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Unit 1

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  1. Unit 1 • Functional Groups • Depicting Structures of Organic Compounds • Lewis Structures • Condensed structural formulas • Line angle drawings • 3-dimensional structures • Resonance Structures • Acid-Base Reactions • Curved Arrows

  2. Classes of Organic Compounds • Organic compounds are commonly classified and named based on the type of functional group present. • An atom or group of atoms that influences the way the molecule functions, reacts or behaves. • The center of reactivity in an organic compound

  3. Classes of Organic Compounds • You must be able to recognize and draw the functional groups listed in your syllabus and on the following slides. • Use the following slides and the tables given in the front of your text and in the chapter to learn these.

  4. Functional Groups

  5. Functional Groups

  6. Functional Groups

  7. Alkanes • Contain C-C single bonds • no functional group • Nonpolar covalent bonds • electrons shared equally • Tetrahedral electron domain geometry • sp3 hybridized carbons

  8. Cycloalkanes • Contain C – C with at least 3 of the carbons arranged in a cyclic (ring) structure • No functional group • Nonpolar • Tetrahedral • sp3 hybrid orbitals

  9. Alkenes • Contain C=C (carbon-carbon double bonds) • 1 sigma bond & 1 pi bond • Non-polar • Trigonal planar geometry • sp2 hybridized carbons Which atoms must be coplanar in an alkene?

  10. Alkenes • The C=C present in an alkene is composed of 1 sigma (s) bond and 1 pi (p) bond. • A sigma (s) bond forms when two orbitals overlap end to end. • electron density is centered along the internuclear axis • cylindrically symmetrical Internuclear axis

  11. Alkenes • A pi (p) bond forms when two p orbitals overlap side to side • electron density is located above and below the internuclear axis • oriented parallel to the internuclear axis

  12. Alkenes • Free rotation occurs around single bonds. • Double and triple bonds are rigid. • Cannot rotate freely. • Rotation would cause loss of overlap of the p orbitals, destroying the p bond.

  13. Alkynes • Contain C Ctriple bonds • 1 sigma bond • 2 pi bonds • Nonpolar • Linear electron domain geometry • sp hybridized carbons Which atoms must be co-linear in an alkyne?

  14. Aromatic Ring • Planar ring system with alternating single and double bonds • does not react like an alkene • Nonpolar • Trigonal planar • sp2 hybridized carbons • Benzene ring is a very common aromatic ring.

  15. Alkyl Halides • Contain C-halogen bond • F, Cl, Br, or I • Polar covalent C-halogen bond • electrons shared unequally • Many (but not all) alkyl halides are polar molecules as well.

  16. Polar Covalent Bonds • Two ways to indicate bond polarity • partial charges • d- on more electronegative element • d+ on less electronegative element • direction of dipole moment • + on less electronegative element • arrow pointing toward more electronegative element

  17. Polar Covalent Bonds • Dipole moment: A measure of the separation and magnitude of the positive and negative charges in polar bonds or polar molecules. • Important Polar Covalent Bonds: C-O C-N C-halogen O-H N-H H-halogen • Important Nonpolar Covalent Bonds: C-C C-H halogen-halogen

  18. d- d+ d+ Polar Molecules • Polar Molecule: • a molecule with a non-zero molecular dipole moment • net negative end, net positive end • contains 1 or more polar covalent bonds arranged asymmetrically within the molecule • use molecular geometry to determine polarity • exhibit dipole-dipole interactions • How does the BP of a polar molecule compare to the BP of a nonpolar molecule with similar molar mass?

  19. Nonpolar Molecules • Nonpolar molecules: • a zero (or very small) molecular dipole moment • contain either: • only nonpolar covalent bonds • 2 or more polar covalent bonds arranged symmetrically • London dispersion forces • found in all molecules • only IMF found in nonpolar molecules • strength increases as surface area increases • What variables make SA increase? • What happens to the BP as SA ?

  20. Polar/Nonpolar Molecules Example: Which of the following are polar molecules?

  21. Alcohols • Contain C-O-H bond • hydroxyl group • Polar covalent bonds • C-O bond • O-H bond • Polar molecule • dipole-dipole interactions • hydrogen bonding • What causes hydrogen bonding? • How does hydrogen bonding affect BP?

