1. 1 Chemistry 343- Spring 2008 Chapter 1- Carbon Compounds and Chemical Bonds Basic Concepts of Structure Valence= # of bonds an atom can typically form in a neutral compound Carbon = tetravalent; forms 4 bonds Nitrogen = trivalent; forms 3 bonds Oxygen = Divalent; forms 2 bonds, etc. Molecular Formula (e.g.) C 4 H 10, C 6 H 12 O 6 Isomerism= different compounds may have the same formula Isomers: e.g., n-butanol and diethyl ether Structural (constitutional) isomer Shape= geometry around a C atom depends on bonding arrangement 4 single bondstetrahedral 2 single bonds, 1 double bond trigonal planar 1 single 1 triple OR 2 double bondslinear Directly related to hybridization of C-atom orbitals

2. 2 Review of Some Relevant Principles From General Chemistry 1.Lewis dot structures: atoms, molecules, and ions Examples: atoms in first row of periodic table Li Be B C N O F Ne molecules: C H H H H O H H Cl H CC H H H H N H H H H O H H CO O H ions: H Rules:-show only valence e - - show any formal charge present - obey octet rule where appropriate Formal Charge: a method of “e - bookkeeping” NH 3 + H + NH 4 + Atom Formal Charge = group # - net# of e - “belonging to nucleus” (lone pair counts as 2; shared pair contributes 1) e.g., NH 4 + ; N’s group # = 5; net # of e - = 4 formal charge = +1 For NH 3 ; N’s group # = 5; net # of e - = 5 formal charge = 0

3. 3 2. Octet Rule: An atom or ion is especially stable when its valence “shell” has a noble gas configuration (i.e., 8 e - in most cases; 2e - for H, He, or Li + ) some examples C H H H H H H F Li Once the outer shell of any atom is filled, no more e- can be added Larger atoms (3 rd period and higher elements) may accommodate more outer level e- due to their d and f orbitals, e.g., H 2 SO 4, PCl 5 ) Species (atoms, ions, molecules) that do not fulfill the octet rule tend to be much less stable, and therefore much more reactive, e.g., CH 3 +, BF 3, Li. 3. Electronegativity: affinity of an atom for e - ’s For Na ; electronegativity very low For ; electronegativity very high Cl Note extreme opposite positions in periodic table Combination of electronegativity and octet rule suggests that : Nawill lose 1 e - very readily to form Na + will gain 1 e - readily to give Cl - Cl Thus: Na + Na + + Cl - Cl Stability is increased ! 4. Ionic Bonding: bonding based an electrostatic attraction e.g., in solid NaCl: Cl - Na +

4. 4 Each Na+ is equally attracted to 6 neighboring Cl- ions Each Cl- is equally attracted to 6 neighboring Na+ ions Explains physical and chemical properties (BP 1413 o C; aqueous solution, conducts electricity, etc.) 5. Covalent Bonding: give rise to very different structure types Consider 2 Hydrogen atoms: H. H. Electronegativities are identical E-shared to form a covalent bond (shared equally  “pure” covalent) The octet rule is still obeyed for each atom Consider CH 4 (methane) C,H electronegativties almost identical C-H is considered a covalent bond What if electronegativity difference is intermediate? “polar covalent” or polar bond H H not H+H-H+H- C H H H H e.g. H 2 O 2H. + O2H + + O 2- O H H

5. 5 The O-H bond is considered a “polar bond” Consider their respective positions on the periodic table Electronegativity increasing The electronegativity difference causes bonding e- to be “polarized”  + and  - are symbols used to depict partial + or - character ++ ++ Molecule has a “net dipole moment” and is considered polar Other examples: -- -- -- ++ -- ++ A polar molecule Has polar bonds, but dipoles cancel; hence, not a polar molecule

6. 6 All organic molecules have covalent bonds of various types; some have ionic bonds too Sometimes, atoms are bound by 2 or 3 covalent bonds CC H H H H CO H H These have “double” bonds 6. Relevance of bonding type to chemical and physical properties e.g. acetone -- ++ Dipoles can align to increase intermolecular attraction: influences BP, MP, other physical properties Also, if a reactive reagent comes along “seeking” a positive charge, it may attack a  + site and react; Conversely, if a reactive reagent comes along “seeking” a site of e - density, it may attack a  - site and react We will observe this phenomenom many, many times !

