4.6 Quantum Mechanics and Bonding: Hybridization

How are Bonds Formed? THE STORY SO FAR… The Lewis structure of a molecule provides a simplified view of bond formation and allows us to determine the number of bonding and lone electron pairs VSEPR theory can then be used to predict the molecular geometry We can even use electronegativity values to determine bond polarity and bond dipoles to determine the overall molecular polarity BUT…Lewis structures, VSEPR theory, and bond dipoles do not describe how bonds are actually formed

Valence Bond Theory Valence bond theory: states that atomic orbitals overlap to form a new orbital with a pair of opposite spin electrons A covalent bond forms when 2 atomic orbitals, each with an unpaired electron, overlap When the covalent bond forms, the lowest energy state is obtained when participating electrons are of opposite spin To achieve this, some orbitals are hybridized. Why hybridized orbitals? Hybridization: the process of forming hybrid orbitals from the combination of at least 2 different orbitals (s, p and d)

Forming Hybrid Orbitals The covalent bond in hydrogen gas (H 2 ) is formed when the 1s orbital of one hydrogen atom overlaps with the 1s orbital of another hydrogen atom The covalent bond in HF is formed when hydrogen’s 1s orbital and one of fluorine’s 2p orbitals overlap

Forming Hybrid Orbitals The covalent bond in F 2 is formed when one of fluorine’s 2p orbitals overlaps with one of the 2p orbitals in another fluorine atom

The Forming of sp 3 orbitals Example: Determine the hybridization in CH 4 Problem: Methane has 4 covalent bonds, but Carbon has only 2 available orbitals for overlap! Energy level diagram for Carbon shows only 2px and 2py available for bonding! Methane has 4 covalent bonds

Solution: a) Only 1 orbital is available for covalent bonding in the Carbon atom! b) A paired electron from the s-orbital is promoted to the unfilled 2pz orbital. c) The one s-orbital and three p-orbitals merge to form a set of equivalent sp3 hybrid orbitals. There are now 4 orbitals, each with 1 unpaired orbitals available for covalent bonding

The forming of sp 3 orbitals in Methane

Examples 1.Determine the hybridization for the H 2 O molecule.

The forming of sp 2 orbitals Example: Determine the hybridization for BF 3

The forming of sp orbitals Example: Determine the hybridization for BeCl 2

Summary of Hybrid Orbitals

Determine the hybridization for the ammonia molecule Practice:

Sigma and Pi Bonds A sigma bond is formed when the lobes of two orbitals overlap end to end (electron density is between the nuclei) A pi bond is formed when the lobes of two orbitals overlap side by side (electron density is above and below the nuclei) σ bond

Double Bonds Consider the molecule ethene: Since each carbon is bonded to 3 other atoms (trigonal planar), the carbon atom undergoes partial hybridization to form sp 2 orbitals There is one remaining p orbital with 1 electron in it

Double bonds contain: 1 sigma bond (between the two sp 2 orbitals of carbons) 1 pi bond (between the p orbitals of two carbons) Double Bonds

Triple Bonds Consider the molecule ethyne: Since each carbon is attached to 2 other atoms (linear shape), the carbon undergoes partial hybridization to form sp orbitals There are 2 remaining p orbitals that both hold a single electron

Triple Bonds Triple bonds contain: 1 sigma bond (between the two sp orbitals of carbons) 2 pi bonds (between the two p orbitals of each carbon)

Summary of Types of Bonds

Homework P. 238 #1-6