CH 101 CHAPTER 7 NOTES!!
CH 101 CHAPTER 7 NOTES!! CH 101
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This 16 page Class Notes was uploaded by Allie Newman on Saturday October 31, 2015. The Class Notes belongs to CH 101 at University of Alabama - Tuscaloosa taught by Professor John McDuffie in Fall 2015. Since its upload, it has received 50 views. For similar materials see General Chemistry in Chemistry at University of Alabama - Tuscaloosa.
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Date Created: 10/31/15
CHEM 101 Chapter 7 Notes A Magnetic Liquid 02 Oxygen is paramagnetic Paramagnetic material has unpaired electrons Neither Lewis theory nor valence bond theory predict this result Problems with Lewis Theory Lewis theory generally predicts trends in properties but it does not give good numerical predictions Example bond strength and bond length Lewis theory gives good rst approximations of the bond angles in molecules but it usually cannot be used to get the actual angle Lewis theory cannot write one correct structure for many molecules where resonance is important Lewis theory often does not predict the correct magnetic behavior of molecules For example 02 is paramagnetic although the Lewis structure predicts it is diamagnetic Valence Bond Theory Orbital Overlap and the Chemical Bond 0 Applying quantum mechanics to molecules 0 Bonds between atoms should occur when the orbitals on those atoms interacted to make a bond 0 The kind of interaction depends on whether the orbitals align along the axis between the nuclei or outside the axis Interaction Energy of Two Hydrogen Atoms Bond length 4 r H H distance gt Summarizing Valence Bond Theory The valence electrons of the atoms in a molecule reside in quantummechanical atomic orbitals The orbitals can be the standard 5 p d and forbitals or they may be hybrid combinations of these A chemical bond results from the overlap of two half lled orbitals and spinpairing of the two valence electrons or less commonly the overlap of a completely lled orbital with an empty orbital The geometry of the overlapping orbitals determines the shape of the molecule Orbital Diagram for the Formation of H25 H 1 Half lled 15 orbitals overlap H l 15 8 ll llll 35 SP Filled p orbital Bonds formed Filled s orbital Valence Bond Theory Main Concepts 1 The valence electrons of the atoms in a molecule reside in quantummechanical atomic orbitals a The orbitals can be the standard 5 p d and forbitals orthey may be hybrid combinations of these 2 A chemical bond results when these atomic orbitals interact and there is a total of two electrons in the new molecular orbital a The electrons must be spin paired 3 The shape of the molecule is determined by the geometry of the interacting orbitals Bonding with Valence Bond Theory According to valence bond theory bonding takes place between atoms when their atomic or hybrid orbitals interact Overlap To interact the orbitals must do either of the following Be aligned along the axis between the atoms Be parallel to each other and perpendicular to the interatomic axis Valence Bond Theory and Hybridization One of the issues that arises is that the number of partially lled or empty atomic orbitals did not predict the number of bonds or orientation of bonds C 2522pX12py12pz0 would predict two or three bonds that are 90 apart rather than four bonds that are 1095 apart 0 To adjust for these inconsistencies it was postulated that the valence atomic orbitals could hybridize before bonding took place One hybridization of C is to mix all the 25 and 2p orbitals to get four orbitals that point at the corners of a tetrahedron Unhybridized Carbon Orbitals in CH4 Predict the Wrong Bonding and Geometry ls Theoretical prediction C 1 1 25 2P Hybridization Some atoms hybridize their orbitals to maximize bonding o More bonds more full orbitals more stability 0 Hybridizing is mixing different types of orbitals in the valence shell to make a new set of degenerate orbitals 0 5p spz sp3 sp3d 302 0 The same type of atom can have different types of hybridization o C 5p spz 5p3 Large lobe allows for gireater overlap Hybrid orbital Hybrid Orbitals The number of standard atomic orbitals combined the number of hybrid orbitals formed Example for carbon Combining a 25 with