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  • 1. Chemistry: A Molecular Approach, 2nd Ed. Nivaldo Tro Chapter 10 Chemical Bonding II Roy Kennedy Massachusetts Bay Community College Wellesley Hills, MA Tro: Chemistry: A Molecular Approach, 2/e Copyright © 2011 Pearson Education, Inc.
  • 2. Taste • The taste of a food depends on the interaction • • • between the food molecules and taste cells on your tongue The main factors that affect this interaction are the shape of the molecule and charge distribution within the molecule The food molecule must fit snugly into the active site of specialized proteins on the surface of taste cells When this happens, changes in the protein structure cause a nerve signal to transmit Tro: Chemistry: A Molecular Approach, 2/e 2 Copyright © 2011 Pearson Education, Inc.
  • 3. Sugar & Artificial Sweeteners • Sugar molecules fit into the active site of taste cell • • • receptors called Tlr3 receptor proteins When the sugar molecule (the key) enters the active site (the lock), the different subunits of the T1r3 protein split apart This split causes ion channels in the cell membrane to open, resulting in nerve signal transmission Artificial sweeteners also fit into the Tlr3 receptor, sometimes binding to it even stronger than sugar  making them “sweeter” than sugar Tro: Chemistry: A Molecular Approach, 2/e 3 Copyright © 2011 Pearson Education, Inc.
  • 4. Structure Determines Properties! • Properties of molecular substances depend on • the structure of the molecule The structure includes many factors, such as:  the skeletal arrangement of the atoms  the kind of bonding between the atoms ionic, polar covalent, or covalent  the shape of the molecule • Bonding theory should allow you to predict the shapes of molecules Tro: Chemistry: A Molecular Approach, 2/e 4 Copyright © 2011 Pearson Education, Inc.
  • 5. Molecular Geometry • Molecules are 3-dimensional objects • We often describe the shape of a molecule • • with terms that relate to geometric figures These geometric figures have characteristic “corners” that indicate the positions of the surrounding atoms around a central atom in the center of the geometric figure The geometric figures also have characteristic angles that we call bond angles Tro: Chemistry: A Molecular Approach, 2/e 5 Copyright © 2011 Pearson Education, Inc.
  • 6. Lewis Theory Predicts Electron Groups • Lewis theory predicts there are regions of • • electrons in an atom Some regions result from placing shared pairs of valence electrons between bonding nuclei Other regions result from placing unshared valence electrons on a single nuclei Tro: Chemistry: A Molecular Approach, 2/e 6 Copyright © 2011 Pearson Education, Inc.
  • 7. Using Lewis Theory to Predict Molecular Shapes • Lewis theory says that these regions of electron groups should repel each other  because they are regions of negative charge • This idea can then be extended to predict the shapes of molecules  the position of atoms surrounding a central atom will be determined by where the bonding electron groups are  the positions of the electron groups will be determined by trying to minimize repulsions between them Tro: Chemistry: A Molecular Approach, 2/e 7 Copyright © 2011 Pearson Education, Inc.
  • 8. VSEPR Theory • Electron groups around the central atom will be most stable when they are as far apart as possible – we call this valence shell electron pair repulsion theory  because electrons are negatively charged, they should be most stable when they are separated as much as possible • The resulting geometric arrangement will allow us to predict the shapes and bond angles in the molecule Tro: Chemistry: A Molecular Approach, 2/e 8 Copyright © 2011 Pearson Education, Inc.
  • 9. Tro: Chemistry: A Molecular Approach, 2/e 9 Copyright © 2011 Pearson Education, Inc.
  • 10. Electron Groups • The Lewis structure predicts the number of valence • • electron pairs around the central atom(s) Each lone pair of electrons constitutes one electron group on a central atom Each bond constitutes one electron group on a central atom  regardless of whether it is single, double, or triple •• •O • •• N •• O• • •• Tro: Chemistry: A Molecular Approach, 2/e there are three electron groups on N three lone pair one single bond one double bond 10 Copyright © 2011 Pearson Education, Inc.
  • 11. Electron Group Geometry • There are five basic arrangements of electron groups around a central atom  based on a maximum of six bonding electron groups  though there may be more than six on very large atoms, it is very rare • Each of these five basic arrangements results in five different basic electron geometries  in order for the molecular shape and bond angles to be a “perfect” geometric figure, all the electron groups must be bonds and all the bonds must be equivalent • For molecules that exhibit resonance, it doesn’t matter which resonance form you use – the electron geometry will be the same Tro: Chemistry: A Molecular Approach, 2/e 11 Copyright © 2011 Pearson Education, Inc.
  • 12. Linear Electron Geometry • When there are two electron groups around the • • central atom, they will occupy positions on opposite sides of the central atom This results in the electron groups taking a linear geometry The bond angle is 180° Tro: Chemistry: A Molecular Approach, 2/e 12 Copyright © 2011 Pearson Education, Inc.
  • 13. Linear Geometry Tro: Chemistry: A Molecular Approach, 2/e 13 Copyright © 2011 Pearson Education, Inc.
  • 14. Trigonal Planar Electron Geometry • When there are three electron groups around • • the central atom, they will occupy positions in the shape of a triangle around the central atom This results in the electron groups taking a trigonal planar geometry The bond angle is 120° Tro: Chemistry: A Molecular Approach, 2/e 14 Copyright © 2011 Pearson Education, Inc.
  • 15. Trigonal Geometry Tro: Chemistry: A Molecular Approach, 2/e 15 Copyright © 2011 Pearson Education, Inc.
  • 16. Tetrahedral Electron Geometry • When there are four electron groups around the • • central atom, they will occupy positions in the shape of a tetrahedron around the central atom This results in the electron groups taking a tetrahedral geometry The bond angle is 109.5° Tro: Chemistry: A Molecular Approach, 2/e 16 Copyright © 2011 Pearson Education, Inc.
  • 17. Tetrahedral Geometry Tro: Chemistry: A Molecular Approach, 2/e 17 Copyright © 2011 Pearson Education, Inc.
