25-1209 Examining Thermochemistry.Qxd

chemistry Examining Thermochemistry

Teacher’s Manual and Student Guide

251209 Examining Thermochemistry

CREDITS Chemistry Advisory Panel: Mike Jones Pisgah High School Haywood County School System, NC

Dr. Angela King Department of Chemistry Wake Forest University

Maureen Miller The Westminster Schools Atlanta, GA

Dr. Sam Powell Adjunct Professor of Chemistry (ret.) Alamance Community College, Graham, NC

Tonya Slawson Williams High School Alamance-Burlington School System, NC

Jane Smith Centennial High School Frisco, TX

David Vernon Western Alamance High School Alamance-Burlington School System, NC

Dr. Daniel Wright Department of Chemistry Elon University

Differentiated Instruction Reviewer: Marcee M. Steele Professor, Watson School of Education University of North Carolina Wilmington

Carolina Biological Supply Company Staff: Dr. Erin Krellwitz Writer/Developer

Robert Mize Editor

Charlie Johnson Designer

TABLE OF CONTENTS FOR TEACHER’S MANUAL OPTIONAL REPRODUCIBLE MASTERS

LEARNING GOALS PAGE 4 EXPERIMENTAL DESIGN TEMPLATE PAGE S-23 CONTENT STANDARDS PAGE 4 TIME REQUIREMENTS PAGE 5 MATERIALS PAGE 5 SAFETY PAGE 6 INSTRUCTIONAL APPROACH PAGE 6 SCIENCE CONTENT PAGE 9 PREPARATION PAGE 12 IMPLEMENTATION PAGE 13 HELPFUL HINTS PAGE 18 STUDENT MISCONCEPTIONS PAGE 19 DIFFERENTIATED INSTRUCTION PAGE 20 EXTENSION ACTIVITIES PAGE 20 RESOURCES PAGE 21 RELATED PRODUCTS PAGE 22 RUBRIC PAGE 22

TABLE OF CONTENTS FOR STUDENT GUIDE

ENGAGE PAGE S-1 EXPLORE PAGE S-2 ACTIVITY 1 PAGE S-2 ACTIVITY 2A PAGE S-4 ACTIVITY 2B PAGE S-6 ACTIVITY 3 PAGE S-7 EXPLAIN PAGE S-10 EXTEND PAGE S-19 4 Examining Thermochemistry

NOTES LEARNING GOALS Students will use a guided-inquiry technique (explained fully in the “Instructional Approach” section) to discover how chemical and physical changes can absorb and release heat energy. After measuring the heat of combustion for a common fuel and measuring the heats of solution for common salts, student groups will design a hot pack and a cold pack that could be used as a first aid device.

Students will

• develop the skills necessary to design and perform scientific investigations.

• write thermochemical equations for endothermic and exothermic reactions.

• interpret energy diagrams for endothermic and exothermic reactions.

• determine the heat of combustion for a common fuel.

• measure the heats of solution for common salts. • design and test a prototype of a hot pack or cold pack.

CONTENT STANDARDS This kit is appropriate for high school students and addresses the following National Science Education Standards:

Grades 9–12

Science as Inquiry • Abilities necessary to do scientific inquiry

• Understandings about scientific inquiry

Physical Science • Structure and properties of matter

• Chemical reactions

• Conservation of energy and increase in disorder

• Interactions of energy and matter

Science and Technology • Abilities of technological design

• Understandings about science and technology Examining Thermochemistry 5

TIME REQUIREMENTS NOTES

Preparation | 30 minutes

Engage | 10 minutes

Explore Activity 1 | 15 minutes Activity 2A | 20 minutes Activity 2B | 10 minutes Activity 3 | 45 minutes

Explain | 60 minutes (can be done as homework) Extend | 55 minutes

MATERIALS This kit contains materials for 30 students working in 10 groups of 3. Adjustments for group size may be made on the basis of class size.

Included in this kit: ammonium chloride, 500 g* calcium chloride, anhydrous, 500 g* 30 cups, styrene, 8 oz, with 15 lids 10 white candles* HeaterMeal®* 16 steel wool pads* 2 500-mL bottles of white vinegar* 100 pieces of filter paper, 9 cm diameter* Teacher’s Manual and reproducible Student Guide *Items included in the Examining Thermochemistry refill set (RN-251259)

Needed, but not supplied (for 10 groups) 10 Celsius laboratory thermometers or temperature probes 10 empty aluminum soda cans, 12 oz, with opening tab 10 graduated cylinders, 100 mL 10 250-mL beakers 10 400-mL beakers balance, 0.1 g sensitivity or greater 10 index cards, 3 × 5” 10 rubber bands 10 stirring rods or pencils 6 Examining Thermochemistry

NOTES 10 ring stands 10 iron rings access to water

Recommended for student-designed inquires: graph paper or a computer graphing program 2 1-qt resealable freezer bags 2 1-pt resealable bags

SAFETY Ensure that students follow safe laboratory practices when performing any activity in the classroom or lab. Demonstrate the protocol for correctly using the instruments and equipment necessary to complete the activities, and emphasize the importance of proper usage. Use personal protective equipment such as safety glasses or goggles, gloves, and aprons when appropriate. Model proper laboratory safety practices for your students, and require them to adhere to all laboratory safety rules. Clean all laboratory equipment after each use. Appropriate MSDS (Material Safety Data Sheets) are included in this kit.

Use caution when opening the heater bag of the HeaterMeal. The steam produced from the exothermic reaction can cause severe burns.

Allow the heater inside the heater bag to cool to room temperature before discarding it in the trash. Used heaters are EPA approved for discarding in regular trash.

Disposal Follow all local and state recommendations for chemical disposal. All chemicals used in this kit may be washed down the sink with copious amounts of water, except for the heater in the HeaterMeal, which must be discarded as a complete unit in the trash.

INSTRUCTIONAL APPROACH

About Inquiries in Science® Inquiries in Science is a series of kit-based lab activities that takes a guided-inquiry approach to teaching essential high school science topics. Designed around a learning cycle of engage, explore, explain, extend, and evaluate, each kit contains the necessary equipment and supplies for at least 30 students to perform the lab Examining Thermochemistry 7 “Men love to wonder, and that is the seed of science.” Ralph Waldo Emerson

activities. Each kit also includes a comprehensive Teacher’s Manual with a The underlying theme of the reproducible Student Guide. Inquiries in Science Chemistry series is “Properties, Changes, The topics comprising each major Inquiries in Science discipline have been organized and Interactions of Matter.” and selected on the basis of data gathered from numerous sources, including the As we interact with materials National Academy of Sciences’ National Science Education Standards, AAAS’s in our world, we question Project 2061 Science Benchmarks, NAEP’s Science Framework, and state science their properties and education standards from across the U.S. Compiling and analyzing information composition, and the way in from these sources led to the identification of fundamental learning goals within which they undergo physical each discipline—the essential science content that provides a solid foundation for and chemical changes. all high school students. The activities conducted in each Inquiries in Science kit provide students the opportunity to develop and improve their understanding of Evolving from such “how and key science topics through guided-inquiry experiences. Inquiries in Science kits are why” understandings, designed to assist and support teachers as they guide their students in mastering chemistry is now also an this essential knowledge—fostering science literacy by modeling the scientific applied science that has process. improved our standard of living through the creation of What is Inquiry? new materials, including pharmaceuticals, fertilizers, Inquiry is the process of observing, questioning, exploring, examining facts, and pesticides, plastics, and conducting a systematic investigation. In the discipline of science, inquiry is the way textiles. With these new in which information is gained, developed, revised, and verified. In essence, the materials comes the scientific process is inquiry. responsibility to ensure that In science education, it is critical that students understand the scientific process. The they do not harm our most effective way for students to gain this understanding is through inquiry-based environment or us. Thus, instruction, using the scientific process to gain knowledge and understanding of understanding chemistry is an scientific concepts, ideas, and events. Inquiry-based instruction creates a rich important part of being an learning environment; it actively engages students in the scientific process and helps informed citizen, and it is vital them internalize scientific concepts by doing science. for making good decisions Inquiry-based instruction can happen at many levels. At its simplest, it might take about the materials that the form of structured, step-by-step, teacher-derived questions and instructions that affect our health, our students follow to come to a predictable conclusion. At its most complex, it might take environment, our resources, the form of open-ended, long-term, student-developed and directed investigations. and our daily needs for food, clothing, and shelter. Inquiries in Science facilitates a guided-inquiry approach to science instruction, utilizing teacher-directed activities and student-centered investigations to help students grasp important science concepts by using the scientific process.