a 3p gives two sp3 hybrid orbita5 H m hybridize lts valence shell only has one orbital I6 Hybridization Lid J l Foar sp hybrid orbitals J Energy The number and type of standard atomic orbitals combined determines the shape of the hybrid orbitals o The particular kind of hybridization that occurs is the one that yields the lowest overall energy for the molecule Formation of 503 lHybrid Orbitallls One 5 orbital and three 0 orbitals combine to form foiur spi3 orbitals l Z Z Jquot If I n 1quot x x f I 3 r x I r If I I It fl 1 l E ff 7 39 I if xxx if 5 orbital pX orbital a 3 rr 5P L x v 3 quot Hybridization 5p quot r s i Z Z II 5p hybrid orbitals 39 I shown together f l t 39 J r I If gquot I f a l t y x I 3quot x x 3 lb f x x39 I I x py orbital pz orbital I Unhybridized 39 atomic orbitals r39 5p3 hybrid orbitals shown separately sp3 Hybridization Atom with four electron groups around it Tetrahedral geometry 1095 angles between hybrid orbitals Atom uses hybrid orbitals for all bonds and lone pairs C l i ll sp3 Hybridization sp2 Atom with three electron groups around it Trioonal planar svstem C trigonal planar N trigonal bent O linear 120 bond angles Flat 0 Atom uses hybrid orbitals for abonds and lone pairs and uses nonhybridized p orbita for 1 bond EEEEEEEEE l l l I 2P I I I Unlhybridized p orbital I I Hybridization Energy I p j I I l l Three 5p2 hybrid orbitals J l 25 7 Standard atomic orbitals Formation of 502 Hybrid Orbit als One 5 orbital and two p orbitals combine to form three 5p2 orbitals y r x x Hybridization b gt s orbital px orbital Ity orbital Sailg f g igf Unhybridized atomic 5 orbitals l 3 2 5p2 hylbrid orbitals shown separately Types of Bonds 0 A sigma a bond results when the interacting atomic orbitals point along the axis connecting the two bonding nuclei 0 Either standard atomic orbitals or hybrids sto s p to p hybrid to hybrid sto hybrid etc o A pi 1 bond results when the bonding atomic orbitals are parallel to each other and perpendicular to the axis connecting the two bonding nuclei 0 Between unhybridized parallel p orbitals o The interaction between parallel orbitals is not as strong as between orbitals that point at each other therefore 0 bonds are stronger than 1 bonds Sigma and Pi Bonding Half tilled Half filled pyorpZ orbital pyrorpzorbltal W on 0 D a 0 Halftilled Halffilled px orbital px orbital Orbital Diagrams of Bonding Overlap between a hybrid orbital on one atom with a hybrid or nonhybridized orbital on another atom results in a abond Overlap between unhybridized p orbitals on bonded atoms results in a nbond Hybrid orbitals overlap to form a abond Unhybridized p orbitals overlap to form a 1 bond Carbon unhybridized Carbon 5p2 p orbital hybrid orbitals Oxygen 2p orbitals Illydrogen s orbitals E Onewbondl One 0 bornle K J Y Double bondl Bond Rotation Because of the orbitals that form the abond point along the internuclear axis rotation around that bond does not require breaking the interaction between the orbitals But the orbitals that form the 1 bond interact above and below the internuclear axis so rotation around the axis requires the breaking of the interaction between the o rb I ta s H I l Free rotation Rotation restr ictedl H H H C C H about single bond by double bond cc Sigmal sigma pl 21 Cl Cl C1 crC5p3Hl c l H 7 0 SPquot i 3 039 copquot c1p aquot Cllsp 2 Cspg a Clspa Clip Cis amp Trans lsomers of 12Dichloroethane H g H H CC cc C1 H j C Cis12Dichloroethane trans12Dichloroethane sp Hybridization and Triple Bonds Atom with two electron groups Linear shape 180 bond angle 0 Atom uses hybrid orbitals for abonds or lone pairs and uses nonhybridized p orbitals for 1 bonds Unhybriclizeci p orbitals J Energy Two 511 hybrid orbitals J l llybridization Standard atomic orbitals Formation of 5p Hybrid Orbital Formation of sp llybridl Orbitals One 3 orbital and one p orbital combine to form two Sp orbitals j y Hybridization r I u x l x 7 I 5p hybrid orbitals 5 O39rb39llal px orb39tal 5p hybrid orbitals shown together shown separately Unbylaridizedl atomic orbitals Formation of 5p Hybrid Orbital in Ethyne 7T Q CP Valence lborid model sp3d Hybridized Orbital Atom with ve electron groups around it Trigonal bipyramid electron geometry Seesaw Tshape linear 120 and 90 bond angles 0 Use empty