Materials Included In Kit

Additional Materials Required

Prelab Preparation

1. Measure the magnesium ribbon to the nearest 0.1 cm.
2. Weigh the strip in grams and divide this number by its length for the conversion factor for students. Record the number.
3. Cut the magnesium metal ribbon into 15 equal pieces, about 7 cm in length.

Safety Precautions

Hydrochloric acid is toxic by ingestion and inhalation and is corrosive to skin and eyes. Magnesium metal is a flammable solid. Keep away from flames. Do not handle magnesium metal with bare hands. Wear chemical splash goggles and chemical-resistant gloves and apron. Remind students to wash hands thoroughly with soap and water before leaving the lab. Consult current Safety Data Sheets for additional safety, handling and disposal information.

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 final solutions may be disposed of down the drain with excess water according to Flinn Suggested Disposal Method 26b.

Lab Hints

• Enough materials are included in this kit for 30 students working in pairs or for 15 groups of students. The small-scale calorimeters are reusable.
• The experimental work for this lab can reasonably be completed in one 50-minute period. Each trial should take no more than 10 minutes. If time is a concern, consider performing the experiment as a cooperative class activity, in which each group performs two trials of either Reaction A or Reaction B. Students calculate the heats of reaction for their two trials and record their results, along with those of the rest of the class, on the board or overhead projector. The class results for both Reactions A and B are averaged and the average heats of reaction are used to calculate the heat of reaction for Equation 1.
• The small-scale calorimeters used in this lab are available from Flinn Scientific, Inc. (Catalog No. AP5928). These calorimeters are manufactured from dense, 2"-thick polystyrene and lined with a specially formulated coating that permeates the pores of the foam. The lining maximizes heat efficiency and limits water absorption, which means that the calorimeter constant is small and does not change. The maximum temperature recommended for these small-scale calorimeters is 50 °C. Small-scale calorimeters work well with the microscale quantities used in this experiment.
• The thermometer may be used as the stirrer with small-scale calorimeters because the soft foam is not likely to damage any type of thermometer. Advise students, however, not to punch holes or indentations into the bottom or sides of the calorimeter. These indentations may trap liquid and thus interfere with both mixing the solution and drying the calorimeter for multiple trials.
• Because of the small-scale nature of this experiment, the maximum time required for any given trial, including measuring and weighing the reagents, is less than 10 minutes. This factor allows students to carry out several trials in a single class period. The maximum temperatures in this experiment are usually reached within one minute after mixing and are stable for at least 20–30 seconds. The temperatures then begin to decrease at a rate of about 0.1 °C every 20 seconds.
• The best thermometers for small-scale calorimetry are digital electronic thermometers (Flinn Scientific Catalog No. AP8716) or temperature sensors connected to a computer- or calculator-based interface system such as LabPro or CBL. Digital thermometers are reasonably inexpensive, update every second, and are precise to the nearest 0.1 °C. Temperature measurements may be a significant source of error in calorimetry experiments. Supplementary Information contains instructions for adapting this experiment to the use of technology for computer-interfaced data collection and analysis.
• For greater accuracy in the results, add a correction term to each enthalpy calculation to account for the heat loss to the calorimeter. See the Supplementary Information in the Extensions section.

