Materials Included In Kit

Additional Materials Required

Safety Precautions

The cold pack solid is ammonium nitrate. It is slightly toxic by ingestion, harmful if swallowed, and a body tissue irritant. Ammonium nitrate in solid form is a strong oxidizer and may explode if heated or in contact with combustible materials. Avoid contact of the dry solid with organic compounds. Avoid contact of all chemicals with eyes and skin. Wear chemical splash goggles and chemical-resistant gloves. Please consult current Safety Data Sheets for additional safety, handling and disposal information. Remind students to wash hands thoroughly with soap and water before leaving the laboratory.

Disposal

Please consult your current Flinn Scientific Catalog/Reference Manual for general guidelines and specific procedures, and review all federal, state and local regulations that may apply, before proceeding. The cold pack solutions generated in this experiment may be rinsed down the drain with water according to Flinn Suggested Disposal Method #26b.

Lab Hints

• The laboratory work for this inquiry-based, “discovery” experiment can reasonably be completed in a typical 2-hour lab period. Obtaining information from the cold pack label takes about 10 minutes. The actual experimental measurements for the guided-inquiry are also very quick—three independent trials may be completed within 20 minutes.
• The amount of solid required depends on the number of runs or trials that each group does. A minimum of 2–3 runs is recommended to average the effects of random error.
• The most important element for success in an inquiry-based activity is student preparation. Sufficient time should be alotted for students to think through the measurements that must be made, how they will be made, the variables that will influence the measurements, and how the variables can be controlled, if necessary. The Prelaboratory Assignment provides leading questions and guidelines to stimulate class discussion. The questions may be used as the basis of a small-group activity prior to lab or assigned as homework in preparation for lab. Encourage students to work together to devise a procedure for the calorimetry experiment.
• Temperature measurements may be made using digital thermometers, glass-bulb thermometers, or computer-interfaced temperature sensors. Digital thermometers are preferred over glass thermometers because they provide direct readings, update every second, and have a precision of ± 0.1 °C. Glass thermometers are fragile and easily broken, especially if the solutions are vigorously stirred, as suggested in the Sample Procedure. In addition, the 1 °C divisions that are marked on most glass thermometers make them less precise (± 0.5 °C) than digital thermometers. Never allow students to use a glass thermometer as a stirring rod.
• The minimum temperatures recorded in the sample data were generally achieved within one minute after mixing and were usually stable for an additional minute. Both of these factors ensure that temperature is easily and precisely measured and that students will feel confident about their measurements.
• Two Styrofoam cups nested together provide better insulation and thermal stability than one cup. If two cups are used, students can easily run two trials without rinsing and drying the cup between trials. Simply have students interchange the actual solution cup and the bottom cup between measurements. In addition, we recommend that students nestle Styrofoam cups in a beaker for added stability when the thermometer is placed in the cup.
• The volume of liquid is a variable that must be controlled in the experiment, but there is no obvious or absolute value that should be used. The minimum volume of water needed to ensure that the thermometer is completely immersed in liquid is about 20–30 mL. If too large a liquid volume is used, the observed temperature change will be small and may be less reliable. Working with 5 g of cold pack solid and a water volume of 25–30 mL gave three significant figures in the measured temperature difference.
• One of the more stubborn student misconceptions is the idea that if the reaction mixture gets cold, it must have lost heat, therefore the reaction must be exothermic. This misconception may be traced to a lack of understanding of the system versus the surroundings. The temperature change that is measured in a typical coffee-cup calorimeter experiment is that of the surroundings. A heat of solution experiment is probably more confusing on this point than a heat of neutralization or heat of combustion experiment, because water is involved in the reaction. Also, using the combined mass of the solute and the solvent in the heat equation to calculate the heat change tends to blur the traditional distinction between the reactants and products versus the solvent.
• Assuming that the specific heat of the solution is equal to that of water (4.18 J/g • °C) is a source of error in the calculations and results. The specific heat of a solution is equal to the weighted average of the specific heats of the solute and solvent. Based on this definition, the specific heat of a solution containing 5 g NH₄NO₃ and 30 g H₂O is estimated to be 3.8 J/g • °C.
• Students may question why the total mass of the solution is used in the heat energy equation to calculate q. (See Prelab Question 1 and the Background section.) Using the total mass of the solution allows one to approximate the differential heat of solution as opposed to the integral heat of solution.
• Students may compare their experimental results with the literature value for the heat of solution of ammonium nitrate. Assuming a literature value of 25.7 kJ/mole (CRC Handbook of Chemistry and Physics, 82nd Edition, Section 5, pg. 105), the results in the Sample Data section (26.4 kJ/mole) give a calculated percent error of 3%.

