1. McMurry Chapter 1 Structure & Bonding Organic Chemistry I S. Imbriglio

2. Organic Chemistry: What is it? 1770: Organic chemistry referred to compounds isolated from living things. __________ belief that a “magic” force, present in plants and animals, is required to make organic compounds 1789: Lavoisier observed that organic compounds are composed primarily of __________________________

3. Organic Chemistry: What is it? 1816: Michel Chevreul showed that one organic compound could be converted into others. 1828: Friedrich Wohler converted an inorganic compound into an organic compound.

4. Organic Chemistry: the study of carbon compounds ________of all known compounds are composed of carbon (~30 million known compounds) ________of chemists define themselves as organic Organic chemistry is crucial to our way of life: Pharmaceutical, Petroleum, Materials/Polymer, OUR BODIES!

5. Why Carbon? What makes carbon so special? (3 things)

6. Why Carbon? Carbon atoms form __________________ bonds to other atoms (including other carbon atoms). Incredible Structural Diversity!

7. Review of the Atom Atomic Structure: –___________ charged nucleus – dense and small –_______________ charged electrons in a cloud around the nucleus –Diameter approximately 2  10 -10 m

8. Review of the Atom Orbitals: - Three-dimensional shapes indicating where the electron is most likely to be found – s, p, d, f - We will focus on s and p orbitals – carbon atoms do not have d or f orbitals

9. Review of the Atom Orbitals: - Orbitals are organized into different shells and subshells - In the ground state of an atom, electrons occupy the lowest energy orbitals

10. Review of the Atom Electron Configurations: –_________ electrons per orbital (max) –When two electrons occupy the same orbital, they must have opposite spins (Pauli Exclusion Principle) –If two or more empty orbitals of equal energy are available, electrons occupy each with spins parallel until all orbitals have one electron (Hund’s Rule)

11. Write the electron configuration for oxygen. O

12. Review of Chemical Bonding Atoms bond together because the compound that results is more _________ than the separate atoms – bond formation is always exothermic. The Octet Rule: In most cases, atoms will _______ __________________ electrons to gain an octet.

13. Review of Chemical Bonding An _____________ is the electrostatic attraction between a cation (____) and an anion (____). The two bonded atoms do not share electrons.

14. Review of Chemical Bonding ___________________ occurs between atoms of similar electronegativity – very important in organic chemistry. eg. CCl 4 is covalent... Why? Large thermodynamic penalty for ionization of carbon to C 4+.

15. Review of Chemical Bonding Instead of transferring electrons, each chlorine atom shares one valence electron with carbon so that every atom has a filled valence shell (an octet). Lone Pairs: ____________________________________________ Lone Pairs

16. Review of Chemical Bonding Covalently bonded compounds are represented by Lewis or Kekulé structures. Lone pairs will be included in this class.

17. Review of Chemical Bonding The number of covalent bonds that a main group element must form to achieve an octet equals eight minus its group number.

18. Keeping the common bonding patterns in mind, draw a valid Kekulé structure for each of the following formulas. Include all lone pairs. CH 3 NH 2 H 2 O 2 C 3 H 8

19. Review of Chemical Bonding Multiple Covalent Bonds Atoms can share more than one electron pair to gain a full octet. –_________________: two electron pairs shared between two atoms –_________________: three electron pairs shared between two atoms

20. Review of Chemical Bonding Atoms will form combinations of single, double and triple bonds to gain a full octet. 4 total bonds3 total bonds2 total bonds1 bond

21. Draw a valid Kekulé structure for each of the following formulas. Include all lone pairs. CH 3 CO 2 H HCN

22. Review of Chemical Bonding Valence Bond Theory Covalent bond forms when two atoms approach each other closely so that a singly occupied orbital on one atom overlaps a singly occupied orbital on the other atom Electrons are paired in the overlapping orbitals and are attracted to nuclei of both atoms

