1. 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. Chemical Bonds Chemical Bonds Octet rule
The attractive forces that hold atoms or ions together to form molecules or crystals
Octet rule
atoms tend to gain, lose, or share valence electrons to get an octet.
Everything wants to be like a noble gas.
exceptions
near He obey duet rule
Transition metals
n = 3 and above

4. A general comparison of metals and nonmetals
Figure 9.1
A general comparison of metals and nonmetals

5. Types of Chemical Bonding
Typically:
1. Metal with nonmetal:
electron transfer leads to ionic bonding
2. Nonmetal with nonmetal:
electron sharing leads to covalent bonding
3. Metal with metal:
electron pooling leads to metallic bonding

6. The three models of chemical bonding
Figure 9.2
The three models of chemical bonding

7. Lewis Electron-Dot Symbols
- A method for depicting valence electrons and interactions of atoms
For main group elements -
The A group number gives the number of valence electrons.
Place one dot per valence electron around the four sides of the element symbol. Do not pair dots until all four sides have an electron.
Example:
Nitrogen, N, is in Group 5A and therefore has 5 valence electrons. N : . . N : : N . : N .

8. Lewis electron-dot symbols for elements in Periods 2 and 3
Figure 9.3
Lewis electron-dot symbols for elements in Periods 2 and 3
Nonmetals - The number of unpaired dots indicates the number of electrons it gains, or the number of covalent bonds it usually forms.
Metals – The total number of dots is the maximum number of electrons it may lose when forming a cation.

9. Sec 9.2 Ionic Bonding
In ionic bonding, electrons are gained or lost, the resulting bonds are based on electrostatic attraction.
ex Na has 1 valence e, Cl has 7
If Na could only get rid of 1, if Cl could only gain 1….
When sodium metal is placed in Cl2 gas, they react by transferring 1 e- from Na to Cl to form Na+ and Cl-. Now each has an octet. Because both now have a charge they are attracted to each other to form NaCl.

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

12. Energy in Ionic Bonding
Li(g)  Li+(g) + e- IE1 = 520 kJ
F(g) + e-  F-(g) EA = -328 kJ
So the process would appear to be endothermic
Li(g) + F(g)  Li+(g) + F-(g) E = 192kJ
Overall the process is very exothermic, this is because of the lattice energy.
the enthalpy change of gaseous ions coalescing into a crystalline solid.
Indicates the strength of the two ions attraction
Influences melting point, hardness, and solubility
Ionic solids exist only because the lattice energy drives the unfavorable electron transfer.

13. Calculating lattice energy
Lattice energy cannot be directly measured, so it is found by using Hess’s Law.
The enthalpy change for an overall reaction is the sum of the enthalpy changes of the reactions which make it up.
Lattice energies are calculated by using a Born-Haber Cycle
A series of chosen steps from elements to ionic compounds for which all the enthalpies are known.
The steps are hypothetical and not the actual steps of the process

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

15. 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
a DH0lattice
electrostatic energy a
cation radius + anion radius

16. Trends in lattice energy
Effect of ion size.
Increasing the size of the ions decreases lattice energy, therefore attraction between cations and anions decreases down in a group
Effect of ionic charge.
Increasing the charge of the ions increases the lattice energy.

17. Trends in lattice energy
Figure 9.7
Trends in lattice energy

18. Properties of ionic compounds
Ionic compounds are hard, rigid, and brittle
This is a result of ions being held in specific positions in a crystal. So a crystal retains it’s shape until enough energy is applied to shift positions and crack the crystal.

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

20. Properties of ionic compounds
Do not conduct electricity in the solid state
Ions in fixed positions
Do conduct when melted or dissolved
Ions can move independently

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

22. Properties of ionic compounds
High melting and boiling points (all solid at RT)
Enough energy must be supplied to free ions from the attractions of the surrounding ions
Ionic compounds vaporize as ion-pairs even though no “molecules” exist in the crystal

23. Table 9.1 Melting and Boiling Points of Some Ionic Compounds
mp (0C)
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
MgO
2852
3600

24. Vaporizing an ionic compound.
Figure 9.10
Vaporizing an ionic compound.

25. Sec 9.3 Covalent Bonding
Elements can also form octets by sharing e- between them, the bonds that result are called covalent bonds.
usually occurs in nonmetal/nonmetal compounds.
More compounds are covalent than ionic.
A single Cl atom has 7 valence electrons, in a sample of pure Cl one atom cannot steal an electron from another, so they share to form Cl2.
Molecule – a compound formed by 2 or more atoms joined by covalent bonds that behaves as a single particle.