Answers to Prelab Questions

1. Based on Equation 4 in the Background section, identify the information (data) needed to calculate the enthalpy change for a reaction. How will these data be obtained in this experiment?

   In order to calculate the enthalpy change for a reaction, data for the three terms involved in the heat energy equation (q = m x s x ΔT) must be known or measured. The mass (m) is the total mass of the solution after the cold pack solid has dissolved. The specific heat (s) is assumed to be the same as the specific heat capacity of water (4.18 J/g • °C). The temperature change (ΔT) is equal to the difference between the final and initial temperatures (T_final – T_initial).

2. Identify all possible variables that will influence the experimental data.

   Critical variables include: the mass of the solute (cold pack solid), the volume (mass) of the solvent, whether all of the solute dissolves in the solvent, the heat insulating properties of the reaction container, how well the reaction mixture is stirred and the stability of the initial temperature reading.

3. The independent variable in an experiment is the variable that is manipulated by the experimenter, while the dependent variable responds to (depends on) changes in the independent variable. Choose the dependent and independent variable for your guided-inquiry procedure to determine the heat of solution for a cold pack.

   In a calorimetry experiment, the mass of the solute in grams is the independent variable and will be varied in different trials. The temperature change that results depends on the mass of the solute and is thus the dependent variable in a calorimetry experiment.

4. What variables should be controlled (kept constant during the procedure)?

   The following variables should be held constant during the procedure: the volume (mass) of the solvent, the type of reaction container that is used (two insulating foam cups nestled one inside the other will provide better insulation than one cup) and continuous stirring of the reaction mixture.

5. 1. Discuss the factors that will affect the precision of the experimental results.

      Many factors will influence the precision of the results:

      • The precision of the balance used to measure the mass of solute.
      • The precision of the graduated cylinder used to measure the volume of solvent.
      • The precision of the thermometer used to measure the temperature of the reaction mixture.
      • The number of times the experiment is repeated to average the effects of random errors.
      • The type of vessel that is used as the calorimeter—how much heat is gained or lost by the calorimeter itself.
   2. Discuss the equipment and glassware that should be used to obtain precise measurements of the quantities that must be measured in a calorimetry experiment.

      The measurements should be made with the most precise glassware and equipment available in the lab—centigram balances (at least), appropriate size graduated cylinders, and digital thermometers.

   3. How can the precision due to random experimental error be improved?

      One important way to improve the precision of the experimental results is to average data obtained over several runs or trials. A minimum of 2–3 trials is recommended. Alternatively, class data may be averaged to eliminate outlying results.

Sample Data

Answers to Questions

Laboratory Report

What Is an Instant Cold Pack?
The following information was obtained for an AFAS-COLD™ Chemical Cold-Pack, manufactured by Afassco, Minden, Nevada.

