Science Unit: 12 Lesson: 02 Suggested Duration: 7 days High School Science Unit 12 Exemplar Lesson 02: Heat Energy in Chemical Reactions This lesson is one approach to teaching the State Standards associated with this unit. Districts are encouraged to customize this lesson by supplementing with district-approved resources, materials, and activities to best meet the needs of learners. The duration for this lesson is only a recommendation, and districts may modify the time frame to meet students needs. To better understand how your district is implementing CSCOPE lessons, please contact your child s teacher. (For your convenience, please find linked TEA Commissioner s List of State Board of Education Approved Instructional Resources and Midcycle State Adopted Instructional Materials.) Lesson Synopsis This lesson introduces thermochemistry, the study of heat energy in chemical reactions. Students use calorimetry and the STAAR Reference Materials to investigate and calculate enthalpy of reactions that are endothermic and exothermic. Additionally, students study the heating curve for water and calculate the heat of fusion and heat of vaporization. TEKS The Texas Essential Knowledge and Skills (TEKS) listed below are the standards adopted by the State Board of Education, which are required by Texas law. Any standard that has a strike-through (e.g. sample phrase) indicates that portion of the standard is taught in a previous or subsequent unit. The TEKS are available on the Texas Education Agency website at http://www.tea.state.tx.us/index2.aspx?id=6148. Scientific Process TEKS C.11 Science concepts. The student understands the energy changes that occur in chemical reactions. The student is expected to: C.11C Use thermochemical equations to calculate energy changes that occur in chemical reactions and classify reactions as exothermic or endothermic. Readiness Standard C.11E Use calorimetry to calculate the heat of a chemical process. Supporting Standard C.2 Scientific processes. The student uses scientific methods to solve investigative questions. The student is expected to: C.2E Plan and implement investigative procedures, including asking questions, formulating testable hypotheses, and selecting equipment and technology, including graphing calculators, computers and probes, sufficient scientific glassware such as beakers, Erlenmeyer flasks, pipettes, graduated cylinders, volumetric flasks, safety goggles, and burettes, electronic balances, and an adequate supply of consumable chemicals. C.2F Collect data and make measurements with accuracy and precision. C.2I Communicate valid conclusions supported by the data through methods such as lab reports, labeled drawings, graphs, journals, summaries, oral reports, and technology-based reports. GETTING READY FOR INSTRUCTION Performance Indicators High School Science Unit 12 PI 02 Use calorimetry to measure an exothermic and an endothermic enthalpy of reaction. Write a summary report that includes your procedures, data collected, calculations, a discussion of the reactions, and an error analysis. Standard(s): C.2E, C.2F, C.2I, C.11C, C.11E ELPS ELPS.c.3D, ELPS.c.3E, ELPS.c.5B Key Understandings Chemical reactions are accompanied by energy transformations. What energy transformations occur in chemical reactions? Exothermic reactions release energy. How can energy transfer be observed in an exothermic reaction? Endothermic reactions absorb energy. How can energy transfer be observed in an endothermic reaction? Vocabulary of Instruction thermochemistry exothermic endothermic vaporization heat of fusion heat of vaporization heat of solution heating curve/phase diagram of water Last Updated 01/15/13 page 1 of 19
enthalpy fusion heat of reaction heat of combustion Science Unit: 12 Lesson: 02 Suggested Duration: 7 days Materials chemicals selected by students (see Advance Preparation, varies by group) crushed ice (for demonstration, per teacher) distilled or deionized water (approximately 300 ml per group) electronic balance (1 per group) foam cup calorimeters (and/or other preferred types, 1 per group) glue or tape (per group) graduated cylinder (100 ml, 1 per group) Handout: Enthalpy of Reaction (from previous activity) hot plate (for demonstration, 1 per teacher) hot water (per group) ice (per group) M citric acid solution (75 ml per group) M hydrochloric acid (50 ml per group) ring stands with clamps (for demonstration, 2 per teacher) safety goggles (1 pair per teacher) safety goggles (1 per student) sodium hydrogen carbonate (baking soda, 2.0 g per group) solution B (see Advance Preparation, for demonstration, 1.00 M citric acid in labeled flask per teacher) solution C (see Advance Preparation, for demonstration, 1.00 M calcium chloride in labeled flask per teacher) spatula or scoopula (for demonstration, 1 per teacher) temperature probes or thermometers (1 per group) temperature probes or thermometers (for demonstration, 2 per teacher) test tube rack (for demonstration, 1 per teacher) test tubes (for demonstration, 4 per teacher) thermometer or temperature probe (1 per group) water (for demonstration, per class) weighing paper (1 sheet per group) white solid A (sodium hydrogen carbonate (baking soda), for demonstration, 4 scoops per class) zinc metal pieces (2.0 g per group) Attachments All attachments associated with this lesson are referenced in the body of the lesson. Due to considerations for grading or student assessment, attachments that are connected with Performance Indicators or serve as answer keys are available in the district site and are not accessible on the public website. Handout: Heating Curve/Phase Changes of Water (1 per student and 1 for projection) Handout: Heat of Fusion of Water (1 per student) Handout: Enthalpy of Reaction (1 per student) Teacher Resource: Enthalpy of Reaction KEY Handout: Enthalpy Change Calculations (1 per