1. Bonding
Chapters 7-8

2. Octet Rule
Atoms tend to lose or gain electrons to achieve a full valence shell (8)
Exception: First Energy Level is full with 2 electrons

3. Electron Dot Structures
Diagrams that show valence electrons, usually as dots
AKA Lewis Electron Dot Diagrams
Rules
Start on any side
First two get paired together
Next three are separated
Fill in as needed O

4. Examples H He F Ne Ar N Na Cl

5. Ions
Atoms that have gained or lost electrons, and now have a charge
Must show charge
Na+ F-
O-2

6. Practice
Draw Lewis Electron Dot Structures for the following atoms and ions:
Li, B, Mg, Al, P, S, Cl, Br, Kr
H+, Li+, Mg2+, Al3+, O2-, F-, Cl-, S2-, P3-

7. Practice Al Li B Mg P S Kr Cl Br

8. Practice H+
Li+
Mg2+
Al3+
P-3
O-2
Cl- F-
S-2

10. Compounds Two Main Types of Compounds
Ionic
Molecular (Covalent)
Based on type of bonding involved

11. Bonding Bond Three Main Types
Shared or exchanged electrons that hold two atoms together
Three Main Types
Covalent
Ionic
Metallic

12. Covalent Bonds
Electrons are shared between two atoms to hold them together
Each atom will try to achieve a full valence shell
2 nonmetals
Two types of covalent bonds
Non-Polar Covalent – Shared equally
Polar Covalent – Shared unequally

13. Covalent Bonding H2 H H H
Single Bond

14. Covalent Bonding
H2O
Bond O H H
Bond H H

15. More Examples O2 O O O
Double Bond

16. More Examples N2 N N N
Triple Bond

17. More Examples Cl H
HCl
NH3 N H

18. More Examples
CH4
CO2 C H C C O

19. Determining Bond Type
Whether electrons are shared or exchanged is based on electronegativity difference between two bonding atoms
Nonpolar Covalent Bond
2 same Nonmetals (no difference in electronegativity)
Polar Covalent Bond
2 different Nonmetals (small difference in electronegativity)

20. Determining Bond Polarity
The larger the difference in electronegativity, the more polar the bond.

21. Determining Bond Polarity
Which is more polar?
ΔEN H F
Most Polar
Biggest
1.8 H Cl
1.0 H Br
0.8 H I
0.5

23. Bonding
Ionic Bond
Electrons are transferred from one atom to another (one gives, one takes)
Metal and nonmetal, NaCl
Large electronegativity difference

24. Properties Ionic Compounds
Most ionic compounds are hard, crystalline solids at room temperature
High melting points
Mostly soluble in water
Can conduct an electric current when melted or dissolved in water(aq).

25. Properties Covalent Compounds
Most molecular compounds tend to have relatively lower melting and boiling points than ionic compounds.

26. Ionic Compounds Formula Unit
is the lowest whole-number ratio of ions in an ionic compound
Ionic Compounds are repeating lattices of positive and negative ions

27. Ionic Compounds
NaCl

28. Ionic Compounds
Ionic compounds are electrically neutral, even though they are composed of charged ions
Total positive charge equals total negative charge

29. Determining Formulas Must be electrically neutral
Total positive charge must equal total negative charge
Use oxidation numbers from Periodic Table
Group 1  +1 Group 2  +2
Group 13  +3 Group 15  -3
Group 16  -2 Group 17  -1

30. Determining Formulas
Determine number of each ion to balance out charge
Use as subscript for element symbol
Ex: CaCl2, Na3PO4, Mg(NO3)2
Write Positive Ion First
Formula must be smallest whole-number ratio

31. Example
Sodium and Chlorine
Na+
Cl-
Na1Cl1
NaCl

32. Example
Calcium and Fluorine
Ca+2 F- F-
Ca1F2
CaF2

33. Examples
Potassium and Oxygen K+ K+
O-2
K2O1
K2O

34. Polyatomic Ions
Group of atoms that collectively have gained or lost electrons (Table E)
Sodium and Nitrate
Na+
(NO3)-
Na1(NO3)1
NaNO3

