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Chapter 5 : Theory of Chemical Bonding - Coggle Diagram
Chapter 5 : Theory of Chemical Bonding
Valence Bond Theory
Atomic Orbital Overlap
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.
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
Several Important Points
The theory may be used to describe (or predict) the kind of covalent bonds present in a molecule:
●Sigma (σ) bond
●pi (π) bond
The Sigma bond may be produced in 3 ways:
●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
The pi bond is produced by the sideways overlapping of one p orbital with another p orbital.
Pi () Bonds
Pi bonds are characterized by
•Side-to-side overlap.
•Electron density above and below the internuclear axis.
Molecular Orbital Theory
Molecular orbital theory is a theory of the electronic structure of molecules in terms of molecular orbitals, which may spread over several atoms or the entire molecule.
•Orbital Mixing – combination of atomic orbitals forming molecular orbitals.
•Have similar energy levels.
•Overlap well (similar orientation).
•Be close together
Paramagnetic
Paramagnetism 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
Molecular Orbital Configurations
The number of molecular orbitals (MOs) formed is always equal to the number of atomic orbitals combined.
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.
Electronic Configuration
Group IA-VA :
σ1s < σ1s
< σ2s < σ2s
< π2py = π2pz < σ2px < π2py
= π2pz
< σ2px*
Group VIA-VIIIA :
σ1s < σ1s
< σ2s < σ2s
< σ2px < π2py = π2pz < π2py
= π2pz
< σ2px*
The number of electrons is written as superscripts
Electronic configuration of B2 (total 10 electrons):
(σ1s)2 (σ1s
)2 (σ2s)2 (σ2s
)2 (π2py)1 (π2pz)1
The σ1s and σ1s* are generally not shown or written as KK, thus the configuration for B2 may be written as
KK (σ2s)2 (σ*2s)2 (π2py)1 (π2pz)1
Bond Order
This concept is used to indicate:
➢bond strength.
•Larger bond order means greater bond strength
➢number of bonds that exist between two atoms.
➢existence of molecules or ions
bond order = ½ (number of electrons in bonding Mos - number of electrons in antibonding Mos)
●Bond order equal to zero suggests that the molecule does not exist because the bonds are unstable
●Bond order of greater than zero suggests a stable molecule
●A fractional bond order suggests that the bond is relatively unstable compared to a whole number
There are definite correlations between bond order, bond energy, and bond length.
➢As the bond order increases, the bond energy increases and the bond length decreases
➢Large bond energy associated with N2 molecule, which the MOT model predicts will have a bond order of 3, a triple bond
VSEPR Theory
Valence Shell Electron Pair Repulsion Theory
Two geometries for each molecule
*Electron-group geometry is determined by the locations of regions of high electron density around the central atom(s)
*Molecular Geometry or shape of a molecule is determined by the arrangement of atoms around the central atom(s)
Electron Group Geometry
Ideal molecular geometry is based on the idea that pairs of valence electrons in bonded atoms repel one another
An electron group is any collection of valence electrons, localized in a region around a central atom, that repels other groups of valence electrons
The mutual repulsions among electron groups lead to an orientation of the groups that is called electron-group geometry
VSEPR Notation
The effect of nonbonding electrons
•Since electrons in a bond are attracted by two nuclei, they do not repel as much as lone pairs.
•Therefore, the bond angle decreases as the number of lone pairs increases
•Lone pair-lone pair repulsion > Lone pair-bond pair repulsion > Bond pair-bond pair repulsion
Multiple Bonds and Bond Angles
•Double and triple bonds place greater electron density on one side of the central atom than do single bonds.
•Therefore, they also affect bond angles.
Electron Pair and Molecular Geometry
•Draw the Lewis structure,
•Count the total number of electron pairs around the central atom,
•Arrange the electron pairs in one of the above geometries to minimize e--e- repulsion, and count multiple bonds as one bonding pair.
Geometry of Molecules of More Than One Central Atom
●Shapes of complex covalent molecules can be determined by analyzing one central atom at a time
Molecular Shape and Molecular Polarity
•When there is a difference in electronegativity between two atoms, then the bond between them is polar.
•It is possible for a molecule to contain polar bonds, but not be polar.
For example, the bond dipoles in CO2 cancel each other because CO2 is linear.
•In water, the molecule is not linear, and the bond dipoles do not cancel each other. Therefore, water is a polar molecule.
•The overall polarity of a molecule depends on its molecular geometry.
Hybrid Orbital Theory
●One might expect the number of bonds formed by an atom would equal its unpaired electrons.
○Chlorine generally forms one bond and has one unpaired electron.
○Oxygen, with two unpaired electrons, usually forms two bonds.
○However, carbon, with only two unpaired electrons, generally forms four bonds. For example, methane, CH4, is well known.
●One might expect the number of bonds formed by an atom would equal its unpaired electrons.
○Four unpaired electrons are formed as an electron from the 2s orbital is promoted (excited) to the vacant 2p orbital.
○More than enough energy is supplied for this promotion from the formation of two additional covalent bonds.
Hybrid Orbitals and Multiple Covalent Bonds
Covalent bonds formed by the end-to-end overlap of orbitals, regardless of orbital type, are called sigma (s) bonds. All single bonds are sigma bonds
A bond formed by parallel, or side-by-side, orbital overlap is called a pi (p) bond