What is a Learning Cycle? A learning cycle is a series of steps by which people learn through experience as they build knowledge. The experiential, constructivist approach to education is founded on the ideas and research of many educational philosophers and practitioners, including John Dewey, Jean Piaget, David Hawkins, Robert Karplus, and David Kolb. 8 Examining Thermochemistry

A lesson designed around the stages of a learning cycle provides students the ENGAGE opportunity to explore a particular concept or topic, reflect on the experience, and then E v transfer the knowledge gained to a new or similar situation. Learning cycles can be a EXTEND constructed around different numbers of stages, but probably the most widely used l Real-World u learning cycle in science education is a five-stage cycle known as the 5E Instructional Situation

a EXPLORE Model, formulated by Roger Bybee for the development and promotion of the t e Biological Sciences Curriculum Study (BSCS). The stages of this learning cycle are engagement, exploration, explanation, elaboration, and evaluation. EXPLAIN Inquiries in Science has been developed around a learning cycle adapted from the BSCS 5E Instructional Model. This learning cycle includes four cyclical stages, engage, explore, explain, and extend, and a fifth stage, evaluate, which is integrated throughout. Following is a description of each step of the Inquiries in Science learning cycle.

Engage: During the Engage step, the teacher helps students consider their prior knowledge about a particular topic and relate that knowledge to questions or problems in the context of a real-world scenario. In this stage, the teacher can evaluate the students’ current understanding of the topic and identify the misconceptions that students bring with them.

Explore: During the Explore component of the learning cycle, students gain hands-on experience relevant to the question or problem at hand. The teacher guides students through investigations and models the scientific process using a guided-inquiry approach. The teacher evaluates understanding during this stage by observing students’ progress and their interactions with their teammates.

Explain: In the Explain portion of the learning cycle, students develop and fine- tune their grasp of scientific concepts by analyzing, interpreting, and communicating data. They may assimilate information from additional sources (selected readings, for example). The teacher facilitates and advises, guiding students in their development and helping them confront misconceptions. During this stage, the teacher evaluates students’ understanding of the topic through their answers to critical-thinking questions and their analysis of the data gathered during the Explore stage. Then, the teacher knows whether the students are ready to apply knowledge of the topic during the Extend stage.

Extend: In the Extend stage of the learning cycle, students demonstrate their mastery of the concepts they have studied by applying what they have learned in the context of the original real-world question or problem. The teacher observes, assesses, and directs the students’ progress toward the desired learning goals. Evaluation at this stage generally involves assessing students’ design and completion of additional experiments, or their application of science concepts to another situation. Examining Thermochemistry 9

SCIENCE CONTENT IMPORTANT TERMS Many physical changes involve release and absorption of heat. Examples include the calorimeter dissolving of salts, known as heat of solution, changes in state (melting and freezing, evaporation and condensation), and the compression and expansion of gases. ΔH In addition, almost every chemical change either releases or absorbs heat as chemical bonds are broken and formed during the course of a reaction. Energy is endothermic consumed when bonds are broken, and energy is released as new bonds are formed. The net total energy change for a reaction will either be a negative value exothermic with the release of heat (defining an exothermic reaction) or a positive value with the absorption of heat (defining an endothermic reaction). heat The study of the heat energy associated with chemical changes is known as thermochemistry. Changes in heat energy are measured in a device called a heat of combustion calorimeter. Heat flow is measured indirectly, by placing the material to be measured (the system) in a reaction chamber that is surrounded by water (the heat of formation surroundings). As the material reacts in the chamber, heat may be transferred to the water or absorbed from the water. By monitoring the temperature change of the heat of solution water, one can tell whether the reaction is endothermic (water temperature goes down) or exothermic (water temperature goes up). specific heat Heat and temperature are often perceived to be the same quantity, but they are not. capacity Temperature is a property determined by the average kinetic energy of the particles in matter and can be measured temperature Table 1: Specific Heats at 25°C (298 K) directly with a thermometer. Substance Specific Heat (J/g•°C) Temperature measurements in thermochemistry Aluminum (s) 0.897 thermochemistry are in degrees Ammonia (g) 2.090 Celsius or in the absolute temperature scale of Kelvin. Heat Calcium (s) 0.647 is defined as the spontaneous flow Carbon, graphite (s) 0.709 of thermal energy from an object Copper (s) 0.385 at a higher temperature to one at Ethanol (g) 1.420 a lower temperature. The unit for Ethanol (l) 2.440 heat is calories in the English Gold (s) 0.129 system and joules in the SI system. One calorie = 4.18 J. The Law of Iron (s) 0.449 Conservation of Energy states that Lead (s) 0.128 in an insulated system, the heat Mercury (l) 0.140 lost by one material will equal the Water (g) 1.870 heat gained by another material. Water (l) 4.180 The amount of heat transferred between two substances depends Water (s) 2.060 upon the temperature difference, the mass of the materials, and the specific heat capacities of the materials. The specific heat capacity is the amount of heat required to raise one gram of a substance 1°C. The units for specific heat capacity are calories per gram per degree Celsius (cal/g•°C) in the English system, or joules per gram per degree Celsius (J/g•°C) in the SI system. Table 1 lists the specific heats of some common materials at 25°C or 298 K. 10 Examining Thermochemistry

NOTES The equation that links specific heat, the mass of material, and the change in temperature for calculating the heat absorbed or released by a material is as follows: Δ q = (m)(cp)( T), where q is the heat transferred, cp is the specific heat capacity at constant pressure, m is the mass of the material, and ΔT is the change in temperature. The value ΔT is found by subtracting the initial temperature of the

material from the final temperature (Tf – Ti). If the sign of q is negative, heat will be lost by the system (exothermic) and if the sign is positive, heat will be gained by the system (endothermic). Look at the sample problem below for using this equation: How much heat is required to raise 10.0 g of copper from 20.0°C to 50.0°C? Δ q = (m)(cp )( T) q = (10.0 g)(0.385 J/g•°C)(50.0°C – 20.0°C) q = (10.0 g)(0.385 J/g•°C)(30.0°C) q = 116 J

Heat of Reaction for Chemical Changes The energy absorbed or released during a chemical reaction is termed the heat of reaction. Exothermic reactions spontaneously produce products that have lower energy than the reactants and, as a result, release energy. An example would be the combustion of carbon to form carbon dioxide. C(s) + O2(g) CO2(g) + 393.5 kJ Note that the equation for this reaction has heat as a product, which indicates that

393.5 kJ will be released per mole of CO2 formed. Another way of expressing the heat of this reaction is the term ΔH, which means change in enthalpy or heat. The symbol H represents the enthalpy or the heat content of a system at constant pressure. The heat content (H) of a system cannot be measured directly, but the change in enthalpy can. Enthalpy change is signified by ΔH. The Greek letter delta (Δ), meaning “change,” is followed by H, for “heat.” By convention, all exothermic reactions will have a negative sign, indicating a loss of energy. Mathematically, the ΔH can be expressed as follows: Δ H = Hproducts – Hreactants

Figure 1 shows the energy diagram for this exothermic reaction.

Energy Reactants

ΔH = – value Products

Reaction Pathway

Figure 1 Examining Thermochemistry 11

If we reverse the CO2 equation, then energy would have to be absorbed to break NOTES the bonds of the CO2 molecule. The heat (393.5 kJ) appears as a reactant in this thermochemical equation. 393.5 kJ + CO2(g) C(s) + O2(g)

The ΔH will be positive because heat is added in this reaction, giving the products more energy than the reactants. Figure 2 shows the absorption of heat energy for this reaction.

