Heat Energy in Chemical Reactions
Chemistry HS/Science Unit: 12 Lesson: 02 Suggested Duration: 7 days
Heat Energy in Chemical Reactions Lesson Synopsis: This lesson introduces thermochemistry, the study of heat energy in chemical reactions. Students use calorimetry and the STAAR Chemistry 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: C.11 C.11C C.11E
Science concepts. The student understands the energy changes that occur in chemical reactions. The student is expected to: Use thermochemical equations to calculate energy changes that occur in chemical reactions and classify reactions as exothermic or endothermic. Readiness Standard Use calorimetry to calculate the heat of a chemical process. Supporting Standard
Scientific Process TEKS: 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. Collect data and make measurements with accuracy and precision.
C.2F 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 Indicator(s): • Use calorimetry to measure an exothermic and endothermic enthalpy of reaction. Write a summary report that includes your procedures, data collected, calculations, a discussion of the reactions, and an error analysis. (C.2E, C.2F, C.2I; C.11C, C.11E) 3D, 3E; 5B
Key Understandings and Guiding Questions: • • •
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 enthalpy
• fusion • vaporization • heat of fusion • heat of vaporization
• heat of reaction • heat of combustion • heat of solution • heating curve/phase diagram of water
Materials: Refer to Notes for Teacher section for materials.
Attachments: ©2012, TESCCC
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• • • • • • •
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 Teacher Resource: Performance Indicator Instructions KEY
Advance Preparation: 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 Laboratory Investigation. • Prepare all solutions needed for the investigation. • Review all of the MSDS for chemical use and disposal. • Be sure to read the Enthalpy of Reaction Teacher Notes and Key and to perform the activity to make adjustments as needed. 4. Prior to Day 6: • Preview the practice calculations to identify student obstacles. • Find a thermodynamic reference table for students to find ∆Hf values. 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. • 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 = Hf – Hi 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. © 2008, TESCCC
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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 100oC to vapor at 100oC 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 0oC into one mole of water at 0oC 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 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 ΔHrxn = ΔHof (products) – ΔHof (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 Chemistry 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) + O2 (g) C (s) + ½O2 (g) C (s) + O2 (g)
2CO2 (g) CO (g) CO2 (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 O2 (g) CO (g) ΔHof CO (g) = -110.5 kJ/mol ΔHof O2 (g) = 0 kJ/mol (Note: H for ALL elements is zero.) © 2008, TESCCC
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C (s) + O2 (g) CO2 (g) ΔHof CO2 = -393.5 kJ/mol Step 3: Find the sum of the ΔHof of the reactants. The number of moles in the balanced equation must be considered. ΔHof CO = 2 mol CO x -110.5 kJ/mol = -221.0 kJ ΔHof O2 = 1 mol O2 x 0 kJ/mol = 0 kJ ΔHof (reactants) = -221.0 kJ + 0 kJ = -221.0 kJ Step 4: Find the ΔHof of the product. ΔHof CO2 = 2 mol CO2 x -393.5 kJ/mol = -787.0 kJ Step 5: Find the difference between the enthalpy of products and enthalpy of reactants. ΔHrxn = ΔHof (products) – ΔHof (reactants) ΔHrxn = (-787.0 kJ) – (-221.0 kJ) ΔHrxn = -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 CO2 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.
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 the reaction activation energy, the additional energy needed to initiate a chemical reaction. A catalyst (enzyme) 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. NOTE: Activation energy and catalysis are beyond the scope of the TEKS, but may be included for above-level chemistry students.
