1. Chapter 9
Models of Chemical Bonding

2. A general comparison of metals and nonmetals.
Figure 9.1
A general comparison of metals and nonmetals.
Astatine may be a metalloid but it is radioactive.

3. BASIC CHEMICAL BONDING
Atoms form bonds in order to complete their outer shells of electrons
Main group elements try for 8. (Transition metals are more complex.)
Atoms can give up or gain electrons to form ions which then form ionic bonds, or they can share electrons to form covalent bonds
(Or in metals, share electrons throughout all atoms)

4. BASIC CHEMICAL BONDING
Atoms/ions seek lower energy states:
- ions in a bond have a lower energy than the separated ions
Data indicate lowest energy level is for atoms to share electrons at a certain distance = bond length = covalent radii

5. Types of Chemical Bonding
1. Metal with nonmetal:
electron transfer and ionic bonding
2. Nonmetal with nonmetal:
electron sharing and covalent bonding
3. Metal with metal:
electron pooling and metallic bonding
4. Cation with anion: ionic bonding (ammonium ion and nitrate ion)

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

7. “NEW” CONCEPT OF CHEMICAL BONDING: LEWIS THEORY
Valence electrons have fundamental role in chemical bonding
If e-(s) "transferred" ionic bond results
If e-(s) "shared" covalent bond results
If e-(s) “shared as a pool” metallic bond results
Trying to achieve noble gas configuration
Lewis symbol: Atomic symbol with valence electrons for neutral atom, correct number of electrons for ion

8. Silberberg Chp Two
Figure 2.14 from 4th ed.
The relationship between ions formed and the nearest noble gas. (from Chp 2)

9. Lewis Electron-Dot Symbols
For main group elements -
The group number digit 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 15 and therefore has 5 valence electrons. N : . . N : : N . : N .

10. Lewis electron-dot symbols for elements in Periods 2 and 3.
Figure 9.4
Lewis electron-dot symbols for elements in Periods 2 and 3.
How will these Lewis symbols look when these atoms form ions?

11. Lewis Symbols for Respresentative Elements

12. Lewis Structures For Ions:
Put individual ion structures together as group or use coefficient:
CaF2 = [:F:]-[Ca]2+[:F:]- or [Ca]2+2[:F:]-
Note that Ca:  [Ca] e-
each F + e- --> [F]-

13. SAMPLE PROBLEM 9.1
Depicting Ion Formation
PROBLEM:
Use Lewis symbols to depict the formation of Na+ and O2- ions from the atoms, and determine the formula of the compound. Draw in the most correct way when I show you below. Ions should have brackets.
SOLUTION: : Na
+ O .
2 [Na]++[ O ]2- :

14. Figure 9.5
Three ways to represent the formation of Li+ and F- through electron transfer.
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: focus on this way . + F : Li
[Li]+ +
F - :
[They should both have brackets.]

15. Lewis Structures For Ions:
Practice drawing Lewis structure for ions and then use them to make ionic compounds of the two elements listed. Be prepared to show your work on the document camera for the class to see. Draw large structures in ink and use brackets where required!
Mg and O, Al and Cl, Na and S

16. LATTICE ENERGY
When ionic bonding occurs, energy is released – exothermic – because the positive and negative ions achieve their lowest energy state when surrounded by the opposite charge. However, IE is endothermic and usually larger than EA, so the transfer of electrons is endothermic. In chp 6 we will learn about Hess’ Law, which is how we actually determine the value of lattice energy.

17. LATTICE ENERGY
For now: the energy released when forming an ionic compound arranged in a crystal lattice from the elements is called enthalpy of formation. If we know the value of IE and EA, then we can determine the energy released when the ions are moved into the crystal lattice.
Lattice energy is affected by:
Ion size: the larger the radius, the lower the lattice energy
Ionic charge(s): the greater the charge, the higher the lattice energy

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

19. Ionic compound dissolved in water
Figure 9.9
Electrical conductance and ion mobility.
Solid ionic compound
Molten ionic compound
Ionic compound dissolved in water

20. 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

21. Covalent Bonding A shared pair of e-s between two atoms
Other e-s in valence shell become Lone Pairs
Can have single, double or triple bonds (Bond Order)

22. Covalent Bonding Energy is released when atoms join to form bonds
Energy must be absorbed to break bonds - called bond dissociation energy
Measured for gaseous species in kJ/mol
Increases with multiple bonding
Increases with decreased bond length
See Bond Energy and Bond Length tables

23. Formation of a covalent bond between two H atoms. (from Chp 2)
Silberberg Chp Two
Formation of a covalent bond between two H atoms. (from Chp 2)
Figure 2.13
Covalent bonds form when elements share electrons, which usually occurs between nonmetals.

24. BOND ENERGIES & DHrxn:
DHrxn approx same as difference between bond breaking energy and bond formation energy
REMEMBER THIS IS AN APPROXIMATION BECAUSE WE ARE USING AVERAGES!
Estimate DHrxn for N2 + 3 H2  2 NH3
First find kinds of bonds there are
Look up average bond energies in table

26. Draw Lewis structure, look up bond energies, count number of bonds, DHrxn ~ BEprod – BEreact

29. 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: (NO CALCS, JUST THINK!)
(a) S - F, S - Br, S - Cl
(b) C = O, C - O, C O
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

30. Section 9.4
We can’t really calculate bond energies until we can draw Lewis electron dot structures, which is in chp 10, and learn about enthalpy of reaction, which is in chp 6. We will return to this after chp 6.

31. Strong forces within molecules and weak forces between them.
Strong covalent bonding forces within molecules
Figure 9.13
Weak intermolecular forces between molecules

32. Covalent bonds of network covalent solids.
Figure 9.14

33. CONCEPT OF ELECTRONEGATIVITY: from Linus Pauling
I define it as greediness for electrons in a covalent bond
Look at table – a scale with no units
Metals are low and nonmetals are high
Difference in their electronegativities:
If DEN is < .5, bond is covalent (0 to 0.4)
If > 2.0, it's ionic,
and in between is polar covalent (0.5 to 1.9)
Determine something called "% ionic character" using the table in text

34. The Pauling electronegativity (EN) scale.
Figure 9.19
The Pauling electronegativity (EN) scale.

35. SAMPLE PROBLEM 9.4
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.
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

36. Figure 9.21
A. Boundary ranges for classifying ionic character of chemical bonds. B. Gradation in ionic character.
3.0
2.0
0.0
DEN
> 2.0 Ionic
0.5 to 1.9 Polar covalent
0 to 0.4 Nonpolar covalent
Use my simplified version!

37. Percent ionic character of electronegativity difference (DEN).
Figure 9.21 continued
Percent ionic character of electronegativity difference (DEN).

38. Properties of the Period 3 chlorides.
Figure 9.22
Properties of the Period 3 chlorides.

39. Metallic Bonding (not in text but YOU have to know it)
Sea of electrons shared throughout the metal piece, however large it is
Explains both electrical and thermal conductivity, since electrons can move throughout the piece
A mixture of metals, an alloy, also share electrons throughout
Metallic bond strength varies, as you can see from range of melting points, but most metals are solids (only Hg is liquid at room temperature)

40. Table 9.7 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

41. Melting points of the Group 1A(1) and Group 2A(2) elements.
Figure 9.26
Melting points of the Group 1A(1) and Group 2A(2) elements.

42. The reason metals deform.
Figure 9.27
The reason metals deform.
metal is deformed