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by: Brady Spinka


Brady Spinka
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This 66 page Class Notes was uploaded by Brady Spinka on Monday September 7, 2015. The Class Notes belongs to CH 310M at University of Texas at Austin taught by Staff in Fall. Since its upload, it has received 49 views. For similar materials see /class/181888/ch-310m-university-of-texas-at-austin in Chemistry at University of Texas at Austin.

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Date Created: 09/07/15
Alkynes I Chemistry 3183 lOM Sessler Lecture 31 Note Class on Friday Nov 12th is Cancelled If you Wish you may want to attend the research lecture presented by Dr Ravi Pandey of the Roswell Park Memorial Cancer Institute to be held in ACES Room 2402 This is a small lecture hall and standing room only is to be expected Dr Pandey is an expert in photodynamic therapy a modality involving the use of light and drugs to treat cancer and other diseases Initial Hints for Exam 111 still being made up Exam III will cover Chapters 810 essentially in their entirety However Sections 87 and 1010 will NOT be included All reactions presented in class and in the book are fair game The same is true for the mechanisms Top problem picks 86 88 810816 819822 825 828830 833 836839 842844 846847 850 97910 915916 922 924 927 928 930937 939 940 941943946 10310510710810131018 Alkynes Structure Acetylene a m HCECH The overlap of Zpy bonds and 2192 bonds forms a pair of 7 bonds in alkynes C C distance is very short 121 pm or 1 21 A The 6 system is made up of carbons in sp hybridization Nomenclature l Alkynes Change the of alkane to 1m 2 Number chain so that lst carbon of alkyne gets lowest possible number Trivial systems to learn H2 HCEC C HCECH Propargyl Acetylene 3 Number and substituents in usual way examples 6 9H3 H2 g g 1 H CCHCCx E H Terminal Alkyne 3 1 HCquot 0 CH 5 3 5 H23 H2 4 3 5methyI1hexyne 139439Pentadiyne CH3 CH4 3 2 1 Internal Alk ne 2 H30 5 gCEC CHs y Note Alkenes 6 have priority of 5methyI2hexyne alkynes more examples a 1 CH3 0070 TH 2 3 5 1 42 1 2 2 Hlt32CCHFCHa O CHa H 2 CH3 a 2 V l H CH3 1Elutyne transtHexen dyne 33DimelhyI1butyne notabutyne not ransAhexenZyne no 22 dimathylzavbutyne 042 Ethynylcyclohexane Pr0pynylcyclopentane ov cyclohexylace ylene or pFOPargybyClopemane Cycloalkynes Cyclooctyne Cyclononyne Cyclooctyne is the smallest cycloalkyne isolable It is highly strained with the CC bond angle is 15 5 It polymerizes at room temperature Cyclononyne is strained CC angle of 1600 but has normal alkyne reactivity Physical Properties The properties of alkynes are similar to those of alkenes and alkanes Lower weight alkynes are gases at room temp Those that are liquids are less dense than water 1 gmL They are nonpolar compounds soluble in each other and non polar organic solvents Table 101 Physical Properties of Some LowMolecular Weight Alkynes Melting Boiling Density Point Point at 20 C Name Formula C C gmL Ethyne HCECH 81 i 84 a gas Propyne CHgCECH 102 i 23 a gas lButyne CHgCHzCECH 126 8 a gas QBntyne C1113CECCHS 32 27 0691 lPentync CHSCHZ ZCECH 90 40 0690 lHexyne CH3CHZ3CECH 132 71 0716 lOctyne CHSCHZ5CECH 79 125 0746 lDecjnle CHgCHg7CECH 7 36 174 0766 Heats of formation of Alkynes TABLE 34 Heats of Formation for Some Small Hydrocarbons l Alkanes AH kcalmol Ethane 20 1 Propane 250 Butane 802 Alkenes Ethylene Propylene 1 Butene 2Bu tene All kcallmnl 125 48 01 1 9 015 49 trans Alkyn es Acetylene Propyne 1 Butyne 2 Butyne l AHo kcallmol 545 446 395 347 2283 kJmol 1868 kJmol 1654 1Jmol 1453 kJmol Heats of formation show that alkynes are less stable than alkanes or alkenes Trends indicate that the more 7 bonds in a molecule the higher the heat of formation In addition internal alkynes are more stable by about than terminal alkynes In the case of 2butyne vs lbutyne there is about a 20 kJmol difference Acidity of Alkynes B39 HQH3 X EH CH3 