CHAPTERS 1-3 STUDY GUIDE Aliphatic—straight/branched chain organic compounds Cyclic—organic compounds in ring form Heterocyclic—cyclic compounds with an atom other than carbon in the ring Alkanes—organic molecules with all single bonds, general formula CnH2n+2

Alkenes—organic molecules with a double bonds, general formula CnH2n Alkynes—organic molecules with a triple bonds, general formula CnH2n-2 Hydrocarbons—organic molecules with hydrogen and carbon only Isomers Constitutional Isomers Ionic/Covalent Electronegativity Resonance Lewis Acid/base Aromatics Olefin Saturated/Unsaturated Functional Group Polarity Dipole Moment Substitution/Addition/Elimination/Rearrangement Heterolysis/Homolysis Carbocations Carbanions Solubility Hybridization Inductive Effect Solvation Protic/Aprotic Solvents Sterics Leveling effect Alkoxide ion

1

HISTORY/CHAPTER 1 History of life—air-sea-soup-nucleics(sugar, pyrimidines, purines)-DNA-cells- life-humans Organic chemicals around for years—Egyptians(alizarin, indigo), Phoenicians(royal purple) Today—ceramics, plastics, synthetics, fibers, paints, food, drink, fuel…. Vitalism—vital force; all organic compounds must be made from other organic compounds Wohler and the synthesis of urea NH4CNO NH2CONH2 Lavoisier, Leibig, Berzelius, Dumas, Cannizaro, Kekule, Couper, Butlerov—qual/quant/ef/mf/structure Organic vs synthetic Organic vs. Inorganic—flammable, mp, bp, solubility, bonding, complexity, reactivity Organics—drugs, medicine, plastics, polymers, ceramics, superconductors, fabrics, cloths, synthetics, foods, drinks, cleansers, detergents, plants/animals/humans…. Polyester/nylons…cheap clothing, easy to make…allergic reactions, waste Adhesives/sealants…makes building cheaper easier…waste, fumes Polymers, plastics…convenience, cost, multiple uses…non biodegradable Herbicides, pesticides, insecticides…helps prevent illness, increases productivity…environmental BHA, BHT, MSG…flavor enhancer, preservative…allergies Saccharin…sugar substitute…cancer causing Antibiotics, Sulfa, Steroids…saves millions of lives…side effects, can be abused CFC’s…cheap, nonflammable, saved lives and money…harmful to the ozone layer STRUCTURE OF ORGANIC COMPOUNDS Organic compounds contain covalent bonds. You need to know Lewis Structures, Formal Charge and Resonance theories to draw correct structures(see handout). Carbon—always 4 bonds Nitrogen—always 3 bonds Oxygen—always 2 bonds Hydrogen, halogens—always 1 bond Carbon-carbon connections can be single, double or triple bonds.

2

Isomers—connectivity/structure EF vs MF vs Structural Formula Van’t Hoff/Le Bel 4 bond arrangement is a tetrahedral Lewis/Kossel—ionic vs covalent bonding, octet rule… Linus Pauling—electronegativity Linus Pauling won 2 nobel prizes. One for chemistry and a peace prize. His chemistry prize was for his groundbreaking work on the bonding of chemicals(esp. orbital overlap, sigma, pi… bonds) Who else won 2? Lewis Structure Exceptions to the octet Rule—all molecules can violate by having less than 8, 3rd period can violate by having more than 8, d-orbitals Formal Charge = valence e – bonds – nonbonding e Carbon with 3 bonds is positive—carbocation Carbon with 3 bonds and lone pair is negative—carbanion Nitrogen with 4 bonds is positive, with 2 bonds(and 2 lone pairs) is negative Oxygen with 3 bonds is positive, with 1 bond(and 3 lone pairs) is negative RESONANCE Curved arrow notation Electrons move from (-) to (+) Can only move electrons(lone pairs or double/triple bonds) Can not break or violate octet rules unless allowed by d-orbitals Use correct reaction arrows to indicate resonance Know Rules pages 16-18

3

Rules for Resonance Structures 1. Only move electrons(lone pairs or double bonds). 2. Electrons can move only from element to bond or vice versa. 3. Move electrons from electron dense(negative…) to electron poor(positive) region. 4. When moving a pair of electrons if a violation of the octet occurs by having more than 8 electrons on a period 1 or 2 element, that is an illegal resonance structure. 5. When a pair of electrons moves into a bond and creates a violation of the octet rule, other electrons must be moved away to stabilize the structure. 6. All elements can violate the octet rule by having less than 8 electrons. 7. Benzene structures have resonances by rotating the 3 double bonds inside. 8. Resonances with all closed shells and no charges are the best. 9. Charge separation lowers the stability. 10. Negative charges on less electronegative atoms and positive charges on more electronegative atoms destabilize the resonance structure. 11. Open octets are the most destabilizing factor. Rules for Orbital Diagrams. 1. Replace each double bonded(sp2) element with a pi-orbital. 2. Replace each triple bonded element with 2 pi-orbitals at 90 degrees to each other. 3. Connect double and triple bonded elements with 2 lines(one from top lobe and one from bottom lobe). 4. Leave all single bonds as is.

