Materials Included In Kit

Additional Materials Required

Safety Precautions

This activity requires the use of hazardous components and/or has the potential for hazardous reactions. Please review current Safety Data Sheets for additional safety, handling and disposal information.

Teacher Tips

• The zipper-lock bags are heavy-duty, 6 mil bags and should be reusable for many experiments. To extend the life of the bags, never poke the bag with sharp objects. Inspect the bags before each experiment to check for tears, holes, or weak seams. Due to the heat and caustic materials in The Hot Bag, make sure the bags are in good shape before performing the reaction.
• All three reactions can be performed during one 45-minute lab period.
• There are enough chemicals to perform each experiment 60–80 times. Extra bags have been provided for large class sizes or if they break. Heavy-duty zipper-lock bags can also be purchased at local stores.

  Instant Cold Pack

• This reaction is very simple and thus makes a good introductory lab activity.
• The amount of ammonium nitrate and water used is not critical. A 2:1 weight ratio of water to ammonium nitrate seems to work best. Using about 3 g of ammonium nitrate and 5 mL of water will result in a final temperature around 0 °C. Tap water works fine.
• If the bag gets too cold, simply place the bag on the counter or in a sink and allow it to warm up.
• Disposal: Please consult your current Flinn Scientific Catalog/Reference Manual for general guidelines and specific procedures governing the disposal of laboratory waste. The ammonium nitrate solution may be flushed down the drain with excess water according to Flinn Suggested Disposal Method #26b. Check with local authorities before disposing of any laboratory chemicals.

  Cool Reaction

• The cool reaction takes 20–30 seconds to initiate but within about 2 minutes the two solids become a liquid and reach a temperature below 0 °C.
• Temperatures as low as –5 to –8 °C can be reached. Please warn the students not to squeeze the bag or handle the bag too long after it becomes cold. These temperatures are cold enough to freeze skin.
• Ammonia gas is produced in this reaction. Therefore, ammonia fumes will be present during this experiment. If ammonia fumes become too strong, increase ventilation of the room or evacuate the room. Take the following steps to reduce ammonia fumes:
  1. Make sure the room has good ventilation.
  2. Make sure the bags are well sealed before starting the reaction.
  3. Have a disposal container inside an operating fume hood. Tell students not to open the bag until it is inside the fume hood.
  4. Use only the amount of chemicals prescribed in the laboratory procedure.
  5. Have the class perform the experiments in different order. Therefore, the whole class will not perform this reaction at the same time. This will spread the ammonia vapor out over the entire class period and hopefully it can be handled by the ventilation system.
• Disposal: Please consult your current Flinn Scientific Catalog/Reference Manual for general guidelines and specific procedures governing the disposal of laboratory waste. Use Flinn Suggested Disposal Method #26c for the products of this reaction.

  The Hot Bag

• The Hot Bag is the most difficult and potentially dangerous of the three reactions to perform; therefore, students must closely follow all safety procedures. Demonstrating the proper technique and possible problems to your students is recommended. This reaction can be safely performed if the instructions are followed but understanding safety issues is important.
• The Hot Bag reaction is sometimes difficult to get started. Use only one lump of CaO and 1 mL of water to start the reaction. The generated heat and additional water will complete the reaction. If the lump has a coating of Ca(OH)₂ or if too much water is added initially, the reaction will not proceed. If the lump of CaO is crushed and does not react with the water (i.e., no heat is generated), simply dispose of the mixture, rinse the bag out, and try again.
• It is important to only use 0.85 to 1.5 g of CaO in this reaction. Too much CaO will generate excessive heat which will weaken the bag and lead to bag rupture. Most of the lumps are between 0.85 and 1.1 g and these are perfect for the reaction. With 0.95 g of CaO, the temperature will rise to over 60 °C! If the bag gets too hot to handle, simply place the bag in the sink and run cold water over it.
• Do not tightly seal the bag. The bag will expand due to the generation of hot air and steam. If the bag begins to inflate, “unzip” the bag and release the pressure. If the student cannot “unzip” the bag, simply place the bag in the sink and run cold water over it.
• Do not crush the CaO with any sharp or hard object. The CaO lumps are very hard and in the process of crushing them to initiate the reaction, the lump can puncture the wall of the bag. Using a pen, fingernail, or any other sharp or hard object may lead to a puncture or weakening of the wall of the bag. Use only the soft part of fingers to break apart the lumps.
• As soon as the reaction begins and heat is generated, immediately add the remaining water to prevent the reaction from overheating.
• Disposal: Please consult your current Flinn Scientific Catalog/Reference Manual for general guidelines and specific procedures governing the disposal of laboratory waste. The calcium oxide suspension should be diluted with excess water, neutralized with 6 Molar hydrochloric acid, and then flushed down the drain with excess water accordign to Flinn Suggested Disposal Method #10. Check with local authorities before disposing of any laboratory chemicals.