Teacher Tips

• This experiment is designed as an advanced lab activity for students who have completed a basic, introductory calorimetry experiment, such as “Measuring Energy Changes,” “Discovering Instant Cold Packs” or “Measuring Calories” in the Flinn ChemTopic^™ Labs, Thermochemistry manual. For a review of the basics of calorimetry, see the Supplementary Information section.
• Consider leading into this experiment with a visual demonstration of the combustion of magnesium ribbon in a Bunsen burner flame. This simple demonstration will arouse interest and provides a counterbalance against the abstract nature of Hess’s Law. When performing the demonstration for the students, instruct students not to look directly or stare at the bright white flame. The bright light emitted by the burning magnesium is UV light and can damage the eyes. Students should look at the flame “out of the corners of their eyes,” using their peripheral vision.
• Using the total mass of the reaction mixture (hydrochloric acid solution plus magnesium or magnesium oxide) in the heat equation calculations (q = m x s x ΔT) may be confusing to some students. Students may argue that they are measuring the temperature increase in the surroundings, not the system, and thus they should not include the mass of the reactants and products. Using the combined mass of the reaction mixture is traditional in these types of calorimetry exercises and may compensate for the fact that the specific heat of the solution is assumed to be equal to 4.18 J/g•°C, the same as that of water.
• What makes Hess’s Law a law? This may be a good time to review with students the definition of a natural law. A law is not engraved in stone in nature—it is the expression of the results of many experiments repeated for many different systems. The “law” is a generalization that has been widely tested and has been found to be true for every reaction that has been tested. Hess’s Law is also known as the Law of Additivity of Reaction Heats.

Further Extensions

Supplementary Information

Calorimetry Basics
Calorimetry experiments are carried out by measuring the temperature change in water that is in contact with or surrounds the reactants and products. In a typical calorimetry experiment, the reaction of a known mass of reactant is carried out either directly in or surrounded by a known quantity of water and the temperature increase or decrease in the surrounding water is measured. The temperature change (ΔT) produced in the water is related to the amount of heat energy (q) absorbed or released by the reaction system according to the following equation:

q = m x s x ΔT

where m is the mass of the aqueous solution, s is the specific heat of water, and ΔT is the observed temperature change. The specific heat of water is defined as the amount of heat required to increase the temperature of one gram of water by 1 °C and is equal to 4.18 cal/g•°C. It is also represented by the symbol C.

Measuring the Calorimeter Constant of the Calorimeter
When equal volumes of hot and cold water are combined, the new temperature should be the average of the two initial temperatures if there is no heat loss to the calorimeter. In actual practice, the new temperature will be slightly less than the average due to heat loss to the calorimeter. Use the following procedure to determine the calorimeter constant in J/°C.

1. Label two calorimeters “cold water” and “warm water.” Mass each calorimeter to the nearest 0.01 g. Add 8 mL of cold tap water to the cold water calorimeter and find its mass to the nearest 0.01 g. Add 8 mL of warm tap water to the warm water calorimeter and find its mass to the nearest 0.01 g.
2. Stir the water in each calorimeter and record the initial temperature in each calorimeter to the nearest 0.1 °C (the initial temperature should be stable for at least 20 sec).
3. Pour the cold water into the warm water calorimeter and record the resulting final temperature of the mixture to the nearest 0.1 °C.
4. Rinse and dry the calorimeter and repeat at least once.
5. Calculate the mass of water in each calorimeter.
6. Subtract the initial temperature of the cold and warm water from the final temperature to determine the temperature change for both the cold and warm water (ΔT_cold and ΔT_warm).
7. Calculate the energy gained by the cold water: q_cold = m_cold x s x ΔT_cold.
8. Calculate the energy lost by the warm water: q_warm = m_warm × s × ΔT_warm.
9. Determine the absolute difference between the energy lost by the warm water and the energy gained by the cold water. Calculate the calorimeter constant in J/°C using the following equation:
   {13889_Extensions_Equation_3}
10. For each trial in Parts A and B, multiply the observed temperature change ΔT by the calorimeter constant to determine the heat absorbed by the calorimeter.
11. In calculating the heat released by the exothermic reactions in Parts A and B, add a term to correct for the heat absorbed by the calorimeter: heat released by reaction = – (heat absorbed by solution + heat absorbed by calorimeter)

Alternative Procedure (Computer-Interfaced Data Collection and Analysis)
The following instructions are provided for adapting the experiment to the use of technology (computer-interfaced data collection and analysis).

