Orbitals Flashcards
promotion energy
energy required to promote e- from ground state to valence state (= one type of excited state configuration used for bonding)
localised
e- are confined to a particular bond or atom
Lewis approach to bonding
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
octet rule
most main group atoms will tend to end up with ns2 np6 electronic configuration
importance of Lewis model
1st structure to describe bonding in structures of diatomics/polyatomics
introduced ideal of e- being shared
introduced bond order
problems with Lewis model
can’t easily explain radicals
difficult to explain expansion of octet
can’t explain why O2 = paramagnetic
paramagnetism
contains unpaired e-
weakly attracted by externally applied magnetic field
valence bond theory (VBT)
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-)
sigma bond
symmetrical about internuclear axis
pi bond
have node on internuclear axis + sign changes across axis
double bond
σ bond and π bond
triple bond
σ bond and 2 π bonds
diamagnetism
all e- paired
repelled by magnetic field
formation of π bond
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
bond length and bond order
higher bond order = shorter bond length = higher bond energy/stronger bond
sp3 hybridisation
occurs in molecules with tetrahedral shape
exhibits sp3 hybridisation on central atom
sp2 hybridisation
occurs in molecules with trigonal planar configuration
hybridised orbitals = σ bonds
unhybridised orbitals = π bonds
sp hybridisation
linear geometry
[examples]
BeH2
CO2 - sp on C and sp2 on O
PF5 hybridisation
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
octahedral geometry
d2sp3 hybridisation - 6 equivalent bonds 90 degrees from each other
MOs
orbitals that encompass entire molecule (constructed from AOs)
basic rules for molecular orbital theory
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
energy of bonding and anti-bonding MO
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)
forces in anti-bonding orbital
attractive forces outweighed by repulsive forces
net of raising energy (reduced stability) relative to 2 atoms being apart