Energy Products

ΔH = + value Reactants

Reaction Pathway

Figure 2

Heat of Formation A heat of reaction in which a mole of a compound is produced from its elements is called heat of formation, or molar heat of formation. The values for heats of formation are usually measured at their standard states and at standard conditions (1 atmosphere of pressure and room temperature, considered to be 25°C, or 298 K).

For example, the standard state of H2O at 1 atmosphere and 25°C is the liquid state. The degree symbol (°) in enthalpy symbols indicates that the measurements are made for standard states at standard conditions. The subscript f (f) indicates that the value is for the heat of formation. By definition, the heat of formation for a Δ pure element is zero ( H°f = 0). One can use a table of heats of formation to find the value for many compounds. The more negative the enthalpy value, the more stable the compound since energy was lost as it was formed from its elements. Compounds with positive enthalpies are unstable and will often decompose back into their elements.

Calculating the heat of a reaction from molar heats of formation One can calculate the heat of a reaction for a chemical change by referring to a standard thermochemistry table to find the heat of formation for the various compounds in the reaction. For example, calculate the heat of reaction for the following chemical change and predict whether it is exothermic or endothermic: CO2(g) + 2H2O(l) CH4(g) + 2O2(g) 12 Examining Thermochemistry

Δ Hint: Coefficients greater than one must be multiplied by the H°f value to determine the total amount of heat for that given number of moles.

Heats of formation for each compound taken from a standard thermochemistry table are as follows: Δ Δ CO2(g): H°f = –393.5 kJ/mole H2O(l): H°f = –285.8 kJ/mole Δ Δ CH4(g): H°f = –74.9kJ/mole O2(g): H°f = 0.0 kJ/mole Calculation:

Δ ΣΔ ΣΔ Σ means “sum of” H° = H°f of products – H°f of reactants Δ Δ Δ Δ Δ H° = { H°f CH4 + 2[ H°f O2(g)]} – { H°f CO2(g) + 2[ H°f H2O(l)]} ΔH° = [(–74.9 kJ/mole) + 2(0.0 kJ/mole)] – [–393.5 kJ/mole + 2(–285.8 kJ/mole)] ΔH° = (–74.9 kJ/mole) + (965.1 kJ/mole) ΔH° = 890.2 kJ/mole

The reaction is endothermic due to the positive sign.

Heat of Combustion Another type of heat of reaction is heat of combustion, in which a substance reacts with oxygen gas to produce energy, carbon dioxide, and water. The combustion of fossil fuels is an exothermic reaction that provides civilization with most of its current energy needs. The burning of propane gas in a gas grill is a good example of a combustion reaction.

C3H8(g) + 5O2(g) 3CO2(g) + 4H2O(l) + 2219 kJ ΔH° = –2219 kJ/mol

PREPARATION 1. Familiarize yourself with the content of the Student Guide, including the activity instructions and assessments. 2. Collect 10 empty, 12-oz aluminum soda cans, with the opening tab still intact. Rinse out any remaining soda with water and allow the cans to dry overnight. 3. Take five steel wool pads and divide them equally into halves. 4. Set up 10 student lab stations. Each station should have the following materials: 2 styrene foam cups, 8 oz white candle

1/2 of a steel wool pad Examining Thermochemistry 13

laboratory thermometer or temperature probe aluminum soda can, 12 oz, with tab graduated cylinder, 100 mL beaker, 250 mL beaker, 400 mL index card, 3 × 5” rubber band stirring rod or pencil ring stand iron ring access to water 5. Set up a weighing station with the following materials for groups to share:

balance, 0.1 g sensitivity or greater pack of filter paper, 9 cm ammonium chloride, 500 g calcium chloride, 500 g 2 bottles of white vinegar, 500 mL 6. Photocopy the Student Guide for each student in the class. At your discretion, photocopy the Experimental Design Template (one for each group) for students to use in planning Activity 4, the Extend activity.

IMPLEMENTATION

Engage ENGAGE Students focus on these inquiry activities by considering how hydrocarbon Help students combustion reactions supply a major portion of our energy needs. They then consider current and potential uses for thermochemical reactions. consider their prior

Lead the class in a brainstorming session about alternative chemical reactions that knowledge about a could be used for producing energy in lieu of hydrocarbon combustion. Have them topic and relate list any commercial products that make use of exothermic or endothermic chemical that knowledge to reactions and explain how they are used. questions or problems in the context of a real-world scenario. 14 Examining Thermochemistry

EXPLORE Explore Students gain The Explore section of this lesson gradually increases the level of student inquiry. hands-on • In Activity 1, student groups calculate the heat of solution for an endothermic experience relevant and exothermic salt as each is dissolved in 100 grams of water in a “coffee cup” calorimeter. to the question or problem at hand. • In Activity 2A, students determine whether the oxidation of iron from a steel wool pad soaked in vinegar is endothermic or exothermic. Guide students • In Activity 2B, students observe a demonstration of how a small precooked meal through can be heated from an oxidation reaction by 8 grams of powdered magnesium investigations and iron in a water-absorbent pouch. If you allow your students to eat a portion and model the of this warmed meal, make sure there are no students with food allergies for the scientific process. listed ingredients. • In Activity 3, students calculate the heat of combustion for a hydrocarbon fuel, paraffin wax, and write a balanced equation for the exothermic reaction including the heat of combustion as a product.

Note that in the Extend activity, students design either a hot pack or cold pack for treating injuries such as joint sprains. Lab groups first design an experiment to test 5-g increments of a compound dissolved in 100 g of water, to enable them to predict how much will be required to heat 100 g of water to a temperature of 55°C or cooled to a temperature of 3°C. Then, lab groups collaborate and build a prototype of one hot pack and one cold pack for the entire class. They activate each one as a class demo, measure the highest and lowest temperatures reached, and then compare these to the target maximum and minimum temperatures.

Answers to Questions in the Explore Activities

Activity 2B

1. Why does this 8-g absorbent pad produce higher temperatures than the steel wool pad (98%–99.5% iron) of about the same mass did in Activity 2A? Hint: Compare the location of Mg and Fe on the periodic table. Magnesium is a more active metal than iron, and the magnesium is powdered, providing more surface area for reaction with saltwater.

2. The magnesium metal in the pad reacts with water to produce a magnesium hydroxide compound and a flammable gas, with the release of heat energy. Write a balanced chemical equation for this reaction. Mg + 2H2O Mg(OH)2 + H2 + heat

3. Why use saltwater rather than freshwater for activating the pad? Ions in the saltwater react with the metals, speeding up the corrosion reaction. Examining Thermochemistry 15

Activity 3

1. Can one assume that all of the heat from the burning candle is absorbed by the aluminum can and the water inside it? No. Some escapes to the surrounding air.

2. Why is it necessary to stir the water in the can while monitoring the temperature change of the water? To distribute the heat evenly to all of the water in the can, which will reduce error in temperature monitoring.

3. Why must you continue to monitor the temperature of the water even after the flame is extinguished? Heat is still being transferred to the water from the can.

4. Write a complete balanced equation for this reaction and include your experimental value for the heat of combustion. C25H52 +38O2 25CO2 + 26H2O + ______J

Activities 1 through 2B can be completed in one standard class period (45 minutes needed). Activity 3 requires 45 minutes.

Move around the room, making sure that students are on task and understand exactly what each activity requires. Encourage students to think independently.

Explain EXPLAIN Use the Explain section as it works best for your class. You may read it together as Students develop a class or have students read it on their own. It may be assigned as a post-activity reading selection or as homework. Higher-level students may be able to answer the and fine-tune their questions at the end of the Explain section without reading the selection. grasp of scientific

You may choose to have a class discussion at the end of each activity or wait until concepts by all of the activities are completed. analyzing, interpreting, and Answers to Questions communicating 1. Compare the heat content of the products to that of the reactants in the data. Advise following two types of reactions and comment on the chemical stability of the product(s). students and help a. exothermic – The heat content is lower for the products, and products are them confront more stable. misconceptions. b. endothermic – The heat content is higher for the products, and products are less stable. 16 Examining Thermochemistry

NOTES 2. Explain the difference between heat of formation, heat of reaction, and heat of combustion. Heat of formation is the amount of energy released or absorbed in a reaction in which a mole of a compound is produced from its elements. Heat of reaction is the amount of energy released or absorbed during a chemical reaction. Heat of combustion is the amount of heat released when a substance reacts with oxygen gas.

