1. Vapor and it Pressure

2. Vapor Definition: Gas phase of a substance that is normally a liquid at room temperature. Some particles have enough KE to “escape” from the surface of the liquid. Generally, the easier a substance evaporates, the weaker the intermolecular forces. Some evaporation occurs at all temperatures.

3. Evaporation (con’t) Reverse of this process is called Particles with lower KE are recapture by the liquid. Rate of Evaporation depends on: 1. Nature of the liquid (polar or nonpolar) 2.Temperature of the liquid 3.Surface Area of the liquid

4. Open System Fig a: Open System Evaporation continues until all liquid is gone. Closed System Fig b: Closed System Eventually, rate of evaporation = rate of condensation. System appears static.

5. escape The molecules in the liquid phase have a range of kinetic energies. Some of them have enough kinetic energy to “escape” from the liquid phase.

6. Equilibrium Vapor Pressure

7. The pressure exerted by the molecules of a vapor which are in equilibrium with its corresponding liquid. Solids can also have an equilibrium vapor pressure since some solids change directly to vapor.  Examples: I 2 and Dry Ice (CO 2 ) Depends only on temperature!

8. Vapor Pressure and Temperature The equilibrium vapor pressure has an exponential temperature dependence for any given substance. We can see this from the liquid/vapor equilibrium curve:

9. Vapor Pressure Pressure exerted by a vapor over its liquid. How is vapor pressure affected by temperature? The higher the temperature, the higher the vapor pressure.

10. For all liquids, vapor pressure increases as temperature increases.

11. Vapor Pressure Vapor pressure does NOT depend on how much liquid is present. As long as some liquid is present, can be a teaspoon or a gallon. Vapor pressure depends only on…

12. Intermolecular Forces in Liquids WEAK FORCES High vapor pressure High rate of evaporation Low boiling point Small H v STRONG FORCES Low vapor pressure Low rate of evaporation High boiling point Large H v

13. Phase Changes Melting Boiling (occurs throughout, constant temperature)  Evaporation (occurs at surface; temperature drops) Sublimation Freezing Deposition Condensation

14. Phase Changes are accompanied by energy changes. Which phase changes absorb energy? Melting, Boiling Melting, Boiling Which phase changes release energy? Condensation, Freezing Condensation, Freezing

15. System’s Perspective... Endothermic Process Exothermic Process

16. Boiling Point Temperature at which the vapor pressure of a liquid = external or atmospheric pressure. Normal Boiling Point Normal Boiling Point is temperature at which vapor pressure of the liquid = 1 atm. (or standard pressure conditions) Substances with high boiling points have strong molecular interactions.

17. Boiling or Vaporization

18. Melting Point Temperature at which the liquid phase and solid phase of a substance can coexist at equilibrium. The higher the melting point, the stronger the molecular interactions.

19. Freezing Point Temperature at which a liquid is converted to a crystalline solid. How does freezing point compare to melting point?

20. Strong or Weak attractive forces? High vapor pressure Large H v High boiling point Low boiling point Low vapor pressure Small H v Evaporates rapidly Evaporates slowly Weak Strong Weak Strong Weak Strong

21. Boiling and Pressure If you increase the external pressure (say you are camping in Death Valley), the boiling point is ________ than 100 o C. If you decrease the external pressure (say you are eating Raman noodles at the top of Mt. Whitney), the boiling point is ________ than 100 o C. > <

22. Boiling vs. Evaporation BOILING Occurs throughout the liquid Occurs at a definite temperature Occurs at a constant temperature Only spontaneous when the vapor pressure equals the external pressure EVAPORATION Occurs at the surface of the liquid (or solid) Occurs to some extent at all temperatures The liquid cools off as some of it evaporates Spontaneous all the time