Hot and Cold Packs: A Thermochemistry Activity
A discussion of chemical hot and cold packs can really warm up a classroom lesson on thermochemistry. In the following activity, students use a coffee cup calorimeter to measure the heat of solution of a chemical salt using 3 different masses. It’s recommended for safety—and for a more green chemistry experience—that students use ammonium chloride for the cold packs and calcium chloride for the hot packs. If chemicals are limited, consider having half the students work with 1 chemical and half with the other. After students have measured the heat of solution for the 3 masses, they graph their data (change in temperature vs. mass) and draw a best fit line. The best fit line can be used to determine what mass of chemical is needed to achieve a specific temperature. Students then design a hot and/or cold pack that utilizes 100 mL of water and can be activated when needed. Ensure that you review student designs before allowing them to perform the activity and that students understand and follow appropriate safety protocols.
Calorimetry is the science of measuring heat. Many chemical and physical transformations involve energy transfer in the form of heat. The magnitude and direction of heat may be determined using a calorimeter. In reactions that occur in aqueous solutions, the energy is transferred to or taken away from the water. A calorimeter is an apparatus that is insulated and prevents heat from flowing in or out of the system. Since the experiment is run under constant pressure (atmospheric), the change in water temperature that is measured is due to the enthalpy of reaction (heat of reaction). The heat of reaction may be calculated using the values measured for change in the water temperature.
The heat transfer, or change in enthalpy in a reaction (qrxn), is related to the mass of the solution (m), the specific heat capacity of the solution (c), and the temperature change (ΔΤ = Τfinal — Τinitial).
qrxn= – (m × c × ΔΤ)
The specific heat capacity of a substance is the amount of energy required to raise 1 g of the substance 1° C. The specific heat capacity of water is 4.186 J / (° C × g). In experiments conducted in aqueous solution, the specific heat capacity of water is generally used.
Instant Hot and Cold Packs
Many instant hot and cold packs function by dissolving a salt into water. As the salt disassociates, heat is either released in an exothermic reaction or absorbed in an endothermic reaction. Commercial instant cold packs typically use either ammonium nitrate or urea as their salt component; hot packs often use magnesium sulfate or calcium chloride. These reactions happen in a similar manner. When the salt is dissolved in water, the ionic bonds of the salt separate. This process requires energy, which is obtained from the surroundings. The ions then form bonds with the water, a process that releases energy. If more energy is released than taken in, then the process is exothermic, making the solution feel warmer. If more energy is taken in than released, then the process is endothermic, making the solution feel cooler.
Commercially, there are 2 other commonly sold types of instant hot packs. One heats up when exposed to air. This hot pack functions as iron reacts with oxygen to form iron (III) oxide, an exothermic reaction. The other type relies on the super cooling of sodium acetate. Upon heating the solution, it can become supersaturated. Without a seed crystal, the sodium acetate will remain in solution as it cools. This type of hot pack typically contains a metal disk that provides a site for crystallization when depressed. As the sodium acetate forms a regular crystal arrangement, heat is released. This hot pack is reusable as it can be regenerated in boiling water to once again form the supersaturated solution.
Use this activity only in accordance with established laboratory safety practices, including appropriate personal protective equipment (PPE) such as gloves, chemical splash goggles, and lab coats or aprons. Ensure that students understand and adhere to these practices. Know and follow all federal, state, and local regulations as well as school district guidelines for the disposal of laboratory wastes. Students should not eat, drink, or chew gum in the lab and should wash their hands before and after entering or exiting the lab.
Materials (per student group)
- Calcium Chloride
- Ammonium Chloride
- 3 Polystyrene Cups with Lids
- Beaker, 400 mL
- Beaker, 150 mL
- Graduated Cylinder, 100 mL
- Hot Plate
- Weigh Boats
- Various Materials for Making Thermal Packs (as specified in the student designs)
- Place 1 polystyrene cup inside another and then place both cups inside a 400-mL beaker. This is your calorimeter for measuring changes in temperature.
- Measure 100 mL of water with a graduated cylinder and pour it into the top polystyrene cup of the created calorimeter.
- Place the lid on the calorimeter, pulling back the tab to form an opening. Insert a thermometer into the calorimeter through the opening in the lid.
- Stir the water with the thermometer, monitoring the temperature until it is stable. Record this temperature (±0.1° C) as the initial temperature.
- Measure 5 g of the chemical salt. Record the exact value used.
- Remove the calorimeter lid, add the 5 g of chemical salt, and stir. Replace the lid and thermometer.
- Continue to stir and monitor the temperature for 2 minutes. Record the highest or lowest temperature obtained (±0.1° C) as the final temperature.
- Discard the solution as directed by your teacher and rinse the inner cup. Thoroughly dry the calorimeter apparatus before reusing.
- Repeat steps 1 to 8 twice more with 10 g and 15 g of chemical salt, respectively.
- Determine the change in temperature.
ΔT = Tf − Ti
- Graph the change in temperature of chemical salt compared to the mass of chemical salt. Draw a best fit line for the points.
- Optional data analysis: Calculate the heat absorbed by the water for each mass of the chemical salt and the heat absorbed by the calorimeter for each mass of chemical salt. The specific heat of water is 4.186 J/° C × g.
qω = – [cω × mω × ΔΤ]
- Calculate the enthalpy of the solution for each mass of chemical salt.
Design a portable, 1-time-use hot pack or cold pack for treating injuries. The pack must have 100 g of water separated from a solid chemical and be activated only when the user does something to the pack to mix the 2 components. Your job is to determine how many grams of the chemical are required to achieve the following temperatures: hot pack, 55° C (131° F); cold pack, 3° C (37° F).
The following is a guide for creating your thermal pack. Ensure that you understand how your teacher expects you to present your designs.
- Diagram your hot or cold pack. Include labels to indicate sizes and quantities of materials used.
- List all materials and quantities needed to create your thermal pack.
- Explain the steps that you will follow to build your thermal pack.
- Describe the safety precautions you will use when creating and testing the thermal pack.
- Explain how you will test your hot or cold pack. Consider how you will know if you were successful.
- Get approval from your instructor to construct your thermal pack.
- Construct your thermal pack. Record any modifications that you make to your design during the construction phase.
- Test your thermal pack. Record any data that you collect.
- Explain whether your design was successful and what changes or modifications you would need to make if conducting this activity again.
- Inquiries in Science®: Examining Thermochemistry Kit (item #251209)
- Carolina Investigations® for AP* Chemistry: Fundamentals of Calorimetry Kit (item #840592)
*AP is a registered trademark of the College Board, which was not involved in the production of, and does not endorse, this product.