General Class Information.

Size: px

Start display at page:

Download "General Class Information."

Brett Hill
5 years ago
Views:

1 General Class Information Instructors: Lectures: Recitations: Text: B. Imperiali & S. 'Connor utline, Syllabus & Suggested Reading on Website Start Second Week; See andout for Policy on Changes "rganic Chemistry" 7th Edition by McMurray Molecular Models: ighly Recommended! Allowed in Exams Problem Sets: Exams: Grading: Posted each wednesday (except exam weeks) Key posted following Monday Not graded, but doing them is required for good exam performance Four 50 minute Exams (Drop Lowest Score) & Final 65% In-Class Exams; 35% Final Exam

2 Study Skills (ow to Pass 5.12 with Flying Colors) Read before coming to lecture. Attend all lectures. Attend all recitations. Complete and understand all problem sets. Do suggested problems and practice exams without the key. Confused? Ask questions immediately.

3 rganic Chemistry: What is it? 1780: rganic compounds obtained from living sources (complex) Vitalism: Belief that a "magic" force, present in plants and animals, is necessary for the synthesis of organic compounds 1789: Antoine Laurent Lavoisier observed that organic compounds are composed primarily of carbon and hydrogen 1828: Friedrich Wohler synthesized urea from inorganic compounds!! Pb + NC N + 4 N + 4 NC heat C 2 N N 2 lead ammonium ammonium cyanate hydroxide cyanate urea inorganic organic Modern Definition: rganic chemistry is the chemistry of carbon compounds.

4 "The Age of rganic Chemistry" > 95% of All Known Compounds Composed of Carbon Crucial to ur Way of Life: Pharmaceuticals, Petroleum, Materials, Polymers (Clothing), UR BDIES STRUCTURE Determining the Way in Which Atoms Are Put Together in Space to Form Complex Molecules MECANISM Understanding the Reactivity of Molecules: ow and Why Chemical Reactions Take Place METDLGY Discovering, Developing, and Improving Chemical Reactions SYNTESIS Building Complex Molecules From Simple Molecules Using Chemical Reactions

5 Why Carbon? Carbon forms a variety of strong covalent bonds to itself and other atoms. INCREDIBLE STRUCTURAL DIVERSITY! N 2 C 3 C 3 N N N C 3 N N DNA Bases N ormones 2 N C 3 2 N C C C 2 2 N C C C 3 Carbohydrates Amino Acids

6 5.12 rganic Chemistry Fall 2008 utline: Lectures 1 and 2 I. Chemical Bonding A. ybridization B. ow to draw molecules and mechanisms C. Electronegativity and formal charge D. Resonance Reading: McMurray, 7th Edition Chapter 1 (review), Problem Set 1 Posted September 3 Key Posted Monday September 8

7 I. Chemical Bonding: Review from 5.111, 5.112, Lewis Bonding Theory: Atoms transfer or share electrons to gain a full valence. (Carbon wants to form four bonds to gain a full octet.) Molecular rbital Theory: Atomic orbitals on different atoms combine to form stable molecular bonding orbitals. (Carbon forms stable bonds to hydrogen, nitrogen, oxygen, sulfur, halides, etc.) VSEPR Theory: Electron repulsion favors certain geometries (divalent = linear, trivalent = trigonal planar, tetravalent = tetrahedral). (Carbon atoms can be linear (180 ), trigonal planar (120 ), or tetrahedral (109 ).

8 I. Chemical Bonding in 5.12 ybridization: atomic orbitals (s and p orbitals) combine or hybridize to form new atomic orbitals that can bond or overlap with one another more effectively We are concerned with only s and p orbitals for majority of 5.12 course

9 A. ybridization Tetravalent carbon forms bonds using four sp 3 -hybridized orbitals. hybridization 2s 2p x 2p y 2p z 4 sp 3 orbitals Trivalent carbon forms bonds using three sp 2 -hybridized orbitals. hybridization 2s 2p x 2p y 2p z 3 sp 2 orbitals 2p z Divalent carbon forms bonds using two sp-hybridized orbitals. hybridization 2s 2p x 2p y 2p z 2 sp orbitals 2p y 2p z ybrid orbitals are more directional (better overlap to form molecular orbitals). ybrid orbitals minimize electron-electron repulsion (think VSEPR). ybridization does not just apply to carbon. This theory can be applied to molecules containing any elements in the periodic table. So, what do these hybridized orbitals look like?

