Transcript Chapter 9 PowerPoint

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Chapter 9
Models of Chemical Bonding
9-1
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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 Bond Energy and Chemical Change
9.5 Between the Extremes:
Electronegativity and Bond Polarity
9-2
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Figure 9.1
9-3
A general comparison of metals and nonmetals.
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Types of Chemical Bonding
1. Metal with nonmetal:
electron transfer and ionic bonding
2. Nonmetal with nonmetal:
electron sharing and covalent bonding
3. Metal with metal:
electron pooling and metallic bonding
9-4
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Figure 9.2
9-5
The three models of chemical bonding.
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Lewis Electron-Dot Symbols
For main group elements The A group number gives the number of valence electrons.
Place one dot per valence electron on each of the four sides
of the element symbol.
Pair the dots (electrons) until all of the valence electrons are used.
Example:
Nitrogen, N, is in Group 5A and therefore has 5 valence electrons.
. N. .
.
. N:
.
.
. N.
:
9-6
:
.
: N .
.
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Figure 9.4
Lewis electron-dot symbols for elements in Periods 2 and 3.
9-7
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SAMPLE PROBLEM 9.1
PROBLEM:
PLAN:
Depicting Ion Formation
Use partial orbital diagrams and Lewis symbols to depict the
formation of Na+ and O2- ions from the atoms, and determine
the formula of the compound.
Draw orbital diagrams for the atoms and then move electrons to
make filled outer levels. It can be seen that 2 sodiums are
needed for each oxygen.
SOLUTION:
O2-
Na
2s
O
2 Na+
.
Na
3s
3p
+ : O:
.
Na
9-8
:
2p
2Na+ +: O:2-
:
2s
Na
2p
.
3p
.
3s
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Figure 9.5
Three ways to represent the formation of Li+ and Fthrough electron transfer.
Electron configurations
Li 1s22s1
F 1s22s22p5
+
Li+ 1s2
+
F- 1s22s22p6
2s
2p
Orbital diagrams
Li+
Li
1s
2s
1s
2p
+
+ F
1s
2s
F1s
2p
2s
Lewis electron-dot symbols
:
9-9
Li+
+
: F: -
:
+
:
Li .
.
:F:
2p
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Periodic Trends in Lattice Energy
Coulomb’s Law
charge A X charge B
electrostatic force a
distance2
energy = force X distance
therefore
charge A X charge B
electrostatic energy a
distance
cation charge X anion charge
electrostatic energy a
9-10
cation radius + anion radius
a DH0lattice
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Figure 9.7
9-11
Trends in lattice energy.
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Figure 9.8
Electrostatic forces
9-12
and the reason ionic
compounds crack.
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Figure 9.9
Solid ionic
compound
9-13
Electrical conductance and ion mobility.
Molten ionic
compound
Ionic compound
dissolved in water
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Table 9.1 Melting and Boiling Points of Some Ionic Compounds
Compound
bp (0C)
CsBr
636
1300
NaI
661
1304
MgCl2
714
1412
KBr
734
1435
CaCl2
782
>1600
NaCl
801
1413
LiF
845
1676
KF
858
1505
2852
3600
MgO
9-14
mp (0C)
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Figure 9.10
9-15
Covalent bond formation in H2.
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Figure 9.11
9-16
Distribution of electron density of H2.
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Table 9.2
back to previous slide
9-17
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Figure 9.12
Internuclear distance
(bond length)
Bond length and covalent radius.
Covalent
radius
Internuclear distance
(bond length)
72 pm
Internuclear distance
(bond length)
Covalent
radius
100 pm
9-18
Covalent
radius
114 pm
Internuclear distance
(bond length)
Covalent
radius
133 pm
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Table 9.3
9-19
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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.
(b) Using bond orders we get
Bond length: S - Br > S - Cl > S - F
Bond length: C - O > C = O > C
Bond strength: S - F > S - Cl > S - Br
Bond strength: C
9-20
O
O>C=O>C-O
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Figure 9.13
Strong forces within molecules and weak forces between them.
Strong covalent bonding forces within molecules
Weak intermolecular forces between molecules
9-21
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Figure 9.14
9-22
Covalent bonds of network covalent solids.
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Figure 9.15
9-23
The infrared (IR) spectra of diethyl ether and 2-butanol.
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Figure 9.16
Using bond energies to calculate DH0rxn.
Enthalpy, H
DH0rxn = DH0reactant bonds broken + DH0product bonds formed
BOND BREAKING
DH01 = + sum of BE
DH02 = - sum of BE
BOND FORMATION
DH0rxn
9-24
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Figure 9.17
Using bond energies to calculate DH0rxn of methane.
BOND BREAKING
Enthalpy,H
4BE(C-H)= +1652kJ
2BE(O2)= + 996kJ
DH0(bond breaking) = +2648kJ
2[-BE(C O)]= -1598kJ
4[-BE(O-H)]= -1868kJ
DH0(bond forming) = -3466kJ
DH0rxn= -818kJ
9-25
BOND FORMATION
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SAMPLE PROBLEM 9.3
PROBLEM:
Calculating Enthalpy Changes from Bond
Energies
Use Table 9.2 (button at right) to calculate DH0rxn for the
following reaction:
CH4(g) + 3Cl2(g)
PLAN:
CHCl3(g) + 3HCl(g)
Write the Lewis structures for all reactants and products and
calculate the number of bonds broken and formed.
SOLUTION:
H
Cl
H C H
+
3
Cl
H C Cl
+
3 H
Cl
H
bonds broken
9-26
Cl
bonds formed
Cl
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SAMPLE PROBLEM 9.3
Calculating Enthalpy Changes from Bond
Energies
continued
bonds broken
4 C-H
bonds formed
= 4 mol(413 kJ/mol) = 1652 kJ
3 C-Cl = 3 mol(-339 kJ/mol) = -1017 kJ
3 Cl-Cl = 3 mol(243 kJ/mol) = 729 kJ
1 C-H = 1 mol(-413 kJ/mol) = -413 kJ
DH0bonds broken = 2381 kJ
3 H-Cl = 3 mol(-427 kJ/mol) = -1281 kJ
DH0bonds formed = -2711 kJ
DH0reaction = DH0bonds broken + DH0bonds formed = 2381 kJ + (-2711 kJ) = - 330 kJ
9-27
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Table 9.4
9-28
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Figure 9.20
9-29
The Pauling electronegativity (EN) scale.
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SAMPLE PROBLEM 9.4
PROBLEM:
Determining Bond Polarity from EN Values
(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 9.19(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
9-30
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Figure 9.21
9-31
Electron density distributions in H2, F2, and HF.
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Figure 9.22
Boundary ranges for classifying ionic character
of chemical bonds.
3.0
DEN
2.0
0.0
9-32
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Figure 9.23
9-33
Properties of the Period 3 chlorides.