Chemical Bonding, cont’d
Covalent bonding
Chemical bond formed by sharing a pair of eEach atom gains an octet of e- by sharing electron
pairs, e.g.
H +H
F + F
H H
F F
These are known as Lewis structures for H2 and F2;
Show the shared e- pair as a line between atoms, i.e.,
as a single bond
In the Lewis model, valence is the number of e- needed to
complete an octet: valence = 8 – group number
What are the valences of C, N, O? What are the formulas
of their hydrides?
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Multiple bonds
Atoms may share more than 1 e- pair to form double
or triple bonds
e.g., how can each N in N2 gain an octet?
Triple bond makes N2 very unreactive - also, triple
bond is shorter than single N-N bond
e.g., draw the Lewis structure of O2
Bond Polarity and Electronegativity
The two cases of bonding we've studied thus far...
covalent
ionic
F-F
Na+ Cl-
e- are shared
e- transferred
Is there anything in between these extremes?
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Bond polarity
Nonpolar bond: e- shared equally
e.g., O2, N2, etc
Polar covalent: one atom exerts
greater attraction for e- than other(s)
Two nonmetals with different IE
and EA
Ionic: formed if difference in ability to
attract e- is large enough (e.g., metal
and nonmetal)
How to estimate whether a bond is nonpolar, polar
covalent, or ionic?
Electronegativity (EN)
Ability of an atom in a molecule to attract e- to
itself
Compare to EA….. what’s the difference?
Electronegativity is related to ionization energy &
electron affinity; how?
165
Periodic trends in electronegativity
Where are the most/least electronegative atoms?
Electronegativity & bond polarity
Use the difference in electronegativity between atoms to
gauge the polarity of a bond
The greater the difference in electronegativity EN the
more polar the bond, e.g.,
compound
F2
HF
LiF
 EN
0
1.9
3.0
bond type
cov
pol cov
ionic
HF: F more electronegative than H; F attracts the e- from H
but it is not transferred completely as in LiF
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e.g., Based solely on position in the periodic table, which
bond is more polar, C-N, B-F, or N-O?
The charge distribution of a molecule is determined by
both the shape of the molecule and the polarity of the
bonds within the molecule
A molecule is polar if its centers of positive and negative
charge do not coincide
e.g., HCl, O2
Any diatomic molecule with a polar bond is a polar
molecule
if two nonmetals differ in EN, what happens to the
distribution of electrons in the molecule?
One end of a polar molecule has a slight positive charge
and the other has a slight (and equal) negative charge
this separation of positive and negative charge creates a
dipole
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Quantitatively, the degree of polarity of a molecule is
measured by its dipole moment 
If a distance r separates two equal and opposite charges
q+ and q-, then the magnitude of the dipole moment is the
product of q and r:
 = Qr
Dipole moments have units of charge x distance
Charge Q is measured in units of electronic charge, e
e = 1.60 x 10-19 C
distances are generally measured in Å (1Å =1 x 10-10 m)
E.g., suppose that two charges 1+ and 1- are separated by
a distance of 1.00 Å: calculate 
Note that dipole moments are generally reported in units of
debyes (D): 1 D = 3.34 x 10-30 C-m
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e.g., the H-Cl bond length in HCl is 1.27 Å.
Calculate the dipole moment (in D) that would result if the
charges on H and Cl were +1 and -1, respectively
Experimentally, the dipole moment of HCl(g) is 1.08 D.
Calculate the charges on H and Cl.
Important points: if a molecule has a nonzero dipole
moment, it is polar
If its dipole moment is zero it is nonpolar
The larger the dipole moment, the more polar the
molecule
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Dipole moment trends in the hydrogen halides
Compound
, D
r(Å)
HF
1.82
0.92
HCl
1.08
1.27
HBr
0.82
1.41
HI
0.44
1.61
Calculate the charges on H and the halogens for the series
above...
how to explain this trend? how do these data correlate
with properties such as electronegativity?
