2s and 2p orbitals in N to form four sp3 hybrid orbitals. One of the sp3 has two paired electrons - the lone pair on N The three other sp3 orbitals form s-orbitals with each of the three H 1s orbitals

  12. H2O H: Is1 O: Is2 2s2 2p4 Hybridize the 2s and 2p orbitals in O to form four sp3 hybrid orbitals. O   2p 2s O  sp3 Two of the sp3 have two paired electrons - the two lone pairs The two other sp3 orbitals overlap with H 1s orbitals

  13. An s orbital and two p orbitals can hybridize to form three sp2 hybrid orbitals which point to the corners of an equilateral triangle - trigonal planar geometry Example: BF3

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

  15. PCl5 P: [Ne] 3s2 3p3 Cl: [Ne] 3s2 3p5 P  Cl  3p 3p  3s 3s Promote a 3s electron to the 3d orbital P   _ _ _ _ sp3d empty 3d Valence shell expansion - expansion to include d orbitals along with s and p orbitals

  16. One 2, three p, and one d orbital form five sp3d hybrid orbitals, each pointing towards a corner of a trigonal bipyramid

  17. One 2, three p, and two d orbital form six sp3d2 hybrid orbitals, each pointing towards a corner of a octahedron

  18. SF6 S: [Ne] 3s2 3p4 F: [He] 2s2 2p5 S  3p  3s Include two 3d orbital and hybridize one s, three p and two d S    _ _ _ sp3d2 empty 3d

  19. Multiple Bonds Ethylene: CH2CH2 Experimental data: all six atoms lie in the same plane and the H-C-H and C-C-H bond angles are 120o. Trigonal planar geometry indicates that each C is sp2 hybridized For each C: two of the sp2 orbitals bond with two H 1s orbitals to form s-bonds, The third Csp2 bond on each bond with each other to form a C-C s-bond

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

  21. Acetylene: C2H2 Linear molecule; each C is sp hybridized, leaving two pure p orbitals on each C http://www.whfreeman.com/chemicalprinciples/

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

  23. Formic acid: HCOOH

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