DE Chemistry Dr Walker Hybridization and the Localized Electron Model Hybridization The mixing of two or more atomic orbitals of similar energies on the same atom to produce new orbitals of equal ID: 1010547
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1. Covalent Bonding: OrbitalsDE ChemistryDr. Walker
2. Hybridization and the Localized Electron ModelHybridizationThe mixing of two or more atomic orbitals of similar energies on the same atom to produce new orbitals of equal energiesMixes s, p, and/or d orbitals when covalently bondedHybrid OrbitalsOrbitals of equal energy produced by the combination of two or more orbitals on the same atom
3. Hybridization - The Blending of Orbitals Poodle++Cocker Spaniel====++s orbitalp orbitalCockapoosp orbital
4. Why Hybridization?Think of the electrons involved in bonding….Typically, most bonding occurs between electrons in s and p orbitalsThese orbitals are all valence electrons, so all are involved in bonding
5. Why Hybridization?Think about carbon….We know carbon needs to make four covalent bonds to complete its octet and become stableTwo electrons are at one energy level and two electrons are at another energy level
6. Things we know….In methane, the electrons in hydrogen’s 1s orbital bond with the orbitals in the second energy level of carbon
7. Houston, we have a problem….We know that all of the C-H bonds in methane are identicalIf two bonds are made with electrons in 2s and two with electrons in 2p, these bonds SHOULD be at different energies….How can we explain this?
8. The RationalizationHow can we explain this?Hybridization is the combining of two or more orbitals of nearly equal energy within the same atom into orbitals of equal energy.
9. In the case of methane, they call the hybridization sp3, meaning that an s orbital is combined with threep orbitals to create four equal hybrid orbitals.These new orbitals have slightly MORE energy thanthe 2s orbital…… and slightly LESS energy than the 2p orbitals.1s2sp32sp32sp32sp3
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11. Notice the tetrahedral arrangement of sp3 orbitals on theright correspond with the molecular geometry of methane and other AX4 type molecules, with bond angles of 109o (assuming no unbonded electron pairs).
12. sp2 HybridizationAnother hybrid is the sp2, which combines two orbitals from a p sublevel with one orbital from an s sublevel.One of the p orbitals is not used in hybridizationEnergy diagram of sp2 orbitals
13. sp2 HybridizationThis hybridization results in bond angles of 120o, consistent with trigonal planar geometries
14. sp2 Hybridization3 effective pairs of electrons surround the carbon (double bond treated as one effective pair)In boron compounds, the unused p orbital is empty
15. s and p bondss bondsOccupies space directly between atomsp bondsOccupies space above and below atoms from the overlap between two p orbitalsThe picture on the left represents a double bond, which uses one s and one p bond
16. Notice the sp2 geometry on the left. The single bonds in ethene are s bonds. Notice the unhybridized p orbital onthe left, which forms the p bond shown in the previous slide.
17. sp HybridizationAs you can probably guess this is a hybrid of one s and one p orbital, with 2 p orbitals remaining unhybridized.In the previous example, the unhybridized p orbitals made the double bond. In this case, two unhybridized p orbitals results in a triple bond between carbons (one s bond, two p bonds.
18. sp HybridizationThe three different p orbitals (px, py, pz) are oriented on each axis of space. If only one p orbital hybridizes, it results in a 180o angle. This is consistent with linear molecules.
19. sp Hybridizationsp Hybridization of carbon in ethyne.
20. sp Hybridizationsp Hybridization of carbon in carbon dioxide.
