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GS105 Chapter 15: Bonding

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  1. GS105 Chapter 15: Bonding • Valence Electrons & e- Dot Structures • Octet Rule & Ions • Ionic Compounds & Formulas • Covalent Compounds & Formulas • Polyatomic Ions • Molecular Shapes & Polarity • Attractive Forces

  2. 24 12 Mg Electron arrangement Electrons fill layers around nucleus Low  High 32 18 8 2 Shells = Energy levels

  3. Inner vs. valence electrons • Valence electrons • This is where • most chemical • reactions • occur. • Inner electrons • Not much happens • here under normal • conditions.

  4. 1 1 4 2 H He 9 4 Be 20 10 7 3 Ne Li 40 18 Ar 24 12 Mg 23 11 Na IA IIA VIIIA 2, 1 2, 2 2, 8 2, 8, 8 2, 8, 1 2, 8, 2

  5. 1 1 1 H 9 4 11 5 Be B 7 3 Li 24 12 Mg 27 13 23 11 Al Na Valence electrons Where most chemical Reactions occur. 2 3 2, 3 2, 1 2, 2 2, 8, 3 2, 8, 1 2, 8, 2

  6. 1 1 1 4 2 H He 9 4 20 10 7 3 Be Ne Li 40 18 Ar 24 12 Mg 23 11 Na 8 Octet Rule 2 2, 1 2, 2 2, 8 2, 8, 8 2, 8, 1 2, 8, 2

  7. 1 1 H H 7 3 Li Li Na 23 11 Na K Lewis Structures Show only Valence Electrons

  8. Ge Ga Be H He Ca Mg Al B Li Na Si C P N As S O F Se Cl Br Kr Ar Ne K 1 8 2 3 4 5 6 7

  9. Na 23 11 Na Ions 11 +’s 11 -’s 0 Metals give e-s to make Cations 11 +’s 10 -’s 1 + Na1+ 2, 8 = [Ne]

  10. 35 17 Cl Cl 1- Cl Ions 17 +’s 17 -’s 0 Nonmetals take e-s to make Anions 17 +’s 18 -’s 1 - = Cl1- 2, 8, 8 = [Ar]

  11. Formation of NaCl e-moves from Metal  Nonmetal Stable octets _ Na + Cl Na+ +Cl Metal Cation Nonmetal Anion + and - ions attract to form an ionic bond.

  12. Ionic compounds • Not individual molecules • Form crystal arrays • Ions touch many others • Formula represents the average ion ratio NaCl sodiumchloride Na Cl Cl Na Cl Na

  13. H He Li Be B C N O F Ne Na Mg Al Si P S Cl Ar K Ca Sc Ti V Cr Mn Fe Co Ni Cu Zn Ga Ge As Se Br Kr Rb Sr Y Zr Nb Mo Tc Ru Rh Pd Ag Cd In Sn Sb Te I Xe Cs Ba Ls Hf Ta W Re Os Ir Pt Au Hg Tl Pb Bi Po At Rn Fr Ra Ac Ce Pr Nd Pm Sm Eu Gd Tb Dy Ho Er Tm Yb Lu Th Pa U Np Pu Am Cm Bk Cf Es Fm Md No Lr Common ions Representative Elements 1+ 4+ 4- 2+ 3+ 3- 2- 1- 4 - 6

  14. Ionic Formulas Metal Cations +NonmetalAnions Na1+ Cl1- Al3+ Cl1- Cl1- Cl1- NaCl AlCl3 SodiumChloride Aluminum Chloride

  15. Sc Ti V Cr Mn Fe Co Ni Cu Zn Y Zr Nb Mo Tc Ru Rh Pd Ag Cd Ls Hf Ta W Re Os Ir Pt Au Hg Ac Ce Pr Nd Pm Sm Eu Gd Tb Dy Ho Er Tm Yb Lu Th Pa U Np Pu Am Cm Bk Cf Es Fm Md No Lr Common ions Representative Elements 1+ 4+ 4- H He 2+ 3+ 3- 2- 1- Transition Elements Li Be B C N O F Ne Variable Na Mg Al Si P S Cl Ar K Ca Ga Ge As Se Br Kr Rb Sr In Sn Sb Te I Xe Cs Ba Tl Pb Bi Po At Rn Fr Ra