4

N

O

N

O

N

N

O

O

N

O

N

O

N

N

O

O

5

O

NN

N

OCH3

Rank = 2

Rank = 2

Rank = 1

O

NN

N

OCH3

O

NN

N

OCH3

O

NN

N

OCH3

O

NN

N

OCH3

O

NN

N

OCH3

O

NN

N

OCH3

Rank = 2

O

NN

N

OCH3

6

MOLECULAR/ATOMIC ORBIT THEORY MO theory is helpful for understanding the physics behind bonding of organic molecules and some of their reactivity. Don’t get bogged down in theory though. You can know everything you need to know about organic chemistry without understanding orbital theory. Quantum mechanics—Schrodinger—wave mechanics, wave equations, wave functions The graph of the solutions for Schrodinger’s wave functions(squared) lead to a region of space where electrons have the highest probability of being found in relation to the nucleus(zero, zero, zero). This plot gives us the orbital shapes for s, p, d, f. Remember orbital represent a 3-d graph of space where it is likely to find electrons. There is NOT a physical shell around the electrons. Recall, the Aufbau principle along with Hund’s Rule and Pauli’s Principle allow us to create a very simple easily understood model of the electron density around the nucleus. As two atoms approach each other, atomic orbitals of each combine to form molecular orbitals. # of atomic orbitals = # of molecular orbitals. Bonding molecular orbits are always lower than atomic orbitals(in Energy). They represent a positive interference, when the two atomic orbitals enter together in phase and reinforce each other, representing a large probability that electrons will be found in that region of space between the two nuclei. Electrons found here increase or stabilize the newly formed covalent bond. Nonbonding molecular orbitals are always higher than atomic orbitals(in Energy). They represent a negative interference, when the two atomic orbitals enter together out of phase and cancel each other, representing a small probability that electrons will be found in that region of space between the two nuclei. Electrons found here destabilize the bond strength of the new covalent bond. HOMO—highest occupied molecular orbital LUMO—lowest unoccupied molecular orbital Hybridization of orbitals—In organic chemistry orbitals are often hybridized. For ex, singly bonded carbons exhibit sp3 hybridization. The 2s and the three 2p orbitals mix to form 4 new(s + p + p + p) sp3 orbitals that exhibit characteristics of both. These new orbitals are 25% s and 75% p in character. Tetrahedral geometry. Angles = 109.5 o

7

Double-bonded carbons exhibit sp2 hybridization. The 2s and two of the three 2p orbitals mix to form 3 new(s + p + p) sp2 orbitals that exhibit characteristics of both. These new orbitals are 33% s and 67% p in character. This leaves one p orbital unhybridized. This unhybridized p orbital make the “double” bond. Trigonal planar geometry. Angles = 120 o Triple-bonded carbons exhibit sp hybridization. The 2s and one of the three 2p orbitals mix to form 2 new(s + p) sp orbitals that exhibit characteristics of both. These new orbitals are 50% s and 50% p in character. This leaves two p orbital unhybridized. These unhybridized p orbitals make the “triple” bond. Linear geometry. Angles = 180o s-orbitals(and sp, sp2 and sp3) make single “sigma” bonds. p-orbitals make double/triple “pi” bonds. A double bond consists of a sigma and a pi bond. A triple bond consists of a sigma and 2 pi bonds. The more s in a bond the shorter the bond. I.E. the shortest bonds are sp, the longest are sp3. Or in other words, triple bonds are shortest, single bonds are the longest and double bonds are in between in size. sp3 orbitals make sigma(single) bonds. They are the longest and have a tetrahedral geometry. sp2 orbitals make sigma(single) bonds. They are middle in length and have a trigonal planar geometry. sp orbitals make sigma(single) bonds. They are the shortest and have a linear geometry. p orbitals make pi(double/triple) bonds. The double bonds are trigonal planar, the triple are linear. Single bonds can rotate. Double bonds can not rotate. Double bonds can exist as cis/trans isomers. Cis isomers have the groups on same side of double bond. Trans have them on opposite sides. Know summary 36-38. VSEPR—tetrahedral geometry includes trigonal pyramidal(ammonia) and angular(water) Structural Formulas—dash, condensed, bond line