student) Teacher Resource: Enthalpy Change Calculations KEY Resources Suggested Websites: Standard Enthalpies of Formation: http://chemed.chem.wisc.edu/chempaths/genchem-textbook/standard-enthalpies-of-formation- 551.html Heat of Formation: http://www.khanacademy.org/science/physics/thermodynamics/v/heat-of-formation Specific Heat, Heat of Fusion, and Vaporization: http://www.khanacademy.org/science/chemistry/states-of-matter/v/specific-heat--heat-offusion-and-vaporization Background Information: https://www.khanacademy.org/science/chemistry/reaction-rates/v/introduction-to-kinetics Instructor Background: http://www.youtube.com/user/chemflicks/videos?view=0 Energy Diagrams: http://www.emsb.qc.ca/laurenhill/science/exo.html http://misterguch.brinkster.net/energydiagram.html State Resource: Texas Education Agency, STAAR End-of-Course Success Training: Lesson: Heat Transfer District approved local resources including textbooks Last Updated 01/15/13 page 2 of 19
Advance Preparation Science Unit: 12 Lesson: 02 Suggested Duration: 7 days 1. Prior to Day 1, collect materials for the Phase Change and Heat of Fusion of Water investigations. Be sure to perform the demonstration prior to class and make adjustments as needed. Reuse the foam cup calorimeters from Lesson 01. Repair and/or replace as needed. Make arrangements to store ice. 2. Prior to Day 3, collect and organize materials for Demonstrating Endothermic and Exothermic Reactions. Materials needed include 4 scoops per class of white solid A: sodium hydrogen carbonate (baking soda), solution B: 1.00 M citric acid in labeled flask, and solution C: 1.00 M calcium chloride in labeled flask. Review all of the MSDS for chemical use and disposal. Be sure to perform the activity and make adjustments as needed. 3. Prior to Day 4, collect and organize materials for the Enthalpy of Reaction investigation. Prepare all solutions needed for the investigation. Review all of the MSDS for chemical use and disposal. Be sure to perform the activity to make adjustments as needed. 4. Prior to Day 6, preview the practice calculations to identify student obstacles. 5. Prior to Day 7, preview the Instructional Procedures for the Performance Indicator in order to determine the types of quantities of chemicals and equipment to be used. Some additional endothermic processes beyond those conducted in the lesson are: sodium hydrogen carbonate and hydrochloric acid ammonium nitrate and water (Heat of solution = 25.7 kj/mol ) Some additional exothermic processes beyond those conducted in the lesson are: reaction of Mg with hydrochloric acid calcium chloride and water (Heat of solution = -82.8 kj/mol ) Review all of the MSDS for chemical use and disposal. 6. Prepare attachment(s) as necessary. Background Information As discussed in Lesson 01 of this unit, in a chemical system, we think about energy first in terms of kinetic molecular theory. Moving particles (translation, rotation, and vibration) in gases, liquids, and solids have kinetic energy. Heat is a measure of the total amount of kinetic energy in the system. The temperature of a system is a measure of the average kinetic energy of the system. But, we also have to think about the stored or potential energy in a system. There is potential energy stored in the bonds within particles and in attractions among particles (such as hydrogen bonds, van der Waals forces, etc.). When bonds are broken and/or made and particles move farther apart or closer together, then potential energies change. The total energy content (kinetic and potential) of any system at constant pressure is called enthalpy, and it is symbolized by H. Enthalpy cannot be measured directly. Instead, scientists measure changes in enthalpy. enthalpy change in the system = enthalpy final enthalpy initial or ΔH = H f H i If an enthalpy change, ΔH, is negative, the total energy of the final system is less than that of the initial system. So, energy is released to the environment. If an enthalpy change, ΔH, is positive, the total energy of the final system is more than that of the initial system. So, energy is absorbed from the environment. An example of a simple enthalpy change is a beaker of water being heated on a hot plate. As the water absorbs heat from the hot plate, its molecules move faster and faster, so they have more kinetic energy and its temperature goes up. Therefore, its ΔH is positive. Eventually, when the water is boiling, water molecules have absorbed enough energy to enter the gas phase they become water vapor. But, there is no temperature change during a phase change. The amount of energy, ΔH, needed to change one mole of water from liquid at 100 o C to vapor at 100 o C is the molar heat of vaporization of water, about +41 kilojoules (kj) per mole. Of course, energy is also required to melt ice. The amount of energy, ΔH, required to turn one mole of ice at 0 o C into one mole of water at 0 o C is the molar heat of fusion (melting) of water, +6 kilojoules (kj) per mole. When the processes are reversed (i.e., freezing and condensing) the same amounts of energy are released per mole. Therefore, the ΔH values are negative. Energy flows from the system as water vapor changes to the liquid or when liquid changes to solid. Note that heats of fusion and vaporization are sometimes termed latent heats because, again, there is no temperature change associated with the change of phase the temperature remains constant during the process. So therefore, determination of ΔH for a phase change requires using the law of conservation of energy and calorimetry to measure changes in the temperature of a mass of water due to the phase change (as was done with specific heat). Thermochemistry examines the enthalpy changes that occur in chemical reactions. This enthalpy change can often be observed by carefully touching the reaction vessel. For Last Updated 01/15/13 page 3 of 19
