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Chapter 6 Thermochemistry

Chapter 6 Thermochemistry. Contents and Concepts. Understanding Heats of Reaction The first part of the chapter lays the groundwork for understanding what we mean by heats of reaction. Energy and Its Units Heat of Reaction Enthalpy and Enthalpy Changes Thermochemical Equations

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Chapter 6 Thermochemistry

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  1. Chapter 6Thermochemistry

  2. Contents and Concepts Understanding Heats of Reaction The first part of the chapter lays the groundwork for understanding what we mean by heats of reaction. • Energy and Its Units • Heat of Reaction • Enthalpy and Enthalpy Changes • Thermochemical Equations • Applying Stoichiometry to Heats of Reaction • Measuring Heats of Reaction

  3. Using Heats of Reaction Now that we understand the basic properties of heats of reaction and how to measure them, we can explore how to use them. 7. Hess’s Law 8. Standard Enthalpies of Formation 9. Fuels—Foods, Commercial Fuels, and Rocket Fuels

  4. Learning Objectives Understanding Heats of Reaction • Energy and Its Units • Define energy, kinetic energy, and internal energy. • Define the SI unit of energy (joule) as well as the common unit of energy (calorie). • Calculate the kinetic energy of a moving object. • State the law of conservation of energy.

  5. Heat of Reaction • Define a thermodynamic system and its surroundings. • Define heat and heat of reaction. • Distinguish between an exothermic process and an endothermic process. • Enthalpy and Enthalpy Changes • Define enthalpy and enthalpy of reaction. • Explain how the terms enthalpy of reaction and heat of reaction are related. • Explain how enthalpy and internal energy are related.

  6. Thermochemical Equations • Define a thermochemical equation. • Write a thermochemical equation given pertinent information. • Learn the two rules for manipulating (reversing and multiplying) thermochemical equations. • Manipulate a thermochemical equation using these rules. • Applying Stoichiometry to Heats of Reaction • Calculate the heat absorbed or evolved from a reaction given its enthalpy of reaction and the mass of a reactant or product.

  7. Measuring Heats of Reaction a. Define heat capacity and specific heat. b. Relate the heat absorbed or evolved to the specific heat, mass, and temperature change. c. Perform calculations using the relationship between heat and specific heat. d. Define a calorimeter. e. Calculate the enthalpy of reaction from calorimetric data (temperature change and heat capacity).

  8. Using Heats of Reaction • Hess’s Law a. State Hess’s law of heat summation. b. Apply Hess’s law to obtain the enthalpy change for one reaction from the enthalpy changes of a number of other reactions.

  9. 8. Standard Enthalpies of Formation a. Define standard state and reference form. b. Define standard enthalpy of formation. c. Calculate the heat of a phase transition using standard enthalpies of formation for the different phases. d. Calculate the heat (enthalpy) of reaction from the standard enthalpies of formation of the substances in the reaction.

  10. Fuels—Foods, Commercial Fuels, and Rocket Fuels • Define a fuel. • Describe the three needs of the body that are fulfilled by foods. • Give the approximate average values quoted (per gram) for the heat values (heats of combustion) for fats and for carbohydrates. • List the three major fossil fuels. • Describe the processes of coal gasification and coal liquefaction. • Describe some fuel-oxidizer systems used in rockets.

  11. Thermodynamics • The science of the relationship between heat and other forms of energy. • Thermochemistry • An area of thermodynamics that concerns the study of the heat absorbed or evolved by a chemical reaction

  12. Energy • The potential or capacity to move matter. • One form of energy can be converted to another form of energy: electromagnetic, mechanical, electrical, or chemical. • Next, we’ll study kinetic energy, potential energy, and internal energy.