dorbitals from valence shell dorbitals used to make 1 bonds Valence bond model Unhybridized d orbitals l Hybridization l l Five Spsd orbitals l Standard atomic orbitals Energy a 5p3d hybrid orbitals shown together 13 sp3d2 Atom with six electron groups around it Octahedral electron geometry Square pyramid square planar 90 bond angles Use empty dorbitals from valence shell to form hybrid dorbitals used to make 1 bonds bondmodel sp3d2 Hybridized Orbital I 3d i Unhybriciized d orbitals L 1 i Six 517312 hybrid orbitals Standard atomic orbitals 3 513352 hybrid orbitals shown together bi 731 Hybridization Scheme from Electron Genrnetry Number of Electron Geometry Hybridization Orbital Shapes and Electron Groups from VSEE R Theory Scheme Relative Orientation 2 Li near 3p 1 3 Trigonal planar still2 1 20 x 1 09 5 4 Tetra h ed ra l Sp3 a 90 3 5 Trlgonal blpyramldal Sp 039 4 1 20 x 90 6 Uctahedral Spadz Predicting Hybridization and Bonding Scheme 1 Start by drawing the Lewis structure 2 Use VSEPR theory to predict the electron group geometry around each central atom 3 Use Table 71 to select the hybridization scheme that matches the electron group geometry 4 Sketch the atomic and hybrid orbitals on the atoms in the molecule showing overlap of the appropriate orbitals 5 Label the bonds as aor 11 Problems with Valence Bond VB Theory 0 VB theory predicts many properties better than Lewis theory 0 Bonding schemes bond strengths bond lengths bond rigidity However there are still many properties of molecules it doesn t predict perfectly 0 Magnetic behavior of 02 o In addition VB theory presumes the electrons are localized in orbitals on the atoms in the molecule it doesn t account for delocalization Molecular Orbital MO Theory Electron Delocalization o In MO theory 0 Applies Schrodinger s wave equation to the molecule to calculate a set of molecular orbitals The equation solution is estimated The estimated solution is evaluated and adjusted until the energy of the orbital is minimized In this treatment the electrons belong to the whole molecule so the orbitals belong to the whole molecule 0 Delocalization LCAO Linear Combination of Atomic Orbitals The simplest guess starts with the atomic orbitals of the atoms adding together to make molecular orbitals this is called the linear combination of atomic orbitals LCAO method Weighted sum 1 T Energy Bonding orbital Because the orbitals are wave functions the waves can combine either constructively or destructively Node l f o i o O O 39 1 5 15 Molecular Orbitals 0 When the wave functions combine constructively the resulting molecular orbital has less energy than the original atomic orbitals o Called a bonding molecular orbital o Designated a n 0 Most of the electron density between the nuclei 0 When the wave functions combine destructively the resulting molecular orbital has more energy than the original atomic orbital o Called an antibonding molecular orbital o Designated 0 11 0 Most of the electron density outside the nuclei 0 Nodes between nuclei Interaction of 15 Orbitals Destructive interference is is 012 0 is Antibonding molecular orbital Bonding molecular orbital Constructive interference 15 ls Molecular Orbital Theory Electrons in bonding MOs are stabilizing 0 Lower energy than the atomic orbitals o Electrons in antibonding MOs are destabilizing 0 Higher in energy than atomic orbitals 0 Electron density located outside the internuclear axis 0 Electrons in antibonding orbitals cancel stability gained by electrons in bonding orbitals MO and Properties 0 Bond order difference between number of electrons in bonding and antibonding orbitals 0 Only need to consider valence electrons 0 May be a fraction 0 Higher bond order stronger and shorter bonds 0 If bond order 0 then the bond is unstable compared to individual atoms and no bond will form 0 A substance will be paramagnetic if its MO diagram has unpaired electrons o If all electrons are paired it is diamagnetic Bond order 12 Bonding Electrons Antibonding Electrons I quot Antibonding 15 r 15 l l I H atom Bonding Iquot H atom ll Ll Energy H2 molecule Bond order 12 Bonding Electrons Antibonding Electrons Bond Order H2 12 2 0 1 Why Doesn39t the Molecule H62 Exisi Antibondlng l Energy lll lt l ll 15 I He atom x Bonding If He atom ll is