1. Chapter 6: Chemical Bonding

2. Why do atoms bond?
To achieve the most favorable (lowest energy) arrangement

3. Octet Rule
Atoms will give, take, or share electrons in order to have 8 valence electrons
Exception: Hydrogen wants only 2 & Helium has only two

4. Valence Electrons Ar Ar Na Mg Al N O Cl
Electrons in the outermost energy level (s, p)
Ex. Na – 1s22s22p63s1 – 1 valence e-
Ex. Cl – 1s22s2p63s23p5 – 7 valence e-
Shown in dot diagrams (1 dot per valence e-): Na Mg Al N O Cl Ar Ar or

5. What’s special about 8? The number of valence e-s in the noble gases
Indicates stability (low energy)

6. Types of bonding
Determined by the difference in electronegativity between atoms involved
Covalent - sharing electrons
Ionic – transfer of electrons
Metallic - sea of electrons

7. Covalent Bonding
Occurs when 2 atoms SHARE valence electrons so that both follow the octet rule
No ions involved
Nonmetal + Nonmetal

8. Example …
F – 7 valence e-s, wants 8 F 3p 3s
One unpaired e-

9. F F Example, cont. Another F also has 7 valence e-s 3p 3p 3s 3s
One unpaired e-

10. Example, cont.
By sharing electrons… F F

11. Example, cont.
By sharing electrons… F F

12. Example, cont.
By sharing electrons… F F

13. Example, cont.
By sharing electrons… F F

14. Covalent bonding
By sharing electrons… F F

15. F F Example, cont. …both end with full orbitals 3p 3p 3s 3s
One unpaired e-

16. Example, cont. F F
8 Valence electrons

17. Example, cont. F F
8 Valence electrons

18. F F Lewis Structure Covalent bond – 2 e-s
Lone pair – not in bond
aka unshared pair F F
Covalent bond – 2 e-s
Molecule – group of atoms held together by covalent bonds

19. Structural Formula F F
Covalent bond – 2 e-s

20. Covalent Compounds
Bond length – distance between 2 covalently bonded atoms
Bond energy – energy required to break a covalent bond
ex.H2 = 436 kJ/mole, N2 = 163 kJ/mole

21. Polar Bonds
Sometimes, the electrons in a covalent bond aren’t shared equally, causing partial charges
Electronegativity difference determines polarity: – nonpolar, – polar, ionic H Cl
Electronegativity = 2.10
Electronegativity = 3.16

22. Properties of Covalent Compounds
Any state (solid, liquid, or gas)
Nonconductors
Low melting point and boiling point

23. Drawing Dot Structures
Count total number of valence e-s
Arrange atoms with bonds between them
-Symmetry, C in center
3. Add lone pairs to fulfill octet rule
-H never gets lone pairs
4. Count total number of e-s in structure
5. Add extra bonds if necessary
-If there are too many e-s in structure

24. Example: CH4 C H Total valence = 4 + 1 + 1 + 1 + 1 =8 Arrange atoms:
Add lone pair(s)
Count e-s in structure = 8
*No extra bonds needed* H C

25. Example CH3I H I C Total valence = 4 + 1 + 1 + 1 + 7 = 14
Arrange atoms:
Add lone pair(s)
Count e-s in structure = 14
*No extra bonds needed* H I C

26. Example C2H4 H H H H C C C C H H H H Count : 4+4+1+1+1+1 = 12 Draw:
3. Add “dots”
4. Re-count: 14
5. Add extra bond(s): = 2 extra e-s = add 1 bond
Double bond (4 e-s) H H H H C C C C H H H H

27. Double/Triple bonds CONS like to multiple bond
Determine number of bonds to add
For each bond you add, you must subtract TWO lone pairs (one from each atom the bond touches)
**Double bond- shorter/stronger than single
***Triple bond – shorter/stronger than double

28. Example N2 N N N N Count: 5+5 = 10 Draw: Add “dots” Re-count = 14
14-10 = 4  2 extra bonds needed N N N N

29. Example HCN H C N H C N Count: 1+4+5 = 10 Draw: Add “dots”
Re-count = 14
14-10 = 4  2 extra bonds needed H C N H C N

30. Example CO2 O C O O C O Count: 4+6+6 = 16 Draw: Add “dots”
Re-count = 20
20-16 = 4  2 extra bonds needed O C O O C O

31. Polyatomic Ions
Group of atoms covalently bonded together that has an overall charge
Add or subtract electrons from initial count