10 1. sp ybridization (Linear) s p 2 sp two sp-orbitals X 180 bond angle minimizes electron repulsion (VSEPR) enhanced e density in bonding regions But, we only used one s- and one p-orbital! There are two more p-orbitals. hybridization 2s 2p x 2p y 2p z 2 sp orbitals 2p y 2p z

11 Complete rbital Picture of an sp ybridized Atom p Y sp X p Z sp 2 sp-orbitals 2 p-orbitals What do these orbitals look like in a molecule? e.g. Acetylene C C! y! z 2 " C " C " C " For simplicity, draw lines connecting p-orbitals to represent!-bonds. hybridization 2s 2p x 2p y 2p z 2 sp orbitals 2p y 2p z

12 2. sp 2 ybridization (Trigonal Planar) s 2 p 3 sp 2 three sp 2 -orbitals Complete rbital Picture p Y 120 X sp 2 X sp 2 enhanced electron density in bonding regions For simplicity, leave out small back lobes. Draw the molecular orbital picture for ethylene (C 2 4 ). hybridization 2s 2p x 2p y 2p z 3 sp 2 orbitals 2p z

13 3. sp 3 ybridization (Tetrahedral) s 3 p 4 sp 3 X 109 No p orbitals left over! e.g. Methane (C 4 ) e.g. Ammonia (N 3 )! C N!!!!!! hybridization 2s 2p x 2p y 2p z 4 sp 3 orbitals ybridized rbitals: All σ-bonds (single bonds) and lone pairs Unhybridized p-rbitals: All π-bonds (double/triple bonds) sp3: 0 π-bond (single bond) sp2: 1 π-bond (double bond 1 σ-bond + 1 π-bond) sp: 2 π-bonds (triple bond 1 σ-bond + 2 π-bonds)

14 4. Assigning ybridization to Atoms in a Molecule (You need to be able to do this!) Count the hybrid atomic orbitals. # of hybrid orbitals = # of!-bonds + # of lone pairs # hybrid orbitals hybridization geometry 4 sp 3 tetrahedral 3 sp 2 trigonal planar 2 sp linear approx. bond angles e.g. What is the hybridization of each non-hydrogen atom in acetonitrile? Draw the bonding orbitals (leaving out the small back lobes). 3 C C N a b

15 4. Assigning ybridization to Atoms in a Molecule (You need to be able to do this!) Count the hybrid atomic orbitals. # of hybrid orbitals = # of!-bonds + # of lone pairs # hybrid orbitals hybridization geometry 4 sp 3 tetrahedral 3 sp 2 trigonal planar 2 sp linear approx. bond angles e.g. What is the hybridization of each non-hydrogen atom in acetonitrile? Draw the bonding orbitals (leaving out the small back lobes). 3 C C N a b C a : sp 3 (4!-bonds)! C!!!

16 4. Assigning ybridization to Atoms in a Molecule (You need to be able to do this!) Count the hybrid atomic orbitals. # of hybrid orbitals = # of!-bonds + # of lone pairs # hybrid orbitals hybridization geometry 4 sp 3 tetrahedral 3 sp 2 trigonal planar 2 sp linear approx. bond angles e.g. What is the hybridization of each non-hydrogen atom in acetonitrile? Draw the bonding orbitals (leaving out the small back lobes). 3 C C N a b C a : sp 3 (4!-bonds) C b : sp (2!-bonds)!! C!! C!