Science Unit: 12 Lesson: 02 Suggested Duration: 7 days example, reactions of many metals with hydrochloric acid are exothermic, and the test tube that contains the reacting substances will become warm or hot to the touch as the reaction proceeds. Conversely, the reaction of sodium hydrogen carbonate (baking soda) and acetic acid solution (vinegar) is endothermic the reaction vessel gets cooler. In the laboratory, the ΔH for a reaction is determined by using the law of conservation of energy and calorimetry to measure the change in the temperature of a mass of water due to the energy involved in the chemical reaction. enthalpy of reaction = (enthalpy of products) (enthalpy of reactants) or ΔH rxn = ΔH o f (products) ΔHo f (reactants) Chemical reactions and phase changes that absorb heat energy (ΔH values are positive) are termed endothermic (or endergonic). Those that release heat energy (ΔH values are negative) are called exothermic (or exergonic). In an endothermic process, energy flows from the environment into the system, and in an exothermic process, energy flows from the system to the environment. Note: The ΔH equation as above is included in the STAAR Reference Materials. The TEKS use the terms exothermic and endothermic; some textbooks may use other terms as well. Scientists have measured the enthalpy changes for many reactions and produced tables, such as those found in the references, as Standard Heats (enthalpies) of Formation. Hess s law of the additivity of reaction heats states that when a reaction can be expressed as the sum of two reactions, then the enthalpy of the reaction is simply the algebraic sum of the enthalpies of the two component reactions. For example, consider the enthalpy of reaction for the formation of carbon dioxide from carbon monoxide and oxygen: Step 1: Write the balanced equation, and identify the component reactions: 2CO (g) + O 2 (g) 2CO 2 (g) C (s) + ½O 2 (g) CO (g) C (s) + O 2 (g) CO 2 (g) Step 2: Using a table of standard enthalpies of formation, identify the standard enthalpy of formation for each of the reactants and the product. C (s) + ½ O 2 (g) CO (g) ΔH o f CO (g) = -110.5 kj/mol ΔH o f O 2 (g) = 0 kj/mol (Note: H for ALL elements is zero.) C (s) + O 2 (g) CO 2 (g) ΔH o f CO 2 = -393.5 kj/mol Step 3: Find the sum of the ΔH o f of the reactants. The number of moles in the balanced equation must be considered. ΔH o f CO = 2 mol CO x -110.5 kj/mol = -221.0 kj ΔH o f O 2 = 1 molo 2 x 0 kj/mol = 0 kj ΔH o f (reactants) = -221.0 kj + 0 kj = -221.0 kj Step 4: Find the ΔH o f of the product. ΔH o f CO 2 = 2 mol CO 2 x -393.5 kj/mol = -787.0 kj Step 5: Find the difference between the enthalpy of products and enthalpy of reactants. ΔH rxn = ΔH o f (products) ΔHo f (reactants) ΔH rxn = ( 787.0 kj) ( 221.0 kj) ΔH rxn = -566.0 kj (exothermic) Thus, the calculated enthalpy for the formation of CO 2 in the given reaction would be -566.0 kj/2 = -283.0 kj/mole of CO 2 since there are 2 moles of CO 2 in the balanced equation. Reaction energy or reaction pathway diagrams are useful for conveying information about exothermic and endothermic reactions and the changes in potential energy as a reaction progresses. They are based on the idea that chemical reactions occur when there are collisions of sufficient energy between reacting species. Below are simplified model diagrams that show only initial and final potential energy states for the two types. Last Updated 01/15/13 page 4 of 19
Science Unit: 12 Lesson: 02 Suggested Duration: 7 days If the potential energy of the reaction system increases, energy is absorbed and the reaction is endothermic. If the potential energy of the reaction system decreases, energy is released and the reaction is exothermic. More detailed energy pathway diagrams also include activation energy, the additional energy needed to initiate a chemical reaction. A catalyst lowers the activation energy and is not consumed, so it does not affect the final energy state of the products, either lower (exothermic) or higher (endothermic). Catalysis can be shown in an energy diagram as well. See Energy Diagrams in the Resources and References for additional information. NOTE: Activation energy and catalysis are beyond the scope of the TEKS, but may be included for above-level chemistry students. INSTRUCTIONAL PROCEDURES Instructional Procedures ENGAGE I Demonstrating Phase Changes 1. Display a beaker of boiling water on a hot plate and a beaker of crushed ice to the side, each with a temperature probe or thermometer mounted in it. 2. Ask students to diagram the setup in their notebooks and record data and their answers to your questions during the demonstration. 3. Direct student attention to the boiling water. Ask: What is the temperature of the boiling water? (100 o C) Verify student predictions. 4. Point out that the hot plate is turned on, the water is boiling, and the temperature is staying at 100 o C. Ask: What does it mean that the temperature doesn t rise while the water is boiling? All of the heat energy from the hot plate is being used to boil the water. 5. Direct attention to the crushed ice. Ask: Notes for Teacher NOTE: 1 Day = 50 minutes Suggested Day 1 Materials crushed ice (for demonstration, per teacher) temperature probes or thermometers (for demonstration, 2 per teacher) water (for demonstration, per class) hot plate (for demonstration, 1 per teacher) beakers (220 ml, for demonstration, 2 per teacher) ring stands with clamps (for demonstration, 2 per teacher) safety goggles (1 pair per teacher) Safety Note Wear safety goggles. Do not touch hot liquids. Science Notebook Students record temperature data and the phases of water in their science notebooks. What is the temperature in the crushed ice? (0 o C) Verify student predictions. 6. Point out that the ice is melting and the temperature is staying at 0 o C. Ask: What does it mean that the temperature doesn t rise while the ice is melting? All of the heat energy from the surroundings is being used to melt the ice. What would happen if I were to put the beaker of crushed ice on the hot plate? Why? Heat energy from the hot plate would be used to melt the ice faster. The temperature would stay at 0 o C as long as there was ice. 7. Inform students that they will investigate the heating of water in more detail. EXPLORE/EXPLAIN I Phase Changes of Water Suggested Day 1 (continued) and 2 Last Updated 01/15/13 page 5 of 19