  16. Information in the table Atomic number Atomic mass (weight) 26 55.845 Oxidation states (Valence) 2,3 Fe Elemental Symbol Electronic Configuration [Ar]3d64s2 Iron Name of the element

  17. Transition Metal Ions Fe2+ Cl1- Fe3+ Cl1- Cl1- Cl1- Cl1- FeCl2 FeCl3 Iron (II)Chloride Iron (III) Chloride FerrusChloride Ferric Chloride

  18. Ionic compounds Anions Br1- O2- N3- NaBr Na2O Na3N Na1+ Sodium Bromide Sodium Oxide Sodium Nitride Mg2+ MgBr2 MgO Mg3N2 Cations Magnesium Bromide Magnesium Oxide Magnesium Nitride AlBr3 Al2O3 AlN Al3+ Aluminum Bromide Aluminum Oxide Aluminum Nitride FeBr3 Fe2O3 FeN Fe3+ Iron (III) Bromide Ferric Bromide Iron (III) Oxide Ferric Oxide Iron (III) Nitride Ferric Nitride Cu1+ CuBr Cu2O Cu3N Copper (I) Bromide Cuprous Bromide Copper (I) Oxide Cuprous Oxide Copper (I) Nitride Cuprous Nitride

  19. Metallic Bonds Metal ions in electron “fluid” • An alloy is a mixture of metallic elements.

  20. Ge Ga Be H He Ca Mg B Al Li Na Si C P N As S O F Se Cl Br Ne Ar Kr K Nonmetals Share e-s with other nonmetals 1 8 2 3 4 5 6 7 Metals give e-s to nonmetals

  21. H H H H Cl Cl Cl Cl O O O O N N N N Covalent Bonds + + + +

  22. H H Cl Cl O O N N N N Covalent Bonds H-H H2 Cl-Cl Cl2 O=O O2 N2

  23. C O Covalent Bonds CO Carbon monoxide C O CO2 Carbon dioxide O C O O=C=O May modify rules to improve sound. ie - monoxide not monooxide.

  24. Polyatomic Ions PO43- Na1+ SO42- NH41+ Na1+ NH41+ NH41+ Na2SO4 (NH4)3PO4 SodiumSulfate Ammonium Phosphate

  25. Polyatomic Ions NO21- C2H3O21- Ca2+ Sn2+ C2H3O21- NO21- Ca(C2H3O2)2 Sn(NO2)2 Tin (II) Nitrite Calcium Acetate Stannous Nitrite

  26. Naming Practice AlF3 Aluminum Flouride FeF3 Iron (III) Flouride or Ferric Flouride Nitrogen Triflouride NF3 Sulfur Dioxide SO2 K2SO3 Potassium Sulfite Calcium Carbide Ca2C Calcium Carbonate CaCO3 H2CO3 Dihydrogen Carbonate

  27. Cl H Bond Polarity, Electronegativity H Cl Electrons in covalent bonds rarely get shared equally.

  28. H Cl Bond Polarity, Electronegativity • This unequal sharing results in polar bonds. + - • Slight positive side • Smaller electronegativity • Slight negative • Larger electronegativity

  29. H Li Be B C N O F Na Mg Al Si P S Cl K Ca Ga Ge As Se Br Rb Sr In Sn Sb Te I Cs Ba Tl Pb Bi Po At Electronegativity • Relative ability of atoms to attract e-. 4 - 50

  30. H Li Be B C N O F Na Mg Al Si P S Cl K Ca Ga Ge As Se Br Rb Sr In Sn Sb Te I Cs Ba Tl Pb Bi Po At Electronegativity Relative ability of atoms to attract e-. 2.20 2.55 3.04 3.44 3.98 2.19 2.58 3.16 2.96 4 - 50