26. Covalent Bonding
atoms share electrons by overlapping orbitals so electrons can exist in the orbitals of both atoms at once.
These shared or bonding pairs of electrons are represented by lines in structures.
Other valence electrons that are not involved in bonding are called unshared or lone pairs.
Every pair of electrons shared between atoms is a bond
1 pair – single bond
2 pairs – double bond , stronger
3 pairs – triple bond, strongest
aka bond order
Lewis structures of Cl2, O2, N2

27. Covalent bond formation in H2.
Figure 9.11
Covalent bond formation in H2.

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

29. Bond Energy
Bond energy or Bond Strength - the energy required to overcome the attraction of covalently bonded atoms.
It is defined as energy required to break bonds in 1 mole of gaseous atoms.
Bond energy depends on the specific elements involved.
It can vary from molecule to molecule so table values are averages.

32. Bond length and covalent radius.
Figure 9.13
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

33. Bond length, bond energy, and bond order are closely related
Higher bond order is shorter, and stronger for a given set of atoms
With a constant bond order, longer bonds are usually weaker.

35. SAMPLE PROBLEM 9.2
Comparing Bond Length and Bond Strength
PROBLEM:
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
PLAN:
(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 O
Bond strength: S - F > S - Cl > S - Br
Bond strength: C O > C = O > C - O

36. Properties of covalent cmpds
The physical properties of molecular compounds, are not related to the strength of their covalent bonds.
Most covalent compounds have low m.p. and b.p. because the strong covalent bonding is typically isolated within molecules. The attractions between separate molecules, called intermolecular forces, are what must be overcome to melt or boil these covalent substances.
The physical properties of network covalent solids, are related to the strength of their covalent bonds.
In these substances there are no individual molecules, the covalent bonding extends in 3-D throughout the substance.
Ex Quartz (SiO2) very hard, mp 1550°C
Diamond (C) hardest known substance, mp 3550°C

37. Strong forces within molecules and weak forces between them.
Strong covalent bonding forces within molecules
Figure 9.14
Weak intermolecular forces between molecules

38. Covalent bonds of network covalent solids.
Figure 9.15
Covalent bonds of network covalent solids.

39. Properties of covalent cmpds
Most covalent substance are poor electrical conductors, when solid, liquid, or dissolved.

40. Sec. 9.4 Between the extremes
Most real bonds fall somewhere between the ideal of ionic or covalent bonding theory.
The type of bond atoms form depends on electronegativity
some atoms attract e- more strongly than others, we say these are more electronegative
electronegativity increases going right and up the table
Covalent bonds in which e- are not shared equally because of electronegativity differences are called polar covalent bonds

41. The Pauling electronegativity (EN) scale.
Figure 9.16
The Pauling electronegativity (EN) scale.

42. Electronegativity and atomic size.
Figure 9.17
Electronegativity and atomic size.

43. Representation of Polar Bonds

44. 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(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

45. Nonpolar, Polar, or Ionic
In general if the electronegativity difference between two bonded atoms is:
0, usually between identical nonmetal atoms, called nonpolar covalent
< .4 , mostly covalent
.4 to 1.7, 2 different nonmetals called polar covalent
> 1.7 , usually nonmetals and reactive metals, is mostly ionic
Note: there is no perfect ionic bond.

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

47. Percent ionic character of electronegativity difference (DEN).
Figure 9.19
Percent ionic character of electronegativity difference (DEN).

48. Figure 9.20 Li F

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

50. Solid Metals and metal alloys have metallic bonding Electron Sea Model
Sec. 9.5 Metallic Bonding
Solid Metals and metal alloys have metallic bonding
Electron Sea Model
All the metal atoms contribute their valence electrons to a delocalized pool of electrons. The metal cations are held together by attraction to the delocalized electrons.

51. Properties of Metals
Most are solid at RT, with moderate to high mp, and very high b.p.
m.p. are not very high because the metallic bonds don’t have to be broken to become liquid
b.p. are very high because the cation and it’s electrons must be separated from the others
m.p. are higher for metals with more valence electrons
cation charges are higher resulting in greater cation-electron sea attractions

52. 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

53. Melting points of the Group 1A(1) and Group 2A(2) elements.
Figure 9.23
Melting points of the Group 1A(1) and Group 2A(2) elements.

54. Properties of Metals
Metals are good conductors of electricity when solid, or liquid.
The delocalized electrons are able to move under an electric field
Metals are good conductors of heat.
The delocalized electrons disperse heat more quickly
Metals are malleable and ductile, not brittle
The cations are able to slide past each other and still retain their attraction to the electron sea.

55. The reason metals deform.
Figure 9.24
The reason metals deform.
metal is deformed

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
Figure B9.1
Infrared Spectroscopy
Some vibrational modes in general diatomic and triatomic molecules.

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

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