Chapter 9 Chemical Bonding

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					Chapter 9 Chemical Bonding

        Bonding of Atoms
                           Bonding Extremes
• In chapter 4 you learned that there were two types of
• Ionic Bonds – Electrons are transferred from one atom
  to another forming ions.
• Covalent (molecular) Bonds – Electrons are
shared between atoms
• This is simple, but not all bonds are one type or
the other. Instead, there is a continuum from one
type to the other.
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                Examples of the Extremes

            •Na + Cl            [Na]+ + [ Cl ]-
            •This is an example of the ionic

            • C + O + O                 O C O
            • This is an example of the
            covalent extreme.
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                A tug of war between atoms

   • Picture two atoms, each one is represented by a person
     below. Both pull on the other’s valence electrons. Neither is
     strong enough to pull the electrons totally away, but one can
     pull a bit harder and the electrons move toward the harder
d-                                                              d+

                           e- e-
                The electrons are shared, but not equally, the
                resulting molecule has a somewhat negative
                end and a somewhat positive end.
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• The relative attraction of an atom for the
  electrons in a bond is called electronegativity
• Electronegativity is a periodic property and
  each element has been assigned a quantitative
  number as shown in table 9.2 on page 304.
• The difference in the electronegativity values
  between any two elements will determine the
  bonding type that will occur between them.

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                                   Some Examples
• Take a look at Hydrogen         • Look at Hydrogen and
  and Chlorine                      Fluorine
• The value for H is 2.1
                                  • H is still 2.1
• The value for Cl is 3.0
                                  • F is 4.0
• The electronegativity
  difference (DEN) is .9            DEN = 4-2.1 or 1.9
• Using figure 9.1 on page          It is also polar, but to a
  303, the bond type is read        greater degree than is
  as polar covalent since .9
                                    the HCl.
  is between 0.5 and 2.0

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                                       For you to try
•   Lithium and Bromine           DEN = 1.8 is classified as
•   Cesium and Fluorine            ionic, but has less ionic
                                   character than K and Cl
•   Fluorine and Fluorine
                                  DEN = 3.3 Ionic
•   Sodium and Chlorine
                                  DEN = 0 Pure Covalent
•   Potassium Chlorine            DEN = 2.1 Ionic
•   Oxygen and Oxygen             DEN = 2.2 Ionic
•   Silicon and Carbon            DEN = 0 Pure Covalent
•   Phosphorus and                DEN = .7 Polar
    Bromine                       DEN = .7 Polar
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                         Periodic Trends in

   Electronegativity Increases

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               Variation in % Ionic Character

Character to
the bond

               0         0.5    1.0      1.5    2.0   2.5     3.0
                   Electronegativity difference between atoms
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                               From the Graph

• Once an electronegativity difference of 1.7 is
  reached, the bond has better than 50% ionic
• The bond might be labeled correctly as polar
  covalent, but for the most part the rules you
  learned about metals and nonmetals forming ionic
  compounds and two nonmetals forming covalent
  compounds remains a good rule of thumb for
  deciding which bonding type is present.
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                                 Polar Bonding

• We’ve looked at the extremes which are a
  bond with a DEN so large that it is
  completely ionic or with a DEN that is zero
  where the bond is completely covalent.
• Now we will look closer at the unequal
  sharing that occurs in polar bonds and what
  it means.

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                              Unequal Sharing
• Hydrogen and Chlorine are 2 nonmetals and by
  our earlier rule they will form a covalent bond.
• Their DEN is .9 which puts them in the polar
  covalent range.
• Since Cl is more electronegative than H, it will
  have a stronger attraction for the electrons and the
  Cl end of the molecule will have a somewhat
  negative charge while the H end will have a
  somewhat positive charge.
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                  Hydrogen Chloride

                                         The electron from
                                         H is pulled toward
            Cl                           the Cl.

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            Water is a polar molecule

• Water is H2O, but let’s look at each H-O
  bond separately.
• The DEN for H-O is 1.4, which makes it
• The Oxygen is most electronegative and the
  electrons will be pulled in that direction
• Let’s look at a model like the one for HCl

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                 The Hydrogen – Oxygen bond

d-                                          The electrons are
                 O     H        d+          pulled harder by the
                                            oxygen than the
                                            hydrogen atom.

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           A look at the Lewis structure of

                                                  Oxygen has 6
                                            O     electrons and
                                                  Hydrogen has 1.
                 O   H        d+
                                                  H        H
             H                          H
                                              O H
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          Water’s Polarity and it’s shape

    • Water is bent because of it’s shared and
      unshared pairs of electrons.
                                                   These two pair
                                                   of electrons are
These two pair       H                             shared.
of electrons are          O H

  The unshared pair repel (like charges repel) and the
  molecule is “bent”. The angle is 105 o.
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      Water Polarity and it’s shape 2

• We’ll talk more about water in a later
  chapter, but if the 4 pair of electrons
  surrounding water in the H2O molecule
  were all equal the bond angle would be that
  of a tetrahedron (109.5o)
• The nonbonding electrons require more
  room and “squeeze” the bonding pair closer
  and thus reduce the bond angle to 105 o.
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       Coffee Filter Chromatography

• The mini lab on page 312 will help to
  illustrate some of water’s unique properties.
• It will also give you some insight into polar
  and non-polar substances and how they
  might be expected to behave.
• Follow the directions closely and answer
  the analysis questions.

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    Lewis Structures and Molecular
• The following steps may be used to determine the correct
  Lewis Structure.
• 1.) Decide which atoms are bonded
• 2.) Count ALL valence electrons
• 3.) Place 2 electrons in each bond.
• 4.) Complete the octets of the atoms attached to the central
  atom by adding e- in pairs.
• 5.) Place any remaining electrons on the central atom in
• 6.) If the central atom does not have an octet, form double
  bonds. If necessary, form triple bonds.
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            Examples of Lewis Structures
                  and Molecular Shape
• SO3 - and ClO4 – Draw Lewis Structures
• First SO3 – S 1 atom x 6 electrons = 6
              O 3 atoms x 6 electrons = 18
  Total = 24 e-
                          There are no electrons left
                          and S does not have 8
     O S O                electrons so one pair from
                          one of the oxygen atoms
                          needs to be pulled in.
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                Sulfur trioxide’s shape
The molecule will have one double bond
 between one Oxygen and the Sulfur
 and single bonds joining the other two
 Oxygen atoms to the sulfur.
                               The shape of the
     O S        O              molecule will be
                               somewhat triangular.
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            Websites with 3D Models

• Virtual Chemistry site
• If you use this site at home, you’ll need to
  get a Chime Plugin The MDL Chime Site
• You will be building gumdrop models and
  will need to consider the spacing and
  direction some of the parts of the molecules
  are directed.

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            Formal Charge and Lewis
• [Formal Charge] = [Number of e- in valence
  shell of the isolated atom] – [(Number of
  bonds to the atom) + (Number of unshared
• Let’s look at the structures that are possible
  for Sulfuric acid. H2SO4.
• There is a structure that obeys the octet rule,
  but it is not the structure that is found by
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