chapter 9 (silberberg 3 rd edition) models of chemical bonding 9.1 atomic properties and chemical...
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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
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 Bond1. Electrostatic force of attraction between
oppositely charged ions2. 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?
Figure 9.1
Conditions Needed for Ionic Bond Formation
1. Chemical Bonding occurs only if it results in a decrease in PE» i.e. The process is exothermic
2. Cation formation is Endothermic (PE increases)....Why?» Relate to Ionization Energy
3. 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» PE lowering due to the attraction of
anions to cations» Highly Exothermic
2. Ionic bonding will only result when......» Lattice Energy is more exothermic than
E. A. + I.E. is endothermicE.g Li (s) + ½ F2 (g) LiF (s)
Figure 9.6
9-10
Figure 9.6 The Born-Haber cycle for lithium fluoride
Figure 9.7
Factors that affect Lattice Energy
1. Lattice energya. 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?
9-11
Periodic Trends in Lattice Energy
Coulomb’s Lawcharge A X charge B
electrostatic force distance2
But: Energy = Force x Distance, therefore
charge A X charge Belectrostatic energy
distance
cation charge X anion chargeelectrostatic energy
cation radius + anion radius
H0lattice
Periodic Trends in Lattice Energy
1. Down a groupa. Down group IAb. Down group IIAc. Down group IIIA
2. Across a perioda. Across period 2
Electron Configurations of Ions
1. Octet RuleAtoms 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» Lose electrons until reach Noble gas configuration
2. Nonmetals» Gain electrons until reach Noble gas configuration
3. 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
» Do NOT obey the Octet Rule!!» More than one ion is often possible
2. Transition Metals» Lose s-Sublevel electrons, then d-electrons
e.g. Fe 2+, Fe 3+ , Zn 2+ , Cu1+ , Cu2+ ,
3. 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 -
VIIANote: 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 solidsc. Conduct electricity in as liquids, but
not as solids
Covalent Bonding
1. Involve the sharing of one or more PAIRS of electrons between atoms of nonmetallic elements
2. 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
a) Large distance between atoms
b. Atoms approach each other
c. Covalent bond formation
H H H H H2
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. Amount of energy released during bond formation
2. Amount of energy needed to break a bond
9-23
SAMPLE PROBLEM 9.2 Comparing Bond Length and Bond Strength
PROBLEM:
PLAN:
SOLUTION:
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 (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.
(a) Atomic size increases going down a group.
Bond length: S - Br > S - Cl > S - F
Bond strength: S - F > S - Cl > S - Br
(b) Using bond orders we get
Bond length: C - O > C = O > C O
Bond strength: C O > C = O > C - O
9-24
Figure 9.14Strong covalent bonding forces within molecules
Weak intermolecular forces between molecules
Strong forces within molecules & weak forces between them.
9-25
Figure 9.15 Covalent bonds of network covalent solids
In Quartz: each Si atom is covalently bonded to 4 O atom. Each O atom is bonded to 2 Si atoms
In Diamond: each C atom is covalently bonded to 4 other C atoms.
Fig. 9.15
Network Covalent solids have very high melting points
Illustrating Covalent Bonding with Lewis Structures
1. Apply the Octet Rule» Atoms tend to share electrons until their
valence shell contains 8 electrons
2. Use Lewis Structures to illustrate bond formation for.....
» H2, F2, H2O, NH3, CH4
3. Multiple Bonds» N2, SiO2 , NO3
-
Guidelines for writing Lewis Structures
1. Decide which atoms are bonded2. Count all valence electrons3. Place 2 electrons in each bond4. Complete the octets of the atoms attached to
the central atom by adding electrons in pairs5. Place any remaining electrons on the central
atom in pairs6. 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– Pure nonpolar bonds are quite uncommon....Why??
2. 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. Pure Nonpolar Covalent: 02. More Nonpolar than Polar: < 0.53. Polar Covalent: ~ 0.5 to 1.7 4. More Ionic than Polar Covalent: > 1.7
9-28
SAMPLE PROBLEM 9.3 Determining Bond Polarity from EN Values
PROBLEM:
PLAN:
SOLUTION:
(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.
(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.
(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
Some Examples
1. Indicate the kind of bonding in.....a) Water
b) Ammonia
c) Carbon dioxide
d) Aluminum Chloride
e) Methane
f) 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
Figure 9.24
metal is deformedThe reason metals deform.
Why metals deform: Metal atoms slide past each other when a force is applied
Why do metals conduct electricity?
Figure 9.24
Explaining the
Properties of Metals
Table 9.5 Melting and Boiling Points of Some Metals
Element 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 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: Infrared Spectroscopy
Figure B9.1
Some vibrational modes in a triatomic molecule
Tools of the Laboratory: Infrared Spectroscopy
Figure B9.1
The infrared (IR) spectrum of acrylonitrile.