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Chapter 10 Bonding Theory and Molecular Structure

Chapter 10 Bonding Theory and Molecular Structure. General Chemistry: An Integrated Approach Hill, Petrucci, 4 th Edition. Mark P. Heitz State University of New York at Brockport © 2005, Prentice Hall, Inc. Molecular Geometry.

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Chapter 10 Bonding Theory and Molecular Structure

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  1. Chapter 10Bonding Theory and Molecular Structure General Chemistry: An Integrated ApproachHill, Petrucci, 4th Edition Mark P. Heitz State University of New York at Brockport © 2005, Prentice Hall, Inc.

  2. Molecular Geometry The molecular geometry, or the shape of a molecule is described by the geometric figure formed when the atomic nuclei are imagined to be joined in straight lines. EOS Chapter 10: Bonding Theory...

  3. The mutual repulsions among electron groups lead to an orientation of the groups that is called electron-group geometry EOS Valence-Shell Electron-Pair Repulsion (VSEPR) Ideal molecular geometry is based on the idea that pairs of valence electrons in bonded atoms repel one another An electron group is any collection of valence electrons, localized in a region around a central atom, that repels other groups of valence electrons Chapter 10: Bonding Theory...

  4. 2 electron groups – linear 3 electron groups – trigonal planar 4 electron groups – tetrahedral 5 electron groups – trigonal bipyramidal 6 electron groups – octahedral EOS Electron-Group Geometries How can electrons on the central atom be arranged so they are as far apart as possible? Chapter 10: Bonding Theory...

  5. In the VSEPR notation used to describe molecular geometries, the central atom in a structure is denoted as A, terminal atoms as X, and the lone pairs of electrons as E EOS VSEPR Notation Table 10.1 in the text summarizes various possibilities for molecular geometries in relation to electron-group geometries Chapter 10: Bonding Theory...

  6. Geometry is unaffected by lone pairs of electrons EOS A VSEPR Summary The optimal repulsion arrangement is opposite ends of a line Chapter 10: Bonding Theory...

  7. The optimal repulsion arrangement is the corners of a triangle—note the planarity With a lone pair, the geometry is essentially the same EOS A VSEPR Summary Chapter 10: Bonding Theory...

  8. The optimal repulsion arrangement is the corners of a regular tetrahedron—four equal triangular faces There can be one or two lone pairs on the central atom EOS A VSEPR Summary Chapter 10: Bonding Theory...

  9. EOS A VSEPR Summary Chapter 10: Bonding Theory...

  10. The optimal repulsion arrangement is a triangle with a bisecting line through the triangle face Up to three lone pairs can be included here EOS A VSEPR Summary Chapter 10: Bonding Theory...

  11. A VSEPR Summary EOS Chapter 10: Bonding Theory...

  12. One or two lone-pair geometries are most common EOS A VSEPR Summary The optimal repulsion arrangement is at the face centers of a cube Chapter 10: Bonding Theory...

  13. A VSEPR Summary Illustration EOS Chapter 10: Bonding Theory...

  14. The dipole moment (m) of a molecule is the product of the magnitude of the charge (d) and the distance (d) that separates the centers of positive and negative charge m = dd EOS Polar Moleculesand Dipole Moments A molecule with separate centers of positive and negative charge is called a polar molecule Chapter 10: Bonding Theory...

  15. Dipolar molecules align with electric fields EOS Polar Moleculesand Dipole Moments Dipole moments are generally expressed in a quantity called a debye, D Chapter 10: Bonding Theory...

  16. All polar covalent bonds have a bond dipole; a separation of positive and negative charge centers in an individual bond A charge separation in the molecule as a whole, considering all the bonds, is a molecular dipole These concepts explain why CO2 is linear with no dipole moment (m = 0 D) and water is bent (bond angle = 104.5o) with a dipole moment of m = 1.84 D EOS Bond and Molecular Dipoles Chapter 10: Bonding Theory...