  • 18. Trigonal Bipyramidal Electron Geometry • When there are five electron groups around the central • • • • • atom, they will occupy positions in the shape of two tetrahedra that are base-to-base with the central atom in the center of the shared bases This results in the electron groups taking a trigonal bipyramidal geometry The positions above and below the central atom are called the axial positions The positions in the same base plane as the central atom are called the equatorial positions The bond angle between equatorial positions is 120° The bond angle between axial and equatorial positions is 90° Tro: Chemistry: A Molecular Approach, 2/e 18 Copyright © 2011 Pearson Education, Inc.
  • 19. Trigonal Bipyramid Tro: Chemistry: A Molecular Approach, 2/e 19 Copyright © 2011 Pearson Education, Inc.
  • 20. Trigonal Bipyramidal Geometry Tro: Chemistry: A Molecular Approach, 2/e 20 Copyright © 2011 Pearson Education, Inc.
  • 21. Tro: Chemistry: A Molecular Approach, 2/e 21 Copyright © 2011 Pearson Education, Inc.
  • 22. Octahedral Electron Geometry • When there are six electron groups around the central • atom, they will occupy positions in the shape of two square-base pyramids that are base-to-base with the central atom in the center of the shared bases This results in the electron groups taking an octahedral geometry  it is called octahedral because the geometric figure has eight sides • All positions are equivalent • The bond angle is 90° Tro: Chemistry: A Molecular Approach, 2/e 22 Copyright © 2011 Pearson Education, Inc.
  • 23. Octahedral Geometry Tro: Chemistry: A Molecular Approach, 2/e 23 Copyright © 2011 Pearson Education, Inc.
  • 24. Octahedral Geometry Tro: Chemistry: A Molecular Approach, 2/e 24 Copyright © 2011 Pearson Education, Inc.
  • 25. Tro: Chemistry: A Molecular Approach, 2/e 25 Copyright © 2011 Pearson Education, Inc.
  • 26. Molecular Geometry • The actual geometry of the molecule may be • • different from the electron geometry When the electron groups are attached to atoms of different size, or when the bonding to one atom is different than the bonding to another, this will affect the molecular geometry around the central atom Lone pairs also affect the molecular geometry  they occupy space on the central atom, but are not “seen” as points on the molecular geometry Tro: Chemistry: A Molecular Approach, 2/e 26 Copyright © 2011 Pearson Education, Inc.
  • 27. Not Quite Perfect Geometry Because the bonds and atom sizes are not identical in formaldehyde, the observed angles are slightly different from ideal Tro: Chemistry: A Molecular Approach, 2/e 27 Copyright © 2011 Pearson Education, Inc.
  • 28. The Effect of Lone Pairs • Lone pair groups “occupy more space” on the central atom  because their electron density is exclusively on the central atom rather than shared like bonding electron groups • Relative sizes of repulsive force interactions is Lone Pair – Lone Pair > Lone Pair – Bonding Pair > Bonding Pair – Bonding Pair • This affects the bond angles, making the bonding pair – bonding pair angles smaller than expected Tro: Chemistry: A Molecular Approach, 2/e 28 Copyright © 2011 Pearson Education, Inc.
  • 29. Effect of Lone Pairs The bonding electrons are are localized on nonbonding electrons shared by two atoms, so some of the negative charge is the central atom, so area of negative charge removed from the central atom takes more space Tro: Chemistry: A Molecular Approach, 2/e 29 Copyright © 2011 Pearson Education, Inc.
  • 30. Bond Angle Distortion from Lone Pairs Tro: Chemistry: A Molecular Approach, 2/e 30 Copyright © 2011 Pearson Education, Inc.
  • 31. Bond Angle Distortion from Lone Pairs Tro: Chemistry: A Molecular Approach, 2/e 31 Copyright © 2011 Pearson Education, Inc.
  • 32. Bent Molecular Geometry: Derivative of Trigonal Planar Electron Geometry • When there are three electron groups around • the central atom, and one of them is a lone pair, the resulting shape of the molecule is called a trigonal planar — bent shape The bond angle is less than 120°  because the lone pair takes up more space Tro: Chemistry: A Molecular Approach, 2/e 32 Copyright © 2011 Pearson Education, Inc.
  • 33. Pyramidal & Bent Molecular Geometries: Derivatives of Tetrahedral Electron Geometry • When there are four electron groups around the central atom, and one is a lone pair, the result is called a pyramidal shape  because it is a triangular-base pyramid with the central atom at the apex • When there are four electron groups around the central atom, and two are lone pairs, the result is called a tetrahedral—bent shape  it is planar  it looks similar to the trigonal planar—bent shape, except the angles are smaller • For both shapes, the bond angle is less than 109.5° Tro: Chemistry: A Molecular Approach, 2/e 33 Copyright © 2011 Pearson Education, Inc.
  • 34. Methane Tro: Chemistry: A Molecular Approach, 2/e 34 Copyright © 2011 Pearson Education, Inc.
  • 35. Pyramidal Shape Tro: Chemistry: A Molecular Approach, 2/e 35 Copyright © 2011 Pearson Education, Inc.
  • 36. Pyramidal Shape Tro: Chemistry: A Molecular Approach, 2/e 36 Copyright © 2011 Pearson Education, Inc.
  • 37. Tetrahedral–Bent Shape Tro: Chemistry: A Molecular Approach, 2/e 37 Copyright © 2011 Pearson Education, Inc.
  • 38. Tetrahedral–Bent Shape Tro: Chemistry: A Molecular Approach, 2/e 38 Copyright © 2011 Pearson Education, Inc.
  • 39. Derivatives of the Trigonal Bipyramidal Electron Geometry • When there are five electron groups around the central atom, • and some are lone pairs, they will occupy the equatorial positions because there is more room When there are five electron groups around the central atom, and one is a lone pair, the result is called the seesaw shape  aka distorted tetrahedron • When there are five electron groups around the central atom, • • • and two are lone pairs, the result is called the T-shaped When there are five electron groups around the central atom, and three are lone pairs, the result is a linear shape The bond angles between equatorial positions are less than 120° The bond angles between axial and equatorial positions are less than 90°  linear = 180° axial–to–axial Tro: Chemistry: A Molecular Approach, 2/e 39 Copyright © 2011 Pearson Education, Inc.
  • 40. Replacing Atoms with Lone Pairs in the Trigonal Bipyramid System Tro: Chemistry: A Molecular Approach, 2/e 40 Copyright © 2011 Pearson Education, Inc.
