Valence bond theory Electrons are not simply dots

Valence bond theory Electrons are not simply dots And bonds are not sticks

Learning objectives § Describe principles of valence bond theory § Predict hybridization of orbitals based on Lewis dot structures and electronic geometry § Describe difference between sigma and pi bonding

Taking it to the next level: acknowledging orbitals § VSEPR is quite successful in predicting molecular shapes based on the simplistic Lewis dot approach § But our understanding of the atom has the electrons occupying atomic orbitals § How do we reconcile the observed shapes of molecules with the atomic orbital picture of atoms

Valence bond theory § Valence bond theory is the simplest approach to an orbital picture of covalent bonds § Each covalent bond is formed by an overlap of atomic orbitals from each atom § The individual orbital identity is retained § The bond strength is proportional to the amount of orbital overlap

Overlap of two 1 s orbitals in H 2

§ Overlap of two 2 p orbitals directed along the bond axis (sigma bond) § Overlap of p and s orbitals

Problems with tetrahedral bonds § In CH 4 the bonds are all equivalent and at angles of 109. 5° § The 2 p orbitals in C are at 90° - far from optimum for overlap § The ground state configuration is 2 s 22 p 2 § Reconcile these facts with the known structure

Hybridization § The wave mechanics permits mixing of the atomic orbital set to produce “hybrid” orbitals § Hybridization alters the shape and energy of the original § In the case of C, the differences between the 2 s and 2 p are smoothed out and a homogeneous collection of four sp 3 hybrid orbitals is produced

sp 3 hybridization § Formally, one of the 2 s electrons is promoted to the empty 2 p orbital (an energy cost, which is repaid on bond formation) § The four basis orbitals are then “hybridized” to yield the set of four sp 3

Tetrahedral directions and sp 3 hybrids

Valence bond picture of CH 4 § Each C sp 3 hybrid contains one electron § Each H 1 s contains one electron

Lone pairs occupy sp 3 hybrid orbitals § Valence bond picture of the tetrahedral electronic geometry provides same results for the molecules with lone pairs

Notes on hybridization § The total number of orbitals is unchanged § Four atomic orbitals (s + 3 x p) give four hybrid orbitals (4 x sp 3) § The electron capacity remains unchanged § There is one hybridization scheme for each of the five electronic geometries § The same hybridization scheme is always used for a given electronic geometry

sp hybridization for linear geometry § One s and one p orbital

sp 2 hybridization for trigonal planar § One s and two p orbitals

Sigma and pi bonding § The hybridized orbitals describe the electronic geometry: bonds along the internuclear axes (sigma bonds) § The “unused” p orbitals overlap in a parallel arrangement above and below the internuclear axis (pi bonds)

Comparison of pi and sigma bonding

Pi bonding accounts for bond multiplicity § Two unused p orbitals in sp hybrid (linear geometry) § Two pi bonds § N≡N triple bond (one sigma, two pi) § One unused p orbital in sp 2 hybrid (trigonal planar geometry § One pi bond § C=C double bond (one sigma, one pi)

Valence bond picture of ethylene H 2 C=CH 2 § Sigma bonds between C and H (blue/red) and C (blue) § Six electrons around C § Pi bond between C and C (green) § Two electrons around C

Valence bond picture of acetylene HC≡CH § Sigma bonds between C and H (red and blue) and C (blue) § 4 electrons around C § Two pi bonds between C and C (green) § 4 electrons around C

Beyond coordination number 4 § Invoke empty d orbitals (impossible for second row elements) § One d orbital for trigonal bipyramidal § Two d orbitals for octahedral § Number of orbitals in hybrid always equals number of charge clouds

Trigonal bipyramid – sp 3 d

Octahedral –sp 3 d 2

Shortcomings of valence bond § § § The orbitals still maintain atomic identity Bonds are limited to two atoms Cannot accommodate the concept of delocalized electrons – bonds covering more than two atoms § Problems with magnetic and spectroscopic properties