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Atomic and Molecular Orbitals

Atomic and Molecular Orbitals. Structure and Properties of Organic Molecules. Electrons act as both particles and waves (duality) An orbital can be described as a 3D standing wave To understand this, we can look at wave properties of guitar strings. Wave Harmonics with Dr. Jimi.

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Atomic and Molecular Orbitals

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  1. Atomic and Molecular Orbitals

  2. Structure and Properties of Organic Molecules • Electrons act as both particles and waves (duality) • An orbital can be described as a 3D standing wave • To understand this, we can look at wave properties of guitar strings

  3. Wave Harmonics with Dr. Jimi http://www.rollingstone.com/music/lists/100-greatest-guitarists-of-all-time-19691231/jimi-hendrix-19691231; Wade, Organic Chemistry, 2013

  4. Wave Harmonics with Dr. Jimi http://www.rollingstone.com/music/lists/100-greatest-guitarists-of-all-time-19691231/jimi-hendrix-19691231; Wade, Organic Chemistry, 2013

  5. Wave properties of electron in s orbitals Wade, Organic Chemistry, 2013

  6. Wave properties of electron in p orbitals Wade, Organic Chemistry, 2013

  7. Linear Combination of Atomic Orbitals • Atomic orbitals combine and overlap to produce more complex standing waves • Any wave overlap can be constructive or destructive. Wade, Organic Chemistry, 2013; http://animals.howstuffworks.com/insects/butterfly-colors1.htm

  8. Linear Combination of Atomic Orbitals • When orbitals on different atoms interact (think of covalent bonds), they produce molecular orbitals that lead to bonding or antibonding interactions. • The stability of covalent bond results from increased electron density in the bonding region (i.e., the space between the 2 nuclei where orbitals overlap). • There is an optimal distance between these nuclei where charge attraction and repulsion are balanced. • This leads to a somewhat consistently fixed bond length. Electrostatic potential map of H2 Wade, Organic Chemistry, 2013

  9. Sigma Bonding in Hydrogen Molecules • Formation of cylindrically symmetrical sigma (σ) bond as a result of constructive addition • e- density in bonding region increases • Forms bonding MO Wade, Organic Chemistry, 2013

  10. Sigma Bonding in Hydrogen Molecules • The 2 out of phase 1s hydrogen orbitals overlap destructively to cancel out part of the wave, producing a node • Forms antibonding MO with (σ*) bond Wade, Organic Chemistry, 2013

  11. Describing Chemical Bonds: Molecular Orbital Theory • A molecular orbital (MO): where electrons are most likely to be found (specific energy and general shape) in a molecule • Additive combination (bonding) MO is lower in energy • Subtractive combination (antibonding) MO is higher energy

  12. Molecular Orbitals in Ethylene • The  bonding MO is from combining p orbital lobes with the same algebraic sign • The  antibonding MO is from combining lobes with opposite signs • Only bonding MO is occupied

  13. Formation of MO’s Destructive interference of orbitals Constructive interference of orbitals Wade, Organic Chemistry, 2013

  14. Sigma Bonding in Hydrogen Molecules Formation of the bonding MO requires less E to maintain, i.e., it is more stable Wade, Organic Chemistry, 2013

  15. Sigma Bonding between p and s orbitals Wade, Organic Chemistry, 2013

  16. Pi (π) Bonds • Pi (π) bonds formed by overlap of 2 parallel p orbitals. • Not cylindrically symmetrical like σ bond • Not as strong as σ bond • Important for chemical reactions • Pi bonds form double and triple bonds. Wade, Organic Chemistry, 2013

  17. Orbital Hybridization • When orbitals on the same atom interact, they produce hybrid atomic orbitals that define bond geometry. • Consider the problem with carbon in the energy diagram below: 1s22s22p2 2p Valence shell electrons 2 Energy 2s 1 C 1s How does carbon form 4 bonds? Wade, Organic Chemistry, 2013

