Presentation on theme: "Molecular Orbital Theory"— Presentation transcript:
1Molecular Orbital Theory The goal of molecular orbital theory is to describe molecules in a similar way to how we describe atoms, that is, in terms of orbitals, orbital diagrams, and electron configurations.
2Atomic Orbitals Orbital Mixing When atoms share electrons to form a bond, their atomic orbitals mix to form molecular bonds. In order for these orbitals to mix they must:Have similar energy levels.Overlap well.Be close together.This is an example of orbital mixing. The two atoms share one electron each from there outer shell. In this case both 1s orbitals overlap and share their valence electrons.
7Atomic and Molecular Orbitals In molecule, electrons occupy molecular orbitals which surround the molecule.The two 1s atomic orbitals combine to form two molecular orbitals, one bonding (s) and one antibonding (s*).This is an illustration of molecular orbital diagram of H2.Notice that one electron from each atom is being “shared” to form a covalent bond. This is an example of orbital mixing.
8Molecular Orbital Theory Each line in the diagram represents an orbital.The molecular orbital volume encompasses the whole molecule.The electrons fill the molecular orbitals of molecules like electrons fill atomic orbitals in atoms.
9Molecular Orbital and Valence Bond Theory Electrons go into the lowest energy orbital available.The maximum number of electrons in each molecular orbital is two.One electron goes into orbitals of equal energy, with parallel spin, before they begin to pair up.Aufbau PrinciplePauli exclusion principleHund's Rule
11Bond Order electrons electrons Bond Order = 2 # bonding - # antibondingelectrons electronsBond Order =2The bond order is an indication of bond strengthGreater bond order = Greater bond strengthThe bond order equals the number of bonds formed in the molecule
12Predicting Species Stability Using MO Diagrams PROBLEM:Use MO diagrams to predict whether H2+ and H2- exist. Determine their bond orders and electron configurations.bond order = 1/2(1-0) = 1/2bond order = 1/2(2-1) = 1/2ssssH2- does existH2+ does exist1sAO of H1sAO of HMO of H2-MO of H2+
13The paramagnetic properties of O2 Figure 11.21The paramagnetic properties of O2
15MO vs. VBAtoms or molecules in which the electrons are paired are diamagnetic.Those that have one or more unpaired electrons are paramagnetic, attracted to a magnetic field.Liquid oxygen is attracted to a magnetic field and can actually bridge the gap between the poles of a horseshoe magnet.The molecular orbital model of O2 is therefore superior to the valence-bond model, which cannot explain this property of oxygen.
19MO diagram for He2+ and He2 Figure 11.16MO diagram for He2+ and He2Energys*1ss1ss*1sAO of He1sAO of He+1sAO of He1sAO of He1sEnergys1sMO of He+MO of He2He2 bond order = 0He2+ bond order = 1/2He2 does not exist!
20Relative MO energy levels for Period 2 diatomic molecules. MO energy levels for B2, C2, and N2MO energy levels for O2, F2, and Ne2
21SAMPLE PROBLEM 11.4Using MO Theory to Explain Bond Propertiesantibonding e- lostcontinuedN2N2+O2O2+2ps2ss2ss2pp2p2p2pbonding e- lost2p2p2ps2ss2sbond orders1/2(8-2)=31/2(7-2)=2.51/2(8-4)=21/2(8-3)=2.5
22Bonding in s-block homonuclear diatomic molecules. s*2ss2sFigure 11.17s*2ss2s2s2s2sBonding in s-block homonuclear diatomic molecules.Be2Li2Energy1ss*1ss1s1ss*1ss1sLi2 bond order = 1Be2 bond order = 0
23Conclusions Bonding electrons are localized between atoms. Atomic orbitals overlap to form bonds.Two electrons of opposite spin can occupy the overlapping orbitals.Bonding increases the probability of finding electrons in between atoms.It is also possible for atoms to form ionic and metallic bonds.