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Organic Chemistry Midterm 1 Study Guide

by: Colin

Organic Chemistry Midterm 1 Study Guide Chem241

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Chem 241 Exam 1 Friday January 29, 2016 at 9:00 AM in Randell 326 (during lecture) You will have the entire class period for the exam. Coverage: Chapters 1, 2, and 3 Lectures 1-8 100 poin...
Organic Chemistry
Ms. Susan Rutkowsky
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This 33 page Study Guide was uploaded by Colin on Sunday March 6, 2016. The Study Guide belongs to Chem241 at Drexel University taught by Ms. Susan Rutkowsky in Winter 2016. Since its upload, it has received 68 views. For similar materials see Organic Chemistry in Chemistry at Drexel University.

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Date Created: 03/06/16
Chapter 1: Introduction and Review  Definition of Organic Chemistry o Chemistry of carbon compounds  Atoms o Number of protons determines identity of element o Atomic Number (Z) Number of protons in an element  o Mass Number  (A) Number of protons and neutrons in an element o Isotope  Atoms with the same atomic number but different atomic mass  Due to different number of neutrons o Written as:  AZX  Electronic Structure o Electrons  Outside the nucleus  In orbitals  Wave properties  Density is the probability of finding an electron in a particular part of an orbital  Orbitals are grouped into shells at different distances away from the nucleus  Dense positively charged nucleus surrounded by sea of electrons  Density highest at nucleus and exponentially decreases as distance increases o Shells  (n)  Shell 1  Holds 2 electrons  Shell 2  Three P orbitals  One for each axis o Configurations of Shells  Assigned in shells (n)  Start with lowest energy available building up to higher energies  Aufbau Principle  Must fill the lowest energy orbitals first  Hund's Rule  When there are two or more orbitals of the same energy, one electron will go into each orbital before pairing up in the same orbital  Energy Level diagram  Figure 1  Shows relative energies of orbitals  Shows order of how electrons fill orbitals  Configuration tells arrangement of electrons in orbitals  Completely filled shells  Are stable  Electrons called core electrons  Can be abbreviated with noble gasses  Don’t contribute to chemical behavior  Electrons past core called valence electrons  Atoms on the outermost shell o Lewis Dot Structure  Simple way to represent electrons in valence shells  Octet Rule  8 Electrons  4 Pairs of electrons  Represents completely filled stable configuration  (H and He only have 2)  Elements gain, lose, and share electrons with other elements in order to complete their octets  Elements in the 3rd or lower rows can have more than 8 electrons due to "d" orbitals  Chemical Bond o Net attractive force between two atoms o Occurs due to interactions with their valence electrons o Ionic Bond  Charge-to-charge attraction between ions  One or more valence electrons are transferred from one atom to another  Atoms with opposite charges attract  Electrons are transferred  Figure 2 o Covalent Bond  Sharing of valence electrons  Common in inorganic compounds  Relatively uncommon in organic compounds  Want to complete octet  Nonpolar covalent  When electrons are shared evenly  Equal sharing  Polar covalent  When electrons are not shared evenly  Un-equal sharing  Figure 3  Lewis Dot Structure o Shows  Connectivity of atoms  Distribution of electrons Bonding in covalently bounded compounds   Bond Pairs  Single bonds  One pair of electrons shared between two atoms  Lone Pairs  Non-bonding pairs  A pair of electrons not shared between two atoms  Double Bonds  Two -lectron pairs shared between two atoms  4 e  Triple Bonds  Three electron pairs shared between two atoms -  6 e  Writing Lewis Dot Structure  Decide on central atom The lease electronegative atom is usually  central  C, N, P and S are often central  Determine total number of valence electrons in the molecule or ion  For a neutral molecule, add the valence electrons from each atom  For a cation, subtract the ion's charge from the sum of valence electrons  For an anion, add the ion's charge to the sum of valence electrons  Place single bonds between the central atom and each terminal atom  Place remaining electron pairs on terminal atoms so they each have an octet  If any electrons are left