The methyl anion emide BgHLCECH A 3H ZCEC H pk 25 The acetylide ion It is generally difficult to remove a proton from a hydrocarbon Strong bases B can remove the terminal hydrogen of acetylene Water soluble bases eg hydroxide are not strong enough bases to remove a proton from acetylene In fact acetylide ion reacts with water rapidly to produce acetylene and hydroxide anion Alkane and alkene hydrogens have very high pka s Alkali metal hydroxides sodium hydride and lithium diisopropylamide are not strong enough to cause deprotonation Still relative to alkanes terminal acetylenes are reasonably strong acids Why H H L base quot 3H H gt Hm G H H Methane Methide c is hybridized approximately spa therefore the electrons are in an orbital of about 25 5 character base E H CEC H gt E C H BH Acetylene Acelylide Now C is hybridized approximately 5p d the eleclrons are in an orbital Di approximately 50 5 character In an acetylide ion the two nonbonding electrons are accommodated in an Sp orbital Reactions of Alkynes Alkylation of Acetylide Anions Formation of terminal alkynes H g2 g2 i HCC C CEC CH3 HoseNa Hsc Cz 3 H2 H Na Br H2 2 sodium acetylide 1br0mobutane 1 hexyne Because of the ease of which acetylene is converted into a good nucleophile alkylation of the acetylide anion is the most convenient method to prepare terminal alkynes While a great nucleophile sodium acetylide is also a very strong base So that means this reaction is generally only useful for 10 systems with little B branching Otherwise E2 Formation of internal alkynes H2 8N2 CH3 SOdium propynide 1bromoethane 239pentyne QuickTimeT39V39 and a Sorenson Video decompressor are needed to see this picture As a general rule the acetylide ion reacts with 2 carbocations to give products from E2 elimination not SN2 substitution E2 HC cN HHQCBH HCECH Q NaBr E 39 a r acetylene Cyclopentene sodium acetylide bromocyclopentane Acetylides are strong bases and will generally alkylate effectively only primary halides tosylates and mesylates Preparation of Alkynes In addition to alkylation alkynes can be prepared by dihalogenation of an alkene followed by two successive B elimination reactions The two halogens may be on the same carbon geminal or gem dihalides or on adjacent carbons Vicinal or Vicdihalide RCHZCHR Br2 CCI4 Alkene good starting materials to produce alkynes E E 2 Na39NH2 R C C R R CEC R 2NH3 2NaBr H H NH3 vicinal halide 9H3 2 H3c c oNa Br H cH 9H3 I 2 3 RCC R gt RCCR 2H30CIZOH 2KBr r DMSO CH3 geminal halide With a weaker base KOH the dehalogenation can stop at the haloalkene H2CCHX X Cl Br I Allenes Sideproducts of dehalogenation of haloalkenes Na39NH2 RCzcCH3 RCECCH3 RCCCH2 NH3 NaBr39 E3r NH3 H alkyne allene Allenes have two adjacent double bonds The end carbons are Sp2 hybridized While the center is Sp hybridized c5 spzsp H20CCH2 Allene Allenes are usually minor products as they are more unstable than their isomeric alkynes H3CCECCH3 AH0 167 kJ mol HZCC3CH3 12butadiene 2butyne Catalytic Reduction Treatment of an alkyne with hydrogen H2 using a transition metal catalyst Pt Pd or Ni results in the addition of two moles of hydrogen to the alkyne converting it to an alkane The process is two step 1st reduction of the alkyne to an alkene and then alkene to alkane Usually it is not possible to stop the reduction at the alkene H H 2 2 Pd H2 H2 H2 H2 H3C C 39CC C 39CH3 gt HBCC C C 0 CH3 339heXyne 3 atm H2 hexane However by carefully choosing the catalyst it is possible to stop the reaction at the alkene stage Lindlar s catalyst Finely powdered palladium deposited on solid calcium carbonate that has been poisoned with lead salts This modifies the surface of the catalyst rendering it less active Alternative versions of this eXist including the use of barium sulfate and quinoline In any event the idea is to lower the reactivity of the surface such that the intermediate alkene can escape before undergoing further reduction hydogenation The