17 4. Assigning ybridization to Atoms in a Molecule (You need to be able to do this!) Count the hybrid atomic orbitals. # of hybrid orbitals = # of!-bonds + # of lone pairs # hybrid orbitals hybridization geometry 4 sp 3 tetrahedral 3 sp 2 trigonal planar 2 sp linear approx. bond angles e.g. What is the hybridization of each non-hydrogen atom in acetonitrile? Draw the bonding orbitals (leaving out the small back lobes). 3 C C N a b C a : sp 3 (4!-bonds) C b : sp (2!-bonds) N: sp (1!-bond, 1 lone pair)!! C!! C! N

18 4. Assigning ybridization to Atoms in a Molecule (You need to be able to do this!) Count the hybrid atomic orbitals. # of hybrid orbitals = # of!-bonds + # of lone pairs # hybrid orbitals hybridization geometry 4 sp 3 tetrahedral 3 sp 2 trigonal planar 2 sp linear approx. bond angles e.g. What is the hybridization of each non-hydrogen atom in acetonitrile? Draw the bonding orbitals (leaving out the small back lobes). 3 C C N a b C a : sp 3 (4!-bonds) C b : sp (2!-bonds) N: sp (1!-bond, 1 lone pair)!! C!! C! " y " z N

19 B. Short-and for Chemists: Easy Communication 1. Line-Angle Formulas: Drawing Complex Molecules Quickly C C C C 3 C C C C C 3 C C C C C C C C C C C C C 3 C 3 C Propane (pretty easy) Testosterone (not so easy) much easier! Just imagine... N N N C 2 Me Me N C Vincristine Ac C 2 Me

20 Rules for Drawing Line-Angle Formulas Bonds are represented by lines (one line = two shared electrons) Do not draw carbon or hydrogen atoms, except at termini (for aesthetics) Assume carbon atoms are at ends of lines and where they meet Assume enough C bonds to give each neutral carbon atom four bonds (an octet) Draw heteroatoms and attached hydrogen atoms (N,,S,P,F,Cl,Br,I, etc.) e.g. isopropanol: C 3 C()C 3 C C or C 3 C C 3 e.g. cyclohexanone C C C C C C

21 2. Using Dashes and Wedges: Molecules Are Not Flat! Tetra-Substituted Carbon Is Tetrahedral (more on this later). C C methane lines dash wedge lines: in the plane of the paper dashes: going back into the paper (away from you) wedges: coming out of the paper (toward you) e.g. Propane e.g. Isomers of 1,2-Cyclohexanediol

22 2. Using Dashes and Wedges: Molecules Are Not Flat! Tetra-Substituted Carbon Is Tetrahedral (more on this later). C C methane lines dash wedge lines: in the plane of the paper dashes: going back into the paper (away from you) wedges: coming out of the paper (toward you) e.g. Propane e.g. Isomers of 1,2-Cyclohexanediol

23 Representing Molecules Lewis/Kekulé Structures: Represent atoms sharing electrons to form bonds Line-Angle Structures: Simplify the drawing of complex molecular structures Dashes and Wedges: Allow chemists to draw molecules in 3-D BUT! These simplified structures do not always represent the electronic nature or reactivity of organic molecules! Electrons are not stationary objects, so it sometimes helps to think about electrons "in motion"... BUT W D WE REPRESENT ELECTRNS IN MTIN?

24 3. Curved Arrow Formalism (Arrow Pushing) Chemists use arrows to represent the motion of electrons within and between molecules. double arrow: 2 electrons moving fishook arrow: 1 electron moving 1. The tail starts at the electrons that are moving (lone pair or bond). 2. The head shows where the electrons end up (lone pair or bond). e.g. electron motion in a substitution reaction (much more detail later) (l.p. to bond) 2 N 3 C Cl 2 N C 3 Cl (bond to l.p.)

25 Sample Problem: Using What You Know Use what you know about Lewis bonding to predict the product of the following reaction. Remember to indicate formal charge. Use curved arrows to show the mechanism (movement of electrons). 3 N B 3? N 8 electrons (1 lone pair) B 6 electrons (wants 2 more) Nitrogen atom is nucleophilic (Lewis basic). nucleophile: electron-rich atom, often negatively charged, with a free lone pair to donate to another atom Boron atom is electrophilic (Lewis acidic). electrophile: electron-poor atom with a low-lying vacant or easily vacated orbital; wants to accept electrons from a nucleophile 3 N B 3 3 N B 3

26 C. Electronegativity and Formal Charge covalent bonds have a degree of ionic character to them polar covalent bonds bonding electrons are attracted more strongly to one atom than the other more electronegative atoms attract bonding electrons more strongly (see table 2.2 page 36) dipole moment! " C! +! + Li! C " oxygen EN = 3.5 carbon EN = 2.5 lithium EN = 1.0 carbon EN = 2.5

