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Cold Pack Chemistry: Where Does the Heat Go?

Difficulty
Time Required Average (6-10 days)
Prerequisites None
Material Availability A few specialty items, including instant cold packs containing ammonium nitrate, are needed.
Cost Average ($50 - $100)
Safety Adult supervision is required. Ammonium nitrate is potentially hazardous, wear gloves and safety goggles when working with it. Avoid contact with skin and eyes.

Abstract

Instant cold packs are popular with coaches and parents for treating minor bumps and bruises. The instant cold packs are not pre-cooled—you just squeeze the cold pack and its starts to get cold. So how does it work? In this chemistry science fair project, you will investigate the chemical reaction that occurs in instant cold packs.

Objective

The objective of this chemistry science fair project is to determine how the temperature of a mixture of water and ammonium nitrate changes with the amount of ammonium nitrate dissolved in the water.

Credits

David B. Whyte, PhD, Science Buddies

  • StyrofoamTM is a registered trademark of The Dow Chemical Company.

Cite This Page

MLA Style

Science Buddies Staff. "Cold Pack Chemistry: Where Does the Heat Go?" Science Buddies. Science Buddies, 5 Sep. 2013. Web. 1 Oct. 2014 <http://www.sciencebuddies.org/science-fair-projects/project_ideas/Chem_p081.shtml>

APA Style

Science Buddies Staff. (2013, September 5). Cold Pack Chemistry: Where Does the Heat Go?. Retrieved October 1, 2014 from http://www.sciencebuddies.org/science-fair-projects/project_ideas/Chem_p081.shtml

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Last edit date: 2013-09-05

Introduction

Have you ever used a hot pack to warm your hands or a cold pack on an injury? How can something produce heat or cold without any microwaving or refrigeration involved? The answer is: chemistry. Chemical reactions that produce heat are called exothermic. The burning of gasoline in a car engine is an example of an exothermic reaction.

Reactions that are accompanied by the absorption of heat are called endothermic. As an example of an endothermic reaction, when the chemical ammonium nitrate is dissolved in water, the resulting solution is colder than either of the starting materials. This kind of endothermic process is used in instant cold packs. These cold packs have a strong outer plastic layer that holds a bag of water and a chemical, or mixture of chemicals, that result in an endothermic reaction when dissolved in water. When the cold pack is squeezed, the inner bag of water breaks and the water mixes with the chemicals. The cold pack starts to cool as soon as the inner bag is broken, and stays cold for over an hour. Many instant cold packs contain ammonium nitrate. Ammonium nitrate is a white crystalline substance. When it is dissolved in water, it splits into positive ammonium ions and negative nitrate ions. In the process of dissolving the crystal, the water molecules "donate" some of their energy. As a result, the water cools down. How much heat energy is "lost" when ammonium nitrate dissolves in water? You can measure the amount of heat that is involved using Equation 1.

Equation 1:

q =   c m (T1 -T2)


  • q = energy, measured in joules (J)
  • c = heat capacity, measured in joules per gram per degree Celsius, J/(g°C)
  • m = mass of solution, measured in grams (g)
  • J = joules (J), unit of energy
  • g = grams (g) of water
  • °C = degrees Celsius
  • T1 = starting temperature, in degrees Celsius
  • T2 = lower temperature after ammonium nitrate has dissolved, in degrees Celsius

Equation 1 states that "the amount of heat energy that is lost when water changes from temperature T1 to the lower temperature, T2, equals the difference in the two temperatures, times the heat capacity, times the mass of the solution."

The heat capacity of a substance tells you how much the temperature will change for a given amount of energy exchanged. For water at 25°C, the heat capacity is 4.18 J/(g°C).

Equation 2:

c (water) =   4.18 J
(g°C)


  • c = heat capacity, measured in joules per gram per degrees Celsius, (J/g°C)
  • J = joules (J), unit of energy
  • g = grams (g) of water
  • °C = degrees Celsius

Equation 2 says that "the heat capacity of water is 4.18 joules per gram of water per degree Celsius." What this means is that if you add 4.18 J of heat energy to 1 g of water, its temperature will increase by 1.0°C. Substances other than water have different heat capacities.

In this chemistry science fair project, you will determine how the amount of ammonium nitrate that is dissolved in water affects the magnitude of the temperature change that occurs. In the procedure, you will dissolve different amounts of ammonium nitrate in water and measure the before and after temperatures, then graph the temperature change vs. grams of ammonium acetate added.

Terms and Concepts

  • Exothermic
  • Endothermic
  • Ammonium nitrate
  • Ion
  • Heat energy
  • Heat capacity
  • Joule (J)
  • Entropy

Questions

  • Based on your research, what are some examples of exothermic and endothermic reactions?
  • What is the chemical formula for ammonium nitrate dissolving in water?
  • How does the heat capacity of water compare with the heat capacity of other common materials, such as aluminum, glass, or air?
  • Based on your research, what is the definition of entropy? Does it increase or decrease when ammonium nitrate dissolves in water?

Bibliography

These more-advanced sites discuss the thermodynamics of the ammonium nitrate reaction. The W.W. Norton site has nice Flash tutorials that explain the basic chemistry.

