CHEM 221 section 01 LECTURE #02 Thurs., Sept.8, 2005
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1 EM 221 section 01 LETURE #02 Thurs., Sept.8, 2005 ASSIGNED READINGS: TODAY S LASS: 1.6 Introduction to molecular orbital theory Bonding: single, double, triple bonds 1.10 Bonding in cations, anions, radicals Bonding in simple compounds 1.14 Summary: Bond lengths, strengths & angles NEXT LASS: , (1)
2 Typical ways to draw organic structures Kekulé structures: show all bonds but not lone pairs ondensed structures: formulas that show connectivity Kekulé structure ondensed structures (2)
3 Kekulé structure ondensed structures Two or more identical groups considered bonded to the first atom on the left can be shown (in parentheses) to the left of that atom, or hanging from the atom. N or ( 3 ) 2 N N (3)
4 (4) Kekulé structure ondensed structures An oxygen doubly bonded to a carbon can be shown hanging off the carbon or to the right of the carbon. O O 3 O or or 3 (=O) O O O
5 1.5 REVIEW: Atomic orbitals The 1s & 2s Orbitals Spherical where s core e - sreside Part of s valence (5)
6 1.5 REVIEW: Atomic orbitals The 2p Orbitals Dumbell-shaped Opposite mathematical sign on either side of node (shown as +/- or colours) Lobes actually rather diffuse 3 degenerate orbitals in 3 orientations: (6)
![2-5] + overlap = Overlap localized between the atoms only = Valence Bond Theory description](/docs-images/88/116080373/images/7-2.jpg "(simpler view) overlap results in new orbital spread over entire molecule = Molecular Orbital")
7 1.6 Molecular Orbital ( M.O. ) Theory [Note: Gen. hem. texts also helpful - Kotz & Treichel 10.3; Zumdahl 9.2-5] + overlap = Overlap localized between the atoms only = Valence Bond Theory description (simpler view) overlap results in new orbital spread over entire molecule = Molecular Orbital Theory description (more realistic) a σ-bond: formed by the overlap of two s orbitals (7)
8 Bond strength descrobed by bond dissociation energy = energy required to break a bond (always endothermic!) = energy released when a bond forms (always exothermic!) Bruice Figure 1.2 (8)
9 Remember that orbitals are WAVES Like waves: orbitals have phase & interfere if overlapped onsider a simplified view of 2 s-orbitals: In-phase overlap: Out-of-phase overlap: See Fig constructive interference destructive interference (9)
10 In-phase overlap forms a bonding MO out-of-phase overlap forms an antibonding MO Important: BOT types of overlap occur every time 2 orbitals overlap! Figure 1.4 An MO diagram ON OUTSIDES: relative energies, names & occupancies of the orbitals of the 2 atoms coming together & ON INSIDE: the resulting MO s Note: both e - s will occupy the bonding orbital (since lower in E!) (10)
11 IMPORTANT: Bonding occurs if it is energetically favourable If a bond s σ AND σ* orbital are simultaneously filled: the bond will not actually exist. Example on board: Justify why e 2 does not form. (11)
12 A sigma bond (s) can also be formed by end-on overlap of two p orbitals: So, what makes a bond a sigma-bond? if the lobes of the orbitals that are overlapping are pointing at each other along the bond axis (i.e., the space between the nuclei) the bond can rotate without diminishing the overlap (12)
13 A pi bond (π) is formed by side-on overlap of 2 p orbitals: So, what defines a π-bond? the overlapping lobes are parallel to each other the bond cannot rotate without diminishing the overlap Note: A σ bond is stronger than a π bond. WY? Larger parts of the orbitals actually overlap when end-on vs. when oriented side-on. i.e., higher e - density in region of overlap (13)
14 1.7 Bonding in Methane & Ethane: Single Bonds Recall VSEPR model for predicting geometries: draw Lewis structure: determine # of single bonds & lone pairs electron pairs (bonding & lone) will spread themselves out in space as far from each other as possible Atomic orbitals hybridize to achieve necessary geometries Valence shell of : Kotz&Treichel Figure 10.6 (See Bruice Fig ) (14)
