1. Unit 6: Ionic & Covalent Compounds, Lewis Structures
Chapter 5 & 6

2. Key Concepts Ionic, Covalent, Metallic Compounds
Polar/Nonpolar Covalent Bonds
Bond Length, Bond Strength
Resonance Structures
Lewis Dot Structures
VSEPR Theory
Molecular Shapes
Molecular Polarity

3. Compounds Ionic (M/NM): Covalent (NM/NM): Metallic (M/M):
e- are transferred from a metal to a nonmetal forming charged ions that stick together
Covalent (NM/NM):
e- are shared in overlapping molecular orbitals
Metallic (M/M):
e- move freely between atoms

4. Octet Rule
Octet Rule: atoms gain or lose e- to attain an e- configuration of the nearest noble gas
Duet Rule: 2 val e- for hydrogen
Octet Rule: 8 val e- for all others

5. Ions Ion: atom/molecule with a charge+1 +2 ±4 -3 +3 -2 -1
TM form ions with various charges
Elements always gain or lose the same number of e- when they form ionic compounds
(except TM)
When an atom or molecule gains a negative electron, it becomes a negative ion
When an atom or molecule loses a negative electron, it becomes a positive ion

6. Ionic Compounds (M/NM)

7. Ionic Compounds (M/NM)
Ionic (M/NM):
e- are transferred from a metal to a nonmetal forming charged ions that stick together
Cation: ion with a positive charge
Anion: ion with a negative charge
All ionic compounds are salts
Soome common salts: potassium chloride, potassium iodide
Cation
Anion
[Overall charge is neutral]
Ionic compounds (salts)
form a crystal lattice structure

8. Covalent Compounds (NM/NM)
Covalent (NM/NM):
e- are shared in overlapping molecular orbitals

9. Metallic Compounds Metallic (M/M):
Delocalized e- move freely between atoms
Metals are conductive (they conduct electricity) because their e- move.
Metals are malleable/not rigid because their e- move.

10. Ionic Compounds (M/NM)
Ionic (M/NM):
e- are transferred from a metal to a nonmetal forming charged ions that stick together
Cation: ion with a positive charge
Anion: ion with a negative charge
All ionic compounds are salts
Soome common salts: potassium chloride, potassium iodide
Cation
Anion
[Overall charge is neutral]
Ionic compounds (salts)
form a crystal lattice structure

11. Unknown Ionic Compounds Lab
Identify the metal cation and nonmetal anion components of an unknown ionic compound M NM
*Every group will receive a different set of unknown compounds to identify

12. Unknown Ionic Compounds LabNM

13. Unknown Ionic Compounds LabNM
Flame Test

14. Unknown Ionic Compounds LabNM
Flame Test
AgNO3 Test

15. Unknown Ionic Compounds LabNM
Flame Test
HCl Test
AgNO3 Test

16. Covalent Compounds (NM/NM)
Covalent (NM/NM):
e- are shared in overlapping molecular orbitals

17. Covalent Bonds Nonpolar Covalent Bond: e- are shared equally
Polar Covalent Bond: e- are attracted to the atom with higher electronegativity
Electronegativity increases as you move up and to the right on the periodic table.

18. Covalent Bonds
Ex. Determine polar or nonpolar. Draw a dipole arrow towards the more negative side
H-Cl
O-H
C-C
O-F
S-O + -
Dipole arrow points towards the more electronegative atom.
Electronegativity increases as you move up and to the right on the periodic table.
H-C +
O-H +
C-C =
O-F +
S-O +

19. Polarity C - C C - Cl Na - Cl Difference in Electronegativity None
High
Nonpolar
Covalent
Polar
Covalent
Ionic
C - C
C - Cl
Na - Cl

20. Bond Length & Bond Strength
Atoms form single, double, or triple bonds depending on what’s needed to make an octet
Longer/Weaker
Shorter/Stronger

21. Resonance
Resonance: more than one Lewis structure can be drawn because there’s an extra bond that moves between two or more locations
The bonds have the same length and strength! They’re not like one single bond and one double bond, but more like two 1 ½ bonds
Ex.
SO2
CO32-
Laboratory measurements show that all the bonds are equal in length, and between single and double bond length.
None of these Lewis structures drawn are actually correct. The actual structure has equal bonds between single and double bond length and strength.
The structures are changing so quickly that all we can measure is an average blur instead of being able to detect individual structures.