  22. Ethers • Contain C-O-C bond • Polar covalent bonds • tetrahedrale.d. geometry • bent molecular geo. • Polar molecules • dipole-dipole interactions

  23. primary secondary tertiary Amines • Contain C-N-R R’ • Polar covalent bonds • Polar molecules • Dipole-dipole interactions • Hydrogen bonding (1o and 2o) • Common organic bases • lone pair of e- on N

  24. O C - H Aldehydes • Contain (-CHO) • Carbonyl (C=O) • always on the 1st or last carbon in a chain • trigonal planar geometry • sp2 hybrid orbitals • Polar covalent bond • Polar molecule • dipole-dipole interactions

  25. O C-C-C Ketones • Contain • Carbonyl attached to middle of chain • Polar covalent bond • Polar molecule • dipole-dipole interactions • Trigonal planar e.d. geo. • sp2 hybridized C

  26. Carboxylic Acids • Contain carboxyl group • Polar covalent bonds • Polar molecules • dipole-dipole • hydrogen bonding • trigonal planar • sp2 hybridized carbon

  27. Acid Chloride • Contain • Polar molecule • dipole-dipole forces • Trigonal planar geo. • sp2 hybridized C • Lachrymators

  28. Esters • Contain • Polar covalent bonds • Polar molecules • dipole-dipole interactions • trigonal planar • sp2 hybridized

  29. Amides • Contain • where R and R’ = H or alkyl • Polar covalent bonds • Polar molecules • dipole-dipole interactions • hydrogen bonding (1o and 2o) • C=O is trigonal planar & sp2 hyrbridized

  30. Nitriles • Contain • Polar covalent bond • Polar molecule • dipole-dipole • Linear, sp hybridized C

  31. Functional Groups Example: Identify the functional groups present in the following compounds. testosterone Vanillin Lisinopril

  32. Depicting Structures of Organic Compounds • Organic compounds can be described using a variety of formulas: • Empirical formula • Molecular formula • Lewis structure • Full structural formula • Three dimensional drawings • Condensed structural formula • Line angle drawings

  33. Depicting Structures of Organic Compounds • Ethyl acetate is an organic molecule with: • empirical formula = C2H4O • lowest whole number ratio • molecular formula = C4H8O2 • actual number of each type of atom present

  34. Depicting Structures of Organic Compounds • Ethyl acetate is an organic molecule with: • Lewis structure: • depicts all covalent bonds using a straight line and shows all nonbonding pairs of electrons • What are covalent bonds? • Full structural formula: • a Lewis structure without the nonbonding electrons

  35. Depicting Structures of Organic Compounds • Ethyl acetate is an organic molecule with: • 3-d drawing: • Condensed structural formula • Line angle drawing

  36. Lewis Structures • Lewis structures are used to represent the covalent bonds present in a molecule. • Symbol for each atom • Covalent bonds between atoms depicted using a solid line • Unshared electrons are shown around the appropriate atom

  37. Lewis Structures • To draw a Lewis structure: • Count the number of valence electrons • For a cation (+), subtract 1 electron for each positive charge • NH4+ : 5 + 4 (1) -1 = 8 e- • For an anion (-), add 1 electron for each negative charge • CN-: 4 + 5 + 1 = 10 e-

  38. Lewis Structures • Draw a skeleton structureshowing the chemical symbol for each atom. Connect the appropriate atoms using a single bond. • Skeletons for organic compounds: • The backbone generally contains C-C bonds • N,O, and S can either be part of the backbone or attached to one of the carbons as a substituent. • H will be attached to C,N,S and/or O • Halogens will be attached to C as substituents.

  39. Lewis Structures • Add pairs of electrons to the atoms giving each one an octet • H only gets 2 electrons • Try filling the octets of N, O, S, and halogens first • There will generally NOT be any “leftover” electrons for organic compounds. • Organic ions, however, may have leftover electrons. • Put them on an atom that needs an octet.

  40. N N Lewis Structures • If there are not enough electrons to give all atoms an octet, share electrons to form multiple bonds. • Single bond: • one pair of electrons shared • double bond: • two pairs of electrons shared • triple bond: • three pairs of electrons shared

  41. Lewis Structures Common neutral bonding patterns CNOHHalogens total bonds 4 3 2 1 1 lone pairs 0 1 2 0 3

  42. Lewis Structures Example: Draw the Lewis structure for C2H3I.

  43. Lewis Structures Example: Draw the Lewis structure for a ketone with the molecular formula, C5H9BrO.

  44. Lewis Structures Example: Draw two possible Lewis structures for CH2COCH3-. These two Lewis structures are resonance structures.

  45. Formal Charge • Which atom in each of the previous Lewis structures is negatively charged? • Formal charge provides a method for keeping track of electrons in a compound. • determines which atom(s) in a structure bear(s) the charge in a polyatomic ion • identifies charged atoms within a molecule that is neutral overall.

  46. Formal Charge • Formal charge: • a calculated value that compares the number of valence electrons for a particular atom to the number of electrons assigned to that atom in a Lewis structure • FC = group # - nonbonding e- - 1/2 (bonding e-) -

  47. Positive Neutral Negative C C+ C C - N N+ N N- O O+ O O - Formal Charge Common Organic Bonding Patterns and Formal Charges:

  48. Formal Charge Example: Calculate or determine the formal charge on N and O in (CH3)3NO. All Lewis structures you draw from now on should include any non-zero formal charges that are present.