9-1

Chapter 9

Models of Chemical Bonding

9-2

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

9-3

A general comparison of metals and non-metals

Figure 9.1

9-4

Types of Chemical Bonding

1. Metal with non-metal

electron transfer and ionic bonding

2. Non-metal with non-metal

electron sharing and covalent bonding (localized)

3. Metal with metal

electron pooling and metallic bonding (delocalized)

9-5

Figure 9.2

The three models of chemical bonding

9-6

Lewis Electron-Dot Symbols

For Main Group elements:

Example:

Nitrogen (N) is in Group 5A and therefore has 5 valence electrons.

N:.

..

:

N .. ..N :.

. :N ...

The group number gives the number of valence electrons.

Place one dot per valence electron on each of the four sidesof the element symbol.

Pair the dots (electrons) until all of the valence electrons are used.

9-7

Figure 9.3

Lewis electron-dot symbols for elements in Periods 2 and 3

9-8

General Rules

For a metal, the total number of dots equals the maximum numberof electrons it loses to form a cation.

For a non-metal, the number of unpaired dots equals the numberof electrons that become paired either through electron gain orelectron sharing. The number of unpaired dots equals either thenegative charge of the anion an atom forms or the number ofcovalent bonds it forms.

9-9

The Ionic Bonding Model

Involves the transfer of electrons from metal to non-metal to formions that come together in a solid ionic compound

The Octet RuleWhen atoms bond, they lose, gain or share electrons to attain a filled

outer shell of eight (or two) electrons

In ionic bonding, the total number of electrons lost by the metal atomsequals the total number of electrons gained by the non-metal atoms.

9-10

SAMPLE PROBLEM 9.1 Depicting Ion Formation

PLAN:

SOLUTION:

PROBLEM: 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 two sodiums are needed for each oxygen.

3s 3p

Na

3s 3p

Na2s 2p

O2s 2p

O2-

2 Na+

:Na

Na+ O

.:

..

.2Na+ + O 2-

:: : :

9-11

1. Electron configurations

Li 1s22s1

2. Orbital diagrams

3. Lewis electron-dot symbols

+ F 1s22s22p5 Li+ 1s2 + F- 1s22s22p6

Three ways to represent the formation of Li+ and F- through electron transfer

Figure 9.4

Li

1s 2s 2p

F

1s 2s 2p

+

Li+

1s 2s 2p

F-

1s 2s 2p+

.+ F: ::Li . Li+ + F -::

::

9-12

Ionic Bonding and Lattice Energy

The electron transfer process is an endothermic process, but ionic compoundformation is an exothermic process.

Li(g) Li+ + e- IE1 = 520 kJ

F(g) + e- F-(g) EA = -328 kJ

Li(g) + F(g) Li+(g) + F-(g) IE1 + EA = 192 kJ

But ∆Hfo for solid LiF = -617 kJ/mol!

Li+(g) + F-(g) LiF(g) ∆Ho = -755 kJ

(an exothermic process due to the attraction ofoppositely charged ions)

9-13

Even more energy is released when the gaseous ions coalesceinto a crystalline solid. Thus….

Li+(g) + F-(g) LiF(s) ∆Holattice of LiF = lattice energy = -1050 kJ

The lattice energy is the enthalpy change that occurs when gaseousions coalesce into an ionic solid.

How do we measure lattice energy experimentally? UseHess’s law in a Born-Haber cycle

9-14

Figure 9.6 The Born-Haber cycle for lithium fluoride

9-15

Working the Numbers

STEP 1: Enthalpy of Li atomization = 161 kJ

STEP 2: 1/2 the bond energy of F2(g)= 0.5(159 kJ) = 79.5 kJ

STEP 3: IE1 for Li(g) = 520 kJ

STEP 4: EA of F(g) = -328 kJ

The enthalpy change for the overall process, ∆Hfo, = -617 kJ

Only the lattice energy is unknown, and it is equal to the enthalpychange of the overall process minus the sum of the above foursteps = -1050 kJ

9-16

Central Point

Ionic solids exist only because the lattice energy drives theenergetically unfavorable electron transfer.

