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CHEM:341 Organic Chemistry I Notes Week 2

by: Raffasarru

CHEM:341 Organic Chemistry I Notes Week 2 CHEM341

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These are the Organic Chemistry I notes for week 2. They build off of my week 1 notes so if you don't have those, they're free on my page. Thanks!
Organic Chemistry
Anna Elizabeth Allen
Class Notes
Organic Chemistry, Chemistry
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This 14 page Class Notes was uploaded by Raffasarru on Saturday August 27, 2016. The Class Notes belongs to CHEM341 at Colorado State University taught by Anna Elizabeth Allen in Fall 2016. Since its upload, it has received 6 views.


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Date Created: 08/27/16
CHEM:341 Organic Chemistry I Week 2 Notes  The Resonance Hybrid o Individual resonance structures are not real and do not fully represent a molecule alone o Resonance structures are not in equilibrium with each other there is no movement of electrons between resonance forms o Resonance hybrid: combination of all possible resonance structures and is accurate representation of molecule  Example: ???????? ????3???? −   Each structure contributes 50% to hybrid o Examples of hybrids using partial charges and partial double bonds   Partial double bonds between ???? and ???? and ????  Partial charges on ???? and ???? 1   Major and Minor Contributors (Resonance Hybrid) o When resonance structures of a molecule are not the same, they will have different contributions to the hybrid structure o Hybrid structure is a weighted average of contributing resonance structures based on their stability (more stable = larger contribution o To determine relative stability of resonance structures, consider the following:  Fewer atoms with formal charges    B has more formal charges than A, so A is more stable  Has fewer formal charges  Atoms with full octets of electrons  More stable if every atom has a full octet of electrons  Major: all atoms have full octet  Charges placed on atoms with appropriate electronegativity (tie-breaker) 2  Put negative charges on more electronegative atoms and positive charges on less electronegative atoms  Major: negative charges on more electronegative, positive on less electronegative o In general, most stable resonance structure has more bonds and few charges  Bonding in Organic Molecules o Understanding molecular structures and reactivity depends on reasonable model to explain bonding in molecules (orbitals) o Electrons occupy orbitals (area of space an electron can be found) o o p orbitals are higher energy because electrons can move farther from the nucleus (negative is farther from positive) o All s orbitals have electron density at the nucleus  Orbitals and Bonding: Molecular Orbital Theory 3 o In molecules, it’s often not sufficient to consider only atomic orbitals, need more sophisticated bonding model to explain structure and reactivity o Atomic orbital: orbitals belonging to single atom  Unhybridized: s and p 3 2  Hybridized: ???????? ,???????? ,???????? o Molecular orbital: orbital that belongs to entire molecule; electrons are no longer localized on a single atom o Electrons can be described as both particles and waves and their movement can be described by the function of a wave o Constructive interference and destructive interference:  o Apply same concept to s orbitals  Overlap of like-phase orbitals creates bonding molecular orbital 4  Majority of electron density is between the two nuclei, which is what keeps them together  Overlap of opposite phase orbitals creates an antibonding molecular orbital   Majority of electron density is on the far sides of the nuclei so they don’t bond  Waves cancel between the two nuclei so there is no electron density at node  When two orbitals overlap, both bonding and antibonding orbitals are formed  Example: Hydrogen   Bond exists 5 # bonding ???? −# anti-bonding ????  Bond order: 2 2−0 o = 1 2  Example: Helium  2−2  Bond order: = 0 2  Bond doesn’t exist  Antibonding orbitals are there to break their corresponding bond orbital o Apply the same concept to p orbitals  When perpendicular, the bonding and antibonding interactions exactly cancel so there is no interception and no molecular orbitals are formed  When parallel, bonding occurs   ???? − bonding: parallel overlap (side-on) of p-orbitals, hold electrons above/below plane of atoms 6   General Structure  o General Guidelines for molecular orbital construction  Number of molecular orbitals produced is always equal to number of orbitals you start with  Need to be able to have same number of electrons overlap of two makes two molecular orbitals  Better overall of orbitals, stronger they interact, closer in energy orbitals are, the stronger they interact (better overlap = stronger bond)  When two orbitals interact, always make a bonding orbital of