35. More Examples
Potassium and Sulfate
Ammonium and Sulfur
K2SO4
(NH4)2S

36. Short-cut (criss-cross method)
Magnesium and Phosphate
Mg+2
PO4-3
Mg3(PO4)2

37. Short-cut (criss-cross method)
Magnesium and Carbonate
Mg+2
CO3-2
Mg2(CO3)2
Must Simplify
MgCO3

39. Na2SO4 Ca3(PO4)2 Do This Now Write out formulas for: Sodium Sulfate
Calcium Phosphate
Na2SO4
Ca3(PO4)2

40. [ ] Cl Na+ Cl- - Dot Structures Shows valence electrons
Must show charge for Ions
NaCl Cl
[ ] -
Na+
Cl-

41. Dot Structures
MgO
Mg+2
O-2

42. Dot Structures
CaF2
Ca+2 F-

44. Polyatomics
Compounds with polyatomic ions contain BOTH ionic and covalent bonds
Example: NaNO3 - N O
Na+

45. Network Solids
All atoms in a network solid are covalently bonded together
Network solids have very high melting and boiling points, since melting requires the breaking of many bonds throughout the compound.
Some of the strongest materials known to man are network solids.

46. Network Solids Diamonds ( C ) Graphite ( C ) Silicon Dioxide (SiO2)
Silicon Carbide (SiC)

47. Metallic Bonding
Bonding within metallic samples is due to highly mobile valence electrons
Free flowing valence electrons
“Sea of Electrons”

48. Metallic Substances
High melting points
Conducts as a solid

49. Bond Energy When two atoms form a bond, energy is released
Example: Cl + Cl  Cl2 + energy
Energy needs to be added to break a bond
Example: Cl2 + energy  Cl + Cl

51. O H N H O H Cl Structural Formulas
Shared electrons are written as a line, unshared electrons are not written
Each line represents 2 electrons O H N H O H Cl

52. Molecular Polarity Polar Molecule Nonpolar Molecule
one end of a molecule is slightly negative(δ-) and the other end is slightly positive(δ+).
Asymmetrical charge distribution
Nonpolar Molecule
Can not be separated into different ends
Symmetrical charge distribution

53. O H Polar Molecule H2O Polar Covalent Bond Electrons shared Unequallyδ- O H δ+

54. More Examples
HCl
NH3 H Cl δ- δ+ δ- N H N H δ+

55. Another Example
CH4
Nonpolar Molecule δ+ C H δ- δ+ δ+ δ+

56. Polarity Ionic Compounds are Ionic
Nonpolar Covalent Bonds always indicate Nonpolar Molecules
Polar Covalent Bonds
Determine Symmetry

57. “Like Dissolves Like”
Polar and Ionic substances will dissolve in other Polar Substances
Nonpolar substance will dissolve in other nonpolar substances

59. Intermolecular Forces
Intermolecular Forces of Attraction
attraction between two molecules or ions that hold them together (not a bond)
Determines melting and boiling points of compounds
Stronger intermolecular forces, higher melting and boiling points

60. Intermolecular Forces
Van der Waals
Dispersion
Dipole-Dipole
Molecule-Ion
Hydrogen Bonding
Weakest
Strongest

61. Van der Waals Dispersion
Electrons of one atom are attracted to the Protons of the next atom.
Also called an induced dipole
Attraction increases with increasing mass e e p p p p e e

62. H Cl H Cl Van der Waals Dipole-Dipole
negatively charged end of polar molecule is attracted to positively charged end of another polar molecule δ- δ+ δ+ δ- H Cl H Cl

63. Molecule Ion
Attraction between polar molecules and ions in solution O H O H
Na+
Cl-

64. Hydrogen Bonding
Hydrogen bonded to N, O, or F, is attracted to the N, O, or F of another molecule.
Not actual bond, just attraction
Hydrogen “Bond” H F

65. Boiling Point of H compounds

66. Boiling Point of H compounds