Valence Bond Theory vs. MO Theory - PowerPoint PPT Presentation

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Valence Bond Theory vs. MO Theory
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Valence Bond Theory vs. MO Theory

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  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 wholemolecule. (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 symmetrymust 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 BeH2       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 unhybridizedatomic 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 BH3          BH3 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 orbitalspointing to the corners of a triangle

  4. VB Theory of CH4             CH4 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 andpointing to the corners of a tetrahedron

  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) = sp2 (s+p+p with one leftover p orbital) iii) Tetrahedral (109.5º) = sp3 (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 sp3 hybrid orbitals on each C. The H atoms bond using their 1satomic orbitals In all there are 14 electrons or 7 electron pair bonds in the molecule 1s C-Csp3-sp3 single bond and 6s C-H sp3-s single bonds are formed       

  7.      Change perspective to show the π bond! Double bonds: ethene If we treat ethane by the VSEPR theory, we find that both carbon atoms are trigonal planar The molecule is planar. Why ? sp2 hybrid orbitals on each carbon atom, which leaves one atomic p orbital unused on each C atom, while H atoms use their 1satomic orbitals There are 6 e’ pair bonds in the molecule, 5 in σ orbitals, 1 in the π orbital 1 s sp2-sp2 C-C bond, 1ppx-px C-C bond , and 4s sp2-s C-H bonds The sigma skeleton of ethene The pi manifold of ethene  2px 2px

  8. VB Theory of H2O O 2 H 2 LP’s       sp hybrids 3   s s 1 1   2 C-H H2O is bent and belongs to the tetrahedral family with 2 BP and 2 LP The s and p orbitals combine sp3 hybrids The 2 sp3orbitals combine with 21s orbital to form 2 C-H bonds The 6 e’sfrom O, singly occupy 2 sp3 orbitals and doubly occupy the remaining 2 as LP’s

  9. Triple Bonds: Ethyne The Lewis structure for ethyne (C2H2) 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. Formaldehyde The Lewis structure for formaldehyde (CH2O) . VSEPR theory: Carbon atom is trigonal planar hence sp2 hybridized: i) C–H  bonds: combination of an sp2 orbital from C and a 1s orbital from H. ii) C–O  bond: combination of an sp2 orbital from C and a sp2orbital from O. iii) C–O  bond: combination of 2pz orbitals from C and O. iv) Lone pairs: remaining sp2 hybrid AOs of O.

  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 sporbital 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 2px and 2py orbitals from C and N. iv) Lone pair: remaining sp hybrid atomic orbital (2s).

  12. Bonding in large molecules sp2 sp3 sp2 sp3 sp3 sp3 sp3 s s-sp3 H -O bond 2 s s-sp3 N-H bonds s sp3-sp3 O-C bond 2 s s-sp3 H-C bonds s sp2-sp2 O-C bond 1 sp3 LP p p-p O-C bond s sp3-sp2 C-C bond s sp3-sp3 C-C bond s sp2-sp3 C-O bond s s-sp3 H-C bonds 2 sp2 LP’s s sp3-sp3 C-N bond s s-sp3 H-O bond