over, assign them to the central atom  Count the electrons on the central atom  If it has fewer than 8 electrons, move 1 or more lone pairs from the terminal atoms to form multiple bonds with the central atom to satisfy the octet rule  H can never form double or triple bonds (neither can halides)  Double or triple bonds often involve C, N, O, or S  Figure 4 o Nonbonding electrons  Also called lone pairs  Valence shell electrons that are not shared between atoms  Dipole Moment o When electrons are more heavil- attracted to one side of a bond o Give positive partial charge  ) on one side and negative partial charge (  on the other side o Symbolized with arrow with positive sign at the positive end and arrow head at the negative end  Arrow then has extra electron  Plus then is missing an electron o Dipole moment    Measure of bond polarity  Amount of charge separation multiplied by bond length  Charge separation shown by electrostatic potential map (EPM)  Red indicates partially negative region  Blue indicates partially positive region  Electronegativity o Used to predict whether a bond will be polar o Greater change in electronegative energy means greater bond polarity o Figure 5 o Formal Charge  Used to keep track of electrons  May or may not correspond to actual charges in the molecules 1  = [ Group Number] - [Non-bonding electrons ] - / [ 2hared Electrons] o Atoms in molecules usually have charge o Figure 6: Common Bonding Patterns  Resonance o When the Lewis structures can be interconverted by moving electrons only o True structure is a hybrid of the two o Delocalization  When two atoms each take half of a negative or positive charge  Due to two resonance structures  Charge stays between two atoms  Stabilizes the ion  Has bond order of 1.5 o Major  All atoms have complete octets  More stable  Lower energy  No charge separation  Electrons can flow through easier o Minor  Not all octets  Less stable  Higher energy  Charge separation o When two resonance obey octet rule:  Major contributor is the one with the negative charge on the most electronegative atom  Electronegativity increases up and to the right of the periodic table o Non-equivalent resonance  Opposite charges should be on adjacent atoms  More stable is the one with the smallest separation of oppositely charged atoms  Structural Formulas o Full  Show all bonds and atoms in order  Figure 7 o Condensed  Written without showing individual bonds  Atoms bonded to central atoms are listed after the central atom  Only double and triple bonds shown  Figure 8 o Line-Angle  Also called skeletal structure  Also called stick figure  Bonds represented by lines  Carbons are present where a line begins, ends or where two meet  Hydrogens attached to carbons are not shown  Hydrogens attached to other atoms are shown  Nitrogen, oxygen, and halides must be shown  Atoms other than carbon must be shown  Double and triple bonds shown  Often used for cyclic compounds  Figure 9 o Lewis Dot Structure  Dashed wedge represents bond going away from you  Solid wedge represents bond coming towards you  Empirical Formulas o Experimentally measured mass percent composition and molar mass of a compound used to determine empirical and molecular formulas o Given percent composition for each element o Assume 100 grams of the element o Empirical Formula Steps:  Convert grams of each element to moles  The percent is the number of grams of the element  Divide this by the molar mass  Divide by the smallest moles to get ratio for empirical formula  Greater number of moles of a element divided by the lesser amount  Ex. 5mol H/2 mol B is B H 2 5  Gives empirical formula  Molecular formula may be a multiple of the empirical formula  Divide the molar mass of the larger element by the molar mass of the compound  Multiply that number by the Empirical formula for molecular formula  Ex. 2/1 so -> B H4 10  Arrhenius Acids and Bases o Definition 1:  Substances that dissociate in water to give [H O ] ions 3  Stronger acids dissociate to a greater degree than weaker acids  Arrhenius bases are substances that dissociate in water to give [OH ]  Stronger bases dissociate more than weaker bases  Strong acids and bases are 100% dissociated  Acidity or basicity is measured by the [H O ]3 which is - related to [OH ] by the water ion-product constant  K w [H O 3 [OH ] -  Neutral solutions: + - -7  [H3O ] = [OH ] = 1.0 * 10 M  Acidic solutions: + -7 - -7  [H3O ] > 