Chapter 9 (Silberberg 3rd Edition) Models of Chemical Bonding

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Chapter 9 (Silberberg 3rd Edition)Models of
Chemical Bonding

• 9.1 Atomic Properties and Chemical Bonds
• 9.2 The Ionic Bonding Model
• 9.3 The Covalent Bonding Model
• 9.4 Between the Extremes Electronegativity
  and Bond Polarity
• 9.5 An Introduction to Metallic Bonding

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Types of Chemical Bonding

• Whats a Chemical Bond?
• Attraction that holds atoms or ions together in
  compounds
• Ionic Bonding vs Covalent Bonding
• Whats the difference?
• Kinds of atoms involved?
• Metallic Bonding
• Kinds of atoms involved?

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Ionic Bond

1. Electrostatic force of attraction between
   oppositely charged ions
2. Ions result from the transfer of one or more
   electrons from a metal to a nonmetal (Trans of
   NaCl)
3. Why do metals lose electrons to form cations?
4. Why do nonmetals gain electrons to form anions?

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Figure 9.1
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Conditions Needed for Ionic Bond Formation

• Chemical Bonding occurs only if it results in a
  decrease in PE
• i.e. The process is exothermic
• Cation formation is Endothermic (PE
  increases)....Why?
• Relate to Ionization Energy
• Anion formation is Exothermic (PE
  decreases)......Why?
• Relate to Electron Affinity

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Conditions Needed for Ionic Bond Formation

• Cation formation is usually more endothermic than
  Anion formation is exothermic
• Why then is Ionic Bond formation EXOTHERMIC?

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Must Consider Lattice Energy

• Lattice Energy
• PE lowering due to the attraction of anions to
  cations
• Highly Exothermic
• Ionic bonding will only result when......
• Lattice Energy is more exothermic than
  E. A. I.E. is endothermic
• E.g Li (s) ½ F2 (g) ? LiF (s)

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Figure 9.6
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Figure 9.7
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Factors that affect Lattice Energy

• Lattice energy
• Depends on the charge, size and distance between
  the ions involved Why??
• Due to the electrostatic attractions between
  cations and anions
• Electrostatic attractions depends on
• Charge and size of ionsWhy?
• Distance between ionsWhy?

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Periodic Trends in Lattice Energy

• Down a group
• Down group IA
• Down group IIA
• Down group IIIA
• Across a period
• Across period 2

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Electron Configurations of Ions

• Octet Rule
• Atoms of many elements tend to gain, lose, or
  share electrons until their valence shell
  contains 8 electrons

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Rules for Writing Electron Configurations of
Ions...

• Group IA , IIA Metals and Aluminum
• Lose electrons until reach Noble gas
  configuration
• Nonmetals
• Gain electrons until reach Noble gas
  configuration
• Write the electron configurations for the ions
  in......
• KCl, CaCl2, AlCl3, CaO, Na2 O, Al2O3

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Rules for Writing Electron Configurations of
Ions...

• Transition and Post-transition Metals
• Do NOT obey the Octet Rule!!
• More than one ion is often possible
• Transition Metals
• Lose s-Sublevel electrons, then d-electrons
• e.g. Fe 2, Fe 3 , Zn 2 , Cu1 , Cu2 ,
• Post Transition Metals
• Lose p-sublevel electrons, then s-electrons
• e.g. Sn 2 , Sn 4 , Pb 2 , Pb 4

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Lewis Symbols

• Symbol of element surrounded by valence electrons
• Used to represent bond formation
• Write Lewis Symbols for....
• Representative Elements, Groups IA - VIIA
• Note Group Number number of valence electrons

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Using Lewis Symbols to Illustrate Ionic Bond
Formation

• Use Lewis Symbols to diagram the reaction that
  produces the following compounds.....
• KCl, CaCl2, AlCl3, CaO, Na2 O, Al2O3
• ZnCl2

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Explaining the Properties of Ionic compounds

• Ionic compounds
• Have high melting points and boiling points
• (all are solids at room temp.)
• Hard, but brittle solids
• Conduct electricity in as liquids, but not as
  solids

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Covalent Bonding

• Involve the sharing of one or more PAIRS of
  electrons between atoms of nonmetallic elements
• Occurs when ionic bond formation is not favored
  energetically
• i.e. when .... I.E. E.A. is more endothermic
  than the lattice energy is exothermic

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Bond formation between two Hydrogen Atoms
H
H
H
H
H2

1. Atoms approach each other

1. Covalent bond formation

1. Large distance between atoms

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Bond Length

• Determined by a balance between the
  following......
• Attractions of shared electrons to both nuclei
• Causes a decrease in PE
• Repulsion between both nuclei
• Causes an increase in PE

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Figure 9.12
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Figure 9.11
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Figure 9.13
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Bond Energy

1. Amount of energy released during bond formation
2. Amount of energy needed to break a bond

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Fig. 9.15 Network Covalent solids have very high
melting points
In Diamond each C atom is covalently bonded to 4
other C atoms.
In Quartz each Si atom is covalently bonded to 4
O atom. Each O atom is bonded to 2 Si atoms
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Illustrating Covalent Bonding with Lewis
Structures

• Apply the Octet Rule
• Atoms tend to share electrons until their valence
  shell contains 8 electrons
• Use Lewis Structures to illustrate bond formation
  for.....
• H2, F2, H2O, NH3, CH4
• Multiple Bonds
• N2, SiO2 , NO3-

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Guidelines for writing Lewis Structures

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 electrons in pairs
5. Place any remaining electrons on the central atom
   in pairs
6. If the central atom does not have an octet, form
   double bonds, or if necessary, a triple bond.

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Nonpolar vs Polar Covalent Bonding

• Nonpolar Covalent Bond
• Involves equal sharing of an electron pair
  between two nuclei
• Pure nonpolar bonds are quite uncommon....Why??
• Polar Covalent Bond
• Unequal sharing of electrons
• Results from the electronegativity difference
  between atoms of different elements

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Figure 9.16
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Figure 9.17
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Electronegativity Differences and Bond Types

1. Pure Nonpolar Covalent 0
2. More Nonpolar than Polar lt 0.5
3. Polar Covalent 0.5 to 1.7
4. More Ionic than Polar Covalent gt 1.7

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Some Examples

• Indicate the kind of bonding in.....
• Water
• Ammonia
• Carbon dioxide
• Aluminum Chloride
• Methane
• Fatty Acids

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Polar Bonds vs Polar Molecules

1. Why are water molecules polar, whereas carbon
   dioxide molecules are nonpolar?

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Figure 9.21
Properties of the Period 3 chlorides.
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Explaining the Properties of Metals

• Have high melting points
• (all but Hg are solids at room temp.)
• Malleable (deform when a force is applied)
• Conduct electricity

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Explaining the Properties of Metals
Figure 9.24
Why metals deform Metal atoms slide past each
other when a force is applied Why do metals
conduct electricity?
Figure 9.24
The reason metals deform.
metal is deformed
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Table 9.5 Melting and Boiling Points of Some
Metals
Element
mp(0C)
bp(0C)
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Melting points of the Group 1A(1) and Group 2A(2)
elements. Figure 9.23
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Tools of the Laboratory Infrared Spectroscopy
Figure B9.1
Some vibrational modes in a diatomic molecule
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Tools of the Laboratory Infrared Spectroscopy
Figure B9.1
Some vibrational modes in a triatomic molecule
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Tools of the Laboratory Infrared Spectroscopy
Figure B9.1
The infrared (IR) spectrum of acrylonitrile.