Chapter 8 - Covalent Bonding - PowerPoint by JnNi205

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									    Chapter 16

    Covalent bonding




1
             How does H2 form?
       The nuclei repel




                  +        +

2
           How does H2 form?
     The nuclei repel
     But they are attracted to electrons
     They share (or fight over) the electrons




                 +              +

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              Covalent bonds
     Bonding between two nonmetals.
     Nonmetals hold onto their valence
      electrons, so they can’t give away
      electrons to bond.
     They still want noble gas configurations.
     Share (or fight over) valence electrons
      with each other.
     By sharing both atoms get to count the
      electrons toward noble gas
      configuration.
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              Covalent bonding
       Fluorine has seven valence electrons




                F
5
            Covalent bonding
     Fluorine has seven valence electrons
     A second atom also has seven




              F              F
6
            Covalent bonding
     Fluorine has seven valence electrons
     A second atom also has seven
     By sharing electrons




              F              F
7
            Covalent bonding
     Fluorine has seven valence electrons
     A second atom also has seven
     By sharing electrons




               F            F
8
            Covalent bonding
     Fluorine has seven valence electrons
     A second atom also has seven
     By sharing electrons




                F F
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             Covalent bonding
      Fluorine has seven valence electrons
      A second atom also has seven
      By sharing electrons




                  F F
10
             Covalent bonding
      Fluorine has seven valence electrons
      A second atom also has seven
      By sharing electrons




                   F F
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             Covalent bonding
      Fluorine has seven valence electrons
      A second atom also has seven
      By sharing electrons
      Both end with full orbitals



                                     8 Valence
                   F F               electrons


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             Covalent bonding
      Fluorine has seven valence electrons
      A second atom also has seven
      By sharing electrons
      Both end with full orbitals



 8 Valence
 electrons         F F
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           Single Covalent Bond
      A sharing of two valence electrons.
      Only nonmetals and hydrogen.
      Different from an ionic bond because
       they actually form groups called
       molecules.
      Two specific atoms are joined in
       covalent bonding.
      In an ionic solid you can’t tell which
       atom the electrons moved from or to.
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     How to show how they formed
      It’s like a jigsaw puzzle.
      I have to tell you what the final formula
       is.
      You put the pieces together to end up
       with the right formula.
      For example- show how water is formed
       with covalent bonds.



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              Water
         Each hydrogen has 1 valence

     H    electron
         Each hydrogen wants 1 more

         The oxygen has 6 valence
          electrons

     O   The oxygen wants 2 more
         They share to make each other
          happy
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                     Water
      Put the pieces together
      The first hydrogen is happy
      The oxygen still wants one more




                 HO
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                     Water
      The second hydrogen attaches
      Every atom has full energy levels




                 HO
                  H
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19
             Diatomic Molecules
      Two atoms of same element joined
       together
      There are elements that come in pairs;
       they NEVER go out alone.
      These are hydrogen, nitrogen, oxygen,
       fluorine, chlorine, bromine, iodine.
      1 + 7 on periodic table

        Double your fun with HOFBrINCl

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               Multiple Bonds
      Sometimes atoms share more than one
       pair of valence electrons.
      A double bond is when atoms share two
       pair (4) of electrons.
      A triple bond is when atoms share three
       pair (6) of electrons.




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         Carbon dioxide
           CO2 - Carbon is central
            atom ( I have to tell you)

     C     Carbon has 4 valence
            electrons
           Wants 4 more
           Oxygen has 6 valence

     O      electrons
           Wants 2 more


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                 Carbon dioxide
        Attaching 1 oxygen leaves the oxygen 1
         short and the carbon 3 short




                       CO
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                 Carbon dioxide
        Attaching the second oxygen leaves
         both oxygen 1 short and the carbon 2
         short




                OC O
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                 Carbon dioxide
        The only solution is to share more




               O CO
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                 Carbon dioxide
        The only solution is to share more




              O CO
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                 Carbon dioxide
        The only solution is to share more




              O CO
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                 Carbon dioxide
        The only solution is to share more




              O C O
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                 Carbon dioxide
        The only solution is to share more




              O C O
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               Carbon dioxide
      The only solution is to share more
      Requires two double bonds
      Each atom gets to count all the atoms in
       the bond



            O C O
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               Carbon dioxide
      The only solution is to share more
      Requires two double bonds
      Each atom gets to count all the atoms in
       the bond
                   8 valence
                   electrons

            O C O
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               Carbon dioxide
      The only solution is to share more
      Requires two double bonds
      Each atom gets to count all the atoms in
       the bond
          8 valence
          electrons

            O C O
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               Carbon dioxide
      The only solution is to share more
      Requires two double bonds
      Each atom gets to count all the atoms in
       the bond
                               8 valence
                               electrons

            O C O
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                How to draw them
        Add up all the valence electrons.
        Write the symbols of all elements in their
         places.
        Place two electrons (dots) between all atoms
         (they MUST be bonded).
        Subtract the electrons you have used from
         the total.
        Place remaining electrons (dots) on individual
         atoms to make an octet.
        If you can’t do this with the remaining
         electrons, add another bond and try again.

