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Molecular Geometry: Shape Determines Function

9. Molecular Geometry: Shape Determines Function. Chapter Outline. 9.1 Molecular Shape 9.2 Valence-Shell Electron-Pair Repulsion Theory (VSEPR) 9.3 Polar Bonds and Polar Molecules 9.4 Valence Bond Theory 9.5 Shape and Interactions with Large Molecules

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Molecular Geometry: Shape Determines Function

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  1. 9 Molecular Geometry: Shape Determines Function

  2. Chapter Outline 9.1 Molecular Shape 9.2 Valence-Shell Electron-Pair Repulsion Theory (VSEPR) 9.3 Polar Bonds and Polar Molecules 9.4 Valence Bond Theory 9.5 Shape and Interactions with Large Molecules 9.6 Chirality and Molecular Recognition 9.7 Molecular Orbital Theory

  3. Molecular Shape • Chemical/physical properties are related to molecular shape. • Lewis structures • Show atoms and bonds, but not spatial orientations (3D). • Molecular models • Show orientations and bond angles; help us understand physicochemical properties.

  4. Lewis Structures vs. Models

  5. Bond angle: Angle (in degrees) defined by lines joining the centers of two atoms to the center of a third atom to which they are covalently bonded Not always predictable from Lewis structures Molecular Shape (cont.)

  6. Chapter Outline • 9.1 Molecular Shape • 9.2 Valence-Shell Electron-Pair Repulsion Theory (VSEPR) • Central Atoms with No Lone Pairs • Central Atoms with Lone Pairs • 9.3 Polar Bonds and Polar Molecules • 9.4 Valence Bond Theory • 9.5 Shape and Interactions with Large Molecules • 9.6 Chirality and Molecular Recognition • 9.7 Molecular Orbital Theory

  7. Valence-Shell Electron-Pair Repulsion Theory (VSEPR) • VSEPR Theory: • A model predicting that the arrangement of valence electron pairs around a central atom minimizes repulsion to produce the lowest-energy orientation. • Electron-pair geometry: • Three-dimensional arrangement of bonding e– pairs and lone pairs electrons about a central atom • Molecular geometry: • 3-dimensional arrangement of atoms in a molecule.

  8. VSEPR: Electron-Pair Geometry • To determine electron-pair geometry: • Draw Lewis structure (see Chapter 8). • From Lewis structure, determine steric number (SN): • Determine optimal spatial arrangement of electron pairs (bonding + nonbonding) to minimize repulsion.

  9. Molecular Geometry: Central Atom with No Lone Pairs • Molecular geometry = electron-pair geometry • Determine steric number (SN): • SN = 2 (two atoms bonded to central atom) • geometry  linear • SN = 3 (three atoms bonded to central atom) • geometry  trigonal planar • SN = 4  tetrahedral • SN = 5  trigonal bipyramidal • SN = 6  octahedral

  10. Central Atoms with No Lone Pairs

  11. Central Atoms with No Lone Pairs (cont.)

  12. Geometric Forms • Examples: • CO2 BF3 CCl4 PF5 SF6

  13. Practice: Molecular Geometry (No Lone Pairs) • Collect and Organize: We are given molecular formulas and asked to predict their molecular geometry. Determine the molecular geometry of: a) H2CO (C is central atom) b) CH4

  14. Practice: Molecular Geometry (No Lone Pairs) • Analyze: We can use the periodic table to determine the number of valence electrons for each atom. From the molecular formula and valence electrons, we can draw Lewis structures. From the Lewis structures we can determine the SN. From the SN we can predict the electron-pair geometry. Since there are no lone pairs, the electron-pair geometry is the same as the molecular geometry. Determine the molecular geometry of: a) H2CO (C is central atom) b) CH4

  15. Practice: Molecular Geometry (No Lone Pairs) • -Solve: H2CO  C is the central atom, with single bonds to each H atom and a double bond to O, SN = 3. Molecular geometry = trigonal planar. • CH4 C is central atom, with single bonds to each H atom, SN = 4. Molecular geometry = tetrahedral. Determine the molecular geometry of: a) H2CO (C is central atom) b) CH4

