Chapter 9 Chemical Bonding
Bonding of Atoms
• 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
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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
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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• 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
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
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Variation in % Ionic Character
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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• 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
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• 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
• 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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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
• 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
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A look at the Lewis structure of
Oxygen has 6
O electrons and
Hydrogen has 1.
O H d+
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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
• 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
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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
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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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