1. Connect the interface (e.g., LabPro, CBL system) to a computer or calculator.
2. Plug a temperature probe into the interface.
3. Open and set up a graph in your data collection software so that the y-axis reads temperature in degrees Celsius. Set the minimum and maximum temperature values at 20 degrees and 50 degrees, respectively.
4. The x-axis should be set for time in minutes. Set the minimum and maximum time values at 0 seconds and 240 seconds, respectively.
5. The time interval should be set so a temperature reading is taken every 10 seconds.
6. Obtain 15 mL of hydrochloric acid in a graduated cylinder and carefully transfer the acid to the calorimeter.
7. Place the temperature probe in the acid solution. Allow the probe to equilibrate at the initial temperature for 1 minute, then press Start to begin collecting temperature data.
8. Stir the solution with the temperature sensor and add the preweighed amount of magnesium ribbon or magnesium oxide in Parts A and B, respectively.
9. Continue stirring the solution and collecting temperature data for 4 minutes (240 seconds).
10. The system will automatically record data for the allotted time (240 seconds), then stop.
11. Print the computer-generated data table and graph, if possible, and use the data to complete the data table and the Post-Lab Calculations.

Answers to Prelab Questions

1. Review the Background section. Arrange Equations A, B and C in such a way that they add up to Equation 1.
   {13889_PreLabAnswers_Equation_A}
   {13889_PreLabAnswers_Equation_B}
   {13889_PreLabAnswers_Equation_C}
   {13889_PreLabAnswers_Equation_1}
2. Use Hess’s Law to express the heat of reaction for Equation 1 as the appropriate algebraic sum of the heats of reaction for Equations A–C.

   ΔHA – ΔHB + ΔHC = ΔH₁

3. The heat of reaction for Equation C is equal to the standard heat of formation of water. The heat of formation of a compound is defined as the enthalpy change for the preparation of one mole of a compound from its respective elements in their standard states at 25 °C. Chemical reference sources contain tables of heats of formation for many compounds. Look up the heat of formation of water in your textbook or in a chemical reference source, such as the CRC Handbook of Chemistry and Physics.

   The standard heat of formation of liquid water at 25 °C is equal to –285.8 kJ/mole at 25 °C. (CRC Handbook of Chemistry and Physics, 82nd Edition, 2001).

Sample Data

Results Table

§ The calculated enthalpy changes are all negative values. These are both exothermic reactions—the heat released by the system resulted in a temperature increase in the surroundings. The sign of the enthalpy change is a common source of student error.

Answers to Questions

Construct a results table to summarize the results of all calculations. For each reaction and trial, calculate the:

1. Mass of hydrochloric acid solution.

   Sample calculation for Reaction A, Trial 1: 16.97 g – 2.38 g = 14.59 g. See Sample Results Table for results of all other calculations.

2. Total mass of the reactants.

   Sample calculation for Reaction A, Trial 1: 14.59 g + 0.03 g = 14.62 g. See Sample Results Table for results of all other calculations.

3. Change in temperature, ΔT = Tfinal – Tinitial.

   Sample calculation for Reaction A, Trial 1: 30.0 °C – 21.2 °C = 8.8 °C. See Sample Results Table for results of all other calculations.

4. Heat (q) absorbed by the solution in the calorimeter. Note: q = m x s x ΔT, where s is the specific heat of the solution in J/g °C. Use the total mass of reactants for the mass (m) and assume the specific heat is the same as that of water, namely, 4.18 J/g °C.

   Sample calculation for Reaction A, Trial 1: q = 14.62 g x 4.18 J/g °C x 8.8 °C = 540 J. See Sample Results Table for results of all other calculations.

5. Number of moles of magnesium and magnesium oxide in Reactions A and B, respectively.

   Sample calculation for Reaction A, Trial 1: 0.028 g Mg x (1 mole/24.3 g) = 1.2 x 10–3 moles Mg. See Sample Results Table for results of all other calculations.