Answers to Questions

Instant Cold Pack

1. Record your general observations.

   Student answers may vary.

2. Is this reaction exothermic or endothermic? Why?

   The process is endothermic because it gets cold. It is taking heat from its surroundings.

3. How would the temperature of the bag be affected if twice as much water is added?

   The same amount of energy will be produced because it is based on the system—ammonium nitrate. However, if excess water is added, it becomes part of the surroundings and absorbs the energy. Therefore, the more water present, the temperature drop will be less because the energy stays constant, but the mass of the immediate surroundings increases.

Cool Reaction

1. Record your general observations.

   Student answers may vary.

2. Is this reaction exothermic or endothermic? Why?

   The reaction is endothermic because it gets cold. It is taking heat from its surroundings.

3. Where does the water come from to create the final solution?

   The water comes from two sources. Barium hydroxide is an octahydrate so eight waters of hydration are released as it goes into solution. Two additional water molecules are generated during the reaction of the hydroxide (OH^–) and proton (H⁺) during the reaction.

The Hot Bag

1. Record your general observations.

   Student answers may vary.

2. Is this reaction exothermic or endothermic? Why?

   The reaction is exothermic because it gets warm. The reaction is releasing heat to its surroundings.

3. How would a coating of calcium hydroxide affect the reaction?

   Calcium hydroxide has limited solubility in water. If there is a coating of calcium hydroxide on the calcium oxide, it will not dissolve off and will form a protective coating on the oxide. The water will not be in contact with calcium oxide and the reaction will not proceed.

Introduction

Feeling the heat generated from an exothermic process or the loss of heat during an endothermic process will make it easier to understand the concepts of heat, temperature and enthalpy. All three activities are performed in a heavy-duty zipper-lock bag and demonstrate a different type of thermodynamic process or reaction. 

Instant Cold Pack 
This activity demonstrates an endothermic process. The Instant Cold Pack is not a chemical reaction but merely the physical act of dissolving a solute in a solvent. When ammonium nitrate is dissolved in water, heat is required to break apart the forces holding the ammonium nitrate together as a crystal. The process is endothermic and is represented by the following chemical and enthalpy equations.

NH₄NO₃(s) + heat → NH₄NO₃(aq)
ΔH = ΔH_f(products) – ΔH_f(reactants)
ΔH = ΔH_f(NH₄NO₃(aq)) – ΔH_f(NH₄NO₃(s))
ΔH = –339.9 kJ/mole – (–365.6 kJ/mole) = 25.7 kJ/mole

Cool Reaction
This activity is an example of an endothermic reaction. Many reactions produce heat since products usually have less energy than the reactants in a spontaneous reaction. However, there are many examples of endothermic reactions. One of the more striking examples of a reaction that requires heat is when solid barium hydroxide and ammonium thiocyanate are mixed together. The resulting products, a solution of barium thiocyanate and ammonia in water, contain more energy than the reactants. The solution will be very cold. The driving force for this reaction is not the increase in enthalpy but rather the increase in entropy (or randomness). The overall free energy change (ΔG) for the reaction is negative so the reaction is spontaneous. Thermodynamic data is not available for barium thiocyanate or ammonium thiocyanate so the ΔH for the reaction cannot be calculated. However, it will be obvious by the observations made that ΔH is positive.

Ba(OH)₂•8H₂O(s) + 2NH₄SCN(s) + heat → Ba(SCN)₂(aq) + 2NH₃(aq) + 10H₂O(l)

The Hot Bag
This activity is an exothermic reaction. Calcium oxide is also known as lime or quicklime and is used to make plaster, mortar, bricks and many other construction materials. Calcium oxide is produced by heating limestone (calcium carbonate) in air. Calcium oxide readily absorbs and reacts with carbon dioxide and water to form calcium carbonate (CaCO₃) and calcium hydroxide (Ca(OH)₂) respectively. When calcium oxide is added to water, a large amount of heat and calcium hydroxide are produced. Calcium hydroxide is used to treat acidic soils, soften water, and in the preparation of many building materials such as plaster, mortar, and bricks. The reaction of water and calcium oxide produces heat and is an exothermic reaction. The solubility of calcium hydroxide is very low, about 1.6 g/L so the product is Ca(OH)₂(s), not Ca(OH)2(aq).

