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Chapter 12

Chapter 12. 12.1 Types of Chemical Bonds 12.2 Electronegativity. 12.1 Types of Chemical Bonds. Objectives: To learn about ionic and covalent bonds and explain how they are formed. To learn about the polar covalent bond.

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Chapter 12

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  1. Chapter 12 12.1 Types of Chemical Bonds 12.2 Electronegativity

  2. 12.1 Types of Chemical Bonds Objectives: To learn about ionic and covalent bonds and explain how they are formed. To learn about the polar covalent bond.

  3. Molecular Bonding and Structure play the central role in determining the course of chemical reactions, many of which are vital to our survival.

  4. Structure plays a central role in our senses of smell and taste.Substances have a particular odor because they fit into the specially shaped receptors in our nasal passages.

  5. 12.1 Types of Chemical Bonds Bond: force that holds groups of 2 or more atoms together and makes them function as a unit. Bond energy: energy required to break the bond. This gives information about the strength of a bond.

  6. 12.1 Types of Chemical Bonds Ionic bonding: attraction between closely packed opposite charged ions. Metal Nonmetal Ionic Compound M + X M+X- Ionic compound: metal reacts with a nonmetal e-

  7. 12.1 Types of Chemical Bonds • Covalent bonds: electrons are shared by nuclei

  8. 12.1 Types of Chemical Bonds Polar covalent bonds: electrons are unequally shared by nuclei. When 2 nonmetals react.

  9. 12.2 Electronegativity • Objectives: To understand the nature of bonds and their relationship to electronegativity. Electronegativity: relative ability of an atom in a molecule to attract shared electrons to itself. Chemists determine this by measuring the polarity of the bonds between various atoms.

  10. Polarity of a bond depends on the difference between the electronegativity values of the atoms forming the bond. If this difference is small relatively nonpolar If this difference is large more polar >2 IONIC

  11. Table 12.1

  12. Covalent bond between 2 identical atoms. Polar covalent bonds Ionic Bond with no electron sharing

  13. Let’s Practice Which is the more polar pair? 1) H-S or H-F H-F O-S 2) O-S or O-F C-Cl 3) C-S or C-Cl Homework Self-check Exercise 12.1 8-12 p. 394-395

  14. 12.3 Bond Polarity and Dipole Moment Objective: To understand bond polarity and how it is related to molecular polarity

  15. 12.3 Bond Polarity and Dipole Moments Dipole moment: a molecule that has a center of positive charge and a center of negative charge. Dipole character is often represented by an arrow.

  16. 12.3 Bond Polarity and Dipole Moment The fact that water is a polar allows water to attract both positive and negative ions. This allows many things to dissolve in water.

  17. 12.4 Stable Electron Configurations and Charges on Ions Objectives: To learn about stable electron configurations. To learn to predict the formulas of ionic compounds.

  18. Table 12.2 When a metal reacts with a nonmetal the metal loses electrons and the nonmetal Gains these electrons.

  19. 12.4 Stable Electron Configurations and Charges on Ions In almost all stable chemical compounds Of the representative elements, all of the atoms have achieved a noble gas electron configuration.

  20. Table 12.3

  21. 12.4 Stable Electron Configurations and Charges on Ions So what do we need to know this for Ca: [Ar]4s2 O: [He]2s22p4. Calculate the electronegativity Oxygen is 3.5, Ca is 1.0 Difference is 2.5 So how many electrons are transferred? CaO is the empirical formula

  22. 12.5 Ionic Bonding and Structures of Ionic Compounds Objectives: To learn about ionic structures To understand factors governing ionic size.

  23. Anion is always larger than the parent atom. Cation is always smaller than the parent atom

  24. WHY?

  25. A note on polyatomics Polyatomic ions NH4+ and NO3- are held together by covalent bonds.

  26. 12.6 Lewis Structures Objective: To learn to write Lewis structures

  27. 12.6 Lewis Structures Bonding involves just the valence electrons. Lewis structure: representation of a molecule that shows how the valence electrons are arranged among the atoms in the molecule. Name after G.N. Lewis who used it as a tool to teach Chemistry students.

  28. 12.6 Lewis Structures Example: K+ [:Br:]- Ionic Bonds Covalent Bonds Duet rule: for Hydrogen H:H Octet rule: surrounded by 8 electrons :F:F: .. .. Bonding pair .. .. .. .. Lone pairs or Unshared pairs

  29. H:P:H H .. .. .. .. .. .. 12.6 Lewis Structures Example: Write the Lewis structures of the following molecules.CCl4PH3 .. .. :Cl: :Cl:C:Cl: :Cl: STEPS: 1) Get the sum of the valence electrons from all of the atoms. 2) Use one pair of electrons to form a bond between each pair of bound atoms. 3) Arrange the remaining electrons to satisfy the octet rule or duet rule

  30. 12.7Lewis Structures of Molecules w/Multiple Bonds Objective: To learn how to write Lewis structures for molecules with multiple bonds.

  31. .. :O:::C-O: or :O:C:::O: .. 12.7Lewis Structures of Molecules w/Multiple Bonds Example: Write the Lewis structures of the following molecules.CO2 4+ 6+ 6= 16 valence electrons .. .. .. .. :O:C:O: Single Bond shares 1 pair of electrons But this is not correct because: ? .. .. :O::C::O: DOUBLE BOND- 2 PAIRS of e- BUT .. .. Triple Bond-3 e- pairs RESONANCE: more than one Lewis structure can be drawn for the molecule.

  32. 12.7Lewis Structures of Molecules w/Multiple Bonds Some exceptions to the Octet Rule Boron tends to form compounds in which the boron has fewer Than 8 electrons around it. F VERY REACTIVE w/ NH3 and H2O H F B H-N-B-F F F H F Another exception: .. .. HOWEVER, Oxygen is paramagnetic Which suggests an unpaired electrons. :O=O:

  33. 12.8 Molecular Structure Molecular or geometric structure: 3-Dimensional arrangement of the atom in a molecule. “BENT” or “V-shaped” Bond angle about 105o

  34. O C O 180o 12.8 Molecular Structure Molecular or geometric structure: 3-Dimensional arrangement of the atom in a molecule. Linear structure Bond angle 180o F 120o Trigonal planar B F F Tetrahedral structure

  35. 12.9 Molecular Structure: The VSEPR Model Objective: To learn to predict molecular geometry from the number of electron pairs.

  36. 12.9 Molecular Structure: The VSEPR Model VSEPR: valence shell electron pair repulsion model. -is that the structure around a given atom is determined by minimizing repulsions between electron pairs. -bonding and nonbonding electron pairs are positioned as far apart as possible.

  37. Predicting the Molecular Structureusing the VSEPR model Step 1: Draw the Lewis structure for the molecule. Step 2: Count the electron pairs and arrange them in the way that minimizes repulsions. (as far apart as possible) Step 3: Determine the positions of the atoms from the way the electron pairs are shared. Step 4: Determine the name of the molecular structure from the positions of the atoms.

  38. .. H:N:H .. H Example: NH3

  39. Table 12.4

  40. 12.10:Molecular Structure with Double Bonds Objective: To learn to apply the VSEPR model to molecules with double bonds.

  41. 12.10:Molecular Structure with Double Bonds When using the VSEPR model to predict the molecular geometry of a molecule, a double bond is counted the same as a single electron pair. SOOOO apply the same rules as for single bonds.

  42. 12.10:Molecular Structure with Double Bonds CO2-planar NO3-trigonal planar

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