lower energy and antibonding orbital higher energy  Lower the energy of bonding orbitals, the stronger the bond  Higher energy antibonding orbitals, stronger the bond 7  Orbitals produced for a given molecule can be ordered by energy according to number of nodes present  Antibonding orbitals always have one more node than bonding orbitals  Bond Strength o Stabilization of bonding molecular orbital, relative to the non-bonding atomic orbitals, is related to strength of bond o Making a bond is exothermic: bond is lower energy than two original molecules o Bond dissociation energy: amount of energy it takes to break a bond o Breaking a bond is endothermic o Ways bonds break:  Homolytic bond cleavage:   Unpaired electron: radical  Heterolytic bond cleavage:   Both electrons go to same atom (more electronegative, in this case, A)  Introduction to Reactivity: Lewis Acids and Bases 8 o Two possible ways that orbitals can interact to create a filled, stabilized bonding orbital o One electron from each atom, or a pair to an empty orbital (Lewis acid-base reaction) are both possible and both make bonds o Lewis Acid: electron pair acceptor (empty orbital) o Lewis Base: electron pair donor (full orbital electrons available to donate) o Lewis acid-base reaction, new bond is formed from based sending electron pair to acid    Curved arrows from Lewis base to Lewis acid o General reactivity pattern in organic chemistry:  Lewis acid: electrophile (electron deficient)  Lewis base: nucleophile (electron rich) o Remember: 9  Boron does not have a full octet when it’s neutral, generally a Lewis acid  ???????? 3Lewis acid) is isoelectronic with ????????3(Lewis base)  Organic Compounds and Functional Groups o Reactivity of organic molecules determined by functional groups o Functional group: atom/group of atoms with characteristic physical and chemical properties o Since reactivity is determined by functional groups, carbon backbone of molecules is abbreviated by R o Hydrocarbon Functional Groups  Alkanes  Contain only ???? and ???? in single covalent bonds, no rings, all 3 carbons ???????? hybridized (tetrahedral)  General chemical formula ???????????? 2????+2  ???????? 4 methane, tetrahedral  Expanded, Condensed and Line structure Molecular Condensed Line/Skeletal Formula Structure Structure ???? = 2, ethane ???? ???? ???????? ???????? Line 2 6 3 3 ???? = 3, propane ???? ???? ???????? ???????? ???????? Lines connected at 3 8 3 2 3 vertex   Alkenes  Contain at least one double bond, no rings 10  Generic molecular formula: ???????????? 2????  Additional double bonds reduce ???? number by 2  They’re unsaturated hydrocarbons   Alkynes  At least 1 triple bond, no rings  Generic formula = ???? ???? (1 triple bond) ???? 2????−2  Additional triple bonds reduce H number by 4  Like alkenes, they’re unsaturated hydrocarbons   Aromatics  Contains a benzene ring   Examples: (Don’t need to know) 11 o  Heteroatom-Containing Organic Functional Groups  Alkyl Halide o ???? − ???? where ???? = ????,????????,????????,????  Alcohol o ???? − ????????  Ether o ???? − ???? − ????  Thiol o ???? − ????????  Sulfide o ???? − ???? − ????  Amine o ???????? ,3???? ????,2 − ???????? 2  Next semester functional groups: Aldehyde, Ketone, Carboxylic Acid, Ester, Acid Chloride, Amide, Anhydride, Carbonyl  Classification of Carbons and Functional Groups 12 o Carbon atoms (and other compounds) are classified based on number of other carbons directly bonded to them:  Primary carbon (1° carbon): bonded to one other ???? atoms  Secondary carbon (2° carbon): bonded to two other ???? atoms  Tertiary carbon (3° carbon): bonded to three other ???? atoms  Quaternary carbon (4° carbon): bonded to four other ???? atoms  o Hydrogen atoms (based on type of carbon)  Primary hydrogen (1° hydrogen): bonded to a 1° carbon  Secondary hydrogen (2° hydrogen): bonded to a 2° carbon  Tertiary hydrogen (3° hydrogen): bonded to a 3° carbon o Alcohols  Primary alcohol (1° alcohol): bonded to a 1° carbon  Secondary alcohol (2° alcohol): bonded to a 2° carbon  Tertiary alcohol (3° alcohol): bonded to a 3° carbon o Alkyl Halide  Primary halide (1° halide): bonded to a 1° carbon  Secondary halide (2° halide): bonded to a 2° carbon  Tertiary halide (3° halide): bonded to a 3° carbon o Amines  Primary amine (1° amine): bonded to one ???? atoms (lone pair)  Secondary amine (2° amine): bonded to two ???? atoms (lone pair)  Tertiary amine (3° amine): bonded to three ???? atoms (lone pair)  Quaternary amine (4° amine): bonded to four ???? atoms (positive charge) Chapter Two 13  Alkanes o Simple hydrocarbon molecules with only ???? − ???? and ???? − ???? bonds o Methane (???????? ) 4  Four identical bonds, tetrahedral, from 2????,???????? ,???????? ,2????  How to create four identical bonds from s and p orbitals? Orbital Hybridization  Use ???????? orbitals making it ????????   How:  14


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