Chemical Bonding Lewis Theory-VSEPR Valence Bond Theory Molecular Orbital Theory
Problems with Valence Bond Theory VB theory predicts properties better than Lewis theory bonding schemes, bond strengths, lengths, rigidity There are still properties it doesn t predict perfectly magnetic behavior of certain molecules strength of bonds VB theory presumes the electrons are localized in s doesn t account for delocalization
Molecular Orbital Theory In MO theory, we apply Schrödinger s wave equation to the molecule to calculate a set of molecular s. The equation solution is estimated. We start with good guesses as to what the s should look like, then test the estimate until the energy is minimized The electrons belong to the whole molecule s are delocalized
LCAO The simple guess starts with s of the atoms adding together to make molecular s, the Linear Combination of Atomic Orbitals. The waves can combine either constructively or destructively.
Molecular Orbitals When wave functions combine constructively, the resulting molecular has less energy than the original s it is called a Bonding Molecular Orbital σ, π most of the electron density between the nuclei Amplitudes of wave functions added
Molecular Orbitals When wave functions combine destructively, the resulting molecular has more energy than the original s it is called an Antibonding Molecular Orbital σ*, π* most of the electron density outside the nuclei nodes between nuclei Amplitudes of wave functions subtracted.
Interaction of 1s Orbitals destructive interference - + σ1s* antibonding molecular Increasing energy 1s 1s constructive interference σ1s bonding molecular
Molecular Orbital Theory Electrons in bonding MOs are stabilizing. lower energy than the s Electrons in antibonding MOs are destabilizing. higher in energy than s electron density located outside the internuclear axis electrons in antibonding s cancel stability gained by electrons in bonding s
Molecular Orbitals and Properties Bond Order = difference between number of electrons in bonding and antibonding s only need to consider valence electrons may be a fraction higher bond order = stronger and shorter bonds If bond order = 0, then bond is unstable compared to individual atoms and no bond will form. A substance will be paramagnetic if there are unpaired electrons in molecular s.
A Molecular Orbital Diagram - H2 antibonding MO σ* H H 1s σ bonding MO 1s
A Molecular Orbital Diagram - H2 LUMO lowest unoccupied molecular σ* H H 1s 1s σ HOMO highest occupied molecular
A Molecular Orbital Diagram - H2 σ* H H 1s σ 1s BO = ½( be abe) Because more electrons are in bonding s than are in antibonding s, there is a net bonding interaction. = 1
X A Molecular Orbital Diagram - He2 σ* He: 1s 1s He: σ BO = ½( be abe) Because as many electrons are in bonding s as in antibonding s, NO NET BONDING INTERACTION. = 0
A Molecular Orbital Diagram - Li2 σ* Li σ Li σ* 1s σ 1s
A Molecular Orbital Diagram - Li2 σ* Li Li σ = 1 Because more electrons are in bonding s than are in antibonding s, there is a net bonding interaction.
MO Diagram - Be2 σ* Be: σ :Be Because as many electrons are in bonding s as in antibonding s, NO NET BONDING INTERACTION. = 0
Interaction of p Atomic Orbitals Each time we combine two s, we obtain two molecular s: one bonding and one antibonding. In the case of p s, overlap occurs either "head-to-head" to form σ and σ* molecular s or sideways" to form π and π* molecular s.
Atomic s + 2p z 2p z σ 2p * Molecular s antibonding σ 2p bonding
Atomic s Molecular s antibonding bonding + + 2p x 2p x π 2p * π 2p + + 2p y 2p y π 2p * π 2p
Molecular Orbitals B2, C2, N2, O2, F2, Ne2,
MO Diagram - B2 σ2p* π2p* 2p σ2p 2p π2p BO = ½(6 be 4 abe) BO = 1 σ* B B σ
MO Diagram - C2 σ2p* π2p* 2p σ2p 2p π2p BO = ½(8 be 4 abe) BO = 2 σ* C C σ
MO Diagram - N2 σ2p* π2p* 2p σ2p 2p π2p BO = ½(10 be 4 abe) BO = 3 σ* N N σ Dinitrogren ( N2 ) is Diamagnetic!!
MO Diagram - C2 2- (carbide ion) σ2p* π2p* 2p σ2p 2p π2p BO = ½(10 be 4 abe) BO = 3 σ* C C σ
Molecular Orbitals B2, C2, N2 σ2p* π2p* 2p σ2p 2p π2p Molecular Orbitals O2, F2, Ne2 σ2p* π2p* 2p π2p 2p σ2p
MO Diagram - O2 σ2p* π2p* 2p π2p 2p σ2p BO = ½(10 be 6 abe) BO = 2 σ* O O σ
Dioxygen ( O2 ) is Paramagnetic!!
MO Diagram - F2 σ2p* π2p* 2p π2p 2p σ2p BO = ½(10 be 8 abe) BO = 1 σ* F F σ
X MO Diagram - Ne2 σ2p* π2p* 2p π2p 2p σ2p BO = ½(10 be 10 abe) BO = 0 σ* Ne Ne σ
Using MO Theory to Explain Bond Properties As the following data show, removing an electron from N2 forms an ion with a weaker, longer bond than in the parent molecules, whereas the ion formed from O2 has a stronger, shorter bond: These facts can be explained by examining diagrams that show the sequence and occupancy of MOs.
N2 N2 + O2 O2 + σ2p* σ2p* π2p* π2p* σ2p π2p π2p σ2p σ* σ* σ* σ* σ σ σ σ B.O.=3 B.O.=2.5 B.O.=2 B.O.=2.5 Electron is lost from bonding molecular. Electron is lost from antibonding molecular.