Chemical Bonding. Lewis Theory-VSEPR Valence Bond Theory Molecular Orbital Theory

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1 Chemical Bonding Lewis Theory-VSEPR Valence Bond Theory Molecular Orbital Theory

2 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

3 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

4 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.

5 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

6 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.

7 Interaction of 1s Orbitals destructive interference - + σ1s* antibonding molecular Increasing energy 1s 1s constructive interference σ1s bonding molecular

8 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

9 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.

10 A Molecular Orbital Diagram - H2 antibonding MO σ* H H 1s σ bonding MO 1s

11 A Molecular Orbital Diagram - H2 LUMO lowest unoccupied molecular σ* H H 1s 1s σ HOMO highest occupied molecular

12 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

13 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

14 A Molecular Orbital Diagram - Li2 σ* Li σ Li σ* 1s σ 1s

15 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.

16 MO Diagram - Be2 σ* Be: σ :Be Because as many electrons are in bonding s as in antibonding s, NO NET BONDING INTERACTION. = 0

17 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.

18 Atomic s + 2p z 2p z σ 2p * Molecular s antibonding σ 2p bonding

19 Atomic s Molecular s antibonding bonding + + 2p x 2p x π 2p * π 2p + + 2p y 2p y π 2p * π 2p

20 Molecular Orbitals B2, C2, N2, O2, F2, Ne2,

21 MO Diagram - B2 σ2p* π2p* 2p σ2p 2p π2p BO = ½(6 be 4 abe) BO = 1 σ* B B σ

22 MO Diagram - C2 σ2p* π2p* 2p σ2p 2p π2p BO = ½(8 be 4 abe) BO = 2 σ* C C σ

23 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!!

24 MO Diagram - C2 2- (carbide ion) σ2p* π2p* 2p σ2p 2p π2p BO = ½(10 be 4 abe) BO = 3 σ* C C σ

25 Molecular Orbitals B2, C2, N2 σ2p* π2p* 2p σ2p 2p π2p Molecular Orbitals O2, F2, Ne2 σ2p* π2p* 2p π2p 2p σ2p

26 MO Diagram - O2 σ2p* π2p* 2p π2p 2p σ2p BO = ½(10 be 6 abe) BO = 2 σ* O O σ

27 Dioxygen ( O2 ) is Paramagnetic!!

28 MO Diagram - F2 σ2p* π2p* 2p π2p 2p σ2p BO = ½(10 be 8 abe) BO = 1 σ* F F σ

29 X MO Diagram - Ne2 σ2p* π2p* 2p π2p 2p σ2p BO = ½(10 be 10 abe) BO = 0 σ* Ne Ne σ

30 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.

31 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.

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