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Chemical Bonding

Chemical Bonding. Valence Electrons. The outermost electrons. Lost, gained, or shared during chemical bonding. 1. 18. 2. 16. 14. 13. 15. 17. 2. 3. 5. 7. 1. 8. 4. 6. Vary. Terminology. ATOMS Neutral Particles Na, O, Al, I. IONS Charged Particles

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Chemical Bonding

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  1. Chemical Bonding

  2. Valence Electrons The outermost electrons Lost, gained, or shared during chemical bonding 1 18 2 16 14 13 15 17 2 3 5 7 1 8 4 6 Vary

  3. Terminology ATOMS Neutral Particles Na, O, Al, I IONS Charged Particles Na+1, O-2, Al+3, I-1

  4. Octet Rule An atom loses, gains, or shares electrons in order to get 8VALENCE electrons GAIN ELECTRONS LOSE ELECTRONS Metals + charge Nonmetals - charge OXIDATION NUMBER Charge and # (electrons lost/gained) (+/-)

  5. Oxidation Numbers 1 18 2 13 14 15 16 17 -4 -1 +1 0 +2 -2 -3 +3 +1, +2, +3, +4, +5, +6 +4

  6. IONS Charged Particles + Ions Metals NAME - Same as on P.T. Transition metal, Sn, Pb get a Roman Numeral to indicate their charge Ag ALWAYS +1, Zn ALWAYS +2 - Ions Nonmetals NAME - Change ending on P.T. to ”–ide” Polyatomic Ions Group of elements that are charged Name, symbols, & charge have to be learned

  7. Ionic Bond • The attraction of a + ion to a – ion • The # of + & - charges come together to balance (cancel) each other out Example: Na+1 and Cl-1 vs. Na+1 and O-2 Need 1 Na Need 1Cl FORMULA = NaCl Need 2 Na Need 1 O FORMULA = Na2O Use SUBSCRIPTS (in the formula) to indicate the number of ions used to make the charges cancel out

  8. Formulas and Names Formulas Symbols and Subscripts NaCl Cu(OH)2 Fe2O3 Names Full name of each ion in the compound Na – sodium, Cl – chloride NaCl – sodium chloride Cu – copper (II), OH – hydroxide Cu(OH)2 – copper (II) hydroxide Fe – iron (III), O – oxide Fe2O3 – iron (III) oxide

  9. Ionic Bonds • Electrons areTRANSFERREDbetween the elements +lose electrons, -gain the electrons lost + = metal - = nonmetal/polyatomic

  10. Properties of Compounds Ionic Compounds Metal + Nonmetal High melting points Most dissolve in water Solutions - high conductivity Low volatility Covalent Compounds Nonmetal + nonmetal Low melting points Some dissolve in water Little to no conductivity More volatile Metallic Bonds Metal atoms only Melting points range Do not dissolve in water Conducts electricity in metal form Not volatile

  11. Electron Dot Diagrams Represents the number of VALENCE ELECTRONS in an atom C 4 valence electrons C F 7 valence electrons F

  12. Covalent Compounds • SHARE valence electrons in order to get 8 • Typically between 2 NONMETALS F Shared (pair) Electrons F C F Lewis Structure Lone (pair) Electrons F

  13. Lewis Structures Bond so that elements get 8 electrons • Exceptions ---- H – gets 2 electrons Be – gets 4 electrons B – gets 6 electrons • Sharing 2 electrons ----- Single bond • Sharing 4 electrons ----- Double bond • Sharing 6 electrons ----- Triple bond **Double and triple bonds typically occur between C, N, O, and S

  14. VSEPR – Ms. J’s http://mrsj.exofire.net/chem/notes.htm#molecularstructure

  15. Polarity • Some elements like to hang out with electrons more than others • Based upon their electronegativities • Even distribution of electrons = NONPOLAR • Uneven distribution of electrons = POLAR Bent, pyramidal = ALWAYSPOLAR Linear, planar, tetrahedral = POLAR if outside elements are different = NONPOLAR if outside elements are the same

  16. For the following compounds:1. Draw the Lewis Structure2. Identify the shape using VSEPR3. Determine if the molecule is polar or nonpolar NH3 H2O H2S BF3 CO2 Cl2 CO SiF4 CH4 HF CH2O PCl3 H2 CH2Cl2 SiO2 NO2-1

  17. Naming and Writing Formulas Covalent Compounds First element – name stays same gets a prefix to indicate MORE THAN one element (comes from subscript) Second element – change ending to “ide” ALWAYS put a prefix to indicate how many elements are in the compound (from subscript) Prefixes

  18. Practice MgSO4 CO N4O8 LiCl P2O5 Cu3(PO4)2 Boron trifluoride Iron (III) sulfide Calcium carbonate Dinitrogen trioxide Strontium nitrate Ammonia

  19. Review • Rags to Riches Game • Ionic Bonding Practice • Ionic Bonds using polyatomic ions

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