  • 41. Seesaw Shape Tro: Chemistry: A Molecular Approach, 2/e 41 Copyright © 2011 Pearson Education, Inc.
  • 42. T–Shape Tro: Chemistry: A Molecular Approach, 2/e 42 Copyright © 2011 Pearson Education, Inc.
  • 43. T–Shape Tro: Chemistry: A Molecular Approach, 2/e 43 Copyright © 2011 Pearson Education, Inc.
  • 44. Linear Shape Tro: Chemistry: A Molecular Approach, 2/e 44 Copyright © 2011 Pearson Education, Inc.
  • 45. Derivatives of the Octahedral Geometry • When there are six electron groups around the • central atom, and some are lone pairs, each even number lone pair will take a position opposite the previous lone pair When there are six electron groups around the central atom, and one is a lone pair, the result is called a square pyramid shape  the bond angles between axial and equatorial positions is less than 90° • When there are six electron groups around the central atom, and two are lone pairs, the result is called a square planar shape  the bond angles between equatorial positions is 90° Tro: Chemistry: A Molecular Approach, 2/e 45 Copyright © 2011 Pearson Education, Inc.
  • 46. Square Pyramidal Shape Tro: Chemistry: A Molecular Approach, 2/e 46 Copyright © 2011 Pearson Education, Inc.
  • 47. Square Planar Shape Tro: Chemistry: A Molecular Approach, 2/e 47 Copyright © 2011 Pearson Education, Inc.
  • 48. Tro: Chemistry: A Molecular Approach, 2/e 48 Copyright © 2011 Pearson Education, Inc.
  • 49. Predicting the Shapes Around Central Atoms 1. Draw the Lewis structure 2. Determine the number of electron groups around the central atom 3. Classify each electron group as bonding or lone pair, and count each type  remember, multiple bonds count as one group 4. Use Table 10.1 to determine the shape and bond angles Tro: Chemistry: A Molecular Approach, 2/e 49 Copyright © 2011 Pearson Education, Inc.
  • 50. Example 10.2: Predict the geometry and bond angles of PCl3 1. Draw the Lewis structure a) 26 valence electrons 2. Determine the Number of electron groups around central atom a) four electron groups around P Tro: Chemistry: A Molecular Approach, 2/e 50 Copyright © 2011 Pearson Education, Inc.
  • 51. Example 10.2: Predict the geometry and bond angles of PCl3 3. Classify the electron groups a) three bonding groups b) one lone pair 4. Use Table 10.1 to determine the shape and bond angles a) four electron groups around P = tetrahedral electron geometry b) three bonding + one lone pair = trigonal pyramidal molecular geometry c) trigonal pyramidal = bond angles less than 109.5° Tro: Chemistry: A Molecular Approach, 2/e 51 Copyright © 2011 Pearson Education, Inc.
  • 52. Practice – Predict the molecular geometry and bond angles in SiF5− Tro: Chemistry: A Molecular Approach, 2/e 52 Copyright © 2011 Pearson Education, Inc.
  • 53. Practice – Predict the molecular geometry and bond angles in SiF5─ Si least electronegative 5 electron groups on Si Si is central atom 5 bonding groups 0 lone pairs Si = 4e─ F5 = 5(7e─) = 35e─ (─) = 1e─ total = 40e─ Shape = trigonal bipyramid Bond angles Feq–Si–Feq = 120° Feq–Si–Fax = 90° Tro: Chemistry: A Molecular Approach, 2/e 53 Copyright © 2011 Pearson Education, Inc.
  • 54. Practice – Predict the molecular geometry and bond angles in ClO2F Tro: Chemistry: A Molecular Approach, 2/e 54 Copyright © 2011 Pearson Education, Inc.
  • 55. Practice – Predict the molecular geometry and bond angles in ClO2F Cl least electronegative 4 electron groups on Cl Cl is central atom 3 bonding groups 1 lone pair Cl = 7e─ O2 = 2(6e─) = 12e─ F = 7e─ Total = 26e─ Shape = trigonal pyramidal Bond angles O–Cl–O < 109.5° O–Cl–F < 109.5° Tro: Chemistry: A Molecular Approach, 2/e 55 Copyright © 2011 Pearson Education, Inc.
  • 56. Representing 3-Dimensional Shapes on a 2-Dimensional Surface • One of the problems with drawing molecules is • • • • trying to show their dimensionality By convention, the central atom is put in the plane of the paper Put as many other atoms as possible in the same plane and indicate with a straight line For atoms in front of the plane, use a solid wedge For atoms behind the plane, use a hashed wedge Tro: Chemistry: A Molecular Approach, 2/e 56 Copyright © 2011 Pearson Education, Inc.
  • 57. Tro: Chemistry: A Molecular Approach, 2/e 57 Copyright © 2011 Pearson Education, Inc.
  • 58. SF6 F F S F F F F Tro: Chemistry: A Molecular Approach, 2/e 58 Copyright © 2011 Pearson Education, Inc.
  • 59. Multiple Central Atoms • Many molecules have larger structures with many • • interior atoms We can think of them as having multiple central atoms When this occurs, we describe the shape around each central atom in sequence • • shape around left C is tetrahedral shape around center C is trigonal planar shape around right O is tetrahedral-bent Tro: Chemistry: A Molecular Approach, 2/e 59 Η Ο • • | || • • Η − Χ − Χ − Ο − Η | • • Η Copyright © 2011 Pearson Education, Inc.
  • 60. Describing the Geometry of Methanol Tro: Chemistry: A Molecular Approach, 2/e 60 Copyright © 2011 Pearson Education, Inc.
  • 61. Describing the Geometry of Glycine Tro: Chemistry: A Molecular Approach, 2/e 61 Copyright © 2011 Pearson Education, Inc.
  • 62. Practice – Predict the molecular geometries in H3BO3 Tro: Chemistry: A Molecular Approach, 2/e 62 Copyright © 2011 Pearson Education, Inc.
  • 63. Practice – Predict the molecular geometries in H3BO3 oxyacid, so H attached to O 3 electron groups on B 4 electron groups on O B least electronegative O has B has 2 3 bonding groups 2 lone pairs 0 ponepairs B Is Central Atom B = 3e─ O3 = 3(6e─) = 18e─ H3 = 3(1e─) = 3e─ Total = 24e─ Shape on B = trigonal planar Shape on O = tetrahedral bent Tro: Chemistry: A Molecular Approach, 2/e 63 Copyright © 2011 Pearson Education, Inc.