  13. Kinetic Energy, EK • The energy associated with an object by virtue of its motion. • m = mass (kg) • v = velocity (m/s)

  14. The SI unit of energy is the joule, J, pronounced “jewel.” • The calorie is a non-SI unit of energy commonly used by chemists. It was originally defined as the amount of energy required to raise the temperature of one gram of water by one degree Celsius. The exact definition is given by the equation:

  15. A person weighing 75.0 kg (165 lbs) runs a course at 1.78 m/s (4.00 mph). What is the person’s kinetic energy? m = 75.0 kg V = 1.78 m/s EK = ½ mv2

  16. Potential Energy, EP • The energy an object has by virtue of its position in a field of force, such as gravitaitonal, electric or magnetic field. • Gravitational potential energy is given by the equation • m = mass (kg) • g = gravitational constant (9.80 m/s2) • h = height (m)

  17. Internal Energy, U • The sum of the kinetic and potential energies of the particles making up a substance. • Total Energy • Etot = EK + EP + U

  18. Law of Conservation of Energy • Energy may be converted from one form to another, but the total quantity of energy remains constant.

  19. Thermodynamic System • The substance under study in which a change occurs is called the thermodynamic system (or just system). • Thermodynamic Surroundings • Everything else in the vicinity is called the thermodynamic surroundings (or just the surroundings).

  20. Heat, q • The energy that flows into or out of a system because of a difference in temperature between the thermodynamic system and its surroundings. • Heat flows spontaneously from a region of higher temperature to a region of lower temperature. • • q is defined as positive if heat is absorbed by the system (heat is added to the system) • • q is defined as negative if heat is evolved by a system (heat is subtracted from the system)

  21. Heat of Reaction • The value of q required to return a system to the given temperature at the completion of the reaction (at a given temperature)

  22. Endothermic Process • A chemical reaction or process in which heat is absorbed by the system (q is positive). The reaction vessel will feel cool. • Exothermic Process • A chemical reaction or process in which heat is evolved by the system (q is negative). The reaction vessel will feel warm.

  23. In an endothermic reaction: The reaction vessel cools. Heat is absorbed. Energy is added to the system. q is positive. • In an exothermic reaction: The reaction vessel warms. Heat is evolved. Energy is subtracted from the system. q is negative.

  24. Enthalpy, H • An extensive property of a substance that can be used to obtain the heat absorbed or evolved in a chemical reaction. • Extensive Property • A property that depends on the amount of substance. Mass and volume are extensive properties.

  25. A state function is a property of a system that depends only on its present state, which is determined by variables such as temperature and pressure, and is independent of any previous history of the system.

  26. The altitude of a campsite is a state function. • It is independent of the path taken to reach it.

  27. Enthalpy of Reaction • The change in enthalpy for a reaction at a given temperature and pressure: • DH = H(products) – H(reactants) Note: D means “change in.” Enthalpy change is equal to the heat of reaction at constant pressure: DH = qP

  28. The diagram illustrates the enthalpy change for the reaction • 2Na(s) + 2H2O(l)  2NaOH(aq) + H2(g) The reactants are at the top. The products are at the bottom. The products have less enthalpy than the reactants, so enthalpy is evolved as heat. The signs of both q and DH are negative.

  29. Enthalpy and Internal Energy • The precise definition of enthalpy, H, is • H = U + PV • Many reactions take place at constant pressure, so the change in enthalpy can be given by • DH = DU + PDV • Rearranging: • DU = DH –PDV • The term (–PDV) is the energy needed to change volume against the atmospheric pressure, P. It is called pressure-volume work.

  30. For the reaction • 2Na(s) + 2H2O(l)  2NaOH(aq) + H2(g) –PDV The H2 gas had to do work to raise the piston. For the reaction as written at 1 atm, -PDV = -2.5 kJ. In addition, 368.6 kJ of heat are evolved.

  31. Thermochemical Equation • The thermochemical equation is the chemical equation for a reaction (including phase labels) in which the equation is given a molar interpretation, and the enthalpy of reaction for these molar amounts is written directly after the equation. • For the reaction of sodium metal with water, the thermochemical equation is: • 2Na(s) + 2H2O(l)  • 2NaOH(aq) + H2(g); DH = –368.6 kJ

  32. Sulfur, S8, burns in air to produce sulfur dioxide. The reaction evolves 9.31 kJ of heat per gram of sulfur at constant pressure. Write the thermochemical equation for this reaction.

