orbitals & hybridisation

• when atoms approach each other, their separate sets of atomic orbitals merge to form a single set of molecular orbitals
• some molecular orbitals, known as 'bonding molecular orbitals', occupy the region between two nuclei
• the attraction of positive nuclei to negative electrons occupying bonding molecular orbitals is the basis of bonding between atoms
• the molecular bonding theory describes how atomic orbitals merge and overlap when atoms bond to form molecules

  it shows how the valence of electrons are distributed within molecular orbitals spread across the molecule
• the bonding orbital is where the electrons of a different spin spend most of their time between two nuclei
• the bonding orbitals are where the electrons are most likely to be found
• the shape of a bonding orbital is determined by quantum mechanics
• the bonding molecular orbital has a lower energy than the separate atomic orbitals

  e.g. hydrogen molecule is more stable than the hydrogen atom
• the number of molecular orbitals is the same as the number of atomic orbitals that combine

  e.g. if we have 2 atomic orbitals that we have 2 molecular orbitals
• antibonding orbitals have a higher energy than the bonding orbitals, they also have a different shape and do not contain electron
• there are 3 different types of molecular bonding orbitals that come about by end-on overlapping

  these are sigma bonds
• s orbital + s orbital

  = sigma bond (s-s overlap)
• s orbital + p orbital

  = sigma bond (s-p overlap)
• p orbital + p orbital

  = sigma bond formed by end-on-end overlap of 2 p orbitals
• a covalent bond is formed when two half-filled orbitals overlap, and if they overlap along the axis of the bond (end on), it is known as a sigma bond
• p orbitals can also overlap side on, which is a pi bond
• pi bonds arise when an atom makes multiple bonds
• a pi bonds occurs when atomic orbitals lie perpendicular to the bond and overlap side on
• end-to-end overlap is more efficient than side-on overlap, therefore, sigma bonds are stronger than pi bonds
• in it's ground state, carbon has the electron configuration of 1s2 2s2 2p2

  it has two half-filled orbitals in the 2p subshell, which leads us to expect it to form 2 bonds instead of 4 as it only has 2 unpaired electrons

  but there's only a small energy difference between 2s and 2p subshells

  it makes it easy for an electron to be promoted from the 2s to the 2p, which creates 4 atomic orbitals with unpaired electrons

  therefore, carbon has a valency of 4 and can form 4 bonds
• spectroscopy shows molecular orbitals are all the same shape, which must be due to the mixing of the two types of orbitals, s and p, to form hybrid orbitals this is hybridisation
• sp3

  this hybrid is named this because its made by combining 1 s and 3 p orbitals
• the 4 sp3 hybrid orbitals are degenerate and identical to one another
• sp3 hybridisation is found in saturated hydrocarbons
• in alkenes, sp2 hybridisation is found
• sp2

  this is when 1 s orbital is hybridised with 2 of the p orbitals, and 1 of the p orbitals is not involved in hybridisation

  3sp2 orbitals are degenerate and singly filled

  the sp2 orbitals form sigma bonds, the unhybridised p orbitals overlap to form a pi bond
• in alkynes a triple carbon-carbon bond exists which is called an sp hybridisation
• an sp hybridisation is formed by mixing one of the 2s orbitals with one 2p orbital to form two sp orbitals

  which leaves two of the 2p orbitals unhybridized

  there are two p orbitals on each carbon that are not hybridised, which overlap to form pi bonds
• carbon compound: c-c

  hybridisation: sp3

  bond type: sigma
• carbon compound: c=c

  hybridisation: sp2

  bond type: 1 sigma and 1 pi
• carbon compound: c-=c

  hybridisation: sp

  bond type: 1 sigma and 2 pi