  • 64. Polarity of Molecules • For a molecule to be polar it must 1. have polar bonds   electronegativity difference - theory bond dipole moments - measured 2. have an unsymmetrical shape  vector addition • Polarity affects the intermolecular forces of attraction  therefore boiling points and solubilities  like dissolves like • Nonbonding pairs affect molecular polarity, strong pull in its direction Tro: Chemistry: A Molecular Approach, 2/e 64 Copyright © 2011 Pearson Education, Inc.
  • 65. Molecule Polarity The H─Cl bond is polar. The bonding electrons are pulled toward the Cl end of the molecule. The net result is a polar molecule. Tro: Chemistry: A Molecular Approach, 2/e 65 Copyright © 2011 Pearson Education, Inc.
  • 66. Vector Addition Tro: Chemistry: A Molecular Approach, 2/e 66 Copyright © 2011 Pearson Education, Inc.
  • 67. Tro: Chemistry: A Molecular Approach, 2/e 67 Copyright © 2011 Pearson Education, Inc.
  • 68. Molecule Polarity The O─C bond is polar. The bonding electrons are pulled equally toward both O ends of the molecule. The net result is a nonpolar molecule. Tro: Chemistry: A Molecular Approach, 2/e 68 Copyright © 2011 Pearson Education, Inc.
  • 69. Molecule Polarity The H─O bond is polar. Both sets of bonding electrons are pulled toward the O end of the molecule. The net result is a polar molecule. Tro: Chemistry: A Molecular Approach, 2/e 69 Copyright © 2011 Pearson Education, Inc.
  • 70. Predicting Polarity of Molecules 1. Draw the Lewis structure and determine the molecular geometry 2. Determine whether the bonds in the molecule are polar a) if there are not polar bonds, the molecule is nonpolar 3. Determine whether the polar bonds add together to give a net dipole moment Tro: Chemistry: A Molecular Approach, 2/e 70 Copyright © 2011 Pearson Education, Inc.
  • 71. Example 10.5: Predict whether NH3 is a polar molecule 1. Draw the Lewis structure and determine the molecular geometry a) eight valence electrons b) three bonding + one lone pair = trigonal pyramidal molecular geometry Tro: Chemistry: A Molecular Approach, 2/e 71 Copyright © 2011 Pearson Education, Inc.
  • 72. Example 10.5: Predict whether NH3 is a polar molecule 2. Determine if the bonds are polar a) electronegativity difference b) if the bonds are not polar, we can stop here and declare the molecule will be nonpolar Tro: Chemistry: A Molecular Approach, 2/e 72 ENN = 3.0 ENH = 2.1 3.0 − 2.1 = 0.9 therefore the bonds are polar covalent Copyright © 2011 Pearson Education, Inc.
  • 73. Example 10.5: Predict whether NH3 is a polar molecule 3) Determine whether the polar bonds add together to give a net dipole moment a) vector addition b) generally, asymmetric shapes result in uncompensated polarities and a net dipole moment Tro: Chemistry: A Molecular Approach, 2/e 73 The H─N bond is polar. All the sets of bonding electrons are pulled toward the N end of the molecule. The net result is a polar molecule. Copyright © 2011 Pearson Education, Inc.
  • 74. Practice – Decide whether the following molecules are polar EN O = 3.5 N = 3.0 Cl = 3.0 S = 2.5 Tro: Chemistry: A Molecular Approach, 2/e 74 Copyright © 2011 Pearson Education, Inc.
  • 75. Practice – Decide whether the following molecules Are polar Trigonal Bent Trigonal Planar 2.5 1. polar bonds, N-O 2. asymmetrical shape 1. polar bonds, all S-O 2. symmetrical shape nonpolar polar Tro: Chemistry: A Molecular Approach, 2/e 75 Copyright © 2011 Pearson Education, Inc.
  • 76. Molecular Polarity Affects Solubility in Water • Polar molecules are attracted • to other polar molecules Because water is a polar molecule, other polar molecules dissolve well in water  and ionic compounds as well • Some molecules have both polar and nonpolar parts Tro: Chemistry: A Molecular Approach, 2/e 76 Copyright © 2011 Pearson Education, Inc.
  • 77. Problems with Lewis Theory • Lewis theory generally predicts trends in properties, but does not give good numerical predictions  e.g. bond strength and bond length • Lewis theory gives good first approximations of • • the bond angles in molecules, but 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  e.g. O2 is paramagnetic, though the Lewis structure predicts it is diamagnetic Tro: Chemistry: A Molecular Approach, 2/e 77 Copyright © 2011 Pearson Education, Inc.
  • 78. Valence Bond Theory • Linus Pauling and others applied the principles • • of quantum mechanics to molecules They reasoned that bonds between atoms would occur when the orbitals on those atoms interacted to make a bond The kind of interaction depends on whether the orbitals align along the axis between the nuclei, or outside the axis Tro: Chemistry: A Molecular Approach, 2/e 78 Copyright © 2011 Pearson Education, Inc.
  • 79. Orbital Interaction • As two atoms approached, the half-filled valence atomic orbitals on each atom would interact to form molecular orbitals  molecular orbtials are regions of high probability of finding the shared electrons in the molecule • The molecular orbitals would be more stable than the separate atomic orbitals because they would contain paired electrons shared by both atoms  the potential energy is lowered when the molecular orbitals contain a total of two paired electrons compared to separate one electron atomic orbitals Tro: Chemistry: A Molecular Approach, 2/e 79 Copyright © 2011 Pearson Education, Inc.
  • 80. Orbital Diagram for the Formation of H2S H 1s ↑↓ ↑ + ↑↓ 1s ↑ 3s ↑ ↑ ↑↓ S 3p H─S bond ↑↓ H─S bond H Predicts bond angle = 90° Actual bond angle = 92° Tro: Chemistry: A Molecular Approach, 2/e 80 Copyright © 2011 Pearson Education, Inc.
