What is Bond Hybridisation Theory and why does it matter?
In a previous post I discussed the bond angles in molecules like methane CH4.
Bond hybridisation theory seeks to offer explanations for these bond angles and bond lengths in a way that tries not to violate the actual observed facts about molecules like methane and ethene.
So let's start with what we can observe about methane and ethene and then try and relate those facts to the atomic electron structures of carbon and hydrogen.
Here are in picture form some of those observed features of both methane and ethene molecules.
Methane
Ethene
Ethene has a trigonal planar shape about each carbon atom i.e. it is a flat molecule.
Lets look closely now at Methane.
A couple of things stand out.
Methane has four equivalent covalent bonds.
Each is 108.70pm long and therefore each is of the same energy.
Each is a C-H bond and it has been calculated and shown experimentally that a mole of these bonds takes 416 kJ to break.
How can we explain this data?
When we examine the electron arrangement of carbon.
The lowest energy carbon atom has its 6 electrons in the following configuration:
Carbon 1s2 2s2 2p2
The four electrons in the second energy level are in two sub shells of slightly different energy.
Now there's the issue.
These four electrons to remain in these two subshells would not make four covalent bonds with hydrogen electrons that were of equal energy.
The four electrons would have to be in the same energy level for that to happen - but they are not.
In hybridisation theory, hybrid atomic orbitals are assumed to form from the 2s and 2p atomic orbitals when the hydrogen atoms bond to the carbon atom and create the methane molecule.
This effectively puts the carbon atom in an energy level higher than its ground state ( or lowest energy level.)
Diagrammatically its often drawn like this:
In this model an electron is promoted to the empty 2p orbital and then the 2s and the three 2p orbitals "hydridise" to form four degenerate or equal energy atomic orbitals in the carbon outer shell.
These four atomic orbitals are designated sp3 because each one is formed from an s orbital and three p orbitals.
Each of these sp3 orbitals can now accept one electron from each hydrogen atom and the molecule of methane is formed.
Thus this model accounts for the formation of four equivalent C-H bonds around the carbon atom in methane.
The model is not the reality it is one of the best representations of the reality of the methane molecule.
The model's capacity to predict what we observe about the properties of methane is what makes it very useful.
The shapes of these sp3 orbitals have been calculated from studies of the quantum mechanics of the methane molecule.
So a spatial representation of the carbon atom in methane looks like this:
This diagram shows us a model of the methane molecule that is the right shape with the right bond angles between the four C-H bonds.
Notice too that combining the sp3 atomic orbitals of the carbon atom and the four 1s hydrogen atomic orbitals leads to the structure of the methane molecule.
In the model molecular orbitals (MO) result from the overlap of atomic orbitals.
As the ends of these atomic orbitals overlap the resulting molecular orbital is called a sigma σ
molecular orbital.
As the diagram notes the molecular orbital has σ symmetry i.e. it is symmetrical about the central axis that connects both hydrogen and carbon atoms.
In a previous post I discussed the bond angles in molecules like methane CH4.
Bond hybridisation theory seeks to offer explanations for these bond angles and bond lengths in a way that tries not to violate the actual observed facts about molecules like methane and ethene.
So let's start with what we can observe about methane and ethene and then try and relate those facts to the atomic electron structures of carbon and hydrogen.
Here are in picture form some of those observed features of both methane and ethene molecules.
Methane
As we know a methane has a tetrahedral shape
Ethene
Lets look closely now at Methane.
A couple of things stand out.
Methane has four equivalent covalent bonds.
Each is 108.70pm long and therefore each is of the same energy.
Dr Carr’s Rescue Box
This box occurs just when
you need it to explain some unusual appearance on one of my blogs. In this case the bond
length of the C-H bond is 108.7 pm.
What is a "pm"? It is a pico metre. Divide a meter into a
million million bits and each bit is a picometer wide or 10-12 m.
Hope that helps. It is pretty small!!
Each is a C-H bond and it has been calculated and shown experimentally that a mole of these bonds takes 416 kJ to break.
How can we explain this data?
When we examine the electron arrangement of carbon.
The lowest energy carbon atom has its 6 electrons in the following configuration:
Carbon 1s2 2s2 2p2
The four electrons in the second energy level are in two sub shells of slightly different energy.
Now there's the issue.
These four electrons to remain in these two subshells would not make four covalent bonds with hydrogen electrons that were of equal energy.
The four electrons would have to be in the same energy level for that to happen - but they are not.
In hybridisation theory, hybrid atomic orbitals are assumed to form from the 2s and 2p atomic orbitals when the hydrogen atoms bond to the carbon atom and create the methane molecule.
This effectively puts the carbon atom in an energy level higher than its ground state ( or lowest energy level.)
Diagrammatically its often drawn like this:
In this model an electron is promoted to the empty 2p orbital and then the 2s and the three 2p orbitals "hydridise" to form four degenerate or equal energy atomic orbitals in the carbon outer shell.
These four atomic orbitals are designated sp3 because each one is formed from an s orbital and three p orbitals.
Each of these sp3 orbitals can now accept one electron from each hydrogen atom and the molecule of methane is formed.
Thus this model accounts for the formation of four equivalent C-H bonds around the carbon atom in methane.
The model is not the reality it is one of the best representations of the reality of the methane molecule.
The model's capacity to predict what we observe about the properties of methane is what makes it very useful.
The shapes of these sp3 orbitals have been calculated from studies of the quantum mechanics of the methane molecule.
So a spatial representation of the carbon atom in methane looks like this:
This diagram shows us a model of the methane molecule that is the right shape with the right bond angles between the four C-H bonds.
Notice too that combining the sp3 atomic orbitals of the carbon atom and the four 1s hydrogen atomic orbitals leads to the structure of the methane molecule.
In the model molecular orbitals (MO) result from the overlap of atomic orbitals.
As the ends of these atomic orbitals overlap the resulting molecular orbital is called a sigma σ
molecular orbital.
As the diagram notes the molecular orbital has σ symmetry i.e. it is symmetrical about the central axis that connects both hydrogen and carbon atoms.
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