6. Enthalpy change for each reaction in units of kilojoules per mole (kJ/mole).

   Sample calculation for Reaction A, Trial 1: –540 J/(1.2 x 10–3 moles) x 1 kJ/1000-J = –450 kJ/mole. Notice the negative sign for the enthalpy change for the reaction. Reactions A and B are both exothermic reactions. The heat absorbed by the solution in the calorimeter is equal in magnitude but opposite in sign to the heat released by the reaction. See Sample Results Table for results of all other calculations.

7. Average enthalpy change (heat of reaction, ΔH_rxn) for Reactions A and B. Note: The enthalpy change is positive for an endothermic reaction, negative for an exothermic reaction.

   Sample calculation for Reaction A: [– 450 + (– 420)]/2 = – 440 kJ/mole. See Sample Results Table for results of all other calculations.

8. Use Hess’s Law to calculate the heat of reaction for Equation 1. Hint: See your answer to PreLab Question 2.

   ΔH₁ = ΔHA – ΔHB + ΔHC
   ΔH₁ = [–440 – (– 130) + (– 286)] kJ/mole = – 600 kJ/mole (rounded to two significant figures)

9. The heat of reaction for Equation 1 is equal to the heat of formation of solid magnesium oxide.
   1. Look up the heat of formation of magnesium oxide in your textbook or a chemical reference source.
   2. Calculate the percent error in your experimental determination of the heat of reaction for Equation 1.
      1. The standard heat of formation of solid magnesium oxide at 25 °C is equal to –601.6 kJ/mole. (CRC Handbook of Chemistry and Physics, 82nd Edition, 2001).
      2. {13889_Answers_Equation_2}

References

This activity is from Flinn ChemTopic™ Labs, Volume 10, Thermochemistry; Cesa, I., Ed; Flinn Scientific: Batavia, IL, 2002.

Introduction

The reaction of magnesium metal with air in a Bunsen burner flame provides a dazzling demonstration of a combustion reaction. Magnesium burns with an intense flame that produces a blinding white light. This reaction was utilized in the early days of photography as the source of “flash powder” and later in flashbulbs. It is still used today in flares and fireworks. How much heat is produced when magnesium burns?

Concepts

• Heat of reaction
• Heat of formation
• Hess’s law
• Calorimetry

Background

Magnesium reacts with oxygen in air to form magnesium oxide, according to Equation 1.

As mentioned above, a great deal of heat and light are produced—the temperature of the flame can reach as high as 2400 °C. The amount of heat energy produced in this reaction cannot be measured directly using standard methods in the high school lab. It is possible, however, to determine the amount of heat produced by an indirect method, using Hess’s Law.

The heat or enthalpy change for a chemical reaction is called the heat of reaction (ΔH_rxn). The enthalpy change—defined as the difference in enthalpy between the products and reactants—is equal to the amount of heat transferred at constant pressure and does not depend on how the transformation occurs. This definition of enthalpy makes it possible to determine the heats of reaction for reactions that cannot be measured directly. According to Hess’s Law, if the same overall reaction is achieved in a series of steps, rather than in one step, the enthalpy change for the overall reaction is equal to the sum of the enthalpy changes for each step in the reaction series. There are two basic rules for calculating the enthalpy change for a reaction using Hess’s Law.

• Equations can be “multiplied” by multiplying each stoichiometric coefficient in the balanced chemical equation by the same factor. The heat of reaction (ΔH) is proportional to the amount of reactant. Thus, if an equation is multiplied by a factor of two to increase the number of moles of product produced, then the heat of reaction must also be multiplied by a factor of two.
• Equations can be “subtracted” by reversing the reactants and products in the balanced chemical equation. The heat of reaction (ΔH) for the reverse reaction is equal in magnitude but opposite in sign to that of the forward reaction.