Materials and Equipment

A project kit containing most of the items needed for this science project is available for puchase from AquaPhoenix Education. Alternatively, you can gather the materials yourself using this shopping list:

  • StyrofoamTM cups, 12-ounce (oz.) (5)
  • Permanent marker
  • Water, distilled
  • Graduated cylinder, 100mL or metric measuring cup, liquid, with 100mL mark
  • Latex or nitrile gloves
  • Safety goggles
  • Scissors
  • Instant cold packs containing ammonium nitrate and water (4). Instant cold packs are available at most sporting goods stores. The water is in a separate bag within the cold pack.
  • Plastic bowl, disposable
  • Weigh boats or wax paper, cut into 10-cm squares (15, more as needed); used as a surface for weighing out the ammonium nitrate crystals
  • Plastic spoons (5)
  • Newspaper, scrap to cover your work surface
  • Digital scale, accurate to at least 1 g, such as the Fast Weigh MS-500-BLK Digital Pocket Scale, 500 by 0.1 G, available from Amazon.com
  • Digital thermometer
  • Lab notebook
  • Digital timer
  • Helper
  • Graph paper

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Experimental Procedure

Performing the Experiment

  1. Label the Styrofoam cups with numbers 1 to 5.
  2. Add 100 mL of distilled water to each of the five cups.
  3. Cover the work surface with newspaper.
  4. Collect the ammonium nitrate from the instant cold pack, as follows:
    1. Put on your safety goggles and latex gloves.
    2. Shake the instant cold pack gently to move the water bag and crystals to the bottom.
    3. Cut the top of the bag off.
    4. Pour the ammonium nitrate crystals into the plastic bowl.
    5. Dispose of the water bag.


Cold pack ammonium nitrate
Figure 1. Ammonium nitrate from an instant cold pack.


    1. Place one of the wax paper squares on the scale.
    2. Zero the scale.
    3. Use a plastic spoon to add 10.0 g of ammonium nitrate on the square/weigh boat.
  1. Record the starting temperature of the water in cup #1 in a data table in your lab notebook.
    1. Point the digital thermometer at the water's surface and push the "on" button to get the temperature.
  2. Add 10.0 g of ammonium nitrate to the water in cup #1.
  3. Start the timer.
  4. Stir the contents with a plastic spoon.
  5. Record the temperature every 15 seconds (sec) until it stabilizes.
    1. You can record at longer intervals, such as 30 sec, if you choose.
    2. You might want a helper to write the times and temperatures down as you take the readings.
    3. Stir the contents of the cup gently between each reading.
    4. Remove the spoon when taking the temperature, as it may cause an error in the reading.
    5. Stop taking readings when the temperature stops decreasing.
  6. Dispose of the ammonium nitrate solution down the sink.
  7. Repeat steps 6–10, with new and clean materials and equipment, adding the following amounts of ammonium nitrate. Note: You may need to divide the ammonium nitrate onto two or more pieces of wax paper/weigh boats if it will not all fit on one piece. Be sure to record the starting temperature, the intermediate temperatures, and the final temperature for each sample.
    1. Cup #2: 20 g
    2. Cup #3: 30 g
    3. Cup #4: 40 g
    4. Cup #5: 50 g
  8. Repeat steps 1–12 two more times, so you have at least three trials. This ensures that your results are accurate and repeatable. Create a new data table for each trial. You can reuse the spoons and cups for the new trials, after rinsing them thoroughly with water.
  9. Dissolve any leftover ammonium nitrate in water and dispose of it in a sink.

Analyzing Your Results

  1. Look at your data tables. Subtract the ending temperature from the beginning temperature for each cup.
  2. Add the mass of ammonium nitrate to the data table for samples 1–5.
  3. Graph the data, with the grams of ammonium nitrate on the x-axis and the temperature difference on the y-axis.
  4. How does the final temperature change as more ammonium nitrate is added?
  5. Using Equation 1 from the Introduction, calculate the heat energy, q, in joules, that each sample lost.
    1. Use the heat capacity value from Equation 2 for "c."
    2. Use the total mass (water plus ammonium nitrate) for "m": 110 g for cup #1, 120 g for cup #2, etc.
    3. Use the temperature change from the table for T1 - T2.
  6. Graph the results, with the amount of ammonium nitrate in grams on the x-axis and the heat energy (q) on the y-axis.

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Variations

  • Use different amounts of ammonium nitrate. What happens when the solution becomes saturated (that is, no more ammonium nitrate can be dissolved)?
  • Graph the temperature vs. time for each sample. Look at the slope of the graphs. Do they all have the same initial rate of change, or does the initial rate of change increase as the amount of crystal increases?
  • Calculate q/gram for each sample. Graph the results.
  • What are some sources of error in the procedure? Devise a new procedure that reduces or eliminates the sources of error you identified.
  • Devise a procedure to determine how varying the starting temperature affects the size of the temperature change.
  • Compare ammonium nitrate with other chemicals, such as ammonium chloride, calcium chloride (exothermic), and sodium chloride.
  • Research the thermodynamics of the ammonium nitrate dissolution and calculate the enthalpy and entropy changes that occur.

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