7. 7 7. Resonance: A term used to describe a compound that cannot be represented by a single Lewis dot structure e.g. CH 3 NO 2 Experimentally: Both N-O bond lengths are same! Each of the resonance forms above are equally possible CH 3 NO 2 behaves like Thus, CH 3 NO 2 is a “resonance hybrid” of these two forms shown This is NOT an equilibrium!! Symbol used to show that structures are related by resonance Note how one can show the conversion of one resonance form to the other by showing only the movement of e - (curved arrows below) The use of this kind of arrow to show e - movement in a more general sense (= “curved arrow formalism”) will be a critical tool used throughout this course!!!

8. 8 Molecular structures of covalently-bonded compounds Q: How do we know them ? A: X-ray crystallography, spectral analysis 8. Molecular Geometry A.Bond Length: Dependent on: Bond order (single, double, etc.)C=C vs. C-C Size of atoms C-C vs. H-H Electronegativity differenceC-C vs. C-F B. Bond Angles: Generally, angles are arranged so that e - pairs (lone or bonding) are as far apart as possible (VSEPR Theory) If any combination of 4 atoms or lone pairs is associated with the central atom: Shape is tetrahedral Tetrahedral Appears Pyramidal Appears Bent

9. 9 If any combination of 3 atoms or lone pairs is associated with the central atom: Shape is trigonal planar If any combination of 2 atoms or lone pairs is associated with the central atom: Shape is linear All of this is based on… 9. Orbitals and Hybridization A.Background: Electrons possess characteristics of particles and waves - - must be dealt with accordingly Mathematical way of doing this involves Quantum Mechanics Heisenberg Uncertainty Principle: We can only calculate the probability that an electron is in a certain region; hypothesizes bounded areas for electron existence “orbitals”

10. 10 Definitions: Atomic orbitals: mathematical wave functions (  ) describing energy states and corresponding regions in space in which a given e - can exist In organic chemistry terms, an orbital is a region of space where the probability of finding an electron is high Orbitals are also known as energy levels… two different orbitals will correspond to two distinct energy levels having certain discreet values Some Key Principles: Aufbau Principle: lowest energy orbitals filled first Pauli Exclusion Principle: no more than two e - per orbital; spins must be paired (+1/2 and -1/2) Hund’s Rule: for orbitals of equal energy, fill one e - at a time, only add 2 nd e - to an orbital once all have one unpaired e - each. Review quantum numbers: must know 1s, 2s, and 2p orbitals and their shapes (Figure 1.5, pg 21) B. Simplest Case: the 1s orbital. Spherical shape, nucleus at center

11. 11 e - in a H atom occupy the 1s orbital- 90% probability that the e - is within 1Å of nucleus; not uniform; probability goes up as r decreases (r is the distance from the center of nucleus) This e- can be depicted as a “smear” or “cloud” of e- density, or e- probability about the nucleus Wave function=  ; magnitude decreases with increasing r; mathematical sign can be + or - Electron density=  2 ; magnitude decreases more rapidly with increasing r; must always be + Mathematical sign is always + for 1s orbital (not true for any other) 1s orbital C. 2s orbital: larger; also spherical, but  is + beyond 1Å; - inside 1Å there is a spherical ‘node’ at 1Å;  2 =0 at that radius (e- cannot exist within that space) e - can cross such nodes by virtue of their wavelike nature Higher energy than 1s; avg. distance to nucleus is greater 2s orbital

12. 12 The 2p orbitals: 3 of them (p x, p y, p z ) Each has two lobes of equal size but opposite mathematical sign; dumbbell shape Cannot occupy same space mutually perpendicular (2p x, 2p y, 2p z ) Lobes of each are separated by a planar node going through nucleus Higher energy than 1s and 2s These (1s, 2s, and three 2p) are all the orbitals a carbon atom has !!