34
          Examples
          NH3

     N    N - has 5 valence
           electrons
          H – each has 1

     H     valence electron
          NH3 has 5+3(1) = 8

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                  Examples
      Draw in the bonds
      8e- - 6e- = 2e-
      Place remaining 2 electons on N, since
       it is not full.
                   H
                 H NH
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                   Examples
      All 8 electrons are accounted for
      Everything is full




                   H
                 H NH
37
                   Examples
      HCN C is central atom
      N - has 5 valence electrons
      C - has 4 valence electrons
      H - has 1 valence electron
      HCN has 5+4+1 = 10




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                      HCN
      Put in single bonds
      Remaining 6 electrons will not complete
       both octets.
      Add a bond - must go between C and N,
       because H is full.



                 HC N
39
                       HCN
      Put in double bond
      Leaves 4 more electrons – still won’t fill
       C and N
      Add another bond - must go between C
       and N



                   HC N
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                       HCN
      Put in triple bonds
      2 electrons left
      Must go on N, C is full now
      Molecule is complete (all atoms full, 10
       electrons used.)


                  HC N
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         Another way of indicating
                  bonds
      Often use a line to indicate a bond
      Called a structural formula
      Each line is 2 valence electrons

                      is the

     HOH H O H        same
                      as


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      Structural Examples
                  C has 8 electrons
                   because each

     H C N         line is 2 electrons
                  Ditto for N


     H            Ditto for C here
       C O        Ditto for O


     H
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     Polar Molecules

     Molecules with ends




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              Polar Molecules
      Molecules with a positive and a
       negative end
      Requires two things to be true
      The molecule must contain polar bonds
      This can be determined from differences
       in electronegativity.
     Symmetry can not cancel out the
       effects of the polar bonds.
       Must determine geometry first.

45
                  Polar Bonds
      One atom in the bond is pulling harder
       on the electrons that are shared than
       the other. As a result, the electrons are
       shared unequally.
      In a nonpolar bond, both atoms have
       equal or nearly equal strength, so
       electrons are shared equally.



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                         Polarity
        Measured in terms of differences in
         electronegativity
           Remember that electronegativity is the
            amount of pull on the electrons!
        Find electronegativity of elements on periodic
         table.
        Subtract to find difference:
          – 0.0 to 0.4 = nonpolar covalent bond
          – 0.5 and up = polar covalent bond
        Ionic bonds will have greater differences yet,
         usually above 1.7

47
                Showing Polarity
      There are two ways to show that a bond
       is polar
      Use an arrow with a positive end:


        You can use a partial positive symbol
         and a partial negative symbol:

                  +        -
        Both mean the same thing!!!
48
             Is it polar?
      HF
      H2O
      NH3
      CH4
      CO2
      N2


49
     Intermolecular Forces

     What holds molecules to each
                 other



50
          Intermolecular Forces
      These are the forces BETWEEN
       molecules that hold them together.
      INTRAmolecular bonding is what holds
       the atoms within the molecule together
      They are what make solid and liquid
       molecular compounds possible.




51
           van der Waal Forces
      The weakest are called van der Waal’s
       forces
      EVERY molecule has these forces – no
       exceptions
      There are two kinds
        – Dispersion forces
        – Dipole Interactions
      You just need to know the main group


52
              Disperson forces
      Temporary dipole is formed by moving
       electrons
      Bigger molecules = more electrons
      More electrons means stronger forces,
       since electrons are more likely to shift
       and create temporary dipole
          • Fluorine is a gas
          • Bromine is a liquid
          • Iodine is a solid
53
              Dipole interactions
      Occur when polar molecules are
       attracted to each other.
      Slightly stronger than dispersion forces.
      Opposites attract but not completely
       hooked like in ionic solids.

         +   -           +    -
         H F               H F
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          Dipole Interactions




            +   -



     +     -
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             Hydrogen bonding
      Are the attractive force caused by
       hydrogen bonded to F, O, or N.
      F, O, and N are very electronegative so
       it is a very strong dipole.
      The hydrogen partially share with the
       lone pair in the molecule next to it.
      The strongest of the intermolecular
       forces.


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      Hydrogen Bonding



     + -
     H O
        H +



57
      Hydrogen bonding


     H O
        H




58
     Circle those elements with
         Hydrogen Bonding:

     H2O          HCl
     CH4          HF
     NH3          H2S

59
     Circle those elements with
         Hydrogen Bonding:

     H2O          HCl
     CH4          HF
     NH3          H2S

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