catalyst gives the usual syn addition as when transition metal catalysts are used Thus producing cisalkenes from alkynes H2 H2 PdCaCOsle HBC CQz HZC CHs Hsc C 39CEC C 39CH3 H2 CC 3hexyne Lindlar catalyst H H cis3hexene Chemical Reduction Alkynes can be reduced in liquid ammonia by lithium or sodium metal The metal is oxidized in the process to M and the addition is stereoselective Chemical reduction of alkynes proceeds by antiaddition to give transalkenes H2 H2 H2 H H2 E C 0 CH HSCCC C C CCCH3 2 Na T H30 C 3 H H2 H2 H2 H2 4Octyne trans4octene mechanism resonance stabilized RCECR 39Na gt REZCZR lt gt RECR Na A R R 39 H N H cc NH239 39Na gt CC Na R CC R Ill R H R H alkenyl radical alkenyl anion A H Er C R H r H Cc NHZ 39 R R H H H trans alkene Lecture 3 Review of Lewis Dot Theory and Related Topics Organic Chemistry I 310318M Pre Health Professionals and Chemical Engineering Majors Unique numbers 52560 and 52785 Rules for Lewis Structures All valence electrons are shown Total e39 Sum of all valence electrons on all atoms involved Total number of e39 is sum of valence electrons for each atom Add amp subtract electrons for anionic and cationic charges respectively Determine connectivity either given or by trial and error Complete octet for each atom to fullest extent possible 9 o 9 Zee 921C229 QICZIQ QCQ 0k somewhat wrong wrong Important Considerations for Drawing Lewis Structures H will have a duet 3rd period atoms P amp S can expand their octet to 10 or even 12 electrons Double and triple bonds shown by 2 and 3 electron pairs respectively Species without a complete octet are very reactive H H H NjH H2CIH HIC1H H H 9H He I methyl cation methyl radical H H39 39 39H H H H Complete Lewis Structure Minimize Formal Charges Assign to an atom all its unshared electrons and half of its bonding electrons Compare this number against the number of valence electrons in the neutral unbonded atom Number of valence electrons all unshared 12 of all Formal Charge in the neutral unbonded atom lt electrons shared electrons I 1 9 e e Q H3NH 0C0 H H quot 4e39 for N here 1 7e39 for 0 here 3e for C here 1 5e39 for N atom 6e39 for O atom 46 for C atom 7e39 for 0 here 6e39 for O atom II I p t You Will Often See People Use A Slightly Different but Analogous Way of Calculating Formal Charge 6 Positive charges n 2 is Electra s r electrons 6 Negative Charges P m on 8 POS Wquot Charges Neutral 8 Negative Charges Net neutral electrons B POSlllVe charges 8 Protons 2 1s Electrons 1 Shared Nitromethane H electron 6 Nonbonding 9 Negative charges electrons Net 1 7 Positive charges 7N 215 Electrons ared electrons 6 Negative Charges Net 1 Formal Charge Signi cance and Utility Formal charges obtained are not real unlike ionic charges Sum of formal charges ionic charge of the ionmolecule Formal charges are useful in predicting Polarity in many cases and ionic charge in a few cases Reactivity Most stable geometries Resonance This is the nal aspect of Lewis dot theory we need to discuss before considering simple molecular geometry Note Molecular Geometry is a COMPLICATED topic in organic chemistry and is a subject we ll Visit and revisit often Resonance Some Molecules cannot be described by a single Lewis dot structure Example CO3 Three Lewis dot structures can be written The actual structure is a resonance hybrid of all 3 structures O 5 C 9 j 1 Cl 2 a j CPO m WWW a 9 Picture resonance structures as if looking through a stack of transparencies Resonance Double arrow does NOT imply a reaction Rather it is internal fast movement of electrons Can be more than one double arrow Hybrid resonance structure like a mule from donkey and horse Again picture resonance structures as if looking through a stack of transparencies benzene Chloronitrile e H H H f a a s 39 O H H H 39e o E Rules for Resonance Structures Nuclei never change positions Only electrons in nonbonding con gurations or multiple bonds change If it can resonate it