1. Chapter 9
Models of Chemical Bonding

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

3. A general comparison of metals and non-metals
Figure 9.1
A general comparison of metals and non-metals

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)

5. The three models of chemical bonding
Figure 9.2

6. Lewis Electron-Dot Symbols
For Main Group elements:
The 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 . : N .

7. Lewis electron-dot symbols for elements in Periods 2 and 3
Figure 9.3

8. General Rules
For a metal, the total number of dots equals the maximum number
of electrons it loses to form a cation.
For a non-metal, the number of unpaired dots equals the number
of electrons that become paired either through electron gain or
electron sharing. The number of unpaired dots equals either the
negative charge of the anion an atom forms or the number of
covalent bonds it forms.

9. The Ionic Bonding Model
Involves the transfer of electrons from metal to non-metal to form
ions that come together in a solid ionic compound
The Octet Rule
When 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 atoms
equals the total number of electrons gained by the non-metal atoms.

10. SAMPLE PROBLEM 9.1
Depicting Ion Formation
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.
PLAN:
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.
SOLUTION: 2s 2p
O2- 3s 3p Na 2s 2p O
2 Na+ 3s 3p Na : Na
+ O .
2Na+ + O 2- :

11. Figure 9.4
Three ways to represent the formation of Li+ and F- through electron transfer
1. Electron configurations
Li 1s22s1 +
F 1s22s22p5
Li+ 1s2 +
F- 1s22s22p6
2. Orbital diagrams
Li+ 1s 2s 2p Li 1s 2s 2p F 1s 2s 2p + F- 1s 2s 2p +
3. Lewis electron-dot symbols . + F : Li
Li+ +
F - :

12. Ionic Bonding and Lattice Energy
The electron transfer process is an endothermic process, but ionic compound
formation 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 = kJ/mol!
Li+(g) + F-(g) LiF(g) ∆Ho = kJ
(an exothermic process due to the attraction of
oppositely charged ions)

13. Even more energy is released when the gaseous ions coalesce
into a crystalline solid. Thus….
Li+(g) + F-(g) LiF(s) ∆Holattice of LiF = lattice energy = kJ
The lattice energy is the enthalpy change that occurs when gaseous
ions coalesce into an ionic solid.
How do we measure lattice energy experimentally? Use
Hess’s law in a Born-Haber cycle

14. The Born-Haber cycle for lithium fluoride
Figure 9.6

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 enthalpy
change of the overall process minus the sum of the above four
steps = kJ

16. Central Point
Ionic solids exist only because the lattice energy drives the
energetically unfavorable electron transfer.

17. Periodic Trends in Lattice Energy
Coulomb’s Law
charge A x charge B
electrostatic force a
distance2
But energy = force x distance. Therefore,
charge A x charge B
electrostatic energy a
distance
cation charge x anion charge
a DHolattice
electrostatic energy a
cation radius + anion radius

18. Trends in lattice energy
Figure 9.7

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

20. Does the ionic model explain the properties of ionic compounds?

21. Electrostatic forces and the reason ionic compounds crack
Figure 9.8

22. Electrical Conductance and Ion Mobility
Molten ionic compound
Solid ionic compound
Ionic compound dissolved in water
Electrical Conductance and Ion Mobility
Figure 9.9

23. Table 9.1 Melting and Boiling Points of Some Ionic Compounds
mp (oC)
bp (oC)
CsBr
636
1300
NaI
661
1304
MgCl2
714
1412
KBr
734
1435
CaCl2
782
>1600
NaCl
801
1413
LiF
845
1676 KF
858
1505
MgO
2852
3600

24. Vaporizing an ionic compound
Figure 9.10

25. The Covalent Bonding Model
Each atom in a covalent bond “counts” the shared electrons
as belonging entirely to itself.
An electron pair that is part of an atom’s valence shell but not
involved in bonding is called a lone pair, or unshared pair.
Bond order: the number of electron pairs being shared between
any 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

26. Covalent bond formation in H2
Figure 9.11

27. The attractive and repulsive forces in covalent bonding
Figure 9.12

28. Properties of Covalent Bonds
Bond energy (bond enthalpy or bond strength): the energy required
to overcome the mutual attraction between the bonded nuclei and
the shared electrons.
Bond breakage is an endothermic process; bond energy is always
positive.
Bond formation is an exothermic process.

