1. Intermolecular Forces

2. In your notes: Draw the Lewis structure of CF4.
Draw the Lewis structure for CH3F.
What is a dipole?
How do they occur?
How do we show dipoles?

3. Intra vs. Inter Intermolecular force—holds molecules together.
Examples: dipole-dipole, London dispersion, etc.
Intramolecular force—holds atoms together in a molecule (any chemical bond).
Example: Metallic, ionic, and covalent bonds.

4. Intermolecular Forces
What are they?
The forces of attraction between particles or molecules of a substance.
Vary in strength
Related to how polar or nonpolar the substance is.
Still weaker than actual covalent, ionic, or metallic bond.

5. London Dispersion Forces
Result from the constant motion of electrons and the creation of instantaneous dipoles in a molecule.
Weakest force of attraction
Exists in all molecules and compounds
Only intermolecular force for noble gases and nonpolar compounds.
Increase with increasing atomic or molecular mass.
Why? Polarizability

6. Dipole-Dipole Forces
Strong force of attraction, but weaker than ionic or metallic forces of attraction.
Exists between polar molecules
Why so strong?
Positively charged regions of one molecule are attracted to negatively charged regions of the other.

7. Hydrogen Bonding
H is bonded to a highly electronegative atom is attracted to the unshared electrons of another electronegative atom in a nearby molecule.
Force is characterized by compounds having unusually high boiling points.
Happens with compounds containing O, F, and N.

8. Ion-Dipole Exist between an ion and a polar molecule.
Driving force behind dissociation of ionic compounds in polar liquids.

9. Relative Strengths

10. In order of increasing strength:
Dispersion forces
Dipole-dipole
Hydrogen bonding
Ion-dipole

11. Chapter 10
States of Matter

12. Section 1
Gases

13. State of Matter
Chemical and Physical properties determined by two things:
Composition—what atoms make up the substance.
Structure—how the atoms are arranged.

14. Kinetic Molecular Theory (KMT)
What is it?
The idea that particles of matter are constantly in motion.
Used for describing the properties of solids, liquids, and gases in terms of energy

15. gases
2 types
Ideal gas—hypothetical gas that fits all assumptions of the KMT.
Real gas—gas that does not behave like KMT assumptions.

16. For Ideal Gases: 5 assumptions
1.) Gases have large numbers of tiny particles that are far apart relative to their size.
Most volume occupied by a gas is empty space
Explains why gases have lower densities than liquids or solids.
Explains why gases can be compressed easily

17. 4.) There are no forces of attraction between particles.
2.) Collisions between gas particles are elastic (no net loss of energy).
Energy is transferred between particles during collisions
3.) Particles are in continuous, rapid motion, and therefore possess kinetic energy.
Kinetic energy of particles overcomes the attraction between them.
4.) There are no forces of attraction between particles.

18. 5.) Temperature of a gas is dependent on the average kinetic energy of the gas.
KE = ½ mv2
m = mass
v = velocity
All specific gases contain particles that have the same mass. Therefore, KE only depends on speed of particles.
Speed and KE depend on temp.
All gases at the same temperature have the same KE. Therefore, gases with different masses have different speeds.

19. Rank the following in order of increasing boiling point.
MgCl2 F2
NaOH
CH2Cl2

20. The Nature of Gases Expansion—gases fill their container.
KMT explains why 3 4

21. The Nature of Gases
Fluidity—the ability of gases to behave like liquids
Particles flow past one another.
Why? Assumption 4

22. The Nature of Gases
Low Density—particles are spread farther apart than those of solids and liquids.
Assumption 1

23. The Nature of Gases
Compressibility-gas particles are initially far apart, and after compression, get much closer together.
Why? Assumption 1—particles are far apart.

24. The Nature of Gases
Diffusion—spontaneous mixing of the particles of two substances caused by random motion.
Effusion—gas particles passing through a small opening.
Rate of effusion is directly proportional to velocity of particles.
Inversely proportional to square root of the molar mass. (the smaller the particles, the quicker the rate of effusion).

25. Deviations from Ideal Behavior
All gases have some degree of variation from ideal behavior.
When does KMT hold true? Usually for gases whose particles have no attraction (noble gases).
Polar gases deviate more than nonpolar gases.
NH3 > F2 > Ne
High temperature and pressure also make gases deviate from KMT.

27. Liquids

28. Liquids
Although most abundant, it is the least common state of matter on earth
Why?
Narrow range of temperature and pressure
What makes something a liquid or a gas?

29. Properties of Liquids Particles are in constant motion
Particles are closer together than gases Why?
Stronger attractive forces
More ordered—strong IMF = lower mobility of particles
Particles are not bound together in fixed positions

30. Properties of Liquids Fluidity—particles flow past one another.
Relatively High Density
Relative Incompressibility
Particles are more packed together
Ability to diffuse
Much slower than diffusion of gases
Why? Attractive forces!!!!

31. Surface tension
Attractive forces tend to pull adjacent parts of a liquid’s surface together. Why?
Surface tension = the energy required to increase the surface area of a liquid by a unit amount.
Particles are drawn to body of liquid, creating sphere (smallest surface area).
Surfactants—substances used to lower surface tension. (Ex—soap or detergent.)

32. Capillary Action Liquids being drawn up narrow cylinders
Cohesion forces—attractive force between identical molecules
Adhesion forces—attractive forces between molecules that are different
Adhesion forces are greater, so water climbs up the glass.

33. Viscosity The measure of the resistance to flow
The stronger the IMF between molecules the greater the resistance to flow
Larger mass increases viscosity
Higher temperature decreases viscosity