1. Structure and Bonding A Review of Needed Material

2 Organic Chemistry “Organic” – until mid 1800’s referred to compounds from living sources (mineral sources were “inorganic”) Wöhler in 1828 showed that urea, an organic compound, could be made from a inorganic materials Organic compounds are those based on covalently bonded carbon and study of their structures and reactions.

3 Atomic Structure Structure of an atom Nucleus = assigned a positive charged, very dense, protons and neutrons in it and small ( m), deflects an alpha particle shot at it. Electrons = therefore are negatively charged, occupy allowed energies (sometimes called an electron cloud) ( m across) around nucleus … and are not “orbiting” the nucleus … Ahhhhh see page 4! Diameter is about 2  m (200 picometers (pm)) [the unit angstrom (Å) is m = 100 pm]

4 Isotopes Isotopes are atoms of the same element that have different numbers of neutrons and therefore different mass numbers Nuclear “spin” occurs when an odd number of nuclear particles are present. The spin has a specific energy when placed in a magnetic field. This energy is characteristic of the isotope and its environment … and can be detected, characterized, and used to determine which atoms are bonded to each other in a pure substance! NMR - Yaaahooo

5 Atomic Structure: Orbitals Quantum mechanics: describes electron energies by a wave equation A Wave function, , is a solution of wave equation A plot of  2 describes probable electron density The Shapes that come from all this math are called orbitals (  2 ). We pick the solutions that work and through the rest out. We can mathematically combine these solutions to make other orbitals at our whim – again, we use the ones that work and leave the rest out: HYBRID ORBITALS!

6 It’s the Filled Orbitals – Stupid! Orbitals have an energy order to them … I hate the word “SHELL” Each primary quantum number denotes increasing energy of the electrons within them and contain different numbers and kinds of orbitals Each orbital can be occupied by two electrons n =1 contains one s orbital, denoted 1s, holds only two electrons n = 2 contains one s orbital (2s) and three p orbitals (2p), eight electrons n = 3 contains an s orbital (3s), three p orbitals (3p), and five d orbitals (3d), 18 electrons Eliminate your idea of “filled shells” because n = 3 shell is very happy not being filled and only having 8 electrons. It is the filled orbitals that cause stability. Filled s and p orbitals are the ones we are most interested in with organic chemistry.

7 Shapes of Atomic Orbitals Four different kinds of orbitals are useful to discuss s, p, d, and f s and p orbitals most important in organic chemistry

8 Chemical Bonding Theory Observation: Carbon makes four bonds and has yet to be observed to ever have 5 bonds that are isolable. Observation: 2-chlorobutane has two isomers Explanation: Carbons with 4 bonds = tetrahedral Note that a wedge indicates a bond is coming forward Note that a dashed line indicates a bond is behind the page

9 Chemical Bonds Atoms form bonds because the compound that results is more stable than the separate atoms. Bond Energy = Bond Length

10 Valence Bond Theory Valence electrons are what make bonds. … Lewis structures show valence electrons of an atom as dots First bond cylindrically symmetrical, sigma (  ) bond Second and third bonds … called pi (  ) bonds …

11 Number of Covalent Bonds Atoms with one, two, or three valence electrons form one, two, or three bonds and the empty p orbital does not bond directly but is available to accept electron density – Lewis Acid!

12 Number of Covalent Bonds Atoms with four or more valence electrons form as many bonds as they need electrons to fill the s and p orbitals = octet rule

13 Valence Electrons of Boron Boron has three valence electrons (2s 2 2p 1 ), forming three bonds (BF 3 )

14 Valence Electrons of Carbon Carbon has four valence electrons (2s 2 2p 2 ), forming four bonds (CH 4 )

15 Valence of Nitrogen Nitrogen has five valence electrons (2s 2 2p 3 ) but forms only three bonds (NH 3 )

16 Valence Electrons of Oxygen Oxygen has six valence electrons (2s 2 2p 4 ) but forms two bonds (H 2 O)

17 Hybrid orbitals – Reorganization of orbital energies! In CH 4, all C–H bonds are identical (tetrahedral) Each C–H bond has a strength of 438 kJ/mol and bond length of 110 pm Bond angle: each H–C–H is 109.5°, the tetrahedral angle. Explanation: sp 3 hybrid orbitals: s orbital and three p orbitals combine to form four equivalent orbitals (sppp = sp 3 ), Pauling (1931) arranged in a tetrahedral array.

18 Pictures – hybrid orbitals

19 Bonding Ethane

20 Bonding Ethene

21 Bonding in Ethyne

22 What about lone pairs?

23 Molecular Orbital Theory Helps us render the idea that the electrons are lowering in energy when bonds form. Best in 2 electron scenerios! Bonding Anti-bonding But really … I have never had to use these ideas to remember or derive anything in all 10+ organic chemistry classes I have taken or with any of the organic synthesis I have accomplished. 2 patents 32 anticancer compounds synthesized 4 new synthetic pathways 7 novel reactions yielding >93% yield from others work

24 What works! Thermodynamics, Bond Energies - Will it even work Equilibrium – Can you get enough of it Kinetics Is the energy barrier to high? Which of the “X” possible reactions will win? Polarity in molecules Who likes whom? Consonant vs. Dissonant synthesis

25 Summary Organic chemistry – chemistry of carbon compounds Atom: positively charged nucleus surrounded by negatively charged electrons Electronic structure of an atom described by wave equation Electrons occupy orbitals around the nucleus with specific energies. Different orbitals have different allowed energies and different shapes s orbitals are spherical, p orbitals are dumbbell-shaped Covalent bonds - electron pair shared between two atoms are lower in energy Valence bond theory – Valence electrons to the bonding Sigma (  ) bonds - Circular cross-section and are formed by head-on interaction Pi (  ) bonds – “dumbbell” shape from sideways interaction of p orbitals Atoms use hybrid orbitals to form bonds in organic molecules. tetrahedral geometry has four sp 3 hybrid orbitals planar geometry uses three equivalent sp 2 hybrid orbitals and one unhybridized p orbital linear uses two equivalent sp hybrid orbitals with two unhybridized p orbitals