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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 explains that a covalent bond forms when two electrons, each with opposite spins, come together in the area where atomic orbitals from two atoms overlap. This overlapping region has a high concentration of electrons. The stronger the overlap between these orbitals, the stronger the bond between the two atoms.
In a molecule, most electrons stay in their original positions from separate atoms.
Bonding electrons are concentrated where atomic orbitals overlap.
For orbitals with specific directional shapes, the strongest and most stable bond occurs when they overlap end to end. Greater overlap means a stronger bond.
Valence Bond Theory describes types of covalent bonds in molecules:
Sigma (σ) bond
Pi (π) bond
Sigma bonds are formed in three ways:
Overlapping of one s orbital with another s orbital.
Overlapping of one s orbital with a p orbital.
Head-on overlapping of two p orbitals.
Pi bonds result from sideways overlapping of two p orbitals.
Pi (π) Bonds
are characterized by :
Side-to-side overlap.
Electron density above and below the internuclear axis
Overlap of two half-filled orbitals leads to the formation of a covalent bond.
Overlap of an empty orbital with a fully-filled orbital leads to the formation of a co-ordinate covalent bond or dative bond
Molecular Orbital Theory (MOT)
:
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. In order for these orbitals to mix they must:
Have similar energy levels.
Overlap well (similar orientation).
Be close together
VBT
A molecule is a group of atoms bound together through localized overlap of valence-shell atomic orbitals
Each atoms in a molecule retain its AO in a molecule
MOT
A molecule is a collection of nuclei with the electron orbitals delocalized over the entire molecule
Mixing of AO forming MO with a given energy and shape that are occupied by the molecule’s electrons
Molecular Orbital (MO) 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.
Antibonding MO = region of diminished electron density
Bonding MO = enhanced region of electron density
Antibonding MOs, σ* = higher energy
Lower stability than AO
Bonding MOs, σ = lower energy
Higher stability than AO
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 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
PARAMAGNETISM
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
VSEPR Theory
Electron groups around the central atom will be most stable when they are as far apart as possible. This is called the
V
alence
S
hell
E
lectron
P
air
R
epulsion Theory, (VSEPR) Theory
Two geometries for each molecule :
Electron-group geometry is determined by the locations of regions of high electron density around the central atom
Molecular geometry or shape of a molecule is determined by the arrangement of atoms around the central atom
ELECTRON GROUP GEOMETRY
:
Ideal molecular geometry concept relies on valence electrons' repulsion between bonded atoms.
Electron groups consist of valence electrons localized around a central atom, and they repel other electron groups.
These mutual repulsions among electron groups determine the orientation of groups, known as electron-group geometry.
Electron group
2 (linear)
3 (trigonal planar)
4 (tetrahedral)
5 (trigonal bipyramidal)
6 (octahedral)
VSEPR NOTATION
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
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
Assume the C on the left is the 1st central atom, the C on the right is the 2nd and the O atom as the 3rd one
The 1st C has the AL4 electron arrangement, hence a tetrahedral geometry
The 2nd C has the AL3 electron arrangement, it has a trigonal planar geometry
The O atom has the AX2E2 electron arrangement, thus a bent geometry
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.
Hybridization of Atomic Orbitals
:
Based on ground-state electron configuration, carbon should have only two bonds
If a 2s electron is promoted to an empty 2p orbital, then four unpaired electrons can give rise to four bonds
These four orbitals become mixed, or hybridized to form bonds
Hybrid Orbitals
Note that there is a relationship between the type of hybrid orbitals and the geometric arrangement of those orbitals.
Thus, if you know the geometric arrangement, you know what hybrid orbitals to use in the bonding description.
sp3 hybridization
Occurs most often for central atom only
The total number of hybrid orbitals is equal to the number of atomic orbitals combined
Hybrid orbitals may overlap with pure atomic orbitals or with other hybrid orbitals
sp2 hybridization
This hybridization scheme is useful in describing double covalent bonds
Comprised of one 2s orbital and two 2p orbitals to produce a set of three sp2 hybrid orbitals
Determining Molecular Geometry
The geometric distribution of the three sp2 hybrid orbitals is within a plane, directed at 120o angles
This distribution gives a trigonal planar molecular geometry, as predicted by VSEPR
sp hybridization
This hybridization scheme is useful in describing triple covalent bonds
The geometric distribution of the two sp hybrid orbitals is on a line, directed at 180 angles
d Subshells Hybrid Orbitals
This hybridization allows for expanded valence shell compounds – typical for group 5A elements
This hybridization allows for expanded valence shell compounds – typically group 6A elements
Predicting Hybridization Schemes
:
In hybridization schemes, one hybrid orbital is produced for every simple atomic orbital involved
Write a plausible Lewis structure for the molecule or ion
Use the VSEPR method to predict the electron-group geometry of the central atom
Select the hybridization scheme that corresponds to the VSEPR prediction
Describe the orbital overlap and molecular geometry
A Problem to Consider
Describe the bonding in H2O according to valence bond theory. Assume that the molecular geometry is the same as given by the VSEPR model
From the Lewis formula for a molecule, determine its geometry about the central atom using the VSEPR model.
Lewis formula for H2O
From this geometry, determine the hybrid orbitals on this atom, assigning its valence electrons to these orbitals one at a time.
Note that there are four pairs of electrons about the oxygen atom.
According to the VSEPR model, these are directed tetrahedrally, and from the previous table you see that you should use sp3 hybrid orbitals
Each O-H bond is formed by the overlap of a 1s orbital of a hydrogen atom with one of the singly occupied sp3 hybrid orbitals of the oxygen atom
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
A double bond is made up of one sigma bond and one pi bond
A triple bond is made up of one sigma bond and two pi bonds
Compounds Containing Double Bonds (sp2)
The single 2p orbital is perpendicular to the trigonal planar sp2 lobes.
The fourth electron is in the p orbital.
Two sp2 hybridized C atoms plus p orbitals in proper orientation to form C=C double bond.
The portion of the double bond formed from the head-on overlap of the sp2 hybrids is designated as a s bond.
The other portion of the double bond, resulting from the side-on overlap of the p orbitals, is designated as a p bond.
Thus a C=C bond looks like this and is made of two parts, one and one bond
Compounds Containing Triple Bonds – C2H2
:
A bond results from the head-on overlap of two sp hybrid orbitals.
The unhybridized p orbitals form two p bonds.
Note that a triple bond consists of one and two p bonds.