1. Instruct students to sketch in their notebooks a prediction of what the temperature curve look like as water goes through phase changes from solid, liquid, to gas. Materials Science Unit: 12 Lesson: 02 Suggested Duration: 7 days 2. Provide the Handout: Heating Curve/Phase Changes of Water to students. Instruct students to compare and contrast their sketch to the one on the handout individually and then again with a partner. 3. Project a copy of a handout on the board or overhead. 4. Facilitate a discussion of the temperature changes of water observed in the heating curve. Guide students to identify what is happening to the temperature in each part of the curve. Ask students to discuss the energy required for a phase change ice to water and water to water vapor. Remind students to take notes. 5. Divide students into groups of 2 4. Distribute the Handout: Heat of Fusion of Water to each student. 6. Emphasize safety concerns, and answer any questions that students may have about the procedures. 7. Instruct students to prepare a data table in their notebooks before beginning. 8. Question, monitor, and assist students as they complete the investigation. 9. Instruct students to post their data. Facilitate a discussion in which students reflect on the results to compare and contrast group data. 10. Compare experimental values for the heat of fusion of water with the accepted value, and discuss sources of error in the experimental procedure. Students should list sources of error in their science notebooks. 11. Guide students to develop a definition for heat of fusion and again for heat of vaporization. Ask students to record the definitions in their science notebooks. 12. Provide this value for the molar heat of vaporization of water: 41 kj/mole. Discuss this with students. 13. Model how to work a few additional problems related to the heat of fusion and heat of vaporization for water. Discuss as a class. 14. Instruct students to affix the Handouts: Heating Curve/Phase Changes of Water and Heat of Fusion of Water in their science notebooks. thermometer or temperature probe (1 per group) hot water (per group) ice (per group) foam cup calorimeters (see Advance Preparation, 1 per group from Lesson 1) graduated cylinder (100 ml, 1 per group) safety goggles (1 pair per student) glue or tape (per group) Attachments: Handout: Heating Curve/Phase Changes of Water (1 per student and 1 for projection) Heat of Fusion of Water (1 per student) Safety Note Wear goggles. Do not touch hot liquids. Instructional Note If you have time, student teams could heat ice water to collect data and make their own heating curves for water. If temperature probes are available, the laboratory investigation should be modified to use these tools. Procedures are included for students to implement, rather than plan, investigation procedures in the Handout: Heat of Fusion of Water. Although TEKS C.2E was addressed in Lesson 01 in which students planned and implemented an investigation, you may wish to modify the handout so that students have the opportunity to plan this investigation as well. STAAR Notes The STAAR Reference Materials include the formula for Q and for Enthalpy of Reactions. It and also includes heat constants and conversions. State Resources Texas Education Agency, STAAR End-of-Course Success Training: Lesson: Heat Transfer Science Notebook Students attach the heating curve handout, write the procedure, and draw a data table in their science notebooks. Explore/Explain II Demonstrating Endothermic and Exothermic Changes Suggested Day 3 Students then collect data and make calculations in their science notebooks. Definitions and sources of error are also included. 1. Prior to class, prepare and label solutions for demonstration (see Advance Preparation). Display solid A and solutions B and C. 2. Ask students to diagram the setup in their notebooks and record data and responses during the demonstration. 3. Put on safety goggles. Place a small scoop of the white solid A in each of four test tubes in a test tube rack. 4. Pour a half test tube full of solution B into two of the test tubes of solid A and mix. Pass the test tubes to students to feel the bottom of the test tubes. Ask students to share their observations with the class. Materials test tubes (for demonstration, 4 per teacher) test tube rack (for demonstration, 1 per teacher) spatula or scoopula (for demonstration, 1 per teacher) white solid A (sodium hydrogen carbonate (baking soda), for demonstration, 4 scoops per class) solution B (see Advance Preparation, for demonstration, 1.00 M citric acid in labeled flask per Last Updated 01/15/13 page 6 of 19