  16. Practice: Molecular Geometry (No Lone Pairs) • Think About It: The molecular geometries are consistent with VSEPR theory for 3 and 4 electron clouds. It is worth noting that the C atom always has 4 bonds, but a double bond counts as only one electron cloud, resulting in a trigonal planar geometry. Determine the molecular geometry of: a) H2CO (C is central atom) b) CH4

  17. Central Atoms with Lone Pairs • Molecular geometry  electron-pair geometry • Replace bonding pair(s) with lone pair(s). • Example: SO2 (SN = 3) • Three electron pairs (2 bonding + 1 lone pair)

  18. Central Atoms w/ Lone Pairs (cont.) • Bond angles less than predicted • Electron pair repulsion! • Lone pair–lone pair = greatest repulsion. • Lone pair–bonding pair • Bonding pair–bonding pair = least repulsion. • Multiple bonds > single bonds

  19. Molecular Geometry: SN = 4 Note: bond angles decrease as # of lone pairs increases.

  20. Molecular Geometry: SN = 4 (cont.) Two lone pairs = greater repulsion, decreased bond angle.

  21. Molecular Geometry: SN = 5 Note: lone pairs occupy equatorial positions.

  22. Molecular Geometry: SN = 6 Note: bond angles = 90 (geometries w/ more than 2 lone pairs are possible.)

  23. Practice: Molecular Geometry • Collect and Organize: We are given the molecular formulas for two polyatomic ions and asked to predict the molecular geometries. What are the molecular geometries of the ions: SCN– and NO2– ?

  24. Practice: Molecular Geometry • Analyze:We can use the periodic table to determine the number of valence electrons for each atom. From the molecular formula and valence electrons, we can draw Lewis structures. From the Lewis structures, we can determine the SN. From the SN, we can predict the electron-pair geometry. Making note of the number of bonding pairs and lone pairs, we can identify the molecular geometry. What are the molecular geometries of the ions: SCN– and NO2– ?

  25. Practice: Molecular Geometry • Solve: • SCN– As the least electronegative element and the one with the greatest bonding capacity, C is the central atom. Although there are several possible resonance structures, they all have C with SN = 2 and no lone pairs. Molecular geometry = linear. • NO2– With N as the central atom, the Lewis structure has N with SN = 3 (two bonding pairs and one lone pair). Again, although there are two possible resonance structures, they both have the same SN value. Molecular geometry = bent, with bond angle <120 due to the extra repulsive energy of the lone pair on N. What are the molecular geometries of the ions: SCN– and NO2– ?

  26. Practice: Molecular Geometry • Think About It: It is worth noting that both molecular structures have two bonding electron clouds, but different molecular geometries due to the differences in steric number. What are the molecular geometries of the ions: SCN– and NO2– ?

  27. Chapter Outline • 9.1 Molecular Shape • 9.2 Valence-Shell Electron-Pair Repulsion Theory (VSEPR) • 9.3 Polar Bonds and Polar Molecules • What Makes a Molecule Polar? • Dipole Moments • 9.4 Valence Bond Theory • 9.5 Shape and Interactions with Large Molecules • 9.6 Chirality and Molecular Recognition • 9.7 Molecular Orbital Theory

  28. Polar Bonds and Polar Molecules • Requirements for polar molecule: • 1. Polar bonds (i.e. covalent bond between atoms with ΔEN). • 2. Nonuniform distribution of polar bonds.

  29. Polar Molecules (cont.) • Bond dipole: • Separation of electrical charge created when atoms with different EN form a covalent bond. • Polar molecule: • Vectors of bond dipoles whose sum > zero.

