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Chapter Eight

Object: Bonding and Structure of materials (chemicals). Comparing diamond & graphite : The bounding of substances (chemical) has a profound effect on chemical and physical properties. Comparing silicon & carbon: Group 4A CO 2 vs SiO 2

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Chapter Eight

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  1. Object: Bonding and Structure of materials (chemicals) Comparing diamond & graphite: The bounding of substances (chemical) has a profound effect on chemical and physical properties. Comparing silicon & carbon: Group 4A CO2 vs SiO2 Molecular bonding and structure play the central role in understand the properties of matter and in the determining the course of all chemical reactions. Chapter Eight

  2. Questions to Consider • What is meant by the term “chemical bond?” • Why do atoms bond with each other to form molecules? • How do atoms bond with each other to form molecules?

  3. Figure 10.27: Examples of silicate anions, all of which are based on SiO44- tetrahedra.

  4. Key Ideas in Bonding Chemical Bond: energy , force ＃: intramolecular and intermolecular bond Types of Chemical Bonds: • Ionic Bonding: Electrons are transferred, NaCl, Ionic compounds, Coulomb’s force • Covalent Bonding: Electrons are shared equally, H2 • What about intermediate cases? • HF, polar covalent bonding

  5. Figure 8.1 a & b (a) The Interaction of Two Hydrogen Atoms (b) Energy Profile as a Function of the Distance Between the Nuclei of the Hydrogen Atoms

  6. Figure 8.2 The Effect of an Electric Field on Hydrogen Fluoride Molecules

  7. 8.2 Electronegativity • If lithium and fluorine react, which has more attraction for an electron? Why? • In a bond between fluorine and iodine, which has more attraction for an electron? Why?

  8. Figure 8.3 The Pauling Electronegativity Vaules Δ= (H–X)act - (H–X)exp

  9. Table 8.1 The Relationship Between Electronegativity and Bond Type

  10. 8.3 Bond Polarity and Dipole Moments Figure 8.4 An Electrostatic Potential Map of HF

  11. Figure 8.5 a-c The Charge Distribution in the Water Molecule

  12. Figure 8.6 a-c The Structure and Charge Distribution of the Ammonia Molecule

  13. Figure 8.7 a-c The Carbon Dioxide Molecule

  14. Table 8.2 Types of Molecules with Polar Bonds but No Resulting Dipole Moment

  15. e.p. Diagram HCL

  16. e.p.Diagram SO3

  17. e.p. Diagram CH4

  18. e.p. Diagram H2S

  19. Question Which of the following bonds would be the least polar yet still be considered polar covalent? Mg-O C-O O-O Si-O N-O

  20. Ions: Electron Configurations and Sizes What we can “read” from the periodic table: • Trends for • Atomic size • Ion radius • Ionization energy • Electronegativity • Electron configurations • Predicting formulas for ionic compounds • Ranking polarity of covalent bonds

  21. Table 8.3 Common Ions with Noble Gas Configurations in Ionic Compounds

  22. 8.4 Ions: Electron Configurations and Sizes Figure 8.8 Sizes of Ions Related to Positions of the Elements on the Periodic Table

  23. QUESTION Arrange the following, without consulting any specific listed radii, in order from smallest to largest size (radius) S2– ; Cl–; K+; Rb+ 1. S2– ; Cl–; K+; Rb+ 2. K+; Rb+; Cl–; S2– 3. S2– ; Cl–; Rb+; ; K+ 4. S2– ; Cl–; Rb+; K+

  24. Formation of an Ionic Solid 1. Sublimation of the solid metal • M(s)  M(g) [endothermic] 2. Ionization of the metal atoms • M(g)  M+(g) + e [endothermic] 3. Dissociation of the nonmetal • 1/2X2(g)  X(g) [endothermic] 4. Formation of X ions in the gas phase: • X(g) + e X(g) [exothermic] 5. Formation of the solid MX • M+(g) + X(g)  MX(s) [quite exothermic]

  25. Figure 8.9 The Energy Changes Involved in the Formation of Lithium Fluoride from Its Elements

  26. Figure 8.11 Comparison of the Energy Changes Involved in the Formation of Solid Sodium Fluoride and Solid Magnesium Oxide

  27. Figure 8.10 a & b The Structure of Lithium Fluoride

  28. Figure 8.12 a-c The Three Possible Types of Bonds

  29. Figure 8.13 The Relationship Between the Ionic Character of a Covalent Bond and the Electronegativity Difference of the Bounded Atoms

  30. Models • Models are attempts to explain how nature operates on the microscopic level based on experiences in the macroscopic world.

  31. Fundamental Properties of Models • A model does not equal reality. • Models are oversimplifications, and are therefore often wrong. • Models become more complicated as they age. • We must understand the underlying assumptions in a model so that we don’t misuse it.

  32. Model of Chemical Bond • Individual bond occurs between pair electron. • Average individual bond energy • Bond energy values can used to calculate the reaction energies.

  33. Table 8.4 Average Bond Energies (kj/mol)

  34. Table 8.5 Bond Lengths for Selected Bonds

  35. Localized Electron Model • A molecule is composed of atoms that are bound together by sharing pairs of electrons using the atomic orbitals of the bound atoms. • bonding pairs, lone pairs

  36. Localized Electron Model • Description of valence electron arrangement (Lewis structure). • Prediction of geometry (VSEPR model). • Description of atomic orbital types used to share electrons or hold long pairs.

  37. Lewis Structure • Shows how valence electrons are arranged among atoms in a molecule. • Reflects central idea that stability of a compound relates to noble gas electron configuration.

  38. Lewis Structures • Sum the valence electrons. • Place bonding electrons between pairs of atoms. • Atoms usually have noble gas configurations.

  39. 5)

  40. Resonance • More than one Lewis structure for a given molecule. Ex. NO3- • Resonance structure, Ex. NO2- • Formal charge: (# of VEs on free atom)-(# of VEs assigned to the atom in the molecule), Ex. SO42- • Rules of formal charge • Assigning the structure based on formal charge (predicted structure should be judged by expt.

  41. Molecular Structure: VSEPR Model • Molecular structure: 3D arrangement of the atoms in a molecule. • VSEPR: valence shell electron-pair repulsion • Assume: The structure around a given atom is determined principally by minimizing electron pair repulsions. (Bonding and nonbonding pairs around a given atom should be positioned as far as possible.) • Ex. BeCl2, BF3, CH4

  42. Figure 8.15 The Molecular Structure of Methane

  43. Balloons Tied Together Naturally Form Tetrahedral Shape

  44. Predicting a VSEPR Structure • Draw Lewis structure. • Put pairs as far apart as possible. • Determine positions of atoms from the way electron pairs are shared. • Determine the name of molecular structure from positions of the atoms.

  45. Figure 8.16 a-c The Molecular Structure of Ammonia is a Trigonal Pyramid

  46. Figure 8.17 a-c The Tetrahedral Arrangement of Oxygen In a Water Molecule

  47. Figure 8.18 The Bond Angles In the CH4, NH3, and H2O Molecules

  48. Figure 8.19 a & b In a Bonding Pair of Electrons the Electrons are Shared by Two Nuclei (b) In a Lone Pair, Both Electrons Must Be Close to a Single Nucleus

  49. Predicting a VSEPR Structure • Draw Lewis structure. • Put pairs as far apart as possible. • Determine positions of atoms from the way electron pairs are shared. • Determine the name of molecular structure from positions of the atoms. • Lone pairs require more room than bonding pairs

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