5. Pour a half test tube full of solution C into two of the test tubes of solid A, and mix them together. Pass the test tubes among students to feel the bottom of the test tubes. Ask students to share their observations with the class. 6. Guide students to develop a definition of endothermic and exothermic changes. Students should record these definitions in their notebooks. 7. Apply the definitions to chemical reactions and phase changes. Relate these changes to changes in potential energy in a system. 8. Instruct students to discuss their notes with a shoulder partner and to revise their notes if needed. Monitor and assist. Ask: What energy transformations occur in chemical reactions? Energy is released or absorbed. How can energy transfer be observed in an exothermic reaction? The temperature will increase it feels warmer. How can energy transfer be observed in an endothermic reaction? The temperature decreases it feels cooler. Science Unit: 12 Lesson: 02 Suggested Duration: 7 days teacher) solution C (see Advance Preparation, for demonstration, 1.00 M calcium chloride in labeled flask per teacher) safety goggles (for demonstration, 1 pair per teacher) Instructional Note Exothermic and endothermic apply to chemical and physical changes. Refer to Background Information to help guide student thinking. Check For Understanding Monitor and assist as student partners review and revise their notes. Science Notebook Students make observations of reactions and energy transfer and write definitions of endothermic and exothermic reactions in their science notebooks. EXPLORE III Enthalpy of Reaction Suggested Day 4 1. Divide students into groups of 2 4. Distribute the Handout: Enthalpy of Reaction to each student. 2. Instruct students to read the investigation and prepare a data table in their notebooks. 3. Stress all of the safety concerns, and answer any questions that students may have regarding the instructions. 4. Monitor students for safety and disposal as students complete the investigation. Assist students as necessary. 5. Instruct students to affix the handout in their notebooks. Materials foam cup calorimeters (see Advance Preparation, 1 per group) temperature probes or thermometers (1 per group) electronic balance (1 per group) sodium hydrogen carbonate (baking soda, 2.0 g per group) 1.00 M citric acid solution (75 ml per group) zinc metal pieces (2.0 g per group) 1.00 M hydrochloric acid (50 ml per group) graduated cylinder (100 ml, 1 per group) safety goggles (1 per student) glue or tape (per group) Attachments: Handout: Enthalpy of Reaction (1 per student) Teacher Resource: Enthalpy of Reaction KEY Safety Note Wear safety goggles. Review all of the MSDS sheets for safe handling and disposal of chemicals. Science Notebook Students prepare a data table, record data and observations, and make calculations in their science notebooks. EXPLAIN III Calculating Enthalpy Suggested Days 4 (continued) and 5 1. Review the observations and results from the Handout: Enthalpy of Reaction. 2. Facilitate a class discussion on the Calculations and Analysis sections provide assistance to students as needed. 3. Continue the discussion, including energy transfer and the conservation of energy. Ask: What energy transformations occur in chemical reactions? Energy is absorbed from the environment or flows into the environment. How can energy transfer be observed in an exothermic reaction? (The temperature of the system increases.) How can energy transfer be observed in an endothermic reaction? (The temperature of the system decreases.) Materials Handout: Enthalpy of Reaction (from previous activity) Instructional Note Use the example from the Background Information at the beginning of this lesson to introduce the table(s) of standard heats of formation and illustrate how these will be used in solving problems involving the heat of reaction. Last Updated 01/15/13 page 7 of 19
4. Generalize enthalpy changes to both physical and chemical changes. Introduce heat of combustion and heat of solution to go along with heat of reaction, heat of fusion, and heat of vaporization. 5. Introduce thermochemical equations using a simple equation, such as CaO (s) + H 2 O (l) Ask: Ca(OH) 2 + 65.2 kj. Science Unit: 12 Lesson: 02 Suggested Duration: 7 days Use the Background Information to present simple endothermic and exothermic reaction pathways to students. Science Notebook Students record definitions and make calculations in their science notebooks. What are the products in the equation as it is written? (Calcium hydroxide and energy) What energy transfer occurred in the reaction? (Energy was released.) Is this an endothermic or exothermic reaction? (Exothermic) What is the enthalpy of reaction? (ΔH = -65.2 kj) What is shown in a thermochemical equation? The heat absorbed or released in a reaction is shown as a reactant or product. How can this equation be represented using ΔH? CaO (s) + H 2 O (l) Ca(OH) 2 ΔH = 65.2 kj 6. Guide students in writing the thermochemical equation and determining the quantity of heat absorbed or produced for a given number of moles or mass of reactant for several simple reactions: decomposition, combustion, and solution. Heat of Decomposition Sample Problem: Given the equation, 2NaHCO 3 (s) + 129 kj Na 2 CO 3 (s) + H 2 O (g) + CO 2 (g) determine the kj of heat needed to decompose 2.24 moles of sodium hydrogen carbonate. From the given, 129 kj is needed to decompose 2 moles of solid NaHCO 3, so a simple ratio is used to find the heat required to decompose 2.4 moles. ΔH = 2.4 mol x (129 kj/2 mol) = 144 kj Ask: How much heat would be required to decompose 100 g of NaHCO 3 (s)? (100 g / 84.0 g/mol) x 129 kj/2 mol = 76.9 kj 7. Instruct students to practice (moles and grams) with thermochemical equations representing heat of combustion and heat of solution (accepted values listed). Heat of Combustion: CH 4 (g) + 2O 2 (g) = CO 2 (g) + 2H 2 O (l) ΔH = 890kJ/mol Heat of Solution: NaOH (s) Na + (aq) + OH - (aq) ΔH = 445.1 kj/mol 8. Introduce the table(s) of standard heats of formation, and illustrate how these will be used in solving problems involving the heat of reaction. Use the example from the Background Information at the beginning of this lesson. 9. Use the Background Information to present simple endothermic and exothermic reaction pathways to students. Then, make pathway diagrams for the reactions in steps 6 and 7. ELABORATE Practice Problems Suggested Day 6 1. Review how to solve problems calculating the heat of reaction using standard heats of formation and the relationship: Enthalpy of reaction = (enthalpy of products) (enthalpy of reactants) ΔH = ΔH o f(products) ΔH o f (reactants) 2. Distribute the Handout: Enthalpy Change Calculations, and assist students as necessary. 3. Assign additional problems as needed, from a locally adopted textbook and/or other Attachments: Handout: Enthalpy Change Calculations (1 per student) Teacher Resource: Enthalpy Change Calculations KEY STAAR Notes The STAAR Reference Materials includes the formula for Last Updated 01/15/13 page 8 of 19
