AP Chemistry: Designing an Effective Hand Warmer Student Guide INTRODUCTION AP and the Advanced Placement Program are registered trademarks of the College Entrance Examination Board. The activity and materials in this kit were developed and prepared by Aldon Corporation, which bears sole responsibility for their contents. The majority of familiar chemical reactions occur at a constant pressure. Chemical reactions are all around. Many life-sustaining chemical reactions occurring in biological systems, the combining and mixing of substances during the manufacture of products, reactions performed in the chemical laboratory, etc. are performed at ambient, or atmospheric, pressure. During these reactions chemical bonds may be broken or chemical bonds may be formed. The reaction may require the input of energy or the reaction may release energy. Thermodynamics is the branch of science that examines/ analyzes the transformation of energy during physical and chemical processes. The energy change that occurs during a reaction is defined as enthalpy (H). At constant pressure enthalpy is a measure of the energy either gained from or lost to the surroundings in which the process is occurring. If the process requires the input of energy the change in enthalpy (ΔH) is positive and if the process releases energy the change in enthalpy is negative. Heat of Solution The breaking of bonds or attractions between chemical species during a reaction absorbs energy from the surroundings and the formation of bonds or attractions between chemical species releases energy to the surroundings. For example, when ionic salts are dissolved in water the ionic bonds between the ions (cations and anions) of the salt as well as hydrogen bonds between some of the water molecules are broken. In turn new attractions between the anions and some water molecules as well as the cations and some water molecules are formed. The amount of energy required to break these bonds varies depending on the properties of the particular ions from which the salt is formed. If the amount of energy required for the breaking of the bonds between the ions is greater than the amount of energy released during the formation of new attractions between the ions and water molecules the overall process of dissolving the salt results in the absorption of energy in the form of heat. Conversely if the amount
of energy required for the breaking of the bonds between the ions is lesser than the amount of energy released during the formation of new attractions between the ions and water molecules the overall process of dissolving the salt results in the release of energy in the form of heat. The change in enthalpy when a solute is dissolved in a solvent is referred to as the heat of solution (ΔH soln ). At constant pressure the heat of solution is equivalent to the heat flow (q) that occurs during the process, or: ΔH soln = q If the overall process absorbs heat energy it is an endothermic reaction and the heat flow is positive (q>0). If the overall process releases heat energy it is an exothermic reaction and the heat flow is negative (q<0). Due to the relationship between the heat flow and heat of solution (change in enthalpy) when the heat flow is positive the change in enthalpy will be positive and when the heat flow is negative the change in enthalpy will be negative. Calorimetry The examination and quantification of the heat flow that occurs during a process, such as the dissolution of a solute in a solvent, is usually measured using a calorimeter. A calorimeter is an insulated container that reduces atmospheric influence on the data being collected. For example in the case of dissolving an ionic salt in water the dissolution of the ionic salt may release heat to or absorb heat from the resulting aqueous solution. If heat is released some of the heat from the solution could also be quickly released to the atmosphere and if heat is absorbed it could cause the solution to also absorb some heat from the atmosphere. By performing the process in a calorimeter this possible atmospheric influence is minimized. While the calorimeter limits the effects of the surroundings on the measurements it does not eliminate them. They can still be affected by factors such as the insulating ability of the calorimeter or the material from which the calorimeter is composed for example. To compensate for any such issues the calorimeter is typically calibrated prior to use. According to the law of conservation of energy the total amount of energy in a closed system is constant. In the system energy can be converted, such as heat energy being converted to work energy, but overall the amount of energy must remain the same; energy can be neither created nor destroyed. If hot water is mixed with cold water, the hot water will transfer some of its heat energy to the cold water. There is heat flow, or change in enthalpy, when the two are combined. As energy must remain constant it stands to reason that the amount of heat energy lost by the hot water (q hot ) would be identical to the amount of heat energy gained by the cold water (q cold ) or: q hot + q cold = 0
Assuming the values of q hot and q cold are identical the only way the two could total zero is if they are opposite in sign. Hot water releases heat energy to cold water and the release of heat is an exothermic process. Remember, in exothermic processes the heat flow is negative so the value of q would be negative. The proceeding expression could be rewritten as: - q hot = q cold The change in enthalpy for a substance is directly related to the mass of the substance (m), the specific heat capacity of the substance (c), and the change in temperature change that occurs (ΔT). The relationship between these is expressed by the heat energy equation: q = mcδt Specific heat capacity is a constant specific to an individual substance. Water, for example, has a specific heat capacity of 4.184 J/g C. Prior to collecting data the calorimeter must be calibrated to determine the calorimeter constant. The calorimeter constant represents the effect of the calorimeter on heat flow. As previously noted q hot and q cold should be identical in value (and opposite in sign) so if two equal volumes of hot and cold water are mixed the resulting temperature should end up being exactly half way between the initial temperature of the hot water and the initial temperature of the cold water. While theoretically this should be the case, in reality it most likely will not be. Any difference between q hot and q cold is attributed to heat flow to the calorimeter (q cal ). With the effect of the calorimeter factored in and still observing the law of conservation of energy the overall heat flow in the system could be rewritten as: q hot + q cold + q cal = 0 or rearranged for the heat flow to the calorimeter: q cal = - (q hot ) - q cold q hot and q cold can each be calculated using the heat energy equation: q hot = m hot cδt hot q cold = m cold cδt cold and the value for each can be used to determine the enthalpy change of the calorimeter (q cal ).
