A Molecular Approach, 4e - Notes for Chapter (6).doc

Chapter 6. Thermochemistry Chapter 6. Thermochemistry Chapter 6. Thermochemistry Student Objectives 6.1 Chemical Hand Warmers Define thermochemistry. Understand the idea of heat exchange as a flow of energy. 6.2 The Nature of Energy: Key Definitions Define energy, work, heat, kinetic energy, thermal energy, potential energy, and chemical energy. Understand the difference between kinetic and potential energy and know why thermal energy and chemical energy are examples of each, respectively. Know the law of conservation of energy. Understand the difference between the system and the surroundings. Know and convert between the common units of energy: joule, calorie, kilocalorie, kilowatt-hour. 6.3 The First Law of Thermodynamics: There Is No Free Lunch Define and understand thermodynamics and the first law of thermodynamics. Understand why a perpetual motion machine violates the first law of thermodynamics. Define and understand internal energy and state function. Understand the flow of energy, as work or heat, from the standpoint of the system and the surroundings. Understand that energy lost by the surroundings is equal to the energy gained by the system and vice versa. Understand the mathematical definition of the first law of thermodynamics in terms of the change in internal energy, heat, and work. Understand the sign conventions for heat, work, and the change in internal energy. 6.4 Quantifying Heat and Work Define and understand thermal equilibrium. Define and understand heat capacity, specific heat capacity, and molar heat capacity. Understand the equation that relates heat flow to the amount of substance, the specific heat capacity, and the temperature change. Calculate using the equation relating heat flow and temperature change in terms of both the system and the surroundings. Define and understand pressure-volume work. Understand and use the equation for work in terms of the pressure and volume change. 6.5 Measuring E for Chemical Reactions: Constant-Volume Calorimetry Understand how constant-volume calorimetry is used to measure the exchange of energy between system and surroundings in terms of the heat flow. Calculate the heat released from reactions in a bomb calorimeter. 6.6 Enthalpy: The Heat Evolved in a Chemical Reaction at Constant Pressure Define enthalpy in terms of internal energy, pressure, and volume. Understand that the change in enthalpy for a reaction or process is equal to the heat flow under constant pressure. Define and understand endothermic and exothermic reactions. Understand the molecular view of endothermic and exothermic reactions. Define enthalpy of reaction. Understand and calculate enthalpy changes with respect to the stoichiometry of chemical equations. 6.7 Constant-Pressure Calorimetry: Meassuring Hrxn Describe and understand the coffee-cup calorimeter and how it can measure the enthalpy change for a chemical reaction or physical process. Understand and calculate enthalpy changes for reactions in a coffee-cup calorimeter. Understand the difference between constant-volume and constant-pressure calorimetry. 6.8 Relationships Involving Hrxn Know, understand, and calculate using reaction enthalpies: multiplying a chemical equation by a factor, reversing a chemical equation, and summing a series of chemical equations. Understand the use of Hess’s law to calculate the enthalpy change for a reaction from a series of steps. 6.9 Determining Enthalpies of Reaction from Standard Enthalpies of Formation Define and understand standard state and standard enthalpy of formation. Write thermochemical equations for the formation of compounds. Understand the equation for calculating the enthalpy of reaction from enthalpies of formation as an illustration of Hess’s law. Calculate the enthalpy of reaction using enthalpies of formation of products and reactants. 