  • 81. Valence Bond Theory – Hybridization • One of the issues that arises is that the number of partially filled or empty atomic orbitals did not predict the number of bonds or orientation of bonds C = 2s22px12py12pz0 would predict two or three bonds that are 90° apart, rather than four bonds that are 109.5° apart • 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 2s and 2p orbitals to get four orbitals that point at the corners of a tetrahedron Tro: Chemistry: A Molecular Approach, 2/e 81 Copyright © 2011 Pearson Education, Inc.
  • 82. Unhybridized C Orbitals Predict the Wrong Bonding & Geometry Tro: Chemistry: A Molecular Approach, 2/e 82 Copyright © 2011 Pearson Education, Inc.
  • 83. Valence Bond Theory Main Concepts 1. The valence electrons of the atoms in a molecule reside in quantum-mechanical atomic orbitals. The orbitals can be the standard s, p, d, and f orbitals, or they 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 Tro: Chemistry: A Molecular Approach, 2/e 83 Copyright © 2011 Pearson Education, Inc.
  • 84. Hybridization • Some atoms hybridize their orbitals to • maximize bonding • more bonds = more full orbitals = more stability Hybridizing is mixing different types of orbitals in the valence shell to make a new set of degenerate orbitals  sp, sp2, sp3, sp3d, sp3d2 • Same type of atom can have different types of hybridization  C = sp, sp2, sp3 Tro: Chemistry: A Molecular Approach, 2/e 84 Copyright © 2011 Pearson Education, Inc.
  • 85. Hybrid Orbitals • The number of standard atomic orbitals combined = the number of hybrid orbitals formed  combining a 2s with a 2p gives two 2sp hybrid orbitals  H cannot hybridize!! its valence shell only has one orbital • The number and type of standard atomic orbitals • combined determines the shape of the hybrid orbitals The particular kind of hybridization that occurs is the one that yields the lowest overall energy for the molecule Tro: Chemistry: A Molecular Approach, 2/e 85 Copyright © 2011 Pearson Education, Inc.
  • 86. Carbon Hybridizations Unhybridized ↑↓ ↑ ↑ 2p 2s sp hybridized ↑ ↑ ↑ 2sp ↑ 2p sp2 hybridized ↑ ↑ ↑ 2sp2 sp3 hybridized ↑ ↑ ↑ ↑ 2p ↑ 2sp3 Tro: Chemistry: A Molecular Approach, 2/e 86 Copyright © 2011 Pearson Education, Inc.
  • 87. sp3 Hybridization • Atom with four electron groups around it  tetrahedral geometry  109.5° angles between hybrid orbitals • Atom uses hybrid orbitals for all bonds and lone pairs Tro: Chemistry: A Molecular Approach, 2/e 87 Copyright © 2011 Pearson Education, Inc.
  • 88. Tro: Chemistry: A Molecular Approach, 2/e 88 Copyright © 2011 Pearson Education, Inc.
  • 89. Orbital Diagram of the sp3 Hybridization of C Tro: Chemistry: A Molecular Approach, 2/e 89 Copyright © 2011 Pearson Education, Inc.
  • 90. sp3 Hybridized Atoms Orbital Diagrams • Place electrons into hybrid and unhybridized valence orbitals as if all the orbitals have equal energy • Lone pairs generally occupy hybrid orbitals Unhybridized atom 2s ↑↓ 2s ↑ ↑ 2p C ↑ ↑ ↑ ↑ 2p N ↑ Tro: Chemistry: A Molecular Approach, 2/e 90 ↑ ↑ ↑ 2sp3 ↑ ↑↓ sp3 hybridized atom ↑ ↑ ↑ 2sp3 Copyright © 2011 Pearson Education, Inc.
  • 91. Practice – Draw the orbital diagram for the sp3 hybridization of each atom Unhybridized atom 3s ↑↓ 2s ↑↓ ↑↓ ↑ 3p Cl ↑↓ ↑↓ ↑↓ ↑ 3sp3 ↑↓ ↑ ↑ 2p O ↑ ↑↓ ↑ ↑ 2sp3 Tro: Chemistry: A Molecular Approach, 2/e 91 ↑ ↑↓ sp3 hybridized atom Copyright © 2011 Pearson Education, Inc.
  • 92. 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 either be aligned • along the axis between the atoms, or The orbitals must be parallel to each other and perpendicular to the interatomic axis Tro: Chemistry: A Molecular Approach, 2/e 92 Copyright © 2011 Pearson Education, Inc.
  • 93. Methane Formation with sp3 C Tro: Chemistry: A Molecular Approach, 2/e 93 Copyright © 2011 Pearson Education, Inc.
  • 94. Ammonia Formation with sp3 N Tro: Chemistry: A Molecular Approach, 2/e 94 Copyright © 2011 Pearson Education, Inc.
  • 95. Types of Bonds • A sigma (σ ) bond results when the interacting atomic orbitals point along the axis connecting the two bonding nuclei  either standard atomic orbitals or hybrids  s–to–s, p–to–p, hybrid–to–hybrid, s–to–hybrid, etc. • A pi (π ) bond results when the bonding atomic orbitals are parallel to each other and perpendicular to the axis connecting the two bonding nuclei  between unhybridized parallel p orbitals • The interaction between parallel orbitals is not as strong as between orbitals that point at each other; therefore σ bonds are stronger than π bonds Tro: Chemistry: A Molecular Approach, 2/e 95 Copyright © 2011 Pearson Education, Inc.
  • 96. Tro: Chemistry: A Molecular Approach, 2/e 96 Copyright © 2011 Pearson Education, Inc.
  • 97. Orbital Diagrams of Bonding • “Overlap” between a hybrid orbital on one atom • with a hybrid or nonhybridized orbital on another atom results in a σ bond “Overlap” between unhybridized p orbitals on bonded atoms results in a π bond Tro: Chemistry: A Molecular Approach, 2/e 97 Copyright © 2011 Pearson Education, Inc.
  • 98. CH3NH2 Orbital Diagram H σ σ C ·· σ H H σN σ σ σ ↑ σ ↑ ↑ ↑ 1s H 1s H Tro: Chemistry: A Molecular Approach, 2/e ↑ σ ↑ ↑ σ ↑ ↑ σ 1s H 98 ↑ ↑↓ sp3 N σ ↑ sp C 3 H ↑ H 1s H 1s H Copyright © 2011 Pearson Education, Inc.