The calorimeter constant (C cal ) is the heat energy absorbed by the calorimeter per degree of temperature change. Once q cal is determined the calorimeter constant can be calculated as: C cal = q cal ΔT cal where ΔT cal is the actual observed temperature change in the calorimeter after the hot and cold water have been mixed. Once the calorimeter constant has been determined, the heat flow effect to the calorimeter on further reactions performed in the calorimeter can be compensated for, such as during the previously discussed dissolution of ionic salts in water. Dissolving an ionic salt in water may be exothermic or may be endothermic. Heat may be transferred to or from the resulting solution (q soln ) as a result of the difference of thermal energy between the solution and the individual reactants (q rxn ) which in this case would be the solid ionic salt and water. Like the hot and cold water, the heat flow between the two should be same, just opposite in direction, or: - q soln = q rxn However the enthalpy of the calorimeter must also be taken into account so the overall thermal flow can be expressed as: q soln = - (q rxn + C cal ΔT) with the value of C cal being the previously determined calorimeter constant and q rxn calculated as: q rxn = mcδt where m is the mass of all reactants (solid salt and water). Once q soln is determined, the molar heat of solution, or overall change in enthalpy (ΔH soln ), can be calculated: ΔH soln = ( q soln ) X molar mass solute mass solute (g)
The Investigation Several commercial products take advantage of the exothermic and endothermic nature of some chemical reactions. One example is the hand warmer. There are several types of hand warmers commercially available that take advantage of different chemical reactions to provide a gentle heat for an extended period of time. One type of hand warmer contains a mixture of dry chemicals, including among other things powdered iron and salt. When exposed to air and moisture the iron oxidizes. The oxidation reaction is exothermic and gives off a gentle heat to warm hands. The salt is a catalyst and increases the amount of heat given off as oxidation of iron without a catalyst is a slow process. The powdered materials are sealed in a pouch that allows air to penetrate and initiate the reaction. Prior to use they are stored in air-tight packaging to prevent them from reacting until needed. This type of hand warmer usually produces a low, gentle heat for an extended period of time. Another type of hand warmer takes advantage of the exothermic nature of dissolving ionic salts in water. These hand warmers contain a small sealed plastic pouch inside another larger sealed pouch. The inner pouch contains a salt and the outer pouch contains water. The hand warmer is struck in a manner that ruptures the inner pouch, releasing the ionic salt into the water of the outer pouch. The salt dissolves and the water warms. This type of hand warmer tends to produce a more vigorous heat than the dry powder type of hand warmer but does not produce heat for quite as long. Several salts will be tested and the heat (if any) generated compared using calorimetry.
Pre-lab Discussion 1. When some ionic salts are dissolved in water the reaction is exothermic. When others are dissolved in water the reaction is endothermic. Why? 2. Why must the calorimeter be calibrated before use? What is the purpose of the calibration? 3. Suppose 50ml of 61.7 C water was added to a calorimeter containing 50ml of 21.3 C water. If there was no heat flow to the calorimeter what would the resulting temperature of the combined water be?
4. The above was performed in a calorimeter and in reality the resulting temperature of the combined water samples was 40.9 C. Calculate the calorimeter constant. The specific heat capacity of water is 4.18 J/g C and the density of the water is 1.0g/ml. 5. Suppose the above was performed in a different calorimeter that had a calorimeter constant of 27.8 J/ C. What should the resulting temperature of the combined water samples?
