Organic Chemistry Reviews Chapter 1

1. Organic Chemistry ReviewsChapter 1 Cindy Boulton August 30, 2009

2. What is Organic Chemistry? • Organic Chemistry: • Carbon based molecules • Tetravalent Carbon • Forms 4 bonds with 4 valence electrons • Vitalism • Concept of organic compounds before 1828 • Compounds that came from living organisms • Could not be man-made in a laboratory • “Vital force” was necessary to synthesis an organic compound

3. Important People in History • Wöhler (1828) • Disproved vitalism • Made urea from non-living substances • Stanley Miller (1953) • Imitated basic earth elements: water, methane, ammonium, hydrogen, and heat • Found Amino Acids -> Amino Acids can form life • Lavoisier (1780s) • Named “acids” • Conservation of mass in combustion and burning • Organic compounds burned -> CO₂ + H₂O • Organic compounds contain C, H, O

4. Important People in History cont. • Avogadro (1811) • Mole – how to count molecules and atoms • 6.02 × 10²³ • Kekulé (1861) • Predicted structure of Benzene • Ring shape • Structural Theory • Octet Rule: An element wants 8 valence electrons • Atoms have a fixed number of bonds using valence electrons • C-tetravalent (4), N-trivalent(3), O-divalent (2), H-monovalent (1) • Multiple bonds are allowed

5. Isomers • Isomers: • Different Compounds • Atoms arranged differently • Same Molecular Formula • Different Properties • Ex: • Ethanol • Dimethyl ether • Constitutional isomers • Different compounds with same molecular formula but differ in connectivity or sequence the atoms are bonded • Different properties

6. Bonding • Ionic Bonds • Transfer of one or more electrons from one atom to another to create ions • Ex: Na-Cl • Covalent Bonds • Sharing of electrons between atoms • Ex: C-H • Polarity of Bonds and Molecules • Atoms differ in electronegativity – ability of an atom to attract electrons • Partial charge • Ex: H2O

7. Polarity of Bonds and Molecules cont. • Electronegativity: • How much an atom wants electrons • Atoms in bonds • Not measured directly • Increase left to right and bottom to top of Periodic Table • Most electronegative element is Fluorine (not Helium) • Dipole • Vector • Direction of more electronegative element in bond • Net Dipole = overall direction of all dipoles in molecule

8. Lewis Structures • 2-Dimensional • Show connections between atoms only using valence electrons • Satisfy the Octet Rule • Exceptions • Each bond line represents 2 electrons • Covalent bonding • If necessary use multiple bonds • Example: CH4 , NH3 , and H2O • Count Valance Electrons • Know how the atoms are connected

9. VSEPR • Valance Shell Electron Pair Repulsion • 3-Dimensional • Electron clouds repel each other • Move as far apart as possible • Example: CH4 , NH3 , and H2O • 2 bonds on the plane of paper • 1 bond forward out of the paper • 1 bond backward in to the paper • Type: AX4, AX3E, AX2E2 • Angles: 109.5o, 107o, and 104o • Geometry Names: tetrahedral, trigonal pyramid, and bent

10. Formal Charge • Formal Charge = valance electrons of neutral unbonded atom – valance electrons assigned to bonded atom • Valance electrons of neutral unbonded atom found on periodic table • Valance electrons assigned to bonded atom • One electron per covalent bond + unshared electrons • Add formal charge for each atom to get overall charge of molecule • Ex: [NH4]+

11. Resonance and Resonance Hybird • A molecule or ion can be represented by two or more Lewis Structures that differ only in the positions of electrons • Hybrid – average of Lewis Structures • Partial charge and partial bond • Ex: CO32- • Three different but equivalent structures • Use curved arrows to show movement of electron pairs • Resonance arrow : ↔(not equilibrium arrow)

12. Electron Configuration • Atomic Orbitals • A region of space where the probability of finding an electron is large • 1s-2 electrons: sphere shaped • 2s-2 electrons: sphere shaped • 2p-6 electrons: dumbell (x, y, and z) • Aufbau Principle- fill lowest energy levels first • Hund’s rule- add electron to an empty orbital first • Pauli Exclusion Principle- two electrons in same orbital have different spins

13. Hybridization • Only hybridize atomic orbitals with valence electrons • Only hybridize atomic orbitals that form the molecules geometry shape. • Molecular orbitals are formed • One electron in each orbital, no pairing • All orbitals are the same energy-degenerate • Energy decreases • sp3 hybridized- use all 3 p orbitals • sp2hybirdized- 2 p orbitals hybridized and 1 p orbital unhybridized • sp hybridized- 1p orbital hybridized and 2 p orbitalsunhybridized

14. Sigma and Pi Bonds • Sigma Bond (σ) • All single bonds • Hybridized orbitals of two atoms overlap to share electrons in covalent bond • Free rotation • Pi Bond (π) • A double or triple bond • Parallel unhybirdized p orbital of two atoms overlap in both regions to share electrons in covalent bond • Restricted Rotation

15. Alkanes, Alkenes, and Alkynes • Alkanes • sp3 hybridized, tetrahedral • All C-C single bonds and all sigma bonds ( C-C and C-H) • Angles are 109.5o • Example: Methane • Alkenes • sp2 hybridized, trigonal planar • One C-C double bond: 1 sigma bond + 1 pi bond • Angles are 121o and 118o • Example: Ethene • Alkynes • sp1 hybridized, linear • One C-C triple bond: 1 sigma bond + 2 pi bonds • Angles are 180o • Example: Ethyne

16. Saturated vs. Unsaturated • Saturated • No multiple bonds- all single/sigma bonds • Maximum number of Hydrogens possible • Unsaturated • At least one pi bond (double or triple bonds) • Trans- • Opposite side • Cis- • Same side