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        Measuring Bond Energy of an Ionic Compound

        In this media-rich lesson, students investigate bond energy and the law of conservation of energy. They examine the chemistry behind instant cold packs by using a calorimeter to study the endothermic dissociation of ammonium chloride in water.

        Lesson Summary


        Many students are familiar with cold packs—plastic bags that when manipulated to break an internal compartment exhibit a sudden drop in temperature, making them useful as first aid for sports injuries. One of the chemical compounds used in this type of product is ammonium chloride. In this lesson, students study the endothermic dissociation of ammonium chloride in water and, by taking careful measurements of its heat of solution with a calorimeter, determine its bond energy in kJ/mole.


        • Be able to calibrate a calorimeter
        • Understand how calorimeters are used to measure energy transfers in endothermic reactions
        • Understand what is bond energy
        • Understand the concept of dissociation
        • Understand the difference between endothermic and exothermic reaction
        • Apply the law of conservation of energy to endothermic reactions
        • Learn how scientists measure bond energy, and how they minimize sources of error between their measurements and standard reference measurements

        Grade Level: 9-12

        Suggested Time

        • Four 45-minute class periods or two 1.5-hour lab periods
        • TIP: The lab activity may be adapted to be a classroom demonstration, or the calorimeters can be calibrated ahead of time, allowing more time for the dissociation experiment.

        Multimedia Resources


        Lesson Preparation

        • Preparing the Calorimeters PDF Document
        • 8-oz plastic thermos jar (one for each lab group)
        • Electric drill with 1/2 in (12 mm) bit
        • One-holed rubber stoppers (to fit hole made by drill)
        • Disposable stir rods


        Part II

        • Calibrating the Calorimeters PDF Document
        • Determining a Calorimeter’s Heat Capacity PDF Document
        • Cold pack, such as those used in sports first-aid kits
        • Celsius thermometers, capable of 0.1°C measurements in the range of 5°C-105°C (Digital thermometers are best, but students can estimate tenths of a degree on a liquid-filled glass thermometer if they hold their eye perpendicular to the scale. Students should practice measurements with their thermometers if they are not already able to obtain consistent results. Students should follow safety procedures with glass thermometers and other glassware.)
        • Laboratory scale (accurate to 0.1 g)
        • Stopwatches/timers (one for each group, or use a clock with a second hand)
        • 100 g aluminum cylinder
        • Laboratory tongs
        • Laboratory heater
        • Several 1 L Pyrex beakers (fill two with tap water, and leave out for several hours to equilibrate with room temperature)
        • 250 mL graduated cylinders
        • Safety goggles


        Part III

        Before the Lesson

        The Lesson

        Part I: Thinking About Chemical Reactions

        1. Begin by eliciting students' prior understanding about chemical bonds. Write the following questions on the board and ask students to answer them in their science notebooks:

        1. What is a force? Name as many different kinds of forces as you can.
        2. What is the relationship between force and energy?
        3. Give an example of how forces might be active on the microscopic scale of atoms.
        4. What are some ways forces might change how matter behaves at the atomic scale?

        2. Have students work through the Dissolving Salts in Water Flash Interactive in pairs or small groups. Ask students to discuss the questions on the board as they work through the activity. When they have finished, ask them if they would like to revise their examples of how forces might be active on the microscopic scale. In what ways might the arrangement of atoms in an ionic compound reflect these forces?

        3. Bring the class back together and review the terms "endothermic reaction" and "exothermic reaction." Ask students how they think they could determine if a reaction they are observing and measuring is endothermic or exothermic. What do they think they would need to know to be able to make this determination? Lead the class to the understanding that changes in temperature will help them determine if a reaction is endothermic (it absorbs energy and thus a temperature drop occurs) or exothermic (it releases energy in the form of heat and thus a temperature increase occurs).

        4. Show students the first-aid instant cold pack and ask them to describe it. What is it for? How is it used? Record their answers on the board, then choose a volunteer to demonstrate how the cold pack works. After the student activates the cold pack, pass it around the room so that everyone can feel the change in temperature. Then ask students to jot down in their notebooks whether they think an exothermic or endothermic reaction is occurring, and their initial ideas about the role energy plays in the cold pack becoming cold. When they are finished, have them compare their ideas in their small groups, and then discuss the following questions:

        1. Why do you think the pack became cold when the internal compartment inside the bag was broken? How is this different from other reactions (i.e., instant heat packs) that raise the temperature of the reactants and their surroundings?
        2. What do you think is happening at the molecular level? How would you apply the law of conservation of energy to the situation with the cold pack? (According to the law of conservation of energy, the total energy at the end of the experiment is equal to the energy at the beginning of the experiment.)
        3. Predict what you think you'd find if you could measure the total energy in the bag before and after the bag was activated.

        5. When they are finished, ask representatives of each group to share their ideas with the class. Some students will say that there was a chemical reaction, and that bonds were made and/or broken, and that these contributed to the energy equation. Explain that in the case of the cold pack, they observed the large-scale result of the breaking of billions of chemical bonds of a compound called ammonium chloride (NH4Cl) as it dissociates into ammonium and chloride ions when it dissolves in water.

        At this point, make sure students understand the connection between the breaking of the bonds and the temperature change. Remind them that if the temperature went down, then heat energy was required, and therefore the chemical reaction required energy, which it took from its environment. Since the environment is quite large, and heat can be transferred from it in many ways, measuring the temperatures of the cold pack before and after activation would allow them to determine heat of dissociation. But how could they accurately measure how much energy is required to break these bonds? Have students throw out some initial ideas before you introduce the calorimeter.

