Chapter 8 - Molecular Orbital Theory! Flashcards
MO theory general description
treats a covalent bond as a pair of electrons w/ bonding orbital derived from the overlap of orbitals of adjacent atoms
paramagnetic
unpaired electrons, interacts with magnetic fields
diamagnetic
no unpaired electrons, weakly influenced by magnetic field
valence bond theory
localized application of MO theory to specific bonds
superposition of waves
combined to produce wave with amplitude equivalent to sum of other two - or changing electron density
covalent bond forms when
waves superimpose and increased electron density between nuclei
opposite vs same phases and what time of orbitals are formed
2 wavefunctions w/ same phase (++ or –), constructive intereference forms a BONDING ORBITAL with increased electron density between the nuclei
wavefunctions w opposite phases, destructive interference forms ANTIBONDING ORBITAL, reduces electron density so works against bonding. nodal planes could form between nuclei
bonding orbital
wavefunctions w same phase from constructive interference, increased electron density
when drawing, combine so pretty much two same phase circles slowly combine to form one bigger circle
lower energy
antibonding orbital
wavefunctions w different phases destructively interfere to reduce electron density
nodal planes can form between nuclei
when drawing they dont really fully combine, like two diff coloured circles approaching, remain different just change shape, more electron density away from the other
higher energy
where does optimal bond length occur?
at the distance with lowest possible energy, depends on identity of bonding atoms
what does greater overlap in orbitals mean?
greater difference in energy between bonding and antibonding orbitals
increases and decreases in energy w distance, antibonding and bonding
bonding orbitals - energy decreases as internuclear distance decreases
antibonding - energy increases as internuclear distance decreases
affect repulsion between nuclei on PE
repulsion increases PE at short internuclear distances
what would a representation of MOS using cross sections of 90% probability surface look like?
draw two atoms with 90% prob surface, ex for 1s in Hydrogen it would be two circles, with phase noted + or -, draw a + sign between and an arrow pointing to result, if same phase they would connect to form a kind of oblong orbital with two dots for the two nuclei, and indicate if + or -, if opposite phases, would end up funky and asymetrical, part of the 90% prob surface would be super close to the nuclei, and space in between
rules for forming MOs
- total MOs formed is equal to # atomic orbitals interaction
- when 2 AOs combine, resulting bonding MO fewer nodes than antibonding MO
- the more the AOs overlap the larger the energy gap bonding and antibonding orbitals
- only AOs of similar energy interact significantly (ex 1s & 1s)
- prob of finding an electron @ x = wavefunction squared
- phase changes at each node
- prob of finding electron @ node = 0
how do bonding and antibonding orbitals differ from atomic orbitals?
bonding MO is lower in energy comp to AO, antibonding has greater energy by pretty much the same amount
MO with more nodes higher or lower energy?
higher energy than those with few nodes
construction of an orbital interaction diagram
atomic orbitals of seperate atoms shown on right and left side, the orbital name written under a line.
bonding and antibonding orbitals are drawn between them at relative energy levels, * rep antibonding, bonding below AO, antibonding above
dotted lines connect ao to bonding and antibonding, showing where electrons come from
labelling of orbital diagrams
- symmetry of orbital, the o with a line out the right top (sigma) for symmetric w respect to rotation about internuclear axis, pi if asymmetrical (one nodal plane containing internuclear axis)
- whether bonding or antibonding (* for antibonding)
- AO which MO originates from in subscript ex pi1s, 1s in subscript
how to write electronic configuration of a molecule
ex for hydrogen it would be (sigma symbol)1s in subscript^2
how does MO theory and orbital interaction diagrams show why H2 exists instead of H
energy level of bonded electrons in H2 is lower than the energy of the electrons unbonded, therefore energetically favourable
equation for bond order of diatomic molecule w particular molecular configuration
Bond order = (e bonding - e antibonding)/2
where e bonding is # of electrons in all bonding orbitals
e antibonding is # of electrons in all antibonding orbitals
steps for filling orbital interaction diagram w electrons
- calculate total # electrons in AOs
- Fill MOs w/ number of electrons calculated in step 1
- fill lowest energy orbitals first (aufbau)
- place max of 2 spin paired electrons each MO (pauli)
-electrons in degenerate (same energy) MOs singly, w/
same spin before pairing (hund’s rule)
why doesn’t He2 exist?
antibonding = 2, # bonding = 2, therefore bond order = 0 so no bonds form and the element doesn’t exist
homonuclear diatomic species
molecules made of exactly 2 identical atoms
second row homonuclear diatomic species energy level diagrams
same sort of set up as fist row, first set up energy levels of seperate atoms, big space between 1s and 2s, do the dotted line thing, write proper amount of electrons into original AOs then transfer to MOs
molecular orbitals that include p block
s block set up the same as normal, space between 1s 2s etc
in p block, start the same, set up ao configuration diagrams for the seperate atoms starting with 2p
assume x is the internuclear axis, therefore px would be symmetrical and sigma, py and pz would be asymmetrical and therefore pi
signma px would have the lowest energy of the ps, then pi pz and py, then pi star pz and py then sigma star px
recap - what leads to greater energy difference between bonding and antibonding?
more overlap of AOs, orbitals with higher n value would have more overlap and therefore a greater energy difference
HOMO
highest occupied molecular orbital - highest energy MO that contains any electrons
LUMO
lowest unoccupied molecular orbital - molecular orbital next higher energy after highest occupied MO
order that MOs fill up for F2
sigma 2s then sigma star 2s then sigma 2p then pi 2p then pi star 2p then sigma star 2p
order that MOs fill up for O2
sigma 2s then sigma star 2s then sigma 2p then pi 2p then pi star 2p then sigma star 2p
order that MOs fill up for N2
sigma 2s then sigma star 2s then pi 2p then sigma 2p then pi star 2p then sigma star 2p
order that MOs fill up for C2
sigma 2s then sigma star 2s then pi 2p then sigma 2p then pi star 2p then sigma star 2p
order that MOs fill up for B2
sigma 2s then sigma star 2s then pi 2p then sigma 2p then pi star 2p then sigma star 2p
MO energy levels change across a period
energy levels decrease from left to right across a period bc increase in Z and Zeff
rationale for change in relative energies of sigma 2p and pi 2p orbitals between N2 and O2
explained partially from looking at relative energies of 2s and 2p orbitals. when energies of 2s and 2p orbitals similarly mixing could happing, meaning sigma orbitals not derived from “pure” 2s-2s and 2px-2px but some 2px-2s
unsettles energy of all sigma orbitals therefore energy of sigma 2p is higher than pi 2p when 2s and 2p are similar in energy