resources. Enthalpy of Reaction. Science Unit: 12 Lesson: 02 Suggested Duration: 7 days Thermochemical equations in chemical reactions will be tested as a Readiness Standard under Reporting Category 4: Gases and Thermochemistry. Science Notebook Students should work problems in their science notebooks. EVALUATE Performance Indicator Suggested Day 7 High School Science Unit 12 PI 02 Use calorimetry to measure an exothermic and an endothermic enthalpy of reaction. Write a summary report that includes your procedures, data collected, calculations, a discussion of the reactions, and an error analysis. Standard(s): C.2E, C.2F, C.2I, C.11C, C.11E ELPS ELPS.c.3D, ELPS.c.3E, ELPS.c.5B 1. Instruct students to review their Enthalpy of Reactions investigation and notes. Explain your requirements for question/hypotheses, procedures, materials, safety precautions, etc. in the student investigations and their summary reports. 2. Provide students with one or more endothermic and exothermic changes from which to choose one of each to plan and investigate (see Instructional Notes). 3. Instruct students to obtain approval of their team s procedures before beginning the investigation. 4. Inform students that they may work as teams, but must submit individual summary reports. 5. Monitor for safety and disposal as students complete the activity. 6. The summary report may be assigned for homework if necessary. Materials foam cup calorimeters (and/or other preferred types, 1 per group) temperature probes or thermometers (1 per group) graduated cylinder (100 ml, 1 per group) chemicals selected by students (see Advance Preparation, varies by group) electronic balance (1 per group) weighing paper (1 sheet per group) distilled or deionized water (approximately 300 ml per group) Attachments: Handout: Laboratory Activity: Enthalpy of Reaction (from previous activity) Instructional Note You may wish to provide additional equipment for students to choose from as part of their investigative planning. You may want students to choose one of the endothermic and exothermic processes from this lesson. Some additional endothermic processes are: sodium hydrogen carbonate and hydrochloric acid ammonium nitrate and water (Heat of solution = 25.7 kj/mol ) Some additional exothermic processes are: reaction of Mg with hydrochloric acid calcium chloride and water (Heat of solution = -82.8 kj/mol ) Standard Heats of Formation: Last Updated 01/15/13 page 9 of 19
Heating Curve/Phase Changes of Water Unit: 12 Lesson: 02 Diagram courtesy A. Venegas 2012, TESCCC 01/15/13 page 1 of 1
Heat of Fusion of Water Laboratory Investigation Unit: 12 Lesson: 02 Introduction: The amount of energy required to change a solid to a liquid at constant pressure is called the heat of fusion of the substance. In this investigation, you will use calorimetry to determine the molar heat of fusion of ice. The heat required to melt a given mass of ice will be determined using the formula Q = mc p T. Remember that the specific heat of water (c p ) is 4.18 J/g o C. The mass of melted ice and water will be determined by assuming that 1.00 ml of water has a mass of 1.00 g since the density of water is 1.00 g/ml. The heat of fusion of ice will be calculated from the experimental data. Note: The heat of fusion of ice = heat absorbed by ice / mass of ice. Materials: foam cup calorimeter temperature probe or thermometer 100 ml graduated cylinder hot water ice cubes Safety Notes: Wear safety goggles. Be careful with hot liquids. Procedure: (Record all data and observations in your science notebooks.) 1. Fill a 100 ml graduated cylinder with hot water, and let it stand for about one minute. Carefully pour the water into the sink. 2. In the heated graduated cylinder, measure 70.0 ml of hot water, and pour it into the foam cup calorimeter. Measure the temperature of the water, and record the data. 3. Add an ice cube to the hot water, and cover it with the foam lid containing the temperature probe or thermometer. Swirl it gently, but DO NOT STIR with the thermometer or probe. When the ice cube is completely melted and the temperature is stable, record the lowest temperature. To ensure an accurate reading, be sure the thermometer is not resting on the bottom of the cup. 4. Carefully pour all of the water from the calorimeter into the graduated cylinder, and measure the total volume. Record your data. Calculations: 1. From your data, determine the change in temperature of the hot water and final mass of the hot water. Calculate the heat lost by the hot water using Q = mc p T. 2. From the volume of the cooled water, determine the mass of ice that melted. 3. Using the law of conservation of energy, calculate the heat of fusion of ice in joules/g. 4. Calculate the molar heat of fusion of ice from your experimental results. 2012, TESCCC 01/15/13 page 1 of 2
Analysis: Unit: 12 Lesson: 02 1. Why was the hot water placed in the graduated cylinder and then poured out into the sink during first step of the investigation? 2. The accepted heat of fusion of ice is 334 J/g. The molar heat of fusion of ice is 6.01 kj/mole. Compare your results to the accepted values by calculating the percent error. 3. Discuss sources of error in this investigation. Suggest ways that some errors might be minimized. 2012, TESCCC 01/15/13 page 2 of 2