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Image Sources: Venegas, A. (Artist) (2012). Endothermic reaction pathway. [Print Drawing]. Used with permission. Venegas, A. (Artist) (2012). Exothermic reaction pathway [Print Drawing]. Used with permission. Mann, J. C. (Artist). (2006, November 26). Activation2 updated [Web Drawing]. Retrieved from http://commons.wikimedia.org/wiki/File:Activation2_updated.svg
GETTING READY FOR INSTRUCTION SUPPLEMENTAL PLANNING DOCUMENT Instructors are encouraged to supplement and substitute resources, materials, and activities to differentiate instruction to address the needs of learners. The Exemplar Lessons are one approach to teaching and reaching the Performance Indicators and Specificity in the Instructional Focus Document for this unit. Instructors are encouraged to create original lessons using the Content Creator in the Tools Tab located at the top of the page. All originally authored lessons can be saved in the “My CSCOPE” Tab within the “My Content” area.
INSTRUCTIONAL PROCEDURES Instructional Procedures
Notes for Teacher
ENGAGE I – Demonstrating Phase Changes
NOTE: 1 Day = 50 minutes Suggested Day 1
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. Pose the following questions as you conduct the demonstration: • What is the temperature of the boiling water? (100oC) Verify student predictions. 4. Point out that the hot plate is turned on, the water is boiling, and the temperature is staying at 100oC: • 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.
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) • apron ( 1 for teacher)
5. Direct attention to the crushed ice: • What is the temperature in the crushed ice? (0oC) © 2008, TESCCC
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Verify student predictions. 6. Point out that the ice is melting and the temperature is staying at 0 oC: • 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 oC 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 1. Instruct students to sketch in their notebooks a prediction of what the temperature curve looks like as water goes through phase changes from solid, liquid, to gas. 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 the class 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 © 2008, TESCCC
Revised 09/29/08
Safety Notes: Wear safety goggles and an apron. Do not touch hot liquids.
Science Notebooks: Students record temperature data and the phases of water in their science notebooks.
Suggested Days 1 (continued) and 2
Materials: • 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) • apron (1 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 Notes: Wear goggles and an apron. Do not touch hot liquids. Instructional Notes: 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. page 6 of 11
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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.
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 Chemistry 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-ofCourse Success Training: Lesson: Heat Transfer
Science Notebooks: Students attach the heating curve handout, write the procedure, and draw a data table in their science notebooks. Students then collect data and make calculations in their science notebooks. Definitions and sources of error are also included.
Explore/Explain II – Demonstrating Endothermic and Exothermic Changes 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. Review safety procedures, put on safety goggles and a lab apron. 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. 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. © 2008, TESCCC
Revised 09/29/08
Suggested Day 3
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 teacher) • solution C (see Advance Preparation, for demonstration, 1.00 M calcium page 7 of 11
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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. Pose the following questions in closing: • 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.
• •
chloride in labeled flask per teacher) safety goggles (1 pair per teacher) apron (1 per teacher)
Instructional Notes: Exothermic and endothermic apply to chemical and physical changes. Refer to Background Information to help guide student thinking.
Safety Notes: Wear goggles and an apron. Review all MSDS for safe handling and disposal of chemicals.
Check for Understanding: Monitor and assist as student partners review and revise their notes.
Science Notebooks: 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 the class 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) • apron (1 per student) • glue or tape (per group) Attachments: • Handout: Enthalpy of Reaction (1 per
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Instructional Procedures
Notes for Teacher •
student) Teacher Resource: Enthalpy of Reaction KEY
Safety Notes: Wear safety goggles. Review all of the MSDS sheets for safe handling and disposal of chemicals.
Science Notebooks: Students prepare a data table, record data and observations, and make calculations in their science notebooks.
EXPLAIN III – Calculating Enthalpy 1. Review the observations and results from the Handout: Enthalpy of Reaction. • How can energy transfer be observed in an exothermic (The temperature the system increases.) • reaction? How can energy transfer beof observed in an endothermic reaction? (The temperature of the system decreases.) 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) + H2O (l) Ca(OH)2 + 65.2 kJ. Ask: • 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) + H2O (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:
Suggested Days 4 (continued) and 5 Attachments: • Handout: Enthalpy of Reaction (from previous activity) Instructional Notes: 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. Use the Background Information to present simple endothermic and exothermic reaction pathways to students. Add in activation energy and catalysis diagrams as appropriate. Consider using Internet resources to help in presenting this section.
