Honors Chemistry ch 9

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Honors Chemistry
Chapter 9: Chemical Bonding I
9.1 Lewis Dot Diagrams
• Symbol surrounded by dots for valence e• Separate dots as much as possible
9.2 Ionic Bonding
• Electrostatic force holding ions together
• Ions formed by electron transfer
• Low Eion loses e-, high e.a. gains e-
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Na + Cl  Na+ ClMg + S  Mg2+ S2Li + S  2 Li+ S2-
9.3 Lattice Energy
• Energy that holds ionic compounds
together in a crystal lattice
• Transfer of e- requires energy (Eion) and
releases energy (e.a.)
• In general, the cation requires more
energy than the anion releases, which
makes bond formation unstable
• Lattice energy releases additional energy,
making bond formation stable
9.4 The Covalent Bond
• Covalent bond = shared pair of electrons
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F + F  F–F
• Shared pair – shared electrons, bond
• Lone pair – electrons not involved in bond
9.4 Lewis Structures
• Representation of covalent compounds
using dots for e- and lines for bonds
• Octet rule
• atoms bond in such a way as to gain 8 e- in
valence shell
• Exceptions – H and He
• Multiple Bonds
• Double bond – share 2 pairs; eg, O2
• Triple bond – share 3 pairs; eg, N2
9.4 Bonding Summary
9.4 Ionic / Covalent Properties
• Intermolecular attractive forces
• Ionic – strong, covalent – weak
• Consider phase, density, solubility,
conductivity
• Ionic
Covalent
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Solid
High density
Usually soluble
Good conductor
Liquid or gas
Low density
Often insoluble
Poor conductor
9.5 Electronegativity
• Element’s relative attraction for shared e-
9.5 Pauling Electronegativities
9.5 Electronegativity and Atomic
Radius
9.5 Bond Character
• Degree of sharing of the bonded e• Depends on the difference in electronegativity
• Small electronegativity difference
• Equal sharing of bonded e• True covalent bond (nonpolar covalent)
• Moderate electronegativity difference
• Unequal sharing of bonded e• Polar covalent bond
• Large electronegativity difference
• Ionic bond
9.5 Bond Character
3.0
DEN
2.0
0.0
9.6 Writing Lewis Structures
• Draw a reasonable skeletal structure for
the compound
• Count the total valence electrons available
• Draw single bonds between all atoms and
use remaining electrons to fulfill the octet
rule
• If there are not enough electrons to fulfill
the octet rule, form double or triple bonds
• Draw Lewis structures for NH3, O3, CO32-
9.7 Formal Charges
• Difference between the number of e- an
atom has in a Lewis structure and the
number of e- it has as a free element
• Assigning formal charges
• Atom counts all its nonbonding e-’s
• Atom counts 1 e- from each bond
• Total formal charge must add up to the
total charge of the molecule / ion
9.7 Formal Charges
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The most plausible structures have:
The fewest formal charges
Formal charges of smallest magnitude
Negative formal charges on the most
electronegative elements
• No adjacent charges of the same type
9.8 Resonance
• Some molecules have more than one
plausible Lewis structure
O
O
O
O
O
O
• Resonance –use of both structures to
represent a molecule
• Reality is that bonds are delocalized
O
O
O
9.8 Resonance
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Draw resonance structures for
N2O
HSCN
NO3CO32Extreme resonance – C6H6
9.9 Exceptions to the Octet Rule
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Incomplete Octets
Not enough electrons to make an octet
Usually occurs with Groups IIA and IIIA
Examples: BeH2, BF3
Consider resonance in BF3
BF3 + NH3  F3BNH3
Coordinate covalent bond
• One atom donates both shared electrons
9.9 Exceptions to the Octet Rule
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Odd Electron Molecules
Often called radicals
Examples: NO, NO2
Single e- goes on element with lower EN
Very reactive
Tend to form dimers
9.9 Exceptions to the Octet Rule
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Expanded Octets
More than 8 e- on central atom
Requires the atom to have a d orbital
Can’t happen with periods 1 and 2
Examples: SF6, XeF4, ClO4-

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