Ch 9-section-1
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Ch 9-section-1 Presentation Transcript

  • 1. Chapter 9
    Covalent Bonding
  • 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. What is a covalent bond?
    Chemical bond that results from sharingof valence electrons
    Occurs b/w nonmetal & a nonmetal
    Balance b/w attractive and repulsive forces
    2 Hydrogen Atoms
    Sharing their 1 Ve-
  • 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. H
    Cl
    Unshared or
    Lone pair (LP)
    shared or Bond pair
    Single Covalent Bonds
    A single covalent bond –Atom shares 1 pair (2) electrons.
    Shared pairs – both elements count the electron pair to achieve octet
    Lonepairs– 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.
    . .
    . .
    . .
  • 6. In the Fluorine Molecule…..
    How many bonding pairs are there in each?
    1
    How many lone pairs are there each?
    3
  • 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. Covalent Bond Formation in Hydrogen
    • Increased overlap brings the electrons and nuclei closer together while simultaneously decreasing electron-electron repulsion.
    • 9. However, if atoms get too close, the internuclear repulsion greatly raises the energy.
  • The attractive and repulsive forces in covalent bonding must be balanced.
  • 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. Bond Energy - 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. Comparing Bond Length and Bond Strength
    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. Sigma () Bonds
    Sigma bonds are characterized by
    Head-to-head overlap.
    Cylindrical symmetry of electron density about the internuclear axis.
  • 14. Pi () Bonds
    Pi bonds are characterized by
    Side-to-side overlap.
    Electron density above and below the internuclear axis.
  • 15. Single Bonds vs. Multiple bonds
    Single bonds are always  bonds, because  overlap is greater, resulting in a stronger bond and more energy lowering.
    In a multiple bond one of the bonds is a  bond and the rest are  bonds.
  • 16. Orbital overlap
  • 17. Lewis Dot Structures
    Determine the number of Valence e- for all atoms in the molecule
    Divide the Ve- by 2 to get pairs (2 dots or 1 line)
    Decide on central atom (least electronegative or furthest to the left).
    Hydrogen & halogens are terminal atoms
    Connect all atoms to the central atom by a bonding pair (single line)
    Place remaining pairs around all atoms before moving on to central atom.
    Check for octet (not H)
    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).