3. How much heat is required to change 20.0 g of gold from 25.0°C to 100.0°C? Δ q = mcp T q = (20.0 g)(0.129 J/g•°C)(75.0°C) q = 194 J

4. A mole of powdered magnesium (24.3 g) when mixed with water will produce enough heat to raise 1 L of water from 25°C to boiling. Calculate the heat in kJ that would be produced. 1 L = 1000 mL = 1000 g Δ q = mcp T q = (1000 g)(4.18 J/g•°C)(100°C – 25°C) q = 313,500 J q = 313.5 kJ

5. Blocks of iron heated in a fireplace were often used as bed warmers during the colonial period of American history. How much heat in joules would be lost from a 1.50-kg block of iron that started at 150.0°C and cooled down to 20.0°C? 1000 g Convert 1.50 kg to g: 1.50 kg × = 1500 g 1 kg Δ q = mcp T q = (1500 g)(0.449 J/g•°C)(20.0°C – 150.0°C) q = (1500 g)(0.449 J/g•°C(–130.0°C) q = –87,555 J

6. A 120-g sample of copper is exposed to 20.0 kJ of heat in an insulated chamber. By how many degrees Celsius will the copper sample’s temperature increase? 1000 J First, change 20 kJ to J by the following calculation: 20.0 kJ × = 20,000 J 1 kJ Δ q = mcp T 20,000 J = (120 g)(0.385 J/g•°C)(ΔT) 20,000 J = (46.2 J/°C)(ΔT) 20,000 J = ΔT 46.2 J/°C 433°C = ΔT Examining Thermochemistry 17

7. a. Calculate the enthalpy, or heat of solution (ΔH), in kJ/mole for NaOH, given NOTES the following data: 5.0 g dissolved in 100 mL of water raised the water temperature in an insulated cup from 23.0°C to 36.0°C. 5.0 g moles of NaOH = = 0.125 moles NaOH 40.0 g/mole Assume that 100 mL water = 100 g Δ q = –[mcp T]* q = –[(100 g)(4.18 J/g•°C)(36.0°C – 22.6.0°C)] q = –[(100 g)(4.18 J/g•°C)(13.4°C)] q = –5601 J q = –5.60 kJ –5.60 kJ ΔH = = –44.8 kJ/mole 0.125 moles *Since NaOH loses heat to water, use a minus sign.

b. Is this physical change endothermic or exothermic? It is exothermic, as indicated by the negative sign.

8. For the following reactions and their given ΔH° values, write the equation with the enthalpy value on the product or reactant side, and identify the reaction as endothermic or exothermic. Δ a. CO2(g) + 2H2O(l) CH4(g) + 2O2(g) H° = +891 kJ 891 kJ + CO2(g) + 2H2O(l) CH4(g) + 2O2(g) endothermic

Δ b. 2Mg(s) + O2(g) 2MgO(s) H° = –1200 kJ 2Mg(s) + O2(g) 2MgO(s) + 1200 kJ exothermic

Δ c. 2NO2(g) 2NO(g) + O2(g) H° = +114 kJ 114 kJ + 2NO2(g) 2NO(g) + O2(g) endothermic

9. Ethanol, or ethyl alcohol, is the major component of E85 gasoline (85% ethanol and 15% gasoline). Write the balanced equation for ethanol combustion and calculate the heat of combustion for one mole of ethanol (C2H5OH) using the following standard heats of formation: Δ Δ CO2(g): H°f = –393.5 kJ/mole H2O(l): H°f = –285.8 kJ/mole Δ Δ C2H5OH(l): H°f = –277.0 kJ/mole O2(g): H°f = 0.0 kJ/mole C2H5OH(l) + 3O2(g) 2CO2(g) + 3H2O(l) Δ ΣΔ ΣΔ H° = H°f of products – H°f of reactants Δ Δ Δ Δ Δ H° = {2[ H°f CO2(g)] + 3[ H°f H2O(l)]} – { H°f C2H5OH(l) + 3[ H°f O2(l)]} ΔH° = [2(–393.5 kJ/mole) + 3(–285.8 kJ/mole)] – [(–277.0 kJ/mole) + 3(0.0 kJ/mole)] ΔH°= [(–787.0 kJ/mole) + (–857.4 kJ/mole)] – (–277.0 kJ/mole) ΔH° = –1644.4 kJ/mole + 277.0 kJ ΔH° = –1367.4 kJ/mole 18 Examining Thermochemistry

10. One type of cold pack has 100 g of water that mixes with 20 g of NH4NO3. If the heat of solution for NH4NO3 is 25.4 kJ/mole, what would be the final temperature of the bag in degrees Celsius if the initial temperature were 23°C? 20.0 g moles of NH4NO3 = = 0.25 moles of NH4NO3 80.0 g/mole

(0.25 moles NH4NO3)(25.4 kJ/moles) = 6.35 kJ = 6350 J Δ q = mcp T 6350 J = (100 g)(4.18 J/g•°C)(ΔT) 6350 J = ΔT 418 J/°C 15.2°C = ΔT Final Temperature = 23°C – 15.2°C = 7.8°C

EXTEND Extend Students Student lab groups will be assigned the task of designing a portable hot pack or cold pack for treating injuries. Each group will estimate how much of a chemical is needed demonstrate to mix with 100 g of water to make a hot pack reach a temperature of 55°C or a cold their mastery of pack reach 3°C. After each group predicts the amount of chemical needed, they confer concepts. Observe, with other groups that had the same pack assignment and arrive at a consensus for the assess, and direct amount of chemical needed. These groups collaborate on their pack design and build one class prototype. Each type of pack is activated in a class demo, and the temperature the students’ is monitored to see whether the designs and predictions were satisfactory. progress toward the learning goals.

HELPFUL HINTS • If your copy paper is rationed, you may get by with making one copy per student group of the Engage and Explore sections of the Student Guide and letting the team members share those portions.

• In order to keep all group members engaged in the lab activities, you may want to delegate or have students volunteer for the following job duties in each lab group.

Materials Manager—responsible for assembling apparatus, massing solutes, and preparing solutions. Data Manager—responsible for performing the experiment, taking measurements, and doing calculations. Recorder—responsible for recording data and calculations provided by the Data Manager. Examining Thermochemistry 19

• To facilitate sharing and easy access, choose a centrally located table or counter NOTES for the weighing station, which would include all chemicals, filter paper for massing, and one or more electronic or mechanical balances.

• Students may assume that for water, 1 mL = 1 g. Therefore, 100 mL of water measured with a graduated cylinder will have a mass of 100 g.

• When measuring a change in temperature for a system (ΔT), remind students Δ always to subtract in this order ( T = Tf – Ti), where Tf is the final temperature and Ti is the initial temperature. • In Activity 3, the change in temperature for the aluminum can calorimeter will be the same as the change in temperature for the 100 g of water inside the can.

• In Activity 3, the reason for using chilled water (10–15°C below room temperature) and then heating the water to that same amount above room temperature is to cancel out any error from the transfer of heat to and from the surrounding room air.

STUDENT MISCONCEPTIONS • Students often think that heat and temperature are the same quantities. Use the analogy of a burning log and a burning match. Both can be at the same temperature, but which one would give you the worse burn if you touched it?

• Students may think that cold can be transferred from a cooler object to a warmer one. Heat is always transferred from warmer to cooler objects.

• Students may believe that heat and cold can be exchanged at the same time between two or more objects.

• Students may assume that different materials in the same environment can have different temperatures. For example, metals feel colder to the touch than wood. Actually, metals feel colder because they are good conductors of heat and they remove heat from your hand when grasped.