1.0 * 10 M and [OH ] < 1.0 * 10 M  Basic solutions: + -7 - -7  [H3O ] < 1.0 * 10 M and [OH ] > 1.0 * 10 M  Measure of acidity and basicity is on a logarithmic scale (pH scale)  pH = -log [H O ] + 3 o Definition 2:  Br nsted-Lowry Acids and Bases  NH l3mits Arrhenius definition  Acids donate a proton  Bases accept a proton  Conjugate acid-base pairs  When a base accepts a proton, it becomes an acid  Acid turns to conjugate base  When an acid donates a proton, it becomes a base  Base turns to conjugate acid  Acid-dissociation constant  (Ka)  When an acid ionizes in water, the constant is used to report the degree of ionization  Strength of the acid  [K a = [A ]*[H 3 ] / [HA]  Acid Strength  Stronger acids have [K ] gaeater than 1  Weaker acids have [K ] lass than 1  Most organic acids are weak  [K ] < 10 -4 a  Acid-dissociation constants are often expressed as [pK a  [pK a = -log 10 [Ka]  Base Strength  Measured like that of an acid  Base-dissociation constant  [K b = [HA]*[OH ] / [A ] -  [pK b = -log 10 [Kb]  For conjugate pairs -14  [K a[K b = [K ]w= 10  [pK a + [pK ]b= -log 10 [10 ] = 14  The stronger the acid, the weaker the conjugate base  The weaker the acid, the stronger the conjugate base  Spontaneous acid-base reactions proceed from stronger to weaker  Structural Effects on Acidity  Electronegativity  More electronegative element bears negative charge easier  Making more stable conjugate base  Weaker base  Stronger acid  Proton is attached to atoms of similar size  Increases top right of periodic table  Basicity and acidity are opposite. Acidity follows EN trends  Size  Rows on periodic table are similar in size  Proton is attached to atoms very different in size  As size increases, the H becomes looser held  Bond becomes easier to break  Larger size also stabilizes the anion  Resonance stabilization of conjugate base  Delocalization of the negative charge on the conjugate base will stabilize the anion  Substance is a stronger acid  The more a molecule can delocalize the charge, the more acidic it will be  More resonance structures usually means greater stabilization  Induction effects  Electron-withdrawing atoms and groups can stabilize the conjugate base  Result in stronger acid  Magnitude of inductive effect depends on nature of the group  How close the group is to the site of the negative charge  Multiple groups allow for further stabilization o Definition 3:  Lewis Acids and Bases  Lewis acids accept electron pairs  Electrophile (accept)  Lewis bases donate electron pairs  Nucleophile (donate)  Doesn’t only apply to reactions with proton transfer  Curved arrow formalism to show flow of electrons  When forming a bond  Nucleophile attacks electrophile  Arrow goes from negative to positive  Figure 10  When breaking a bond  More electronegative atom receives the electrons  Figure 10  Lewis structures and VSPER theory give information on atom connectivity and molecular shape Chapter 2: Structure and Properties of Organic Molecules  VSPER Theory o Valence shell electron pair repulsion theory o Most important factor in determine geometry of a molecule is repulsion between electron pairs o Molecules adopt shape that minimizes electron pair repulsion o Used to explain molecular shapes  Molecules adopt shape that minimizes electron-pair repulsion  Molecular orbital Theory o MO o Describe bonding at the orbital level  Waves o Electrons have wave properties o Standing wave vibrates in a fixed location (3D)  Electron in atomic orbital (AO)  Plucked guitar string o Wave Function    Mathematical description of size, shape, orientation of orbital o Electron density is wave function squared o "+" and "-" do not denote charge  Shows instantaneous phase of constantly changing wave o Nodal plane splits wave in the middle across nucleolus o Node  Center point  Zero displacement o Out of phase  When wave vibrates in opposite directions  One positive and one negative  Displacement may be upward or downward  Cancel out o In phase Wave vibrates in same direction   Same sign  Add together o Atomic orbitals  Can combine and overlap for complex waves  Can add and subtract AO wave functions to make new orbital wave functions o Linear combination of atomic orbitals  Different atoms produce molecular orbitals (MOs) that produce bonding (or anti-bonding) interactions  On same atom give hybrid atomic orbitals that define geometry of bonds  2 3 sp, sp , sp , ect  Number of orbitals must