CHM 211 (Organic Chemistry) Summer 2009

1. CHM 211 (Organic Chemistry)Summer 2009 • Dr. Ned H. Martin • Office: Dobo 242E • Telephone: • 962-3453 (campus) • Email: martinn@uncw.edu

2. Texts • Organic Chemistry, 7th edition, McMurry • Optional • Study Guide and Solutions Manual for McMurry's Organic Chemistry, 7th edition • Molecular model kit • Course Website (Syllabus, Grading Policy): http://www.uncw.edu/chem/Courses/Martinn/chm211martin/index.htm

3. Grading Policy • Four 40-minute tests, each worth 60 points. • The final exam will consist of six sections. The first four are like the four tests; the higher grade counts. Section 5 is new material (since the last test). Section 6 is comprehensive. You may take (or not) as many of the first four sections as you want. Everyone must take sections 5 and 6. • There will be no make up exams. • Each of the tests may include at least one problem from the homework assignments. Tests 2- 4 may contain one review question from the previous test. • 93%=A, 90%=A-, 87%=B+, 84%=B, 80%=B-, etc.

4. Attendance & Homework • Attendance is expected, but not officially monitored for grading purposes. • Missing 1 day in the summer is like missing 1 week during a regular semester! • Homework problems are assigned, but not collected. • Actively working the homework problems allows you to test whether you understand the material and serves as a review guide for the exams.

5. Keys to Success in CHM 211 • Memorization alone is not sufficient. • Reasoning alone is not sufficient. • Study three times: • Before the lecture • After the lecture • Before the test • Actively do problems (Keep a notebook). • Cooperate – form study groups.

6. What is Organic Chemistry? • The study of carbon-containing compounds • Important because: • Carbon forms 4 bonds, and can bond to itself in long chains • Carbon has three different geometries giving rise to a variety of structures • Carbon bonds strongly to other common elements: O, N, Cl, etc. • Organic compounds have many applications and uses: dyes, medicines, fabric, plastics, food (protein, carbohydrates, fats, oils), fuel, pesticides, paint, preservatives, hormones, etc. • This PowerPoint covers: Chapter 1. Structure and Bonding

7. C (Carbon) • Carbon’s atomic number = 6, therefore it has 6 protons in its nucleus. • A neutral atom of 12C has 6 protons, 6 neutrons and 6 electrons; its amu = 12 ( = 6p + 6n) • A neutral atom of 13C has 6 protons, 7 neutrons and 6 electrons; its amu = 13 ( = 6p + 7n) • A neutral atom of 14C has ? protons, ? neutrons and ? electrons; its amu = ? ( = ?p + ?n) • Carbon’s atomic weight = 12.011; this is a weighted average of the three isotopes: 12C, 13C, and 14C.

8. Parts of an Atom • Protons (+ charge) and neutrons (0 charge) are in the center or nucleus of the atom • Electrons (- charge) are considered to be a cloud of charge around the nucleus. Orbitals describe where the electrons are. Electrons have very little mass compared to protons and neutrons. • Electrons are found in s orbitals (spherical), p orbitals (dumbbell), or d orbitals (various shapes) • Electrons are grouped in different layers or shells.

9. 1.1 Atomic Structure • Structure of an atom • Positively charged nucleus (very dense, protons and neutrons) and small (10-15 m) • Negatively charged electrons are in a cloud (10-10 m) around nucleus • Diameter is about 2  10-10 m (200 picometers (pm)) [the unit Angstrom (Å) is 10-10 m = 100 pm]

10. 1.2 Atomic Structure: Orbitals • Quantum mechanics: describes electron energies and locations by a wave function,  • A plot of 2 describes the region where electrons are most likely to be • An electron cloud has no specific boundary so we represent its shape by the region of highest probability of finding an electron. • Solutions of the wave equation give rise to regions of electron density on each atom of specific shapes (atomic orbitals)

11. Shapes of Atomic Orbitals • Four different kinds of orbitals occupied by electrons • Denoted s, p, d, and f (listed in increasing energy) • s and p orbitals are most important in organic chemistry • s orbitals: spherical, with the nucleus at center • p orbitals: dumbbell-shaped, with the nucleus at the center

12. p-Orbitals • There are three perpendicular p orbitals, px, py, and pz, of equal energy • Lobes of a p orbital are separated by region of zero electron density, called a node.