8

Molecular Orbital Theory

H C

H

H

C

H

H

C

H

H

HC C

H

H

H

H

π

σ

π

H

O

H

H

H

H

H

H

H

H

C4H10 CH3CH2CH2CH3 H C

H

H

C C

H

H

H

H

C

H

H

H

1

2

3

4

MF Condensed Dash Bond Line

O O

Br

9

(CH3)2CHCHBrCH(CH3)CHCH2H3C

CHCH

CHCH

CH2

CH3

Br

CH3

H3CCH

CHCH

CH

CH2

CH3

Br

CH3

CH3CH2CH(OH)CHCHCH2CCHH3C

H2C

CH

HC

CH

H2C

C

CH

OH

H3C

H2C

CH

HC

CH

H2C

C

OH

CH

Br OH

H3CCH

C

HC

CH

H2C

CH3

H3C CH3

C

HC

CH2

Br OH

(CH3)2CHC(CH3)2CHCHCH2CBrCHCH2OH

10

CH2

C C C

O

OHCH

H2C

HC

HCCH

CH

C

HC C

CH 2

CHC

CH

BrO

sp

sp2

sp3

sp3

sp2

BrO

CH2

C C C

O

OHCH

H2C

CH

CH2H2C C

O

OHCH2C

CH

CH2H2C C

O

OHCH2C

11

EXTRA PRACTICE: Draw at least 2 resonance structures for each and indicate the best resonance structures of the ones you drew.

O

O

N

HO

O

N

O

12

O

H3C

H2C

CH2

CH HC

H2C

CH3

CH3

CH3

H2C

HC

CH2

OCH2

H2C

CH3

H3CCH

CH

CCH2

CHC

H2C

CH3

CH3

H3C CH3

CH3 CH3

13

O O NO

H

O

14

O

O

O

.

O

15

CHAPTER 2 methane

ethane

propane

butane

pentane

hexane

heptane

octane

nonane

decane

CH 4

16

cyclopropane

cyclobutane

cyclopentane

cyclohexane

cycloheptane

cyclooctane

Benzene—Know the Kekule Structure

According to first 2 pictures there should be alternating double and single bonds. But in reality all 6 bonds are of the same length(1.39A). This can be explained by resonance and Molecular Orbital Theory.

What is meant by delocalization? The electrons are dispersed between multiple bonds. In benzene the electrons are delocalized through the 6 pi orbitals of the 6 carbons in the ring structure. These electrons are able to move freely from carbon to carbon. Know polarity/electronegativity. Be able to identify polar bonds and indicate dipole moments and/or partial charges. Polar covalent bonds/dipole moments are responsible for all reactions/mechanism

17

O C O

P o l a r b o n d s , n o n e t d i p o l e P o l a r b o n d s , n e t d i p o l e N o p o l a r b o n d s o r d i p o l e

C lC

C lC l

C l

B r

B r

H 2 O

N H 3

HC

C lC l

C l

HC

C lC l

H

HC

HC l

H

C l C l

C H 4

C 5 H 1 2

C 8 H 1 8

FUNCTIONAL GROUPS Alkyl functional groups—methyl, ethyl, propyl, butyl…isopropyl

1 1

2

1

2

3

1

2

3

4

1

32

1

2

4

3

1

2 3

4

12

4

3

1

2

3

4

5

METHYL ETHYL

PROPYLBUTYL

tert-Butyl ort-butyl

neopentyl

ISOPROPYL

ISOBUTYL sec-Butyl

vinyl

allyl

18

Phenyl

Ph- φ− C6H5

Benzyl

H2C H2C

C6H5CH2- Bn-

Alcohol/halides—F, Cl, Br, I, OH Need to know difference between 1o, 2o and 3o alkyl halide or alcohols.