Measuring the Heat of Solution Sample Procedure

1. Stack two insulating foam cups in a 400-mL beaker for stability.
2. Measure 30.0 mL of distilled water in a graduated cylinder and add it to the insulating foam cup. Record the volume of water used in the data table.
3. Place a digital thermometer in the water in the cup and allow the thermometer and the water to adjust to room temperature for at least 2–3 minutes.
4. Tare a weighing dish on the balance. Add about 5 g of cold pack solid to the weighing dish. Measure and record the precise mass of solid to the nearest 0.01 g.
5. Measure the initial temperature of the water in the cup to the nearest 0.1 °C and record the value in the data table.
6. Stir the water in the cup with a stirring rod and add the massed amount of cold pack solid to the water. Make sure that all of the solid dissolves in the water.
7. Stir constantly and monitor the temperature of the solution. Record the lowest temperature that is reached.
8. Rinse the cup contents down the drain with excess water.
9. Switch the inner and outer foam cups and repeat steps 1–8 at least once more. Record all data.

Post-Lab Questions

1. For each trial, calculate the heat energy change in joules when the cold pack solid dissolved in water. Recall: q = m x s x ΔT, and assume that s (specific heat) is equal to 4.18 J/g • °C. Discuss the meaning of the sign for q.

   Trial 1: q = 35.37 g x 4.18 J/g • °C x (–11.8 °C) = –1740 J
   Trial 2: q = 34.78 g x 4.18 J/g • °C x (–11.0 °C) = –1600 J
   Trial 3: q = 34.89 g x 4.18 J/g • °C x (–11.1 °C) = –1620 J
   The negative sign for the heat change indicates that the solution (the surroundings) lost heat energy to the solute (the system) when the solid dissolved in water. According to the law of conservation of energy, the amount of heat lost by the surroundings must be equal to the amount of energy gained by the system for the reaction to occur. The cold pack solid absorbed energy from the surroundings as it dissolved in water—it is an endothermic process. The heat change for an endothermic process is a positive quantity.

2. Calculate the energy change in joules per gram of solid for the cold pack solid dissolving in water. Note the sign of the energy change.

   Trial 1: ΔH_soln = 1740 J/5.37 g = 324 J/g
   Trial 2: ΔH_soln = 1600 J/4.78 g = 335 J/g
   Trial 3: ΔH_soln = 1620 J/4.89 g = 331 J/g
   Average value = (324 + 335 + 331)/3 = 330 ± 4 J/g
   The heat of solution of ammonium nitrate dissolving in water is a positive quantity, equal in magnitude but opposite in sign to the heat change calculated in Question 1.

3. Calculate the energy change in units of kilojoules per mole of solid for the cold pack solid dissolving in water. To do this:
   1. Convert the heat energy change found in Question 1 to kilojoules.

      Trial 1: 1740 J x (1 kJ/1000 J) = 1.74 kJ
      Trial 2: 1600 J x (1 kJ/1000 J) = 1.60 kJ
      Trial 3: 1620 J x (1 kJ/1000 J) = 1.62 kJ

   2. Convert the grams of solid used in the experiment to moles.

      Trial 1: 5.37 g x (1 mole/80.04 g) = 0.0671 mole
      Trial 2: 4.78 g x (1 mole/80.04 g) = 0.0597 mole
      Trial 3: 4.89 g x (1 mole/80.04 g) = 0.0611 mole

   3. Divide the energy change in kilojoules by the number of moles of solid to determine the energy change in units of kJ/mole. If more than one trial was performed, calculate the average value of the heat of solution also.

      Trial 1: 1.74 kJ/0.0671 mole = 25.9 kJ/mole
      Trial 2: 1.60 kJ/0.0597 mole = 26.8 kJ/mole
      Trial 3: 1.62 kJ/0.0611 mole = 26.5 kJ/mole
      Average value = (25.9 + 26.8 + 26.5)/3 = 26.4 ± 0.3 kJ/mole

4. Using the result from Question 3c and the information obtained for a commercial cold pack, calculate the kilojoules of heat energy transferred when the entire cold pack is activated. How cold will the instant cold pack’s solution become?