29. Bond Length
For a given pair of atoms, a higher bond order results in a
shorter bond length and a higher bond energy.
A shorter bond is a stronger bond.

32. Bond length and covalent radius
internuclear distance
(bond length)
covalent radius
internuclear distance
(bond length)
covalent radius
internuclear distance
(bond length)
covalent radius
internuclear distance
(bond length)
covalent radius
Figure 9.13

34. SAMPLE PROBLEM 9.2
Comparing Bond Length and Bond Strength
PROBLEM:
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
PLAN:
(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.
SOLUTION:
(a) Atomic size increases moving down a group.
(b) Using bond orders we get:
Bond length: S - Br > S - Cl > S - F
Bond length: C - O > C = O > C O
Bond strength: S - F > S - Cl > S - Br
Bond strength: C O > C = O > C - O

35. Properties of Covalent Compounds
Weak forces between molecules, not the strong covalent
bonds within each molecule, are responsible for the
physical properties of covalent compounds.
Covalent compounds have relatively low melting and boiling
points.
Most covalent compounds are poor electrical conductors.

36. Strong forces within molecules, weak forces between them
Strong covalent bonding forces within molecules
Figure 9.14
Weak intermolecular forces between molecules

37. Network Covalent Solids
No separate molecules; held together by covalent bonds that
extend throughout the sample
quartz: melts at 1550 oC.
diamond: melts at 3550 oC.
These examples illustrate the strength of covalent bonds.

38. Covalent bonds of network covalent solids
Figure 9.15

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 unequal
sharing of electrons, with F bearing a partial negative
charge and H bearing a partial positive charge. The
attraction between the partial charges strengthens the
bond.

40. The Pauling electronegativity (EN) scale
Figure 9.16

41. Trends in Electronegativity
In general, electronegativity is inversely related to
atomic 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 of
the Periodic Table).

42. Electronegativity and atomic size
Figure 9.17

43. Electronegativity and Oxidation Number
(a) The more electronegative atom in a bond is assigned
all of the shared electrons; the less electronegative atom is
assigned none of the shared electrons.
(b) Each atom in a bond is assigned all of its unshared
electrons.
(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 valence
electrons; has an O.N. of = -1
H has 1 valence electron; has an O.N. of = +1

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 ( ) that
points toward the negative pole
H2 and F2: examples of nonpolar covalent bonds

45. 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 to find EN values; the arrow should point toward the negative end.
(b) EN 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

46. Partial Ionic Character of Polar Covalent Bonds
Related directly to the electronegativity difference (∆EN) between
the bonded atoms
The greater the ∆EN, the larger the partial charges and the higher
the partial ionic character (PIC).
Thus LiF has more PIC than HF; HF has more PIC than F2.

47. Boundary ranges for classifying the ionic character of chemical bonds
3.0
DEN
2.0
Boundary ranges for classifying the ionic character of chemical bonds
0.0
Figure 9.18

48. Percent ionic character as a function of electronegativity difference (DEN)
Figure 9.19

49. Li F Charge density of the LiF molecule (an ionic compound)
No bond has 100%
ionic character; electron
sharing occurs to some extent
Figure 9.20

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 becomes
more covalent.

51. Properties of the Period 3 chlorides
Figure 9.21
Properties of the Period 3 chlorides

52. Metallic Bonding The electron-sea model: all metal atoms in the sample
contribute their valence electrons to form an “electron
sea” that is delocalized throughout the substance
The metal atoms are not held in place as rigidly as are
the ions of an ionic solid.

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

54. Melting points of the Group 1A and Group 2A elements
Figure 9.23

55. The reason metals deform
metal is deformed
Figure 9.24

56. Infrared Spectroscopy
Tools of the Laboratory
Infrared Spectroscopy
Figure B9.1
Some vibrational modes in general diatomic and triatomic molecules

57. Tools of the Laboratory
Infrared Spectroscopy
Some vibrational modes in general diatomic and triatomic molecules
Figure B9.1

58. Some vibrational modes in general diatomic and triatomic molecules.
Tools of the Laboratory
Infrared Spectroscopy
Some vibrational modes in general diatomic and triatomic molecules.
Figure B9.1

59. Tools of the Laboratory
Figure B9.2
The infrared (IR) spectrum of acrylonitrile