32. - Example NO2- O N O O N O Count: 5+6+6+1 = 18 Draw: Add “dots”
Re-count = 20
20-18 = 4  1 extra bond needed - O N O O N O

33. PO4-

34. Resonance
Occurs when a molecule has more than one plausible structure In nature, the actual structure is an average of all possible resonance structures - - O N O O N O

35. VSEPR Theory Valence shell electron pair repulsion theory
Predicts the 3D shape of molecules
Caused by repulsion of electrons in bonds and lone/unshared pairs
5 shapes (there are many more you will learn in college chemistry)

36. 1. Linear
-1 bond OR -2 bonds with no lone pair(s) on central atom O C O H F

37. 2. Bent H O H 2 bonds with lone pair(s) on central atom

38. 3. Trigonal Planar
3 bonds with no lone pair on central atom F B F F

39. 4. Trigonal Pyramidal
3 bonds with a lone pair on central atom H N H H

40. 5. Tetrahedral
4 bonds from central atom H H C H H

41. Molecular Models Ammonia – NH3 Carbon Sulfide – CS2 Methane – CH4
Propane – C3H8
Formaldehyde – H2CO
Urea – H4N2CO
Glucose – C6H12O6

42. Polar Molecules
Just like individual bonds can be polar, so can molecules
Shape of molecule helps determine if the molecule is polar
If bonds within molecule are not polar, the molecule CANNOT be polar

43. Example – H2O
Bonds are polar; molecule is polar

44. Example – CO2
Bonds are polar; molecule is NOT polar

45. Example – CCl4
Bonds are polar; molecule is NOT polar

46. Example – NH3
Bonds are polar; molecule is polar

47. Ionic Bonding - Cl Na+ Cl Na NaCl Metal bonded to a nonmetal
One atom TRANSFERS electron(s) to another so both follow the octet rule - Cl
Na+ Cl Na
Cation
1s22s22p6
Wants to lose 1e-
1s22s22p63s1
Wants to gain 1e-
1s22s22p63s23p5
Anion
1s22s22p63s23p6
NaCl
formula unit

48. Other examples… Al Cl -
Al3+ 3 Cl Cl
AlCl3 Cl

49. Try these
Na and O
Al and S

50. Lattice Energy
Energy released when one mole of an ionic compound is formed

51. Properties of Ionic Compounds
High melting point and boiling point
Brittle
Dissolve in water
Conduct electricity when molten or dissolved in water
Exist in crystal lattice structure
Made up of formula units

52. Lattice structure

53. Metallic Bonding
Bonding that results from the attraction between the metal atoms (their nuclei) and the surrounding sea of e-s
Empty p and d orbitals allow electrons to be VERY mobile (“sea” of electrons)

54. Properties caused by “sea” of e-s
High electrical and thermal conductivity
e-s become excited and when return to ground state emit light - luster
Malleable
Ductile

55. Intermolecular forces (IMF)
Forces of attraction between molecules
Weaker than “real” bonds
Cause property differences between compounds (boiling point, surface tension …)
3 types you need to know

56. 1. Dipole-Dipole Must have POLAR molecule
Created by attraction between opposite ends of polar molecules
Example: Br-F (polar) -20°C bp
F-F (not) °C bp ɗ+ ɗ- ɗ+ ɗ- ɗ- ɗ+ ɗ- ɗ+

57. 2. Hydrogen Bonding EXTREMELY strong type of dipole-dipole
Occurs when H is bonded to a highly electronegative atom (FON) and there is a nearby lone pair on another highly electronegative atom (FON)
H2O ɗ- ɗ+
Causes many of water’s properties
(high boiling point, surface tension) ɗ- ɗ+
Hydrogen bond ɗ- ɗ+

58. 3. London Dispersion Forces
Weakest IMF
Exists in ALL molecules
Caused by temporary dipoles created as electrons move
More electrons = stronger force

59. Chapter Review
Pg. 209
#’s 1-6, 10-17, 19-21, 24, 28, 30, 46-48