Enthalpy of Reaction Laboratory Investigation Unit: 12 Lesson: 02 Introduction: Enthalpy is the heat content of a system at constant pressure. Using calorimetry, the change in enthalpy of a chemical reaction can be observed and calculated. In this investigation, you will determine the heats of reaction for both an endothermic and exothermic reaction. Materials: foam cup calorimeter temperature probe or thermometer electronic balance sodium hydrogen carbonate 1.00 M citric acid solution 1.00 M hydrochloric acid zinc metal pieces 100 ml graduated cylinder Safety Notes: Wear safety goggles. Review all MSDS sheets for proper handling and disposal of chemicals. Procedure: Part 1 (Record all data and observations in your science notebooks.) 1. Measure 75.0 ml of 1.00 M citric acid solution in a graduated cylinder, and pour it into the cup calorimeter. Record the temperature of the solution to the nearest 0.1 o C. 2. Measure 2.00 g of sodium hydrogen carbonate. 3. Add the sodium hydrogen carbonate to the citric acid solution, and place the lid with a thermometer or probe over the cup. Swirl it gently. 4. Record the final temperature when the reaction is complete. The temperature should remain constant. 5. Dispose of the contents of the calorimeter as directed by your teacher. Rinse it thoroughly. Procedure: Part 2 (Record all data and observations in your science notebooks.) 1. Measure 50.0 ml of 1.00 M hydrochloric acid using a graduated cylinder, and pour it into the cup calorimeter. Record the temperature of the solution to the nearest 0.1 o C. 2. Obtain 2.00 g of small zinc metal pieces. 3. Add the zinc to the hydrochloric acid solution, and place the lid with a thermometer or probe over the calorimeter cup. 4. Record the final temperature when the reaction is complete. The temperature should remain constant. 5. Dispose of the contents of the calorimeter as directed by your teacher. Rinse it thoroughly. Calculations: (Show your calculations in your science notebooks.) 1. Calculate the heat of reaction for the citric acid and sodium hydrogen carbonate reaction. Assume that the mass of solute in the solution is negligible and that the heat is transferred to or from the 75.0 ml of water. (Hint: Remember that Q = mc p T and the specific heat of water is 4.184 J/g o C.) 2. Calculate the heat of reaction for the zinc and hydrochloric acid reaction. Again, assume that the mass of solute in the hydrochloric acid is negligible and the heat is transferred to or from the 50.0 ml of water. Analysis: 1. Discuss the enthalpy change of each reaction. Classify each reaction as either endothermic or exothermic. Explain your reasoning. 2. How could you determine the molar heat of reaction for sodium hydrogen carbonate? For zinc? 3. Discuss sources of error in this investigation, and suggest ways that some errors might be minimized. 2012, TESCCC 1/15/13 page 1 of 1
Enthalpy of Reaction KEY Laboratory Investigation Unit: 12 Lesson: 02 Introduction: Enthalpy is the heat content of a system at constant pressure. Using calorimetry, the enthalpy of chemical reactions can be observed and calculated. In this investigation, heat of reaction for both endothermic and exothermic reactions will be determined. Part 1: The reaction of citric acid and sodium hydrogen carbonate (baking soda) is endothermic. An equation for the reaction is: H 3 C 6 H 5 O 7 (aq) + 3 NaHCO 3 (s) 3 CO 2 (g) + 3 H 2 O (l) + Na 3 C 6 H 5 O 7 (aq) Part 2: The reaction of zinc and hydrochloric acid is exothermic. An equation of the reaction is: Zn (s) + 2 HCl ZnCl 2 + H 2 (g) Procedure: Part 1 1. Measure 75.0 ml of 1.00 M citric acid solution in a graduated cylinder, and pour it into the cup calorimeter. Record the temperature of the solution to the nearest 0.1 o C. Remind students to record the initial temperature BEFORE adding the sodium hydrogen carbonate. For purposes of this investigation, you will assume that the specific heat of the citric acid solution is the same as that of water and the amount of solute in the solution is very small. Thus, the heat will be transferred from 75.0 g of water (Density of water = l.00 g/ml). 2. Obtain 2.00 g of sodium hydrogen carbonate. The mass does not have to be exactly 2.00 g, but students should determine the mass to the nearest 0.01 g using an electronic balance. 3. Add the sodium hydrogen carbonate to the citric acid solution, and place the lid with a thermometer or temperature probe over the cup. Swirl it gently. Caution students to NOT use the thermometer as a stirring rod. 4. Record the final temperature when the reaction is complete. The temperature should remain constant. Students should record the lowest temperature reached by the reacting substances. The temperature should be stable for a short time. If too much time is allowed, the temperature will begin to change again. 5. Dispose of the contents of the calorimeter as directed by your teacher, and thoroughly clean or replace the cup. Procedure: Part 2 1. Measure 50.0 ml of 1.00 M hydrochloric acid using a graduated cylinder, and pour it into the cup calorimeter. Record the temperature of the solution to the nearest 0.1 o C. Remind students to record the initial temperature BEFORE adding the zinc. For purposes of this investigation, you will assume that the specific heat of the hydrochloric 2012, TESCCC 01/15/13 page 1 of 3
