1. John E. McMurry www.cengage.com/chemistry/mcmurry Paul D. Adams University of Arkansas 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. http://www.rollingstone.com/music/lists/100-greatest-guitarists-of-all-time-19691231/jimi-hendrix-19691231http://www.rollingstone.com/music/lists/100-greatest-guitarists-of-all-time-19691231/jimi-hendrix-19691231; Wade, Organic Chemistry, 2013 Wave Harmonics with Dr. Jimi

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

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

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

7.  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 Linear Combination of Atomic Orbitals

8.  When orbitals on different atoms interact (think of covalent bonds), they produce molecular orbitals that lead to bonding or antibonding interactions. Wade, Organic Chemistry, 2013 Electrostatic potential map of H 2 There is an optimal distance between these nuclei where charge attraction and repulsion are balanced. This leads to a somewhat consistently fixed bond length. 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). Linear Combination of Atomic Orbitals

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

10. Wade, Organic Chemistry, 2013 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 Sigma Bonding in Hydrogen Molecules

11.  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 Describing Chemical Bonds: Molecular Orbital Theory

12.  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 Molecular Orbitals in Ethylene

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

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

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

16.  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 Pi (π) Bonds

17.  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: Wade, Organic Chemistry, 2013 2p 2s 1s Energy Valence shell electron s C 1 2 How does carbon form 4 bonds? 1s22s22p21s22s22p2 Orbital Hybridization

18.  Carbon has 4 valence electrons (2s 2 2p 2 )  In CH 4, all C–H bonds are identical (tetrahedral)  sp 3 hybrid orbitals: s orbital and three p orbitals combine to form four equivalent, unsymmetrical, tetrahedral orbitals (sppp = sp 3 ), Pauling (1931) Orbital HybridizationEnergy

20.  An sp 3 orbital 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 (sp 3 ) 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, 7 th Edition, 2011 Orbital Hybridization

21.  Two C’s bond to each other by  overlap of an sp 3 orbital from each  Three sp 3 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 sp 3 Orbitals and the Structure of Ethane

22.  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. Wade, Organic Chemistry, 2013 trigonal tetrahedrallinear  geometry VSEPR Theory

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

24. 1. One bond (sigma, σ) is formed by overlap of two sp 2 hybrids. 2. The second bond (pi, π) is formed by connecting the unhybridized p orbitals. Seager SL, Slabaugh MR, Chemistry for Today: General, Organic and Biochemistry, 7 th Edition, 2011 The Geometry of Alkenes

25.  The planar geometry of the sp 2 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, 7 th Edition, 2011 The Geometry of Alkenes

26.  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, 7 th Edition, 2011 The Geometry of Alkenes

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

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

29. Summary of Hybridizations

30. Adjacent Double Bonds

31. Comparison of C–C and C–H Bonds

32.  Elements other than C can have hybridized orbitals  H–N–H bond angle in ammonia (NH 3 ) 107.3°  C-N-H bond angle is 110.3 °  N’s orbitals (sppp) hybridize to form four sp 3 orbitals  One sp 3 orbital is occupied by two nonbonding electrons, and three sp 3 orbitals have one electron each, forming bonds to H and CH 3. Hybridization of Nitrogen and Oxygen