6.10 Energy Use and the Environment Know about energy consumption in the U.S.: sources and uses. Know about environmental problems associated with fossil fuel combustion. Understand and perform calculations involving the combustion of fossil fuels and the formation of carbon dioxide. Know about potential alternative fuels, especially renewable energy sources. Section Summaries Lecture Outline Terms, Concepts, Relationships, Skills Figures, Tables, and Solved Examples Teaching Tips Suggestions and Examples Misconceptions and Pitfalls Lecture Outline Terms, Concepts, Relationships, Skills Figures, Tables, and Solved Examples 6.1 Chemical Hand Warmers Thermochemistry Heat flow Intro figure: chemical hand warmers unnumbered figures: billiard balls, energy, work 6.2 The Nature of Energy: Key Definitions Energy work heat kinetic energy thermal energy potential energy chemical energy Law of conservation of energy system surroundings Energy units joule (J) calorie (cal) Calorie or kilocal (Cal or kcal) kilowatt-hour (kWh) Figure 6.1 The Different Manifestations of Energy Figure 6.2 Energy Transformation: Potential and Kinetic Energy I Figure 6.3 Energy Transformation: Potential and Kinetic Energy II Figure 6.4 Energy Transfer unnumbered figure: 100-W lightbulb Table 6.1 Energy Conversion Factors Table 6.2 Energy Uses in Various Units 6.3 The First Law of Thermodynamics: There Is No Free Lunch First Law of Thermodynamics definition perpetual motion machine State function Internal energy system surroundings relationship to heat and work E = q + w Chemistry in Your Day: Redheffer’s Perpetual Motion Machine Figure 6.5 Altitude as a State Function unnumbered figures: energy diagram of CO2 formation and decomposition Table 6.3 Sign Conventions for q, w, and E Figure 6.6 Energy, Work, and Heat Example 6.1 Internal Energy, Heat, and Work Teaching Tips Suggestions and Examples Misconceptions and Pitfalls 6.1 Chemical Hand Warmers Chemical hand warmers are easy to get and provide a simple but useful demonstration. 6.2 The Nature of Energy: Key Definitions Many students will have experience with a billiard table. Some examples can be reproduced on a countertop in a classroom. Drop a tennis ball from waist-high and ask the students to observe and explain all of its behavior, including coming to rest on the floor, in terms of kinetic and potential energy. Energy is always conserved, but one might not be able to identify easily the form(s) into which it is transformed. 6.3 The First Law of Thermodynamics: There Is No Free Lunch Paths to the top of a hill or mountain are common examples of path functions, whereas the change in elevation is a state function. Conceptual Connection 6.1 System and Surroundings Conceptual Connection 6.2 Heat and Work Billiard balls don’t keep moving forever since friction creates the heat loss that accounts for some of the energy. Lecture Outline Terms, Concepts, Relationships, Skills Figures, Tables, and Solved Examples 6.4 Quantifying Heat and Work Thermal equilibrium Heat flow and temperature change heat capacity specific heat capacity (Cs, in J/g°C) molar heat capacity (J/mol°C) q=mCsT Pressure–volume work w = PV Table 6.4 Specific Heat Capacities of Some Common Substances unnumbered figure: photo of San Francisco Example 6.2 Temperature Changes and Heat Capacity unnumbered figure: illustration of metal block in water Example 6.3 Thermal Energy Transfer unnumbered figure: illustration of internal combustion engine Figure 6.7 Piston Moving within a Cylinder against an External Pressure Example 6.4 Pressure–Volume Work 6.5 Measuring E for Chemical Reactions: Constant-Volume Calorimetry Internal energy change w = 0 at constant V q = E at constant V Calorimetry bomb calorimeter Figure 6.8 The Bomb Calorimeter Example 6.5 Measuring Erxn in a Bomb Calorimeter 6.6 Enthalpy: The Heat Evolved in a Chemical Reaction at Constant Pressure Enthalpy definition: H = E + PV H = q at constant pressure Reactions or processes endothermic: H > 0 exothermic: H < 0 Thermochemical equations H and stoichiometry unnumbered figures: illustrations of endothermic and exothermic processes Example 6.6 Exothermic and Endothermic Processes Example 6.7 Stoichiometry Involving H 6.7 Constant-Pressure Calorimetry: Measuring Hrxn Calorimetry