   Recall (Question 6): The number of moles of solid in the instant cold pack is 0.973 mole.
   Amount of heat transfer = 26.4 kJ/mole x 0.973 mole = 17.4 kJ.
   This amount of heat transfer should be sufficient to cool 63 mL of water containing 78 g of ammonium nitrate from 25 to –19 °C! That’s cold!
   Note: We did the experiment by combining the water and solid from a commercial cold pack and measuring the temperature change. It dropped from 23 °C to –9 °C within 2–3 minutes. Not all of the solid dissolved, however. The calculated molarity of the solution (if it all dissolved) is > 15 M.

5. Circle or highlight the correct choices in the following sentence to summarize the heat change that occurs when the commercial cold pack is activated.

   “When the white solid in the commercial cold pack dissolves in water, the pack feels cold because the temperature of the solution decreases. Energy is absorbed from the surroundings during this reaction and the reaction is classified as endothermic. The sign of ΔH for the heat of solution is positive.”

Introduction

Instant cold packs are familiar first aid devices used to treat injuries when ice is unavailable. Most commercial cold packs consist of a plastic package containing a white solid and an inner pouch of water. Firmly squeezing the pack causes the inner pouch to break. The solid then dissolves in the water producing a change in temperature. Can we measure the temperature change that occurs when the cold pack solid dissolves in water and determine the heat change for this process?

Concepts

• Thermochemistry
• Enthalpy change
• Calorimetry
• Heat of solution
• Exothermic vs. endothermic
• Heat and temperature

Background

Thermochemistry is the study of heat changes that accompany a physical process or a chemical reaction—heat may be either absorbed or released. Heat is defined as the energy transferred from one object to another due to a difference in temperature. We do not observe or measure heat directly; we measure the temperature change that accompanies heat transfer.

It is often not possible to measure the temperature of the reactants or products in isolation. The temperature change is determined for the surroundings. When a system of reactants and products releases heat to the surroundings, the temperature of the surroundings increases and the reaction container will feel warm to the touch. This is an exothermic reaction—the prefix exo- means “out of” and the root thermic means heat. Heat flows out of the system. An example of an exothermic reaction is the combustion of propane (C₃H₈) to produce carbon dioxide, water and heat. Equation 1 gives the chemical equation for this reaction.

When a system absorbs heat from the surroundings during a reaction, the temperature of the surroundings decreases and the reaction container will feel cold to the touch. This is an endothermic reaction, where the prefix endo- means “into.” Heat flows into the system. A familiar example of an endothermic process is the melting of ice. Solid water (ice) needs heat energy to break apart the forces holding the molecules together in the solid state. This physical change is represented by Equation 2. The amount of heat transferred in these processes depends on a difference in a quantity called the enthalpy, represented by the symbol H. The enthalpy change for a physical process or a chemical reaction is defined as the heat change that occurs at constant pressure. This is convenient, because most of the reactions that we carry out in the lab are in flasks or containers that are open to the atmosphere—they take place at a constant pressure equal to the external pressure.

Equation 3 shows the equality between the change in enthalpy (ΔH) of a system and the amount of heat transferred, symbolized by q_p, for a reaction carried out at constant pressure. The amount of heat (q_p) transferred to a substance or object depends on three factors: the mass (m) of the object, its specific heat (s) and the resulting temperature change (ΔT). See Equation 4. The specific heat (s) of a substance reflects its ability to absorb heat energy and is defined as the amount of heat needed to raise the temperature of one gram of material by one degree Celsius. The specific heat of water is 4.18 J/g • °C. The temperature change (ΔT) is equal to the difference between the final temperature and the initial temperature (ΔT = T_final – T_initial).