170
Differentiating Ionic and Covalent Bonding
Notice: molecular compounds (covalent bonding)
tend to exist as molecules
Typically have low melting/boiling points and are
nonelectrolytes
Ionic compounds (ionic bonding) tend to be high-melting
solids with extended lattice structures
Is a compound composed of a metal and a nonmetal
ALWAYS ionic?
Not always –
SnCl4 is molecular – it freezes at -33oC and boils
at 114oC
Mn2O7 is a liquid that freezes at 5.9oC; MnO is a
solid that melts at 1842oC
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How to tell when a compound containing a metal (or
metalloid) and a nonmetal Is covalent or ionic?
General principle: when the oxidation number of the metal
becomes highly positive (+4 or larger), then we expect a
significant degree of covalent character when bonding
with a nonmetal
E.g., In the following pairs of binary compounds,
determine which is molecular and which is ionic, and
name them using the appropriate naming convention
CaF2 and TiCl4
SbCl5 and AlF3
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Drawing Lewis structures of polyatomics
e.g., draw the Lewis structures of CO2, PCl3 & ClO3-Sum the valence e- from all atoms (use group number)
-Write the element symbols - show how they are
connected
For binary compounds: central atom written first
-Draw a single bond between each bonded atom pair
Each single bond uses two electrons
-Complete the octets of atoms bonded to central atom;
place leftover e- on central atom
-Make multiple bonds if necessary to complete octet on
central atom
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Formal Charge and Lewis Structures
Concept of formal charge is used as an aid in deciding
between alternative Lewis structures
e.g., how many acceptable Lewis structures can
be drawn for CO2?
Formal charge is the charge an atom would have
if all atoms in a molecule had the same
electronegativity
To calculate the formal charge on any atom in a
Lewis structure, we assign electrons as follows:
formal charge =
group # - [1/2(bonding e-) + nonbonding e-)]
The most stable Lewis structure will be the that in which:
All atoms have formal charges closest to zero
Any negative formal charges reside on the more
electronegative atoms
e.g., based on the concept of formal charge, which Lewis
structure above is more likely for CO2?
How does formal charge differ from oxidation number?
174
Resonance structures
Draw the Lewis structure for
NO3-
Experimentally, all N-O bond lengths are the same
Bond length: single > double > triple
Equivalent Lewis structures: resonance forms
'real' molecule is described by an ‘average’ of the
resonance forms
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Exceptions to the octet rule
Odd # of e- , e.g. NO, ClO2 : cannot acheive an octet about
each atom
Less than an octet: usually in compounds of Be or B; e.g.
draw the Lewis structure of BF3
More than an octet: observed only for elements in period 3
and beyond, e.g., PCl5
Why only period 3 and beyond? Can NH5 exist?
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Covalent Bond Strength
The bond dissociation energy D is defined as the energy
required to break 1 mol of a specific bond in a gaseous
substance, e.g.,
Cl2(g)  2Cl(g)
D(Cl-Cl) =242 kJ/mol
Bond dissociation energies are always positive, i.e.,
energy is always required to break a bond
(a negative energy means that energy is released during
the process)
What about polyatomic species, e.g., NH3? What is the
average bond energy of 1 N-H bond?
Here, we measure the amount of energy required to break
all three N-H bonds, and divide by 3 to get an ‘average’ NH bond energy
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Bond energies can be tabulated & used to determine
energy changes for reactions where bonds are
made/broken, i.e.,
e.g., given the following bond energies,
Bond energy, kJ/mol
H-Cl
431
F-F
155
H-F
567
Cl-Cl
242
determine the energy change for the reaction
2HCl(g) + F2(g)  2HF(g) + Cl2(g)
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Problems du Jour
Write Lewis structures for the following:
H2CO
AsO33-
ICl2-
XeF4
For each of the following, write a single Lewis structure that
obeys the octet rule and calculate the oxidation numbers and
formal charges on all the atoms:
SO2
SO3
SO32-
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