coffee-cup calorimeter q = H at constant pressure Figure 6.9 The Coffee-Cup Calorimeter Example 6.8 Measuring Hrxn in a Coffee-Cup Calorimeter Teaching Tips Suggestions and Examples Misconceptions and Pitfalls 6.4 Quantifying Heat and Work A good conceptual check concerning heat capacity and temperature is to ask a student whether a chalk tray and eraser at the front of the room are the same temperature. Specific heat capacity is a measure of how much energy is required to change the temperature of a substance. Conceptual Connection 6.3 The Heat Capacity of Water Conceptual Connection 6.4 Thermal Energy Transfer Temperature and heat are not the same entity. Temperature is a measure of average kinetic energy, whereas heat is a flow of energy. Heat is a flow of energy. Substances do not contain heat; they lose or absorb heat. “Hot” and “cold” are only relative terms. Students sometimes forget that in this context always means final initial. 6.5 Measuring E for Chemical Reactions: Constant-Volume Calorimetry Heat capacities and heats of reaction are measured quantities; the values in tables come from experiments using calorimeters. Microcalorimeters are used for small quantities, especially for very valuable samples and/or biological samples. Ask the students to estimate the limits of this technique. The bomb calorimeter must be calibrated. Even though the heat capacity for water is well known, the best way to get the values for all the other parts is to measure it; this can be done by electrical heating. 6.6 Enthalpy: The Heat Evolved in a Chemical Reaction at Constant Pressure The direction of heat flow and conventions for numerical signs of exothermic and endothermic processes should be emphasized. Conceptual Connection 6.5 The Difference between H and E Conceptual Connection 6.6 Exothermic and Endothermic Reactions Practice with thermochemical equations will convince students that H can be used as a conversion factor just like the stoichiometric coefficients from the balanced reaction. 6.7 Constant-Pressure Calorimetry: Measuring Hrxn The coffee-cup calorimeter can be used for reactions and physical processes, e.g. dissolving a solid and phase changes. Conceptual Connection 6.7 Constant-Pressure versus Constant-Volume Calorimetry Students may initially miss the fact that an exothermic reaction done in a coffee-cup calorimeter will result in an increased temperature for the aqueous solution even though the reaction itself is decreasing in enthalpy. Lecture Outline Terms, Concepts, Relationships, Skills Figures, Tables, and Solved Examples 6.8 Relationships Involving Hrxn Chemical equations and Hrxn multiplying equations by a factor reversing equations summing a series of equations Hess’s law illustration of state function Figure 6.10 Hess’s Law Example 6.9 Hess’s Law 6.9 Determining Enthalpies of Reaction from Standard Enthalpies of Formation Definitions standard state standard enthalpy change, H° standard enthalpy of formation, H°f Calculating the standard enthalpy change for a reaction Table 6.5 Standard Enthalpies (or Heats) of Formation, Hf°, at 298 K Example 6.10 Standard Enthalpies of Formation Figure 6.11 Calculating the Enthalpy Change for the Combustion of Methane Example 6.11 and Standard Enthalpies of Formation unnumbered figure: photo of thermite reaction Example 6.12 and Standard Enthalpies of Formation 6.10 Energy Use and the Environment Energy consumption U.S. energy use U.S. energy sources Environmental problems related to energy use carbon dioxide production from fossil fuel combustion unnumbered figure: pie–chart of energy consumption Figure 6.12 Energy Consumption by Source Table 6.6 Changes in National average Pollutant Levels, 1980–2010 Figure 6.13 The Rise in Atmospheric Carbon Dioxide Example 6.13 Fossil Fuels and Climate Change Chemistry and the Environment: Renewable Energy Teaching Tips Suggestions and Examples Misconceptions and Pitfalls 6.8 Relationships Involving Hrxn Quantitative manipulations of thermochemical equations require changes in Hrxn: multiplying by a factor, reversing sign, and summing enthalpies. Hess’s law is a powerful and flexible way to generate new thermodynamic data from chemical equations. In order to apply Hess’s law properly, all the other species except the desired ones must cancel. 