1 / 15

Organic Chemistry Reviews Chapter 1

Organic Chemistry Reviews Chapter 1. Cindy Boulton September 7, 2008. 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

taryn
Télécharger la présentation

Organic Chemistry Reviews Chapter 1

An Image/Link below is provided (as is) to download presentation Download Policy: Content on the Website is provided to you AS IS for your information and personal use and may not be sold / licensed / shared on other websites without getting consent from its author. Content is provided to you AS IS for your information and personal use only. Download presentation by click this link. While downloading, if for some reason you are not able to download a presentation, the publisher may have deleted the file from their server. During download, if you can't get a presentation, the file might be deleted by the publisher.

E N D

Presentation Transcript

  1. Organic Chemistry ReviewsChapter 1 Cindy Boulton September 7, 2008

  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. 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

  8. 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 • Angles: 109.5 o , 107 o , and 104 o • Geometry Names: tetrahedral, trigonal pyramid, and bent

  9. 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]+

  10. 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)

  11. 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

  12. 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

  13. 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

  14. 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

  15. 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

More Related