Unit: 12 Lesson: 02 acid solution is the same as that of water and the amount of solute in the solution is very small. Thus, the heat will be transferred to 50.0 g of water (Density of water = l.00 g/ml). 2. Obtain 2.00 g of small zinc metal pieces. The mass does not have to be exactly 2.00 g but students should determine the mass to the nearest 0.01 g using an electronic balance. 3. Add the zinc to the hydrochloric acid solution, and place the lid with a thermometer over the cup. Caution students to NOT use the thermometer as a stirring rod. 4. Record the final temperature when the reaction is complete. The temperature should remain constant. Students should record the highest temperature reached by the reacting substances. The temperature should be stable for a short time. If too much time is allowed, the temperature will begin to change again. 5. Dispose of the contents of the calorimeter as directed by your teacher. Sample Data Table: Part 1 Part 2 Initial Temperature 22 C Initial Temperature 22 C Final Temperature Change in Temperature Final Temperature Change in Temperature Mass of NaHCO 3 2.00 g Mass of Zn 2.00 g Moles of NaHCO 3 2.00 g/84 g/mole Moles of Zn 2.00 g/65.6 g/mole Heat (Q) Transferred Molar Heat of Reaction Heat (Q) Transferred Molar Heat of Reaction Calculations: 1. Calculate the heat of reaction for the citric acid and sodium hydrogen carbonate. Assume that the mass of solute in the solution is negligible and the heat is transferred to or from the 75.0 ml of water. (Hint: Remember that Q = mc p T and the specific heat of water is 4.18 J/g C.) Q = 75.0 g x 4.18/g C x (Final Temperature Initial Temperature) 2. Calculate the heat of reaction for the zinc and hydrochloric acid. Again, assume that the mass of solute in the hydrochloric acid is negligible and the heat is transferred to or from the 50.0 ml of water. Q = 50.0 g x 4.18/g C x (Final Temperature Initial Temperature) 2012, TESCCC 01/15/13 page 2 of 3
Analysis: Unit: 12 Lesson: 02 1. Discuss the enthalpy of each reaction. Was the reaction endothermic or exothermic? Explain your reasoning. Part 1: The final temperature is lower than the initial temperature. This indicates that heat was transferred from the solution (water and citric acid) to the products. This is an endothermic reaction, so the change in enthalpy (ΔH) is positive. Part 2: The final temperature is higher than the initial temperature. This indicates that heat was transferred to the solution (water and hydrochloric acid) as zinc chloride was formed. This is an exothermic reaction, so the change in enthalpy (ΔH) is negative. 2. How could you determine the molar heat of reaction for sodium hydrogen carbonate? For zinc? Molar heat of reaction = Value for Q computed from data / Moles of sodium hydrogen carbonate used 3. Discuss sources of error in this investigation, and suggest ways that some errors might be minimized. Students may suggest that: 1) the calorimeter may not have prevented all heat transfer to/from the environment, 2) human error in making measurements, or 3) loss of products to the environment (gases may escape in the reactions). NOTE: A more quantitative approach to this laboratory activity would be to have students calculate the expected heat of reaction using the standard heats of formation and then determine the percent error using their experimental results. (For both Zn and H 2, heat of formation is 0 kj/mol.) Standard Heats of Formation H 3 C 6 H 5 O 7-2684 kj/mol CO 2 (g) -393.5 kj/mol NaHCO 3 (s) -950.8 kj/mol H 2 O (l) -285.8 kj/mol Na 3 C 6 H 5 O 7-2307 kj/mol HCl (aq) -167 kj/mol ZnCl 2 (aq) -487 kj/mol Standard Heats of Solution NaHCO 3 (aq) H 3 C 6 H 5 O 7, anhydrous (aq) H 3 C 6 H 5 O 7, monohydrate (aq) -18.0 kj/mol 18.2 kj/mol 29.1 kj/mol 2012, TESCCC 01/15/13 page 3 of 3
Enthalpy Change Calculations Unit: 12 Lesson: 01 Use the appropriate formula found in the STAAR Reference Materials to solve the calculations below. Use a reference thermodynamic table to find Δ H f values. Be sure to consider significant digits in your answers. Complete your answers in your science notebook. 1. Explain the difference between specific heat and enthalpy change. 2. What are endothermic (endergonic) and exothermic (exergonic) reactions? Which type of reaction may occur spontaneously? 3. Calculate the enthalpy change in the following reaction: Carbon monoxide + Oxygen Carbon dioxide CO + O 2 2CO 2 4. For each reaction, calculate ΔH and classify each as endothermic or exothermic: a. 2NO (g) + O 2 (g) 2NO 2 (g) b. 4FeO (s) + O 2(g) 2Fe 2 O 3 (cr) 5. Draw an energy pathway diagram that has a H = - 100 kj/mol. 2012, TESCCC 01/15/13 page 1 of 1
Enthalpy Change Calculations KEY Unit: 12 Lesson: 01 Use the appropriate formula found in the STAAR Reference Materials to solve the calculations below. Use a reference thermodynamic table to find Δ H f values. Be sure to consider significant digits in your answers. 1. Explain the difference between specific heat and enthalpy change. Enthalpy change is the energy gained or lost per mole or gram, but specific heat is the energy required to raise the temperature of a gram of substance one degree. 2. What are endothermic (endergonic) and exothermic (exergonic) reactions? Which type of reaction may occur spontaneously? Exothermic reactions have negative ΔH, and heat is released. Endothermic reactions have positive ΔH, and heat is absorbed. (Note: Exothermic reactions may occur spontaneously.) 3. Calculate the enthalpy change in the following reaction: Carbon monoxide + Oxygen Carbon Dioxide CO + O 2 2CO 2 Reactants: 2 mol CO X -110.5 kj = - 221.0 kj 1 mol CO Free elements have enthalpy of 0. Products: 2 mol CO 2 x -393.5 kj = - 787.0 kj 1 mol CO 2-787.0 kj (-221.0kJ) = - 566.0 kj 4. For each reaction, calculate H and classify each as endothermic or exothermic: a. 2NO (g) + O 2 (g) 2NO 2 (g) Reactants: 2 mol NO x 90.25 kj = 180.5 kj 1 mol NO Free elements have enthalpy of 0. Products: 2 mol NO 2 x 33.18 kj = 66.36 kj 1 mol NO 2 66.36 kj 180.5 kj = -114.1 kj exothermic reaction (4 sig. figs) 2012, TESCCC 01/15/13 page 1 of 2
b. 4FeO (s) + O 2(g) 2Fe 2 O 3 (cr) Unit: 12 Lesson: 01 Reactants: 4 mol FeO x -272.0 kj 1 mol FeO = -1,088 kj Free elements have enthalpy of 0. Products: 2 mol Fe 2 O 3 x -824.2 kj = -1648.4 kj 1 mol Fe 2 O 3-1,648 kj (-1,088 kj) = -560.4 kj exothermic reaction 5. Draw an energy pathway diagram that has a H = - 100 kj/mol. 2012, TESCCC 01/15/13 page 2 of 2