  • 99. Formaldehyde, CH2O Orbital Diagram ↑ pC ↑ σ ↑ σ ↑ ↑ pO ↑ ↑↓ ↑↓ sp2 O σ ↑ sp C 2 π ↑ 1s H 1s H Tro: Chemistry: A Molecular Approach, 2/e 99 Copyright © 2011 Pearson Education, Inc.
  • 100. sp2 • Atom with three electron groups around it  trigonal planar system  C = trigonal planar  N = trigonal bent  O = “linear”  120° bond angles  flat • Atom uses hybrid orbitals for σ bonds and lone pairs, uses nonhybridized p orbital for π bond Tro: Chemistry: A Molecular Approach, 2/e 100 Copyright © 2011 Pearson Education, Inc.
  • 101. Tro: Chemistry: A Molecular Approach, 2/e 101 Copyright © 2011 Pearson Education, Inc.
  • 102. sp2 Hybridized Atoms Orbital Diagrams Unhybridized atom 2s ↑↓ 2s ↑ ↑ 2p C 3σ 1π ↑ ↑ ↑ ↑ 2p N 2σ 1π ↑ Tro: Chemistry: A Molecular Approach, 2/e 102 ↑ ↑↓ sp2 hybridized atom ↑ ↑ 2sp2 ↑ 2p ↑ ↑ 2sp2 ↑ 2p Copyright © 2011 Pearson Education, Inc.
  • 103. Practice – Draw the orbital diagram for the sp2 hybridization of each atom. How many σ and π bonds would you expect each to form? ↑↓ 2s ↑↓ 2s ↑ 2p ↑↓ ↑ ↑ 2p Tro: Chemistry: A Molecular Approach, 2/e sp2 hybridized atom B 3σ 0π ↑ ↑ ↑ 2sp2 2p O 1σ 1π ↑ ↑↓ ↑ 2sp2 ↑ 2p 103 ↑ Unhybridized atom Copyright © 2011 Pearson Education, Inc.
  • 104. Hybrid orbitals overlap to form a σ bond. Unhybridized p orbitals overlap to form a π bond. Tro: Chemistry: A Molecular Approach, 2/e 104 Copyright © 2011 Pearson Education, Inc.
  • 105. CH2NH Orbital Diagram ↑ σ ↑ σ ↑ σ ↑ sp C H C H pN ↑ ・・ N ↑ ↑↓ sp2 N σ H ↑ 1s H 1s H Tro: Chemistry: A Molecular Approach, 2/e ↑ ↑ ↑ pC 2 π 1s H 105 Copyright © 2011 Pearson Education, Inc.
  • 106. Bond Rotation • Because the orbitals that form the σ bond point • along the internuclear axis, rotation around that bond does not require breaking the interaction between the orbitals But the orbitals that form the π bond interact above and below the internuclear axis, so rotation around the axis requires the breaking of the interaction between the orbitals Tro: Chemistry: A Molecular Approach, 2/e 106 Copyright © 2011 Pearson Education, Inc.
  • 107. Tro: Chemistry: A Molecular Approach, 2/e 107 Copyright © 2011 Pearson Education, Inc.
  • 108. Tro: Chemistry: A Molecular Approach, 2/e 108 Copyright © 2011 Pearson Education, Inc.
  • 109. sp • Atom with two electron groups  linear shape  180° bond angle • Atom uses hybrid orbitals for σ bonds or lone pairs, uses nonhybridized p orbitals for π bonds π σ π Tro: Chemistry: A Molecular Approach, 2/e 109 Copyright © 2011 Pearson Education, Inc.
  • 110. Tro: Chemistry: A Molecular Approach, 2/e 110 Copyright © 2011 Pearson Education, Inc.
  • 111. Tro: Chemistry: A Molecular Approach, 2/e 111 Copyright © 2011 Pearson Education, Inc.
  • 112. sp Hybridized Atoms Orbital Diagrams ↑↓ 2s ↑↓ 2s ↑ ↑ 2p ↑ ↑ ↑ 2p Tro: Chemistry: A Molecular Approach, 2/e C 2σ 2π N 1σ 2π 112 sp hybridized atom ↑ ↑ 2sp ↑ Unhybridized atom ↑ ↑ 2sp ↑ ↑ 2p ↑ ↑ 2p Copyright © 2011 Pearson Education, Inc.
  • 113. HCN Orbital Diagram ↑ pC sp C ↑ ↑ 2π σ ↑ ↑ ↑ pN ↑ ↑↓ sp N s ↑ 1s H Tro: Chemistry: A Molecular Approach, 2/e 113 Copyright © 2011 Pearson Education, Inc.
  • 114. sp d 3 • Atom with five electron groups around it  trigonal bipyramid electron geometry  Seesaw, T–Shape, Linear  120° & 90° bond angles • Use empty d orbitals from • valence shell d orbitals can be used to make π bonds Tro: Chemistry: A Molecular Approach, 2/e 114 Copyright © 2011 Pearson Education, Inc.
  • 115. Tro: Chemistry: A Molecular Approach, 2/e 115 Copyright © 2011 Pearson Education, Inc.
  • 116. sp3d Hybridized Atoms Orbital Diagrams Unhybridized atom ↑↓ 3s ↑↓ 3s ↑ ↑ ↑ 3p ↑↓ ↑ ↑ 3p sp3d hybridized atom P 3d ↑ ↑ ↑ ↑ 3sp3d S ↑↓ ↑ 3d ↑ ↑ ↑ ↑ 3sp3d (non-hybridizing d orbitals not shown) Tro: Chemistry: A Molecular Approach, 2/e 116 Copyright © 2011 Pearson Education, Inc.
  • 117. SOF4 Orbital Diagram ↑ dS ↑ ↑ σ σ ↑ pO σ ↑ ↑ σ ↑ σ ↑ ↑ ↑↓ ↑↓ sp2 O ↑ sp d S ↑ 3 π ↑ 2p F 2p F 2p F 2p F Tro: Chemistry: A Molecular Approach, 2/e 117 Copyright © 2011 Pearson Education, Inc.
