Chapter 9 (Silberberg 3 Edition)
Models of Chemical Bonding
rd
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
Types of Chemical Bonding
1.
What’s a Chemical Bond?
• Attraction that holds atoms or ions together in
compounds
2.
Ionic Bonding vs Covalent Bonding
• What’s the difference?
• Kinds of atoms involved?
3.
Metallic Bonding
• Kinds of atoms involved?
Ionic Bond
1.
2.
Electrostatic force of attraction between
oppositely charged ions
Ions result from the transfer of one or more
electrons from a metal to a nonmetal (Trans of
NaCl)
3.
4.
Why do metals lose electrons to form cations?
Why do nonmetals gain electrons to form
anions?
Figure 9.1
Conditions Needed for Ionic Bond
Formation
1.
2.
3.
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
Conditions Needed
for Ionic Bond Formation
1.
Cation formation is usually more
endothermic than Anion formation is
exothermic
» Why then is Ionic Bond formation
EXOTHERMIC?
Must Consider Lattice Energy
1.
Lattice Energy
»
»
2.
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)
Figure 9.6
Figure 9.6
9-10
The Born-Haber cycle for lithium fluoride
Figure 9.7
Factors that affect Lattice Energy
1.
Lattice energy
a. Depends on the charge, size and distance
between the ions involved— Why??
b. Due to the electrostatic attractions between
cations and anions
 Electrostatic attractions depends on…
•
•
Charge and size of ions—Why?
Distance between ions—Why?
Periodic Trends in Lattice Energy
Coulomb’s Law
charge A X charge B
electrostatic force

But: Energy = Force x Distance,
electrostatic energy 
distance2
therefore
charge A X charge B
distance
cation charge X anion charge
electrostatic energy 
cation radius + anion radius
9-11
H0lattice
Periodic Trends in Lattice Energy
1.
Down a group
a. Down group IA
b. Down group IIA
c. Down group IIIA
2.
Across a period
a. Across period 2
Electron Configurations of Ions
1.
Octet Rule
Atoms of many elements tend to gain, lose,
or share electrons until their valence shell
contains 8 electrons
Rules for Writing
Electron Configurations of Ions...
1.
Group IA , IIA Metals and Aluminum
»
2.
Nonmetals
»
3.
Lose electrons until reach Noble gas configuration
Gain electrons until reach Noble gas configuration
Write the electron configurations for the ions
in......
»
KCl, CaCl2, AlCl3, CaO, Na2 O, Al2O3
Rules for Writing
Electron Configurations of Ions...
1.
Transition and Post-transition Metals
»
»
2.
Transition Metals
»
3.
Do NOT obey the Octet Rule!!
More than one ion is often possible
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+
Lewis Symbols
1.
Symbol of element surrounded by
valence electrons
» Used to represent bond formation
2.
Write Lewis Symbols for....
» Representative Elements, Groups IA - VIIA
Note: Group Number = number of valence
electrons
Using Lewis Symbols to
Illustrate Ionic Bond Formation
1.
Use Lewis Symbols to diagram the reaction
that produces the following compounds.....
» KCl, CaCl2, AlCl3, CaO, Na2 O, Al2O3
» ZnCl2
Explaining the
Properties of Ionic compounds
1.
Ionic compounds
a. Have high melting points and boiling points
(all are solids at room temp.)
b. Hard, but brittle solids
c. Conduct electricity in as liquids, but not as
solids
Covalent Bonding
1.
2.
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
Bond formation between two
Hydrogen Atoms
H
a) Large distance
between atoms
H
H
H
b. Atoms approach
each other
H2
c. Covalent
bond
formation
Bond Length
1.
Determined by a balance between the
following......
a) Attractions of shared electrons to both nuclei
– Causes a decrease in PE
b) Repulsion between both nuclei
– Causes an increase in PE
Figure 9.12
Figure 9.11
Figure 9.13
Bond Energy
1.
2.
Amount of energy released during bond
formation
Amount of energy needed to break a bond
SAMPLE PROBLEM 9.2
PROBLEM:
Comparing Bond Length and Bond Strength
Using the periodic table, but not Tables 9.2 and 9.3, rank
the bonds in each set in order of decreasing bond length
and bond strength:
(a) S - F, S - Br, S - Cl
PLAN:
(b) C = O, C - O, C
O
(a) The bond order is one for all and sulfur is bonded to
halogens; bond length should increase and bond strength
should decrease with increasing atomic radius. (b) The same
two atoms are bonded but the bond order changes; bond
length decreases as bond order increases while bond
strength increases as bond order increases.