13. 13 E. Electron Configuration: for atoms of individual elements (Figure 1.6, pg 22) F. Molecular Orbitals (MO’s): mathematical combinations of atomic orbitals (  ’s) Rules: Can add or subtract AO’s mathematically to get MO’s If we combine n orbitals, we must end up with n orbitals Overall energy of new orbitals must be same after combination Simplest Example: H 2 Two 1s Atomic Orbitals (AO’s) combine to form two new MO’s (  and  *) One MO arises from addition of the AO’s; the other from subtraction

14. 14 Addition “bonding” MO (s); Subtraction “antibonding” MO (s*) “Orbital Diagram” (Figure 1.10, pg 25) Note that net overall energy of orbitals is the same as before, but the system is more stable because the 2 previously unpaired e - ’s can both go into a lower energy orbital ! Stability is gained through bond formation A bond results from overlap of two s-orbitals = a “  -bond” Note that Hund’s rule and the Pauli Exclusion principle are still obeyed

15. 15 Contrast this to what would happen if two He atoms came together: Two of the four e - would have to occupy the higher energy  * antibonding orbital bond formation not favorable in this instance He exists in monoatomic form Other Implications: H 2 - can exist in theory, but would be unstable because the third e- would have to go into a  * orbital If enough energy were applied to an H 2 molecule, an e - could be “excited” to the  * MO In most common bonding situations, bonding MO’s are filled, antibonding MO’s are not

16. 16 G. Hybrid Atomic Orbitals Let’s take a closer look at CH 4 … Experimentally, we know that: 1.It has four equvalent C-H bonds 2.It is tetrahedral Neither observation is consistent with the results expected from carbon AO’s Quantum Theory explains this: The most favorable way for carbon to form bonds to four H atoms is to mathematically combine 2s and three 2p orbitals to give four new equivalent hybrid atomic orbitals called sp 3 orbitals (each is “one part s to three parts p”) sp 3 hybrid orbital energy is intermediate between s and p valence e- of carbon are redistributed- one e- per sp 3 orbital each sp 3 orbital can then combine with an H atom’s 1s orbital to form MO’s (only the bonding MO’s are filled) form  -bonds MIX A “Free Carbon Atom” The Carbon atom of CH 4

17. 17 The shape of sp 3 orbitals is dictated by the math of the combination (Figure 1.11) This same situation occurs anytime a C atom is bonded to four neighbors e.g., look at ethane (Figure 1.17) Note that free rotation can occur around the C-C bond (CH 3 s can “twirl around” Applies to more complex molecules as well 17

18. 18 Q: What about a situation where a C atom is only bonded to three neighboring atoms, and shares two e- pairs with one of the neighbors ? A: a difference in hybridization occurs! e.g. simplest case: ethene (a.k.a. ethylene, mol. Formula C 2 H 4 ) Experimentally: Ethene has four equivalent C-H bonds; two to each C Each C has trigonal planar geometry All the atoms in the molecule are in a single plane Molecule is rigid (no out of plane rotations or bending) Neither unhybridized AO’s nor sp 3 hybridization can explain this!

19. 19 Quantum Theory states: The most favorable way for carbon to form bonds with three neighbors is to combine 2s and two of the three 2p orbitals to give three equivalent hybrid atomic orbitals called sp 2 orbitals (each is one part s to two parts p) sp 2 energy is intermediate between s and p (lower Energy than sp 3 ) Shape of sp 2 orbitals is dictated by the math of combination (Figure 1.21) Valence e - redistributed; one per sp 2 ; one in the remaining 2p Each sp 2 orbital can then combine with an H atom’s 1s OR the neighboring carbons sp 2 orbital to form MO’s Only the bonding MO’s are filled  -bonds

