Overlap and Bonding

Valence Bond Theory

• Lewis taught us to think of covalent bonds forming through the sharing of electrons by adjacent atoms. • In such an approach this can only occur when orbitals on the two atoms overlap according to Walter Heitler and Fritz London .

Atomic Orbitals are bad… mmmkay Linus Carl Pauling (1901-1994) Nobel prizes: 1954, 1962

Localized Electron Model

Overlap and Bonding

From this description of covalent bonding rises the Localized Electron Model:

• Increased overlap brings the electrons and nuclei closer together while simultaneously decreasing electronelectron repulsion. • However, if atoms get too close, the internuclear repulsion greatly raises the energy.


Orbitals overlap to form a bond between two atoms due to simultaneous attractions to both nuclei.
Two electrons, of opposite spin, can be accommodated in the overlapping orbitals (usually one electron is supplied by each of the two bonded atoms)
Because of orbital overlap, the bonding electrons have a higher probability of being found with along the internucleus axis.

This simple picture of orbital overlap works well to describe the energies associated with the formation of covalent bonds but faces a challenge when you try to explain the geometry and number of bonds expected for multiple atoms.

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For example, CH4 should have a tetrahedral geometry with four bonded hydrogens surrounding the carbon atom; However, carbon only has two unpaired electrons to use in overlapping orbitals, not 4.

So then, how can we account for four bonds to carbon when the atomic orbitals for carbon will only allow for the sharing of two electron pairs according to Localized Electron Model. Also, how can the tetrahedral shape that is found experimentally be attained?

Hybridize the Orbitals!

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Atomic orbitals of C: 2p __ __ __

2s _____

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In 1930, Pauling described Orbital Hybridization to explain the problems that arise between the predicted electron arrangement (L.E. model) and the predicted molecular geometry (VSEPR theory)

4 C atom orbitals Hybridized Orbitals of hybridize to form C: four equivalent __ __ __ __ sp3 hybrid sp3 hybrid orbitals atomic orbitals.

Hybridization • Process that changes properties of valence electrons by mixing atomic orbitals to form special orbitals for bonding

atomic orbitals

Hybridized orbitals

Principles of Hybridization Orbital Hybridization

1. Conservation of orbitals 2. Hybrid orbitals correlate with molecular geometry 3. Energy level of hybrid orbitals is between that of AO’s 4. All bonded atoms hybridize and attain the lowest energy arrangement possible

All hybrid orbitals of an atom are said to be DEGENERATE (of equal energy)

Hybrid Orbitals • Consider beryllium:
In its ground electronic state, it would not be able to form bonds because it has no singly-occupied orbitals.
But if it absorbs the small amount of energy needed to promote an electron from the 2s to the 2p orbital, it can form two bonds.

• Mixing the s and p orbitals yields two degenerate orbitals that are hybrids of the two orbitals.
These sp hybrid orbitals have two lobes like a p orbital.
One of the lobes is larger and more rounded as is the s orbital.

• These two degenerate orbitals would align themselves 180°° from each other. • This is consistent with the observed geometry of beryllium compounds: linear

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Hybrid Orbitals sp hybridization linear species

• With hybrid orbitals the orbital diagram for beryllium would look like this. • The sp orbitals are higher in energy than the 1s orbital but lower than the 2p.

Using a similar model for boron leads to…

sp2 hybridization trigonal planar species

…three degenerate sp2 orbitals.

With carbon we get…

…four degenerate sp3 orbitals.

sp3 hybridization

tetrahedral species

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Orbital Orbital Hybridization Hybridization BONDS BONDS

SHAPE SHAPE

22

linear linear

sp sp

22 pp’’ss

33

trigonal trigonal planar planar

sp sp22

11 pp

44

tetrahedral tetrahedral sp sp33

15 15

HYBRID HYBRID REMAIN REMAIN

none none

2. Identify and draw the hybridization for formaldehyde, CH2O.

Where is the second bond? Remember, a bond is an overlap of electron densities between atoms

Sigma (σ σ) Bonds

1. Identify and draw the hybridization for ammonia, NH3.

Notice all bonds lie on the internuclear axis for each orbital.

Valence Bond Theory • With the inclusion of Resonance and Orbital Hybridization by Pauling, Localized Electron Model grew into the modern Valence Bond Theory.
Hybridization is a major player in this approach to bonding describing two ways orbitals can overlap to form bonds between atoms.

Pi (π π) Bonds • Pi bonds are characterized by

• Sigma bonds are characterized by
Head-to-head overlap.
Cylindrical symmetry of electron density about the internuclear axis.


Side-to-side overlap.
Electron density above and below the internuclear axis.

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Single Bonds Single bonds are always σ bonds, because σ overlap is greater, resulting in a stronger bond and more energy lowering.

Multiple Bonds • In a molecule like formaldehyde (shown at left) an sp2 orbital on carbon overlaps in σ fashion with the corresponding orbital on the oxygen. • The unhybridized p orbitals overlap in π fashion.

3.Draw the Lewis structure for ethene (C2H4), describe its hybridization and identify each bond in the molecule.

Multiple Bonds In a multiple bond one of the bonds is a σ bond and the rest are π bonds.

Multiple Bonds In triple bonds, as in acetylene, two sp orbitals form a σ bond between the carbons, and two pairs of p orbitals overlap in π fashion to form the two π bonds.

Consequences of Multiple Bonding

There is restricted rotation around C=C bond.

This gives rise to cis/trans, or geometric, isomerism for alkenes.

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σ and π Bonding in C2H2

H– C ≡ C –H

Delocalized Electrons: Resonance When writing Lewis structures for species like the nitrate ion, we draw resonance structures to more accurately reflect the structure of the molecule or ion.

Delocalized Electrons: Resonance This means the π electrons are not localized between the nitrogen and one of the oxygens, but rather are delocalized throughout the ion.

Investigate NO3-

In nitrogen trioxide, the electrons of the double bond are delocalized across the molecule. Due to the morphing of atomic orbitals to sp2 hybridized orbitals, the delocalized electrons stabilize the molecule. Therefore, the actual structure should be representative of all three Lewis structures.

Delocalized Electrons: Resonance • In reality, each of the four atoms in the nitrate ion has a p orbital. • The p orbitals on all three oxygens overlap with the p orbital on the central nitrogen.

Resonance The organic molecule benzene has six σ bonds and a p orbital on each carbon atom.

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Resonance • In reality the π electrons in benzene are not localized, but delocalized. • The even distribution of the π electrons in benzene makes the molecule unusually stable.

The Valence Bond theory, by its nature, does not account for resonance structures due to the delocalization of electrons, however, it does predict its existence for some molecules. But… It does provide a simple model to describe a visual picture of molecular structure through the use of: • Lewis Structures • VSEPR Theory • Orbital Hybridization • Resonance • Don’t forget bond polarity and any resulting molecular polarity according to molecular geometry.

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