6.9 Determining Enthalpies of Reaction from Standard Enthalpies of Formation Absolute enthalpy cannot be measured, so the enthalpies of formation of the elements (in their standard states) are set to 0. Each enthalpy of formation value must be multiplied by the stoichiometric coefficient from the balanced reaction. Students sometimes fail to see that the equation for the standard enthalpy change is just an application of Hess’s law. 6.10 Energy Use and the Environment Review the uses and sources of energy in the U.S., and examine trends in consumption. Performing simple calculations concerning CO2 emission makes the emission problem far more obvious. Discuss alternative fuels and sources of energy, both the positives and negatives. Renewable fuels like ethanol still produce carbon dioxide. Additional Problem for Pressure–Volume Work (Example 6.4) Inflating one’s lungs requires the inflator to do pressure–volume work on the surroundings. If the lungs are inflated from a volume of 2.5 L to 5.5 L against an external pressure of 1.10 atm, how much work is done? Sort You are given the initial and final volumes of the lungs and the pressure against which they expand. The lungs and their contents are the system. Given V1 = 2.5 L V2 = 5.5 L P = 1.10 atm Find w Strategize The equation w = PV specifies the amount of work done during a volume change against an external pressure. A conversion factor for converting to joules is necessary. Conceptual Plan P, V w Relationships Used V = V2 – V1 101.3 J = 1 L atm Solve Enter the values into the equation and calculate the value for work. Solution V = 5.5 – 2.5 L = 3.0 L Check The units of the answer (J) are correct. The sign is negative as it should be for expansion; work is being done on the surroundings by expanding the lungs. Additional Problem for Measuring Erxn in a Bomb Calorimeter (Example 6.5) When 1.10 g of ethanol (C2H6O) undergoes combustion in a bomb calorimeter, the temperature rises from 22.32°C to 29.48°C. Find Erxn for the combustion of ethanol in kJ/mol. The heat capacity of the bomb calorimeter, determined in a separate experiment is 4.70 kJ/°C. Sort You are given the mass of ethanol, the heat capacity of the calorimeter, and the initial and final temperatures. You are asked to find the change in internal energy for the reaction. Given 1.10 g ethanol Ti = 22.32°C Tf = 29.48°C Ccal = 4.70 kJ/°C Find Erxn Strategize The conceptual plan has three parts: 1) Use the temperature change and heat capacity of the calorimeter to find qcal. 2) Use qcal to find qrxn, a matter of changing signs. 3) Divide qrxn by the number of moles of ethanol to get Erxn per mole of ethanol. Conceptual Plan Ccal, T qcal qrxn Relationships Used qcal = Ccal T = qrxn T = Tf – Ti molar mass (C2H6O) = 46.07 g/mol Solve Enter the values into the equations for T, qcal and Erxn. Solution T = 29.48°C – 22.32°C = 7.16°C Check The units of the answer (kJ/mol) are correct. The sign is negative and is correct for a reaction that gives off energy. Additional Problem for Stoichiometry involving H (Example 6.7) A cigarette lighter contains 2.90 g of butane, C4H10. Calculate the heat (in kJ) associated with the complete combustion of the butane in the lighter. C4H10(g) + 13/2 O2(g) 4 CO2(g) + 5 H2O(g) Hrxn = 2658 kJ Sort You are given the mass of propane and asked to find the heat evolved in its combustion. Given 2.90 g C4H10 Find q Strategize Starting with mass of butane, convert to moles using the molar mass as a conversion factor. Next, use the stoichiometric relationship between C4H10 and kJ to find the heat evolved. Conceptual Plan g C4H10 mol C4H10 kJ Relationships Used molar mass C4H10 = 58.12 g/mol 1 mol C4H10: 2658 kJ (from equation) Solve Follow the conceptual plan to solve the problem. Begin with the mass and multiply with the appropriate conversion factors to arrive at kJ. Solution Check The units (kJ) are