  17. From the molecular shape, determine whether bond dipoles cancel to give a nonpolar molecule or combine to produce a resultant dipole moment for the molecule EOS Molecular Shapesand Dipole Moments Molecules can be predicted to be polar or nonpolar based on the following three-step approach • Use electronegativity values to predict bond dipoles • Use the VSEPR method to predict the molecular shape Chapter 10: Bonding Theory...

  18. Valence Bond (VB) Theory states that a covalent bond is formed by the pairing of two electrons with opposing spins in the region of overlap of atomic orbitals between two atoms The more extensive the overlap between two orbitals, the stronger is the bond between two atoms EOS Atomic Orbital Overlap This overlap region has a high electron charge density Chapter 10: Bonding Theory...

  19. Bonding in H2S EOS Chapter 10: Bonding Theory...

  20. For orbitals with directional lobes, maximum overlap occurs when atomic orbitals overlap end to end; that is, a hypothetical line joining the nuclei of the bonded atoms passes through the region of maximum overlap EOS Several Important Points Most of the electrons in a molecule remain in the same orbital locations that they occupied in the separated atoms Bonding electrons are localized in the region of atomic orbital overlap Chapter 10: Bonding Theory...

  21. Based on ground-state electron configuration, carbon should have only two bonds If a 2s electron is promoted to an empty 2p orbital, then four unpaired electrons can give rise to four bonds These four orbitals become mixed, or hybridized to form bonds EOS Hybridization of Atomic Orbitals HybridizationVideo Chapter 10: Bonding Theory...

  22. The total number of hybrid orbitals is equal to the number of atomic orbitals combined Hybrid orbitals may overlap with pure atomic orbitals or with other hybrid orbitals EOS sp3 Hybridization Occurs most often for central atom only Chapter 10: Bonding Theory...

  23. sp3 Hybridization The carbon atom in methane (CH4) has bonds that are sp3 hybrids Note that in this molecule carbon has all singlebonds EOS Chapter 10: Bonding Theory...

  24. Bonding in Ammonia Ammonia (NH3) is similar except the lone pair of electrons occupies the 4th hybrid orbital EOS Chapter 10: Bonding Theory...

  25. The empty 2p orbital remains unhybridized EOS sp2 Hybridization This hybridization scheme is useful in describing double covalent bonds Comprised of one 2s orbital and two 2p orbitals to produce a set of three sp2 hybrid orbitals Chapter 10: Bonding Theory...

  26. This distribution gives a trigonal planar molecular geometry, as predicted by VSEPR EOS Determining Empirical Formulas The geometric distribution of the three sp2 hybrid orbitals is within a plane, directed at 120o angles Chapter 10: Bonding Theory...

  27. The geometric distribution of the two sp hybrid orbitals is on a line, directed at 180o angles This distribution gives a linear molecular geometry EOS sp Hybridization This hybridization scheme is useful in describing triple covalent bonds Chapter 10: Bonding Theory...

  28. This hybridization allows for expanded valence shell compounds – typical for group 5A elements, e.g., P A 3s electron can be promoted to a 3d subshell, which gives rise to a set of five sp3d hybrid orbitals EOS d Subshells Hybrid Orbitals Chapter 10: Bonding Theory...

  29. This hybridization allows for expanded valence shell compounds – typically group 6A elements, e.g., S A 3s and a 3p electron can be promoted to the 3d subshell, which gives rise to a set of six sp3d2 hybrid orbitals EOS d Subshells Hybrid Orbitals Chapter 10: Bonding Theory...

  30. Describe the orbital overlap and molecular geometry EOS Predicting Hybridization Schemes In hybridization schemes, one hybrid orbital is produced for every simple atomic orbital involved Write a plausible Lewis structure for the molecule or ion Use the VSEPR method to predict the electron-group geometry of the central atom Select the hybridization scheme that corresponds to the VSEPR prediction Chapter 10: Bonding Theory...