  30. Dipole moment () –a quantitative expression of the polarity of a molecule. Units = debyes(D); 1 D = 3.34 × 10–30C·m ) Measuring Polarity

  31. Chapter Outline • 9.1 Molecular Shape • 9.2 Valence-Shell Electron-Pair Repulsion Theory (VSEPR) • 9.3 Polar Bonds and Polar Molecules • 9.4 Valence Bond Theory • Bonds from Orbital Overlap • Hybridization (sp3, sp2, sp, sp3d, sp3d2) • 9.5 Shape and Interactions with Large Molecules • 9.6 Chirality and Molecular Recognition • 9.7 Molecular Orbital Theory

  32. Atomic Orbitals and Bonds Overlap of 1s orbitals • Valence Bond Theory (Linus Pauling) • Quantum mechanics-based model • Covalent bond = overlap of half-filled orbitals • Sigma () bond: • Covalent bond in which the highest electron density lies between the two atoms along the bond axis.

  33. Hybridization: sp3Orbitals Hybridization – the mixing of atomic orbitals to generate new sets of orbitals that form covalent bonds with other atoms.

  34. Tetrahedral Sigma Bonds Overlap of 1s with sp3orbitals Tetrahedral orientation of sp3 hybridized orbitals = tetrahedral molecular geometry

  35. Other sp3 Hybrid Examples Note: lone pairs (non-bonding)

  36. Trigonal Planar: sp2 Hybridization Unhybridized p orbitals form double bonds.

  37. sp2 Hybridization (cont.) pi () bond: a covalent bond in which electron density is greatest around—not along—the bonding axis.

  38. Linear: sp Hybridization Form triple bond (one  and two πbonds).

  39. TrigonalBipyramidal: sp3d Hybridization Formed by mixing one s, one d, and three p orbitals. Example: PF5 – five sigma bonds

  40. Octahedral: sp3d2 Hybridization Formed by mixing one s, two d, and three p orbitals.Example: SF6 – six sigma bonds

  41. Practice: Hybrid Orbitals • Collect and Organize: Note that these are the same molecules for which we determined molecular geometry back in section 9.2. Using Lewis structures and VSEPR, we determined that the electronic geometry around the central carbon in SCN– was linear (SN = 2), and the electronic geometry around the central N atom in NO2– was trigonal planar (SN = 3). What are the hybridizations of the central atoms of the ions: SCN– and NO2– ?

  42. Practice: Hybrid Orbitals • Analyze: From the steric number of the central atoms and valence bond theory, we can determine the hybridization around the central atom based on electron-pair geometry. What are the hybridizations of the central atoms of the ions: SCN– and NO2– ?

  43. Practice: Hybrid Orbitals • Solve: • SCN– is linear (SN = 2), so the hybridization must be sp. • NO2– is trigonal planar (SN=3), so the hybridization must be sp2. What are the hybridizations of the central atoms of the ions: SCN– and NO2– ?

  44. Chapter Outline • 9.1 Molecular Shape • 9.2 Valence-Shell Electron-Pair Repulsion Theory (VSEPR) • 9.3 Polar Bonds and Polar Molecules • 9.4 Valence Bond Theory • 9.5 Shape and Interactions with Large Molecules • Shape and Molecular Recognition • Delocalized Electrons • 9.6 Chirality and Molecular Recognition • 9.7 Molecular Orbital Theory

  45. Molecular Recognition All atoms in same plane. • Molecular recognition: • The process by which molecules interact with other molecules in living tissues to produce a biological effect. • Example: ethylene (ripening agent)

  46. Delocalization of Electrons a) b) Delocalization: spreading of electrons in alternating single and double bonds over three or more atoms in a molecule

  47. Aromatic Compounds • Aromatic compound: • A cyclic, planar compound with delocalized  electrons above and below the plane of the molecule. • Example: polycyclic aromatic hydrocarbons (PAH) • Planar shape may allow intercalation in DNA

  48. Chapter Outline • 9.1 Molecular Shape • 9.2 Valence-Shell Electron-Pair Repulsion Theory (VSEPR) • 9.3 Polar Bonds and Polar Molecules • 9.4 Valence Bond Theory • 9.5 Shape and Interactions with Large Molecules • 9.6 Chirality and Molecular Recognition • Optical Isomerism • 9.7 Molecular Orbital Theory

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