• Students may believe that energy can be lost or destroyed, such as when a pot of boiling water cools to room temperature. In reality, the heat lost by the pot is transferred to the surrounding air. Energy is always conserved in the universe. 20 Examining Thermochemistry

NOTES DIFFERENTIATED INSTRUCTION • Have English Language Learners and lower level students create flash cards of the Important Terms listed in the margin of the “Science Content” section. They could write a term on one side and its definition on the other side. These cards could also be used for review sessions and remediation.

• Have your students who are poor readers make charts, concept maps, Venn diagrams, or other graphic organizers to reinforce concepts such as conservation of energy, energy flow, energy diagrams for exothermic and endothermic reactions, and schematics for hot pack and cold pack designs.

• Assign your most tactile learners to the position of “Materials Manager” to assemble and monitor lab equipment for each activity. For the other two group positions, make assignments according to students’ abilities for reading, writing, math, art, and communication.

• Advanced learners could create a PowerPoint® presentation on energy, energy flow, energy diagrams, and schematics for hot pack and cold pack prototypes. • Guide and assist the class on heat-of-reaction calculations for chemical changes.

EXTENSION ACTIVITIES 1. If you have temperature sensor probeware, have the students measure temperature and graph their results in real time for all activities. As the temperature/time graph is being drawn, they can interpret the data and the relationship between both variables.

2. Students can measure the calories per gram of peanuts, cashews, potato chips, or corn chips. Be aware of any students’ food allergies before approving this experiment. They can use the same setup as in Activity 3, using the aluminum soda can as a calorimeter. They mass the food item and place it on a stand made of a bent paper clip with one end embedded in a cork. Then, they ignite the food with a match or Bunsen burner and allow it to burn to completion. The heat will be transferred to the 100 g of water in the can. Hint: 1 food calorie (1 C) is 1000 cal (1 kcal) of heat. The specific heat capacity of water is 1.0 cal/g•°C in the English System. Compare your calculation, q = (100 g)(1.0 cal/g•°C)(ΔT) ÷ 1000 cal/Cal, with the calories/g on the wrapper of the food item.

3. Which candle wax is the best fuel? Students can measure the heats of combustion for other types of candles such as beeswax or soy wax and compare these with the heat of combustion of paraffin. Examining Thermochemistry 21

4. Students can research Flameless Ration Heaters (FRH) and do a written and/or oral report on how these heaters work to heat a precooked meal to 140°F.

5. Students can do research on how hot packs and cold packs work and the difference between one-time-use hot packs and reusable hot packs.

RESOURCES

Web Sites ADDITIONAL WEB SITES At the time of this printing, the following Web sites were active. You may wish to perform an independent search for similar sites.

Thermochemistry. ©Erik Epp. This link is a concise summary of thermochemistry. http://eppe.tripod.com/thermchm.htm

The Transfer of Energy 1: Thermochemistry. ©AAAS. This site explains heat and temperature, energy transfer, and energy conservation. Student misconceptions about these topics are also listed. http://www.sciencenetlinks.com/lessons.cfm?DocID=401

Thermochemistry and Calorimetry. ©Stephen Lower. This site covers thermochemistry, enthalpy, and calorimetry and gives examples of calculations. There is an excellent concept map at the end of this section. http://www.chem1.com/acad/webtext/energetics/CE04.html ADDITIONAL PRINTED MATERIAL

Daley, H. 1994. Heat of solution: hot packs. Journal of Chemical Education, 71 (9): 791.

Marsella, G. 1987. Hot and cold packs. ChemMatters, (2): 7.

Shakhashiri, B. 1983. Chemical Demonstrations: A Handbook for Teachers of Chemistry. University of Wisconsin Press: Madison, WI. 22 Examining Thermochemistry

WISH LIST RELATED PRODUCTS Following is a list of related items available from Carolina Biological Supply Company. For more information, please refer to the most recent CarolinaTM Science catalog, call toll free 800-334-5551, or visit our Web site at www.carolina.com.

RN-753565 Measuring Energy Efficiency

RN-753490 Specific Heat Specimen Set

RN-753518 Double-Wall Calorimeter

RN-753526 Food Calorimeter

RN-753700 Heat Transfer Kit RN-753704 Ice Melting Kit

RUBRIC On the following page is an optional rubric for evaluating the student-created inquiries. For students who are new to inquiry or who need extra guidance, an optional reproducible master Experimental Design Template is included following the Student Guide. If students need suggestions, inquiry possibilities include the following:

• How much chemical is needed for the hot pack or cold pack? After having measured the temperature change for increments of 5 g per 100 g of water, students should see the relationship of chemical to change in temperature. If they plot temperature versus grams, they should get a graph that is linear, which they can “extrapolate” to get the predicted amount of chemical needed for the desired final temperature.

• For the one-container design, the water can be placed in a smaller resealable bag and then placed inside a larger resealable bag containing the dry compound. Activation is accomplished by breaking the inner bag to release the water to mix with the compound.

• Other chemicals that could be tested for their use as a hot pack or cold pack

include Epsom Salts (MgSO4•7H2O), NH4NO3, KNO3, NaNO3, Na2CO3. • Students may also vary the amount of water from the suggested 100 g for the hot pack and cold pack prototypes and note any changes in the final temperature.

• Can these one-time-use hot and cold packs be modified for reuse? Could the water be evaporated and the solute recovered and reused? Examining Thermochemistry 23

NOTES TOPIC 1 2 3 4

Question Student did Student Student identified Student not identify identified a a testable independently a question. question that question, with developed a cannot be teacher relevant question tested. assistance that can be tested.

Hypothesis Student did Student Student Student not state a developed a developed a independently hypothesis. hypothesis that relevant developed a did not relate to hypothesis, hypothesis that the question. with teacher pertains directly assistance. to the question.

Procedure Student did Student Student Student not developed a developed a independently develop a procedure that procedure that is developed a logical lacked details sensible and procedure that procedure. and sequence. easy to follow, clearly outlined with teacher each step needed assistance. to complete the activity.

Data No data Data was Data was Data was was incomplete. complete. complete, logically collected. Graphs and Graphs and organized, and tables were tables were well presented. unlabeled or legible. Scale Graphs and tables difficult to read; was not were clearly and scale was consistent with accurately labeled, inconsistent with collected data. and the scale data. matched collected data.

Conclusion No A conclusion A detailed Student conclusion was reached conclusion was independently was with teacher reached with reached a reached. assistance. The teacher detailed conclusion did assistance. conclusion that not include pertained to the reference to the original question original question and hypothesis. or hypothesis. Examining Thermochemistry

Student Guide 1 Engage 2 Explore 3 Explain 4 Extend Examining Thermochemistry | STUDENT GUIDE S-1

1 Engage

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Energy is one of the most critical concerns in our society today. Cost and availability are two important considerations for our society as we search for sources of energy to fuel our industry, provide heat and electricity for our homes, and power our transportation. Much of our current energy comes from the combustion of fossil fuels. As supplies of fossil hydrocarbons decrease worldwide, we will likely have to find other sources of energy that can be produced from similar exothermic chemical reactions. Ideally, these reactions will be more “green,” emitting a smaller quantity of toxins and greenhouse gases.

Both endothermic reactions, which absorb heat, and exothermic reactions, which release it, offer great potential for further invention. For example, portable cooling or heating devices activated by thermochemical reactions would be great in areas with no electricity. An understanding of thermochemistry may help you find new solutions to the age-old problem of producing enough energy for a safe and comfortable environment.

PRIOR KNOWLEDGE 1. Brainstorm with your partners and list some chemical reactions that could produce energy using fuels other than fossil fuels such as coal, gasoline, natural gas, and propane. 2. List commercial products that make use of exothermic or Engageendothermic chemical reactions, and explain their use.