Will Structures With second period atoms eg C N 0 having lled octets are the most important Maximize the number of covalent bonds Minimize the charge separation When formal charges pertain try to put them on the most electronegative for minus formal charge or electropositive for positive formal charge Avoid breaking rules 1 3 to do so Practice Resonance Structures allyl cation 9 a e Q 1s0cyanate anlon ZNCO NECO a G e methyl 1s0cyanate H3CNCo H3CNCO H3CNCO I Best Structure Resonance of Neutral Compounds Note that resonance structures contributors can also be written for compounds for which satisfactory Lewis dot structures can be written These often help explain the reactivity or structure of molecules even though they are not the best Lewis dot structure QZCZQ O CO ch Q MOSt Unimpo am site of high resonance Important Structures reacthity 6e Resonance Structure Practice Resonance 39 H39 HCO 39 quot o e B1carbonate C02 0 3 9CQ e e e e 59 lQ O QQ Q ICI Q 0H 9H gH minor contributor H He H H New Key Concept Protonated Q Formaldehyde H Use of curved a1rows to show H H H e H ow of electron HCOe 9 pairs Theygo H FROM electrons to unimportant no t f 1 tr 1mp0rtant octet for carbon SI es e CC on de c1ency Resonance Explains Reactivity 39 c H H039 HCH 2 H H 39O 39CO 39gt H o O H ll HCH ecN Note that normal CO bond length is 120 A and normal CO bond length is 143 A Contributions from minor resonance structure explain longer bond in protonated form ca 138 A Geometric Structure Isomerism Chapter 1 Introduction In contrast to ionic bonds the formation of covalent bonds imposes a certain geometry on the molecules As soon as more than 2 atoms are involved we must consider the 3D arrangement When GN Lewis proposed dot structures he included Van t Hoff amp Le Bel s ideas and suggested that the pairs of electrons are oriented as far as possible from each other Qualitatively This theory works quite well although more accurate is MO molecular orbital theory we ll discuss next is better Experimentally structures are determined by Xray electron or neutron diffraction analysis or by microwave spectroscopy Examples of structures below plane of H paper Ij 1095o MH above plane of the pa per 109 A Methane H H Hf Q H E H Ethane 109A 1214o 77 u H 11730 134 A H perspective Ethylene H four pairs as far apart as possible H Iioll H 109A 153A H H H 10930 IIOII IIOII I H sawhorse Single bond always 153A H H C 39C C Fr Kekul WH H Double bond ca 134 A Examples of structures 1 80 109A FCC7C H HcCzH 120A Lewis Kekul e 1095o 100A I 096A 39 39 H quot39ll H H WH HeTOH 10730 1045 H e H z H Not linear H Ammonium like methane We will expend much effort on structure we ll discuss it more throughout the course This is enough now to go on to the polarity problem Polarity of Covalent Molecules In molecular ions Without formal charges it is easy to see Polarity o eg 2OH and in symmetrical molecules e g HH it is easy to see nonpolarity When a covalent bond is formed between atoms which differ in their electronegativity a polarization occurs leading to a bond dipole 5 5 I gt example H CI H CI 7 198 Debye Vectors Determine Polarity The vector sum of polarized bonds occurring in a molecule leads to the dipole moment We re not concerned with numbers just effects r 1r 39 I39 X t x Ill42W H H Big Br y2184D y179D yOD Nonbonding Interactions Between Molecules Dipoledipole These are usually weak 13 kcalmole covalent bonds 30110 kcalmol these are important in establishing orientations in liquids and solids Hydrogen Bonding We are alive because of it H can be bound to O F N makes polar socalled hydrogen bonds Such H atoms can interact with nonbonding electrons on other 0 F N atoms 310 kcalmole per Hbond Hydrogen Bonding A Most Important Example water MW 18 bp 100 C hydrogen bonds hydrogen sul de MW 34 bp 62 C no hydrogen bonds Hydrogen Bonds Acceptor vs Donor Molecules may have H bond acceptor or donor or both sites Increases H20 solubility 