20. 20 The unhybridized p-orbitals can also overlap to form a (weaker) bond; the 2 nd bond of the “double bond” a “  -bond” p-orbitals must be parallel and inline with one another to overlap for bond formation to occur; explains rigidity; barrier to rotation around C=C without breaking  -bond ! Fig 1.23, 1.24

21. 21 This restricted rotation leads to a special type of isomerism: cis-trans isomerism “cis-2-butene” “trans-2-butene” “cis-2-butene is a different compound from “trans-2-butene” !! These do not interconvert! Cis and trans isomers are also known as “geometrical isomers”. This is one type of “stereoisomerism” (the connectivity is the same, but the arrangement of atoms in space is somehow different) Ok, but there are also molecules wherein a C atom is bound to only two others! e.g. simplest case: ethyne (a.k.a. “acetylene”)

22. 22 Experimentally: Ethyne has two equivalent C-H bonds; one to each C Each C has linear geometry all the atoms form a line Neither unhybridized AO’s, sp 2, nor sp 3 hybridization can explain this ! Quantum Theory states: The most favorable way for a C atom to form bonds with these two neighbors is to combine the 2s and one of the three 2p orbitals to give two equivalent hybrid atomic orbitals called sp orbitals (each is “one part s to one part p”) sp orbital energy intermediate between s and p (lower E than sp 2 ) Shape of sp orbitals- again dictated by the math (Figure 1.27)

23. 23 Valence e - redistributed- 1 per sp orb.; 1 in each remaining 2p orb. Each sp orbital can then combine with an H atom’s 1s OR the neighboring carbon’s sp orbital to form MO’s Only the bonding MO’s are filled  -bonds The two pairs of unhybridized p-orbitals can also overlap to form two p-bonds; the other two bonds of the “triple bond” Each pair of p-orbitals must again be parallel and inline for  -bond to form Figures 1.28, 1.29

24. 24 Q: Are there any consequences of hybridization besides shape? A: yes! Bond strengths and lengths Double bonds are shorter and stronger than single bonds Orbitals having more “s-character” form single bonds that are stronger and shorter (e.g., order of strength: sp-sp > sp-sp 2 > sp-sp 3 ~ sp 2 -sp 2 > sp 2 -sp 3 > sp 3 -sp 3 ) in general: more “s-character” lower energy How does all this related back to VSEPR theory ? Recall the shapes of ammonia and water; these can be rationalized by hybridization! Consider NH 3 : N bound only to 3 atoms, but there is also a lone pair! Molecule hybridizes to get the e- as far apart as possible! Pyramid shape, but if lone pair is “counted”, it is tetrahedral!! H 2 O does the same thing! It has two lone pairs and two bonds “bent” shape, but if you count lone pairs; it is ~ tetrahedral!!

25. 25 This way of thinking applies to other molecules too. Contrast NH 3 with BF 3 : (NH 3 has lone pair, BF 3 does not; therefore NH 3 has pyramidal shape; BF 3 is flat (trigonal planar) Another example: consider CO 2 (only two atoms around the central carbon, but each forms a double bond with the carbon, hence the shape is linear) H. Structural Formulas of Organic Chemistry Several different ways to draw organic molecules; Figure 1.36 Another common practice: lone pairs on O, N or other atoms are sometimes omitted, they’re just assumed to be there.

26. 26 Condensed and bond-line (“shorthand”) presentations are especially important; structures are complex!! Also, need to distinuish many isomers Some structures are cyclic; easiest to draw as bond-line structures Cyclohexane Benzene In condensed formulas, it’s easier to draw structures with C’s and H’s absent, known as “implicit” carbons and hydrogens Also, will need a way to represent the three-dimensionality (3-D) of these structures; use “wedges and dashes” (Figure 1.37) - (wedges) Depict atoms coming out of the plane of the paper - (dashes) Depicts atoms going into the plane of the paper It is ESSENTIAL to learn to visualize organic structures in 3D !!