be conserved  Bonding o Region  Stability of covalent bond results from large amount of electron density in bonding region  Bond length  Inter-nuclear distance between two bonded nuclei  Attraction and repulsions are balanced  Gives minimum energy (strongest bond)  Figure 11 o MO theory states that atomic orbitals combine to form molecular orbitals  Since atomic orbitals represent electrons as a wave, combinations may produce constructive or destructive interference  Constructive  Doubles the height of the wave  Doubles (a)  Less energy than original atomic orbitals  Called bonding molecular orbital     Destructive  Halves height of wave  Halves (a)  More energy than original atomic orbitals  Called anti-bonding molecular orbital   *,*  Most electron density outside nuclei  Nodes between nuclei o Sigma Bonding    Formed by head-to-head overlap of two atomic orbitals  Characteristics  All single bonds are sigma bonds  Have bonding electrons concentrated along axis of bond  Free rotation is allowed along axis of sigma bond  May be formed by overlap of a variety of orbital types  Overlap types  S-S  P-P  S + P  Hybridized orbital overlap  Higher energy to anti-boding MO o Pi Bonding    Formed by side-to-side overlap of two atomic p orbitals  Formed after sigma bonds  Sideways overlap of parallel p orbitals  Most of the electron density is centered above and below axis connecting nuclei  Not as strong as most sigma bonds o Multiple Bonds  Double bond  2 pairs of shared electrons  One sigma bond and One pi bond  Rigid  Cannot rotate  Triple bond  3 pairs of shared electrons  One sigma bond and Two pi bonds  Orbital Hybridization o Result when orbitals in the same atom combine o Lower energy due to electron pairs being farther apart  Farther from the nucleus, lower energy o Requirements:  Electron pair geometry around central atom determines orbital hybridization required  Required for every atom and lone pair around central atom  The number of hybrid atomic orbitals created is equal to the number of pure atomic orbitals used o Orbital Geometry:  2 orbitals: 180°, linear  3 orbitals: 120°, Trigonal planar  4 orbitals: 109.5°, Tetrahedral  Isomerism o Isomers:  Same molecular formula but different arrangement of atoms  Number of isomers increases as the number of carbons increases o Constitutional isomers  Also called structural isomers  Differ in bonding sequence  Same chemical formula but connected in different orders  Different properties o Stereoisomers  Differ only in arrangement of atoms in space  Same molecular formula and sequence of bonded atoms  Only different in 3D arrangement of atoms  Geometric isomers  Also called Cis-trans isomers  Molecules differ in arrangement of atoms on either side of C=C bond  Cis-  Groups are on the same side  Trans-  Groups are on the opposite sides  Optical isomers  Also called enantiomers  Dipole Moments o Bond dipole moments  (  Polarity of an individual bond is measured as its bond dipole moment  Result of differences in electronegativity  () depends on the amount of charge ( ) nd distance (d) between the charges  Measure in debyes (D)  Organic bonds range from 0 to 3.6 D o Molecular Dipole Moments  Dipole moment of the whole molecule  Vector sum of the bond dipole moments  Depends on bond polarity and angles o Effect of Lone Pairs  Contribute to dipole moment  Negative is red and arrow head o Intermolecular forces  Strength of attractions between molecules influence boiling point and melting point  Solubility of compounds also influenced  Classification depends on structure  Dipole-dipole forces  Between Polar molecules  Positive end aligns with negative for attraction  Opposite ends repulse  Mostly oriented positive to negative  High boiling point for strongly polar compounds  London dispersion  Between nonpolar molecules  Temporary dipole moment can induce temporary dipole moment in a nearby molecule  Stronger with large atoms  Larger surface area  Branching lowers boiling point  Decreased surface contact between molecules  Arrange randomly  Hydrogen bonding  Strong dipole-dipole attraction  In all molecules  Organic molecules must have N-H or O-H  Hydrogen from one is strongly attracted to lone pair on oxygen or nitrogen of another  Increases boiling points  Polarity o Solubility  Likes dissolve likes  Polar solutes dissolve in polar solvents  Hydration releases energy  Entropy increases  Nonpolar solutes dissolve in