Chemical Bonding I:
Basic Concepts
Lecture 3 of x, Chapter 9
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Activity 1
The Concept of Resonance
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Resonance Structures
❖
Resonance is the use of two
or more Lewis structures to
represent a particular
molecule.
❖
Example: Draw Lewis
structure of Ozone (O3)
❖
Always use ↔ to indicate
structures are resonance
structures.
❖
Are the oxygen-to-oxygen
bond lengths the same in O3?
❖
Even though, the bonds
lengths are truly the same,
the best way to describe a
molecule like O3 is to draw
two or more resonance
structures.
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Resonance Structures
❖
Draw the structure (s) for the polyatomic carbonate ion.
❖
Do you remember the chemical formula? Don’t forget
about its overall charge?
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Resonance Structures
❖
What are the resonance
structures for benzene, a
well known organic
molecule?
❖
Hexagonal structure of
benzene was 1st proposed by
August Kekule.
❖
Is there a simply way to
describe the structure of
benzene?
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Practice Exercise 9.8
❖
Draw three resonance structures for the thiocyanate ion,
SCN-. Rank the structures in decreasing order of
importance.
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Activity 2
Exceptions to the Octet Rule
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3 Exceptions to Octet Rule
Incomplete
Octet
❖
Odd number of
electrons
Expanded
Octet
The octet rule applies mainly to the second-period elements
(C, N, O, F).
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Incomplete Octet
❖
Elements that tend to form compounds in which they are
surrounded by fewer than 8 electrons have an incomplete
octet.
❖
Examples: Beryllium hydride (BeH2) and boron
trifluoride (BF3)
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Example Problem
❖
Draw the Lewis structure for aluminum triiodide (AlI3)
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Odd-Electron Molecules
❖
Some molecules have an odd number of electrons, so octet
cannot be satisfied for all atoms.
❖
Example: Nitrogen dioxide (NO2)
❖
Odd electron molecules are called radicals, which are highly
reactive molecules because their unpaired electron wants to
form a covalent bond with another molecule.
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Expanded Octet
❖
Elements in and beyond the 3rd period can have more than 8
electrons around them, which is an expanded octet.
❖
Example: Sulfur hexafluoride (SF6)
❖
Sulfur’s 6 valence electrons form a covalent bonds with each
fluorine, so there are 12 electrons around the central sulfur atom.
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Example Problem
❖
Draw the Lewis structure for phosphorus pentafluoride
(PF5), in which all five F atoms are bonded to the central P
atom.
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Expanded Octet
❖
Sometimes we find that the octet rule is satisfied for all
atoms but there are still valence electrons left to place.
❖
Example: Draw a Lewis structure of the noble gas
compound xenon tetrafluoride (XeF4) in which all F atoms
are bonded to the central Xe atom.
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Activity 3
Bond Enthalpy
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Bond Enthalpy
❖
A measure of the stability of a molecule is its bond enthalpy,
which is the enthalpy (or heat) change required to break a
particular bond in 1 mole of gaseous molecules.
❖
H2(g) → H(g) + H(g)
❖
Cl2(g) → Cl(g) + Cl(g) ΔHo = 242.7 kJ/mol
❖
HCl(g) → H(g) + Cl(g) ΔHo = 431.9 kJ/mol
❖
O2(g) → O(g) + O(g)
ΔHo = 498.7 kJ/mol
❖
N2(g) → N(g) + N(g)
ΔHo = 941.4 kJ/mol
ΔHo = 436.4 kJ/mol
Note bond enthalpy is
greatest for triple bonds,
then double bonds, thus
showing stability of
triple and double
bonded molecules.
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Enthalpy of Reaction
❖
Imagine a reaction proceeding by breaking all bonds in the
reactants and then using the gaseous atoms to form all the
bonds in the products.
❖
The enthalpy of reaction (ΔHo) can be calculated as follows:
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Example Problem
❖
Use Equation 9.3 to
calculate the enthalpy of
reaction for the process
❖
H2(g) + Cl2(g) → 2HCl(g)
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Example Problem
❖
Estimate the enthalpy
change for the combustion
of hydrogen gas:
❖
2H2(g) + O2(g) → 2H2O(g)
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