© Prentice Hall 2001Chapter 11 Atomic Orbitals We cannot know the exact course of electrons as they...

© Prentice Hall 2001 Chapter 1 1

Atomic Orbitals We cannot know the exact course of electrons as

they orbit the nucleus - Heisenberg uncertainty principle

The wave functions, however, give approximate shapes to orbitals s orbitals are spherical p orbitals are shaped like dumbbells


1s and 2s Orbitals


2p Orbitals


Molecular Orbitals for Hydrogen

Allow the 1s orbital of one hydrogen atom to overlap with the 1s orbital of a neighboring hydrogen atom



The electrons in a molecule are everywhere

The shape we draw is only the surface enclosing a certain percentage of the electron density

Highest electron density is directly between the hydrogen nuclei



The result defines a (sigma) bond between the atoms


Molecular Orbital From p Electrons Molecular orbitals also can be formed

from p orbitals


End-On Overlap of p Orbitals Leads to Bonding


Side-by-Side Overlap of p Orbitals Leads to Bonding


Hybrid Orbitals

Methane, CH4, has four equivalent carbon-hydrogen bonds


Hybrid Orbitals Ground state for carbon is 1s 22s 22p 2

There are only two partially filled orbitals that are capable of participating in bonding

How can four bonds be made? Problem can be solved by promoting one

2s electron to an empty 2p orbital

px py pz

s

px py pz

s


Hybridization If four bonds were formed from the 2s2p2p2p

configuration, we might expect three of the bonds to be at exactly 90o and the other at any angle


Hybridization Theory: Mix the 2s orbital with the three

2p orbitals to form four equivalent hybrid orbitals


Hybridization - Tetrahedral Carbon

In methane, the four sp3 hybrid orbitals overlap with s orbitals from four hydrogen atoms to form four equivalent carbon-hydrogen bonds

A carbon atom that has sp3 hybrid orbitals is called a tetrahedral carbon


Hybridization - Tetrahedral Carbon

The sp3 hybrid orbital on carbon also can bond with another sp3 hybrid orbital from a neighboring carbon to form a carbon-carbon single bond


sp2 Hybridization in Ethene

Each carbon in ethene (H2C=CH2) is bonded to only three atoms

To do this, each carbon hybridizes only three orbitals to give sp2 hybridization


sp2 Hybridization in Ethene The three sp2

hybrid orbitals all lie in a plane and are oriented at angles of 120o

Such carbons are called trigonal planar carbons



The carbon-carbon bond formed from the overlap of an sp2 orbital on one carbon with an sp2 orbital on a neighboring carbon atom results in an orbital which is cylindrically symmetric about the carbon-carbon axis


sp2 Hybridization in Ethene A second bond is formed between the two

carbon atoms via the side-by-side overlap of the remaining (un-hybridized) p orbitals

Electron density accumulates above and below the carbon-carbon axis


sp2 Hybridization in Ethene This second carbon-carbon bond is called a

bond The two p orbitals which overlap side-by-side,

must be parallel to each other The four hydrogen atoms therefore must be in

the same plane No rotation about a carbon-carbon double bond


sp Hybridization in Ethyne Each carbon in ethyne (HCCH) is bonded to

only two atoms To do this, each carbon hybridizes only two

orbitals to give sp hybridization


sp Hybridization in Ethyne The overlap of the sp hybrid orbitals forms

a bond


sp Hybridization in Ethyne The remaining p orbitals overlap side-by-

side, forming bonds with electron density above and below the carbon-carbon axis as well as in front and in back


sp Hybridization in Ethyne

The two bonds together form a cylinder of electron density around the bond


Summary of Orbital Hybridization

© Prentice Hall 2001Chapter 11 Atomic Orbitals We cannot know the exact course of electrons as they...

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