  • 118. sp3d2 • Atom with six electron groups around it  octahedral electron geometry  Square Pyramid, Square Planar  90° bond angles • Use empty d orbitals from • valence shell to form hybrid d orbitals can be used to make π bonds Tro: Chemistry: A Molecular Approach, 2/e 118 Copyright © 2011 Pearson Education, Inc.
  • 119. Tro: Chemistry: A Molecular Approach, 2/e 119 Copyright © 2011 Pearson Education, Inc.
  • 120. sp3d2 Hybridized Atoms Orbital Diagrams Unhybridized atom ↑↓ ↑↓ ↑ ↑ 3s 3p ↑↓ ↑↓ ↑↓ ↑ 5s 5p sp3d2 hybridized atom S 3d ↑ ↑ ↑ ↑ ↑ ↑ 3sp3d2 I 5d ↑↓ ↑ ↑ ↑ ↑ ↑ 5sp3d2 (non-hybridizing d orbitals not shown) Tro: Chemistry: A Molecular Approach, 2/e 120 Copyright © 2011 Pearson Education, Inc.
  • 121. Tro: Chemistry: A Molecular Approach, 2/e 121 Copyright © 2011 Pearson Education, Inc.
  • 122. 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 10.3 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 σ or π Tro: Chemistry: A Molecular Approach, 2/e 122 Copyright © 2011 Pearson Education, Inc.
  • 123. Example 10.7: Predict the hybridization and bonding scheme for CH3CHO Draw the Lewis structure Predict the electron group C1 = 4 electron areas geometry around inside ∴ C1= tetrahedral atoms C2 = 3 electron areas ∴ C2 = trigonal planar Tro: Chemistry: A Molecular Approach, 2/e 123 Copyright © 2011 Pearson Education, Inc.
  • 124. Example 10.7: Predict the hybridization and bonding scheme for CH3CHO Determine the hybridization of the interior atoms C1 = tetrahedral ∴ C1 = sp3 C2 = trigonal planar ∴ C2 = sp2 Sketch the molecule and orbitals Tro: Chemistry: A Molecular Approach, 2/e 124 Copyright © 2011 Pearson Education, Inc.
  • 125. Example 10.7: Predict the hybridization and bonding scheme for CH3CHO Label the bonds Tro: Chemistry: A Molecular Approach, 2/e 125 Copyright © 2011 Pearson Education, Inc.
  • 126. Practice – Predict the hybridization of all the atoms in H3BO3 H = can’t hybridize B = 3 electron groups = sp2 O = 4 electron groups = sp3 Tro: Chemistry: A Molecular Approach, 2/e 126 Copyright © 2011 Pearson Education, Inc.
  • 127. Practice – Predict the hybridization and bonding scheme of all the atoms in NClO •• •O • •• N •• Cl • • •• σ:Osp2─Nsp2 ↑↓ N = 3 electron groups = sp O = 3 electron groups = sp2 Cl = 4 electron groups = sp3 2 O ↑↓ N ↑↓ ↑↓ N π:Op─Np 127 σ:Nsp2─Clp Cl ↑↓ O Tro: Chemistry: A Molecular Approach, 2/e ↑↓ Cl Copyright © 2011 Pearson Education, Inc.
  • 128. Problems with Valence Bond Theory • VB theory predicts many properties better than Lewis theory  bonding schemes, bond strengths, bond lengths, bond rigidity • However, there are still many properties of molecules it doesn’t predict perfectly magnetic behavior of O2 • In addition, VB theory presumes the electrons are localized in orbitals on the atoms in the molecule – it doesn’t account for delocalization Tro: Chemistry: A Molecular Approach, 2/e 128 Copyright © 2011 Pearson Education, Inc.
  • 129. Molecular Orbital Theory • In MO theory, we apply Schrödinger’s wave equation to the molecule to calculate a set of molecular orbitals  in practice, the equation solution is estimated  we start with good guesses from our experience as to what the orbital should look like  then test and tweak the estimate 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  delocalization Tro: Chemistry: A Molecular Approach, 2/e 129 Copyright © 2011 Pearson Education, Inc.
  • 130. LCAO • 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 method  weighted sum • Because the orbitals are wave functions, the waves can combine either constructively or destructively Tro: Chemistry: A Molecular Approach, 2/e 130 Copyright © 2011 Pearson Education, Inc.
  • 131. Molecular Orbitals • When the wave functions combine constructively, the resulting molecular orbital has less energy than the original atomic orbitals – it is called a Bonding Molecular Orbital  σ, π  most of the electron density between the nuclei • When the wave functions combine destructively, the resulting molecular orbital has more energy than the original atomic orbitals – it is called an Antibonding Molecular Orbital  σ*, π*  most of the electron density outside the nuclei  nodes between nuclei Tro: Chemistry: A Molecular Approach, 2/e 131 Copyright © 2011 Pearson Education, Inc.
  • 132. Interaction of 1s Orbitals Tro: Chemistry: A Molecular Approach, 2/e 132 Copyright © 2011 Pearson Education, Inc.
  • 133. Molecular Orbital Theory • Electrons in bonding MOs are stabilizing  lower energy than the atomic orbitals • Electrons in antibonding MOs are destabilizing  higher in energy than atomic orbitals  electron density located outside the internuclear axis  electrons in antibonding orbitals cancel stability gained by electrons in bonding orbitals Tro: Chemistry: A Molecular Approach, 2/e 133 Copyright © 2011 Pearson Education, Inc.
  • 134. Energy Comparisons of Atomic Orbitals to Molecular Orbitals Tro: Chemistry: A Molecular Approach, 2/e 134 Copyright © 2011 Pearson Education, Inc.
  • 135. MO and Properties • Bond Order = difference between number of electrons in bonding and antibonding orbitals  only need to consider valence electrons  may be a fraction  higher bond order = stronger and shorter bonds  if bond order = 0, then bond is unstable compared to individual atoms and no bond will form • A substance will be paramagnetic if its MO diagram has unpaired electrons  if all electrons paired it is diamagnetic Tro: Chemistry: A Molecular Approach, 2/e 135 Copyright © 2011 Pearson Education, Inc.
  • 136. Hydrogen Atomic Orbital Dihydrogen, H2 Molecular Orbitals σ* 1s Hydrogen Atomic Orbital 1s σ Because more electrons are in bonding orbitals than are in antibonding orbitals, net bonding interaction Tro: Chemistry: A Molecular Approach, 2/e 136 Copyright © 2011 Pearson Education, Inc.
