lecture 3 - CHEM 231 LECTURE 3 Bonding and nomenclature of...

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Unformatted text preview: CHEM 231 LECTURE 3 Bonding and nomenclature of HYDROCARBONS LECTURE 3 1 Expressing organic structure through drawings Lecture 2 2 The 2 major types of hydrocarbons AlLIPHATIC: alkanes; alkenes and alkynes AROMATIC: things like benzene H H C H H C H H C H H H2 H2C C CH2 H2C C H2 CH3 Lecture 3 3 The physical chemistry of organic bonding Lewis structure model Valence model Molecular orbital model Lecture 3 4 The Lewis Structure model of organic bonding An accounting method with valence considered to be 4 pairs of electrons `Octet rule' Unpaired electrons are available to make bonds An atom bonds until there are no more unpaired electrons Lecture 3 5 The valence model of organic bonding Think of orbitals as waves Waves can collide constructively to make a larger wave Waves can collide destructively to make a smaller wave Lecture 3 6 The valence model of organic bonding Equivalent to bonding Equivalent to antibonding Energy savings Lecture 3 Energy cost 7 The valence model of organic bonding Dihydrogen High Energy 2 hydrogen atoms Bonding is an energy saving process; otherwise bonding woudn't occur Lecture 3 Low energy 8 Orbitals treated as mathematical functions that can be added or subtracted The Molecular orbital model of organic bonding Bonding is additive and energy saving Antibonding is subractive and energy costing Lecture 3 9 Hybridization a construct for organic chemistry First row main group elements; B,C,N,O,F; each have 4 valence orbitals: one s-type orbital and 3 p-type orbitals For any quantum level an s-type orbital is lower in energy than a p-type orbital From experiments, it was learned that all the bonds in methane, CH , are equal in strength Hybridization is a construct to justify how equal strength bonds can come from orbitals with unequal energies Lecture 3 10 2 phases 2 phases s and p atomic orbitals have different shapes and energies sp3 Hybrid energy is average of AO energies Hybrid shape is average of AO shapes Lecture 3 11 Sigma bonds and lone pairs can come from sp hybridized The natural angle of sp The shape of the sp hybridized orbital is like a stylized straw orbitals Sigma bonds made 109.5 hybridized orbitals is from sp hybridized orbitals are mushroom Lecture 3 12 H If there are 4 bonds to the central sp atom then the shape is If there are 3 bonds to the central sp atom then and one orbital has a If there are 2 bonds to the central TETRAHEDRAL sp atom and 2 orbitals have a lone lone pair the shape to PYRAMIDAL If there are 1 bond is the central sp H C H H H N shape H H H O H F H 13 Lecture 3 Tetrahedral shape in detail Points behind Points ahead H H H C Points behind Points ahead H Tetrahedral shape is three dimensional Lecture 3 14 One p orbital remains it is higher in energy sp2 orbitals are the average of the s and 2 of the p orbitals in shape and energy Since the Z orbital does not participate in the hybridization all sp centers give at the most 2Lecture 3 15 Sigma bonds and lone pairs can come from sp hybridized The natural angle of sp The shape of the sp hybridized orbital is like a stylized orbitals hybridized orbitals is 120 Sigma bonds made from sp hybridized orbitals are rounder straw mushroom Lecture 3 16 What the leftover p orbital can do The orbital can be singly occupied and be part of a pi bond H C H C H H The p orbitals are in phase, indicating bonding The orbital can be singly occupied and not bonded- this is a radical The orbital can be empty Neutral boron compounds always have an empty p orbital Carbocations always have the empty p orbital Lecture 3 H3C C CH3 CH3 F B F F H3C C CH3 CH3 17 What the leftover p orbital can do Lecture 3 18 H If all 3 of the sp hybridized orbitals is H H C C C C H C C H H If 2 of the sp hybridized orbitals is part of a sigma bond then the shape is FLAT (PLANAR) If one of the sp hybridized orbitals is part of a sigma bond and third is in a lone pair, the shape of the molecule is BENT Lecture 3 H2 C N CH3 H2 C O 19 sp Hybridization Unhybridized orbitals don't see the energy savings Hybrid orbitals are lower in energy Lecture 3 20 Properties of sp hybridized orbitals Sigma bonds and lone pairs can come from sp hybridized orbitals The natural angle of sp hybridized orbitals is 180 The shape of the sp hybridized orbital is like a stylized parachute (red) and skydiver (blue) Sigma bonds made from sp hybridized orbitals are formed from the overlap of the red phase of the orbitals with other orbitals of the same phase (red) Lecture 3 21 What the leftover p orbitals can do The orbitals can be singly occupied and each be part of a pi bond H H C H C C H One orbital orbital can be singly occupied and not bonded- this is a radical derived from an alkene H H C H C C H H The orbital can be empty This is a cation derived from an alkene Lecture 3 H H C H C C H H 22 What the leftover p orbitals can do Lecture 3 23 The shape of molecules with an sphybridized orbital is LINEAR O C CH2 N C CH3 Lecture 3 24 Dr. Kissling's method for determining hybridization by looking at a molecule drawing The key bit of knowledge is that single bonds(sigma bonds) and lone pairs of electrons are always formed from hybrid orbitals Count the number of single bonds (sigma bonds) and lone pairs and that gives the number of hybridized orbitals on that atom. For example 4 sigma bonds uses all 4 atomic orbitals together thus sp Lecture 3 25 Possible levels of oxidation in carbon Most reduced Most oxidized Lecture 3 26 Dr. Kissling's method for determining how oxidized an atom is The key bit of knowledge its all about relative electronegativities An element more electronegative than carbon oxidizes, an element less electronegative than carbon reduces A carbon bonded to 3 O and 1 H is more oxidized than carbon bonded to 2 carbons and 2 H. Lecture 3 27 Alkanes Ethane, a simple alkane Carbons and hydrogens only All carbons are sp hybridized Lecture 4 28 Physical Properties of Alkanes-Stability based on structure Relative stability is determined by comparing the heat produced by isomers upon combustion Least stable Produces most heat Most stable Produces least heat Branching increases stability of alkane Lecture 4 29 Lets talk about isomers The simplest alkanes do not have isomers CH4 H3 C CH3 H2 C H3 C CH3 C H is long enough to have isomers H2 C CH3 H C H3C CH3 30 H3C C H2 CH3 Lecture 4 As we increase in carbons the number of isomers increases too C H isomers are shown here H2 C H2 C CH3 C C H2 CH3 CH3 CH3 C H3C CH3 H CH3 H3C H3C C H2 Lets figure out all the isomers for C H 31 Lecture 4 Here are the cycloalkane isomers of C H alkene Here are the alkene isomers of C H Lecture 4 32 ...
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