CaO(s) + H₂O(l) → Ca(OH)₂(s) + heat
ΔH = ΔH_f(products) – ΔH_f(reactants)
ΔH = ΔH_f[Ca(OH)₂(aq)] – [ΔH_f[CaO(s)] + ΔH_f[H₂O(l)]]
ΔH = –986.1 kJ/mole – [–635.1 kJ/mole + –285.8 kJ/mole] = –65.2 kJ/mole

Concepts

• Thermodynamics
• Enthalpy
• Endothermic/exothermic reactions

Background

Thermodynamics is the study of energy changes in a system. In chemistry, thermodynamics is usually the study of heat transfer and work that accompanies chemical reactions or changes of state. Enthalpy is a measure of the heat content of a system. The heat transfer into or out of a system at constant pressure is equal to the enthalpy change and is symbolized by ΔH. The enthalpy change is simply the difference in the enthalpy of the starting materials (usually the reactants) and the ending materials (usually the products). The enthalpy of a system will almost always change any time a reaction occurs or when the physical state of a material is changes (e.g., a solid melting or solute dissolving in a solvent). When looking at the enthalpy change of a system, it is extremely important to understand the difference between the system and the surroundings. Typically, the system is only the reactants and products of the reaction. The rest of the universe (solvent, containers, atmosphere above the reaction, or students) are part of the surroundings and are not included in any enthalpy calculations for the system.

ΔH = q_p — where q_p is the heat transferred to or from a system at constant pressure
ΔH = ΔH_f(products) – ΔH_f(reactants) for a reaction — where ΔH_f is the enthalpy of formation
ΔH = ΔH_f(final) – ΔH_f(initial) — for a process

Exothermic Reactions

When a system releases heat to the surrounding environment during a reaction, the temperature of the surroundings increases and the reaction vessel, which is part of the surroundings, will feel warm. For this type of reaction, the enthalpy of the products is lower than the enthalpy of the reactants and heat is one of the products. This is called an exothermic reaction (exo = “out of”). An example of an exothermic reaction is the reaction of oxygen with carbon compounds to produce carbon dioxide. This type of reaction, called combustion, produces energy in the form of heat and flames. Even though many reactions are spontaneous, just being exothermic does not imply that a reaction is spontaneous. To determine spontaneity, the change in entropy must be considered. 

CH₄(g) + 2O₂(g) → CO₂(g) + 2H₂O(l) + heat
ΔH = ΔH_f(products) – ΔH_f(reactants)
ΔH = ΔH_f[CO₂(g)] + 2ΔH_f[H₂O(l)] – [ΔH_f[CH₄(g)] + ΔH_f[O₂(g)]]
ΔH = –393 kJ/mole + 2(–286) kJ/mole – [–75 kJ/mole + 0] = –890 kJ/mole

A reaction pathway (see Figure 1) may be used to graphically show the relative enthalpies of the reactants and products. If the graph shows the products at a lower enthalpy than the reactants, then the reaction is exothermic. The reaction pathway graph will also show the activation energy or the intitial input of energy required for the reaction to proceed, but this is a topic for further discussions in kinetics. Endothermic Reactions

Sometimes a reaction or process requires heat to proceed and the system will take heat from the surroundings. Since heat is being removed from the surroundings, the reaction vessel will feel cool. This means that the enthalpy of the products is higher than the enthalpy of the reactants and heat is one of the reactants. This is called an endothermic reaction (endo = “into”). A common example of an endothermic process is the melting of ice. Solid water (ice) needs heat energy to help break apart the forces holding it together as a solid. The heat flowing into the ice will cool the surroundings. Another example of an endothermic reaction is the reaction between nitrogen and oxygen gas to form nitric oxide. This reaction does not proceed spontaneously at room temperature. However, just because a reaction is endothermic does not imply that it must also be non-spontaneous. To determine spontaneity, the change in enthalpy must also be considered. 

N₂(g) + O₂(g) + heat → 2NO(g)
ΔH = ΔH_f(products) – ΔH_f(reactants)
ΔH = 2ΔH_f(NO(g)) – ΔH_f(N₂(g)) + ΔH_f(O₂(g))
ΔH = 2(90 kJ/mol) – 0 + 0 = 180 kJ/mol

The reaction pathway for an endothermic reaction (see Figure 2) will show the products at a higher enthalpy than the reactants.

Experiment Overview

In this activity, three reactions or processes that undergo dramatic enthalpy changes are performed inside heavy-duty, zipper-lock bags.

Instant Cold Pack: Dissolve ammonium nitrate in water. This is an endothermic process and creates a cold pack similar to those used commercially by sports trainers or first aid personnel.

