A Types of Chemical Bonds

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					 A. Types of Chemical Bonds
Covalent Bonding




   • A covalent bond results when
     electrons are shared by nuclei
 Formation of a Covalent Bond:

• Nature favors COVALENT bonding
  because most atoms are at lower potential
  energy when bonded to other atoms that
  they are at as independent particles.

  – The electron-proton attraction is stronger than
    the electron-electron and proton-proton
    repulsions,

  – Therefore the atoms are drawn to each other
Covalent Bonding
     • In covalent bonds atoms
       share electrons.

     • There are several
       electrostatic interactions in
       these bonds:

        – Attractions between electrons
          and nuclei

        – Repulsions between electrons
        Covalent Bonding:

• The electrons in each atom are
  attracted to the nucleus of the other.

• The electrons repel each other,

• The nuclei repel each other.

• The reach a distance with the lowest
      H2 Bond Formation:

• When hydrogen atoms are brought
  close together, there are two
  unfavorable potential energy terms:
  – Proton-proton repulsion
  – Electron-electron repulsion


• One favorable term:
  – Proton-electron attraction
       H2 Bond Formation:

• Under what conditions will the hydrogen
  molecule by favored over the separated
  atoms.

• Whichever orientation gives the lowest
  possible energy.
  – The system will act to minimize the sum of the
    positive (repulsive) energy terms and the
    negative (attractive) energy terms.

• The distance where the energy is minimum is
  called the bond length.
As the atoms approach each other, the energy
decreases until the distance reaches 0.074 nm and
then begins to increase again due to repulsions.
 Characteristics of the Covalent
             Bond:

• Bond Length - the distance between two
  bonded atoms at their minimum potential
  energy
  – Average distance between 2 bonded atoms.


• Bond Energy - the energy required to
  break a chemical bond and form neutral
  isolated atoms.
         Electrons & Bonds

• single bond one pair of electrons is
  shared.

• double bond two pair of electrons are
  shared.

• triple bond three pair of electrons are
  shared.
     Covalent Bond Strength



• Most simply, the strength of a bond is
  measured by determining how much energy
  is required to break the bond.

• This is the bond enthalpy.

• The bond enthalpy for a Cl-Cl bond, D(Cl-Cl),
          Average Bond Energy

• This table lists the
  average bond
  enthalpies for many
  different types of
  bonds.



• Average bond
  enthalpies are
  positive, because
          Average Bond Energy

NOTE: These are
 average bond
 enthalpies, not
 absolute bond
 enthalpies; the
 C-H bonds in
 methane, CH4,
 will be a bit
 different than the
 C-H bond in
 chloroform,
    Partial Ionic Character

• There are probably no totally ionic bonds
  between individual atoms.

• Calculate % ionic character using
  electronegativity values.
  A. Types of Chemical Bonds
 Non-Polar Covalent Bonding

• A non-polar covalent bond results when
   electrons are not shared unequally by
                   nuclei
         Covalent Bonding
• Electrons are shared by atoms.

• These are two extremes.

• In between are polar covalent bonds.

• The electrons are not shared evenly.

• One end is slightly positive, the other
 A. Types of Chemical Bonds
Polar Covalent Bonding

 • A polar covalent bond results when
electrons are shared unequally by nuclei
– One atom attracts the electrons more than the other atom
       B. Electronegativity
• The polarity of a bond depends on the
         difference between the
  electronegativity values of the atoms
            forming the bond
        Polar Covalent Bonds

                          • Though atoms often
                            form compounds by
                            sharing electrons, the
                            electrons are not
                            always shared equally.

• Fluorine pulls harder on the electrons it
  shares with hydrogen than hydrogen does.
• Therefore, the fluorine end of the molecule
  has more electron density than the hydrogen
  end.
               Electronegativity
• Electronegativity – the relative ability of an atom in a
  molecule to attract shared electrons to itself
   – Increases from left to right across a period
   – Decreases down a group of representative elements
   The Relationship Between
Electronegativity and Bond Type
       Polar Covalent Bonds




The greater the difference in electronegativity,
the more polar is the bond.
Polar Molecules
                  The Octet Rule:
• Chemical compounds tend to form so that each atom, by
  gaining, losing, or sharing electrons, has 8 electrons in its
  highest occupied energy level.