Consider the following three reactions: It is possible to express the combustion of magnesium (Equation 1) as an algebraic sum of Equations A, B and C. Applying Hess’s Law, therefore, it should also be possible to determine the heat of reaction for Equation 1 by combining the heats of reaction for Equations A, B and C in the same algebraic manner. Note: Chemical equations may be combined by addition, subtraction, multiplication, and division.

Experiment Overview

The purpose of this experiment is to use Hess’s Law to determine the heat of reaction for the combustion of magnesium (Equation 1). The heats of reaction for Equations A and B will be measured by calorimetry. The heats of reaction for these two reactions will then be combined algebraically with the heat of formation of water (Equation C) to calculate the heat of reaction for the combustion of magnesium.

Materials

Prelab Questions

1. Review the Background section. Arrange Equations A, B and C in such a way that they add up to Equation 1.
2. Use Hess’s Law to express the heat of reaction for Equation 1 as the appropriate algebraic sum of the heats of reaction for Equations A, B and C.
3. The heat of reaction for Equation C is equal to the standard heat of formation of water. The heat of formation of a compound is defined as the enthalpy change for the preparation of one mole of a compound from its respective elements in their standard states at 25 °C. Chemical reference sources contain tables of heats of formation for many compounds. Look up the heat of formation of water in your textbook or in a chemical reference source, such as the CRC Handbook of Chemistry and Physics.

Safety Precautions

Hydrochloric acid is toxic by ingestion and inhalation and is corrosive to skin and eyes. Magnesium metal is a flammable solid. Keep away from flames. Do not handle magnesium metal with bare hands. Wear chemical splash goggles and chemical-resistant gloves and apron. Wash hands thoroughly with soap and water before leaving the lab.

Procedure

Record all data for Parts A and B in the data table.

Part A. Reaction of Magnesium with Hydrochloric Acid

1. Obtain a 7-cm strip of magnesium ribbon and cut it into two pieces of unequal length, roughly 3- and 4-cm each. Note: Handle the magnesium ribbon using forceps.
2. Measure the exact length of each piece of magnesium ribbon to the nearest 0.1 cm.
3. Multiply the length of each piece of Mg ribbon by the conversion factor (g/cm) provided by your teacher to obtain the mass of each piece of Mg.
4. Mass a clean, dry calorimeter to the nearest 0.01 g.
5. Using a graduated cylinder, add 15 mL of 1 M hydrochloric acid to the calorimeter and measure the combined mass of the calorimeter and acid.
6. Using a digital thermometer or a temperature sensor, measure the initial temperature of the hydrochloric acid solution to the nearest 0.1 °C.
7. Add the first (shorter) piece of magnesium ribbon to the acid and stir the solution until the magnesium has dissolved and the temperature of the solution remains constant.
8. Record the final temperature of the solution to the nearest 0.1 °C.
9. Rinse the contents of the calorimeter down the drain with excess water.
10. Dry the calorimeter and mass it again to the nearest 0.01 g.
11. Repeat steps 5–9 using the second (larger) piece of magnesium ribbon.

Part B. Reaction of Magnesium Oxide with Hydrochloric Acid

12. Mass a clean, dry calorimeter to the nearest 0.01 g.
13. Using a graduated cylinder, add 15 mL of 1 M HCl to the calorimeter and measure the combined mass of the calorimeter and hydrochloric acid.
14. Tare a small weighing dish and add about 0.20 g of magnesium oxide. Measure the exact mass of magnesium oxide to the nearest 0.01 g.
15. Using a digital thermometer or a temperature sensor, measure the initial temperature of the hydrochloric acid solution to the nearest 0.1 °C.
16. Using a spatula, add the magnesium oxide to the acid. Stir the reaction mixture until the temperature remains constant for several five-second intervals. Record the final temperature of the solution to the nearest 0.1 °C.
17. Pour the reaction mixture down the drain with excess water. Rinse and dry the calorimeter.
18. Repeat steps 12–16 using a second sample of magnesium oxide.

 

Student Worksheet PDF

13889_Student1.pdf