39V O H2 H30C0CCH3 HBCCCCCH3 H2 H2 H2 H2 H20 solubility 759 L H20 solubility 036g L Van der Waals Forces The weak forces of attraction that exist between nonpolar molecules They arise because motion of e39 in bonds gives rise to tiny induced dipoles which result in slight attractions between molecules They are a strong function of distance and hence of geometry of molecules Van der Waals forces explain why bigger molecules have higher mp bp than little ones Elongated molecules have higher boiling points than spherical ones they pack like sausages amp have a lot of contact H3 H2 H3 Iquot2 H2 Iquot2 ch c CH3 ch c ng H3C C C C CH3 CH3 bp 95 00 279 00 361 00 mp 166 C 4599 C 4297 c Melting point re ects ability to form nice crystals here spherical symmetry helps higher mp Not related to Van der Waals forces Qualitative knowledge of boiling points is expected that of melting points is not Ethers Epoxides and Sul des Lecture 1 Chemistry 318310M Professor Sessler Lecture 34 Dr Ian Wasser Guest Lecturer Today New Material Chapter 11 A Short Lecture Friday 111904 Review Exam 111 Answer Key Monday 112204 New Material nish Ch 11 Wednesday 112404 No Class Happy Thanksgiving Monday 112904 Professor Sessler Returns Ethers Structure and Nomenclature An ether is a functional grop with two carbons bound to an oxygen For example dimethyl ether CH3OCH3 a b H H c 9 c H U I ll JR 11o3 1 H2 H2 0 c Diethyl ether quotetherquot Nomenclature By IUPAC rules ethers are named by selecting the longest carbon chain as the parent alkane and naming the OR group attached to as the alkoxy substituent Common names are derived by naming the alkyl groups bonded to the oxygen in alphabetical order and adding the word ether H2 H2 00 C C V H3C 0 CH3 CY Ethoxyethane 2methyI1propoxypropane 2methyI2propoxycyclohexane Dle ihyl ether tert butoxycyclohexane O f E E OOO O O Diethylene glycol dimethyl ether tetrahydrofuran 14dioxane quotdlglymequot quotetherquot Physical Properties of Ethers Ethers are polar molecules where the oxygen has a partial negative charge and the carbons have partial positive charges This leads to dipole interactions in the liquid although these are weak Ethers cannot act as hydrogen bond donors and are less soluble in water than alcohols However they can accept hydrogen bonds which makes them more soluble in water than hydrocarbons Hydrogen 0 rad bonding I H 5quot 8 H 5 H 0 aquot 4 H 8 8CH H H Due to their weak dipole dipole intercations ethers have lower boiling points than the alcohols of corresponding molecular weight They have boiling points much closer to hydrocarbons of similar molecular we1 ght Orly very Weak dipole dipole interaction H H H H s H 051 quot05 4 H H 5 5 H cs H H H Table 111 Boiling Points and Solubilities in Water of Some Ethers and compare Wlth Alcohols of Comparable Molecular Weight Molecular hp Solubility Structural Formula Name Weight C in Water CI IgCHgOH Ethanol 46 78 In nite CHgOCHg Dilnethyl ether 46 i 24 78 g100 g CI IgI IZCI IgCI IZOI I 1Butanol 74 117 74 g 100 g CHgCHgOCHgCHg Diethyl ether 74 35 80 g100 g HOCHgCHgCHgCHgOH 14 Butanediol 90 280 In nite IiirgcngcngcnzcHZOH 1Pentanol 88 138 213 g100 g CHgOCHZCHZOCHg Ethylene glycol 90 84 In nite dim ethyl ether C13jl lgcl lgcl2ocl 13 Butyl methyl ether 88 71 Slight Preparation of Ethers The most common ether preparation is the Williamson ether synthesis It is a general method for ether synthesis involving the SN2 reaction between an alkyl halide and an alkoxide anion In planning a Williamson ether synthesis it is crucial to select conditions and reactants that maximize nucleophilic substitution and minimizes B el1m1natlon 8N2 Y wan CH3 Y NaBr Br QuickTimeTM and a e Ideo decompressor are needed to see this picture Belimination Y Na39o CH3 gt n NaBr39 CH30H Br Yields of the Williamson ether synthesis are the highest with primary and methyl halides Secondary halides give low yields and tertiary halides fail to react both due