23. Carbon uses its four valence electrons to form four covalent bonds. Methane is a perfect tetrahedron – all of the C-H bonds are equivalent. The four orbitals that make those bonds must also be equivalent. Review of Chemical Bonding

24. Hybridization Atomic orbitals (s, p) on the same atom can be combined to form hybrid orbitals (sp, sp 2, sp 3 ) with geometries similar to those observed experimentally. Hybrid Orbitals: – Are more directional (better bonding overlap with orbitals on other atoms) – Minimize electron-electron repulsion (think VSEPR)

25. Review of Chemical Bonding sp 3 Hybridization: One s and three p orbitals combine to form four new sp 3 hybrid orbitals. The large lobes of the four sp 3 orbitals are pointed 109.5  from each other – a tetrahedral geometry. large lobe back lobe

26. Review of Chemical Bonding sp 3 Hybridization in Methane Each sp 3 orbital on carbon overlaps with a 1s orbital on hydrogen to form four sigma bonds. ___________________: cylindrically symmetrical bond resulting from head-on overlap of two orbitals along the bonding axis Orbital Picture Ball & Stick Model Space-Filling Model

27. Review of Chemical Bonding sp 3 Hybridization in Ethane The two carbon atoms are bonded together by sigma overlap between the two sp 3 orbitals. Three sp 3 orbitals on each carbon are used to form C-H sigma bonds.

28. Review of Chemical Bonding sp 2 Hybridization: One s orbital combines with two p orbitals to give three new sp 2 hybrid orbitals One p orbital is left unhybridized. The large lobes of the sp 2 orbitals are pointed 120  from each other – a trigonal planar geometry.

29. Review of Chemical Bonding sp 2 Hybridization in Ethylene Two sp 2 orbitals overlap head-on to form a sigma bond – Electrons in the sigma bond are along the bonding axis Two p orbitals overlap side-by-side to form a pi bond – Electrons in the pi bond occupy regions above and below the bonding axis (not cylindrically symmetrical)

30. Review of Chemical Bonding sp Hybridization: One s and one p orbital combine to form two new sp hybrid orbitals Two p orbitals left unhybridized The large lobes of the sp orbitals are pointed 180  from each other – a linear geometry

31. Review of Chemical Bonding sp Hybridization in Acetylene Two sp orbitals overlap head-on to form a sigma bond Two vertical p orbitals overlap side-by-side to form pi bond Two horizontal p orbitals overlap side-by-side to form pi bond

32. Review of Chemical Bonding Assigning Hybridization to Atoms The hybridization of an atom in a molecule can be determined by counting the number of hybrid orbitals the atom is using. Hybrid orbitals are used to: – Form sigma bonds (not pi bonds) – Hold lone pairs

33. Review of Chemical Bonding Assigning Hybridization to Atoms The number of hybrid orbitals on an atom must be equal to the number of atomic orbitals that were “mixed together” to form the hybrid orbitals. eg. One s orbital plus three p orbitals gives four sp 3 orbitals

34. Review of Chemical Bonding Assigning Hybridization to Atoms Heteroatoms also use hybrid orbitals to form sigma bonds and hold lone pairs. Below are examples of sp 3 -hybridized nitrogen and oxygen.

35. Determine the hybridization of each non-hydrogen atom in the following molecules.

36. Representing Molecules Condensed Structures

37. Representing Molecules Lewis and Kekulé structures are only adequate for very small molecules.

38. Representing Molecules Skeletal Structures (Line-Angle Formulas) Rules for Drawing Skeletal Structures 1.Do not draw carbon atoms. A carbon atom is assumed to be at each intersection and at the end of each line. 2.Do not draw hydrogen atoms bonded to carbon. Assume enough C-H bonds to give each carbon a filled valence. 3.Draw all heteroatoms and attached hydrogens.

39. Representing Molecules

40. Convert the following Kekulé structures to skeletal structures. Show all lone pairs.

41. Representing Molecules Dashes and Wedges: Molecules are not flat.