27 Calculating Formal Charge formal charge = # valence electrons - # bonding electrons/2 - nonbonding electrons 3 N B 3 3 N B 3 N: = +1 B: = 1 Formal Charge: Formal Charge: N: = 0 B: = 0 calculating formal charge for C,, and N should become second nature in this course

28 Common Formal Charges See table 2.2, page 42 C C C N N valence e # bonds # nonbonding e formal charge

29 Chemical reactions generally involve the movement of electrons between two or more molecules, but the concept of electron motion is also important within a molecule. C. Resonance: Electronic Motion Within a Molecule The reactivity of a molecule is not always explained by one Lewis structure. Molecules can be thought of as hybrids or weighted averages of two or more Lewis structures, each with a different placement of electrons. These structures, called resonance structures, are not real or detectable, but they are a useful conceptual tool for understanding the reactivity of molecules.

30 e.g. ow can you predict where a nucleophile (such as 2 N ) will react with formaldehyde (C 2 )? 2 N? Use resonance to better understand the electronic nature of formaldehyde...! major minor!+ The minor resonance structure suggests that the carbon atom is electron-deficient (electrophilic). The nucleophile will react with the electrophilic carbon atom. 2 N!+ 2 N

31 1. Rules for Drawing Resonance Structures 1. nly electrons move! Nuclei and the sigma- (single bond-) framework are unchanged (Resonance occurs in the pi-system: conjugated lone pairs and pi-bonds). 2. Every resonance structure must be a valid Lewis structure. 3. Keep track of lone pairs and formal charges. 4. Use arrow-pushing formalism to interconvert and identify possible resonance structures. 5. Always use double-headed arrow ( ) in between resonance structures. 6. Lower energy resonance structures contribute most to the overall structure of the molecule. ow do you predict the relative energies of resonance structures?

32 2. Guidelines for Predicting Energies of Resonance Structures (In rder of Importance) i) Filled ctets: Second row elements (C, N,, F) want an octet (filled valence shell of electrons). Because C is the least electronegative, structures in which C has 6 electrons, 3 bonds and a positive charge are possible (not possible with N,, F). ii) Minimum # of Formal Charges & Maximum # of Bonds iii) Negative Charges on Most Electronegative Atoms: > N > C iv) Minimize Charge Separation: Try to keep the formal charges close together. 2 C N C 3 2 C N C 3 2 C N C 3 C 3 C 3 C 3 major follows guidelines minor violates i minor violates i

33 3 C C N 3 C C N 3 C C N A B C 3 C C N 3 C C N D E A: follows all guidelines B: violates ii: 2 formal charges C: violates ii & iii: 2 formal charges; negative charge on C D: violates i & ii: 6 electrons on C; 2 formal charges E: violates i, ii & iv: 6 electrons on C; 2 formal charges; more charge separation than D Relative Energy: A << B < C << D < E Relative Importance: A (major) > B > C > D > E minor

34 3. Delocalization of Charge Delocalization of Charge = Stabilization In general, the more resonance structures there are, the greater the stabilization. Equivalent resonance structures provide more stabilization than inequivalent ones. cyclohexanol pk a 17 + Localized charge on oxygen not very stable Not a very stable conjugate base, so cyclohexanol is not very acidic phenol pk a 10 + Negative charge delocalized by resonance Stabilized conjugate base, so phenol is more acidic than cyclohexanol

35 Delocalization of Charge = Stabilization Delocalization (resonance) stabilizes the negative charge in the conjugate base of phenol (phenoxide). The localized charge in the conjugate base of cyclohexanol is much less stable (no resonance!).

Chemical Bonding II. Molecular Geometry Valence Bond Theory Phys./Chem. Properties Quantum Mechanics Sigma & Pi bonds Hybridization MO theory

Chemical Bonding II. Molecular Geometry Valence Bond Theory Phys./Chem. Properties Quantum Mechanics Sigma & Pi bonds Hybridization MO theory Chemical Bonding II Molecular Geometry Valence Bond Theory Phys./Chem. Properties Quantum Mechanics Sigma & Pi bonds ybridization MO theory 1 Molecular Geometry 3-D arrangement of atoms 2 VSEPR Valence-shell