nonpolar solvents  Weak intermolecular attractions of nonpolar substances are overcome by weak attractions of nonpolar solvent  Nonpolar substance dissolves  Polar solutes in nonpolar solvents  Solvent cant break apart intermolecular interactions of solute  Polar solutes will not dissolve in nonpolar solvent  Nonpolar solute in polar solvents  Cannot dissolve as it cannot break hydrogen bonds  Nonpolar substances do not dissolve in water  Molecules with similar intermolecular forces will mix freely  Compounds Classes o Based on functional groups o Hydrocarbons  Composed of only carbon and hydrogen  Alkane: single bond, sp carbons Cycloalkanes: carbon forms a ring  2  Alkene: double bond, sp carbons  Cycloalkenes: double bond in ring  Alkyne: Triple bond, sp carbons  Aromatic: Contains benzene ring o Compounds containing oxygen  Alkyl group represented by R Carbonyl group: C=O   Alcohol: R-OH (Hydroxyl group)  Ether: R-O-R'  Aldehyde: RCHO  Ketone: RCOR'  Carboxylic Acid: RCOOH  Acid Chloride: RCOCl  Ester: RCOOR'  Amide: RCONH 2 o Compounds containing nitrogen  Carbonyl: N  Cyano Group: CH C 3 ≡ N  Amines: RNH , 2NHHR', R N 3  Amides: RCONH , R2ONHR, RCONR 2  Nitrile: RCN Chapter 3: Structure and Stereochemistry of Alkanes  Hydrocarbons o Molecules that are made of only carbon and hydrogen o Alkanes  Single bonds only o Alkenes  Double bond o Alkynes  Triple bond o Aromatics  Rings  Alkanes o CnH 2n+2 o Saturated o Small alkanes have low boiling points, gases o Homologous series  Compounds that differ only in the number of methylene groups  Naming o Isomer: Iso- o Highly branched: Neo- o Find the longest chain of carbons for base of the name o Number the chain beginning with the end that has the nearest substituent (branch) o Name the groups attached to the longest chain as alkyl groups  Give the location of each alkyl group by the number of main-chain carbon atoms to which it is attached o Write the alkyl groups in alphabetical order o When the two longest chains are equal in length, use the chain with the greatest number of substituents  More precise naming o Figure 12: Numbering Main Chain, naming example o Figure 13: Common Alkyl Groups o Figure 14: Degree of Alkyl Substitution o Haloalkanes  Halogen is treated as a substituent  Fluorine: Fluoro-  Chlorine: Chloro-  Bromine: Bromo-  Iodine: Iodo- o Multiple Groups  When two or more of the same substituents are present, use prefixes  Di-  Tri-  Tetra- o Complex substituents  If a branch has a branch, number the carbons from the point of attachment  Name the branch off the branch using a locator number  Parentheses are used around the complex branch name o When substituents are alphabetized, iso- is used as part of the alkyl group naming  The hyphenated prefixes are not  Alkane Physical Properties o Solubility  Do not dissolve in water  Non polar o Density  Less than 1 g/ml o Boiling Points increase with increasing carbons  Decrease with branching  Branches are more compact  Less surface area  Weaker intermolecular forces o Melting Points increase with increasing carbons  Less for odd-number of carbons  Even numbers pack more efficiently into solid structure  Alkane Sources o Obtained from petroleum and petroleum products o Fractional distillation will separate crude oil into mixtures of alkanes  Ranges of boiling points  Reactions o Cracking and Hydrocarbon Cracking  Industrial  After fractional distillation  Less valuable fractions are converted to valuable products  Usually done with heat and a catalyst o Combustion  Usually done with heat o Halogenation  Substitution  Usually done with heat or light  Conformations o Different arrangements of atoms o Caused by rotation about single bond o Conformer  Particle conformation o Pure conformations cannot be isolated in most cases  Molecules are constantly being rotated through all possible conformations o Figure 15 o Eclipsed  C-H bonds of one carbon are directly aligned with C-H bonds on adjacent carbon  60 otation from eclipsed  Electron-electron repulsion between bonds increases energy  Least stable  Higher potential energy o Staggered  C-H bonds on one carbon bisect H-C-H bond angle on adjacent carbon  60 otation from eclipsed  More stable  Lower energy  Lower Potential energy o Dihedral angle  Angle that separates a bond on one atom from the bond on an adjacent atom  Also called torsional angle  Figure 16  Totally eclipsed  When angle is 0  Higher in energy than