2 Explore

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Activity 1. Heat of Solution MATERIALS

In this activity you will calculate the heat of solution for a physical 400-mL beaker change that involves the heat released or absorbed by a salt as it 2 styrene cups dissolves in water. You will measure this indirectly by recording the maximum increase or decrease in the temperature of 100 g of water graduated cylinder, in a coffee cup calorimeter at constant atmospheric pressure. The 100 mL Δ heat flow for the water is calculated by the formula q = mcp T,

where q = heat, m = mass of water, cp = specific heat of water laboratory Δ (4.18 J/g•°C), and T = change in temperature (Tf – Ti). The enthalpy thermometer or change for the reaction ΔH is equal but opposite in sign to the heat temperature Δ flow to or from the water. You will use the formula, H= –(qwater)or probe Δ Δ H = –(mcp T). If the water temperature increases, the compound is exothermic, releasing heat into the water. If the water temperature access to decreases, the compound is endothermic, absorbing heat from the ammonium water. The heat of solution per gram is calculated by taking the total chloride heat (q) and dividing by the number of grams dissolved. The sign of the answer will indicate which type of reaction it is: (–) for exothermic, and access to calcium (+) for endothermic. chloride access to water HEAT OF SOLUTION FOR CALCIUM CHLORIDE 1. Place one styrene cup inside the other and then place both cups filter paper inside a 400-mL beaker. This will be your calorimeter for balance, 0.1 g measuring changes in temperature. sensitivity or 2. Measure 100 mL of water with a graduated cylinder and place better it in the calorimeter.

3. Place a laboratory thermometer in the water and monitor the temperature until it remains constant. Record the temperature in Data Table 1.

4. Tare a piece of filter paper and mass 5.0 g of calcium chloride. The mass does not have to be exactly 5.0 g, but must be accurate to within 0.1 g.

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Activity 1. (continued) 6. While stirring, monitor the change in temperature until the temperature stops changing.

7. Record the final temperature in Data Table 1.

8. Calculate the heat lost or gained from the water by using the Δ formula q = mcp T. The specific heat capacity, cp, of water is 4.18 J/g•ºC. The change in temperature can be represented by the Δ equation T = Tf – Ti (final temperature – initial temperature).

9. Calculate the heat of solution in kJ/g by taking the calculated heat (q), dividing by 5.0 g, and multiplying by 1 kJ/1000 J to convert to kJ/g.

10. Is the heat of solution exothermic or endothermic?

HEAT OF SOLUTION FOR AMMONIUM CHLORIDE

Repeat steps 1–10 above for ammonium chloride (NH4Cl) and record your data in Data Table 1.

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Activity 2A. Heat of a Reaction MATERIALS

In this activity, you will determine whether the oxidation of iron into 250-mL beaker iron(III) oxide (Fe2O3) is an endothermic or exothermic chemical half of a steel reaction. To speed the reaction rate, you will soak the steel wool in wool pad 50 mL of vinegar for a couple of minutes. The acetic acid in the vinegar helps remove the coating on the surface of the filaments, (approximately 8 g) and the acidity increases the rate of oxidation of the iron in the vinegar, 50 mL filaments. The reaction occurs as follows: laboratory 4Fe(s) + 3O (g) 2Fe O (s) 2 2 3 thermometer or Measuring the initial temperature and the final temperature of the temperature pad will help you decide the type of reaction that you have (in terms probe of whether heat is released or absorbed). rubber band 1. Measure the room temperature with a laboratory thermometer and record it in Data Table 2. That should be the same as the 2 styrene cups, temperature of the pad. with lid 2. Add 50 mL of vinegar to a 250-mL beaker. 400-mL beaker 3. Immerse the steel wool pad by pressing down with a glass glass stirring rod stirring rod and keep submerged for 2 minutes.

4. Squeeze all of the vinegar out of the pad back into the beaker.

5. Wrap the steel wool around the bottom portion of the thermometer and secure it with a rubber band.

6. Place the pad/thermometer assembly inside a styrene cup and place the cup inside a 400-mL beaker.

7. Center the lid over the top of the thermometer, punch a hole in the center of the lid, and slide it down the thermometer until it fits over the mouth of the cup.

8. Monitor the temperature until it reaches a maximum or minimum and record that temperature in Data Table 2.

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Activity 2A. (continued) 10. Based upon the temperature change, is the oxidation of iron endothermic or exothermic? Record your answer in Data Table 2.

11. Remove the lid from the cup. In Data Table 2, record a description of the pad.

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Activity 2B. Commercial Application for Metallic Oxidation MATERIALS ® This activity will be a demo by your teacher, showing you how a small HeaterMeal meal can be heated by the oxidation of a metal (a chemical reaction commonly known as corrosion). The heating unit is a small water- absorbent pad filled with 8 grams of powdered magnesium with a small amount of powdered iron acting as a catalyst. The pad is activated by the addition of 2 fl oz of saltwater to the plastic bag that contains the pad. The pre-cooked food in the tray positioned over the pad is heated by this “super corrosion” reaction, creating a ready-to-eat hot meal in about 10 minutes.

QUESTIONS TO CONSIDER FOR 2B: 1. Why does this 8-g absorbent pad produce higher temperatures than the steel wool pad (98–99.5% iron) of about the same mass did in Activity 2A? Hint: Compare the location of Mg and Fe on the periodic table.

3. Why use saltwater rather than freshwater for activating the pad?

2 Explore

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Activity 3. Heat of Combustion MATERIALS

In this activity you will determine the heat of combustion of a 12-oz aluminum

hydrocarbon fuel, paraffin wax, C25H52. All hydrocarbons when soft drink can burned produce carbon dioxide, water, and energy as shown by the with pull tab unbalanced generic equation below: glass stirring rod or CxHy + O2(g) CO2(g) + H2O(g) + energy pencil This chemical reaction is the basis for meeting most of the world’s ring stand energy needs. By burning a given amount of fuel and measuring the heat transferred to a calorimeter filled with water, one can calculate iron ring the heat of combustion for the fuel in kJ/g. The calorimeter in this exercise is an aluminum soft drink can containing 100 g of water. graduated cylinder, 100 mL 1. To make a candle stand, light your candle and hold the burning tip horizontally over the middle of an index card to produce a index card puddle of wax as wide as the base of the candle. Extinguish the paraffin candle flame and seat the candle’s base firmly into the melted wax.

2. Weigh the candle/card assembly and record the value in Data Celsius laboratory Table 3. thermometer or temperature 3. Measure 100 mL of chilled probe water (10–15°C below room temperature) in a graduated balance, 0.1 g cylinder. Carefully pour this sensitivity or water into an empty soft drink greater can. Hint: 1 mL of water = 1 g. chilled water 4. Set up the apparatus shown in Figure 1. The stirring rod or (10–15°C pencil passes through the tab below room of the can and rests on top of temperature) the iron ring. Figure 1

5. Adjust the can so that the bottom is about 2 cm above the candle wick.

6. Before lighting the candle, measure the room temperature and initial water temperature in the can to the nearest 0.1°C. ExploreRecord these values in Data Table 3. ©2009 Carolina Biological Supply Company Examining Thermochemistry | STUDENT GUIDE S-8

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Activity 3. (continued) 7. Light the candle and slide it directly under the center of the can. Adjust the height of the can so that the tip of the flame is touching the bottom of the can. Gently stir the water with the thermometer or temperature probe as the water is heating. 8. Continue heating the can until the water temperature is the same number of degrees above room temperature as it was below room temperature initially. For example, if room temperature is 25°C and your chilled water was at 15°C (10° below room temp.), then you should heat your water to 10° above room temperature (35°C). 9. After the water reaches the desired temperature, extinguish the flame and continue to stir the water until the maximum temperature is reached. Record this value in Data Table 3. 10. Weigh the candle/index card assembly again, including all of the wax drippings, and record the value in Data Table 3. 11. Determine the mass of paraffin combusted and record the value in Data Table 3.

