Introduction to Chemical Bonding

							Introduction to
Chemical Bonding
  Notes – Page 1
Electronegativity – the ability of an atom in a bond
  to attract electrons.
Ion – a charged particle. Obtained the charge by
  gaining or losing electrons.
  Positive ion – achieved by losing electron(s).
  Negative ion – achieved by gaining electron(s).
Ionic Bond – results from a transfer of electrons.
  One atom takes the electrons from the other
  forming ions. Held together by electrostatic
  force. (Opposites attract.)
Covalent Bond – results from the sharing of
 electrons between 2 atoms.
    Non-polar covalent – bond in which electrons are
    equally shared.
   Polar covalent – bonds that have an uneven
    distribution of charge due to unequal sharing of
    electrons, Based on electronegativity.
  Bonds between two unlike atoms are never
    completely ionic and are rarely completely covalent.
   The degree to which bonds are ionic or covalent
    depends on the electronegativity differences of the
    bonded atoms. (Table p. 304)
Bond characteristic:
Electronegativity Chart
Examples:
H - Cl Bond   2.1 – 3.0   = 0.9

H–F           2.1 – 4.0   = 1.9

H–H           2.1 – 2.1   = 0.0
Covalent Bonding
 Molecule – resulting particle when 2 or
  more atoms bond covalently.
 Diatomic Molecule – molecule consisting
  of 2 atoms
 Single bond – covalent bond produced by
  the sharing of one pair of electrons.
 Double bond – sharing of two pairs of
  electrons
 Triple bond – sharing of 3 pairs of
  electrons
Covalent Bonding cont’d
 Bond Length – the average distance
  between 2 bonded atoms.
 Bond Energy – energy required to break a
  chemical bond and form neutral atoms.
 Polyatomic ion – a charged group of
  covalently bonded atoms.
Lewis Structures
   Structural formula – indicated the kind,
    number, arrangement, and bonds of the
    atoms in a molecule.
     Atomic  symbols represent inner-shell
      electrons & nuclei
     Dashes between 2 atomic symbols represent
      shared electron pairs in covalent bonds
     Dots adjacent to only one atomic symbol
      represent unshared or lone electrons.
       -   Represents a bond (2 electrons)
           Represents an unshared electron
Lewis Structures cont’d
 Central atom is the atom furthest to the left
  on the periodic table (Except hydrogen)
 All atoms need to have 8 electrons in their
  outer level to be stable except H
 H needs 2 electrons in its outer level
Drawing Lewis Structures
Step 1:    Find the total number of valence
  electrons that the elements have available.
  Example:CH4

Element    # Atoms    # Valence e-   TOTAL valence
                      per atom       electrons
     C          1           4        1(4)=4
     H          4           1        4(1)=4
  Drawing Lewis Structures
    Step 2 Find the total # of electrons needed to
     have a complete outer shell. (Remember: All
     elements need 8 EXCEPT H, which only needs
     2)
Elem # of atoms         # of e-    Total # of valence e- (#
 ent                  needed in atoms) x (# valence e-
                       o.s. per    per atom)
                         atom
  C         1              8               1(8) = 8

 H          4             2                4(2) = 8
    Drawing Lewis Structures
   Step 3: Subtract the “HAVES” from the
    “NEEDED”. This is the # of electrons that must
    be SHARED.
                                      Needed: 16
                                      Have : -8
    electrons that must be shared:             8
Lewis Structures
   Step 4 Find the number of bonds:
     Dividethe # of SHARED electrons by 2 to get
      the number of bonds.
       8   =   4 bonds (represented by 1 dash per bond)
       2
Step 5 Draw the structure.
Put the atom furthest left on the periodic
 table (Except H) in the center. Fill in the
 bonds (dashes) and lone e- (dots)
Drawing of CH4