  31. Hybrid Orbitals and TheirGeometric Orientations EOS Chapter 10: Bonding Theory...

  32. A triple bond is made up of one sigma bond and two pi bonds EOS Hybrid Orbitals andMultiple Covalent Bonds Covalent bonds formed by the end-to-end overlap of orbitals, regardless of orbital type, are called sigma (s) bonds. All single bonds are sigma bonds A bond formed by parallel, or side-by-side, orbital overlap is called a pi (p) bond A double bond is made up of one sigma bond and one pi bond Chapter 10: Bonding Theory...

  33. Carbon–Carbon Double Bonds A double bond is made up of one sigma bond and one pi bond EOS Chapter 10: Bonding Theory...

  34. EOS Carbon–Carbon Triple Bonds Chapter 10: Bonding Theory...

  35. Each compound is distinctly different in both physical and chemical properties EOS Geometric Isomerism Geometric isomers are isomers that differ only in the geometric arrangement of certain substituent groups Chapter 10: Bonding Theory...

  36. Molecular orbitals (MOs) are built up (Aufbau principle) in the same way as atomic orbitals EOS Molecular Orbitals Molecular orbitals (MOs) are mathematical equations that describe the regions in a molecule where there is a high probability of finding electrons Molecular orbitals (MOs) are essentially combinations of atomic orbitals – two types exist, bonding and antibonding orbitals Chapter 10: Bonding Theory...

  37. Bonding MO = enhanced region of electron density EOS Visualizing MOs The hydrogen molecule Antibonding MO = region of diminished electron density Chapter 10: Bonding Theory...

  38. MOs for the 2p Electrons EOS Chapter 10: Bonding Theory...

  39. Antibonding MOs = higher energy s* and p* MOs EOS Molecular Orbital Diagram Bonding MOs s AOs = s MOs p AOs = p MOs Chapter 10: Bonding Theory...

  40. Second-Period Homonuclear Diatomic Molecules EOS Chapter 10: Bonding Theory...

  41. Kekulé discovered that benzene has a cyclic structure and he proposed that a hydrogen atom was attached to each carbon atom and that alternating single and double bonds joined the carbon atoms together EOS Bonding in Benzene The structure of benzene, C6H6, discovered by Michael Faraday in 1825, was not figured out until 1865 by F. A. Kekulé Chapter 10: Bonding Theory...

  42. EOS Benzene This kind of structure gives rise to two important resonance hybrids and leads to the idea that all three double bonds are delocalized across all six carbon atoms Chapter 10: Bonding Theory...

  43. The s-Bonding Framework EOS Chapter 10: Bonding Theory...

  44. Aromatic compound simply refers to a substance with a ring structure and with bonding characteristics and properties related to those of benzene EOS Aromatic Compounds Many of the first benzene-like compounds discovered had pleasant odors and hence acquired the name aromatic Chapter 10: Bonding Theory...

  45. Summary of Concepts • The VSEPR method is used to predict the shapes of molecules and polyatomic ions • If all electron groups are bonding groups, the molecular geometry is the same as the electron-group geometry • A polar covalent bond has separate centers of positive and negative charge, creating a bond dipole EOS Chapter 10: Bonding Theory...

  46. Unhybridized p orbitals overlap in a side-by-side fashion to form p bonds EOS Summary of Concepts • In the valence bond theory, a covalent bond is formed by the overlap of atomic orbitals of the bonded atoms in a region between the atomic nuclei • Hybridized orbitals include sp, sp2, sp3, sp3d, and sp3d2 Chapter 10: Bonding Theory...

  47. Benzene-like compounds are called aromatic compounds EOS Summary (cont.) • Single bonds are all hybridized s bonds, double bonds have one s bond and one p bond, and triple bonds have one s bond and two p bonds • In molecular orbital theory, atomic orbitals of separated atoms are combined into molecular orbitals • The benzene molecule is usually represented by its resonance hybrid Chapter 10: Bonding Theory...

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