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Chemical Bonding II

Chemical Bonding II. Molecular Shape (VSEPR) Molecular polarity Bonding Theory Valence Bond Theory Molecular Orbital Theory. VSEPR Model. VSEPR (valence shell electron pair repulsion) model Useful for predicting the shapes of species that have main group elements as central atoms

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Chemical Bonding II

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  1. Chemical Bonding II • Molecular Shape (VSEPR) • Molecular polarity • Bonding Theory • Valence Bond Theory • Molecular Orbital Theory

  2. VSEPR Model • VSEPR (valence shell electron pair repulsion) model • Useful for predicting the shapes of species that have main group elements as central atoms • Electron (bonding and nonbonding) pairs repel each other. • Molecular geometry is arranged to minimize repulsions. • The geometry is based on the total number of electron pairs • Determine no. of electron pairs around the central species • Count both bonded and unshared pairs • Multiple bonds are treated as single bonds for geometry

  3. AB2 2 0 linear # of atoms bonded tocentral atom # lone pairs on central atom trigonal planar trigonal planar AB3 3 0 Arrangement of e- pairs Molecular Geometry Class linear B B VSEPR model Electron (bonding and nonbonding) pairs repel each other. Geometry of molecules is arranged to minimize repulsions.

  4. # of atoms bonded tocentral atom # lone pairs on central atom Arrangement ofelectron pairs Molecular Geometry Class tetrahedral tetrahedral AB4 4 0 trigonal bipyramidal trigonal bipyramidal AB5 5 0 VSEPR model

  5. octahedral AB6 6 0 # of atoms bonded tocentral atom # lone pairs on central atom Arrangement ofelectron pairs Molecular Geometry Class octahedral VSEPR model

  6. # of atoms bonded tocentral atom # lone pairs on central atom trigonal planar trigonal planar AB3 3 0 Arrangement ofelectron pairs Molecular Geometry Class tetrahedral octahedral tetrahedral octahedral AB6 AB4 4 6 0 0 trigonal bipyramidal trigonal bipyramidal AB5 5 0 VSEPR model AB2 2 0 linear linear

  7. linear tetrahedral octahedral Examples • BeCl2 • CBr4 • IOF5

  8. lone-pair vs. lone pair repulsion lone-pair vs. bonding pair repulsion bonding-pair vs. bonding pair repulsion > > Effect of Lone Pairs

  9. # lone pairs # of e- pair Arrangement of e- pairs Molecular geometry Ex. Bond angle Class VSEPR AB2 2 0 linear linear CO2 180 AB3 3 0 tri. planar tri. planar BF3 120 AB2E3 1 tri. planar bent NO2– <120

  10. # lone pairs # of e- pair Arrangement of e- pairs Molecular geometry Ex. Bond angle Class VSEPR AB4 4 0 tetrahedral tetrahedral CCl4 109.5 AB3E4 1 trigonal pyramidal NH3 <109.5 AB2E2 4 2 bent H2O <109.5

  11. # lone pairs # of e- pair Arrangement of e- pairs Molecular geometry Ex. Bond angle Class VSEPR AB5 5 0 tri. bipyramidal tri. bipy PF5 90, 120 AB4E5 1 distorted tetrahedron SF4 ~90, 120 (seesaw) AB3E2 5 2 T-shaped ClF3 < 90 AB2E3 5 3 linear XeF2 ~180

  12. # lone pairs # of e- pair Arrangement of e- pairs Molecular geometry Ex. Bond angle Class VSEPR AB6 6 0 octahedral octahedral SF6 90 AB5E6 1 square IF5 ~ 90 pyramidal AB4E2 6 2 square planar XeF4 ~ 90

  13. AB2 AB2E3 AB2E2 AB5E VSEPR Shapes • Try these: HCNI3– NH2– BrF5 • For species with more than one central atoms • Predict VSEPR shape for each individual central atom • Combined the shapes to determine the molecular geometry (a) C2H6 (b) C2H4 - planar (c) C2H2 - linear (d) CH3COOH

  14. F H d- d+ Q is the charge r is the distance between charges 1 D = 3.36 x 10-30 C m Molecular Polarity • Polarity • An important molecular property • Affect the physical and chemical properties of the molecules • Dipole Moment, m • A quantitative measurement of the polarity of a molecule • Polar molecules will align in an electric field • Nonpolar molecules will not align in an electric field electron rich region electron poor region m = Q x r

  15. polar O nonpolar C O O C ABn molecules (B are identical atoms) with the following shapes: linear, trigonal planar, tetrahedral, square planar, trigonal bipyramidal, octahedral are nonpolar Polarity and Molecular Geometry • The degree of molecular polarity is a function of • Molecular geometry • # and type of polar bonds • Examples • CO vs. CO2 • H2O, CH3Cl, CH4, which of these are polar • Which one is more polar: H2O or H2S H2O: polar CH3Cl: polar CH4: nonpolar Similar shape, O-H bond is more polar => H2O is more polar

  16. Bonding Theory • Require quantum mechanics • Two approximate methods • (1) Valence bond (VB) theory • bonds are formed by overlap of atomic orbitals • (2) Molecular orbital (MO) theory • atomic orbitals are combined to form molecular orbitals

  17. + 74 pm Valence Bond Theory • Bonds are formed by overlap of atomic orbitals (AO) • Consider H2 molecule Lewis structure : H-H Bond energy : 436 kJ/mol Bond length: 74 pm = 0.74 Å * H2 is formed by overlap of 1s orbitals from each H atom

