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Bonding Introduction

Bonding Introduction. Octet rule Types of bonds Lewis structures Geometry. Video 5.1. Types of Bonds. Octet Rule Review. Atoms bond with other atoms by sharing or transferring electrons in order to achieve a stable octet (8 valence electrons). Bonding creates stability!

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Bonding Introduction

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  1. Bonding Introduction Octet rule Types of bonds Lewis structures Geometry

  2. Video 5.1 Types of Bonds

  3. Octet Rule Review Atoms bond with other atoms by sharing or transferring electrons in order to achieve a stableoctet(8 valence electrons). Bonding creates stability! *When bonds are formed energy is ___________. *When bonds are broken energy is ___________. released absorbed

  4. Ionic Bonds • Transfer of electrons from the cation to the anion (metal to nonmetal). • High melting point and boiling point • Mostly hardcrystalline solids • Conduct as liquid (either melted or dissolved) due to mobile ions.

  5. Ionic Bonds Sodium Chloride: NaCl (table salt) properties: • Hard • Solid crystals • High melting point, forget boiling! • Liquid phase conducts (electrolytes are salts)

  6. Metallic Bonds • Metals only • All metals lose their valence electrons and form a sea of electrons • High melting point and boiling point • Insoluble in water • Always able to conduct heat and electric due to mobile electrons • Malleable • Ductile

  7. Metallic Bonds Copper (Cu) properties: • Hard solid • High melting point, forget boiling! • Malleable and ductile • Conductor • Can’t dissolve

  8. Metallic BondsSea of electrons Copper (I) ions Copper (II) ions

  9. Covalent Bonds (Molecular) • Nonmetals only • Share electrons between atoms • Low melting point and boiling point • Never conduct heat or electricity • Soft solid or gas

  10. Covalent Bonds Dextrose C6H12O6 (Sugar) properties: • Soft • Melts easily in sauce pans for caramel • Doesn’t conduct (nonelectrolyte)

  11. What type of bond is created? M+ NM = Ionic • Ca + O • K + Br • S + Cl • I + S • Li + Mg • Ba + S M + NM = Ionic NM + NM = Covalent NM + NM = Covalent M + M = Metallic M + NM = Ionic

  12. Video 5.2 Ionic Compounds

  13. Review: Find the ionic formula: - KBr + • K + Br • Mg + Cl • Na + S • Ca + S +2 - MgCl2 + Na2S -2 CaS +2 -2

  14. Draw Lewis structures: KBr MgCl2 Na2S CaS

  15. Which subatomic particle is involved in bonding? Electrons only!

  16. Geometry of ionic crystals Ionic crystal Ions

  17. Video 5.3 Covalent Compounds

  18. Covalent Lewis Structures Rules: CCl4 C: 4 + 4Cl: 7 = 32 valence e- • Add up all valence e- • Draw a skeletal structure with bonds between elements. Least frequent element in the middle. • Subtract 2e- from total for each bond drawn. • Draw in remaining e- to fill each atom’s octet. • Evaluate: each atom should have 8 e- only. Cl Cl—C—Cl Cl 32-8=24

  19. VSEPR “Valence shell electron pair repulsion” is a model for molecules. Lone electron pairs are repelled by one another and should be placed as far apart as possible.

  20. Geometry • Linear: The molecule is on one plane (flat) such as CO2 or H2. • Bent: The molecule is bent at angle like H2O due to unshared electrons and two bonding pairs on the central atom.

  21. Geometry • Pyramidal: The molecule has a triangular shape like NH3 due to a lone pair and three bonding pairs on the central atom. • Tetrahedral: The molecule has four bonding pairs and no lone pairs on the central atom like CH4.

  22. Examples: • Draw the following molecules and identify their geometry: • PCl3 • SiCl2H2 • Br2 • H2S pyramidal tetrahedral linear bent

  23. Video 5.4 Bond Polarity

  24. Bond Polarity The earth has two poles; North and South. A magnet also has two poles. Bonds may have two poles. This means one element is charged different than the other. If a bond is polar, the two elements have different electronegativities. The element with a higher electronegativity will be more negative.

  25. Bond Polarity

  26. Bond Polarity

  27. Nonpolar Bond

  28. Bond Polarity

  29. Ionic, polar or nonpolar? P P • C-Br • Na-S • C-C • H-O • K-O • Be-B • As-O • N-O • C-O • F-F • S-C • N-H P I P NP NP P I NP I P

  30. CovalentBonding • If 2 atoms or more form a bond with the same electronegativity the bonds are nonpolar and they share e- equally. ( F-F ) • If there is an electronegativity difference between bonded atoms, the bonds are polar and e- are pulled toward the more electronegative atom. (H-F) • If a bond is polar, the molecule will have a slightly negative and slightly positive side, like 2 poles of a magnet.

  31. Video 5.5 Molecular Polarity

  32. Molecular Polarity • A polar molecule will be asymmetrical. • A nonpolar molecule will have a symmetrical shape or all nonpolar bonds.

  33. Molecular Polarity Which are polar molecules? Show charges. NP + + P P + + - - - NP NP - + - - - - + -

  34. Molecular Polarity Water is polar, and like dissolves like, so only polar molecules are soluble in water. Polar molecules are also attracted to an electric field.

  35. Molecular Polarity • As you can see, normally polar molecules are unaligned. • When a electric source comes by, the molecules quickly align themselves.

  36. Video 5.6 IMF

  37. IMF • Intramolecular forces is another name for bonds, that keep elements together in compounds. • Intermolecularforces of attraction are weaker than bonds, but are responsible for holding a substance together (multiple molecules in a confined area).

  38. IMF • The stronger the IMF, the tighter the structure (solid). The melting and boiling points will be high. • The weaker the IMF, the looser the structure (gas). The melting and boiling points will be low.

  39. Dipole-Dipole • Dipole-Dipole attractions are strong forces betweenpolarmolecules. It is like static holding the + and – charges together.

  40. Hydrogen Bonding A special case: Hydrogen Bonds are the strongestbonds between Hydrogen and very electronegative atoms such as F, O and N. (H bonds are FON!) For example, H2O and HF, due to their polarity, they will attract each other.

  41. London Dispersion Forces (LDF) The weakest attraction betweennonpolarmolecules occur because electrons temporarily shift creating a temporary + and – charge. The more electrons the compound has, the stronger the force is.

  42. Summary • From weak to strong: • Nonpolar LDF • Polar Dipole Dipole forces • Hydrogen bonds • Covalent Bonds • Ionic Bonds • Metallic Bonds

  43. Class Notes

  44. Show the individual and bonded Lewis structures: • Li and F • Mg and O • Be and S • What did all the cations do? • What did all the anions do? • Which of the subatomic particles were changed and how were they changed?

  45. Type of Bonding? • CaCl2 • CO2 • H2O • BaSO4 • K2O • NaF • Na2CO3 • CH4 • SO3 • LiBr • MgO • NH4Cl • HCl • KI • NaOH • NO2 • AlPO4 • FeCl3 • P2O5 • N2O3

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