Week 8 Molecular geometry and bonding theories Flashcards
Lewis Model Deficiences
With Lewis dot structures, both H2 and F2 are treated equivalently, i.e. both are sharing a single electron pair to form a single bond.
Molecular Orbital Theory
suggests that the shape of a molecule is defined by the protons of the atoms and the electrons then fill molecular orbitals that are defined by the whole molecule.
Valence Bond Theory
VSEPR modelled predicted shapes of molecules but it does not explain why bonds form between atoms. Using quantum mechanics we can consider how atomic orbitals can be used to account for the shape of molecules. Combining Lewis approach of electron sharing with the quantum mechanical idea of atomic orbital lead to a theory called Valence Bond Theory.
Sigma bonds
When two atoms come together to form a bond, there is an overlap of their orbitals which give rise to a chemical bond; the first bond that forms is called a sigma bond. This results from direct overlap of orbitals. The images below show how the quantum mechanical orbitals combine; first two s-orbitals, then two p-orbitals (end-on) and finally an s-orbital and a p-orbital.
Hybridisation
To consider covalent bonds of one atom with many others, need to introduce the idea of hybrid orbitals. Valence bond theory assumes that when covalent bonds are formed atomic orbitals mix or hybridize to form new sets of orbitals. i.e. we see a mixing between atomic orbitals.
Two orbital combine together to form two hybrid orbitals; each consisting of 50% of the s-orbital and 50% of the p-orbital. We call this an sp hybrid orbital.
Hybridisation - sp hybrid
Hybridisation - sp hybrid
If we take 1 electron from the 2s and put it into an empty 2p orbital we now have two unpaired electrons and there is a chance of two bonds being made with these two unpaired electrons
Combining the 2s orbital with one of the 2p orbitals creates two equivalent sp hybrid orbitals into which the electrons from the two hydrogen atoms can be added.
The hybrid orbital has lower energy that the 2p orbital so the hybrid orbitals fill first and the remaining two 2p orbitals remain unfilled.
sp2 Hybridisation
If we mix one s-orbital with two p-orbitals, we obtain three equivalent sp2 hybrid orbitals; allowing for the formation of 3 equivalent bonds. Below is a diagram to demonstrate how the promotion of one s-orbital electron into an empty p-orbital demonstrates how boron can form three identical bonds despite also being an exception to Lewis’ octet rule.
Lewis exceptions - Exceeded octet
Hybridisation can also involve d-orbital when the central atom is in period 3 and beyond of the periodic table. Mixing one s-orbital, three p-orbitals and 1 or more d-orbitals creates sp3d or sp3d2 hybrid orbitals accounting for the trigonal bipyramidal and octahedral molecules determined by VSEPR.
Steps to predict the hybridisation used by the central atom in a molecule
1 Draw the Lewis structure of the molecule or ion,
2
Determine the electron domain geometry using VSEPR theory,
3
List the hybrid orbitals needed to fit the electron pairs according to the shapes in Table 9.4 above.
Multiple bonds
If we have two p-orbitals and we combine them side-on, they form 2 regions of electron density – one above and one below the plane of bonding (the interatomic plane). Remember that end-on overlap of two p-orbitals results in a sigma (σ) bond.
In a molecule this allows a sigma bond to exist along the axis of bonding (with the electron density between the two atoms) and then have a second pi bond where the electron density is above and below the plane of the bond.
Side-on orbital overlap is not as effective as end-on orbital overlap so a sigma (σ) bond will always be stronger than a pi (π) bond. Consequently a pi bond will always be easier to break and therefore double bonds are always more reactive than single bonds.
Double Bonds
Bonds formed by the sideways overlap of 2 p orbitals are called pi (π) bonds.
The sideways overlap of orbitals to form π bonds is not as effective as the end-on overlap of orbitals to form σ bonds. For this reason σ bonds are stronger than π bonds.
