Welcome to Organic Chemistry 234!

1. Welcome to Organic Chemistry 234!

2. How Should I Study? • Do not memorize everything! • Practice writing mechanisms and “talking” yourself through the steps. • Learn to ask the right questions. • Form a small study group (2-3 people). • Work as many problems as you can. • Do not hesitate to visit me during office hours for assistance. • A free tutoring service is available through the LRC.

3. What is Organic Chemistry? • It is the study of carbon-containing compounds

4. Why Carbon? • Carbon neither gives up nor accepts electrons because it is in the center of the second periodic row. • Consequently, carbon forms bonds with other carbons and other atoms by sharing electrons. • The capacity of carbon to form bonds in this fashion makes it the building block of all living organisms.

5. Why Study Organic Chemistry? • Since carbon is the building block of all living organisms, a knowledge of Organic Chemistry is a prerequisite to understanding Biochemistry, Medicinal Chemistry, Chemical Ecology and Pharmacology. • Indeed, Organic Chemistry is a required course for studying Pharmacy, Medicine, and Dentistry. • Admission into these professional programs is highly dependent on your performance in Organic Chemistry.

6. Examples of Organic Compounds Used as Drugs Methotrexate, Anticancer Drug 5-Fluorouracil, Colon Cancer Drug Tamiflu, Influenza Drug AZT, HIV Drug

7. Examples of Organic Compounds Used as Drugs Haldol, Antipsychotic Elavil, Antidepressant Prozac, Antidepressant Viagra, TreatsErectile Dysfunction

8. Fall 2012 Dr. Halligan CHM 234 Chapter 1 • Electronic Structure and Bonding • Acids and Bases

9. “Speaking Organic Chemistry” • What are some of the fundamentals of organic chemistry that we will cover in Chapter 1? • The periodic table • Bonding • Lewis structures • Delocalized electrons and Resonance Structures • Orbital Hybridization • The art of drawing structures and comprehending organic compounds • Trends in electronegativity • Determination of formal charges • The use of molecular models to represent compounds • Acids and Bases

10. Structure and Bonding Note: Sections 1.1 and 1.2 on the structure of an atom can be reviewed in the textbook.

11. Ionic, Covalent, and Polar Bonds • Bonds formed between two oppositely charged ions are considered ionic. These attractive forces are called electrostatic attractions. • In addition to NaCl, what are some examples of compounds with ionic bonds?

12. Covalent Bonding • In covalent bonding, electrons are shared rather than transferred. • Most elements tend to form covalent bonds rather than ionic bonds because a gain or loss of multiple electrons (to achieve the octet) is too high in energy. • e.g. carbon would have to lose 4 electrons or gain 4 electrons in order to participate in ionic bonding. • What are some examples of compounds with covalent bonds?

13. Common Bonding Patterns in Organic Compounds and Ions

14. Equal sharing of electrons: nonpolar covalent bond • (e.g., H2) • Sharing of electrons between atoms of different • electronegativities: polar covalent bond (e.g., HF)

15. A polar covalent bond has a slight positive charge on one end and a slight negative charge on the other

17. (e) : magnitude of the charge on the atom (d) : distance between the two charges A Polar Bond Has a Dipole Moment • A polar bond has a negative end and a positive end dipole moment (D) = m = e x d

18. Molecular Dipole Moment The vector sum of the magnitude and the direction of the individual bond dipole determines the overall dipole moment of a molecule

19. Electrostatic Potential Maps

20. Lewis Structures • Lewis structures are representations of compounds in which lines and dots are used to indicate electrons. A bond line is equal to 2 electrons. • Keep in mind the number of valence electrons that each atom should have (i.e. In which group is the atom located?). • If the atoms in a molecule are to contain charges, think about electronegativity and which atoms will better bear the particular charge.

21. Formal Charge • Formal charge is the charge assigned to individual atoms in a Lewis structure. • By calculating formal charge, we determine how the number of electrons around a particular atom compares to its number of valence electrons. Formal charge is calculated as follows: • The number of electrons “owned” by an atom is determined by its number of bonds and lone pairs. • An atom “owns” all of its unshared electrons and half of its shared electrons.

22. Formal Charge • Determine the formal charge for each atom in the following molecule:

23. Nitrogen has five valence electrons Carbon has four valence electrons Hydrogen has one valence electron and halogen has seven

24. Important Bond Numbers Neutral Cationic Anionic

25. Non-Octet Species • In the 3rd and 4th rows, expansion beyond the octet to 10 and 12 electrons is possible. Sulfuric Acid Periodic Acid Phosphoric Acid • Reactive species without an octet such as radicals, carbocations, carbenes, and electropositive atoms (boron, beryllium). Nitric Oxide Radical, Mammalian Signaling Agent Radical Carbocation Carbene Borane

26. Practice Problems • Count the number of carbon atoms in each of the following drawings.

27. How to Draw Line Angle Structures • Carbon atoms in a straight chain are drawn in a zigzag format. • When drawing double bonds, try to draw the other bonds as far away from the double bond as possible. • When drawing each carbon atom in a zigzag, try to draw all of the bonds as far apart as possible. • In line angle structures, we do draw any H’s that are connected to atoms other than carbon. • It is good practice to draw in the lone pairs for heteroatoms.

28. An orbital tells us the volume of space around the nucleus where an electron is most likely to be found The s Orbitals

29. The p Orbitals

30. Molecular Orbitals • Molecular orbitals belong to the whole molecule. • s bond: formed by overlapping of two s orbitals. • Bond strength/bond dissociation: energy required to • break a bond or energy released to form a bond.

32. In-phase overlap forms a bonding MO; out-of-phase overlap forms an antibonding MO:

33. Sigma bond (s) is formed by end-on overlap of two p orbitals: A s bond is stronger than a p bond

34. Pi bond (p) is formed by sideways overlap of two parallel p orbitals:

35. Bonding in Methane

36. Hybridization of One s and Three p Orbitals

37. The orbitals used in bond formation determine the bond angles • Tetrahedral bond angle: 109.5° • Electron pairs spread themselves into space as far from • each other as possible

38. The Bonds in Ethane

39. Hybrid Orbitals of Ethane

40. Bonding in Ethene: A Double Bond

41. Bonding in Ethyne: A Triple Bond

42. Bonding in the Methyl Cation

43. Bonding in the Methyl Radical

44. Bonding in the Methyl Anion

45. Bonding in Water

46. Bonding in Ammonia and in the Ammonium Ion

47. Bonding in Hydrogen Halides

48. Summary • The shorter the bond, the stronger it is • The greater the electron density in the region of orbital • overlap, the stronger is the bond • The more s character, the shorter and stronger is the • bond • The more s character, the larger is the bond angle

49. Brønsted–Lowry Acids and Bases • Acid donates a proton • Base accepts a proton • Strong reacts to give weak • The weaker the base, the stronger is its conjugate acid • Stable bases are weak bases

50. An Acid/Base Equilibrium Ka: The acid dissociation constant. The stronger the acid, the larger its Ka value and the smaller its pKa value.