CALCULATIONS A. Heat Lost by the Candle

The heat lost by the burning paraffin will be gained by the aluminum can and the water inside it. The Law of Conservation of Energy states that energy can never be created or destroyed. The equation for the calculation will be as follows: heat lost by candle = heat gained by can + heat gained by water candle can water in can Δ Δ q = –[(m1cp1 T1) + (m2cp2 T2)]

q = heat lost from candle combustion The minus sign indicates that the answer will be negative, because heat was lost from the candle.

m1 = mass of aluminum can

cp1 = specific heat of aluminum = 0.90 J/g•°C Δ ExploreT1 = Tf – Ti = change in can temperature (same temp. as water in can) ©2009 Carolina Biological Supply Company Examining Thermochemistry | STUDENT GUIDE S-9

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Activity 3. (continued)

m2 = mass of water in can

cp2 = specific heat of water = 4.2 J/g•°C Δ T2 = Tf – Ti = change in water temperature

B. Heat of Combustion

Assuming that all of the heat absorbed by the can and the water equals the heat released by the candle, one can calculate the heat of combustion for the paraffin fuel in Joules/g and convert to kJ/g by the following equation:

heat of combustion =

heat released by candle (J) 1 kJ × = ______kJ/g mass of paraffin combusted (g) 1000 J

C. Percentage of Error

The accepted value for the heat of combustion of paraffin is 42.0 kJ/g. The percentage of error is calculated as follows:

accepted value – experimental value % error = × 100 accepted value

QUESTIONS TO CONSIDER 1. Can one assume that all of the heat from the burning candle is absorbed by the aluminum can and the water inside it?

2. Why is it necessary to stir the water in the can while monitoring the temperature change of the water?

3. Why must you continue to monitor the temperature of the water even after the flame is extinguished?

4. Write a complete balanced equation for this reaction. ExploreInclude your experimental value for the heat of combustion. ©2009 Carolina Biological Supply Company Examining Thermochemistry | STUDENT GUIDE S-10

3 Explain

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Almost every chemical change either releases or absorbs heat as chemical bonds are broken and formed during the course of a reaction. Energy is consumed (absorbed) when bonds break, and energy is released as new bonds form. The net total for a reaction will either be a negative value with the release of heat (exothermic) or a positive value with the absorption of heat (endothermic).

Physical changes also release and absorb heat. The dissolving of salts involves heat of solution. Changes in state (such as melting and freezing, evaporating and condensing) involve gain and loss of heat. The study of the heat energy associated with chemical and physical changes is known as thermochemistry. Changes in heat energy are measured in a device called a calorimeter. Heat flow is measured indirectly, by placing the material to be measured (the system) in a reaction chamber surrounded by water (the surroundings). As the material reacts in the chamber, heat is transferred to the water or absorbed from the water. By monitoring the temperature change of the water, one can tell whether the material undergoing the reaction is endothermic (water temperature goes down) or exothermic (water temperature goes up).

Heat and temperature are often incorrectly perceived as the same quantity. Temperature is a property determined by the average kinetic energy of the particles in matter and is measured directly with a thermometer. Temperature measurements in thermochemistry are either in degrees Celsius or in the absolute temperature scale (Kelvin).

Heat is defined as the spontaneous flow of thermal energy from an object at a higher temperature to one at a lower temperature. The units for heat are calories in the English system and joules in the SI system. The Law of Conservation of Energy states that in an insulated system, the heat lost by one material will equal the heat gained by another material. The amount of heat transferred between two substances depends upon the temperature difference, the mass of the materials, and the specific heat capacities of the materials. The specific heat capacity is the amount of heat required to raise one gram of a substance 1°C. The SI unit for specific heat is joules per gram per degree Celsius (J/g•°C). Table 1 lists the specific heats of someExplain common materials at 25°C, or 298 K. ©2009 Carolina Biological Supply Company Examining Thermochemistry | STUDENT GUIDE S-11

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Table 1: Specific Heats at 25ºC (298 K) Substance Specific Heat (J/g•ºC) Aluminum (s) 0.897 Ammonia (g) 2.090 Calcium (s) 0.647 Carbon, graphite (s) 0.709 Copper (s) 0.385 Ethanol (g) 1.420 Ethanol (l) 2.440 Gold (s) 0.129 Iron (s) 0.449 Lead (s) 0.128 Mercury (I) 0.140 Water (g) 1.870 Water (l) 4.180 Water (s) 2.060

The equation that links specific heat, the mass of material, and the change in temperature for calculating the heat absorbed or released

by a material is as follows: q = (m)(cp)(ΔT), where q is the heat

transferred, cp is the specific heat capacity at constant pressure, m is the mass of the material, and ΔT is the change in temperature. The value ΔT is found by subtracting the initial temperature of the

material from the final temperature (Tf – Ti). If the sign of q is negative, heat will be lost (exothermic), and if the sign is positive, heat will be gained (endothermic). Look at the sample problem below for using this equation:

How much heat is required to raise 10.0 g of copper from 20.0ºC to 50.0ºC?

Δ q = (m)(cp)( T) q = (10.0 g)(0.385 J/g•ºC)(50.0ºC – 20.0ºC) q = (10.0 g)(0.385 J/g•ºC)(30.0ºC) q Explain= 116 J ©2009 Carolina Biological Supply Company Examining Thermochemistry | STUDENT GUIDE S-12

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HEAT OF REACTION FOR CHEMICAL CHANGES The energy absorbed or released during a chemical reaction is termed the heat of reaction. Exothermic reactions release energy since they spontaneously produce products that have lower energy than the reactants. An example is the combustion of carbon to form carbon dioxide.

C(s) + O2(g) CO2(g) + 393.5 kJ

Note that heat is a product of this reaction and that 393.5 kJ is

released per mole of CO2 formed. Another way of expressing the heat of this reaction is the term ΔH, which means change in enthalpy or heat. The symbol H represents the enthalpy or the heat content of a system at constant pressure. The heat content of a system cannot be measured directly. Instead, the change in enthalpy, ΔH, is measured (the Greek letter delta [Δ] signifies “change”). By convention, all exothermic reactions have a negative sign, indicating a loss of energy. Mathematically, the ΔH can be expressed as follows:

Δ H = Hproducts – Hreactants

Figure 1 shows the energy diagram for this exothermic reaction.

Energy Reactants

ΔH = – value Products

Reaction Pathway

Figure 1

If we reverse the CO2 equation, then energy must be absorbed to

break the bonds of the CO2 molecule. The 393.5 kJ appears as a reactant in this thermochemical equation:

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The ΔH is positive because heat is added in this reaction, giving the products more energy than the reactants. Figure 2 shows the absorption of heat energy for this reaction.

Energy Products

ΔH = + value Reactants

Reaction Pathway

Figure 2

HEAT OF FORMATION A heat of reaction in which a mole of a compound is produced from its elements is called molar heat of formation. The values for heats of formation are usually measured at their standard states and at standard conditions (1 atmosphere of pressure and room temperature, 25°C or

298 K). For example, the standard state of H2O at 1 atmosphere and 25°C is the liquid state. The degree symbol (°) is used with enthalpy symbols to indicate that the measurements are made for standard states

at standard conditions. The subscript f (f) is used to indicate that the value is for the heat of formation. By definition, the heat of formation Δ for a pure element is zero ( H°f = 0). One can use a table of heats of formation to find the value for many compounds. The more negative the enthalpy value, the more stable the compound since energy was lost as it was formed from its elements. Compounds with positive enthalpies are unstable and will often decompose back into their elements.

CALCULATING THE HEAT OF A REACTION FROM MOLAR HEATS OF FORMATION One can calculate the heat of a reaction for a chemical change by referring to a standard thermochemistry table to find the heat of formation for the various compounds in the reaction. Calculate the heat of reaction for the following chemical change and predict whether it is exothermic or endothermic.

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Δ Hint: Coefficients greater than one must be multiplied by the H°f value to determine the total amount of heat for that given number of moles.

Heats of formation for each compound taken from a standard heat of formation table are as follows: Δ Δ CO2(g): H°f = –393.5 kJ/mole H2O(l): H°f = –285.8 kJ/mole Δ Δ CH4(g): H°f = –74.9 kJ/mole O2(g): H°f = 0.0 kJ/mole

Calculation:

Δ ΣΔ ΣΔ H°= H°f of products – H°f of reactants Σ means “sum of” Δ Δ Δ Δ Δ H°= { H°f CH4 + 2[ H°f O2(g)]} – { H°f CO2(g) + 2[ H°fH2O(l)]} ΔH° = [(–74.9 kJ/mole) + 2(0.0 kJ/mole)] – [–393.5 kJ/mole + 2(–285.8 kJ/mole] ΔH° = [–74.9 kJ/mole] + [965.1 kJ/mole] ΔH° = 890.2 kJ/mole

The reaction is endothermic due to the positive sign.

