Orbitals

energy required to promote e- from ground state to valence state (= one type of excited state configuration used for bonding)

e- are confined to a particular bond or atom

pairs of e- are localised in bonds or as non-bonding “lone pairs” on atoms each bond is formed by a pair of e- shared by 2 atoms

most main group atoms will tend to end up with ns2 np6 electronic configuration

1st structure to describe bonding in structures of diatomics/polyatomics introduced ideal of e- being shared introduced bond order

can’t easily explain radicals difficult to explain expansion of octet can’t explain why O2 = paramagnetic

contains unpaired e- weakly attracted by externally applied magnetic field

localised quantum mechanical approach to describe bonding in molecules provides mathematical justification for Lewis model of e- pairs making bonds between atoms thinks of bonds formed between 2 e- waves (bond will form if there is sufficient overlap of appropriate orbitals on 2 atoms + these orbitals are populated by a max. of 2 e-)

symmetrical about internuclear axis

have node on internuclear axis + sign changes across axis

σ bond and π bond

σ bond and 2 π bonds

all e- paired repelled by magnetic field

2px orbitals interact to form a π-bonding orbital (lies above and below plane) 2py orbitals interact in a similar way, giving a π-bonding orbital that lies at 90° to px

higher bond order = shorter bond length = higher bond energy/stronger bond

occurs in molecules with tetrahedral shape exhibits sp3 hybridisation on central atom

occurs in molecules with trigonal planar configuration hybridised orbitals = σ bonds unhybridised orbitals = π bonds

linear geometry [examples] BeH2 CO2 - sp on C and sp2 on O

exceeds octet nS and nP only provide 4 - need nd orbitals (ndz2) = 5 sp3dz2 2 sets - x2 axial (180 degrees from each other) + x3 equatorial (120 from each other and 90 from axial bonds) *not equivalent

d2sp3 hybridisation - 6 equivalent bonds 90 degrees from each other

orbitals that encompass entire molecule (constructed from AOs)

nAOs = nMOs to combine, AOs must have appropriate symmetry and energy each MO must be normal and orthogonal (90) to every other MO MO formed from weighted sums of AOs in-phase = adding atoms out-of-phase = subtracting atoms

bonding MO = lower in energy = more stable (increased e- density between overlap due to overlap reduces repulsion) anti-bonding MO = higher in energy = less stable (decreased e- density between overlap due to overlap increases repulsion)

attractive forces outweighed by repulsive forces net of raising energy (reduced stability) relative to 2 atoms being apart