  31. Cl H Bond Polarity, Electronegativity Electronegativity Difference < 0.5 Nonpolar 0.5-1.7 Polar >1.8 Ionic 2.1 3.0 H Cl d+ d- Polar Covalent

  32. Polarity, Shape CO2 Electronegativity Difference < 0.5 Nonpolar 0.5-1.7 Polar >1.8 Ionic O C O 3.5 2.5 3.5 O=C=O d- d+ d- Polar Covalent Bonds Linear Shape (180o) Nonpolar Compound

  33. Polarity, Shape e-’s in 2 directions = 180o O=C=O Linear d- d+ d- Nonpolar Compound e-’s in 3 directions = 120o d- Trigonal planar Polar Compound d+

  34. F B F F Polarity, Shape BF3 d- 4.0 F d+ 2.0 d- d- F F B 4.0 4.0 (120o) Trigonal Planar Polar Covalent Bonds Nonpolar Compound

  35. H H C H H C H H Cl Cl O H H H-O-H Polarity, Shape e-’s in 4 directions = 109.5o d+ d- Tetrahedral d+ 4 directions = 109.5o d- d+ Bent

  36. N N H H H H H H Polarity, Shape e-’s in 4 directions = 109.5o d- d+ d+ d+ Pyramidal (109.5o) Tetrahedral Configuration of Electrons Trigonal Pyramid Configuration of Atoms

  37. Tetrahedral electron-pair Geometries Tetrahedral Pyramidal Bent

  38. Properties of ionic and covalent compounds • Ionic compounds • Held together by electrostatic attraction • Ionic compounds are solids at room temp. • Exist as 3-D network of ions • Formula is simple average

  39. Attractive Forces Ionic Bonds 150 - 3000 kcal mol Melting Point NaCl 801oC Na2S 920oC MgF2 1248oC

  40. Properties of ionic and covalent compounds • Covalent compounds • Discrete molecular units • Atoms held together by bonds • Covalent compounds exist in all states • (CO2 - gas, H2O - liquid, SiO2 - solid) • Formula represents atoms in a molecule O=O

  41. d+ d+ d+ d+ d- d- d- d- H H H H Cl Cl Cl Cl Attractive Forces Dipole-Dipole 0.1 - 1 kcal mol Melting Point HCl -114oC CH3F -142oC

  42. Attractive Forces Ion-Dipole

  43. Attractive Forces Ion-Dipole

  44. d+ d- F F d+ d- F F d+ d+ d- d- F F F F Attractive Forces Dispersion Forces 0.01 kcal mol Melting Point F2 -220oC CH4 -183oC

  45. Attractive Forces d- d- d- d- O O O O d+ d+ d+ d+ d+ H H H H H H H H d+ d+ d+ Hydrogen Bonds Polar Attraction

  46. Hydrogen Bonding of Water 5 - 10 kcal mol Melting Point H2O 0oC NH3 -78oC Hydrogen Bonds

  47. Hydrogen Bonding of Water Frozen H2O: Slow moving molecules H-Bond in patterns

  48. Boiling and melting points • Chemical Bond Mp Bp • N2Nonpolar -210 -196 • O2 Nonpolar -219 -183 • NH3Polar -78 -33 • H2O Polar 0 100 • NaCl Ionic 804 ? • Melting and Boiling points • Very high for ionic compounds • Typically lower for covalent compounds

  49. Polarity and solubility • Solubility The maximum amount of a solute • that dissolves in a given solvent • Depends on the forces of attraction between molecules - intermolecular • Types of intermolecular attractions most often encountered • Dipole-Dipole • Hydrogen bonding • Van der Wall forces • General rule “Like dissolves like”

  50. Learning Check A. The Group number for sulfur is 1) 4A(14) 2) 8A(18) 3) 6A(16) B. The number of valence electrons in sulfur is 1) 4e 2) 6e 3) 8e C. The change in electrons for an octet requires a 1) gain of 2e2) loss of 2e3) a gain of 4e D. The ionic charge of sulfur is 1) 2+ 2) 2 3) 4