to competing E2 B elimination Br NaOCH3 gt NaBr 0 Ether Synthesis Problems Propose 21 Synthesis for Each Ether Product OCH3 Ether Synthesis Problems O39K OCH2CH3 0 CH3CHgBr gt 0 KBr OCH3 O39K OCH3 CH3 gt A KI AcidCatalyzed Dehydration of Alcohols Some of the commercially available ethers are synthesized on industrial scale using high temperatures and strong acids The alcohol group is activated by the acid and by an SN2 attack of a hydroxyl group it is displaced to give an ether Dehydration is a competitive process but it requires higher heat Therefore dehydration can be reduced by controlling experimental conditions H H so H2 H2 2 2 2 4 C C H20 HCCOH gt H3C 0 CH3 3 140 C mech Wm H AH f 2 2 HsC QOH H39OSO3H Hsc C OH H3C C BTI gt p g2 g2 H2 H2 H20 Hsc O CHs gt H3C C O C CH3 HZQF H H K AcidCatalyzed Addition of Alcohols to Alkenes Under the right conditions alcohols can be added into alkenes to give ethers Limits are alkenes that can form stable carbocations and primary and methyl alcohols Example tbutyl methyl ether H 9H3 Y CHgOH gt H3CCID OCH3 CH3 2methoxy2methylpropane mech39 t butyl methyl ether I HOCH3 Y Hd CH H 39739 A 3 UCHs 9CH3 39 O CH3 HOlCH3 H Sample Problem Propose a mechanism for the following transformation H so OCHs CH3OH A OCH CH30H CH3O H 3 If l CH3 H H G p V OCH3 3 CH3 Reactions of Ethers Ethers are resistant to chemical reaction like hydrocarbons They are stable to potassium dichromate or permanganate Very strong bases are unreactive towards ethers and except for 3 alkyl ethers they are not reactive toward weak acids at moderate temperatures For these reasons among others ethers make good solvents However strong acids such as HX can cleave ethers Ether cleavage requires a strong acid and a good nucleophile Therefore concentrated HI and HBr are the acids of choice for this reaction mech I H MBMM Br gt Ho Br 4 r H8H2 HO H9 Br H Oxidation of Ethers Formation of Hydroperoxides One of the hazards of working with lowmolecular weight ethers other than their high ammability is their reactivity With oxygen Anhydrous ethers react with 02 at the C H adjacent to the ether oxygen to give explosive hydroperoxides H HOO OJ 02 gt OJ A hydroperoxide Hydroperoxidation proceeds by a radical chain mechanism If the C H next to the oxygen is secondary rates of peroxidation are increased due to the stability of the 2 radical intermediate Hydroperoxides are highly explosive They can be detected by a simple starch test Where K1 is oxidized to 12 by the hydroperoxide which in turn reacts With the starch to give a deep blue color Hyrdoperoxides are destroyed by treating with a reducing agent such as FeII sulfate in dilute acid Precipitated peroxides of diisopropyl ether are best left to the BOMB SQUAD Ethers as Protecting Groups In many cases Where two functional groups exist in a molecule it is necessary to protect one in order to do chemistry on the other One such example is with an alcohol and an alkyne pKa 25 pKa 1618 Under the conditions needed to form the acetylide ion the alcohol is deprotonated and thus made more nucleophilic It is necessary to protect the alcohol Protecting groups must 1 Easily added to sensitive group 2 Resistant to the reagents used to transform the unprotected group 3 Easily removed to regenerate the original group All these criteria are met for the alkylation of the acetylide ion when the alcohol is protected as a tert butyl ether 1 n 1 NaNH239 HO OH HO O H2804 2 CH3CHZBr CO C OH T The protecting group ether is added with acid stabile under the basic alkylating conditons and removed with weak acid to regenerate the alcohol Another type of ether protecting group is the silyl ether which is removed by acid or F39 uoride ion CH3 OH CI Si CH3 NO Sliig3 CH3 CH3 3 CH3 Noe CH3 u SI F gt NOH F 339 CH3 CH CH3 3 CH3 Chaps 810 Finishing Up Chemistry 3 183 10M Professor Sessler Lecture 33 Dr Ian Wasser Guest Lecturer


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