Science Notebooks: Students record definitions and make calculations in their science notebooks.
Given the equation, 2NaHCO3 (s) + 129 kJ
Na2CO3 (s) + H2O (g) + CO2 (g)
determine the kJ of heat needed to decompose 2.24 moles of sodium © 2008, TESCCC
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hydrogen carbonate. From the given, 129 kJ is needed to decompose 2 moles of solid NaHCO3, 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 NaHCO3 (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 below). Heat of Combustion: CH4 (g) + 2O2 (g) CO2 (g) + 2H2O (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 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 = ΔHof (products) – ΔHof (reactants) 2. Distribute the Handout: Enthalpy Change Calculations, and project a Thermodynamics Reference table (see Advance Preparation). Assist students as necessary. 3. Assign additional problems as needed from a locally adopted textbook and/or other resources.
Suggested Day 6 Attachments: • Handout: Enthalpy Change Calculations (1 per student) • Teacher Resource: Enthalpy Change Calculations KEY
STAAR Notes: The STAAR Chemistry Reference Materials includes the formula for Enthalpy of Reaction. Thermochemical equations in chemical reactions will be tested as a Readiness Standard under Reporting Category 4: Gases and Thermochemistry.
Science Notebooks: Students should work problems in their science notebooks.
EVALUATE – Performance Indicator © 2008, TESCCC
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Performance Indicator • Use calorimetry to measure an exothermic and endothermic enthalpy of reaction. Write a summary report that includes your procedures, data collected, calculations, a discussion of the reactions, and an error analysis. (C.2E, C.2F, C.2I; C.11C, C.11E) • 3D, 3E; 5B 1. Refer to the Teacher Resource: Performance Indicator Instructions KEY and the Handout: Laboratory Activity: Enthalpy of Reaction for information on administering the assessment.
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: • Teacher Resource: Performance Indicator Instructions KEY • Handout: Laboratory Activity: Enthalpy of Reaction (from previous activity)
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Heat Energy in Chemical Reactions Lesson Synopsis: This lesson introduces thermochemistry, the study of heat energy in chemical reactions. Students use calorimetry and the STAAR Chemistry 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: C.11 C.11C C.11E
Science concepts. The student understands the energy changes that occur in chemical reactions. The student is expected to: Use thermochemical equations to calculate energy changes that occur in chemical reactions and classify reactions as exothermic or endothermic. Readiness Standard Use calorimetry to calculate the heat of a chemical process. Supporting Standard
Scientific Process TEKS: 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. Collect data and make measurements with accuracy and precision.
C.2F 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 Indicator(s): • Use calorimetry to measure an exothermic and endothermic enthalpy of reaction. Write a summary report that includes your procedures, data collected, calculations, a discussion of the reactions, and an error analysis. (C.2E, C.2F, C.2I; C.11C, C.11E) 3D, 3E; 5B
Key Understandings and Guiding Questions: • • •
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 enthalpy
• fusion • vaporization • heat of fusion • heat of vaporization
• heat of reaction • heat of combustion • heat of solution • heating curve/phase diagram of water
Materials: Refer to Notes for Teacher section for materials.
