Molecular orbital theory. Overcoming the shortcomings of the valence bond
Molecular orbital theory Overcoming the shortcomings of the valence bond
Learning objectives Describe basic principles of MO theory Describe dif...
Molecular orbital theory Overcoming the shortcomings of the valence bond
Learning objectives Describe basic principles of MO theory Describe differences between Valence Bond and MO theories Write MO diagrams for some simple diatomic molecules Explain optical and magnetic properties of O2 using MO theory
Shortcomings of valence bond The orbitals still maintain atomic identity Bonds are limited to two atoms Cannot accommodate the concept of delocalized electrons – bonds covering more than two atoms Problems with magnetic and spectroscopic properties
Molecular orbital theory: wavefunctions revisited The wave function describes the path of the electron – ΨA (has no real physical meaning) Wave functions have phase – indicated by “+” and “-” Approach of atoms causes overlap of orbitals + adds to + (constructive interference) + subtracts from – (destructive interference)
Wavefunctions and electron density Ψ describes the electron path Ψ2 describes the electron density
Orbital ΨA and ΨB overlap to form bond Molecular wavefunction (ΨA + ΨB) Joint density is (ΨA + ΨB)2 = ΨA2 + ΨB2 + 2ΨAΨB In molecular orbital the density is greater between the nuclei by an amount 2ΨAΨB
Molecular orbital theory: bonding and antibonding Bonding orbital: additive combination of atomic orbitals σ Antibonding orbital: subtractive combination of atomic orbitals σ*
Linear combination of atomic orbitals Valence Bond theory Hybrid orbitals made using weighted average of different ao’s on the same atom Hybrid orbital confined to that atom
Molecular Orbital theory (LCAO) Weighted average of different ao’s on all atoms of molecule Resulting mo involves all atoms of molecule
Formation of molecular orbitals Bonding orbital More electron density between nuclei More electrostatic attraction Bonding MO at lower energy
Antibonding orbital No density between atoms Lower electrostatic attraction Antibonding MO at higher energy
Bond order BO 1 { bonding elecs - antibonding elecs} 2 Bond order 1 = single bond (1/2 x 2) Bond order 2 = double bond (1/2 x 4) Bond order 3 = triple bond (1/2 x 6)
Summary of important concepts in MO MO’s are formed by linear combination of AO’s Two AO’s combine to give two MO’s: one is higher in energy, one is lower Orbital filling follows aufbau principle: lowest energy orbitals first Maximum occupancy of MO is two (spin-paired) Hund’s rule: degenerate orbitals are singly occupied before pairing Bond order is one half times (number of electrons in bonding MO’s minus number of electrons in anti-bonding MO’s)
On the existence of molecules: MO energy level diagrams H2 (2 electrons) in bonding σ MO; antibonding σ* MO is vacant. Total number of bonds = (+1 – 0) = 1 Configuration (σ1s)2
He2 (4 electrons): two in bonding σ, two in antibonding σ* Total number of bonds = (+ 1 – 1) = 0 Configuration (σ1s)2(σ*1s)2
Second row elements Li2 contains 6 electrons Bonding σ orbitals between 1s and 2s Antibonding σ* orbitals between 1s and 2s Occupied: σ1s,σ2s, and σ*1s Bond order = 2 – 1 = 1 Does Be2 exist?
Formation of π orbitals in MO Defining the internuclear axis as z Overlap of the pz orbitals produces σ bond Overlap of px and py orbitals produces π bonds
General energy level diagram for second-row homonuclear diatomics Assumes no interaction between the 2s and 2p orbitals 2s orbitals lower in energy than 2p orbitals σ2s and σ*2s orbitals lower than σ2p orbital
Overlap of the 2pz is greater than that of the 2px or 2py so σ2p is lower than the π2p orbital The π2p and π*2p are degenerate (2 orbitals with the same energy)
Consequences of interaction between 2s and 2p
The 2s and 2p orbitals do interact σ2s and σ2p orbitals move further apart in energy Strength of interaction changes with atomic number Case A NO interaction: σ2p < π2p Case B STRONG interaction: σ2p > π2p
Second row diatomics: interaction decreases across period B2, C2, and N2 are case B (strong interaction) O2, F2 and Ne2 are case A (weak interaction) Bond order from MO theory matches bond order from Lewis dot diagrams perfectly
Magnetism and electrons Paramagnetism: attracted by a magnetic field Diamagnetism: repelled by a magnetic field Paramagnetic effect is much greater than diamagnetic effect
Electrons have magnetic moments Diamagnetic substances have no unpaired electrons Paramagnetic substances have unpaired electrons
Magnetism of O2 and the limitations of Lewis
O2 is paramagnetic (YouTube) O2 must contain unpaired electrons Lewis dot diagram shows simple lone pairs Lewis predicts diamagnetism Another shortcoming of Lewis dot structures Lewis dot structure
O
O
MO theory to the rescue *
MO theory gives two degenerate π and π orbitals Hund’s rule states that these are singly occupied O2 is paramagnetic If the σ* was below the π* what is the situation?
Correlate magnetic properties with MO diagram
Heteronuclear molecules and NO NO contains 11 electrons implies high 0 0 reactivity -1 +1
N
O N
O
Lewis structure favours unpaired electron on N Experimental bond order appears greater than 2
MO description of NO AOs of more electronegative atom lower in energy (O more electronegative than N) Bonding orbitals have more of more electronegative atom character (O) Antibonding orbitals have more of less electronegative atom character (N)
MO diagram shows bond order 2.5 consistent with experiment Unpaired electron in π* orbital is more N-like (consistent with Lewis dot structure)