other eclipsed conformations  Forces are closer together and electron clouds experience strong repulsion  Figure 17 o 12.6 kJ/mol difference in energy between eclipsed and staggered o The bulkier the chains, the higher energy between stages o Torsional energy  Energy required to twist a bond into a specific conformation o Torsional strain  Resistance to twisting about a bond o Gauche  60 ngle between groups  Higher energy than anti due to steric strain  3.8kJ/mol  Figure 18 o Anti  180 ngle between groups  Figure 19  Drawn in zigzag arrangement  Skeletal structure  Lowest potential energy o Steric Strain  Interference between two bulky groups  Generally when atom is totally eclipsed  Also called steric hindrance  Touching on models o Different conformers have different stabilities  Newman Projections o Best way to judge stability of different conformations of a molecule o Figure 20 o Examine from head on so bonds between carbons cannot be seen  Cycloalkanes o Molecular Formula: C Hn 2n o Contain carbon atom ring o Properties  Non-polar  Relatively inert  Boiling points and melting points depend on molecular weight o Nomenclature  Find the parent cycloalkane  Use prefix "cyclo-"  Prefix immediately precedes name of parent  Name and number substituents  No number is needed to indicate location of single substituent  Begin numbering at one substituent and proceed around the ring  Second substituent should have the lowest number if the same  If different  Assign lower number to substituents alphabetically  If the number of carbons in the ring is greater than or equal to the longest carbon chain, use this naming process o Isomerism  Cis  Like groups on the same side of the ring  Both groups would have to be on either up or down bonds Trans   Like groups on opposite sides of the ring  Rings are restricted  No free rotation  Two distinct faces giving it geometric isomers (cis-trans)  Stereoisomer  Differ only in the way atoms are oriented in space Chapter 3: Structure and Stereochemistry of Alkanes  Stabilities and Structures of Cycloalkanes o 5 and 6 membered rings are most stable  Most common in nature  Bond angle closes to 109.5  o Carbons are sp hybridized o Require bond angle of 109.5  o Angle strain  Also called Baeyer strain  When the carbon has an angle other than 109.5   There will not be optimum overlap  Bond angles are compressed to a different angle instead of regular angle  Leads to non-linear overlap of the sp orbitals and bent bonds  No longer tetrahedral o Torsional strain  When all bonds are eclipsed  All C-C bonds are eclipsed  Contributes to total ring strain o Ring strains  Measured by heats of combustion  Result of two different strain factors (angle and torsional) o Cyclic compounds  Four or more carbons adopt non-planar conformations  Relieve ring strain  Folded or chair conformations  Folded Conformation o Sometimes adopts folded conformation o Also called envelope conformation o Puckered conformation reduces eclipsing of adjacent groups o Lowers torsional strain o Figure 21  Chair Conformation o Sometimes adopts puckered chair conformation o No longer flat o More stable than any other conformation o Eliminates angle strain  All C-C-C angles are 109.5  o Eliminates torsional strain  All hydrogens on adjacent C atoms are staggered o Figure 22 o Half the Carbon atoms pucker up and half pucker down  Solid Black atoms represents up  Non-solid black atoms represent down o Axial Hydrogens  Located above and below the ring  Along perpendicular axis  Vertical lines  Straight up or straight down o Equatorial hydrogens  Located in plane of ring  Around equator  More room than axial  Horizontal lines  Parallel to sides of the chair  Figure 23 o Ring flipping  Conformational change  The up carbons become down and vise versa  Axial and equatorial H atoms interconverted  Axial become equatorial and vise versa  Boat conformation formed during  Destabilized by torsional strain  Hydrogens on the carbon atoms in the plane are eclipsed  Steric stain  Two hydrogens on either side are forced close to each other  Called flag pole hydrogens  Chair form is 7 kcal/mol more stable than boat forms  Figure 24 o Half chair conformation is at highest energy  Footrest planar with side of molecule o Substituents can be in either axial or equatorial position on chair  Larger substituents in equatorial due to more room  More stable o Drawing:  Pick a ring carbon and classify it as