15 ybridization of atomic orbitals IF MIX TESE GET BAK TESE ATOMI ORBITALS YBRID ORBITALS One s & one p One s & two p s One s & three p s ANGLE b/w YBRIDS GEOMETRY 2 sp hybrids 180 Linear 3 sp 2 hybrids 120 Trig. planar 4 sp 3 hybrids Tetrahedral To make a Td carbon: One s & three p s Fat lobes of hybrids used for bonding Kotz & Treichel Figure 10.6 (ompare to Bruice Fig ) (15)
16 Orbital description of the bonding in Methane, 4 4 : requires 4 bonds to tetrahedral geometry sp 3 hybrid set: Tetrahedral bond angle Each σ-bond formed by one sp 3 hybrid overlapping with one 1s orbital Summary: 4 σ-bonds produced by 2sp 3 1s overlap Valence Bond view (easier to see) of the BONDING σ-orbitals but reality is MO (16) Figure 10.6
17 Orbital description of the bonding in Ethane, 3 3 Requires 4 bonds to each each has Td geometry 6 - σ-bonds: 2sp 3 1s overlap & 1 - σ-bond: 2sp 3 2sp 3 overlap Bruice Figure 1.14 MO diagram for a - σ-bond (17)
18 A simple organic reaction: 3 I + F - 3 F + I - Draw the orbitals involved in the - & -I bonds of methyl iodide (both bonding & antibonding): What would happen if a lone pair on F - entered the -I σ* orbital? (18)
19 1.8 Bonding in Ethene: A Double Bond (=1σ+1π) Requires 3 bonds to each each is trigonal sp 2 hybridized Unhybridized (left over) Bruice Fig.1.15 sp 2 -hybridized carbon atom: (19)
20 1.8 Bonding in Ethene: A Double Bond (=1σ+1π) σ-framework: π-bond forms between parallel unhybridized p-orbitals: 1 - σ-bond: 2sp 2 2sp π-bond: 2p - 2p 4 - σ-bonds: 2sp 2 1s TE DOUBLE BOND (20)
21 1.9 Bonding in Ethyne: A Triple Bond (=1σ+2π) Requires 2 bonds to each linear sp hybridized see Bruice Fig unhybridized p-orbitals 1 - σ-bond: 2sp 2sp 2 - π-bonds: 2p - 2p 2 - σ-bonds: 2sp 1s TE TRIPLE BOND (21)
22 1.10 Bonding in methyl cation, radical & anion Methyl cation 3 + (3 pairs of valence e - s) Methyl radical 3 (3 pairs + 1 extra valence e - ) Unpaired e- spread out over whole p-orbital, above & below plane 1. arbocations are planar at their open-shell (e - -deficient) centre. 2. arbon radicals are planar at the radical centre: an unpaired e - does NOT exert enough repulsion to make atom rehybridize (22)
23 Methyl Anion 3-4 pairs of valence e - s 3. arbanions are tetrahedral at the anionic centre: another pair of electrons DOES exert enough repulsion to make the atom rehybridize just like we see for N 3, 2 O, etc based on VSEPR model (23)
24 1.11 Bonding in water 1.12 Bonding in ammonia & ammonium cation (24)
25 1.13 Bonding in the hydrogen halides weakest acid strongest acid -X bonds involve 1s X np overlap larger value of n larger orbital (higher energy) SO WAT? mismatch in orbital sizes results in less efficient overlap weaker bond reflected in bond energy & bond length (25)
26 1.14 Summary of important bonding principles sp 3 -sp 3 sp 3 -s sp 2 -sp 2 & p-p sp 2 -s sp-sp & 2x p-p sp-s 1.) igher bond order stronger bond (for same atom types) 2.) The greater the s-character, the stronger the bond. WY? s-orbitals are closer to nucleus than p s (or hybrids) thus: stronger nucleus-to-electron attraction (26)
27 Summary: important bonding principles When a bonding / antibonding orbital pair are filled, the bond will break (or never form) A π bond is weaker than a σ bond The greater the electron density in the region of orbital overlap, the stronger is the bond The more s character: the shorter and stronger is the bond (27)
28 ASSIGNED READINGS BEFORE NEXT LETURE: Read: rest of Bruice h.1 Review: hem206: acids & bases - conjugate pairs -p, pk a, pk b reactions of acids & bases - effect of structure on acidity Note: supplementary orbital figures were taken from Kotz & Treichel s hemistry & hemical Reactivity 5 th Ed. On reserve at Vanier library for hem 205/206 (28)
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