13. 1.3 Atomic Structure: e- Configuration • The lowest energy electron configuration of an atom of any element can be predicted by following three rules: • The aufbau principle: Electrons are filled into the lowest energy orbitals first (1s, then 2s, then 2p, then 3s, then 3p, then 4s, then 3d) • The Pauli exclusion principle: Only two electrons may occupy an orbital; they must have opposite spin orientations. • Hund’s rule: If there are two or more equal energy (degenerate) orbitals available, the electrons will spread out among the orbitals with parallel spins, only pairing up after the orbitals are half-filled.

14. Examples of Electron Configuration

15. 1.4 The Nature of the Chemical Bond • Atoms form bonds because the compound that results is more stable than the separate atoms. • Ionic bonds in salts form as a result of electron transfers, followed by electrostatic attraction between opposite charges. • Organic compounds form covalent bonds by sharing electrons (G. N. Lewis, 1916). • Lewis structures show valence electrons of an atom as dots. • Hydrogen has one dot, representing its 1s electron. • Carbon has four dots (2s2 2p2). • Stable molecule results in a completed shell, an octet (eight e-) for main-group atoms (two for hydrogen).

16. Number of Covalent Bonds to an Atom • Atoms with one, two, or three valence electrons form one, two, or three bonds. • Atoms with four or more valence electrons form as many bonds as they need electrons to fill the s and p levels of their valence shells to reach a stable octet. • Carbon has four valence electrons (2s2 2p2), therefore forms four bonds (CH4).

17. Valence of Oxygen and Nitrogen • Oxygen has six valence electrons (2s2 2p4), so it forms two bonds (H2O).

18. Valence of Nitrogen • Nitrogen has five valence electrons (2s2 2p3), and it forms three bonds (NH3).

19. Non-bonding electrons • Valence electrons not used in bonding are called nonbonding electrons, or lone-pair electrons. • Consider the nitrogen atom in ammonia (NH3): • N shares six valence electrons in three covalent bonds; the remaining two valence electrons are a nonbonding (lone) pair.

20. 1.5 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. • The H–H bond results from the overlap of two singly occupied hydrogen 1s orbitals. • The H-H bond is cylindrically symmetrical, sigma (s) bond.

21. Bond Energy • The reaction 2 H· H2 releases 436 kJ/mol. • The product has 436 kJ/mol less energy than two H atoms: H–H has bond strength of 436 kJ/mol. (1 kJ = 0.2390 kcal; 1 kcal = 4.184 kJ).

22. Bond Length • Distance between nuclei that leads to maximum stability. • If too close, they repel because both nuclei are positively charged. • If nuclei are too far apart, bonding is weak.

23. 1.6 Hybridization: sp3 Orbitals and the Structure of Methane • Carbon has 4 valence electrons (2s2 2p2) • In CH4, all C–H bonds are identical (tetrahedral) • How can this be explained ??

24. 1.6 Hybridization: sp3 Orbitals and the Structure of Methane • sp3 hybrid orbitals:s orbital and three p orbitals combine to form four equivalent, unsymmetrical, tetrahedral orbitals (s+p+p+p = sp3), Pauling (1931)

25. Tetrahedral Structure of Methane • sp3 orbitals on C overlap with 1s orbitals on 4 H atoms to form four identical C-H bonds • Each C–H bond has a strength of 438 kJ/mol and length of 110 pm • Bond angle:each H–C–H is 109.5°, the tetrahedral angle.

26. 1.7 Hybridization: sp3 Orbitals and the Structure of Ethane • Two C’s bond to each other by s overlap of an sp3 orbital from each C. • The other three sp3 orbitals on each C overlap with H 1s orbitals to form six C–H bonds. • The C–H bond strength in ethane is 420 kJ/mol. • The C–C bond is 154 pm long and its strength is 376 kJ/mol. • All bond angles of ethane are tetrahedral.

27. 120 90 1.8 Hybridization: sp2 Orbitals and the Structure of Ethene (Ethylene) • sp2 hybrid orbitals: A 2s orbital of C combines with two 2p orbitals, giving 3 orbitals (s+p+p = sp2) • sp2 orbitals are in a plane with 120° angles • Remaining p orbital is perpendicular to the plane

28. Carbon-Carbon Bonds in Ethene • Two sp2-hybridized orbitals overlap to form a s bond • Two p orbitals overlap side-to-side to form a pi () bond • sp2–sp2s bond and 2p–2p bond results in sharing four electrons and formation of C=C double bond • Electrons in the s bond are centered between nuclei • Electrons in the  bond occupy regions on either side of a line between nuclei, above and below the plane of the atoms.