21. dsp3 HybridizationForms for central atoms requiring a trigonal bipyramidal arrangementExamples PCl5, other pentavalent atoms that disobey the octet ruleAllows bonding of extra electrons that require more than the s and p orbitals (8 electrons) allow
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23. d2sp3 HybridizationUsed for hexavalent atoms with 12 bonded electrons, requiring an octahedral arrangement
24. Hybridization and Molecular GeometryFormsOverall StructureHybridization of “A”AX2LinearspAX3, AX2ETrigonal Planarsp2AX4, AX3E, AX2E2Tetrahedralsp3AX5, AX4E, AX3E2, AX2E3Trigonal bipyramidaldsp3AX6, AX5E, AX4E2Octahedrald2sp3A = central atomX = atoms bonded to AE = nonbonding electron pairs on A
25. Shortcomings of the Localized Electron ModelAssumes electrons are localized (thus the name…)Doesn’t really explain resonanceDoesn’t work well for unpaired electronsDoesn’t give any indication of bond energy
26. Molecular Orbital (MO) TheoryMolecular OrbitalsSimilar to atomic orbitals, but for moleculesCan hold two electrons with opposite spinsSquare of the orbital's wave function indicates electron probability
27. HydrogenTwo possible bonding orbitals, shapes determine by Y2Bonding takes place in MO1 in which electrons achieve lower energy (greater stability), with electrons between the two nuclei
28. HydrogenBoth orbitals are in line with the nuclei, so they are s (same as in localized model) molecular orbitalsHigher energy orbital is designated as antibonding (*)Electron configuration of H2 can be written as s1s2
29. Bond OrderBond order is the difference between the number of bonding electrons and the number of antibonding electrons, divided by twoLarger bond order =greater bond strengthgreater bond energyshorter bond lengthTypically, molecules with a bond order = 0 don’t actually exist
30. Bonding in Homonuclear Diatomic MoleculesIn order to participate in molecular orbitals, atomic orbitals must overlap in spaceLarger bond order is favoredWhen molecular orbitals are formed from p orbitals, s orbitals are favored over p orbitals (s interactions are stronger than p interactions)Electrons are closer to the nucleus = lower energy
31. Note Regarding MO TheoryIn the localized electron model, hybrid orbitals are formed by mixing s, p, and d orbitals for bondingThere is no hybridization in MO theory. Electrons in s orbitals bond with each other (forming s bonds) and electrons in p orbitals bond with each other (forming s and p bonds)
32. B2 MoleculeNotice that s and p orbitals do not mix when bonding to form molecular orbitalsThere are 4 total bonding electrons and 2 antibonding electronsBond order = (4-2)/2 = 1, making this a stable moleculeThe figure shown on the right is the expected arrangement, but fails to account for certain physical properties of B2.
33. ParamagnetismMagnetism can be induced in some nonmagnetic materials when in the presence of a magnetic fieldParamagnetism causes the substance to be attracted into the inducing magnetic fieldAssociated with unpaired electronsDiamagnetism causes the substance to be repelled from the inducing magnetic fieldAssociated with paired electrons
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35. Observations Regarding ParamagnetismBond order does not necessarily equal number of bonds, as B2 has a triple bond, yet its bond order is 1. If you’re paying attention, the s and p molecular orbital energies are reversed from the expected arrangement shown previously for B2. This alteration to the model explains B2’s paramagnetic behavior. In the expected arrangement shown earlier, B2 should have been diamagnetic….but it isn’t.
36. Bonding in Heteronuclear Diatomic MoleculesSimilar, but not identical atomsUse molecular orbital diagrams for homonuclear moleculesSignificantly different atomsEach molecule must be examined individuallyThere is no universally accepted molecular orbital energy order
37. Bonding in Heteronuclear Diatomic MoleculesExample: NONitrogen has 5 valence electrons, oxygen has 6, giving 11 totalBond order = (8-3)/2 = 2.5With an unpaired electron we expect this to be paramagnetic (and it is)
38. Bonding in Heteronuclear Diatomic MoleculesExample: CNCarbon has 4 valence electrons, oxygen has 6, giving 10 totalBond order = (8-2)/2 = 3With no unpaired electrons we expect this to be diamagnetic (and it is)
39. Combining ModelsResonanceAttempt to draw localized electrons in a structure in which electrons are not localizeds bonds can be described using localized electron modelp bonds (delocalized) must be described using the molecular orbital model
40. Benzenes bonds (C - H and C - C) are sp2 hybridizedUses localized modelp bonds are a result of remaining p orbitals above and below the plane of the benzene ringUses delocalized model (MO)
41. Representations of BenzeneThe s bonding system with carbon’s sp2 orbital system demonstrates the localized bonding model
42. Representations of BenzeneThe non-hybridized p orbitals used to make p bonds above and below the ring show the need for a delocalized model where resonance is necessary