  33. We first write the balanced chemical equation: S8(s) + 8O2(g)  8SO2(g) Next, we convert the heat per gram to heat per mole. Note: The negative sign indicates that heat is evolved; the reaction is exothermic. Now we can write the thermochemical equation: S8(s) + 8O2(g)  8SO2(g); DH = –2.39 × 103 kJ

  34. Manipulating a Thermochemical Equation • • When the equation is multiplied by a factor, the value of DH must be multiplied by the same factor. • • When a chemical equation is reversed, the sign of DH is reversed.

  35. a. CH4(g) + H2O(g)  CO(g) + 3H2(g) • This reaction is identical to the given reaction. • It is endothermic. • DH = 206 kJ CH4(g) + H2O(g)  CO(g) + 3H2(g); DH = 206 kJ

  36. b. 2 CH4(g) + 2H2O(g)  2CO(g) + 6H2(g) • This reaction is double the given reaction. • It is endothermic. • DH = 412 kJ CH4(g) + H2O(g)  CO(g) + 3H2(g); DH = 206 kJ

  37. c. CO(g) + 3H2(g)  CH4(g) + H2O(g) • This reaction is the reverse of the given reaction. • It is exothermic. • DH = -206 kJ CH4(g) + H2O(g)  CO(g) + 3H2(g); DH = 206 kJ

  38. d. 2CO(g) + 6H2(g)  2CH4(g) + 2H2O(g) • This reaction is reverse and double the given reaction. • It is exothermic. • DH = -412 kJ • Equations c and d are exothermic. • Equation d is the most exothermic reaction. CH4(g) + H2O(g)  CO(g) + 3H2(g); DH = 206 kJ

  39. When sulfur burns in air, the following reaction occurs: • S8(s) + 8O2(g)  8SO2(g); • DH = – 2.39 x 103 kJ Write the thermochemical equation for the dissociation of one mole of sulfur dioxide into its elements.

  40. S8(s) + 8O2(g)  8SO2(g); DH = –2.39 × 103 kJ • We want SO2 as a reactant, so we reverse the given reaction, changing the sign of DH: • 8SO2(g)  S8(g) + 8O2(g) ; DH = +2.39 × 103 kJ • We want only one mole SO2, so now we divide every coefficient and DH by 8: SO2(g)  1/8S8(g) + O2(g) ; DH = +299 kJ

  41. Applying Stoichiometry to Heats of Reaction

  42. You burn 15.0 g sulfur in air. How much heat evolves from this amount of sulfur? The thermochemical equation is S8(s) + 8O2(g)  8SO2(g); DH = -2.39 x 103 kJ

  43. S8(s) + 8O2(g)  8SO2(g); DH = -2.39 x 103 kJ • Molar mass of S8 = 256.52 g q = –1.40 × 102 kJ

  44. The daily energy requirement for a 20-year-old man weighing 67 kg is 1.3 x 104 kJ. For a 20-year-old woman weighing 58 kg, the daily requirement is 8.8 x 103 kJ. If all this energy were to be provided by the combustion of glucose, C6H12O6, how many grams of glucose would have to be consumed by the man and the woman per day? C6H12O6(s) + 6O2(g)  6CO2(g) + 6H2O(l); DH = -2.82 x 103 kJ

  45. C6H12O6(s) + 6O2(g)  6CO2(g) + 6H2O(l); • DH = -2.82 x 103 kJ For a 20-year-old man weighing 67 kg: = 830 g glucose required (2 significant figures) For a 20-year-old woman weighing 58 kg: = 560 g glucose required (2 significant figures)

  46. Measuring Heats of Reaction • We will first look at the heat needed to raise the temperature of a substance because this is the basis of our measurements of heats of reaction.

  47. Heat Capacity, C,of a Sample of Substance • The quantity of heat needed to raise the temperature of the sample of substance by one degree Celsius (or one Kelvin). • Molar Heat Capacity • The heat capacity for one mole of substance. • Specific Heat Capacity, s (or specific heat) • The quantity of heat needed to raise the temperature of one gram of substance by one degree Celsius (or one Kelvin) at constant pressure.

  48. The heat required can be found by using the following equations. • Using heat capacity: • q = CDt • Using specific heat capacity: • q = s x m x Dt

  49. A calorimeter is a device used to measure the heat absorbed or evolved during a physical or chemical change. Two examples are shown below.

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