Physical Science Institute
Summer 2013
Melt Away
MATERIALS
ice
(1) hot plate
(1) 600-mL beaker
(2) 12-oz styrofoam cups
(1) 250-mL beaker
(1) 100-mL graduated cylinder
(1) pair of plastic tongs
(1) LabQuest2 data collection device
(1) stainless temperature probe
PROCEDURE
1. Use a 100 mL graduated cylinder to obtain 90.0 mL of water at about 60°C from your hot water bath on the
hotplate. Record this as V1.
2. Obtain a cup of ice.
3. Lower the temperature probe into the warm water (to about 1 cm from the bottom).
4. Begin data collection. Wait until the temperature reaches a maximum (it will take a few seconds for the cold
probe to reach the temperature of the warm water). This maximum will determine the initial temperature, t1, of
the water. Record the maximum temperature of the water, t1, in the data table below.
5. As soon as this maximum temperature is reached, fill the styrofoam cup with ice cubes. Shake excess water from
the ice cubes before adding them (or dry with a paper towel). Use the temperature probe to stir the mixture as
the temperature approaches 0°C. Important: As the ice melts, add more large ice cubes to keep the mixture full
of ice!
6. When the temperature reaches about 4°C, quickly remove the unmelted ice (using tongs).
7. Continue stirring until the temperature reaches a minimum (and begins to rise). This minimum temperature is
the final temperature, t2, of the water. Record t2 in your data table.
8. Use the 100 mL graduated cylinder to measure the volume of water remaining in the Styrofoam cup to the
nearest 0.1 mL. Record this as V2.
DATA
Initial water temperature, t1 (°C)
Final water temperature, t2 (°C)
Initial water volume, V1 (mL)
Final water volume, V2
(mL)
CALCULATIONS AND RESULTS
1. Use the equation ∆t= t2– t1 to determine ∆t, the change in water temperature.
2. Calculate the volume of ice that melted (V2–V1).
3. Find the mass of ice melted using the volume of melted ice (use 1.00 g/mL as the density of water).
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Physical Science Institute
Summer 2013
4. Use the equation Q=mc∆t to calculate the energy (in calories) released by the 90 mL of warm water as it cooled
through ∆t. The specific heat of water is 1.00 cal/g oC.
5. Assuming that the heat energy released by the 90 mL of warm water was used to melt the ice, determine the
heat of fusion of water—the energy required to melt one gram of ice (in cal/g H2O).
6. The accepted value for the heat of fusion for water is 80 cal/g. How does your experimentally determined value
compare with this accepted value?
7. When compared to the accepted value, how would the following experimental errors affect the value for the
heat of fusion of ice—too low, too high, or no effect? Explain your answer.
(a) When ice was added to the warm water at the beginning of the experiment there was excess water on
the ice.
(b) Some of the heat from the warm water transferred to the surroundings, that is, the calorimeter was not
100% insulated.
(c) When measuring the final volume of the water, some of the water is spilled when transferring to the
graduated cylinder.
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