Attachments: ©2012, TESCCC
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• • • • • • •
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 Teacher Resource: Performance Indicator Instructions KEY
Advance Preparation: 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 Laboratory Investigation. • Prepare all solutions needed for the investigation. • Review all of the MSDS for chemical use and disposal. • Be sure to read the Enthalpy of Reaction Teacher Notes and Key and to perform the activity to make adjustments as needed. 4. Prior to Day 6: • Preview the practice calculations to identify student obstacles. • Find a thermodynamic reference table for students to find ∆Hf values. 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. • 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 = Hf – Hi 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. © 2008, TESCCC
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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 100oC to vapor at 100oC 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 0oC into one mole of water at 0oC 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 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 ΔHrxn = ΔHof (products) – ΔHof (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 Chemistry 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) + O2 (g) C (s) + ½O2 (g) C (s) + O2 (g)
2CO2 (g) CO (g) CO2 (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 O2 (g) CO (g) ΔHof CO (g) = -110.5 kJ/mol ΔHof O2 (g) = 0 kJ/mol (Note: H for ALL elements is zero.) © 2008, TESCCC
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C (s) + O2 (g) CO2 (g) ΔHof CO2 = -393.5 kJ/mol Step 3: Find the sum of the ΔHof of the reactants. The number of moles in the balanced equation must be considered. ΔHof CO = 2 mol CO x -110.5 kJ/mol = -221.0 kJ ΔHof O2 = 1 mol O2 x 0 kJ/mol = 0 kJ ΔHof (reactants) = -221.0 kJ + 0 kJ = -221.0 kJ Step 4: Find the ΔHof of the product. ΔHof CO2 = 2 mol CO2 x -393.5 kJ/mol = -787.0 kJ Step 5: Find the difference between the enthalpy of products and enthalpy of reactants. ΔHrxn = ΔHof (products) – ΔHof (reactants) ΔHrxn = (-787.0 kJ) – (-221.0 kJ) ΔHrxn = -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 CO2 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.
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 the reaction activation energy, the additional energy needed to initiate a chemical reaction. A catalyst (enzyme) 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. NOTE: Activation energy and catalysis are beyond the scope of the TEKS, but may be included for above-level chemistry students.
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Image Sources: Venegas, A. (Artist) (2012). Endothermic reaction pathway. [Print Drawing]. Used with permission. Venegas, A. (Artist) (2012). Exothermic reaction pathway [Print Drawing]. Used with permission. Mann, J. C. (Artist). (2006, November 26). Activation2 updated [Web Drawing]. Retrieved from http://commons.wikimedia.org/wiki/File:Activation2_updated.svg
GETTING READY FOR INSTRUCTION SUPPLEMENTAL PLANNING DOCUMENT Instructors are encouraged to supplement and substitute resources, materials, and activities to differentiate instruction to address the needs of learners. The Exemplar Lessons are one approach to teaching and reaching the Performance Indicators and Specificity in the Instructional Focus Document for this unit. Instructors are encouraged to create original lessons using the Content Creator in the Tools Tab located at the top of the page. All originally authored lessons can be saved in the “My CSCOPE” Tab within the “My Content” area.
INSTRUCTIONAL PROCEDURES Instructional Procedures
Notes for Teacher
ENGAGE I – Demonstrating Phase Changes
NOTE: 1 Day = 50 minutes Suggested Day 1
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. Pose the following questions as you conduct the demonstration: • What is the temperature of the boiling water? (100oC) Verify student predictions. 4. Point out that the hot plate is turned on, the water is boiling, and the temperature is staying at 100oC: • 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.
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) • apron ( 1 for teacher)
5. Direct attention to the crushed ice: • What is the temperature in the crushed ice? (0oC) © 2008, TESCCC
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Verify student predictions. 6. Point out that the ice is melting and the temperature is staying at 0 oC: • 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 oC 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 1. Instruct students to sketch in their notebooks a prediction of what the temperature curve looks like as water goes through phase changes from solid, liquid, to gas. 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 the class 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 © 2008, TESCCC
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Safety Notes: Wear safety goggles and an apron. Do not touch hot liquids.
Science Notebooks: Students record temperature data and the phases of water in their science notebooks.
Suggested Days 1 (continued) and 2
Materials: • 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) • apron (1 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 Notes: Wear goggles and an apron. Do not touch hot liquids. Instructional Notes: 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. page 6 of 11
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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.
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 Chemistry 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-ofCourse Success Training: Lesson: Heat Transfer
Science Notebooks: Students attach the heating curve handout, write the procedure, and draw a data table in their science notebooks. Students then collect data and make calculations in their science notebooks. Definitions and sources of error are also included.