Procedure Materials Needed per Group Thermometer or temperature probe Graduated cylinder, 100ml 2 Polystyrene cups, 8oz Lid for 8oz polystyrene cups Magnetic stir plate/stir bar Shared Materials Sample sets #1-5 Hot and cold water Magnesium sulfate Distilled or deionized water Electronic balance Safety Gloves Safety goggles Apron
Part I: Determination of the Calorimeter Constant 1. Construct your calorimeter using two polystyrene cups. Place one cup inside of the other. Nesting one cup inside of the other increases the insulation of the calorimeter. 2. Obtain a plastic lid for the calorimeter. You will need to insert a thermometer or temperature probe through the lid to monitor temperature in the calorimeter during mixing. Depending on the exact thermometer or temperature probe you will be using, create a small hole in the lid, just large enough for the thermometer or probe to fit through. 3. Place the calorimeter on a stir plate and place a stir bar in the calorimeter. 4. Using a graduated cylinder, measure 50ml of cold water. Record the exact measurement of the cold water in the Data Analysis section of the lab. 5. Add the cold water to the calorimeter and turn on the stir plate at low speed to gently stir the cold water. 6. Place the lid on the calorimeter and insert the thermometer. Allow the temperature reading to stabilize and record the exact temperature of the cold water in the Data Analysis section. 7. Using a graduated cylinder, measure 50ml of hot water. Record the exact measurement of the hot water in the Data Analysis section. 8. Remove the thermometer from the calorimeter and record the exact temperature of the hot water in the Data Analysis section. 9. Quickly remove the lid from the calorimeter, add the hot water to the calorimeter, replace the lid, and insert the thermometer. 10. Observe the temperature reading on the thermometer. Continue to monitor the temperature until it reaches its highest reading and then begins to decrease. Record the highest reading reached in the Data Analysis section. 11. Calculate the calorimeter constant in the Data Analysis section. 12. Empty the water from the calorimeter and dry the inside of the calorimeter thoroughly before proceeding.
Part II: Heat of Solution of Magnesium Sulfate 1. Place the calorimeter on a stir plate and place a stir bar in the calorimeter. 2. Using a graduated cylinder, measure 45ml of distilled or deionized water. Record the exact measurement of the water in the Data Analysis section of the lab. 3. Add the water to the calorimeter and turn on the stir plate at low speed to gently stir the water. Place the lid on the calorimeter and insert the thermometer. 4. Using an electronic balance, measure 5g of magnesium sulfate (MgSO 4 ). Record the exact mass of the magnesium sulfate, to the resolution of the balance, in the Data Analysis section. 5. Check the thermometer in the calorimeter and record the exact temperature of the water in the Data Analysis section. 6. Remove the lid and thermometer from the calorimeter, add the magnesium sulfate to the calorimeter, and replace the lid and thermometer. 7. Observe the temperature reading on the thermometer. The temperature may increase or decrease. Continue to monitor the temperature until it reaches its highest or lowest reading. Record the highest or lowest reading reached in the Data Analysis section. 8. Empty the magnesium sulfate solution from the calorimeter. Rinse the inside of the calorimeter several times. Dry the inside of the calorimeter thoroughly before proceeding. 9. Repeat steps #1-8 being sure to record all measurements in the Data Analysis section for the second magnesium sulfate sample. 10. Upon completion of the second trial, empty the magnesium sulfate solution from the calorimeter. Rinse the inside of the calorimeter several times. Dry the inside of the calorimeter thoroughly before proceeding.
Part III: Effectiveness of Materials for Use in a Hand Warmer 1. Have your instructor assign you a sample set of three materials that could potentially be used in a hand warmer. Record the materials in your sample set in the Data Analysis section of the lab. 2. Place the calorimeter on a stir plate and place a stir bar in the calorimeter. 3. Using a graduated cylinder, measure 45ml of distilled or deionized water. Record the exact measurement of the water in the Data Analysis section of the lab. 4. Add the water to the calorimeter and turn on the stir plate at low speed to gently stir the water. Place the lid on the calorimeter and insert the thermometer. 5. Using an electronic balance, measure 5g of the first material in your sample set. Record the exact mass of the material, to the resolution of the balance, in the Data Analysis section. 6. Check the thermometer in the calorimeter and record the exact temperature of the water in the Data Analysis section. 7. Remove the lid and thermometer from the calorimeter, add the material to the calorimeter, and replace the lid and thermometer. 8. Observe the temperature reading on the thermometer. The temperature may increase or decrease. Continue to monitor the temperature until it reaches its highest or lowest reading. Record the highest or lowest reading reached in the Data Analysis section. 9. Empty the solution from the calorimeter. Rinse the inside of the calorimeter several times. Dry the inside of the calorimeter thoroughly before proceeding. 10. Repeat steps #2-9 for the remaining two materials in your sample set. Be sure to record all measurements in the Data Analysis section for the additional samples. 11. Clean up all materials according to your instructor. Be sure to wash your hands before leaving the lab.