29. Carbon-Hydrogen Bonds in Ethene • Each of 4 H atoms form s bonds with four sp2 orbitals • H–C–H and H–C–C bond angles are about 120° • C=C double bond in ethene is shorter and stronger than the C-C single bond in ethane • The ethene C=C bond length is 133 pm (Recall that the C–C bond length in ethane is 154 pm) • The C+C bond strength is 611 kJ/mol, less than twice the strength of a C-C (2 x 376 = 752).

30. 1.9 Hybridization: spOrbitals and the Structure of Acetylene • The in acetylene (ethyne) is a triple bond, with the carbons sharing six electrons • A carbon 2s orbital hybridizes with a single p orbital giving two sp hybrids • The other two p orbitals on each C remain unchanged • sp orbitals are linear, oriented 180° apart (on x-axis) • The two p orbitals are perpendicular, on the y-axis and the z-axis

31. Orbitals of Acetylene • Two sp hybrid orbitals from each C overlap to form an sp–sps bond. • Two pz orbitals from each C form a pz–pz bond by sideways overlap; py orbitals overlap similarly to form a second  bond.

32. Bonding in Acetylene • Sharing of six electrons forms a . • Two sp orbitals form s bonds with hydrogens. • The bond strength is 835 kJ/mol, much less than three times the strength of a C-C (3 x 376 = 1128). The bond length is 120 pm.

33. 1.10 Hybridization of Other Elements • Elements other than C can have hybridized orbitals. • The H–N–H bond angle in ammonia (NH3) is 107.3°, close to the tetrahedral 109.5°. • N’s orbitals (s+p+p+p) hybridize to form four sp3 orbitals. • One sp3 orbital holds two nonbonding electrons, and three sp3 orbitals have one electron each, forming s bonds to three Hs.

34. Hybridization of Oxygen in Water • The oxygen atom is sp3-hybridized. • Oxygen has six valence-shell electrons but forms only two covalent bonds, leaving two lone pairs. • The H–O–H bond angle is 104.5°, slightly smaller than the perfect tetrahedral angle (109.5º) because of electron-electron repulsion between the lone pairs.

35. 1.11 Molecular Orbital Theory • A molecular orbital (MO): where electrons are most likely to be found (specific energy and general shape) in a molecule. • The two (or more) atomic orbitals combine to make two (or more) molecular orbitals. • Additive combination (bonding) MO is lower in energy. • Subtractive combination (antibonding) MO is higher.

36. Molecular Orbitals in Ethene • The  bonding MO results from combining p orbital lobes with the same algebraic sign. • The  antibonding MO comes from combining lobes with opposite signs. • Only the bonding MO is occupied by electrons.

37. Summary • Organic chemistry – chemistry of carbon compounds • Atom: positively charged nucleus surrounded by negatively charged electrons • Electrons occupy orbitals around the nucleus. • Different orbitals have different energy levels and different shapes • s orbitals are spherical, p orbitals are dumbbell-shaped • Covalent bonds - electron pair is shared between atoms • Valence bond theory - electron sharing occurs by overlap of two atomic orbitals

38. Summary, cont’d • Hybrid Atomic Orbital Theory - electron sharing occurs by overlap of two orbitals formed by combining (hybridizing) two or more atomic orbitals (sp, sp2, sp3) • Molecular orbital (MO) theory - bonds result from combination of atomic orbitals to give molecular orbitals, which belong to the entire molecule • Sigma (s) bonds - Circular in cross-section and are formed by head-on interaction • Pi () bonds – “dumbbell” shape, from sideways interaction of p orbitals; located above and below the s bond framework of the molecule

39. Summary, cont’d. • Carbon uses hybrid orbitals to form bonds in organic molecules. • In single bonds with tetrahedral geometry, carbon has four sp3 hybrid orbitals • In double bonds with planar geometry, carbon uses three equivalent sp2 hybrid orbitals and one unhybridized p orbital • Carbon uses two equivalent sp hybrid orbitals to form a triple bond with linear geometry, with two unhybridized p orbitals • Atoms such as nitrogen and oxygen also hybridize to form strong, oriented bonds • The nitrogen atom in ammonia and the oxygen atom in water are sp3-hybridized

40. Quick Review • Carbon • One s and three p orbitals hybridize to form four sp3orbitals • In methane and ethane, C is tetrahedral, with ~109.5° bond angles • In ethene, One s and two p orbitals hybridize to form three sp2 orbitals. The bonds between the nuclei are the  bonds from the overlapped sp2orbitals. The remaining p orbitals overlap side-to-side to form a  bond. C-C p bonds are weaker than C-C s bonds.