Class notes for Chapter 6: Chemical Bonding II in the class General Chemistry: Macroscopic Investigations and Reaction Principles (CHEM 130) at the University of Michigan. Topics covered include the Valence Bond and Molecular Orbital theories, hybridization, and diatomic and polyatomic molecules
Valence Bond Theory
●Valence Bond Theory: advanced bonding model in which electrons reside in hybridized
blends of standard atomic orbitals
○Chemical bonds result from overlapping orbitals ●Unlike Lewis model (electrons are dots), VBT treats VE as residing in atomic orbitals
○Sometimes orbitals are the spdf from chapter 2; sometimes they are hybridized
●In covalent bonds, 2 orbitals overlap and that overlap has 2 electrons of opposite spin
○Potential energy usually negative
●Bond energy is diff between zero energy and energy of the bond
●Atomic Orbitals: explains shape and bonding
●When two atoms interact, the electrons/nucleus of each atom interact with each other
○Valence bond theory calculates the effect of those interactions on orbital energy
■If system energy is lowered, a chemical bond forms
■If system energy is raised, a chemical bond does not form
●Interaction energy calculated as function of distance between nuclei of two bonding atoms
●Bond length (minimum energy) → most stable point
○The two 1s orbitals have significant overlap, interact with both nuclei
○Value of interaction energy is the bond energy
●Interaction energy is usually negative (stabilizing) when interacting atomic orbitals contain two electrons that can orient with opposing spins
○Usually from two half-filled orbitals overlapping, occasionally from one filled orbital overlapping with one empty orbital ○Results in net energy stabilization → covalent bond ●Summary:
○VEs reside in atomic orbitals that can be standard spdf orbitals OR hybrid combinations
○Chemical bonds result from overlap of two half-filled orbitals and spin pairing of the 2 VEs in those orbitals
○Geometry of overlapping orbitals determines shape of molecule
●EX: H2S
●Both have two spots w unpaired electrons
●The two half-filled H orbitals overlap with the two half-filled S orbitals Hybridization
●Orbitals in a molecule are not the same as the orbitals in a single atom
●Hybridization: procedure in which standard atomic orbitals (spdf) are combined to form hybrid orbitals that correspond more closely to electron distribution in bonded atoms
○Hybrid orbitals still localized on individual atoms
●Hybrid orbitals: use # bonding pairs + # lone pairs
○2 → sp → linear
○3 → sp2 → trigonal planar
○4 → sp3 → tetrahedral
○5 → sp3d → trigonal bipyramidal
○6 → sp3d2 → octahedral
●Orbital overlap lowers potential energy of electrons, strengthens bonds
○Hybrid orbitals overlap more → more beneficial
●Hybridization takes energy → the more bonds an atom forms, the more tendency to hybridize
○Energy payback pays off with more hybrid orbitals
○Central and interior atoms hybridize the most
■Carbon always hybridizes
●Orbitals are conserved: number of atomic orbitals “in” = number of hybrid orbitals “out”
○Named by and have characteristics of atomic orbitals from which they form
●Type of hybridization that occurs is the one that has the lowest overall energy for that molecule
●Can predict hybridization using VSEPR
sp3 Hybridization
●sp3 hybrids come from one s and three p orbitals
○All sp3 orbitals are degenerate
●sp3 orbitals are tetrahedral
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