Name: Innovating Science by Aldon Corporation Instructor: Date: Class/Lab Section: DATA ANALYSIS Part I: Determination of the Calorimeter Constant Density of Water 1g/ml Specific Heat Capacity of Water (c) - 4.184 J/g C Volume Cold Water Mass Cold Water Temperature Cold Water Volume Hot Water Mass Hot Water Temperature Hot Water Highest Temperature Reached Change in Temperature (ΔT) Temperature Change, Cold Water (ΔT cold ) Temperature Change, Hot water (ΔT hot ) q hot = m hot cδt hot = q cold = m cold cδt cold = q cal = -(q hot ) - q cold = C cal = q cal ΔT =
Name: Innovating Science by Aldon Corporation Instructor: Date: Class/Lab Section: DATA ANALYSIS Part II: Heat of Solution of Magnesium Sulfate Density of Water 1g/ml Specific Heat Capacity of Water (c) - 4.184 J/g C Molar mass MgSO 4 120.37 g/mol Trial #1 Trial #2 Volume Water Volume Water Mass Water Mass Water Mass MgSO 4 Mass MgSO 4 Initial Temperature Water Highest or Lowest Temperature Reached Change in Temperature (ΔT) Initial Temperature Water Highest or Lowest Temperature Reached Change in Temperature (ΔT) Average Change in Temperature (Trial #1 + Trial #2): 2 q rxn = mcδt = q soln = - (q rxn + C cal ΔT) = ΔH soln = ( q soln ) X molar mass solute = mass solute (g)
Name: Innovating Science by Aldon Corporation Instructor: Date: Class/Lab Section: DATA ANALYSIS Part III: Effectiveness of Materials for Use in a Hand Warmer Density of Water 1g/ml Specific Heat Capacity of Water (c) - 4.184 J/g C Molar mass NH 4 NO 3 80.04 g/mol Molar Mass Na 2 CO 3 106.00 g/mol Molar Mass CaCl 2 111.10 g/mol Molar Mass NaCl 58.45 g/mol Molar Mass LiCl 42.39 g/mol Molar Mass NaC 2 H 3 O 2 82.03 g/mol Material #1: Volume Water Mass Water Mass Solute Initial Temperature Water Highest or Lowest Temperature Reached Change in Temperature (ΔT) q rxn = mcδt = q soln = - (q rxn + C cal ΔT) = ΔH soln = ( q soln ) X molar mass solute = mass solute (g)
Name: Innovating Science by Aldon Corporation Instructor: Date: Class/Lab Section: DATA ANALYSIS Material #2: Volume Water Mass Water Mass Solute Initial Temperature Water Highest or Lowest Temperature Reached Change in Temperature (ΔT) q rxn = mcδt = q soln = - (q rxn + C cal ΔT) = ΔH soln = ( q soln ) X molar mass solute = mass solute (g)
Name: Innovating Science by Aldon Corporation Instructor: Date: Class/Lab Section: DATA ANALYSIS Material #3: Volume Water Mass Water Mass Solute Initial Temperature Water Highest or Lowest Temperature Reached Change in Temperature (ΔT) q rxn = mcδt = q soln = - (q rxn + C cal ΔT) = ΔH soln = ( q soln ) X molar mass solute = mass solute (g)
Name: Innovating Science by Aldon Corporation Instructor: Date: Class/Lab Section: DATA ANALYSIS Post-lab Discussion 1. Were the reactions in the three samples you tested exothermic or endothermic? 2. Which of the materials in your sample set do you think would be most effective for use in a commercial hand warmer? 3. Your group tested three different materials. Other groups may have tested some of the same or some different materials. Obtain results from other groups for materials your group did not test. Based on the results of the materials tested by other groups, would you change your answer to question #2?
Name: Innovating Science by Aldon Corporation Instructor: Date: Class/Lab Section: DATA ANALYSIS 4. Your instructor may have provided you with information about the commercial prices of the materials used in this investigation or your instructor may have asked you to research the cost of each material. Based on the resulting information, which material do you think would be best to use in a commercial hand warmer from a cost of manufacturing perspective? 5. Your instructor may have provided you Safety Data Sheets for each material used or asked you to research and find the information on your own. If you were to consider the manufacturing of a commercial hand warmer with regard to disposal/potential environmental impact of the product used which material would be the most suitable?.
Name: Innovating Science by Aldon Corporation Instructor: Date: Class/Lab Section: DATA ANALYSIS 6. While this investigation focused on chemical reactions that may be used in the production of a commercial hand warmer, there are also commercial cold packs available to consumers. Sometimes these cold packs are frozen in advance and stored cold and removed and used when needed. A different kind of cold pack, the instant cold pack, can be stored at room temperature and then when needed, can rapidly becoming cold. How do you think these types of cold packs work? Would any of the materials you tested potentially be suitable for use in an instant cold pack?