correct. The sign is negative, as it should be for heat evolved by the reaction. Additional Problem for Measuring Hrxn in a Coffee-Cup Calorimeter (Example 6.8) Calcium metal reacts with hydrochloric acid according to the balanced equation: Ca(s) + 2 HCl(aq) CaCl2(aq) + H2(g) A 0.150 g sample of Ca metal is combined with enough HCl to make 100.0 mL of solution in a coffee-cup calorimeter. The calcium reacts and the temperature of the solution rises from 23.5 °C to 25.4 °C. Find Hrxn for the reaction. Use 1.00 g/mL as the density of the solution and Cs,soln = 4.18 J/g °C as the specific heat capacity of the solution. Sort You are given the mass of calcium, the volume of the solution, the initial and final temperatures, the density of the solution, and the heat capacity of the solution. You are asked to find the enthalpy change for the reaction. Given 0.158 g Ca 100.0 mL solution Ti = 23.5 °C Tf = 25.4 °C d = 1.00 g/mL Cs, soln = 4.18 J/g°C Find Hrxn Strategize The conceptual plan has three parts: 1) Use the temperature changes, specific heat capacity, and mass of the solution to solve for qsoln. 2) Use qsoln to find qrxn, which is just a change of sign. At constant pressure qrxn is equivalent to Hrxn. 3) Divide qrxn by the number of moles of calcium to get Hrxn per mole of calcium. Conceptual Plan Cs, soln, msoln, T qsoln qrxn Relationships Used q = m Cs T qrxn = qsoln Solve Enter the values and determine the mass of the solution and the temperature change. Solve for qsoln using the equation. Convert qsoln to qrxn with the sign change. Solve for the enthalpy change by entering the values into the equation. Solution T = Tf – Ti = 25.4 °C – 23.5 °C = 1.9 °C qrxn = qsoln = 779.9 J Check The units (J) are correct. The sign is negative as expected for an exothermic reaction. Additional Problem for H°rxn and Standard Enthalpies of Formation (Example 6.11) Use the standard enthalpies of formation to determine H°rxn for the reaction: Zn(s) + 2 HCl(aq) ZnCl2(aq) + H2(g) Sort You are given the balanced equation and asked to find the enthalpy of reaction. Given Zn(s) + 2 HCl(aq) ZnCl2(aq) + H2(g) Find Strategize To calculate H°rxn from standard enthalpies of formation, subtract the heats of formations of the reactants multiplied by their stoichiometric coefficients from the heats of formation of the products multiplied by their stoichiometric coefficients. Conceptual Plan Solve Begin by looking up (in Appendix IIB) the standard enthalpy of formation for each reactant and product. Remember that the standard enthalpy of formation of pure elements in their standard state is zero. Compute H°rxn by substituting into the equation. Solution Reactant or product Hf° (kJ/mol, Appendix IIB) Zn(s) 0 HCl(aq) 167.2 ZnCl2(aq) 153.39 H2(g) 0 Check The units (kJ) are correct. The answer is positive, which means that the reaction is endothermic. Additional Problem for Fossil Fuels and Climate Change (Example 6.13) In Brazil, vehicles are fueled almost entirely by ethanol C2H6O, produced from sugar cane. Calculate the heat (in kJ) produced by ethanol for each 1.00 kg of CO2 produced. Sort You are given the mass of CO2 emitted and asked to find the energy output for ethanol. Given 1.00 kg CO2 Find kJ Strategize Write the thermochemical equation for the combustion of ethanol. The conceptual plan has two parts: 1) Use the molar mass of CO2 to convert from mass of CO2 to moles of CO2. 2) Use the stoichiometric relationship between moles of CO2 produced and kJ of energy released to calculate the energy output. C2H6O(l) + 7/2 O2(g) 2 CO2(g) + 3 H2O(g) = 1409 kJ Conceptual Plan kg CO2 g CO2 mol CO2 kJ Relationships Used For C2H6O: 2 mol CO2: 1409 kJ 1 kg = 1000 g molar mass CO2 = 44.01 g/mol Solve Convert kg CO2 to mol CO2 and then to kJ. Solution Check The units (kJ) are correct. The value is negative, as expected for exothermic combustion reactions. For comparison, other fuels have the values: For C: 8.94 103 kJ For CH4: 1.82 104 kJ For C8H18: 1.44 104 kJ Ethanol compares favorably. 86 Copyright © 2017 by Education, Inc. 87 Copyright © 2017 by Education, Inc.