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.