Experiment 3: Measuring the Enthalpy of Fusion of Water

As mentioned earlier, the amount of energy required to convert a solid to a liquid at constant pressure and temperature is called the heat of fusion of the substance. The molar heat of fusion is the amount of energy required to completely change one mole of a solid, at its melting point, into a liquid. In this experiment, the molar heat of fusion of ice will be determined. The change of state can be described as:

1 mol H2O (s) + molar heat of fusion → 1 mol H2O (l)

Figure 5: The term “enthalpy of fusion” refers to the change in enthalpy that occurs from heating one mole of a substance to change its state from a solid to a liquid. This phase change occurs at the “melting point”.

The ice will be melted by placing it in a known volume of hot water contained ina calorimeter. No stirring will occur until all the ice has melted. The heat lost by the water will be absorbed by the melting ice. The volume of the ice that melts can be determined by measuring the volume of the water in the cup before the ice is added and after the ice has melted. If the mass of the ice melted and the heat absorbed by the ice are known, the heat required to melt one mole of ice can be calculated.

Materials:
Calorimeter (2 Styrofoam™ cups, 600 mL Glassbeaker, Styrofoam™ cup lid, and thermometer)
Stir rod
100 mL Graduated cylinder
250 mL Beaker
Plastic fork / *Ice
*Water
*Microwave or hot water bathr
*You must provide

Procedure

  1. Use a graduated cylinder to measure and pour approximately 100 mL of tap water to a 250 mL beaker.
  2. Use a microwave or hot water bath to heat the water to 60 °C (use your thermometer to determine the temperature).
  3. While your water is heating, fill the calorimeter (two stacked Styrofoam™ cups) halfway with ice cubes. Place the calorimeter in a 600 mL beaker for support.
  4. When the temperature of the water in Step 1 has reached 60 °C, preheat the 100 mL graduated cylinder. To do this:
  5. Measure 20 mL portion of hot water in a 100 mL graduated cylinder.
  6. Rinse this water in the cylinder by swirling the water inside several times. Try to coat as much of the walls of the cylinder with the hot water as possible.
  7. Discard the rinse down the drain.
  8. Repeat with a second 20 mL portion of hot water.

Hint: It may help to heat the water to a temperature greater than 60 °C and wait for the temperature to drop down to 60 °C exactly; rather than trying to heat the water to 60 °C.

  1. After pre-heating the cylinder, measure and pour 30 mL of the hot water into the graduated cylinder. Record the volume of this water to the nearest mL in Table 4.
  2. Quickly measure and record the temperature of the water to the nearest 0.1 °C in Table 4.
  3. Drain any excess water that may have accumulated from the ice cubes in the Styrofoam™ cup.
  4. Pour the hot water from the graduated cylinder into the calorimeter. Use the stir rod to stir the ice water until the water temperature falls to 2 °C.
  5. If the temperature does not drop to 2 °C within 2 minutes, add an additional piece of ice. If all of the ice melts, add one more piece of ice so that ice is present in the cup for Step 9.
  1. Place the lid on the Styrofoam™ cup and place the thermometer in the lid.
  2. Continue to swirl the Styrofoam™ cup for approximately two minutes, or until the temperature reading on the thermometer stabilizes. Record the lowest temperature of the mixture of ice and water in Table 4.
  3. Use the plastic fork to quickly remove any unmelted ice from the cup.
  4. Carefully pour the cold water from the cup into the graduated cylinder and record the final volume to the nearest mL.

Table 4: Enthalpy Data
Volume of Hot Water / ______mL
Initial Temperature of Hot Water / ______°C
Final Temperature of Water and Melted Ice / ______°C
Final Volume of Water and Melted Ice / ______mL
  1. Calculate the change in the temperature of the hot water.
  2. Calculate the heat lost by the hot water. Heat must be expressed in kilojoules (kJ). It can be calculated by using the following equation:

Q= (4.18 J/g °C) x (1 kJ/1000 J) x mass of water (g) x ΔT

  1. Calculate the volume of ice melted.
  2. Calculate the mass of ice melted. Remember that the density of water is 1 g/mL.
  3. Calculate the molar heat of fusion of ice, i.e., the number of kilojoules of heat per mole ice.
  4. Calculate the percent error in your determination of the value for the molar heat of fusion of ice. The formula for Percent Error = |Experimental - Actual|/ (Actual) x 100 = % Error
  5. In order to do the calculations, you assumed that all the heat lost by the hot water was absorbed by the ice, causing it to melt. Was this assumption correct? Explain.
  6. Write an equation for the melting ice. Include the energy term in kJ on the proper side of the equation.