        Part II: Calibrating a Calorimeter

        6. Tell students that they will be measuring the heat absorbed or released in a chemical reaction using a calorimeter, a device that includes an insulated container, a solution that provides energy or absorbs energy in a reaction, and a thermometer to monitor the temperature. By knowing the mass of the calorimeter and the specific heat of the calorimeter (the solution and container combined)—all of which are constant values—you can determine the heat capacity of the calorimeter. (A bigger calorimeter will have a bigger heat capacity.) And knowing the heat capacity of the calorimeter will allow you to measure the change in temperature during a chemical reaction inside the calorimeter. This temperature change helps determine the amount of energy used to break chemical bonds in an endothermic reaction or the amount of energy released in an exothermic reaction.

        Explain that the first step will be to calibrate the calorimeters. Ask students why they think they need to calibrate the calorimeters.

        After students share some ideas, explain that there is always some loss of heat to or from the calorimeter itself during any reaction. Calibrating the calorimeter accounts for this heat loss so that the reactions that take place inside the calorimeter can still be measured with a level of accuracy. For example, if a certain amount of hot water were added to a known amount of cool water in a calorimeter, the water temperature inside the calorimeter would increase, but less than expected because of heat loss from the calorimeter. Students can calculate the specific heat (Ccalorimeter) of their calorimeters and use that factor to make more accurate measurements.

        7. Have students go to the lab to calibrate their calorimeters. The procedure they will use involves adding a known quantity of hot aluminum to a fixed amount of water in the calorimeter and then measuring the temperature change. Since the specific heat of the aluminum (0.9 J/g-°C at 300°K) and other metals is available from tables, students can calculate the missing energy when a hot cylinder of that metal is added to their calorimeters. Provide students with the necessary supplies and the Calibrating the Calorimeters PDF Document.

        [Note: It is important that the calorimeters are calibrated before students do the activities in Part III, otherwise the experiment will not return consistent results. If time is limited, you can calibrate the calorimeters ahead of time or do it as a teacher-led demonstration.]

        8. When students are finished recording their temperature measurements, have them use the Determining a Calorimeter’s Heat Capacity PDF Document to determine the heat capacity of the calorimeter. Note that a typical heat capacity of the type of calorimeter illustrated in the Preparing the Calorimeters PDF Document was determined to be 54 ± 7 J/°C. If all the calorimeters are made with the same materials, their heat capacities should be similar. It may be useful to take a class average of the values obtained, and use that figure in future calculations for consistency and ease of discussion.

        Part III: Using the Calorimeter to Measure Heat of Dissociation

        9. Tell students that they will now use their calibrated calorimeters to accurately measure the heat of dissociation (bonding energy) of 53.4 g (one mole) of the salt ammonium chloride. Ask students what is special about 53.4 g of that compound. Make the connection to the atomic weight of the compound by looking up the weights on the periodic table. Then provide students with all necessary supplies at their lab tables and distribute the Measuring the Heat of Disassociation of NH4Cl Activity PDF Document.

        10. After students have completed the activity, bring the class back together. Remind students that according to the law of conservation of energy, the total energy of the system before and after the reaction will be the same (leaving out relatively small factors like thermal conduction of the calorimeter and the difference between the C (temperature) of the salt solution and that of the water). Write the following equation on the board:

        • Total energy before = Total energy after
        • Total bond energy (in 53.4g NH4Cl)
          • = qsol + qcal
          • = (200 g x 1 cal/g-°C x (ΔT)) + (54 J/g-°C x (ΔT))
          • = ΔT x (200 J/g-°C + 54 J/g-°C)

        Help students understand that this means that when the ammonium chloride salt dissolves, the ammonium ion (+) becomes dissociated from the chloride ion (-). Because of the microscopic arrangement of the molecules, this requires energy.

        11. Break students up into their small groups, and ask them to predict what the curve of temperature over time would look like for the salt ammonium chloride. Then have students view the Dissolving Salts in Water Flash Interactive. Have them choose NH4Cl, and observe the curve of temperature over time after the salt is introduced. Encourage them to try different combinations of the amount of the salt and the water. How do these affect the curve?

        Part IV: Sharing Results

        12. Have the groups chart their data in a visible way (with a histogram, for example) to see if there is a pattern. Students should record their data on the board or on a piece of chart paper where the whole class can see it. There will most likely be some variation among the results. Use the class data to determine a class average. [Note: While there may be some disparities, the result should be near the experimentally determined standard of 14,780 J.]

        13. Remind students that they dissolved one mole of the compound, so that by dividing the total energy required in the reaction by Avogadro's number (the number of molecules in one mole of a substance), they will determine the individual bond energy of each molecule. Ask students to summarize how energy transfer in exothermic and endothermic reactions is related to breaking or creating chemical bonds in these reactions. What claims can they make about the energy released by using different quantities of the compound in the experiment? What about using the same quantity of a different compound, such as table salt? Would they predict that the post-reaction temperature would be the same, higher, or lower? If lower, would it be the same drop as with the ammonium chloride? Conclude by having students return to their computers and check these other compounds using the Dissolving Salts in Water Flash Interactive.

        Check for Understanding

        Have students discuss the following questions or record their answers in their notebooks:

        1. Write the chemical formula for the dissociation of ammonium chloride and then write the energy reaction underneath.
        2. How might you use the calorimeter to determine the exothermic heat of reaction of combustion? What would be the difference in the measurement of ΔT?
        3. What could you do to improve the accuracy and reliability of your experiment?
        4. Describe how the calorimeter offers an example of the law of conservation of energy.


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