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Organic Chemistry Review Part II
Organic Chemistry: Carbon Atom Structural Classifications Atomic Theory Dipoles & Resonance Isomers Functional Groups Organic Reactions
Organic Chemistry • The chemistry of compounds which contain carbon. • Carbon forms more compounds than any other element, except hydrogen.
Organic Chemistry Major Concepts • Structural Classifications • Hybridization • Charges of Organic Molecules • Dipoles & Dipolar Resonance • Isomers • Functional Groups • Organic Reactions
Structural Classification of Carbon Atoms Three main classifications are: • Primary Carbons • Secondary Carbons • Tertiary Carbons • Quaternary Carbons
Primary Carbons • Denoted as 1° carbons. • Also called terminal or end carbon atoms. • Found at the ends of a straight chains or the branches. • Covalently bonded to one carbon atom.
Secondary Carbons • Denoted as 2° carbons. • Covalently bonded to two other carbon atoms.
Tertiary Carbons • Denoted as 3° carbons. • Covalently bonded to three other carbon atoms.
Quaternary Carbons • Denoted as 4° carbons. • Covalently bonded to four other carbon atoms.
Definitions Valence Bond Theory: • Electrons in a covalent bond reside in a region in which there is overlap of individual atomic orbitals. • For example, the covalent bond in molecular methane (CH4) requires the overlap of valence electrons:
Definitions • Types of valence bond theory overlap:
Definitions Valence Shell Electron Pair Repulsion (VSEPR) • Electron pairs arrange themselves around an atom in order to minimize repulsions between pairs. • Carbon has a valence of fourand must have a tetrahedral geometry. • In methane, each carbon atom must have a bond angle of 109.5⁰. This is the largest bond angle that can be attained between all four bonding pairs at once.
Definitions Hybridization: • Atomic orbitals modify themselves to meet VESPR geometry and valence bond theory. • Three types of hybridization for carbon:
Hybridizations • In sp3 hybridization, an electron is promoted from a 2s orbital into a p orbital. • The 2s orbital and three 2p orbitals form four hybrid orbitals (sp3). • Ground state: 1s2 2s2 2Px1 2Py1 2Pz0 • Excited state: 1s2 2s12Px1 2Py1 2Pz1
Hybridizations • The overlap of each hybrid orbital with a hydrogen atom results in a sigma bond ( σ bond). • Only one σ bond can exist between two atoms.
Hybridizations sp3 hybridization of methane:
Hybridizations sp3 hybridization of ethane:
Hybridizations • In sp2 hybridization, the 2s orbital and two of the 2p orbitals form three hybrid orbitals (sp2). • The Pz orbital of each carbon atom remains unhybridized. • These unhybridized Pzorbitals overlap with one another to form a π-bond.
Hybridizations sp2 hybridization of ethene:
Hybridizations sp2 hybridization and bond rotation:
Hybridizations • In sp hybridization, the 2s orbital and one 2p orbital form two hybrid orbitals (sp). • The triple bond is actually one σ bond and two π bonds.
Hybridizations • sp hybridization of ethyne: No free rotation
Definitions Dipole: • The measure of net molecular polarity. • Formula: the magnitude of the charge Qtimes the distance r between the charges. μ = Q × r • The larger the difference in electronegativities of the bonded atoms, the larger the dipole moment.
Definitions Resonance: • Part of the Valence Bond Theory • Describes the delocalization of electrons within molecules. • Used when Lewis structures for a single molecule cannot describe the actual bond lengths between atoms. • Structures are not isomers of the target molecule, since they only differ by the position of delocalized electrons.
Definitions Resonance Hybrid: • The net sum of valid resonance structures. • Several structures represent the overall delocalization of electrons within the molecule. • A molecule that has several resonance structures is more stable than one with fewer.
Definitions Hyperconjugation: • The interaction of the electrons in a sigma bond (usually C–H or C–C) with an adjacent empty (or partially filled) non-bonding p-orbital, antibonding π orbital, or filled π orbital. • Only electrons in bonds that are β to the positively charged carbon can stabilize a carbocation by hyperconjugation.
Carbon Atom Dipoles • Carbon- Halogen Bonds
Carbon Atom Dipoles • C-O, C-S and C-N Covalent Bonds: δ+ δ- δ+ δ-
Hyperconjugation • A.K.A "no bond resonance". • The delocalization of σ-electrons or lone pair of electrons into adjacent π-orbital or p-orbital. • Overlapping of σ-bonding orbital or the orbital containing a lone pair with adjacent π-orbital or p-orbital. • An α- carbon next to the π bond, carbocation or free radical should be sp3 hybridized with at least one hydrogen atom bonded to it.
Hyperconjugation • Other hydrogens on the methyl group also participate due to free rotation of the C-C bond. • There is NO bond between an α-carbon and one of the hydrogen atoms. • The hydrogen atom is completely detached from the structure. • The C-C bond acquires some double bond character and C=C acquires some single bond character.
Isomers Compounds that have: • The same molecular formula. • Similar or different types of structural formulas. • Different arrangement of atoms.
Isomers: Two main classes are: • Structural or constitutional • Stereoisomers
Structural Isomers • Also known as constitutional isomers
Stereoisomers • Configurational • Geometric or Diastereomers • Optical or Enantiomers • Conformational or Rotamers
