Latent Heat of Fusion

 Latent Heat of Fusion

Supplies: Thermometer, styrofoam cups, scale, towels, ice, water, hammer (or something to smash the ice).

Introduction:

Last week you explored heat energy and learned about specific heat. To review, specific heat is the number of calories it takes to raise the temperature of a gram of a substance 1 °C (or, in SI units, the joules to raise 1 kg 1 °C). This is just about changing the temperature of the object.

However, substances can also experience phase changes. Solid, liquid, gas, and plasma are phases of matter, and at a microscopic level describe how the particles are bound to each other. (For example, atoms in a gas have weak to no bonding with other atoms, but atoms in a solid are strongly bound to other atoms.)

It takes energy to break these bonds, so you have to input energy to transition a solid to a liquid, for example. (If you transition the other way, from liquid to solid, you get energy out.) Notice that this energy is not going to raise the temperature of the substance! (Temperature is about how much the particles are moving around, related to their kinetic energy. The bonds between particles are related to potential energy.)

Breaking these bonds requires a lot of energy, far more than raising the temperature by a degree. The latent heat of fusion (Lf) is the amount of energy it takes for a substance to go from solid to liquid, while latent heat of vaporization (Lv) is the amount of energy it takes for a substance to go from liquid to gas. This week we are going to determine experimentally the Lf for water; specifically, the number of calories it takes to melt one gram of ice.

The formula describing the heat needed to melt a substance is:

Q=mLf

where m=mass (in this lab, we’ll keep the mass in grams) and Lf is calories/gram. Notice that there is no T in the formula because the substance sits at the same temperature the whole time the phase change is taking place.

Procedure:

1. Take the ice out of the freezer a little while before you begin your experiment. We generally assume the ice begins at 0 °C, and letting the ice sit out will help this assumption be accurate.

2. Record room temperature in Table 7.1 below.

3. Record the mass of your empty cupin Table 7.1. You’ll be subtracting this so you can find the mass of the water alone.

4. Put ~150 or so grams of warmwater(6–8 degrees °C above room temperature) into your insulated cup. Record the mass in Table 7.1, showing your calculation here (or explaining what you did to avoid a calculation). Also record its temperature immediately before adding ice.

5. Put twoice cubes into a towel and break them up with a hammer (you can use last week’s rock if you lack a hammer). Add crushed ice to your warm water. Dissolve it, stirring slowly, and add more ice as necessary to bring the water to a final temperature that is around the same amount below room temperature as you started above.

6. Record the final temperature of your water-ice mixturein Table 7.1. Be sure all the ice has melted before doing this.

7. Record the final mass of your water-ice mixturein Table 7.1, being sure to subtract the mass of the empty cup. Show your work here.

8. Calculate the mass of the ice alone, and record that value in Table 7.1. Show your work here.

9. Explain why you would want to have your experiment have equal temperatures above and below room temperature.

Table 7.1

mass of empty cup mass of warm water room temperature initial temperature of warm water final temperature of mixture mass of water + ice mass of ice alone

10. Calculate the temperature change (T) for the warm water.

11. Calculate the temperature change (T) for the cold water that used to be ice.

12. Explain where the heat energy transferred from the water water goes. Hint: it does two things.

13. Calculate how much heat energy left the warm water. (Refer back to Lab 6 as needed.) Show your work.

14. Calculate how much energy went into melting the ice. Show your work.

15. Solve for Lf. Show your work. (Note that Lf is positive.)

Consider:

16. What would contribute to the uncertainty in the experiment? Consider each step of your procedure, and don’t give a lame answer like “experimental error” (or worse, “human error”).

17. Using the same equipment, how could you improve your results? In other words, what could you do differently to improve your results?

Latent Heat of Fusion

Supplies: Thermometer, styrofoam cups, scale, towels, ice, water, hammer (or something to smash the ice).

Introduction:

Last week you explored heat energy and learned about specific heat. To review, specific heat is the number of calories it takes to raise the temperature of a gram of a substance 1 °C (or, in SI units, the joules to raise 1 kg 1 °C). This is just about changing the temperature of the object.

However, substances can also experience phase changes. Solid, liquid, gas, and plasma are phases of matter, and at a microscopic level describe how the particles are bound to each other. (For example, atoms in a gas have weak to no bonding with other atoms, but atoms in a solid are strongly bound to other atoms.)

It takes energy to break these bonds, so you have to input energy to transition a solid to a liquid, for example. (If you transition the other way, from liquid to solid, you get energy out.) Notice that this energy is not going to raise the temperature of the substance! (Temperature is about how much the particles are moving around, related to their kinetic energy. The bonds between particles are related to potential energy.)

Breaking these bonds requires a lot of energy, far more than raising the temperature by a degree. The latent heat of fusion (Lf) is the amount of energy it takes for a substance to go from solid to liquid, while latent heat of vaporization (Lv) is the amount of energy it takes for a substance to go from liquid to gas. This week we are going to determine experimentally the Lf for water; specifically, the number of calories it takes to melt one gram of ice.

The formula describing the heat needed to melt a substance is:

Q=mLf

where m=mass (in this lab, we’ll keep the mass in grams) and Lf is calories/gram. Notice that there is no T in the formula because the substance sits at the same temperature the whole time the phase change is taking place.

Procedure:

1. Take the ice out of the freezer a little while before you begin your experiment. We generally assume the ice begins at 0 °C, and letting the ice sit out will help this assumption be accurate.

2. Record room temperature in Table 7.1 below.

3. Record the mass of your empty cupin Table 7.1. You’ll be subtracting this so you can find the mass of the water alone.

4. Put ~150 or so grams of warmwater(6–8 degrees °C above room temperature) into your insulated cup. Record the mass in Table 7.1, showing your calculation here (or explaining what you did to avoid a calculation). Also record its temperature immediately before adding ice.

5. Put twoice cubes into a towel and break them up with a hammer (you can use last week’s rock if you lack a hammer). Add crushed ice to your warm water. Dissolve it, stirring slowly, and add more ice as necessary to bring the water to a final temperature that is around the same amount below room temperature as you started above.

6. Record the final temperature of your water-ice mixturein Table 7.1. Be sure all the ice has melted before doing this.

7. Record the final mass of your water-ice mixturein Table 7.1, being sure to subtract the mass of the empty cup. Show your work here.

8. Calculate the mass of the ice alone, and record that value in Table 7.1. Show your work here.

9. Explain why you would want to have your experiment have equal temperatures above and below room temperature.

Table 7.1

mass of empty cup mass of warm water room temperature initial temperature of warm water final temperature of mixture mass of water + ice mass of ice alone

10. Calculate the temperature change (T) for the warm water.

11. Calculate the temperature change (T) for the cold water that used to be ice.

12. Explain where the heat energy transferred from the water water goes. Hint: it does two things.

13. Calculate how much heat energy left the warm water. (Refer back to Lab 6 as needed.) Show your work.

14. Calculate how much energy went into melting the ice. Show your work.

15. Solve for Lf. Show your work. (Note that Lf is positive.)

Consider:

16. What would contribute to the uncertainty in the experiment? Consider each step of your procedure, and don’t give a lame answer like “experimental error” (or worse, “human error”).

17. Using the same equipment, how could you improve your results? In other words, what could you do differently to improve your results?