Describe σ and π bonds
According to Lewis theory, covalent bonds consist of shared pairs of electrons, creating an area of electron density between the atoms. Also, atomic orbital theory states that electrons exist in atomic orbitals that are regions of space surrounding the nucleus that can only hold up to two electrons.
Valence bond theory defines a bond as a region of electron density between two atoms formed from two overlapping orbitals. These must be from different atoms, and each contain only one unpaired electron, and merge in the region between the atoms.
Sigma bonds (σ)
Sigma bonds are the result of axial overlap of orbitals, where two s or p orbitals overlap on the same axis. This explains why there is variation in the length on sigma bonds in different molecules, as it depends on the size of the orbitals
Pi bonds (π)
In a pi bond, the atoms have moved close enough for a p orbital to overlap along parallel axes. Therefore, a pi bond is defined as the sideways overlap of parallel p orbitals. However, the amount of overlapping possible here is limited, making pi bonds slightly weaker and sigma bonds.
Pi bonds are usually involved in double or triple bonds. A double bond is made up of one sigma and one pi bond, and a triple bond is made up of a sigma and two pi bonds.
Explain hybridisation in terms of the mixing of atomic orbital to form new orbitals to form new orbitals for bonding
Hybridisation was first noticed when scientists found that all the bonds in a CH4 molecule are the same length, yet a carbon atom
has two s orbitals and two p orbitals, which are different.
For the bonds to form, each orbital can only contain one electron, so the orbitals in the valence shell mix together and make hybrid
orbitals. These orbitals are shaped differently from their constituent orbitals, but there is still the same number of orbitals. Also, there is no energy change involved in forming the orbitals.
Hybridisation is the process of mixing atomic orbitals as atoms approach each other to form bonds
Identify and explain the relationships between Lewis structures, molecular shapes and types of hybridisation (sp, sp2 and sp3)
This is formed from the mixing of three p orbitals and an s orbital, creating four sp3 hybrid orbitals, such as in CH4.
These new hybrid orbitals are made up of two lobes, but one is significantly larger than the other. The four sp3 orbitals arrange into a tetrahedral shape with a bond angle of 109.5o.
Methane – CH4
Ethane – C2H6
Ammonia – NH3
Although it has the same type of hybridisation, ammonia has a bond angle of 107° because of the non-bonding pair of electrons, making it trigonal pyramidal in shape
Hydrazine – N2H4
Water – H2O
There are two non-bonding pairs of electrons, which means that the bond angle is 104.5°, making a V-shaped molecule
This is formed between one s orbital and two p orbitals, allowing for equal bonds in elements like boron (BH3 and BF3). These molecules have a trigonal planar shape. The hybrid orbitals can form sigma bonds with other atoms. This leaves a p orbital available to form pi bonds, making double bonds.
Ethene – C2H4
Only three of the orbitals are hybridised, with the remaining p orbital is able to form a pi bond, allowing for double bonds
Diazene – N2H2
These are formed from an s and a p orbital, making it linear in shape and a bond angle of 180 degrees. This leaves two p orbitals available to form pi bonds in triple bonds.
Beryllium Hydride – BeF2
Ethyne – C2H2