Cool Reaction: React barium hydroxide with ammonium thiocyanate in the solid state. This reaction is an excellent example of how endothermic reactions can still proceed spontaneously if there are other driving forces (i.e., entropy).

The Hot Bag: Add water to calcium oxide to produce calcium hydroxide and heat.

Materials

Safety Precautions

Ammonium nitrate is a strong oxidizer and may explode if heated under confinement. Ammonium nitrate is slightly toxic by ingestion, LD₅₀ 2217 mg/kg, and may be a body tissue irritant. Avoid all body tissue contact. If the Instant Cold Pack bag begins to get too cold to handle, put it down immediately to avoid skin damage. Ammonium thiocyanate is moderately toxic by ingestion, LD₅₀ 750 mg/kg. Barium hydroxide is toxic by ingestion and a body tissue irritant. Avoid all body tissue contact. Ammonia is produced in this reaction and is toxic and irritating by inhalation; avoid inhaling the products. Keep the bag tightly sealed during the Cool Reaction. Do not continue to squeeze the bag if it becomes too cold to touch. Calcium oxide is a corrosive material and a severe body tissue irritant. Avoid all body tissue contact. Reaction of calcium oxide and water will produce large amounts of heat and skin burns are possible. Do not handle the bag if it becomes too hot and do not remove the suspension until it has cooled to room temperature. Do not tightly seal the bag—allow the hot air and water vapor to escape. Wear chemical splash goggles, chemical-resistant gloves and a chemical-resistant apron. Wash hands thoroughly with soap and water before leaving the laboratory. Please review current Safety Data Sheets for additional safety, handling and disposal information.

Procedure

Instant Cold Pack

1. Place one scoopful of ammonium nitrate into a zipper-lock bag. One scoop is about 3 g. Tilt the bag so all the material falls into one corner of the bag.
2. Add about 5 mL of tap water to the ammonium nitrate.
3. Close the zipper-lock bag.

   (Optional) If measuring the temperature change of the reaction, place a thermometer or temperature probe into the mixture and take an initial reading. Monitor the temperature and also record the minimum temperature of the reaction.

4. Gently squeeze the bag to dissolve all the ammonium nitrate in the water.
5. Observe what happens when the ammonium nitrate dissolves. Record observations.
6. After the reaction begins to warm up to room temperature, rinse out the bag with excess water and dispose of the ammonium nitrate according to your teacher’s directions.
7. Dry the inside of the zipper-lock bag using a paper towel before proceeding on to the next experiment.

Cool Reaction

1. Weigh out approximately 5 g of barium hydroxide and 2.5 g of ammonium thiocyanate.
2. Add the barium hydroxide and ammonium thiocyanate to the zipper-lock bag. Tilt the bag so all the solids are in one corner of the bag.
3. Close the zipper-lock bag since ammonia is produced during the reaction; make sure the bag is tightly sealed.

   (Optional) If measuring the temperature change of the reaction, place a thermometer or temperature probe into the mixture and take an initial reading. Monitor the temperature and also record the minimum temperature of the reaction.

4. Gently squeeze the bag to thoroughly mix the two solids.
5. Observe what happens when the solids mix. Record observations.
6. After the reaction begins to warm up to room temperature, rinse out the bag with excess water and dispose of the products into a waste beaker in a fume hood according to your teacher’s directions.
7. Dry the inside of the zipper-lock bag using a paper towel before proceeding on to the next experiment.

The Hot Bag

1. Place one or two lumps of calcium oxide in a zipper-lock bag. Do not use more than 1 g of calcium oxide; the reaction will generate too much heat for the bag to handle and the bag may melt. Tilt the bag so all the material falls into one corner of the bag.
2. Add about 1 mL of tap water to the calcium oxide.
3. Gentle squeeze the bag and try to break the calcium oxide lump apart. Do not use anything hard or sharp to break apart the calcium oxide—fingertips work fine. As soon the calcium oxide begins to heat up, add another 3–4 mL of tap water.
4. Loosely close the zipper-lock bag. Do not tightly seal the bag.

   (Optional) If measuring the temperature change of the reaction, place a thermometer or temperature probe into the mixture and take an initial reading. Monitor the temperature and also record the minimum temperature of the reaction.

5. Keep squeezing the bag to thoroughly mix the water and calcium oxide. The result should be a white suspension of calcium hydroxide.
6. Observe what happens when the calcium oxide reacts with the water. Record observations.
7. After the reaction begins to cool down to room temperature, rinse out the bag with excess water and dispose of the calcium hydroxide suspension according to your teacher’s directions.
8. Dry the inside of the zipper-lock bag using a paper towel before proceeding on to the next experiment.

Student Worksheet PDF

12928_Student1.pdf