• Noble Gas Elements – inert; 8 valence electron provide stability
  (except helium)

• S and P orbitals are completely filled with eight electrons by sharing
  electrons

• Exceptions: Hydrogen, boron, BF3
    – Other elements can be surrounded by more than eight electrons when
      they combine with the highly electronegative elements fluorine, oxygen,
      and chlorine.
       Electron- Dot Notation:

• An electron-configuration notation in which
  only the valence electrons of an atom of a
  particular element are shown, indicated by
  dots place around the element’s symbol.

  – EX: Fluorine, Hydrogen, Nitrogen
        Lewis Dot Structures:

• Formulas in which atomic symbols represent
  nuclei and inner-shell electrons

• Dot-pairs or dashes between 2 atomic symbols
  represent electron pairs in covalent bonds

• Dots adjacent to only one atomic symbol
  represent unshared electrons called Lone Pairs
  – They do not participate in bonding

• EX: H2, F2
                       Lewis Dot Structures:
1. Determine the type & number of atoms in the molecule

2. Write the electron-dot notation for each type of atom in the molecule

3. Determine the total number of valence electrons in the atoms to be combined

4. Arrange the atoms to form a skeleton structure for the molecule, and connect the
   atoms by electron-pair bonds
     1. If carbon is present, then it will be the central atom
     2. Otherwise, the lease electronegative atom is central (except hydrogen)

5. Add unshared pairs of electrons so that each hydrogen atom shares a pair of
   electrons and each other nonmetal is surrounded by eight electrons

6. Count the electrons in the structure to be sure that the number of valence electron
   used equals the number available.

7. If too many electrons have been used, subtract one or more lone pairs until the total
   number of valence electrons is correct. Then move one or more lone electron pairs to
   existing bonds between non-hydrogen atoms until the outer shells of all atoms are
   completely filled.
         Lewis Dot Structures:

• CH3I, NH3, H2S, H2O, CH4
      Resonance Structures:
• Occurs when some molecules and ions
  cannot be represented correctly by a
  single Lewis structure.

  – EX: See Octet Rule/Resonance Structure
    WS
         Molecular Geometry:

• The properties of molecules depend not
  only on the bonding of atoms but also on
  molecular geometry (3D shape)

• Chemical formula doesn’t tell us much
  about the geometry of a molecule.

• Two different theories were developed to
  explain the geometry of a molecule.
  – VSEPR Theory = accounts for molecular bonding
    angles
            VSEPR Theory:
• Can be used to help determine the geometry, or
  shape, of a molecule based on the orbitals that
  contain the valence electrons.

• “Valence-Shell Electron-Pair Repulsion”:
  repulsion between the sets of valence-level
  electrons surrounding an atom causes these
  sets to be oriented as far apart as possible.

• Shared electron pairs are oriented as far away
  from each other as possible
           VSEPR Theory


• VSEPR theory is useful for explaining the
  shapes of molecules.

• However, it does not reveal the
  relationship between a molecule’s
  geometry and the orbitals occupied by its
  bonding electrons (arrangement).
• Use the Molecular Geometry Chart to
  predict the shape and arrangement of
  each of the following molecules and
  ions:
• a. AsF5
• b. SeF6
• c. CF4
• d. NO3-
       II. Hybridization Theory:

• It is used to explain the geometry of
  molecular orbitals occupied by bonding
  electrons (the arrangement).

• Sometimes orbitals become rearranged
  when the atom forms covalent bonds.

• When this happens, a different model is
  used.
            EX: CH4 Methane
• Carbon has 4 valence electrons (2 in the 2s
  orbital and two in 2 separate sp orbitals)
• Tetrahedral geometry
• How does carbon form four equivalent,
  tetrahedrally arranged covalent bonds by
  orbital overlap with four other atoms?
  – 2s and 2p orbitals have different shapes.
  – To achieve four equivalent bonds, carbon’s 2s and
    three 2p orbitals hybridize to form four new, identical
    orbitals called sp3 orbitals.
  – The superscript 3 means that three p orbitals were
             Hybrid orbitals

• Orbitals of equal energy produces by the
  combination of two or more orbitals on the
  same atom.
  – sp = linear geometry, orbitals are 180 apart

  – sp2 = Trigonal planar geometry, orbitals are
    120 apart

  – sp3 = Tetrahedral geometry, orbitals are
    109.5  apart.
   Arrangement of Molecules
• The arrangement is determined by the # of
  regions of electron density. (See the Molecular
  Geometry Chart)

• Put lobes around density regions of electron
  density.

• Count regions to determine arrangement

• Count bonds and lone pairs to determine the
  arrangement.
            Practice WS
• ICA/HW: Types of Lewis Structures

				
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