SOLUTION:
(a) Atomic size increases going down a group.
Bond length: S - Br > S - Cl > S - F
Bond strength: S - F > S - Cl > S - Br
9-23
(b) Using bond orders we get
Bond length: C - O > C = O > C
Bond strength: C
O
O>C=O>C-O
Strong forces within molecules & weak forces between them.
Figure 9.14
Strong covalent bonding forces within molecules
Weak intermolecular forces between molecules
9-24
Fig. 9.15
Figure 9.15
Covalent bonds of network covalent solids
Network Covalent solids have very high melting points
In Quartz: each Si atom is
covalently bonded to 4 O atom.
Each O atom is bonded to 2 Si
atoms
9-25
In Diamond: each C atom is
covalently bonded to 4
other C atoms.
Illustrating Covalent Bonding
with Lewis Structures
1.
Apply the Octet Rule
»
2.
Use Lewis Structures to illustrate bond
formation for.....
»
3.
Atoms tend to share electrons until their valence
shell contains 8 electrons
H2, F2, H2O, NH3, CH4
Multiple Bonds
»
N2, SiO2 , NO3-
Guidelines for
writing Lewis Structures
1.
2.
3.
4.
5.
6.
Decide which atoms are bonded
Count all valence electrons
Place 2 electrons in each bond
Complete the octets of the atoms attached to the
central atom by adding electrons in pairs
Place any remaining electrons on the central atom in
pairs
If the central atom does not have an octet, form
double bonds, or if necessary, a triple bond.
Nonpolar vs Polar Covalent Bonding
1.
Nonpolar Covalent Bond
»
Involves equal sharing of an electron pair between
two nuclei
–
2.
Pure nonpolar bonds are quite uncommon....Why??
Polar Covalent Bond
»
Unequal sharing of electrons
– Results from the electronegativity difference
between atoms of different elements
Figure 9.16
Figure 9.17
Electronegativity Differences
and Bond Types
1.
2.
3.
4.
Pure Nonpolar Covalent: 0
More Nonpolar than Polar: < 0.5
Polar Covalent: ~ 0.5 to 1.7
More Ionic than Polar Covalent: > 1.7
SAMPLE PROBLEM 9.3
Determining Bond Polarity from EN Values
PROBLEM: (a) Use a polar arrow to indicate the polarity of each bond:
N-H, F-N, I-Cl.
(b) Rank the following bonds in order of increasing polarity:
H-N, H-O, H-C.
PLAN:
(a) Use Figure 9.16(button at right) to find EN values; the
arrow should point toward the negative end.
(b) Polarity increases across a period.
SOLUTION: (a) The EN of N = 3.0, H = 2.1; F = 4.0; I = 2.5, Cl = 3.0
N-H
F-N
I - Cl
(b) The order of increasing EN is C < N < O; all have an EN
larger than that of H.
H-C < H-N < H-O
9-28
Some Examples
1.
Indicate the kind of bonding in.....
a)
b)
c)
d)
e)
f)
Water
Ammonia
Carbon dioxide
Aluminum Chloride
Methane
Fatty Acids
Polar Bonds vs Polar Molecules
1.
Why are water molecules polar,
whereas carbon dioxide molecules are
nonpolar?
Figure 9.21
Properties of the Period 3 chlorides.
Explaining the Properties of Metals
a. Have high melting points
(all but Hg are solids at room temp.)
b. Malleable (deform when a force is
applied)
c. Conduct electricity
Explaining the
Properties of Metals
Figure 9.24
Why
deform: Metal atoms slide past each
Figuremetals
9.24
other when a force is applied
The reason metals deform.
Why do metals conduct electricity?
metal is deformed
Table 9.5 Melting and Boiling Points of Some Metals
mp(0C)
bp(0C)
Lithium (Li)
180
1347
Tin (Sn)
232
2623
Aluminum (Al)
660
2467
Barium (Ba)
727
1850
Silver (Ag)
961
2155
Copper (Cu)
1083
2570
Uranium (U)
1130
Element
3930
Melting points of the Group 1A(1) and Group 2A(2) elements.
Figure 9.23
Tools of the Laboratory:
Infrared Spectroscopy
Figure B9.1
Some vibrational modes in a diatomic molecule
Tools of the Laboratory:
Figure B9.1
Some
vibrational
modes in a
triatomic
molecule
Infrared Spectroscopy
Tools of the Laboratory:
Infrared Spectroscopy
Figure B9.1
The infrared (IR) spectrum of acrylonitrile.