Based on the law of conservation of energy, the amount of heat released by the system must be equal to the amount of heat absorbed by the surroundings. The sign convention in Equation 5 reveals that the heat change occurs in the opposite direction. For an exothermic reaction, the heat released by the system results in a temperature increase for the surroundings (ΔT is positive) and the heat absorbed by the surroundings will be a positive quantity. The heat released by the system must have the reverse sign—it must be a negative quantity. According to this convention, the enthalpy change for an exothermic reaction is always a negative value. For an endothermic reaction, in contrast, the heat absorbed by the system results in a temperature decrease for the surroundings (ΔT is negative) and the heat released by the surroundings will be a negative quantity. The heat absorbed by the system must have the reverse sign—it must be a positive quantity. According to this convention, the enthalpy change for an endothermic reaction is always a positive value.

The energy or enthalpy change associated with the process of a solute dissolving in a solvent is called the heat of solution (ΔH_soln). In the case of an ionic compound dissolving in water, the overall energy change is the net result of two processes, the energy required to break the net attractive forces (ionic bonds) between the ions in the crystal lattice, and the energy released when the dissociated (free) ions form ion-dipole attractive forces with water molecules.

Heats of solution and other enthalpy changes are generally measured using calorimetry. A calorimeter is an insulated vessel that reduces or prevents heat loss to the atmosphere outside the reaction vessel. When using a calorimeter, the reagents being studied are mixed directly in the calorimeter and the temperature is recorded both before and after the reaction has occurred. If the mass of the material in the calorimeter is also known, then the amount of heat change occurring in the calorimeter may be calculated using Equation 4. For heat of solution calculations m is the total mass of the solution (solute plus solvent), s is the specific heat of the solution, and ΔT is the observed temperature change. The mass and temperature changes are measured experimentally and the specific heat of the solution is assumed to be the same as that of water, namely, 4.18 J/g • °C.

Experiment Overview

The purpose of this guided-inquiry experiment is to design and carry out a procedure to determine the heat of solution for a “cold-pack solid” dissolving in water.

Materials

Prelab Questions

1. Based on Equation 4 in the Background section, identify the information (data) needed to calculate the enthalpy change for a reaction. How will these data be obtained in this experiment?
2. Identify all possible variables that will influence the experimental data.
3. The independent variable in an experiment is the variable that is manipulated by the experimenter, while the dependent variable responds to (depends on) changes in the independent variable. Choose the dependent and independent variable for your guided-inquiry procedure to determine the heat of solution for a cold pack.
4. What variables should be controlled (kept constant during the procedure)?
5. a. Discuss the factors that will affect the precision of the experimental results. b. Discuss the equipment and glassware that should be used to obtain precise measurements of the quantities that must be measured in a calorimetry experiment. c. How can the precision due to random experimental error be improved?

Safety Precautions

Ammonium nitrate may be harmful if swallowed. It is a body tissue irritant. Avoid contact of all chemicals with eyes and skin. Wear chemical splash goggles and chemical-resistant gloves. Wash hands thoroughly with soap and water before leaving the laboratory.

Procedure

What Is an Instant Cold Pack?
Complete the following activity to become familiar with the materials in a commercial cold pack.

1. Read the label on a commercial cold pack and record the name of the solid used in the pack.
2. Read the warning information on the label and record any hazards associated with the product.
3. Using the known charges of ions, write the formula of the solid.
4. Calculate the molar mass of the solid.
5. Carefully cut open the cold pack without puncturing or activating the inner pouch. Measure the mass of the solid: Tare a large weighing dish on the balance. Transfer the cold pack solid to the tared weighing dish. Record the mass of the solid to the nearest 0.01 g.
6. Calculate the number of moles of solid in the pack.
7. Using a graduated cylinder, measure the volume of water contained in the inner pouch.

Measuring the Heat of Solution
Design and carry out a procedure to determine the enthalpy change (ΔH_soln) that occurs when the cold pack solid dissolves in water. Use a maximum of 5 grams of solid per measurement. Write out the procedure and construct a data table for the Laboratory Report that clearly shows the data that will be collected and the measurements that will be made. Discuss the procedure and data table with your instructor before beginning the experiment. Note: Review Questions 3–5 in the Prelaboratory Assignment.

Student Worksheet PDF

14034_Student1.pdf