Orbitals and Bonding Flashcards
atomic orbital
area of space where an electron is
principle quantum number
shell
1,2,3 etc.
angular momentum quantum number
shape
s,p,d
magnetic quantum number
orientation - which way it points in space
e.g. 2px, 2pz, 2py
probability of finding an e- in 1s orbital
greatest at vicinity of nucleus
decreases exponentially as the distance from the nucleus increases
represented as a spherical cloud of charge
nodal plane
zero probability of finding an electron
orbitals - wave properties
have both amplitudes and phases (+ or -)
what is the density of an electron in an orbital proportional to?
the square of the atomic orbital
probability of finding electron in particular region of space = always positive
how many electrons can each atomic orbital occupy?
2 - must have different spins
atomic core
complete shells
valence shell
partially filled shell
covalent bonding
atomic orbitals in different atoms overlap and electrons pair up with opposite spins
mutual repulsion among themselves and attraction to the nuclei of all atoms
arrange themselves to occupy MOLECULAR ORBITALS
where is the electron density concentrated in covalent bonding?
the region between the 2 nuclei - where atomic orbitals overlap most
what does the overlapping of 2 atomic orbitals produce?
2 molecular orbitals
bonding and antibonding (opposite phases)
antibonding MO subtracts the AO instead of adding them
which has more energy - antibonding MO or bonding MO?
antibonding MO (has highest energy overall - higher than bonding MO and separated atoms)
what are the 2 halves/lobes of the antibonding MO separated by?
nodal plane
which is more stable - H2 or isolated H atoms?
H2 molecules - covalently bonded
sigma bond
formed when 2 AOs overlap along the interatomic axis
pi bond
formed when AOs overlap sideways
which is stronger - sigma or pi bonds?
sigma
given interatomic distance, the overlap of carbon 2p AOs is significantly better when they approach head-to-head rather than side-to-side
x-ray crystallography
enables determination of structure of crystalline material
x-rays are diffracted by interactions with the electrons rather than protons/neutrons
produces electron density map
sum of densities of each electron in each AO =
total electron density of atom
x-ray crystallography - covalent bonds
concentrations of electron density between atoms
shape of molecules - long range
electron density has low values
shape is distorted
only see Van der Waals contact
shape of molecules - short range
electron density has high values
atoms and bonds are more evident
hybridisation
atoms in molecules can change shape and energies of their AOs in order to provide the best overlap and max. bonding with other atoms
why isn’t methane planar?
= regular tetrahedron
repulsive interactions between pairs of electrons in the 4 bonds are minimised by maximising the angles between bonds
= valence-shell electron-pair repulsion (VESPR) theory
sp3 orbitals
identical
directional - pointing towards corners of a tetrahedron
well-suited for good overlap with four 1a AOs of H atoms
why do 4 sp3 orbital hybrids form?
minimises repulsion as shape is symmetrical - electrons position themselves as far away as possible
formation of C-H bonds
overlap of an sp3 hybrid with a hydrogen 1s AO with pairing of electrons with opposite spins
combination of 2 orbitals on C and H atoms = 2 orbitals for each C-H bond:
bonding σ orbital
bonding σ* orbital
energies of the bonding and antibonding orbitals
bonding = closer in energy to the 1s AO
antibonding = closer in energy to sp3 AO
true molecular orbitals for methane
obtained by combining all 4 sp3 hybrids on C with all 4 1s AO on H
delocalised over whole molecule
4 equivalent CH bonds, each with its own bonding σ orbital = valence bond theory
methyl cation bonding - CH3+
C has only 6 e- in its valence shell
CH3+ has planar geometry with 3 equivalent C-H σ bonds => trigonal arrangement (bond angle = 120)
overlap with 3 1s AOs requires 3 equivalent hybrid orbitals on C ∴ C is sp2 hybridised
leaves 1 2p AO on C unused - unoccupied AO perpendicular to plane of CH3
σ and π bonds
σ comes from orbitals with lower energy
π will break before σ bond
ethyne bonding - C2H2
linear geometry (bond angle = 180) with each C forming one C-H σ bond with 2 C-C π bonds
each carbon is diagonal to minimise repulsion between 2 bonds
each carbon is so hybridised