Triple Bonds
Nitrogen Molecule (N2)
One bond is a sigma bond along the axis, one is a pi bond pointing along one axis and the other is a pi bond at right angles to the fist.
There are two electrons in each axis direction (x, y, z) which allows 6 electrons to be in such close proximity.
Hybridisation in carbon molecules
C is a universal building block for chemistry. It can control the structure of the molecules it forms by controlling the number of sigma and pi bonds it forms. In its ground state, the valence shell of carbon contains four electrons 2s2 2p2.
There are three ways in which carbon can form hybrid orbitals; sp, sp2 and sp3.
Summary of Valence Bond Theory
1
Every bonded pair of atoms shares at least one pair of electrons; the number of hybridised orbitals required to form bonds between the central atom and its neighbouring atoms can be determined by the molecular geometry.
2
The first shared pair of electrons between any two atoms are found in the space between the two atoms creating a σ bond.
3
When atoms share more than one pair of electrons, the first pair form a σ bond and additional pairs form π bonds, with the electron density situated above and below the sigma bond.
Molecular Orbital Theory successfully predicts:
Accurate structures (bond lengths, angles)
Bond dissociation energies
Line positions and intensities in molecular spectra
Paramagnetism, diamagnetism and
Quantitative estimates for dipole moments and electronegativity
Further, Molecular Orbital Theory (M.O.T.) does not need resonance structures to describe molecules. M.O.T. treats molecular bonds as a sharing of electrons between nuclei i.e. is treats electrons as being spread out across the entire molecule - the electrons are delocalised.
Determining whether a bond forms
Video on slide “First period diatomic molecules”
In Lewis theory we considered single, double and triple bonds using sigma or pi bonding. In Molecular Orbital Theory we talk about Bond Order instead which similarly determines how atoms bond together.
To calculate bond order we use the equation below. Note: the bond order must be a positive number, greater than zero, for a bond to form.
A bond order of 1 indicates a single bond
A bond order of 2 represents a double bond
A bond order of 3 indicates a triple bond
Because M.O.T. can be used to look at the Lewis exceptions with ‘odd numbers’ of electrons too, it is therefore possible to calculate a bond order which is a fraction.
How to determine electron distribution in diatomic molecules (Has video on slide)
1
Note that the number of molecular orbitals formed equals the number of atomic orbitals combined.
2
Atomic orbitals combine most effectively with other atomic orbitals of similar energy.
3
The effectiveness with which the two atomic orbitals combine is proportional to their overlap. i.e. as the overlap increases, the energy of the bonding molecular orbitals is lowered and the antibonding molecular orbitals is raised.
4
Each molecular orbital can accommodate at most 2 electrons (paired spins).
5
When filling molecular orbitals of the same energy, Hund’s rule applies.
Period 2 Diatomic Elements
The outer most orbitals of the second row elements used in bonding are the 2s orbital (of which there is one) and the 2p orbitals (of which there are three - 2px, 2py, 2pz).
When these orbitals from two second row elements overlap, three bonding and three anti bonding molecular orbitals are formed.
2s and 2p Orbital Interactions
It is possible for interaction between the 2s orbital of one atom and the 2p orbital of another. These interactions cause the σ2s orbital energy to fall and the σ2p energy to rise. This accounts for the changed energetic molecular orbital ordering we see for B2, C2 and N2.
Determine the bond order for the N2 molecule
Check slide “Second period Diatomic molecules”
Heteronuclear diatomic molecules
All the examples we have looked at so far, are homonuclear molecules i.e. both atoms in the molecule are the same. With heteronuclear molecules (two different atoms combined), the electronegativity of the two atoms is important. If the two atoms have similar electronegativity, the energy diagram is similar to the homonuclear diatomic molecule diagram. This would be the case for nitric oxide (NO).
Molecular Orbital Theory can be used to predict the bond order of these molecules in the same way it does for homonuclear molecules. In this case, it is more accurate than that predicted for NO using Lewis structures as NO was an exception to the Lewis Octet Rule (odd number of electrons).