up or down  Draw bonds  Each C has one axial and one equatorial bond  Add substituents  Place larger one equatorial  Label conformer 1  Ring flip  Convert up Cs to down and vise versa  Add substituents  Ring flip changes axial bonds to equatorial and vise versa  Label conformer 2 o Different conformers have different stabilities o Larger axial substituents create unfavorable interactions  Destabilizing conformer o Larger the substituent on a six membered ring  Higher percentage of conformer containing equatorial substituent at equilibrium o With bulky substituents  Less crowding in equatorial position  Makes ring more stable regardless of other groups o When two conformers have one axial and one equatorial group  Same energy  Cis isomer o When two conformers have two axial or two equatorial groups  Trans isomer  Di-equatorial conformation is more stable  Neither group occupies the hindered axial position  Lower energy due to less electron repulsion  Bicyclic Molecules o Two or more rings can be joined into bicyclic molecules o Also called polycyclic systems o Fused rings  Share two adjacent carbons atoms and the bond between them o Bridged rings  Share two nonadjacent carbon atoms and one or more carbon atoms (the bridge) between them o Spirocyclic compounds  Occur when two rings share only one carbon atom  Rare o Figure 25 o Naming  Count the total number of carbons in the entire molecule  This is the parent name  Count the number of carbons between the bridge heads  Place brackets in descending order (ex. [2.2.1])  Place the word "bicyclo-" at the beginning of the name  Ex. "bicyclo[x.y.z]alkane"  X is the number of carbons in the longest path  Y is the number of carbons in the second longest path  Z is the number of carbons in the third longest path  Figure 26 o Naming for Substituent positioning on a ring 1. Start with one of the bridgehead carbons and number it 1 st 2. Number the longest chain of carbons from the 1 bridge head to the 2 nd nd 3. Number the 2 bridgehead carbon then continue numbering along the next longest stain of carbons 4. Once you’ve reached the 1 bridgehead carbon, pass over it and number along the shortest chain of carbons 1. Figure 27 Chapter 5: Stereochemistry  Stereochemistry o Definition:  The study of the three-dimensional structure of molecules o Minor differences in 3D structures can result in vastly different properties o Isomers are different compounds with the same molecular formula o Two major classes: constitutional and stereoisomers o Constitutional Isomers  Also called structural isomers  Have different IUPAC names  Same or different functional groups  Different physical properties  Different chemical properties o Stereoisomers  Differ only in the way the atoms are oriented in space  Have identical IUPAC names  Except for different prefixes like cis or trans  Always have the same functional groups o Figure 28 o Chiral  Molecule or object that is not superimposable on its mirror image  Mirror image is not identical  Ex. hands o Achiral  Molecule or object that is superimposable on its mirror image  Mirror image is identical  Ex. socks  Thalidomide o Enantiomer o Has one chiral center o Exists as two enantiomers  R-thalidomide  Sedative, relieves morning sickness  S-thalidomide  Causes severe birth defects o Can exist as a racemic mixture o Prescribed in 1950s and 1960s to combat morning sickness in women o Thousands of babies born with missing or abnormal arms and legs o Prevents the formation of new blood vessels o Banned in 1962 o Ban lifted in 1985  Used to treat leprosy  Experimentally to treat various cancers Figure 1 Figure 2 Figure 3 Figure 4 Figure 5 Figure 6 Figure 7 Figure 8 Figure 7 Figure 9 More Acidic:  More EN o Increase to top right  Size o Increase Bottom Left  Delocalization  Electron withdrawing atoms Figure 10 Figure 11 Alkyl Groups: Methylene -CH -2 Methyl -CH 3 Ethyl -CH 2H 3 Propyl -CH2CH 2H 3 Butyl -CH 2H 2H C2 3 Isopropyl -CH C3 CH 3 Isobutyl -CH 3HCH 2 CH 3 Sec-butyl -CH C3 CH2 CH 3 CH 3 Tert-butyl -CH 3H CH 3 Branching Groups: Carbonyl Group -C=O Carbonyl -N Cyano Group -CH 3 ≡ N Naming: 1-Methane 2-Ethane 3-Propane 4-Butane 5-Pentane 6-Hexane 7-Heptane 8-Octane 9-Nonane 10-Decane Figure 12 Figure 13 Figure 14 Figure 15 Figure 16 Figure 17 Figure 18 Figure 19 Figure 20 Figure 21 Figure 22 Figure 23 Figure 24 Figure 25 Figure 26 Figure 27 Figure 28


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