  18. Orbital Hybridization • Carbon has 4 valence electrons (2s2 2p2) • In CH4, all C–H bonds are identical (tetrahedral) • sp3 hybrid orbitals:s orbital and three p orbitals combine to form four equivalent, unsymmetrical, tetrahedral orbitals (sppp = sp3), Pauling (1931) Energy

  19. Orbital Hybridization

  20. Orbital Hybridization • An sp3orbital has a two-lobed shape, similar to the shape of a p orbital but with different-sized lobes. • Each carbon-hydrogen bond in methane arises from an overlap of a C (sp3) and an H (1s) orbital. • The sharing of two electrons in this overlap region creates a sigma (σ) bond. Seager SL, Slabaugh MR, Chemistry for Today: General, Organic and Biochemistry, 7th Edition, 2011

  21. sp3Orbitals and the Structure of Ethane • Two C’s bond to each other by s overlap of an sp3 orbital from each • Three sp3 orbitals on each C overlap with H 1s orbitals to form six C–H bonds • C–H bond strength in ethane 421 kJ/mol • C–C bond is 154 pm long and strength is 377 kJ/mol • All bond angles of ethane are tetrahedral

  22. VSEPR Theory • The tetrahedral geometry can be explained by Valence Shell Electron Pair Repulsion (VSEPR) Theory • (VSEPR) Theory: Electron pairs repulse each other in such a way so that they are as far apart from each other as possible. • The bond angles observed in organic compounds can (currently) only be explained by this repulsion of hybridized orbitals. tetrahedral linear  geometry trigonal Wade, Organic Chemistry, 2013

  23. The Geometry of Alkenes • In C=C bonds, sp2 hybrid orbitals are formed by the carbon atoms, with one electron left in a 2p orbital. A representation of sp2 hybridization of carbon. • During hybridization, two of the 2p orbitals mix with the single 2s orbital to produce three sp2 hybrid orbitals. One 2p orbital is not hybridized and remains unchanged.

  24. The Geometry of Alkenes • One bond (sigma, σ) is formed by overlap of two sp2 hybrids. • The second bond (pi, π) is formed by connecting the unhybridized p orbitals. Seager SL, Slabaugh MR, Chemistry for Today: General, Organic and Biochemistry, 7th Edition, 2011

  25. The Geometry of Alkenes • The planar geometry of the sp2 hybrid orbitals and the ability of the 2p electron to form a “pi bond” bridge locks the C=C bond firmly in place. Seager SL, Slabaugh MR, Chemistry for Today: General, Organic and Biochemistry, 7th Edition, 2011

  26. The Geometry of Alkenes • Because there is no free rotation about the C=C bond, geometric isomerism is possible. • cis- isomers have two similar or identical groups on the same side of the double bond. • trans- isomers have two similar or identical groups on opposite sides of the double bond. Seager SL, Slabaugh MR, Chemistry for Today: General, Organic and Biochemistry, 7th Edition, 2011

  27. The Geometry of Alkenes • Geometric isomers have different physical properties. Seager SL, Slabaugh MR, Chemistry for Today: General, Organic and Biochemistry, 7th Edition, 2011

  28. Hybridization/Geometry of Alkynes • Insoluble in water • Less dense than water • Low MP, BP Seager SL, Slabaugh MR, Chemistry for Today: General, Organic and Biochemistry, 7th Edition, 2011

  29. Summary of Hybridizations

  30. Adjacent Double Bonds

  31. Comparison of C–C and C–H Bonds

  32. Hybridization of Nitrogen and Oxygen • Elements other than C can have hybridized orbitals • H–N–H bond angle in ammonia (NH3) 107.3° • C-N-H bond angle is 110.3 ° • N’s orbitals (sppp) hybridize to form four sp3 orbitals • One sp3 orbital is occupied by two nonbonding electrons, and three sp3 orbitals have one electron each, forming bonds to H and CH3.

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