  18. Hybridization • Hybridization- mixing of two or more atomic orbitals to form a new set of hybrid orbitals • Hybrid orbitals: linear combination of AO’s on the same atom # hybrid orbitals= # of AO combined • Hybrid orbitals may have very different shape from original AO • Hybrid orbitals are needed to account for the molecular shape • Covalent bonds are formed by: • Overlap of hybrid orbitals with atomic orbitals • Overlap of hybrid orbitals with other hybrid orbitals

  19. 2p sp3 energy 2s Hybridized orbitals Unhybridized AO sp3 Hybrid Orbitals Remember C: 1s22s22p2 • Consider CCl4 Tetrahedral geometry, why? • sp3 hybrid orbitals (25% s- and 75 % p-orbital character) • 1 2s + 3 2p => 4 sp3 orbitals

  20. 2p sp 2p energy 2s Unhybridized Hybridized sp Hybrid Orbital Remember Be: 1s22s2 • Conside BeH2, BeCl2 • Linear molecule • sp hybrid orbitals

  21. Formation of sp Hybrid Orbitals

  22. p orbital Sp2hybrid 2p 2p energy sp2 2s Unhybridized Hybridized sp2 Hybrid Orbitals Remember B: 1s22s22p1 • BF3 molecule Trigonal planar • sp2 hybrid orbitals (33% s- and 67 % p-orbital character) • 1 2s + 2 2p orbitals = 3 sp2 orbitals

  23. Formation of sp2 Hybrid Orbitals

  24. Hybrid Orbitals • d orbitals can also be involved in the hybrid orbitals (1) sp3d hybrid, example: PF5 (2) sp3d2 hybrid, example: SF6 • Correspondence between shape and hybridization Hybrid # lp+# atom e- pair Shape (Ex) sp 2 Linear (BeCl2) sp2 3 Trigonal planar (BF3) sp3 4 Tetrahedral (CH4,NH3,H2O) sp3d 5 Trigonal bipyramidal (PCl5) sp3d2 6 Octahedral (SF6)

  25. H H H C C H Multiple Bonds • Consider C2H4 Lewis structure: Hybridization : sp2 • Single bond (s bond) • Formed by head-on overlap • p bond • Formed by sidewise overlap • Difficult to rotate along p bond

  26. p orbital sp2hybrid

  27. Bonding in Ethylene (C2H4) 2pz 2pz H H 1s 1s p orbital H H C C H H C C Sp2hybrid H H 1s 1s p H H s s s C C s s H H p Sigma bond (s) – electron density between the 2 atoms Pi bond (p) – electron density above and below plane of nuclei of the bonding atoms sp2orbital

  28. H C C H Bonding in Acetylene (C2H2)

  29. H H H H H H = Delocalized Electron — Benzene • Benzene, C6H6 is a traditional example of delocalized e- • Structure of benzene • p orbitals overlap sidewise all around the ring. No localized double bonds.

  30. Exercises Consider the following molecule H H O 1 | | || H-C = C - C = C - C - O -H 1 2 3 4 5 2 4 10 Total # of s bonds _____, total # of p bonds ____. The bond angle between C4, C5, and O2 is about ______. The hybrid orbitals to describe bonding on C1 is _______. The hybrid orbitals to describe bonding on O2 is _______. 120o sp sp3

  31. Molecular Orbital Method • Bonds are formed from interactions of atomic orbitals (AO) to form molecular orbitals (MO). • # of MO = # of total AO combined • AOs combine most effectively when they have similar energy and proper orientation • Two types of Mos • Bonding orbital - s or p The energy is lower than the AOs and the e- density overlaps • Antibonding orbital - s* or p* The energy is higher than the AOs and the e- density does not overlap • Each MO can only hold two e- (Puali’s exclusion principle) • The filling of MOs proceeds from low to high energies • Hund’s rule

  32. Molecular Orbital Diagram — H2 A bonding MO has lower energy and greater stability than the AOs from which it was formed. An antibonding MO has higher energy and lower stability than the AOs from which it was formed. Bond order = 1/2 (# bonding electrons- # antibonding electrons)=1

  33. s*1s s*1s s*2s 2s 1s 2s 1s energy s1s s2s 1s 1s He He s1s He2 Homonuclear Diatomic Molecule • Diatomics made of the same element • Use energy diagram to determine the bonding • Example : He2(4 e-) and Li2+(5 e-) Bond order = 1/2 (2- 2)=0 Bond order = 1/2 (3- 2)=1/2 e- config= (s1s)2(s*1s)2 e- config= (s1s)2(s*1s)2(s*2s)1

  34. s*2p p*2p p*2p s*2p s2p p*2p p*2p 2px 2py 2pz 2px 2py 2pz p2p p2p 2px 2py 2pz 2px 2py 2pz p2p p2p s2p s*2s 2s 2s s2s s*2s 2s 2s s*1s s2s 1s 1s s1s s*1s 1s 1s s1s Molecular Orbital Diagram O2 N2 Bond order = 1/2 (10- 4)=1/2(6-0)=3 Triple bond, consistent with Lewis structure Bond order = 1/2 (10- 6)=1/2(6-2)=2

  35. Paramagnetism • Paramagnetic • A molecule has unpaired electrons • Align in magnetic field • Example: O2 • Diamagnetic • No unpaired electrons • Example: N2

  36. s*2p p*2p p*2p 2px 2py 2pz 2px 2py 2pz p2p p2p s2px s*2s 2s s2s 2s s*1s 1s s1s 1s Heteronuclear Diatomic Molecule • Heteronuclear - between two nonidentical atoms • Complexity - atomic energy levels are not the same • Bond order = (no. of bonding e- - no. of antibonding e-)/2 • Example : NO

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