9-17

Periodic Trends in Lattice Energy

Coulomb’s Law

charge A x charge B

electrostatic force distance2

But energy = force x distance. Therefore,

charge A x charge B

electrostatic energy distance

cation charge x anion charge

electrostatic energy cation radius + anion radius

Holattice

9-18

Figure 9.7

Trends in lattice energy

9-19

Effect of Ionic Charge on Lattice Energy

Compare LiF and MgO: Li+ and Mg2+ have similar radii, and

F- and O2- have similar radii.

∆Holattice (LiF) = -1050 kJ/mol ∆Ho

lattice (MgO) = -3923 kJ/mol

The nearly four-fold larger value for MgO reflects the difference inthe product of the charges (12 vs 22) in the numerator of the electrostatic energy equation (monovalent vs divalent ions).

9-20

Does the ionic model explain the properties of ionic compounds?

9-21

Figure 9.8

Electrostatic forces and the reason

ionic compounds crack

9-22

Figure 9.9

Electrical Conductance

and Ion Mobility

Solid ionic compound

Molten ionic compound

Ionic compound dissolved in water

9-23

Table 9.1 Melting and Boiling Points of Some Ionic Compounds

Compound mp (oC) bp (oC)

CsBr

661

1300

NaI

MgCl2

KBr

CaCl2NaCl

LiF

KF

MgO

636

714

734

782

801

845

858

2852

1304

1412

1435

>1600

1413

1676

1505

3600

9-24

Figure 9.10

Vaporizing an ionic compound

9-25

The Covalent Bonding Model

Each atom in a covalent bond “counts” the shared electronsas belonging entirely to itself.

An electron pair that is part of an atom’s valence shell but notinvolved in bonding is called a lone pair, or unshared pair.

Bond order: the number of electron pairs being shared betweenany two bonded atoms

single bond (H2) - bond order of 1

double bond (H2C=CH2) - bond order of 2

triple bond (N2) - bond order of 3

9-26

Figure 9.11

Covalent bond formation in H2

9-27

Figure 9.12

The attractive and repulsive forces in covalent bonding

9-28

Properties of Covalent Bonds

Bond energy (bond enthalpy or bond strength): the energy requiredto overcome the mutual attraction between the bonded nuclei andthe shared electrons.

Bond breakage is an endothermic process; bond energy is alwayspositive.

Bond formation is an exothermic process.

9-29

Bond Length

For a given pair of atoms, a higher bond order results in ashorter bond length and a higher bond energy.

A shorter bond is a stronger bond.

9-30

9-31

9-32 Figure 9.13

internuclear distance(bond length)

covalent radius

internuclear distance(bond length)

covalent radius

internuclear distance(bond length)

covalent radius

internuclear distance(bond length)

covalent radius

Bond length and covalent radius

9-33

9-34

SAMPLE PROBLEM 9.2 Comparing Bond Length and Bond Strength

PROBLEM:

PLAN:

SOLUTION:

Using the periodic table, 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) Bond order =1 for all and sulfur is bonded to halogens; bond length should increase and bond strength should decrease with increasing atomic radius. (b) Similar atoms (C) are bonded but bond order changes; bond length decreases as bond order increases, and bond strength increases as bond order increases.

(a) Atomic size increases moving 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

9-35

Properties of Covalent Compounds

Weak forces between molecules, not the strong covalentbonds within each molecule, are responsible for thephysical properties of covalent compounds.

Covalent compounds have relatively low melting and boilingpoints.

Most covalent compounds are poor electrical conductors.

9-36

Figure 9.14 Strong covalent bonding forces within molecules

Weak intermolecular forces between molecules

Strong forces within molecules, weak forces between them

9-37

Network Covalent Solids

No separate molecules; held together by covalent bonds thatextend throughout the sample

quartz: melts at 1550 oC.

diamond: melts at 3550 oC.

These examples illustrate the strength of covalent bonds.