Hez molecule Not stable Bond order 12 Bonding Electrons Antibonding Electrons Bond Order H92 12 2 2 0 Why Does the Molecule He2 Exist Energy 15 H l l l 15 He atom Bonding I He i011 l ll r s He ion Bond order 12 Bonding Electrons Antibonding Electrons Bond Order H92 12 2 1 1 Summarizing LCAOMO Theory Molecular orbitals MOS are a linear combination of atomic orbitals AOs The total number of MOs formed from a particular set of A05 always equals the number of A05 in the set When two AOs combine to form two MOs one MO is lower in energy the bonding MO and the other is higher in energy the antibonding MO When assigning the electrons of a molecule to MOs we ll the lowest energy MOS rst with a maximum of two spinpaired electrons per orbital When assigning electrons to two MOs of the same energy Hund s rule is followed to ll the orbitals singly rst with parallel spins before pairing The bond order in a diatomic molecule is the number of electrons in bonding MOs minus the number in antibonding MOS divided by two Stable bonds require a positive bond order more electrons in bonding MOs than in antibonding MOS Atomic orbital Period Two Homonuclear Diatomic Molecules l Atomic Molecular Atomic Atomic Molecular orbital orbitals orbital orbital orbitals 03925 1 L 03925 I l I i I i l quot quot kl Antibondin quot Antibondmg 2 I l r r l l l 9 i 2 25 I 25 E S 5 1 I l nergy l L x I l L B l t l 139 l B t L1 atom Bonding if L1 atom e a 0m in Bonding I e a om O f l l w 03925 I 25 362 molecule Li molecule Not stable Liz molecule BEz molecule Interaction of p Orbitals Molecular orbitals antib ending I O I C 2p 2p bonding 0 X Molecular orbitals J antibonding Atomic orbitals I l39 bonding 202 2p 1T2 I Molecular Orbital Energy Ordering r Atomic orbitals Molecular Atomic orlbitals orbitals l l l x 1 029 I W 02p i WEI r Energy UZP ll I f 1 0 25 025 I 7 I f ll 2D ll RD 25 x x 25 25 x 25 El D 725 725 Bz C2 N2 02 P2 N62 Molecular Orbital Energy Diagrams for SecondPeriodp Block Homonuclear Diatomic Molecules Small 25 pr interaction D E ll ll liiliil liillil mlillll lllll lillil a2 02 ll ll 1 2 3 Bond energy klmull 290 620 946 498 159 Bond length pm 159 131 110 121 143 Bond order Heteronuclear Diatomic Molecules and Ions When the combining atomic orbitals are identical and of equal energy the contribution of each atomic orbital to the molecular orbital is equal When the combining atomic orbitals are different types and energies the atomic orbital closest in energy to the molecular orbital contributes more to the molecular orbital The more electronegative an atom is the lower in energy are its orbitals Lower energy atomic orbitals contribute more to the bonding MOS Higher energy atomic orbitals contribute more to the antibonding MOS Nonbonding MOs remain localized on the atom donating its atomic orbitals SecondPeriod Heteronuclear Diatomic Molecules Atomic orbitals 2p Energy a 1L My 025 NOmolecule NO molecule N atom 25 0 atom MO and Polyatomic Molecules When many atoms are combined together the atomic orbitals of all the atoms are combined to make a set of molecular orbitals which are delocalized over the entire molecule AtOmic orbitals ls 39 2 Energy quot H atom Atomic orbitals Noribonding orbitals 219y sz X 2p XX1 039 HF molecule F atom HF molecule Gives results that better match real molecule properties than either Lewis or valence bondtheodes Lewis structure Bonding in Metals and Semiconductors The simplest theory of metallic bonding involves the metal atoms releasing their valence electrons to be shared as a pool by all the atomsions in the metal An organization of metal cation islands in a sea of electrons Electrons delocalized throughout the metal structure Bonding results from attraction of cation for the delocalized electrons Valence bond model e sea Semiconductors and Band Theory 0 Band Theory 0 Electrons become mobile when they make a transition from the highest occupied molecular orbital into higher energy empty molecular orbitals oThese occupied molecular orbitals are referred to as the valence band oThe unoccupied orbitals the conduction band Conduction band Conduction lbalrlld conducuon band N0 energy gap Small energy gap Large energy gap l valence band valence band Valence band
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