How do bonds form?

Valence Bond Theory Hybridization and VSEPR

Valence Bond Theory Valence bond theory describes covalent bond in terms of the overlap of atomic orbitals. H2:

This type of end-to end overlap of orbitals produces a sigma,σ bond

1s

1s

HF: 2p

1s

2p

1s

H

H

1s

Overlapping of the 1s orbitals

Other Points about Valence Bond Theory

• In hydrogen fluoride the 1s orbital of the H will overlap with the half-filled 2p orbital of the F forming a covalent bond. 2p

+

+

F

1s

• The newly combined orbital will contain an electron pair with opposite spin just like a filled atomic orbital.

2p

Example HF

H

Example: H2

Covalent Bond H-H

F2 :

1s

• The valence bond model or atomic orbital model was developed by Linus Pauling in order to explain how atoms come together and form molecules. • The model theorizes that a covalent bond forms when two orbitals overlap to produce a new combined orbital containing two electrons of opposite spin. • This overlapping results in a decrease in the energy of the atoms forming the bond. • The shared electron pair is most likely to be found in the space between the two nuclei of the atoms forming the bonds.

Overlapping of the1s and 2p orbitals

+

• This theory can also be applied to molecules with more than two atoms such as water. • Each covalent bond results in a new combined orbital with two oppositely spinning electrons. • In order for atoms to bond according to the valence bond model, the orbitals must have an unpaired electron.

Covalent Bond H-F



Bonding and Molecular Shape • When two atoms like hydrogen come together there is a precise distance between the two orbitals that ensures maximum overlap of the two orbitals • The need for maximum overlap is responsible for the different shapes of molecules found in nature - VSEPR

Valence Bond Theory and Molecular Geometry

1s 2p 1s

2s 2p

2s

2p

Orbital overlap suggests that the bond angle is 90, but we know that the angle is 104.5, therefore there must be a different orbital configuration.

What’s a Hybrid?

• The distance between the two nuclei in a bond is referred to as bond length • The two shared electrons of opposite spins spend most of their time between the two nuclei • Overlap of orbitals can be between like orbitals (s and s) or unlike orbitals (s and p)

Hybrid Orbitals

Consider the H2O molecule: Central atom O:

Bond length

• To make the connection between Quantum theory and VSEPR shapes, we theorize that the central atom hybridizes the available orbitals to achieve the required number of bonds and the correct molecular shape.

Hybrid Orbitals • Hybrid orbitals are mixtures of s, p, and d atomic orbitals with intermediate energies. • The number of s, p, and d orbitals that have combined, equals the number of hybrid orbitals. s

+

p sp sp



sp2 Hybrid Orbitals in BF3

sp Hybrid Orbitals in BeF2 Linear geometry is achieved using two sp hybrid orbitals. Be atom: Promotion: Hybridization:

s

p

p

p

s

p

p

p

sp sp

p

p

Trigonal planar geometry is achieved using three sp2 hybrid orbitals. B atom: Promotion:

sp

sp

Tetrahedral geometry is achieved using four sp3 hybrid orbitals.

Promotion: Hybridization:

s

p

p

p

s

p

p

p

sp3

sp3

sp3 sp3

p

p

s

p

p

p

sp2

sp2 sp2

sp2 sp2 sp2

p

Trigonal bipyramidal geometry is achieved using five sp3d hybrid orbitals. Octahedral geometry is achieved using six sp3d2 hybrid orbitals.

sp3 sp3

sp3

p

sp3d and sp3d2 Hybridization

sp3 Hybrid Orbitals in CH4

C atom:

Hybridization:

s

sp3

Types of Bonds

Types of Bonds - σ

• There are 2 main types of covalent bonds

• End to end overlap of orbitals (s, p, d, f, or hybrid) forms sigma, σ, bonds • The first bond between the central atom and a ligand is a sigma bond

– Single Bonds • Areas of electron density are concentrated between the nuclei of the bonding atoms (along the bond) • Bond is extremely strong

– Double and triple bonds • Areas of electron density are above and below the plane of the molecule • Bonds are highly reactive and easy to break



Covalent Bond Formation A  (sigma) bond results from end-to-end overlap of orbitals. The maximum electron density lies along the bond.

Types of Bonds - π • Double and triple bonds are pi, π, bonds • A single bond is composed of 2 areas of electron density above and below the sigma bond • extremely reactive because they are so far away from the influence of the nucleus and due to their location • π bonds are formed from regular p orbitals

π bond

Covalent Bond Formation

Hybridization in Carbon

A  (pi) bond results from side-to-side overlap of p orbitals. The electron density is zero along the bond.

• Carbon is able to form several different hybrid orbitals depending on how many other atoms it is bound to sp sp2

Overlap of py orbitals in O2:

C

sp3

sp2 hybridization in Carbon, C2H4

sp2 hybridization in Carbon, C2H4

• Recall – you only need hybrid orbitals for the first bond between 2 atoms! • The rest of the orbitals (the ones that will form the double bond) remain as normal p orbitals

• This hybridization forms a planar molecule with 120˚ angles between atoms.

C



sp hybridization in Carbon, C2H2

sp hybridization in Carbon, C2H2

• Each carbon needs 2 hybrid orbitals for the first bond with the C and H atoms. • The double and triple bond is formed from 2 normal p orbitals which will form 2 π bonds above and below the central plane of the molecule

• The unpaired electrons in the two p orbitals of the two adjacent carbon atoms share electrons by forming two π bonds.

sp hybridization in Carbon, C2H2 • This hybridization forms a linear molecule with 180˚ angles between atoms.

UNIT 2 Chapter 4: Chemical Bonding and Properties of Matter Section 4.2

Determining the Hybridization of the Central Atom of a Molecule or Ion