  • 137. H2 σ* Antibonding MO LUMO σ bonding MO HOMO Tro: Chemistry: A Molecular Approach, 2/e 137 Copyright © 2011 Pearson Education, Inc.
  • 138. Helium Atomic Orbital Dihelium, He2 Molecular Orbitals σ* 1s Helium Atomic Orbital 1s σ BO = ½(2-2) = 0 Because there are as many electrons in antibonding orbitals as in bonding orbitals, there is no net bonding interaction Tro: Chemistry: A Molecular Approach, 2/e 138 Copyright © 2011 Pearson Education, Inc.
  • 139. Lithium Atomic Orbitals Dilithium, Li2 Molecular Orbitals σ∗ 2s 2s σ σ∗ BO = ½(4-2) = 1 1s Any fill energy level will generate filled bonding and antibonding MO’s; therefore only need to consider valence shell 1s σ Because more electrons are in bonding orbitals than are in antibonding orbitals, there is a net bonding interaction Tro: Chemistry: A Molecular Approach, 2/e Lithium Atomic Orbitals 139 Copyright © 2011 Pearson Education, Inc.
  • 140. Li2 σ* Antibonding MO LUMO σ bonding MO HOMO Tro: Chemistry: A Molecular Approach, 2/e 140 Copyright © 2011 Pearson Education, Inc.
  • 141. Interaction of p Orbitals Tro: Chemistry: A Molecular Approach, 2/e 141 Copyright © 2011 Pearson Education, Inc.
  • 142. Interaction of p Orbitals Tro: Chemistry: A Molecular Approach, 2/e 142 Copyright © 2011 Pearson Education, Inc.
  • 143. Tro: Chemistry: A Molecular Approach, 2/e 143 Copyright © 2011 Pearson Education, Inc.
  • 144. O2 • Dioxygen is paramagnetic • Paramagnetic material has unpaired electrons • Neither Lewis theory nor valence bond theory predict this result Tro: Chemistry: A Molecular Approach, 2/e 144 Copyright © 2011 Pearson Education, Inc.
  • 145. O2 as Described by Lewis and VB Theory Tro: Chemistry: A Molecular Approach, 2/e 145 Copyright © 2011 Pearson Education, Inc.
  • 146. Oxygen Atomic Orbitals 2p σ∗ π∗ Oxygen Atomic Orbitals 2p O2 MO’s π Because more electrons are in bonding orbitals than are in antibonding orbitals, there is a net bonding interaction σ σ∗ BO = ½(8 be – 4 abe) BO = 2 2s Because there are unpaired electrons in the antibonding orbitals, O2 is predicted to be paramagnetic 2s σ Tro: Chemistry: A Molecular Approach, 2/e 146 Copyright © 2011 Pearson Education, Inc.
  • 147. Example 10.10: Draw a molecular orbital diagram of N2− ion and predict its bond order and magnetic properties Write a MO diagram for N2− using N2 as a base Count the number of valence electrons and assign these to the MOs following the aufbau principle, Pauli principle & Hund’s rule σ∗2p N has 5 valence electrons 2 N = 10e− (−) = 1e− total = 11e− Tro: Chemistry: A Molecular Approach, 2/e π∗2p ↑ ↑↓ σ2p ↑↓ ↑ ↓ π2p ↑↓ σ∗2s ↑↓ σ2s 147 Copyright © 2011 Pearson Education, Inc.
  • 148. Example 10.10: Draw a molecular orbital diagram of N2− ion and predict its bond order and magnetic properties Calculate the bond order by taking the number of bonding electrons and subtracting the number of antibonding electrons, then dividing by 2 Determine whether the ion is paramagnetic or diamagnetic BO = ½(8 be – 3 abe) BO = 2.5 Because this is lower than the bond order in N2, the bond should be weaker Because there are unpaired electrons, this ion is paramagnetic Tro: Chemistry: A Molecular Approach, 2/e 148 σ∗2p π∗2p ↑ ↑↓ σ2p ↑↓ ↑ ↓ π2p ↑↓ σ∗2s ↑↓ σ2s Copyright © 2011 Pearson Education, Inc.
  • 149. Practice – Draw a molecular orbital diagram of C2+ and predict its bond order and magnetic properties Tro: Chemistry: A Molecular Approach, 2/e 149 Copyright © 2011 Pearson Education, Inc.
  • 150. Practice – Draw a molecular orbital diagram of C2+ and predict its bond order and magnetic properties C has 4 valence electrons σ∗2p 2 C = 8e− (+) = −1e− total = 7e− π∗2p BO = ½(5 be – 2 abe) BO = 1.5 Because there are unpaired electrons, this ion is paramagnetic Tro: Chemistry: A Molecular Approach, 2/e 150 σ2p ↑↓ ↑ π2p ↑↓ σ∗2s ↑↓ σ2s Copyright © 2011 Pearson Education, Inc.
  • 151. Heteronuclear Diatomic Molecules & Ions • When the combining atomic orbitals are • identical and 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 Tro: Chemistry: A Molecular Approach, 2/e 151 Copyright © 2011 Pearson Education, Inc.
  • 152. Heteronuclear Diatomic Molecules & Ions • 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 Tro: Chemistry: A Molecular Approach, 2/e 152 Copyright © 2011 Pearson Education, Inc.
  • 153. NO a free radical σ2s Bonding MO shows more electron density near O because it is mostly O’s 2s atomic orbital Tro: Chemistry: A Molecular Approach, 2/e BO = ½(6 be – 1 abe) BO = 2.5 153 Copyright © 2011 Pearson Education, Inc.
  • 154. HF Tro: Chemistry: A Molecular Approach, 2/e 154 Copyright © 2011 Pearson Education, Inc.
  • 155. 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 Gives results that better match real molecule properties than either Lewis or valence bond theories Tro: Chemistry: A Molecular Approach, 2/e 155 Copyright © 2011 Pearson Education, Inc.
  • 156. Ozone, O3 MO Theory: Delocalized π bonding orbital of O3 Tro: Chemistry: A Molecular Approach, 2/e 156 Copyright © 2011 Pearson Education, Inc.