Explore/Explain II – Demonstrating Endothermic and Exothermic Changes 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. Review safety procedures, put on safety goggles and a lab apron. 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. 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. © 2008, TESCCC
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Suggested Day 3
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 teacher) • solution C (see Advance Preparation, for demonstration, 1.00 M calcium page 7 of 11
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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. Pose the following questions in closing: • 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.
• •
chloride in labeled flask per teacher) safety goggles (1 pair per teacher) apron (1 per teacher)
Instructional Notes: Exothermic and endothermic apply to chemical and physical changes. Refer to Background Information to help guide student thinking.
Safety Notes: Wear goggles and an apron. Review all MSDS for safe handling and disposal of chemicals.
Check for Understanding: Monitor and assist as student partners review and revise their notes.
Science Notebooks: 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 the class 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) • apron (1 per student) • glue or tape (per group) Attachments: • Handout: Enthalpy of Reaction (1 per
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student) Teacher Resource: Enthalpy of Reaction KEY
Safety Notes: Wear safety goggles. Review all of the MSDS sheets for safe handling and disposal of chemicals.
Science Notebooks: Students prepare a data table, record data and observations, and make calculations in their science notebooks.
EXPLAIN III – Calculating Enthalpy 1. Review the observations and results from the Handout: Enthalpy of Reaction. • How can energy transfer be observed in an exothermic (The temperature the system increases.) • reaction? How can energy transfer beof observed in an endothermic reaction? (The temperature of the system decreases.) 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) + H2O (l) Ca(OH)2 + 65.2 kJ. Ask: • 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) + H2O (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:
Suggested Days 4 (continued) and 5 Attachments: • Handout: Enthalpy of Reaction (from previous activity) Instructional Notes: 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. Use the Background Information to present simple endothermic and exothermic reaction pathways to students. Add in activation energy and catalysis diagrams as appropriate. Consider using Internet resources to help in presenting this section.
Science Notebooks: Students record definitions and make calculations in their science notebooks.
Given the equation, 2NaHCO3 (s) + 129 kJ
Na2CO3 (s) + H2O (g) + CO2 (g)
determine the kJ of heat needed to decompose 2.24 moles of sodium © 2008, TESCCC
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hydrogen carbonate. From the given, 129 kJ is needed to decompose 2 moles of solid NaHCO3, 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 NaHCO3 (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 below). Heat of Combustion: CH4 (g) + 2O2 (g) CO2 (g) + 2H2O (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 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 = ΔHof (products) – ΔHof (reactants) 2. Distribute the Handout: Enthalpy Change Calculations, and project a Thermodynamics Reference table (see Advance Preparation). Assist students as necessary. 3. Assign additional problems as needed from a locally adopted textbook and/or other resources.
Suggested Day 6 Attachments: • Handout: Enthalpy Change Calculations (1 per student) • Teacher Resource: Enthalpy Change Calculations KEY
STAAR Notes: The STAAR Chemistry Reference Materials includes the formula for Enthalpy of Reaction. Thermochemical equations in chemical reactions will be tested as a Readiness Standard under Reporting Category 4: Gases and Thermochemistry.
Science Notebooks: Students should work problems in their science notebooks.
EVALUATE – Performance Indicator © 2008, TESCCC
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Performance Indicator • Use calorimetry to measure an exothermic and endothermic enthalpy of reaction. Write a summary report that includes your procedures, data collected, calculations, a discussion of the reactions, and an error analysis. (C.2E, C.2F, C.2I; C.11C, C.11E) • 3D, 3E; 5B 1. Refer to the Teacher Resource: Performance Indicator Instructions KEY and the Handout: Laboratory Activity: Enthalpy of Reaction for information on administering the assessment.
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: • Teacher Resource: Performance Indicator Instructions KEY • Handout: Laboratory Activity: Enthalpy of Reaction (from previous activity)
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