Liquids and Solids

( 10 Days )

Tentative Dates

Topic(s)

Reading

Homework

11/13

  • Specific Heat of a Metal Lab
  • Heat Capacity and Specific Heat, pp 295 - 299 
  • Write-Up: Specific Heat of a Metal Lab

11/14

  • Energy and Units
  • Conservation of Energy
  • Specific Heat Calculations
  • Kinetic and Potential Energy
  • Heat Capacity and Specific Heat, pp 295 - 299 
  • Supplemental Problems: 1-3
  • Pre-Lab: Warming and Cooling Curve Lab

11/15

  • Warming and Cooling Curve Lab
  • Warming and Cooling Curve Lab 
  • Pre-Lab: Warming and Cooling Curve Lab

11/16

  • Kinetic Theory Applied to Solids and Liquids
  • Phase Changes (Melting/Freezing and Boiling/Condensing)
  • Energy Transformations, Exothermic and Endothermic Processes, pp 293 - 295
  • Supplemental Problems: 4 - 8

11/19 

  • Energy Needed to Melt Ice Lab (No Pre-lab Required)
  • Energy Transformations, Exothermic and Endothermic Processes, pp 293 - 295
  • Write-Up: Energy Needed to Melt Ice Lab

11/20 

  • Heat of Fusion of Water
  • Exo- and Endo-thermic Reactions
  • Solid-Liquid Phase Change
  • Heats of Fusion and Solidification , pp 307 - 308 
  • Heats of Vaporization and Condensation, p 310
  • Supplemental Problems: 9 - 15

11/21 

  • Liquid-Gas Phase Change
  • Heat of Vaporization of Water
  • A Model for Liquids, Evaporation, pp 274 - 277
  • Supplemental Problems: 16 - 23

11/26 

  • Vapor Pressure and Boiling Point
  • Heats of Vaporization and Condensation, p 310
  • Supplemental Problems: 24 - 32

11/27 

  • Review
  • Chapter 10 and 11 Study Guide (Neglect Section 11.4)
  • Study!

11/28 or 11/29

  • Trimester Final
   

Labs: Specific Heat Lab, Warming and Cooling Curve Lab, Energy to Melt Ice, Energy of Crystallization*

Supplemental Problems

1. 750 calories are equivalent to how many Joules of energy? (Note the conversions on page 296 of your text.)


2. A piece of iron (look up its specific heat in Table 11.2 on page 296) has its temperature increased by 15.0 oC by the addition of 420 Joules of heat. What is the mass of the metal?

3. When a ball is thrown straight up, at what point or points in its trajectory does it have

a) maximum potential energy?

b) maximum kinetic energy?

c) minimum potential energy?

d) minimum kinetic energy?

4. Discuss the potential and kinetic energy that a skier has

a) as she goes aboard a ski lift.

b) as she travels to the top of the ski slope.

c) as she skis down the hill and reaches the bottom of the ski slope.

 5. Iron metal has a specific heat of 0.108 cal/(g oC). How much heat, in calories, would raise the temperature of 45.0 g of iron 30.0 oC? What would this be in Joules?


 6. A 25 g piece of lead (Specific Heat = 3.30 J/(g oC)) is heated with 420 Joules of energy. What will be the temperature increase of the metal?

 

7. Below is an example of an endothermic reaction. (Energy appears as a reactant.) Find the energy absorbed as 1.0 g of H2O (l) is changed to hydrogen and oxygen gas.

68.3 Kcal + H2O (l) H2 (g) + ½ O2 (g)


8. Below is an example of an exothermic reaction. (Energy appears as a product.) How many grams of carbon are required to produce 150 KJ of energy?

2C (s) + 3H2 (g) C2H6 (g) + 84.5 KJ


9. How much energy must be added to 50 g of ice at 0 oC to get the water to melt completely? Report your answer in Kcal and KJ.

 


10. How much heat energy must be removed to freeze an ice tray full of water at 0 oC if the ice tray holds 540 g of water? What happens to this heat energy?

 


11. Compare the potential energy of 10 g of solid water and 10 g of liquid water, both samples at 0 oC.

 

12. When the temperature drops below 0 oC, the fruit on the trees in an orchard can sometimes be protected by flooding the orchard. Why?

