Ch 9-section-1


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  • Exs:SiF4NCl3AsH3SO2CO2CO
  • Ch 9-section-1

    1. 1. COVALENT BONDINGChapter 9
    2. 2. Review…. What is a chemical bond?  Force that holds two atoms together What is an ionic bond?  An electrostatic force that holds oppositely charged particles together in an ionic compound  Compounds formed from metal & nonmetal Forms when….? What are atoms always trying to achieve?  Stability  Complete set of valence electrons… OCTECT
    3. 3. What is a covalent bond? Chemical bond that results from sharing of valence electrons  Occurs b/w nonmetal & a nonmetal  Balance b/w attractive and repulsive forces2 Hydrogen Atoms Sharing their 1 Ve-
    4. 4. Molecules Compound made when 2 or more atoms are bonded covalently Diatomic molecules  In nature, sometimes two atoms of the same element are more stable when they are covalently bonded than the individual atom alone…  BrINClHOF (pronounced “Brinkle Hoff”)  Br2 I2 N2 Cl2 H2 O2 F2
    5. 5. Single Covalent Bonds A single covalent bond –Atom shares 1 pair (2) electrons.  Shared pairs – both elements count the electron pair to achieve octet  Lone pairs – pair of electrons that are not shared b/w the atoms Lewis structures- Use electron dot diagram to show how atoms are arranged in a molecule. .. H Cl .. Unshared or Lone pair (LP) shared or Bond pair
    6. 6. In the Fluorine Molecule….. How many bonding pairs are there in each?  1 How many lone pairs are there each?  3
    7. 7. Multiple Covalent Bonds Double covalent bonds share two pairs of electrons.  CO2 O=C=O Triple covalent bonds share three pairs of electrons.  N2 :N=N:
    8. 8. Covalent Bond Formation in Hydrogen  Increased overlap brings the electrons and nuclei closer together while simultaneously decreasing electron- electron repulsion.  However, if atoms get too close, the internuclear repulsion greatly raises the energy.
    9. 9. The attractive and repulsiveforces in covalent bonding mustbe balanced.
    10. 10. –Bond Length - In general, the closer the electrons are held by the atoms, the shorter the bond length and the higher the bond energy. –Multiple bonds result in stronger, shorter bonds.
    11. 11. –BondEnergy - The amount of energy required to break a bond. The greater the energy, the stronger the bond.–Bond breaking is an endothermic process, so bond breaking enthalpies are positive.
    12. 12. Comparing Bond Length and BondStrength Using the periodic table, but not Tables 9.2 and 9.3, rank the bonds in each set in order of decreasing bond length and bond strength: (a) S F, S Br, S Cl (b) C = O, C O, C O
    13. 13. Sigma ( ) Bonds Sigma bonds are characterized by  Head-to-head overlap.  Cylindrical symmetry of electron density about the internuclear axis.
    14. 14. Pi ( ) Bonds Pi bonds are characterized by  Side-to-side overlap.  Electron density above and below the internuclear axis.
    15. 15. Single Bonds vs. Multiple bondsSingle bonds are always  In a multiple bond one bonds, because of the bonds is aoverlap is bond and the rest aregreater, resulting in a bonds.stronger bond and moreenergy lowering.
    16. 16. Orbital overlap
    17. 17. Lewis Dot Structures1. Determine the number of Valence e- for all atoms in the molecule a. Divide the Ve- by 2 to get pairs (2 dots or 1 line)2. Decide on central atom (least electronegative or furthest to the left). a. Hydrogen & halogens are terminal atoms3. Connect all atoms to the central atom by a bonding pair (single line)4. Place remaining pairs around all atoms before moving on to central atom.5. Check for octet (not H) a. If atom does not have an octet, move lone pairs from a terminal atom to create a double or a triple bond (except grp 7).