1. 1 Valence Bond Theory vs. MO Theory VB Theory begins with two steps: 1) hybridization AOs on atoms participating in bonding 2) Combination of hybrid orbitals to make bonds. Key differences between MO and VB theory: 1) MO theory: has electrons distributed over whole molecule. (molecule centered) VB theory: localizes an electron pair between two atoms. (bond centered) 2) MO theory: combines AOs between different atoms to make MOs (LCAO) VB theory: combines AOs on the same atom to make hybridized atomic orbitals (hybridization) 3) MO theory: the symmetry must be retained in each orbital. VB theory: all orbitals must be viewed simultaneously to see retention of the molecule’s symmetry.

2. VB Theory of BeH 2 Two hybrid atomic orbitals are made to fit the shape of the molecule, in this case linear, using atomic orbitals of the Be atom!    Unused AO are left behind as unhybridized atomic orbitals The energy of the hybrid atomic orbitals are intermediate between those of the original constituent AO’s The hybrid orbitals combine with other orbitals, atomic or hybrid, creating both bonding and anti-bonding molecular orbitals, which are localized molecular orbitals

3. VB Theory of BH 3      BH 3 is trigonal planar with three equal B—H bonds To get this shape the 2s with two 2p AO’s to generate three equivalent hybrid atomic orbitals Combination with the H 1s leads to bonding and anti-bonding molecular orbitals, which are localized molecular orbitals pointing to the corners of a triangle

4. VB Theory of CH 4        CH 4 is tetrahedral with 4 equal C-H bonds To get this shape, we need to combine all the n=2 AO’s to generate four equivalent hybrid atomic orbitals In combination with the H 1s leads to bonding and anti-bonding molecular orbitals localized and pointing to the corners of a tetrahedron

5. 5 Valence Bond Theory Summary Atoms orbitals are hybridized only if it’s necessary to attain the observed geometry and bond lengths. Terminal atoms are not typically hybridized. Geometry determines hybridization : i) Linear (180º) = sp (s+p with two leftover p orbitals) ii) Trigonal planar (120º) = sp 2 (s+p+p with one leftover p orbital) iii) Tetrahedral (109.5º) = sp 3 (s+p+p+p with no leftover p orbitals) Hybrid orbitals combine with each other to make  bonds in which two electrons are localized between two atoms.  bonds are made by combining the unused  -symmetric p orbitals.

6. Ethane  VSEPR theory requires both carbon atoms to be tetrahedral The shape of the molecule, requires that contacts be minimized between the atoms – this is known as the staggered conformation Bonding is explained by using sp 3 hybrid orbitals on each C. The H atoms bond using their 1s atomic orbitals In all there are 14 electrons or 7 electron pair bonds in the molecule 1  C-C sp 3 -sp 3 single bond and 6  C-H sp 3 -s single bonds are formed

7. Double bonds: ethene  2px2px 2px2px The sigma skeleton of etheneThe pi manifold of ethene Change perspective to show the π bond! If we treat ethane by the VSEPR theory, we find that both carbon atoms are trigonal planar The molecule is planar. Why ? sp 2 hybrid orbitals on each carbon atom, which leaves one atomic p orbital unused on each C atom, while H atoms use their 1s atomic orbitals There are 6 e’ pair bonds in the molecule, 5 in σ orbitals, 1 in the π orbital  sp 2 -sp 2 C-C bond, 1  p x -p x C-C bond, and 4  sp 2 -s C-H bonds

8. VB Theory of H 2 O H 2 O is bent and belongs to the tetrahedral family with 2 BP and 2 LP The s and p orbitals combine sp 3 hybrids The 6 e’s from O, singly occupy 2 sp 3 orbitals and doubly occupy the remaining 2 as LP’s O 2 H sp 3 hybrids 1 s 1 s 2 C-H    2 LP’s The 2 sp 3 orbitals combine with 2 1s orbital to form 2 C-H bonds

9. 9 Triple Bonds: Ethyne The Lewis structure for ethyne (C 2 H 2 ) VSEPR theory: each carbon atom is linear hence sp hybridized i) C–H  bonds: combination of an sp orbital from C and a 1s orbital from H. ii) C–C  bonds: combination of sp orbitals from each C. iii) C–C  bonds: combination of 2p orbitals from each C.

10. 10 Formaldehyde The Lewis structure for formaldehyde (CH 2 O). VSEPR theory: Carbon atom is trigonal planar hence sp 2 hybridized: i) C–H  bonds: combination of an sp 2 orbital from C and a 1s orbital from H. ii) C–O  bond: combination of an sp 2 orbital from C and a sp 2 orbital from O. iii) C–O  bond: combination of 2p z orbitals from C and O. iv) Lone pairs: remaining sp 2 hybrid AOs of O.

11. 11 VB Theory of HCN The Lewis structure for hydrogen cyanide (HCN) is. VSEPR theory: C and N atoms are linear hence are sp hybridized. i) C–H  bonds: combination of an sp orbital from C and a 1s orbital from H. ii) C–N  bond: combination of an sp orbital from C and a 2p orbital from sp ybrid orbital on N. iii) C–N  bonds: combination of 2p x and 2p y orbitals from C and N. iv) Lone pair: remaining sp hybrid atomic orbital (2s).

12. Bonding in large molecules sp 3 sp 2  sp 3- sp 3 O-C bond  s - sp 3 H -O bond 2  s - sp 3 H-C bonds  sp 3- sp 3 C-N bond  s - sp 3 H-C bonds 2  s - sp 3 N-H bonds 1 sp 3 LP  sp 3- sp 2 C-C bond  sp 3- sp 3 C-C bond  sp 2- sp 2 O-C bond  p - p O-C bond  sp 2- sp 3 C-O bond  s - sp 3 H-O bond 2 sp 2 LP’s