 

13. Refer to Table 11.5 on page 308, which of the substance listed in the tables would be easiest (require the least amount of energy) to melt? Which would be the most difficult?

 

14. 75 grams of water at 25 oC are cooled and then frozen solid at its freezing point.

a) How much energy is removed as the liquid water cools to 0 oC?

b) How much energy is removed as the liquid water at 0 oC is changed to solid water at 0 oC?

 

15. A metal cylinder, initially at 25 oC, is placed in an insulated cup containing water and ice. After the system has come to thermal equilibrium, there is a little ice remaining.

a) What is the final temperature of the metal cylinder?

b) What happens to the energy given up by the metal cylinder?

c) If 20 g of ice were melted, how much energy was given up by the cylinder?

16. In an experiment, 1.0 x 102 g of water at 80 oC are allowed to interact with ice (at 0 oC). Enough ice is melted to cool the water to 0 oC. The ice and water are placed in an insulated container so that all the energy given up by the water melts ice. After the system reaches equilibrium at 0 oC, there are 2.0 x 102 g of liquid water present. Use this information to calculate the information below.

a) How much ice is melted?

b) How much energy is given up by the 1.00 x102 g of water cooling from 80 oC to 0 oC?

c) How much energy is required to melt 1.0 g of ice?

d) How much energy is required to melt 1.0 mole of ice?

17. A liquid is heated at its boiling temperature. Although energy is added to the liquid, its temperature does not increase. Explain.


18. Which is more likely to cause a severe burn, one gram of H2O (g) at 100 oC or one gram of H2O (l) at 100 oC?


19. Two of the methods suggested for conversion of seawater to fresh water are evaporation and the freezing of water. Which process used more energy?


20. How much energy is given off as 100 g of steam at 100 oC is condensed to 100 g of liquid water at 100 oC? Report your answer in Kcal and KJ.

21. 540 grams of water is at 25 oC. How much energy is required to:

a) heat the water to its boiling point without changing any of it to a gas?

b) change the water at 100 oC to steam at 100 oC?

c) heat the water to its boiling point and change it to steam?

22. Refer to Table 11.5 on page 308, which of the substance listed in the tables would be easiest (require the least amount of energy) to boil? Which would be the most difficult?

23. In a paragraph explain the experiment illustrated in Figure 10.9 on page 277 of your text. Include a definition of the vapor pressure in your paragraph.

24. How would the vapor pressure measured in Figure 10.9 be affected if air molecules were present in the flask? How would the vapor pressure measured in Figure 10.9 be affected if the temperature were greater?

25. Define the normal boiling point of a substance.

26. A sample of steam at 100 oC is condensed to give 1.80 X 10 2 g of water, also at 100 oC. All the heat energy is used to vaporize benzene (C6H6). This procedure yields 1030 g of gaseous benzene at its boiling temperature. (Water's heat of vaporization = 9.7 Kcal/mol)

a) Is the molar heat of vaporization of C6H6 greater or less than that of H2O?

b) What is the value for the molar heat of vaporization for benzene in Kcal/mol?

27. How many grams of ice would be melted by the energy obtained as 18.0 grams of steam are condensed at 100 oC and then cooled to 0 oC? How does the amount of energy obtained as the steam condenses compare to the amount of energy released as the same amount of water is cooled from 100 oC to 0 oC?

28. Name two solids that have such a high vapor pressure at room temperature that you can detect their presence using your sense of smell.

29. What does Table 10.11 on page 279 of you text suggest about the strength of the molecular attractions between chloroform, ethanoic acid, ethanol and water. Try to rank them from lowest attraction to highest attraction.

30. Table 10.11 to estimate an approximate boiling temperature for ethanoic acid, CHCl3.

31. Why is the boiling temperature of water lower in Denver, Colorado (altitude 5280 ft, 1600 m) than in Boston, Massachusetts (at sea level)?


32. Both carbon tetrachloride, CCl4, and mercury, Hg, are liquids whose vapors are poisonous to breathe. If CCl4 is spilled, the danger can be removed merely by airing the room overnight. If Hg is spilled, it is necessary to pick up the liquid droplets with a "vacuum cleaner" device. Explain in terms of the vapor pressures of the two liquids.

Home Page|Homework and Schedule|Safety Rules|Supplies|Tests