9-38

Figure 9.15

Covalent bonds of network covalent solids

9-39

The Concept of Electronegativity (EN)

Defined as the relative ability of a bonded atom to attract shared electrons (not the same as EA)

Bond energy of H2 = 432 kJ/mol

Bond energy of F2 = 159 kJ/mol

Bond energy of HF = 565 kJ/mol, not 296 kJ/mol

The stronger-than-expected HF bond is due to unequalsharing of electrons, with F bearing a partial negativecharge and H bearing a partial positive charge. The attraction between the partial charges strengthens thebond.

9-40Figure 9.16

The Pauling electronegativity (EN) scale

9-41

Trends in Electronegativity

In general, electronegativity is inversely related toatomic size.

For main-group elements, EN generally increases up a group and across a period.

Non-metals are more electronegative than metals.

The least electronegative (most electropositive) non-radioactive element is Cs (lower left-hand corner ofthe Periodic Table).

9-42Figure 9.17

Electronegativity and atomic size

9-43

Electronegativity and Oxidation Number

(a) The more electronegative atom in a bond is assignedall of the shared electrons; the less electronegative atom isassigned none of the shared electrons.

(b) Each atom in a bond is assigned all of its unsharedelectrons.

(c) The oxidation number is given by:

O.N. = # valence e- - (# shared e- + # unshared e-)

e.g.: HCl: Cl more electronegative than H; has 7 valenceelectrons; has an O.N. of 7 - 8 = -1H has 1 valence electron; has an O.N. of 1 - 0 = +1

9-44

Polar Covalent Bonds and Bond Polarity

Covalent bonds involving atoms with different electronegativities: generate partial (+) and (-) charges; defined as polar covalent bonds (e.g., HCl)

Polar covalent bonds: depicted by a polar arrow ( ) thatpoints toward the negative pole

H2 and F2: examples of nonpolar covalent bonds

9-45

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 to find EN values; the arrow should point toward the negative end.

(b) EN 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

9-46

Partial Ionic Character of Polar Covalent Bonds

Related directly to the electronegativity difference (∆EN) betweenthe bonded atoms

The greater the ∆EN, the larger the partial charges and the higherthe partial ionic character (PIC).

Thus LiF has more PIC than HF; HF has more PIC than F2.

9-47

Figure 9.18

EN

3.0

2.0

0.0

Boundary ranges for classifying the ionic

character of chemical bonds

9-48 Figure 9.19

Percent ionic character as a function of electronegativity difference (EN)

9-49

Figure 9.20

Li F

Charge density ofthe LiF molecule

(an ionic compound)

No bond has 100%ionic character; electron

sharing occurs to some extent

9-50

Ionic-To-Covalent Bonding Continuum Across a Period

Consider bonding between a metal and non-metal in Period 3

NaCl, MgCl2, AlCl3, SiCl4, PCl3, S2Cl2, and Cl2

Increasing covalent character (decreasing ionic character) from NaCl to Cl2

Underlying factor: As ∆EN becomes smaller, the bond becomesmore covalent.

9-51

Figure 9.21

Properties of the Period 3 chlorides

9-52

Metallic Bonding

The electron-sea model: all metal atoms in the samplecontribute their valence electrons to form an “electronsea” that is delocalized throughout the substance

The metal atoms are not held in place as rigidly as are the ions of an ionic solid.

9-53

Table 9.5 Melting and Boiling Points of Some Metals

element mp (oC) bp (oC)

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

9-54

Figure 9.23

Melting points of the Group 1A and Group 2A elements

9-55

Figure 9.24

metal is deformedThe reason metals deform

9-56

Infrared Spectroscopy

Tools of the Laboratory

Figure B9.1

Some vibrational modes in general diatomic and triatomic molecules

9-57

Infrared Spectroscopy

Tools of the Laboratory

Figure B9.1

Some vibrational modes in general

diatomic and triatomic molecules

9-58

Infrared SpectroscopyTools of the Laboratory

Figure B9.1

Some vibrational modes in general diatomic and

triatomic molecules.

9-59

Tools of the Laboratory

Figure B9.2

The infrared (IR) spectrum of acrylonitrile