OH

Br Cl

HO F Cl1o

2o 30

3o2o 2o

Ethers—COC

O

O

19

Amines—NR3, if R= 3 H’s, it is ammonia; if R = 2H’s, it is a primary amine; if R = 1 H, it is a secondary amine; if R = 0 H’s, it is a tertiary amine. NH4

+ Ammonium ion

NH3 Ammonia

(H3C)2N N(CH2CH2CH2CH3)2

NHCH2CH3

NHCH3

NH2

3o

1o

2o

2o

3o

Carbonyl—C double bonded to an Oxygen Aldehyde—carbonyl with H attached directly to carbon Ketone—carbonyl with two alkyl groups attached to carbon Carboxylic acid—carbonyl with OH attached to carbon Ester—carbonyl with alkoxy group(OR) attached to carbon Amide—carbonyl with amine attached to carbon Nitrile—carbon triple bonded to nitrogen Imine—nitrogen double bonded to carbon

20

H

O

OH

O

O CH 3

O

O

O

aldehyde

acid

este r

este r

O

NH2

O

C

K etone

am ide

CN

N

n i t rile

NHim ine

NH 2

O

1 o am ide

NHCH 3

O

2 o am ide

N(C H3 )2

O

3 o am ide

PHYSICAL PROPERTIES Melting/Boiling points Melting/boiling points are determined by intermolecular forces. Ion- ion, dipole-dipole forces, hydrogen bonding In general, the larger the MW the larger the melting and boiling points. The more the branching the lower the mp/bp. Increasing the branching decreases the Van der Waal Force’s which decreases mp/bp. Increasing MW increases Van der Waal Force’s which increases mp/fp. In general, the higher the polarity the larger the melting point. Hydrogen bonding increases the melting point. Symmetry increases the melting point. 1-4 Carbons—gas; 5-17 Carbons—liquids;18 and above—solids Solubility Hydrophilic/hydrophobic/lipophilicity

21

In the following pictures, identify the indicated functional groups.

N

NHCH3

O

O

OH

OH NH2

BrOH

N(CH2CH3)2Cl O

H

NH2

O

O

NH

NHCH3

O

O

A

BC

D

E

F

G

H

I J

1

2 3

46 5

A B C

D E F

G H I

J In the above drawings identify as primary, secondary and tertiary the indicated groups(alcohols, halogens, amines and amides).

1 2 3

4 5 6

22

CHAPTER 3 Know differences between substitutions/addition/elimination and rearrangement reactions. Substitutions—1 functional group replaces another Addition—double or triple bond replaced by functional group Elimination—functional group replaced by double or triple bond Rearrangement—internal structure of molecule rearranges Homolysis/Heterolysis—most reactions go through heterolysis. Homolysis is used in radical forming reactions. In general, the more polarized the bond the easier it will be for the bond to undergo heterolysis. ACID/BASE Bronsted-Lowry vs. Lewis definitions Conjugate acid/conjugate base In organic chemistry both forms of acid/base chemistry is important. Bronsted-Lowry—many different situations in the upcoming chapters will involve protonation/deprotonation reactions. Lewis—also we will see acid/base reactions that involve electron pairs and open orbitals. One of the main types of Lewis Acids we will see is group three halides. Boron trifluoride, aluminum trichloride…these have open p orbital which acts as electron pair acceptor.

F

B

F F

NH3

Carbocations(carbons with positive charge and open p orbital) are also good Lewis Acids. Carbanions(carbons with negative charge) are good Lewis Bases. Electrophile/Nucleophile pKa, pKb…

23

The larger the pKa the stronger the base will be that is generated from that substance. The lower the pKa means the substance is a stronger acid. Low pKa strong acid, high pKa weak acid. Increasing base strength is due to the instability of the conjugate base of an acid. The more stable the conjugate base is of an acid, the stronger it will be as an acid. The less stable the conjugate base is of an acid, the weaker it will be as an acid. A stable base is a weak base. An unstable base is a strong base. Insoluble acids can be made soluble in water by adding a stronger base to it. In general, acidity increases as you go down a column; i.e., HI is strongest of halogen acids. In general, basicity increases as you go up a column. In general, acidity increases left to right across periodic table and basicity decreases left to right. The more s character of a bond the more stable the conjugate base will be; therefore, it will be a weaker base. Therefore, alkynes have most s-character and are the best acids and the weakest conjugate base. Alkanes have the least s-character and are the worst acids and the strongest conjugate bases. The more stable a base is the weaker it is(and conversely it is a stronger acid). Stability to a base is increased by the inductive effect and by resonance delocalization of negative charge. Electron withdrawing groups close to the negative charge will pull the negative charge towards them through the bonds. Inducing the negative charge to move. The more(or any) resonance hybrids that can be drawn the more stable the base. The easier it is to ionize(deprotonate) an acid the stronger that acid will be and the weaker its conjugate base will be. The more stable the conjugate base the stronger the acid.

