1. I. Bond Polarity and IMF (237 – 241)
Ch. 8 – Covalent Compounds
I. Bond Polarity and IMF (237 – 241)

2. A. Bond Polarity
Most bonds are a blend of ionic and covalent characteristics
Difference in electronegativity determines bond type

3. A. Bond Polarity Electronegativity
Attraction an atom has for a shared pair of electrons
higher e-neg atom  -
lower e-neg atom +
Draw the Lewis structure for HCl & label partial charges

4. A. Bond Polarity Electronegativity Trend (p. 178)
Increases up and to the right

5. A. Bond Polarity Nonpolar Covalent Bond e- are shared equally
symmetrical e- density
usually identical atoms

6. + - A. Bond Polarity Polar Covalent Bond e- are shared unequally
asymmetrical e- density
results in partial charges (dipole) + -

7. A. Bond Polarity Determine bond polarity: C – O Ca – O PC
Si – Cl
H – F
N – N
3.44 – 2.55 = 0.89
3.44 – 1.00 = 2.44
3.16 – 1.90 = 1.26
3.98 – 2.20 = 1.78
3.04 – 3.04 = 0.00 PC
Ionic
NPC

8. A. Bond Polarity
Nonpolar
Polar
Ionic
View Bonding Animations.

9. B. Molecular Polarity
Polar molecule = one end slightly + and one end slightly –
Molecule with 2 poles = dipolar molecule or dipole

10. B. Molecular Polarity
Shape, symmetry and bond polarity determines molecular polarity
H – O bond is polar and water is asymmetrical, so H2O is polar
C – Cl bond is polar, but CCl4 is symmetrical, so molecule is nonpolar

11. B. Molecular Polarity Identify each molecule as polar or nonpolar SCl2
CS2
CF4
CH2F2
Tetrahedral, bent → polar
Nonpolar bonds → nonpolar
Linear → nonpolar
Tetrahedral → nonpolar
Tetrahedral → polar

12. C. Definition of IMF IMF = Intermolecular Forces
Attractive forces between molecules
Much weaker than chemical bonds within molecules

13. D. Types of IMF
Van der Waals

14. D. Types of IMF
London Dispersion Forces
View animation online.

15. D. Types of IMF
Dipole-Dipole Forces + -
View animation online.

16. D. Types of IMF
Hydrogen Bonding

17. E. Determining IMF PCl3 polar = dispersion, dipole-dipole CH4
nonpolar = dispersion HF
H-F bond = dispersion, dipole-dipole, hydrogen bonding

18. F. Network Solids
Substances in which all atoms are covalently bonded to each other
Very stable
Examples
Diamonds – Carbon covalently bonded to carbon
Quartz – SiO2 covalently bonded and not distinct molecules

19. Ch. 7 – Ionic Bonds & Properties
II. Ions (p. 194 – 200)

20. A. Formula Unit
The lowest whole-number ratio of ions in an ionic compound

21. B. Ionic Bonds
Oppositely charged ions attract, force that holds them together = ionic bond
Electrons are transferred from cations to anions
Bonds formed between metals and nonmetals (or contain a polyatomic ion)

22. B. Properties of Ionic Compounds
Most ionic compounds are crystalline solids at room temp
Ionic compounds generally have high melting points
Large attractive forces result in very stable structures

23. III. Bonding in Metals (p. 201 – 203)
Ch. 7 – Ionic and Metallic Bonding
III. Bonding in Metals (p. 201 – 203)

24. A. Metallic Character
Metals
Nonmetals
Metalloids

25. B. Metals
good conductors because the valence electrons are able to flow freely
Valence electrons of metals can be thought of as
a sea of electrons
Properties can be explained by the mobility of electrons in metals

26. C. Metallic Bond
Metallic Bonding - “Electron Sea”

27. D. Metallic Properties
Properties can be explained by the mobility of electrons in metals
When subjected to pressure , the cations easily slide past each other like a ball bearing immersed in oil

28. E. Summary METALLIC e- are delocalized among metal atoms
Bond Formation
e- are delocalized among metal atoms
Smallest Unit
“electron sea”
Physical RT
solid
Melting
Point
very high