HEAT OF COMBUSTION Another type of heat of reaction is heat of combustion, in which a substance reacts with oxygen gas to produce energy, carbon dioxide, and water. The combustion of fossil fuels is an exothermic reaction that provides the world with most of its current energy needs. The burning of propane gas in a gas grill is a good example of a combustion reaction.

C3H8(g) + 5O2(g) 3CO2(g) + 4H2O(l) + 2219 kJ

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QUESTIONS 1. Compare the heat content of the products to that of the reactants in the following two types of reactions and comment on the chemical stability of the product(s).

a. exothermic

b. endothermic

2. Explain the difference between heat of formation, heat of reaction, and heat of combustion.

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4. A mole of powdered magnesium (24.3 g) when mixed with water will produce enough heat to raise 1 L of water from 25°C to boiling. Calculate the heat in kJ that would be produced.

6. A 120-g sample of copper is exposed to 20.0 kJ of heat in an insulated chamber. By how many degrees Celsius will the copper sample’s temperature increase?

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7. a. Calculate the enthalpy, or heat of solution (ΔH), in kJ/mole for NaOH, given the following data: 5.0 g dissolved in 100 mL of water raised the water temperature in an insulated cup from 23.0ºC to 36.0ºC.

b. Is this physical change endothermic or exothermic?

8. For the following reactions and their given ΔHº values, write the equation with the enthalpy value on the product or reactant side, and identify the reaction as endothermic or exothermic.

Δ a. CO2(g) + 2H2O(l) CH4(g) + 2O2(g) Hº = +891 kJ

Δ b. 2Mg(s) + O2(g) 2MgO(s) Hº = –1200 kJ

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9. Ethanol, or ethyl alcohol, is the major component of E85 gasoline (85% ethanol and 15% gasoline). Write the balanced equation for ethanol combustion and calculate the heat of

combustion for one mole of ethanol (C2H5OH) using the following standard heats of formation:

Δ Δ CO2(g): H°f = –393.5 kJ/mole H2O(l): H°f = –285.8 kJ/mole Δ Δ C2H5OH(l): H°f = –277.0 kJ/mole O2(g): H°f = 0.0 kJ/mole

10. One type of cold pack has 100 g of water that mixes with 20 g

of NH4NO3. If the heat of solution for NH4NO3 is 25.4 kJ/mole, what would be the final temperature of the bag in degrees Celsius if the initial temperature were 23°C?

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Inquiry Activity: Designing a Prototype Hot Pack or Cold Pack

Your teacher will assign your group the task of designing a portable hot pack or cold pack for treating minor injuries, such as sprains. In Activity 1, you measured the change in temperature created by two different solids dissolving in 100 g of water. Use your new knowledge in the following three-part scenario.

A. How Much Chemical? You are a research chemist for an emergency care company. Company management has assigned your team (lab group) to design a portable, one-time-use hot pack or cold pack for treating injuries. This pack must have 100 g of water separated from a solid chemical and be activated only when the user does something to the pack to mix the two components. Your job is to determine how many grams of the chemical are required to achieve the following temperatures: hot pack (55°C, 131°F), cold pack (3°C, 37°F). You are allowed to test only

in increments of 5, 10, 15, and 20 g of either NH4Cl or CaCl2 to predict the final amount of chemical needed for the pack. That predicted amount will allow the business department of your company to do a cost analysis for the chemical needed for each pack.

B. Consensus of Teams Now that you have predicted the amount of chemical needed, you should consult with all the other lab teams that were assigned the same type of pack. After comparing your data, come to a consensus on how much chemical will be needed to achieve the desired final temperature of the pack.

C. Collaboration and Building a Prototype All groups for the same type of pack should brainstorm on the design of the pack, especially focusing on how the two chemicals can remain separated in the pack until needed and then be activated. Once the design is agreed upon, get teacher approval, secure your materials, and build one prototype representing all of the groups, with the 100 mL of water and the predicted amount of chemical. Two students, one representing the hot pack teams and one representing the cold pack teams, will conduct a class demo togetherExtend according to the following steps: ©2009 Carolina Biological Supply Company Examining Thermochemistry | STUDENT GUIDE S-20

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Name Date

1. Place the packs two feet apart on a lab table. Tape a lab thermometer to the top of each pack with transparent tape. Make sure the bulb of the thermometer is taped to the pack, insuring close contact.

2. After a few minutes, have the class record the initial temperature in degrees Celsius for both packs in Data Table 4.

3. Each student should pick up the pack and activate it by mixing the water and chemical.

4. The student for each pack calls out the temperature every minute and the class records it in Data Table 4.

5. Stop recording the data for each bag when its temperature remains constant.

6. Using graph paper or a computer graphing program, plot temperature vs. time, with temperature on the y-axis (in °C) and time on the x-axis (in minutes).

7. How long did it take for each pack to reach its maximum or minimum temperature?

8. Did each pack reach the predicted temperature?

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Data Table 1 Endothermic Heat of solution Compound Mass Final Temp. Initial Temp. ΔT = T – T q = (m)(cp)(ΔT) or f i (J/5.0 g) (1 kJ/1000 J) Exothermic?

CaCl2

NH4Cl

Data Table 2

Endothermic or Element Final Temp. Initial Temp. ΔT = T – T Appearance f i Exothermic?

Data Table 3: Paraffin Heat of Combustion

1. Mass of empty can: ______g

2. Volume of water in can: ______mL = ______g

3. Room temperature: ______ºC

4. Initial temperatue of water in can: ______ºC

5. Final temperature of water in can: ______ºC

6. Initial mass of candle with index card: ______g

7. Final mass of candle with index card: ______g

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Data Table 4: Temperatures of Hot Pack and Cold Pack Prototypes

Minutes Cold Pack Hot Pack 0 1 2 3 4 5 6 7 8 9 10 11 12 13 14 15

Experimental Design Template

Name Date

Question What are you testing in your experiment? What are you trying to find out?

Hypothesis What do you think will happen? Why do you think so?

Materials What are you going to use to find out the answer to the question?

Procedure What are you going to do? How are you going to do it?

Data Collection What data will you record and how will you collect and present it? Show and explain any data tables and graphs that you plan to use.

Data Analysis What happened? Did you observe anything that surprised you? Show and explain any tables and graphs that support your data.

Conclusion What conclusions can you draw based upon the results of your experiment? How does this compare with your initial hypothesis? If given the opportunity, how might you conduct the experiment differently?

©2009 Carolina Biological Supply Company Inquiries in Science® Chemistry Series Reactions: Matter Solutions Organic Chemistry Chemical and Nuclear Strand Strand Strand Strand RN-251301 RN-251302 NP-251303 Understanding Properties Finding Solutions Balancing Chemical Equations Modeling Hydrocarbons of Matter RN-251210 RN-251207 RN-251217 RN-251200

Changing States Observing Colligative Exploring Voltaic and of Matter Properties Electrolytic Cells RN-251201 RN-251211 RN-251215

Expanding on Investigating Reaction Rates the Gas Laws RN-251212 RN-251205

Reconstructing Atomic Attaining Equilibrium Theory RN-251213 RN-251202 Interpreting the Discovering Acids and Bases Periodic Table RN-251214 RN-251203

Bonding Chemically Examining Thermochemistry RN-251204 RN-251209

Determining Simulating Nuclear Chemical Formulas Transformations RN-251206 RN-251216

Calculating with Stoichiometry RN-251208 Inquiries in Science®: Complete Chemistry Series Lab Package Includes kits RN-251200 through RN-251217. ® RN-251300 Carolina Biological Supply Company ISBN 978-1-4350-0399-6 2700 York Road 90000 Burlington, North Carolina 27215 Phone: 800.334.5551 Fax: 800.222.7112 Technical Support: 800.227.1150 www.carolina.com 9 781435 003996 CB311500908