24

Solvent effects Acids are much stronger in solvent than in gas phase. Solutions can also help stabilize the conjugate base anion. Protic solvents are hydrogen bondable and very good at stabilizing conjugate base. Aprotic solvents can not hydrogen bond and do not stabilize the conjugate base. Neutral organic bases Alcohols can protonate to form alklyoxonium ions. Ethers can also do this. Carbonyls can also protonate(or esp. act as Lewis Base). Nitrogen compounds can do this as well. Double/Triple bonds have the ability to protonate. They have large pi electron density in double bond that acts as Lewis base. Leveling effect—In water there can not be a stronger base than OH- or a stronger acid than H3O+ If you need a stronger base you must use a non-water solvent. Sodium amide is often used in liquid ammonia. Hexane and ether bases are often used also. MECHANISM NOTATION Curved arrows are used to show the movement of electrons. Arrows are always drawn from nucleophile(electron pair, negative charge) to the electrophile(positive charge, carbocation, lewis acid…).

HB r

H

B r -

B r

25

Label the acid, base, conjugate acid and conjugate base in the reaction below. Does the reaction go forwards, backwards or is it at equilibrium? Why?

OH

NH 3

+ N H 3

OH

NH 2

+ N H 4

FO R W A R D S O R B A C K W A R D S O R E Q U IL IB R IU M

W H Y ?

H3PO4(aq) + H2O(aq) H3O+ (aq) + H2PO4

-2(aq) Circle the better base from the two below and explain why it is better.

AC

F FF

CH H

H

O R

BO R

O NH

H

E x p l a i n

E x p l a i n

F 3C

H 2C

CH 2

H 2C

O H H 3C

H 2C

CH 2

F 2C

O H

O R

E x p la in :

O H O H

O C H 3O

C H 3

O R

E x p la in :

26

In each pair below, circle the compound that is a stronger acid and explain why it is a stronger acid.

F3C

H2C

CH2

H2C

OH H3C

H2C

CH2

F2C

OH

OR

Explain:

OH OH

OCH3O

CH3

OR

Explain:

Due to the inductive effect, this moleculeis a much stronger acid. The anion is stabilized by the electron withdrawing effectof the two fluorines adjacent to the negativecharge. This one is stronger due to resonance

stabilization. The CH3 can not contributeto the resonance hybrid like the carbonyl can

O-

OCH3O

O

OCH3-O

Negative charge can be delocalized to theoxygen of the carbony. The other molecule can not accomplish this.

27

Using the pKa chart predict and draw the product of this Bronsted acid/base reaction. For each hydrogen indicated by an arrow, list the pKa of that hydrogen.

OH

Br

Br

NH2 O

H2C

O

CH3

O

OH + NH3

1 equivalent

1

2

3

4 5

If you use excess NH3, list the hydrogens(using 1, 2, 3, 4, 5) that will deprotonate. If you use excess CH3CH2O

- list the hydrogens(using 1, 2, 3, 4, 5) that will deprotonate.

OH

Br

Br

NH2 O

H2C

O

CH3

O

OH + NH3

1 equivalent

1

2

3

4 5

10.0

4.2

9.2 19

36

OH

Br

Br

NH2 O

H2C

O

CH3

O

O-

+ NH4+

9.4

If you use excess NH3, list the hydrogens(using 1, 2, 3, 4, 5) that will deprotonate. 2 and 4 If you use excess CH3CH2O

- list the hydrogens(using 1, 2, 3, 4, 5) that will deprotonate. 1, 2 and 4

28

Part I: Connectivity Worksheet—Page 12—bottom of page B, D, G, I, J Worksheet Page 4—Draw 3-D orbital pictures for methyl cyanate and methyl isocyanate Page 5—Draw a resonance form for this cation that has all closed shell atoms Page 7 (page 26) b(i) Page 8 (page 27) A (a,c) Page 11 (page 30) a, b Part II: For the molecule below, list all the functional groups. On the atoms identified with an arrow, indicate if the functional groups are primary, secondary or tertiary (1o, 2o, 3o).

29

Fill in the C’s and H’s for the following:

Cl

FBr

OHOH

30

31

PAGE 5

CC

C

N

C

CH

H

H

HH

H

CC

C

N

C

CH

H

H

HH

H

32

33

H3C

CH

CH

H

C

C

H

CH

CH

H2C

CH3

H2C

CH3

CH3

CH3

H2C

CH3

H3C

CH

CH

CH

CH

CH

C

H

H

C

CH3

CH

H3C CH2

CH3

Cl

FCH3Br

H2C

H

C

C

CH

C

C

C

H2C

CH2

CH3

CH3

OHOH CH3

CH3

H2C

CH3

34