22. Lewis Structures Lewis Structures: drawing valence e- using dots
Count the valence e- in the highest energy level (the outermost orbital)
Covalent bonds
are drawn using lines
2 e- in each bond!
Both need one more e-
to fulfill the octet rule
They share each other’s e-
Shared e- form a covalent bond
Both chlorines have 7 val e-. They share electrons so they both get 8 val e- (octet rule).
If there’s two chlorine atoms and they both have eight val e-, how many total val e- are there?
14, not 16 because they share two of them.

23. Lewis Structures Drawing Lewis structures:
Count the total number of valence e-
Draw the SKELETON Lewis structure
Arrange the atoms with carbon in the middle
Draw single bonds between atoms (2 e- in each bond)
Satisfy octet/duet rule for each atom
Draw 8 e- around each atom
Confirm the total number of valence e-
Compare to the total number of valence e- you counted to the total number you should have (from step 1)
If there are too many e-, add double/triple bonds until they are equal
If the molecule is an ion, enclose the structure in parentheses and put the charge at the top right on the outside of the parentheses.

24. Lewis Structures
Ex: Draw the Lewis structure for O2

25. Lewis Structures Ex: Draw the Lewis structure for O2
6 + 6 = 12 valence e-

26. Lewis Structures
Ex. Draw the Lewis structure for CH3I

27. .. Lewis Structures Ex. Draw the Lewis structure for CH3I
= 14 valence e-

28. Electron Pairs Bond Pairs: e- pairs in a bond (line)
Lone Pairs: e- pairs not in a bond (dots)
Lone Pairs
Bonded Pairs

29. VSEPR Theory
Valence Shell Electron Pair Repulsion Theory: the shape of a molecule is determined by the valence e- around the central atom
Valence e- pairs (bond & lone) repel each other and spread out as far as possible
Balloons: visual demo of orbitals repelling each other (spreading out).

30. Hint: atoms always want 4 pairs (8 e-) according to the octet rule!
Molecular Shapes ..
Linear, trigonal planar, and tetrahedral were the shapes of the balloons.
Trigonal pyramidal is just like tetrahedral except one of the bonded pairs is invisible because it’s a lone pair.
Bent is just like tetrahedral except two of the bonded pairs is invisible because they’re lone pairs.
2 bonded pairs
(two double bonds)
2 bonded pairs
1 or 2 lone pairs
3 bonded pairs
1 lone pair
3 bonded pairs
(one double bond)
4 bonded pairs
Hint: atoms always want 4 pairs (8 e-) according to the octet rule!

31. Molecular Models Activity
Clay = atoms
Toothpicks = bonded pairs (2 e-)
Thumbtack = lone pair (2 e-)
Draw the lewis structures and write the name of the molecular shape.
Then make the molecules in their correct shapes with clay, toothpicks, and thumbtacks.

32. NONPOLAR: there are no poles
POLAR: there are poles
Cut the molecule in half:
If it has different sides (poles), it’s polar.
If it’s the same all around, it’s nonpolar.
NONPOLAR: there are no poles

33. Dipoles reinforce one another
Molecular Polarity
Molecular Polarity: determine whether bond dipoles cancel out or make + and - poles
Dipoles reinforce one another
Polar Molecule
Dipoles cancel out
Non-Polar Molecule
Hint: if all the atoms around the central atom are the same, the molecule is non-polar

34. Lewis Structure, Shape, Polarity
Molecule
Lewis Structure
Molecular Shape
Molecular Polarity
CH4
HCN
SO2
NO3-

35. Lewis Structure, Shape, Polarity
Molecule
Lewis Structure
Molecular Shape
Molecular Polarity
CH4
HCN
SO2
NO3-
8 e-
10 e-
18 e-
24 e-

36. Lewis Structure, Shape, Polarity
Molecule
Lewis Structure
Molecular Shape
Molecular Polarity
CH4
HCN
SO2
NO3-
8 e-
10 e-
18 e-
24 e-

37. Lewis Structure, Shape, Polarity
Molecule
Lewis Structure
Molecular Shape
Molecular Polarity
CH4
Tetrahedral
HCN
Linear
SO2
Bent
NO3-
Trigonal Planar
8 e-
10 e-
18 e-
24 e-

38. Lewis Structure, Shape, Polarity
Molecule
Lewis Structure
Molecular Shape
Molecular Polarity
CH4
Tetrahedral
Non-Polar
HCN
Linear
Polar
SO2
Bent
NO3-
Trigonal Planar
N/A
8 e-
10 e-
18 e-
24 e-