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Chapter 5 (Theory of Chemical Bonding) - Coggle Diagram
Chapter 5 (Theory of Chemical Bonding)
Valence Bond Theory, VBT
Atomic Orbital Overlap
This overlap region has a high electron charge density
The more extensive the overlap between two orbitals, the stronger is the bond between two atoms
Valence Bond (VB) Theory states that a covalent bond is formed by the pairing of two electrons with opposing spins in the region of overlap of atomic orbitals between two atoms.
Types of Bonds
Sigma Bond
How to produce?
Overlapping of one s orbital with another s orbital
Overlapping of one s orbital with a p orbital
Head on overlapping of one p with another p orbital
Pi Bond
How to produce?
The pi bond is produced by the sideways overlapping of one p orbital with another p orbital.
What does it state?
A molecule is a group of atoms bond together through localized overlap of valence-shell atomic orbitals
Each atoms in a molecule retains its atomic orbital, AO in a molecule
Molecular Orbital Theory
Valence Shell Electron Pair Repulsion, VSEPR
Problems faced when using this model
It incorrectly assumes that electron are localized and so the concept of resonance must be added
The model does not deal effectively with molecules containing unpaired electrons.
The model gives no direct information about bonding energies
What does it state?
A molecule is a collection of nuclei with the electron orbital delocalized over the entire molecule
Mixing of atomic orbital, AO forming molecular orbital, MO with a given energy and shape that are occupied by the molecule's electrons
Molecular Orbital Configuration
The number of molecular orbitals (MOs) formed is always equal to the number of atomic orbitals combined.
Example
2s Orbital
2p Orbital
The more stable the bonding MO, the less stable the corresponding antibonding MO.
The filling of MOs proceeds from low to high energies.
Each MO can accommodate up to two electrons.
Use Hund’s rule when adding electrons to MOs of the same energy.
The number of electrons in the MOs is equal to the sum of all the electrons on the bonding atoms.
Bond order
Formula to determine the bond & strength
(No. of electrons in bonding MOs - No. of electrons in antibonding MOs) / 2
Higher number indicates stronger bond, higher stability and no. of bond formed
When calculation = 0, the bond/ molecule does not exist
A fractional bond order suggests that the bond is relatively unstable compared to a whole number
Trend
The stronger the bond, the shorter the bond
Magnetic Property
Paramagnetic causes substance to be attracted into a magnetic field
Associated with unpaired electrons
Diamagnetism causes substances to be repelled from the magnetic field
Associated with paired electrons
Valence Shell Electrons Pair Repulsion, VSEPR
In the VSEPR notation used to describe molecular geometries, the central atom in a structure is denoted as A, terminal atoms as X, and the lone pairs of electrons as E
Electron-group geomeries
4
Tetrahedral, Angle = 109.5
3
Trigonal Planar, Angle = 120
2
Linear, Angle = 180
Not including lone pair elecrons
5
Trigonal bipyramidal, Angle = 120, 90
6
Octahedral, 90
Possible molecular group
2
0 lone pair
Linear
3
0 lone pair
Trigonal Planar
1 lone pair
bent or V-shaped
4
0 lone pair
Tetrahedral
1 lone pair
Trigonal pyramidal
2 lone pair
Bent or V-shaped
5
0 lone pair
Trigonal bipyramidal
1 lone pair
seesaw-shaped
2 lone pair
T-shaped
3 lone pair
Linear
6
0 lone pair
Octahedral
1 lone pair
Square pyramid
2 lone pair
Square planar
3 lone pair
T-shaped
4 lone pair
Linear
Strenght of repulsion order
Lone pair-lone pair repulsion > Lone pair-bond pair repulsion > Bond pair-bond pair repulsion
Hybrid orbital theory
sp
Linear
sp2
Trigonal Planar
sp3
Tetrahedral
sp3d
Trigonal Bipyramidal
sp3d2
Octahedral
